Term
| The Schrodinger Equation and the Atomic Orbital |
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Definition
The electron’s matter wave occupies the space near the nucleus and is continuosly influenced by it. The Schrodinger equation is quite complex but can be represented in simpler form as
HΨ = EΨ
Wher E is energy o fthe atom. THe symbol Ψ (psi) is called a wave function (or atomic orbital), a mathematical description of the electron’s matter-wave in three dimensions. The symbol H, called the Hamiltonian operator, represents a set of mathematical operations that, when carried out with particular Ψ, yield one of the allowed energy states of the atom. Thus each solution of the equation gives an energy state associated with a given atomic orbital.
An important point to keep in mid throughout this discussion is that an “orbital” in the quantum-mechanical model bears no resemblance to an “orbit” in the Bor model: an orbit is an electon’s actual path around the nucleus, whereas an orbital is a mathematical funchtion that describes the electron’s matter-wave but has physical meaning. |
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Term
| The probable location of the Electron |
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Definition
| While we cannot knoe exactly where the electron is at any moment, we can know where it probably is, that is, where it spends most of its time. We get this information by squaring the wave function. Thus, even though Ψ has no physical meaning, Ψ^2 does and is called the probabillity density, a measure of the probability of finding the electron in some tiny volume of the atom. We depict the electron’s probable location in several ways. |
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Term
| Probability of the electron being in some tiny volume of an atom |
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Definition
For each energy level, we can create an electron probaility density diagram, or more simply an electron density diagram. The value of Ψ^2 for a given volume is shown with with dots: the greater the density of dots, the higher the probability of finding the electron in that volume. Note, that for the ground state of the H atom, the electron probability density decreases with distance form the nucleus along a line, r.
These diagrams are also called electron cloud depiction because, if we could take a time-exposure photograph of the electron in wavelike motion around the nucleus, it would appear as a “cloud” of positions. The electron cloud is an imaginary picture of the electron changing its position rapidly over time; it does not mean that an electron is a diffuse cloud of charge. |
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Term
| Total probability density at some distance from the nucleus. |
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Definition
To find radiant probability distribution, that is, the total probability of find the electron at some distance r from the nucleus, we first mentally divide the volume around the nucleus into thin, concentric, spheric layers, like the layers of an onion . Then, we find the sum of Ψ^2 values in each layer to see which is most likely to contain the electron.
The fallout in probability density with distance has an important effect. Near the nucleus, the volume of each layer increases faster than its density of dots decreases. the result of these opposing effects is that the total probaility peaks in a layer is higher than in the first, but this result disappears with greater distance. Figure 7.16D shows this result as a radial probability distribution plot. |
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Term
| Probability contour and the size of the atom |
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Definition
| How far away form the nucleus can we find the electron? this is the same as asking “How big is the H atom?” Recall from Figure 7.16B that the probability of finding the electron far form the nuclus is not zero. Therefore, we cannot assign a definite volume to an atom. However, we can visualize an atom with a 90% probability contour: the electron is somewhere within the volume 90% of the time. |
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Term
| The principal quantum number (n) |
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Definition
| is a positive integer (1,2,3,4...). It indicates the relative size of the orbital and therefore the relative distance from teh nucleus of the peak in the radial probability distribution plot. The principle quantum number specifies the energy level of the H atom: the higher the n value, the higher the energy level. When the electron occupies an orbital with n=1, the H atom is its ground state and its lowest energy. When the electron occupies an orbital with n=2 (first energy state), the atom has more energy. |
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Term
| THe angular momentum quatnum number (l) |
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Definition
| is an integer from 0 to n-1. It is related to the shape of the orbital. Not that the principal quantum number sets a limit on the angular momentum quantum number: n limits l. For an orbital with n = 1, l can have only one value, 0. For orbitals with n = 2, l can have two values, 0 or 1. For with n=3, l can have three values, 0, 1, or 2; and so forth. THus the number of possible l values equals the value of n |
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Term
| The magnetic quantum number m(l) |
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Definition
| is an integer from -1 through 0 to +l. It prescribes the three-dimenesional orientation of the orbital in the space around the nucleus. The angular momentum quantum number sets a limt on the magnetic quantum number: l limits m(l) values, -1, 0 , or +1; there are three possible orbitals with l=1, each is the number of orbitals for a given l. The total number of m(l) values, that is, the total number of orbitals, for a given n value is n^2. |
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Term
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Definition
| Level. The atom’s engergy levels, or shells, are given by the n value: the smaller the n value, the lower the energy level and the greater the probability that the electron is closer to the nucleus. |
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Term
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Definition
THe atom’s levels are divided into sublevels, or subshells, that are given by the l value. Each designates the orbital shape with a letter:
l= 0 is and s sublevel
l= 1 is a p sublevel
l= 2 is a d sublevel
l= 3 is an f sulevel
The letter derive from names of spectroscopic lines: sharp, principal, diffuse, and fundamental. sublevels with l values greater than 3 are designated by consecutive letters after f:g sublevel, h sublevel, and so on. A sublevel is named with its n value and letter disignation: for example, the sublevel (subshell) with n=2 and l=0 is called the 2s sublevel. |
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Term
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Definition
| Each combination of n, l, and m(l) specifies the size (energy), shape, and spatial orientation of one atom’s orbitals. We know the quantum numbers of the orbitals in a sublevel from the sublevel name and the quantum-number hierachy. For example, any orbital in the 2s sublevel has n=2 and l=0, and given that l in the l value, it can have only m(l) = 0; thus, the 2s sublevel has only one orbital. Any orbital in the 3p sublevel has n=3 and l=1, and given that l value, one orbital has m(l)= -1, another has m(l) = 0, and a third has m(l) = +1; thus, the 3p sublevel has three orbitals. |
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Term
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Definition
| An orbital with l = 0 has a spherical shape with nuclues at the center and is called an s orbital. Because a sphere has only one orientation, an s orbital has one m(l) value: for any s orbiral, m(l) = 0 |
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Term
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Definition
| hold the electons in the H atom’s ground state. THe electron probability density is highest at the nucleus. Figure 7.17A shows this graphically (top), and an electron density relief map (inset) depicts the graph’s curve in three dimensions. |
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Term
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Definition
| has two regions of higher electron density. The radial probability distibution of the more distant region is higher than that of the closer one because the Ψ^2 is taken over a much larger volume. Between the two regions is spherical node, where the probability of finding a the electron to zero. Because 2s orbital is larger than 1s, and electron in the 2s spends more time farther from teh nucleus (in the larger of the two regions) thatn it does when it occupies the 1s. |
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Definition
| has three regions of high electron density and two nodes. Here again, the highest radial probability is the greatest distance form the nucleus. This pattern of more nodes and higher probability with distance from the nucleus continues the 4s, 5s, and so forth. |
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Term
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Definition
An orbital with l=1 is called a p orbital and has rwo regions (lobes) of high probability, one on wither side of the nucleus. The nucleus lies at the nodal plane of teh dumbbell-shaped orbital. Since the maximum value of l is n-1, only levels with n = 2 or heger have p orbital: the lowest energy p orbital (the one closest to the nucleus) is the 2p. One p orbital consists of two lobes, and the electron spends equal time in both. Similar to the pattern for s orbitals, a 3p orbital is larger than 2p, a 4p is larger than 3p, and so forth
Unlike s orbitals, p orbitals have different spatial orientaions. The three possible m(l) values of -1,0, and +1 refer to three mutally perpendicular orientations; that is, while identical in size, shape, an denergy, the three p orbtals differ in orientation. We associate p orbital with the x, y, z axes: the p(x) orbital lies along the x axis, the p(y) along teh y axis, and the p(z) along the z axis |
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Term
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Definition
An orbital with l = 2 is called a d orbital. There are five possible m(l) values for l=2: -2,-1,0,1,2. Thus a d orbiral has any one of five orientations. Four of the five d orbitals have four lobes (a cloverleaf shepe) with two mutually perpendicular nodal planes between them and the nucleus at the junction of teh lobes. (the orientation of the nodal planes always lies between the orbtal lobes). Three of these orbitals lie in the xy, xz and yz planes, with their lobes between the axes, and called the d(xy), d(xz), d(yz) orbitals. A fourth the d(x^2-y^2) orbital, also lies in teh xy plane, but its lobes are along the axes. The fifth d orbital, the d(z^2), has two major lobes along the z axes, and a donut-shaped region girdles the cetner. An electron in a d orbital spends equal time in all of its lobes.
In keeping with quantum-number hierachy, a d orbital (l=2) must have principal quantum number of n=3 or higher, so 3d is the lowest energy d sublevel. Orbitals in the 4d sublevels are larger (extended farther form the nucleus) than the 3d, and the 5d larger still. |
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Term
| Orbitals with Higher l Values |
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Definition
| Orbitals with l =3 are f orbitals and have principal quantum number of at least n = 4. Obitals with l = 4 are g orbitals, but they play noknown role in chemical bonding. |
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Term
| The Electron-Spin Quantum Number |
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Definition
The three quantum numbers n, l, and m(l) describe the size (energy), shape, and orientation, respectively, of an atomic orbital. An additional quantum number describe a property called spin, which is a property of the electron and not the orbital.
When a beam of atoms that have one or more lone electrons passes through a mnonuniform magnetic field (created by magnet faces with different shapes), it splits into two beams. Each electron behaves like a spinning charge and generates a tiny magnetic fiel, which can have one of two values of spin. The two lectron fields have opposing directions, so half of the electons are attracted by large external magnetic field while the other half are repelled by it. |
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Term
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Definition
Corresponding to the two directions of the electron’s field, the spin quantum number m(s) has two possible values, +1/2, -1/2. THus, wach electron in an atom is describd completely by a set of four quantum numbers: the first three describe its orbital, and the fourth describes its spin. THe quantum numbers summarized in Table 8.1.
Now we can write a set of four quantum numbers for any electron in tghe ground state of any atomfor example, the set of quantum numbers for the lone electron in hydrogen (H; Z = 1) is n = 1, l = 0, and m(s) +1/2. |
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Term
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Definition
The element after hydrogen is helium (He; Z = 2), the first with atoms having more than one electron. THe first electron in the He ground state has the same ste of quantum numbers as the electron in the H atom, but the second He electron does not. Based on observations of excited states, the Austrian physicist Wolfgang Pauli formulated the exclusion principle; no two electron in the same atom can the same four quntum numbers,. Therefore, the second He electron occupies the same orbital as the first but has an opposite spin: n = 1, l=0, m(l) = 0, and m(s) = -1/2
THe major consequence of the exclusion principle is the that an atomic orbital can hold a maximum of two electrons and they must have opposing spins. We say that the 1s orbital in He is filled and that the electrons have pared spins. Thus, a beam of He atoms is not split in an experiment like that in Figure 8.1 |
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Term
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Definition
| attraction of opposite charges and repulsion of like charges and repulsion of like chatges - play a major role in determining the energy between nucleus and electron and the energy state is determined only by the n value, the energy states of many-electron atoms are also affected by electron-electron repulsions. This gice rise to the splitting of energy levels into sublevels of differing energies: the energy of an orbital in a may-electron atom depends mostly on the n value (size) and to a lesser the extent its l value (shape). |
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Term
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Definition
OUr first encounter with energy-level splitting occurs with lithium (Li; Z = 3). This first two electrons of Li fill its 1s orbital, so the third Li electron must go into the n=2 level. But, this level has 2s and 2p sublevels: which does the third electron enter? For reasons we discuss later, 2s is lower in energy than the 2p, so the ground state of Li has its third electron in the 2s.
This enegy difference arises from three factors - nuclear attraction, electron repulsion, and orbital shape (i.e. radial probability distribution). Their interplay leads to two phenomina - shielding and penetraion - that occur in all atoms except hydrogen.
Remember: more energy is needed to remove an electron from a more stable (lower energy) sublevel than from a less stable (highrer energy) sublevel.] |
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Term
| The Effect of Nuclear Charge (Z) on Sublevel Energy |
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Definition
| Higher charges interact more strongly than lower charges (Coulomb’s law, Section 2.7). Therefore, a higher nuclear charge increases nucleus-electron attraction and thus, lowers sublevel energy (stabilizes the atom). We see this effect by comparing the 1s sublevel energies of three species with one electron - H atom (Z=1), He+ ion (Z=2), and Li2+ ion (Z=3). Figure 8.2 shoes that the 1s sublevel in H is the least stable (hiest energy), so the least energy is needed to remove its electron; and the 1s sublevel in Li2+ is the most stable, so the most energy is needed to remove its electron. |
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Term
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Definition
| The effecto Electron repulsion on Sublevel Energy IN many electrons atoms, each electron “feels” not only the attraction to the nucleus but also repulsion form other electrons. Repulsion conteract the nuclear attraction some what, making of the electron “shielding” the other electrons to some some extent form the nuclear charge. Shielding (also called screning) reduces the full nuclear charge to an effective nuclear charge Z(eff), the nuclear charge an electron actually experiences, and this lower nuclear charge makes the electron easier to remove. |
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Term
| Shielding by other electrons in a given energy level |
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Definition
| Electrons in the same energy level shield each other somewhat. Compare the He atom and He+ ion: both have a 2+ nuclear charge, He has two electrons in the 1s sublevel and He+ has only one (figure 8.3A). It takes less than half as much energy to remove an electron from He than from He+ because the second electron in He repels the first, in effect causing a lower Z(eff). |
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Term
| Shielding by electrons in inner energy levels |
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Definition
| Becaus inner electrons spend nearly all their time between the outer electrons and the nucleus, they cause a much lower Z(eff) than do electrons in the same level. We can see this by comparing the atomic systems with the same nucleus, one with inner electron the other without. THe ground state Li atom has two inner (1s) electrons and one out (2s) electron, while the Li2+ ion has only one electron, which occupies the 2s orbital in teh first 2s electron from the Li atom (520 kJ/mol) as it takes to remove it from the Li2+ ion, because the inner electrons shield very effectively. |
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Term
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Definition
The Effect of Orbital Shape on Sublevel Energy. To see why the third Li electron occupies the 2s sublevel rather than the 2p, we have to consider orbital shapes, that is, radial probability distributions. A 2p orbital is slightly closer the nucleus, on average, than the major portion of the 2s orbital. But a small portion of teh 2s radial probability distribution peaks within the 1s region. Thus, an electron in the 2s orbital spends part of its time “penetrating” bery close to the nucleus. Penetration has two effects:
- It increases the nucleas attraction for a 2s electron over that for a 2p electron
-It decreases the shielding of a 2s electron by the 1s electron
Thus, since it takes more energ to remove a 2s electron than a 2p, the 2s sublevel is lower in energy than the 2p. |
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Term
| Splitting of Levels into Sublevels |
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Definition
| In general, penetration and the resulting effects on shielding cause an energy level split into sublevels of differing energy. The lower the l value of a sublevel, the more its electrons penetrate, and so the greater their attraction to the nucleus. Therefore, for a given n value, a lower l value indicates a more stable (lower energy) sublevel:
Order of sublevel energies: s |
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Term
| The electron configuration |
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Definition
| the shorthand notation consists of the principal energy level (n value), the letter designation of the sublevel (l value), and the number of electrons (#) in the sublevel, written as a superscript: nl^#. |
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Term
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Definition
| An orbital diagram consists of a box for each orbital in a given energy level, grouped by sublevel, with an arrow representing an electron and its spin |
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Term
| Hydrogen: electron configuration |
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Definition
For the electron in H, as you’ve seen, the set of quantum numbers is H (Z=1): n = 1, l=0, m(l) = 0, m(s) = +(1/2). The electron configuration:L
H(Z=1) 1s |
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Term
| Helium: quantum numbers, electron configuration |
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Definition
He (Z=2): n = 1, l=0, m(l)=0, m(s) = -1/2, 1/2
He (Z=2) 1s^2 |
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Term
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Definition
the exclusion principle says an orbital can hold no more than two electrons. Therefore, with He, the 1s orbital, the 1s sublevel, the n=1 level, and Period 1 are filled.
Filling the n =2 level builds up Period 2 and begins with the 2s sublevel, which is the next lowest in energy and consists of only the 3s orbital. When the 2s is filled we proceed to fill the 2p. |
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Term
| Lithium: quantum numbers, electron configuration |
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Definition
n=2 l=0, m(l) = 0, m(s) = +1/2. the electron configuration and orbital diagram are
Li (Z=3) 1s^2 2s^1 |
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Term
| Berilium: quantum numbers, electron configuration |
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Definition
n =2, l = 0, m(l) = 0, m(s) = -1/2
Be (Z=4) 1s2 2s2 |
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Term
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Definition
The next lowest energy sublevel is the 2p. A p sublevel has l = 1, so teh m(l) (orientation) values can be -1, 0, or +1. The three orbirtals in teh 2p sublevel have equal energy (same n and l values) which means that the fifth electron of boron can go into any one of the 2p orbtals.
n=2, l=1, m(l) = -1, m(s) = +1/2
B(Z=5) 1s2 2s2 2p1 |
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Term
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Definition
The Period 3 elements, Na through Ar, lie directly under the Period 2 elements, Li through Ne. That is, even though the n = 3 level plits into 3s, 3p, and 3d sublevels, Period 3 fills only 3s and 3p; as you’ll see shortly, the 3d is filled in Period 4. Table 8.2 introduces two ways to present electron distributions more concisely:
-Partial orbital show only the sublevels being filled, here the 3s and 3p.
-Condensed electron configurations (rightmost column) have the element symbol of teh previous noble gas in brackets, to stand for its configuration, followed by the electron configuration of filled inner sublevels and the energy level being filled. For example, the condensed electron cofiguration of sulfur is [Ne] 3s2 3p4, where [Ne] stands for 1s2 2s2 2p6
In Na and Mg, the electrons are added to the 3s sublevel, which contains only the 3s orbital; this is directly comparable to the filling of the 2s sublevel in Li and Be in Period 2. then, in the same way as the 2p orbitals of B, C, and N in Period 2 are half-filled, the last electrons added to Al, Si, and P in Period 3 half-fill successive 3p orbtals with spins parallel (Hund’s rule). The last electrons added to S, Cl, and Ar then successively pair up to those 3p orbitals, which fills the 3p sublevel. |
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Term
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Definition
The Period 3 elements, Na through Ar, lie directly under the Period 2 elements, Li through Ne. That is, even though the n = 3 level plits into 3s, 3p, and 3d sublevels, Period 3 fills only 3s and 3p; as you’ll see shortly, the 3d is filled in Period 4. Table 8.2 introduces two ways to present electron distributions more concisely:
-Partial orbital show only the sublevels being filled, here the 3s and 3p.
-Condensed electron configurations (rightmost column) have the element symbol of teh previous noble gas in brackets, to stand for its configuration, followed by the electron configuration of filled inner sublevels and the energy level being filled. For example, the condensed electron cofiguration of sulfur is [Ne] 3s2 3p4, where [Ne] stands for 1s2 2s2 2p6
In Na and Mg, the electrons are added to the 3s sublevel, which contains only the 3s orbital; this is directly comparable to the filling of the 2s sublevel in Li and Be in Period 2. then, in the same way as the 2p orbitals of B, C, and N in Period 2 are half-filled, the last electrons added to Al, Si, and P in Period 3 half-fill successive 3p orbtals with spins parallel (Hund’s rule). The last electrons added to S, Cl, and Ar then successively pair up to those 3p orbitals, which fills the 3p sublevel. |
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Term
| Transition elements: Effects of shielding and penetration on sublevel energy. |
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Definition
| The 3d sublevel is filled in Period 4, but the 4s sublevel is filled first. This switch in filling order is due to shielding and penetration effects. based on the 3d radial probability distribution (see figure 7.19, p. 286), a 3d electron spends most of the time outside the filled inner n = 1 and n = 2 lebels, so it is shielded very effectively from the nuclear charge. however, the most outermost 4s electron penetrates close to the nucleas part of teh time, so it is subject to a greater attraction. As a result, teh 4s orbital is slightly lower in energy than the 3d and fills first. In any period, the ns sublevel fills before the (n - 1)d sublevel. Other variations in the filling pattern occur at higer values of n because sublevel energies become very close together. |
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Term
| Transition elements: Filling the 4s and 3d sublevels |
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Definition
| Table 8.3 shows the partial orbital diagrams and full and condensed electron configurations for the 18 elements in Period 4. The first two elements, K and Ca, are the next alkaline earth metals, respectively: the last electron of K half-fills and that of Ca fills the 4s sublevel. The last electron os scandiem (Sc; Z= 21), the first transition element, occupies any one of the five 3d orbitals because they are equal in energy; Sc has the electron configuration [Ar] 4d2 3d1. Filling of the 3d orbitals proceeds one electron at a time, as with p orbitals, except in two cases, chromium (Cr; Z=24) and copper (Cu; Z = 29), discussed next. |
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Term
| Transition elements: Stability of half-filled and filled sublevels |
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Definition
| Vanadium (V; Z =23) has thre half-filled d orbtals (Ar 4s2 3d3). However, the las electron of the next element, Cr, does not eneter a fourth empty d orbital to give Ar 4s2 3d4; instead, Cr has one electron in the 4s sublevel and five in the 3d sublevel, making both sublevels half filled Ar 4s1 3d5 (see margin). In the next element, manganese (Mn; Z = 25), the 4s sublevel is filled again (Ar 4s2 3d5) |
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Term
| Transition elements: Stability of half-filled and filled sublevels 2 |
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Definition
because it follows nickel (Ni; Ar 4s2 3d8), copper would be expected to have the configuration Ar 4s2 3d9. Instead, the 4s sublevel of Cu is half-filled with 1 electron, Cu, we conclude that half-filled and filled sublevels are unexpectedly stable (low in energy); we see this pattern with many other elements.
With zinc (Zn; Z = 30), the 4s sublevel is filled Ar 4s2 3d10, and the first transition series ends. The 4p sublvel is filled by teh next six elements, and Period 4 ends with noble gas krypton (Kr; Z =36). |
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Term
| Similar Outer Electron Configuration Within a Group |
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Definition
| Ang the main-group element (A groups) - the s-block and p-block elements- outer electron configurations within a group are identical. Some variations in the transition elements (B groups, d block) and inner transition elements (f block) occur |
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Term
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Definition
| When the elements are “biult up” by filling levels and sublevels in order of increasing energy, we obtain the sequence in the periodic table. reading the table from left to right, like words on a page, gives the energy order of levels and sublevels (Figure 8.11) |
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Term
| Aid to memorizing sublevels filling order |
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Definition
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Term
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Definition
| are those in the highest energy level 9highest n value). They spend most of their time farthest from the nucleus. |
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Term
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Definition
are those involved in forming compounds.
-For main-groups elements, the valence electrons are the out electrons
- For transition elements, the valence electrons are the outer ns electrons, the (n-1)d electrons are also valence electrons, though the metals Fe (Z = 26) through Zn (Z = 30) may use only a few, if any, of their d electrons in bonding. |
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Term
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Definition
- Among the main-group elements (A groups), the A number equal the number outer electrons (those with the highest n); thus, chlorine (CL; Group 7A) has 7 outer electrons and so on
- The period number is the n value of the highest energy level
-For enter level, the n value squared (n^2) is the number of orbitals, and 2n^2 is the maximum number of electron (or elements). for example, consider the n = 3 level. The number of of orbitals is n^2 = 9: one 3s, three 3p, and five 3d. The number of electrons in 2n^2 = 18: two 3s and six 3p electrons for the eight elemnts. Period 3, nad ten 3d electrons for the ten trransition elements of Period 4. |
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Term
| Intervening Series: transition series |
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Definition
| Periods 4, 5, 6, and 7 incorporate the 3d, 4d, 5d and 6d sublevels , respectively. As you’ve seen, the general pattern is that the (n-1)d sublevel is filled between the ns and np sublevels. Thus, in Period 5, the filling order is 5s, then 4d, and then 5p. |
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Term
| Intervening Series: Inner transition Elements |
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Definition
In period 6, the first of two series of inner transition elements, those in which f orbitals are being filled, appears. The f orbitals have l = 3, so the possible m(l) values are -3, -2, -1, 0, 1, 2, 3, that is, there are seven f orbitals, for a total of 14 elements in each of the two inner transition series:
- The Period 6 inner transition series, called lanthanides (or rare earth), occurs after lanthanum (La; Z =57), and the 4f orbitals are filled
-The Period 7 inner transition series, called the actinides, occurs after actinum (Ac; Z=89), and the 5f orbitals are filled.
Thus i Peiods 6 and 7, the filling sequence is ns, first of the (n-1)d, all (n-2)f, remainder of the (n-1)d, and np
Period 6 ends with the 6p sublevel, Period 7 is incomplete because only five elements with 7p electrons are known at this time; element 117 has not yet been synthesized.
Thus in Periods 6 and 7, the filling sequence is ns, first of the (n-1)d, all (n-2)f, remainder of the (n-1)d, and np.
period 6 ends with 5p sublevel, Period 7 is incomplete because only five elements with 7p electrons are known at this time; element 117 has not yet ben synthesized. |
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Term
| Intervening Series: Irregular filling patterns |
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Definition
| Irregularities in the filling pattern, such as those for Cr and Cu in Period 4, occur in the d and f blocks because the sublevel energies in the larger atoms differ very little. Even though occasional deviations occur in the d block, the sum of ns electrons and (n-1)d electrons always equals the new group number (in parentheses). For instanse, despite variations in Group 6B(6)-Cr, Mo, W, and Sg - the sum of ns and (n-1)d electrons is 6; Groups 8B(10) - Ni, Pd, Pt, and Ds- the sum is 10. |
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Term
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Definition
| Recall from Chapter 7 that we often represent atoms with spherical contours in which the electrons spend 90% of their time. We define atomic size (the extent of the contour) in terms of how closely one atom lies next to another. but, in practice, as we discuss in Chapter 12, we measure the distance between atomic nuclei in a sample of an element and divide that distance in half. Because atoms do not have hard surfaces, the size of an atom in a given compound depends somewhat on the atoms near it. In other words, atomic varies slightly from substance to substance. |
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Term
| Radii - Metallic and Covalent |
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Definition
Metallic radius - Used mostly for metals, it is one-half the distance between nuclei of adjecent, individual atoms in a crystal of the element.
Covalent radius - Used for elements occurring as molecules, mostly nonmetals, it is one-half the shortest distance between nuclei of bonded atoms
Radii measured for some elements are used to determine the radii of other elements from distances between atoms in compounds. For instance, in a carbon-chlorine compound, the distance between nuclei in C-Cl bond is 177 pm. Using the known covaled radius of Cl (100 pm), we find the covalent radius of C (177pm - 100pm = 77pm) |
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Term
| Main-group elements: Changes in n |
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Definition
| As the principal quantum number (n) increases, the probality that outer electron spend most of their time farther from the nucleus increases as well; thus, atomic size increases. |
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Term
| Main-group: Changes in Z(eff) |
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Definition
| As the effective nuclear charge Z(eff) increases, outer electron are pulled closer to the nucleus; thus, atomic size decreases. |
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Term
| Main-groups: down a group a group, n dominates |
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Definition
As we move down a main grouop each member has one more level of inner elecrons that shield the outer electrons very effectively. Even though addditional protons do moderately increase Z(eff) for the outer electrons, the atoms get larger as a result of the increasing n value:
Atomic radius generally increases down a group |
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Term
| Main-groups: Across a period, Z(eff) dominates. |
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Definition
| Across a period from left to right, electrons are added to the same out level, so the shielding by inner electrons does not change. Despite greater electron repulsions, outer electrons shield each ohter only slightly, so Z(eff) rises significantly, and the outer electrons are pulled closer to the nucleus: Atomic radius generally decreases across a period. |
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Term
| Transition elements: Down a transition group |
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Definition
| n increases, but shielding by an additional level of inner electrons results in only a small size increases from Period 4 to 5 and none from 5 to 6. |
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Term
| Transiotion Elements: Across a transition series |
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Definition
| atomic size shrinks through the first two or three elements because of the increasing nuclear charge. But, from then on, size remains relatively constant because shielding by the inner d electrons counteracts the increase in Z(eff). Thus, for example, in Period 4, the third transition element, vanadium (V; Z=23), has the same radius as the last, zinc (Zn; Z = 30). This pattern also appears in Periods 5 and 6 in the transition and both inner transition series. |
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Term
| Transition series affects atomic size in neighbouring main groups |
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Definition
| Shielding by d electrons causes a major size decrease from group 2A(2) to Group 3A(12) in Periods 4 through 6. Because np sublevel has more penetration than the (n-1)d, the first np electron [added in Group 3A(13)] “feels” a musch greater Zeff, due to all the protons added in the intervening transition elements. the greatest decrease occurs in Period 4; calcium (Ca; Z = 20) in Group 2A is nearly 50% larger than gallium (Ga; Z = 31) in 3A(13). In fact, d-robital shielding causes gallium to be slightly smaller than aluminum (Al; Z=13), the element above it! |
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Term
| Trends in Ionization Energy |
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Definition
The ionization energy (IE) is the energy required for the complete removal of 1 mol of electrons from 1 mol of gaseaous atoms or ions. Pulling an electron away from a nucleus requires energy to overcome their electrostatic attraction. Because energy flows into the system, the ionization energy is always positive (like change in H of an endothermic reaction). (Chapter 7 viewed the ionization energy of the H atom s the energy difference between n=1 and n=inf, where the electron is completely removed.) The ioniztion energy is a key factor in an element’s reactivity:
-Atoms with a low IE tend to form cations during reactions
-Atoms with a high IE (except the noble gasses) tend to form anions
Many-electon atoms can lose more than one electron. The first ionization energy (IE1) removes an outermost electron (highest energy sublevel) from a gaseous atom:
Atom(g) -> ion+(g) + e- changeE = IE1>0
The secons ionization energy (IE2) removes a socond electron. Since this electron is pulled away from a positive ion, IE2 is alway larger than IE1:
Ion+(g) -> ion2+ (g) + e- changeE = IE2 (always > IE1) |
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Term
| Periodicity of First Ionization Energy |
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Definition
| Figure 8.15 shows the variation in first ionization energy with atomic number. This up-and-down pattern - IE1 rising across a period to the noble gas (purple) and then dropping down to the next alkali metal - is the inverse of the vatiation in atomic size: as size decreases, it takes more energy to remove an electron because the nucleus is closer, so IE1 increases. |
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Term
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Definition
As we move down a group, the n value increases, so atomic size does as well. As the distance from nucleus to outer electron increases, their attraction lessens, so the electron is easier to remove:
Ionization energy generally decreases down a group.
The only significant exception exception occurs in Group 3A (13): IE1 decreases from boron (B) to alumininum (Al), but not for the rest of the group. Filling the transition series in Periods 4, 5, and 6 causes a much higer Z(eff) and unusually small change in size, so outer electrons in the larger Group 3A members are held tighter. |
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Term
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Definition
As we move left to right across a period, Z(eff) increases and atomic size decreases. The attraction between nucleus and outer electron increases, so the electron is harder to remove:
Ionization energy generally increases across a period.
There are two exceptions to the otherwise smooth increase in IE1 across periods.
-In Periods 2 and 3, there are dips at the Group 3A (13) elements, B and Al. These elements have the first np electrons, which are removed more easily because the resulting ion has a filled (stable) ns sublevel.
- In periods 2 and 3, once again, there are dips at the Group 6A(16) elements, O and S. These elements have a fourth np electron, the first to pair up with another np electron, and electron-electron repulsions raise the orbital energy. The fourth np electron is easier to remove because doing so relieves the repulsions and leaves a half-filled (stable) np sublevel. |
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Term
| Successive Ionization Energies |
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Definition
For a given element IE1, IE2, and so on, increases because each electron is pulled away from a species with a higher positive charge. This increase incluedes an enormous jump after the outer (valence) electrons have been removed because much more energy is needed ro remove an inner (core) electron.
Follow the values of boron (B): IE1 .80MJ is lower than IE2 (2.43 MJ), which is lower (3.66 MJ)which is much lower than IE4 (25.02 MJ). From this jump, we know that boron has three electrons in the highest energy level (1s2 2s2 2p1). Because they are so difficult to remove, core electons are not involved in reactions. |
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Term
| Trends in Electron Affinity |
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Definition
The electron affinity (EA) is the energy change (kJ/mol) accompanying the addition of 1 mol of electrons to 1 mol of gaseous atoms or ions. The first electron affinity (EA1) refers to the formation of 1 mol of monovalent (1-0 gaseaous anions:
Atom(g) + e- -> ion-(g) changeE = EA1
As with ioniztion energy, there is a first electron affinity, a second, and so forth. The first electron is attracted by the atom’s nucleus, so in most cases, EA1 is negative (energy is released), analogous to the negative changeH for an exothermic reaction. But, the second electron affinity (EA2) is always positive because energy must be absorbed to overcome electrostatic repulsions and add another electron to a negative ion.
Factors other than Zeff and atomic size affect electron affinities, so trends are not regular, as are those for size and IE1. The many exceptions arise from changes in sublevel energy and electron-electron repulsion. |
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Term
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Definition
| We might expect a smooth decrease (smaller negative number) down a group because size increases, and the nucleus is farther away from an electron being added. But only Group 1A(1) exhibits this behavior. |
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Term
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Definition
| We might expect a regular (larger negative number() across a period because size decreases and higher Zeff should attract the electron being added more strongly. There is an overall left-to-right increases, but it is not at all regular. |
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Term
| despite these irregulatities. relative values of IE and EA show three general behavior patterns: |
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Definition
1) Reactive nonmetals. Members of Group 6A(16) and especially Group 7A(17) (the halogens) have high IEs and highly negative (exothermic) EAs: these elements lose electrons with difficulty but attract them strongly. Therefore, in their ionic compounds, they form negative ions.
2. Reactive metals. Members of Group 1A(1) and 2A(2) have low IEs and slightly negative (exorthemic) EAs: they lose electrons easily but attract them weakly, if at all. Therefore, in their ionic coumpounds, they form positive ions.
3) Noble gases. have bery high IEs and slightly positive (endothermic) EAs: they tend not to lose or gain electrons. In fact, only the the larger members of group (Kr, Xe, and Rn) form compounds at all. |
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Term
| Trends in Metallic Behavior: Metals |
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Definition
| found in the left and lower three-quaters of the periodic table, are typically shiny solids, hace moderated to high melting points, are good conductors of heat and electricity, can be machined into wires and sheets, and lose electrons to nonmetals. |
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Term
| Trens in Metallic Behavior: Nonmetals |
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Definition
| found in the upper right quarter of the table, are typically not shiny, have relatively low melting points, are poor condutors, are mostly crumbly solids or gases, and tend to gain electrons from metals. |
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Term
| Trends in metallic behavior: metalloids |
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Definition
found between the other two classes, have intermediate properties.
Thus metallic behavior decreases left-to-right across a period and increases hown a group in the periodic table. |
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Term
| Lose or gain Electrons: down a main group |
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Definition
| The increases in metallic behavior doen a group in consistent with an increase in size and a decrease in IE and is most obvious in groups with mre than one class of element, such as Group 5A(15) elements at the top can form anions, and those at the bottom can for cations. nitrogen (N) is a gaseous nonmetal, and phosphorus (P) is soft nonmetal; both occur occasionally as 3- acnions in their compounds. Arsenic (As) and antimony (Sb) are metalloids, with S the more metallic, and neither forms ions redilly. Bismuth (Bi) is a typical metal, forming a 3+ cation in its mostly ionic compounds. Groups 3A(13), 4A(14), and 6A(16) show a similar trend. But even in Group 2A(2), which contains only metals, the tendency to form cations increases down the group: beryllium (Be) forms covalent compounds with nonmetals, whereas all compounds of barium (Ba) are ionic. |
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Term
| Lose or gain Electrons: down a main group |
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Definition
| The increases in metallic behavior doen a group in consistent with an increase in size and a decrease in IE and is most obvious in groups with mre than one class of element, such as Group 5A(15) elements at the top can form anions, and those at the bottom can for cations. nitrogen (N) is a gaseous nonmetal, and phosphorus (P) is soft nonmetal; both occur occasionally as 3- acnions in their compounds. Arsenic (As) and antimony (Sb) are metalloids, with S the more metallic, and neither forms ions redilly. Bismuth (Bi) is a typical metal, forming a 3+ cation in its mostly ionic compounds. Groups 3A(13), 4A(14), and 6A(16) show a similar trend. But even in Group 2A(2), which contains only metals, the tendency to form cations increases down the group: beryllium (Be) forms covalent compounds with nonmetals, whereas all compounds of barium (Ba) are ionic. |
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Term
| Lose or Gain Electrons: across the period |
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Definition
| The decrease in metallic behavior across a period is consistent with a decrease in size, an increase in IE, and a more favorable (more negative) EA. Consider Period 3: elements at the left tend to form cations, and those at the right tend to form anions. Sodium and magnesium are metals taht occur as Na+ adn Mg2+ in seawater, minerals, and organisms. Aluminum is metallic physically and occurs as Al3+ in some compounds, but it bond covalently in most others. silicon (Si) is a shiny metalloid taht does not occur as a monotomic ion. Phosphorus is a white, waxy nonmetal occurs rarely as P3- whereas crumbly, yellow sulfur forms S2- in many compounds, and gaseous, yellow-green chlorine occurs in nature almost always as Cl-. |
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Term
| Redox Behavior of the Main-Grobe Elements |
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Definition
Closely related to an elements’s tendency to lose or gain electrons is its redox behavior - that is, whether it behaves as an oxidizing or reducing agent and the associated changes in its oxidation number (O.N.) . You can find the highest and lowest oxidation numbers of many main-groups elements from the periodic table.
- For most elements, the A-group number is the highest oxidation number (always positive) of any element in the group. The exceptions are O and F.
- For nonmetals and some metalloids, the A-group number minus 8 is the lowest oxidation number (always negative) or any element in the group.
For example, the highest oxidation number of S (group 6A) is +6, as in SF6, and the lowest is 6 - 8, or -2, as in FeS and other metal sulfides. |
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Term
| Redox behavior is closely related to atomic properties: |
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Definition
-With their low IEs and small EAs, the members of Groups 1A(1) and 2A(2) lose electrons readily, so they are strong reducing agents and become oxidized.
-With their high IEs and large EAs, nonmetals in Goups 6A(16) and 7A(17) gain electrons readily, so they are strong oxidizing agents and become reduced. |
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Term
| Acid-Base Behavior of Oxides |
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Definition
Metals are also distinguished from nonmetals by the acid-base behavior of their oxides in water:
- Most main-groups metals transfer electrons to oxygen, so their oxides are ionic. In water, these oxides act as bases, producing OH- ions from O2- and reacting with acids. Calcium oxide is an example.
- Nonmetals share electrons with oxygen, so nonmetals oxides are covalent. They react water to form acids, producing H+ ions and reacting with bases. Tetraphosphorus decaoxide is an example.
Some metals and many metalloids form oxides that are amphoteric: they can act as acids or bases in water
Look at Pg 316 Figure 8.24 |
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Term
| As elements become more metallic down a group (larger size and smaller IE), their oxides become more basic. |
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Definition
In Group 5A(15), dinotrongen pentaoxide, N2O5 forms the strong acid HNO3
N2O3 + H2O -> 2HNO3(aq)
Tetraphosphorus decaoxide, P4o10 forms the weaker acid H3PO4
P4O10 + 6H20 -> 4H3PO4(aq)
The oxide of the matalloid arsenic is wakly acidic, whereas that of the metalloid antimony in weakly basic. Bismuth, the most metallic of the group, forms a basic oxide taht is insoluble in water by reacts with acid to yield a salt and water:
Bi2O3 + 6HNO3 -> 2Bi(NO3)3 + 3H2O |
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Term
| As the elements become less metallic across a period (smaller size and higher IE), their oxides become more acidic. |
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Definition
In Period 3, Na2O and MgO are strongly basic, and amphoteric alumium oxide (AL2O3) reacts with acid or with base:
Al2O3(s) + 6HCl(aq) -> 2AlCl3(aq) + 3H2O(l)
Al2O3(s) + 2NaOH(ap) + 3H2O -> 2NaAl(OH)4
Silcon dioxide is wakly acidic, forming a salt and water base:
SiO2 + 2NaOH -> Na2SiO3 + H2O
the common oxides of phosorus, sulfur, and chlorine from acids of increasing strength H3PO4, H2SO4, and HCIO4. |
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Term
| Ions with a noble gas configuration. |
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Definition
Atoms of the noble gases have very low reactivity because their highest energy level is filled (ns2np6). Thus, when elements at either end of a period form ions, they attain a filled outer level - a noble gas configuration. These elements lie on either side of Group 8A, and their ions are isoelectronic with nearest noble gas.
Elements in Groups 1A and 2A lose electrons and become isoelectronic with the previous noble gas. The Na+ ion, for example, is isoelectronic with neon (NE):
Na(1s2 2s2 3s1) --> e- + Na+(HE 2s2 2p6) [isoelectronic with Ne (He 2s2 2p6)
- Elelments in Groups 6A and 7A gain electrons and become isoelectronic with the next noble gas. The Br- ion, for example, is isoelectronic with krypton (KR):
Br(Ar 4s2 3d10 4p5) + e --> Br- (Ar 4s2 3d10 4p6) |
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Term
| The energy needed to remove electrons from metals or add them to nonmetals determines the charges of the resulting ions |
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Definition
Cations. Remocing antoehr electron from Na+ or from Mg2+ means removing a core electon, which requires too much energy: thus, NaCl2 and MgF3 do not exist.
Anions. Similarly, adding another electron to F- or to O2- means putting it into the next higher energy level (n = 3). with 10 electrons (1s2 2s2 2p6) acting as inner electrons, the nuclear charge would be shielded very effectively, and adding an outer electron would require too much energy: thus, we never see Na2F or Mg3O2. |
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Term
| ions without a noble gas configuration |
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Definition
Except for aluminum, the the metals of Groups 3A to 5A do not form ions with noble gas configurations. Instead, they form cations with two diffferent stable configurations:
- pseudo-noble gas configuration
- Inert pair configuration |
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Term
| Pseudo-noble gas configuration |
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Definition
If the metal atom empties its highest energy level, it attains the stability of empty ns and np sublevels and filled inner (n-1)d sublevel. This (n-1)d10 configuration is called a pseudo-noble gas configuration. For example tin (Sn; Z = 50) loses four electrons to form the tin(IV) ion (Sn4+), which has empty 5s and 5p sublevels and a filled inner 4d sublevel:
Sn (Kr 5s2 4d10 5p2) -> Sn4+ (Kr 4d10) + 4e- |
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Term
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Definition
Alternatively, the metal atom loses just its np electrons and attains a stable configuration with filled ns and (n-1)d sublevels. The retained ns2 electrons are sometimes called an inert pair. For exanple, in the more common tin(II) ion (Sn2+), the atom losses the two 5p electrons and has filled 5s and 4d sublevels:
Sn (Kr 5s2 4d10 5p2) -> Sn2+ (Kr 5s2 4d10) + 2e-
Talllium, lead, and bismuth, the largest and, thus, most metallic members of Groups 3A(13) to 5A(15), form ions that retain the ns2 pair: TI+, Pb2+, and Bi3+. |
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Term
| Electron Configurations of Transition Metal Ions |
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Definition
In contrast to many main groups ions, transition metal ions rarely attain a noble gas configuration. Aside from Period 4 elements scandium, which forms Sc3+, and titanium, which occasionally forms, TI4+, a transition element typically forms more than one cation by losing all of its ns some of its (n-1)d electrons.
The reason, one again, is that energy costs are too high. Let’s condider again the filling of Period 4. At the begining of Period (the same point holds in other periods), penetration makes the 4s sublevel more stable than the 3d. Therefore, the first and second electrons added enter the 4s, which is the outer sublevel. But, the 3d is an inner sublevel, so as it begins to fill, its eletrons are not well shielded from the increasing charge.
A crossover in sublevel energy results: the 3d becomes more stable than the 4s in the transition series. The crossover has a major effect on the formation of Period 4 transtion metal ions: because the 3d electrons are held tightly and shield those in the outer sublevel, the 4s electrons of a transtion metal are lost before the 3d electrons. Thus, 4s electrons are added before 3d electrons to form the atom and are lost before them to form the ion, and so-called “first-in, first-out” rule. |
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Term
| Ion Formation: A Summary of Electron loss or Gain |
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Definition
The various ways that cations form have one point in common - outer electrons are removed first. Here is a sumary of the rules for formation of any main-group or transition metal ion:
- Main-group s-block metals lose all electrons with the highest n value
- Main-group p-block metals lose np electrons before ns electrons.
- Transition (d-block) metals lose ns electrons before (n-1)d electrons.
-Nonmetal gain electrons in the p orbitals of the highest n value. |
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Term
| Magnetic Properties of Tansition Metal Ions |
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Definition
We learn a great deal about an element’s elctron configuration from atom spectra, and magnetic studies provide additional evidence.
Recall that electron spin generates a tiny magnetic field, which causes a beam of H atoms to split in an external magnetic field. Only a beam of a species (atoms, ions, or molecules) with one or more unpaired electrons will split. A beam of silver atoms (Ag; Z=47) was used in the original 1921 experiment
Ag (Z=47) Kr 5s1 4d10
Note the unaired 5s electron. A beam of atoms (Cd; Z = 48) is not split because their 5s electrons are paired Kr 5s2 4d10 |
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Term
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Definition
| A species with unpaired electrons exhibits paramagnetism: it is attracted by an external field. A species with all of its electrons paired exhibits diamagnetism: it is not attracted (and is slightly repelled) by the field. Many transition metals are paramagnetic because their atoms and ions have unpaired electrons. |
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Term
| See examples of Paramegnatism and diamagnatism on pg 319 and 220 |
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Definition
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Term
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Definition
| The ionic radius is a measure of the size of an ion and is obtained from the distance between the nuclei of adjecent ions in a cystalline ionic compound. Form the relation between effective nuclear charge Z(eff) and atomic size, we can predict the size of an ion relative to its parent atom. |
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Term
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Definition
| Cations are smaller than parent atoms. When a cation forms, electrons are removed from the outer level. The resulting decrease in shielding and value of nl allows the nucleus to pull the remaining electrons closer. |
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Term
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Definition
| Anions are larger than parent atoms. When an anion forms, electrons are added to the outer level. The increase in shielding and electron repulsion means the electrons occupy more space |
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Term
| Down and Across a group: Ionic size |
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Definition
Down a group ionic size increase because the n increases
Across a period, for instance, Period 3, the pattern is complex:
- Among cations, the increase in Z(eff) form left to right makes Na+ larger than Mg2+, which is larger than Al3+
- From last cation to first anion, a great jump in size occurs: we are adding electrons rather than removing them, so repulsions increase sharply. For instance, P3- has eight more electrons than Al3+
- Among anions, the increase in Z(eff) from left to right makes P3- larger than S2-, which is larger than Cl-.
- Within an isolelectronic series, these factors have striking results. Within the dashed outline in Figure 8.29, the ions are isoelectronic with neon. Period 2 anions are much larger than Period 3 cations because the same number of eletrons are attracted by an increasing nublear charge. The pattern is
3->2->1->2+>3+ |
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Term
| Cations size decreases with charge. |
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Definition
| when a metal forms more than one cation, the greter the ionic charge, the smaller the ionic radius. WIth the two ions of ron for example, Fe3+ has one fewer electron, so shielding is reduced somewhat, and the same nucleus is attracting fewer electrons. As a result, Z(eff) increases, so Fe3+ is smaller than Fe2+. |
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Term
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Definition
| Metal with nonmetal: electron transfer and ionic bonding. We observe ionic bonding between atoms with large differences in their tendencies to lose or gain electrons. Such differences occur between reactive metals [Groups 1A and groups 2A] and nonmetals [Group 7A and the top of Group 6A]. A metal atom (low IE) loses its one or two valence electrons, and a nonmetal atom (highly negative EA) gains the electrons. Electron transfer from metal to nonmetal occurs, and each atom forms an ion with a noble gas electron configuration. The electrostatic attractions between these positive and negative ions draw them into a three-dimensional array to form an ionic solid. Note that the chemical formula of an ionic compound is the empitical formula because it gives the cation-to-anion ratio. |
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Term
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Definition
| Nonmetal with nonmetal: electron sharing and covalent bonding. when two atoms differ little, or not at all, in their tendencies to lose or gain electrons we observe electron sharing and covalent bonding, which occurs most commonly between nonmetal atoms. Each atom holds onto its own electrons tightly (high IE) and attracts other electrons (highly negative EA). The nucleus of each atom attracts the velence electrons of the other, which draws the atoms together. The shared electron pair is typically localized between the two atoms, linking them in a covalent bond of a particular length and strength. In most cases, seperate molecules result when atoms bond covalently. Note that the chemical formula of a covalent compound is the molecular formula because it gives the actual numbers of atoms of atoms in each molecule. |
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Term
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Definition
| Metal with metal: electron pooling and metallic bonding. Metal atoms are relatively large, and their few outer electrons are well shielded by filled inner levels (core electrons). Thus, they loser outer electrons easily (low IE) and do not gain them readily (slightly negative or positve EA). THese properties lead metal atoms to share their valence electrons, but not by covalent bonding. In the simplest model of metallic bonding, the enormous number of atoms in sample of metal pool their valence electrons into a “see” of electrons that “flows” between and around each metal-ion core (nucleus plus electrons in covalent bonding, electrons in metallic bonding are delocalized, moving freely throughout the entire piece of metal. |
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Term
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Definition
The central idea of the ionic bonding model is the transfer of electrons form metal atoms to nonmetal atoms to form ions that attract each other into a solid compound. In most cases, for the main groups, the ion forms has a filled outer level of either two or eight electrons, the number in the nearest noble gas (octet rule).
The transfer of an electron from a lithium atom to a fluorine atom is depicted in three ways in Figure 9.5. In each, Li loses its single outer electron to fill its n = 2 level (eight e-). In this case, each atom is one electron away form the configuration of its nearest noble gas, so the number of electrons lost by each Li equals the number gained by wach F. Therefore, equal numbers of Li+ and F- ions form, as the formula LiF indicates. That is, in ionic bonding, the total number of electrons lost by the metal atom(s) equals the total number of electrons gained by the nonmetal atom(s) |
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Term
| The electron-transfer process. |
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Definition
Consider the electron-transfer process for the formation of lithium fluoride, which involves a gaseous Li atom losing an electron, and a gaseous F atom gaining it:
- The first ionization energy (IE1) of Li is the energy absorbed when 1 mol of gaseous Li atoms loses 1 mol of valence electrons:
Li(g) --> Li+(g) + e- IE1 = 520 kJ
- The first electron affinity (EA1) of F is the energy released when 1 mol of gaseous F atoms gains 1 mol of electrons:
F(g) + e- --> F-(g) EA1 = -328 kJ
- Taking the sum shows that electron transfer by itself requires energy:
Li(g) + F(g) --> F-(g) IE1 + EA1 = 192 Kj |
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Term
| Other steps that absorb energy. |
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Definition
| The total energy needed prior to ion formation adds to the sum of IE1 and EA1: metallic lithium must be made into gaseous atoms (161 kJ/mol), and fluorine molecules must be broken into separate atoms (79.5 kJ/mol). |
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Term
| Steps that release energy |
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Definition
Despite these endothermic steps, the standard enthalpy of formation (changeH f) of solid LiF is -617 kJ/mol; that is, 617 kJ is released when 1 mol o fLiF(s) forms from its elements. Fomation of LiF is typical of reactions between active metals and nonmetals: ionic solids form readily.
If the overall reaction releases energy, there must be some step that is exothermic enough to outweigh the endothermic steps. This step involves the strong attraction between pairs of oppositely charged ions. When 1 mol of Li+ (g) and 1 mol of F- (g) form 1 mol of gaseous LiF molecules, a large quantity of heat is released:
Li+(g) + F-(g) --> LiF(g) changeH = -755 kJ
But, as you knoe, under ordinary conditions, Li does not exist as gaseous molecules: even more energy is released when seperate gaseous ions coalesce into a crystaline solid because each ion attracts several oppisitely charged ions:
Li+ (g) + F- (g) --> LiF(s) changeH = -1050 kJ
The negative of this enthalpy change is 1054 kJ, the lattice energy of LiF |
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Term
| Lattice energy (changeH lattice) |
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Definition
| is the enthalpy change that accompanies the reverse of this equation - 1 mol of ionic sold separating into gaseous ions. |
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Term
| Determining Lattice Energy with a Born-Haber Cycle |
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Definition
| The magnitude of the lattice energy is a measure of the strength of the ionic interactions and influences macroscopic properties, such as melting point, hardness, and solubility. Despite playing this crucial role in the formation of ionic compounds, lattice energy cannot be measured directly. One way to determine it applies Hess’s law in a Born-Harber cycle, a series of steps from elements to ionic solid for which all the enthalpies are known except the lattice energy. |
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Term
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Definition
Step 1. From solid Li to Li atoms. This step, called atomization, has the enthalpy change changeH-atom. It involves breaking metallic bonds, so it absorbs energy:
Li(s) --> Li(g) changeH-step1 = changeH-atom = 161 kJ
Step 2. Form F2 molecules to F atoms. This step involves breaking a convalent bond, so it absorbs energy; as we discuss later, this is the bond energy (BE) of F2. Since we nee 1 mol of F atoms to make 1 mol of LiF, we start with 1/2 mol of F2:
1/2mF2(g) --> F(g) changeH-step2 = 1/2(BE of F2) = 1/2(159 kJ) = 79.5 kJ
Step 3. From Li to Li+. Removing the 2s electron from li absorbs energy:
Li(g) --> Li+(g) + e- changeH-step3 = IE1 = 520 kJ
Step 4. From F to F-. Adding an electron to F releases energy:
F(g) + e- -> F-(g) changeH-step4 = EA1 = -328 kJ
Steps 5. From gaseous ions to ionic solid. Forming solid LiF from gaseous Li+ and F- releases a lot of energy. The enthalpy change for step is unknown but, by definition, the negative of the lattice energy.
Li+(g) + F-(g) --> LiF(s) changeH-step5 = -changeH-lattice of LiF = ?
The enthalp change of the combination reaction is
Li(s) + 1/2F2(g) --> LiF(s) changeH-overall = changeH-f = -617
Add all the energy together and subtract from -617 kJ = -1050 kJ
changeH-lattice = -(-1050 kJ) = 1050 kJ
The Born-harber cycle shows taht the energy required for elements to form ions is supplied by the attraction among the ions in the solid. And the “take-home lesson” is that ionic solids exist far exceeds the total energy needed to form the ions. |
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Term
| Periodic trends in lattice Energy |
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Definition
| The lattice energy results from electrostatic interactions among ions, so its magnitude depends on ionic size, ionic charge, and ionic arrangement in the solid. Therefore, we expect to see periodic trends in lattice energy. |
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Term
| Explaining the Trends with Coulumb’s Law |
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Definition
Recall from Chapter 2 that Coulomb’s law states that the electrostatic energy between particles A and B is directly proportional to the product of their charges and inversley proportional to the distance between them:
Electrostatic energy = (charge A x charge B)/distance
Latice energy is directly proportional to electrostatic energy. In an ionic solid, cations an danions lie as close to each other as possible, so the distance them is the sum of ionic radii
Electrostatic energy = (cation charge X anion charge) /(cation radius + anion radius) = changeH-lattice |
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Term
| Lattice Energy: Effect of ionic size |
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Definition
| As we move down a group, ionic radii increase, so the electrostatic energy between cations and anions decreases; thus, letice energies should decrease as well. Figure 9.8 on the the next page shows that, for the alkali-metal halides, lattice energy decreases down the group whether we hold the cation constant (LiF to LiI) or the anion constant. |
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Term
| Lattice Energy - Effect of ionic charge |
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Definition
Acros a period, ionic charge changes. For example, lithium fluoride and magnesium oxide have cations and anions of about of about equal radii (Li+ = 76 pm and Mg2+ = 72 pm; F- =133 pm and O2- = 140 pm). The major difference is between singly charge Li+ and F- ions and doubly charged Mg2+ and O2- ions. The difference in the lattice energies of the two compounds is striking:
changeH-lattice of LiF = 1050kJ/mol and changeH-lattice of MgO = 3923kJ/mol
This nearly fourfold increase in changeH-latice relects the fourfold increase in the product charge (1 x 1 vs. 2x2) in the numerator of Equation 9.1 |
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Term
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Definition
we might ask how ionic solids, like MgO, with its doubly charged ions, could even form. After all, forming 1 mol of Mg2+ involves the sum of the first and second ionization energies.
And, while 1 mol of O- ions is exothermic (the electron affinity, EA1), adding a second mole of electrons (EA2) is endothermic because the electron is added to a negative ion.The overal formation of O2- ion is endothermic
There are also endothermic steps for converting Mg(s) (148 kJ/mol and breaking 1/2 mol of O2 molecules into o atoms (498 kJ/mol)l Nevertheless, the 2+ and 2- ionic make the lattice energy so large that solid MgO forms readily whenever Mg burns in air. |
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Term
| Ionic Model: Physical behavior |
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Definition
| As a typical ionic compound, a piece rock sal (NaCl) is had (does no dent), rigid (does not bend), and brittle (crack without deforming). These properties arise form strong attractive forces that hold the ions in specific positions moving them out of position requires overcoming these forces, so rock salt does not dent or bend. If enough for is applied, ions of like are brought net to each other, and repulsion between them crack the sample suddenly. |
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Term
| Electrical conductivity: Ionic model |
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Definition
| Ionic compounds typically do not conduct electricity in the solid state but do conduct when melted or dissolved. According to the model, the solid consists of fixed ions, but when it melts or dissolves, the ions can move and carry a current. |
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Term
| Thermal conductivity: Ionic model |
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Definition
| Large amounts of energy would be needed to free the ions from their positions and seperate them. THus, we would expect ionic compounds to have high melting points and much higher boiling poiints. In fact, the interionic attraction is so strong that the vapor consists of ion pairs, gaseous ionic molecules, rather than individual ions. In their normal state, as you know, ionic compounds are solid arrays of ions, and no separate molecules exist |
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Term
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Definition
arises from the balance between the nuclei attracting the electrons and electrons and nuclei repelling each other
Formation of a covalent bond always rusults greater electron density between the nuclei. Figure 9.13 this fact with a cross section of a space-filling model (A), an electron density contour map (B), and an electron density relief map (C). |
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Term
| Bonding Pairs and Lone Pairs |
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Definition
to achieve a full outer (valence) level of electrons, each atom in a covalent bond “counts” the shared electrons as belonging entirely to itself. Thus, the two shared electron in H2 siultaneously fill the outer level both H atoms. making a bond pair.
An outer-level electron pair that is not involved in bonding is called a lone pair or unshared pair. The bonding pair iin HF fills the outer level of the H atom and, together with three lone pairs, fills the outer level of the F atom as well.
In F2 the bonding pair and three lone pairs fill the outer level of each F atom. |
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Term
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Definition
| is the number of electron pairs being shared by a given pair of atoms |
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Term
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Definition
| as shown before in J2, HF, or F2, is the most common bond and consists of one bonding of electrons. Thus, a single bond has a bond order of 1. |
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Term
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Definition
| Many molecules (and ions) contain multiple bonds, in which more than one pair is shared between two atoms. Multiple bonds usually involve C, O, and/or N atoms. A double bond consists of two bonding electron pairs, for electrons shared between two atoms, the bond order is 2. Ethylene (C2H4) contains a carbon-carbon double bond and four carbon-hydrogen single bonds. |
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Term
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Definition
| consists of three shared pairs: two atoms share six electrons, so the bond order is 3. The N2 molecule has a triple bond, and each N atom also has a lone pair. Six shared and two unshared electrons give each n atom an octet. |
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Term
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Definition
The strength of a covalent bond depends on the magnitude of the attraction between the nuclei and shared electrons. The bond energy (BE) (also called bond enthalpy or bond strength) is the energy needed to overcome this attraction and is defined as the standard enthalpy change for breaking the bond in 1 mol of gaseous molecules. bond breakage is endotermic process, so bond energy is always positive:
The bond energy is the difference in energy between separated and bonded atoms (the potential energy difference between points 1 and 3, the energy “well” in Figure 9.12) THe same quantity of energy absorbed to break the bond is released when the bond forms. Bond formation is an exothermic process, so the sign of its enthalpy change is always negative
-Stronger bonds are lower in energy (have a deeper energy well).
- Weaker bonds are higher in energy (have a shallower energy well |
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Term
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Definition
| A covalent bond has a bond legnth, teh distance between the nuclei of two bonded atoms. In Figure 9.12 bond length is the distance between the nuclei at at the point of minimum energy (bottom of the well), and Table 9.2 shos the lengths of some covalent bonds. Like bond energies, these values are average bond lengths for a bond in different substances. Bond length is related to the sum of the radii of the the bonded atoms. In fact, most atomi radii are calculated from mesured bond lengths. nond lengths for a series of similar bonds as in teh halogens, increase with atomic size. |
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Term
| The order, energy, and length of a covalent bond are interrelated. |
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Definition
Tow nuclei are more strongly attracted to two shared pairs than to one, so double-bonded atoms are drawn closer together and more difficult to pull apart than single-bonded atoms for a given pair of atoms, a higher bond order results in a shorter bond length and for a given pair of atoms, a higher bond order results in a shorter bond length and a higher bond energy. Thus, as Table 9.3 shows, for given pair of atoms, a shorter bond is a stronger bond.
In some cases, we can see a relation among atomic size, bond length, and bond energy by varying one of the atoms in a single bond while holding the other constant:
Variation within a group. the trend in carbon-halogen single bond lengths, C-I > C-Br > C-Cl, parallels the trend in atomic size I>Br>Cl, and is opposite to the trend in bond energy, C-Cl > C-Br > C-I.
- Variation within a period. Looking at single bonds involving carbon, the trend in bond lengths C-N > C-O > C-F, is opposite to the trend in bond energy. |
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Term
| Physical properties of molecular covalent substances. |
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Definition
At first glance, the model seem inconsistent with physical properties of covalent substances. MOst are gases (such as methane and ammonia), liquids (such as benzene and water), or low-melting solids (such as sulfur and paraffin wax). If covalent bonds are so strong (-200 to 500 kJ/mol), why do covalent substances melt and boil at such low temperatures?
To answer this, we’ll focus on two different forces: 1) strong bonding forces hold the atoms together within the molecule, and 2) weak intermolecular forces act between separate molecules in the sample. It is teh eak forces betwee molecules that account for the physical properties of molecular covalent substances. For example, look what happens when pentane (C5H12) boils: weak forces between pentane molecules are overcaome, not the the strong C-C and C-H bonds within each pentance molecule. |
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Term
| Physical properties of network covalent solids. |
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Definition
| Some covalent substances do not consist of separate molecules. Rather, these network covalent solids are held together by covalent bonds between atoms throughout the sample, and their properties do reflect the strength of covalent bonds. Tow examples are quartz and diamond. Quartz (SiO2) has silicon-oxygen covalent bonds in three dimension; no separate SiO2 moleculs exist. It is very hard and melts at 1550 C. Diamond has covalent bons connecting each carbon atom to four others. It is the hardest natural substance known and melts at around 3550 C. thus, covalent bonds are strong, but most covalent substances consist of separate molecules with weak forces between them |
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Term
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Definition
| An electric current is carried by either mobile electrons or mobile ions. MOst substances are poor electrical conductors, whether melted or dissolved, because their electrons are localized as either shared or unshaired pairs, and no ions are present. |
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Term
| Changes in Bond Energy: Where Does changeH-rxn come from |
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Definition
In Chapter 6, we discussed the heat involved in a chemical change but never asked a central question: where does the enthalpy of reaction changeH-rxn come from? For example, wen 1 mol H2 and 1 mol of F2 react to form 2 mol of HF at 1 atm and 296 K,
H2(g) + F2(g) -> 2HF(g) + 546 of kJ
where does the 546 kJ come from? We find the answer by looking closely at the energies of the molecules involved. A system’s total internal energy is composed of its kinetic and potential energy.
The ansewr to “where does changeH-rxn come from?” is that it doesn’t really “come from” anywhere: the heat released or absorbed during a chemical change is due to differences between reactant bond energies an dproduct bond energies. |
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Term
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Definition
| The contibutions to the kinetic energy ate the molecules’ movements in space, rotations, and vibrations. however, since kinetic energy is proportional to temperature, which is constant at 298 K, it doesn’t change during the reaction. |
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Term
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Definition
| THe most important contributions to the potential energy are phase changes and changes in the attraction between vibration atoms, between nucleus and electrons (and between electrons) in each atom, between protons and neutrons in wach nucleas, adn between nuclei and the shared electron pair in each bond. Of these, there are no phase changes, vibrational forces vary only slightly as the bonded atoms change, and forces within the atoms and nuclei don’t change at all. THe only significant change in potential energy comes form changes in the attraction between the nuclei and the shared electron pair - the bond energy |
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Term
| Using Bond Energies to Calculate changeH-rxn |
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Definition
Hass’s law allows us to think of any reaction as a two-step process, whether or not it actually occurs that way:
1) A quantity of heat is absorbed (changeH > 0) to break the reactant bonds and form form separate atoms.
2) A different quantity of heat is then released (changeH < O) when the atoms form product bonds.
The sum of these enthalpy changes is the enthalpy of reaction, changeH-rxn: equation 9.2
In a exothermic reaction, the magntide of changeH-(product bonds formed) is greater than of the changeH-(product bonds broken), so the sum, changeH-rxn, is negative (heat is released).
In endothermic reaction, the opposite situation is true, the magnitude of changeH-(product bonds formed) is smaller than that of changeH-(reactant bond boken), so changeH-rxn is positive (heat is absorbed). |
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Term
| Typically, only certain bonds break and form during a reaction, but with Hess’ law, the following is a simpler method fo calculating changeH-rxn: |
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Definition
1) Breal all the reactant bonds to obtain individual atoms.
2) Use the atoms to form all the product bonds.
3) Add the bond energies, with appropriate signs, to obtain the enthalpy of reaction.
(This method assumes reactants and products do not change physical state; additional heat is involved when phase changes occur) |
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Term
| Formation of HF, changeH-rsn |
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Definition
When 1 mol H-H bonds and 1 mol of F-F bonds absorb energy and break, the 2 mol of H atoms and 2 mol of F atoms form 2 mol of H-F bonds, which releases energy (Figure 9.17). We find the bond energy values in Table 9.2 (p. 340) and use a positive sign for bonds broken and a negative sign for bonds formed:
Bronds broken: (H-H) + (F-F) = 432kJ + 159kJ = 591
Bonds formed: 2(-565kJ/mol) = -1130kJ
Applying Equation 9.2 gives: changeH-rxn = sum changeH-(reactant bonds broken) +sumchangeH-(product bonds formed)
= 591kJ + (-1130 kJ) = -539kJ
The small discrepancy this bond energy balue (-539 kJ) and the value form tabulated changeH values (-546 kJ) is due to variations in experimental method. |
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Term
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Definition
In this more complicated reaction, all the bonds in CH4 and O2 break, and the atoms from all the bonds in CO2 and H2O. One again we used Table 9.2 and approptiate signs for bonds broken and bonds formed:
Bonds broken:
4x(C-H)+2x(O2) = (1652kJ)+ (996 kJ) = 2648 kJ
Bonds formed:
2x(C=O) + 4x(O-H) = (-1598 kJ)+(-1868 kJ) = -3466 kJ
Applying Equation 9.2 gives
changeH-rxn = 2648 kJ + (-3466 kJ) = -818 kJ |
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Term
| Bond Strengths and the Heat Released from Fuels and Foods |
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Definition
| A fuel is a material that reacts with atmospheric oxygen to release energy and is available at a reasonable cost. The most common fuels for machines are hydrocarbons and coal, and the most common ones for organisms are fats and carbohydrates. All of these fuels are composed of large organic molecules many C-C and C-H bonds adn fewer C-O and O-H bonds. According to our two-step approach, when the fuel reacts with O2, all of the bonds break, and the C, H, and O atoms form C=O and O-H bonds in the products(CO2 and H2O). Beacause their combustion is exothermic, the total of the bond energies in the products is greater than the total in the reactants. Weaker bonds (less stable, more reactiver) are easier to break than stronger bonds (more stable, more reactive) are easier to break than stronger bonds (more stable, less reactive) because they are already higher in energy. THerefore, the bonds in CO2 and H2O are stronger (lower energy, more stable) than those in gasoline (or cooking oil) and O2 (weaker, higher energy, less stable). |
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Term
| Fuels with more weak bonds yield more energy than fuels with fewer weak bonds |
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Definition
Why a hydrocaarbon burns, C-C and C-H bonds break; when a alcohol burns, C-O and O-H bonds break as well. Table 9.2 (p. 340) shows that the sum for C-C and C-H bonds (760 kJ/mol) is less than the sum for C-O and O-H bonds (825 kJ/mol). Therefore, it takes more energy to break the bonds of a fuel with a lot of C-O and O-H bonds. Thus, in general, a fuel with fewer bonds to O releases more energy.
Both fats and carbs serve as high-energy foods and consist of chains or of C atoms attaches to H atoms, with some C-O, C=O, and O-H bonds
Carbohydrates have fewer chains of C atoms and bonds and more bonds to O. Fats “contain more Calories” per gram than carbohydrates because fats have fewer bonds to O |
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Term
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Definition
| Electronegativity (EN) is the relative ability of a bonded atom to attract shared electrons. We might expect the H-F bond energy to be the average of an H-H bond (432 kJ/mol) and an F-F bond (159 kJ/mol), or 296 kJ/mol. Nut, the actual HF bond energy is 565 kJ/mol, or 269 kJ.mol higher! to explain this difference, the American chamist Linus Pauling reasoned this way: If F attracted the shared electron pair more strongly than H, that is, if F were more electronegative than H, the electrons would spend more time closer to F, and this unequal sharing would make the F end of the bond partially negative an dthe H end partially positive. The electrostatic attraction bond partially negative and the H end partially positive. The electrostatic attraction between these partial charge bond “poles” would increase the energy required to break the bond. From studies with many other compounds, Pauling derived a scale of relative EN values based on fluorine having the highest EN value, 4.0. |
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Term
| Trends in Electronegativity |
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Definition
In general, electronegativity is inversely related to atomic size because the molecules of a smaller atom is closer to the shared pair than the nuclueus of a larger atom, so it attracts the electrons more strongly.
- Down a main group, electonegativity decreases as size increases.
-Across a period of main-group elements, electronegativity increases
- Nonmetals are more electronegative than metals.
The most electronegative element is fluorine, with oxygen a close second. Thus, except when it bond with fluorine, oxygen always pulls bonding electrons toward itself. The least electronegative element (also referred to as the most electropositive) is francium, in the left corner of the periodic table, but it is radioactive and extremely rare, so for all practical purposes, cesium is the most electropositive. |
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Term
| Eletctronegativity and Oxidation Number |
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Definition
An important use of electronegativity is in determining an atom’s oxidation number
1) the more electronegative atom in a bond is assigned all the shared electrons, the less electronegative atom is assigned none.
2) each atom in a bond is assigned all of its unshared electrons.
3) The oxidation number is given by
O.N. = no. of valence electrons - (no. of shared electrons + no. of unshared electrons)
In HCl, for example, Cl is more electronegative tahn H. Cl has 7 valence electrons and is assigned 8 (2shared + unshared), so its O.N, is 7-8 = -1. TH eH atom has 1 valence electron and is assigned none, so its O.N. its 1-0 = +1 |
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Term
| bond polarity and Partial Ionic Character |
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Definition
Whenever atoms of different electronegativities form a bond, such as H (2.1) and F (4.0) in HF, the bonding pair is shared unequally. This unequal distribution of electron density results in a polar covalent bond. It depicted by a polar arrow (+->) pointing toward the negative pole or by o- and o+ symbols
In the H-H an dF-F bonds, where the atoms are identical, the bonding pair is shared equally, and a nonpolar covalent bond results. In Figure 9.23, relief maps show the distribution of electron density in H2, F2, and HF. |
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Term
| The electonegatibity difference changeEN |
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Definition
the difference between the EN valeus of the bonded atoms, is directly related to a bond’s polarity. It ranges from .0 in a diatomic element, such as H2, O2, or Cl2, all the way up to 3.3, the difference between the most electronegative atom, F(40), and the most electropositive, CS (.7), in the ionic compound CsF
Another parameter closely relatedto changeEN for LiCl(g) is 3.0-1.0 = 2.0; for HCl(g), it is 3.0-2.1 = 0.9; and for Cl2(g), it 3.0-3.0 = 0.0 Thus, the bond in LiCl has more ionic character than the one in HCL, which has more than one in Cl2. |
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Term
| two approaches that quantify character. Both use arbitrary cutoffs, which is not really consistent with the actual gradation in bonding. |
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Definition
1) changeEN range. this approach divides bonds into mostly ionic, polar covalent, mostly covalent, an dnonpolar covalent based on a range of changeEN values.
2) Percent ionic character. This approach is based on the behavior of a diatomic molecule in an electric field. A plot of percent ionic character vs. changeEN for several molecules shows that, as expected, percent ionic character character generally increases with changeEN. A value of 50% divides ionic from covalent bonds. Note that a substance like Cl2(g) has 0% ionic character, but none has 100% ionic character. electron sharing occurs to some extent in every bond, even in an alkali halike. |
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Term
| The Gradation in Bonding Across a Period |
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Definition
A metal an da nonmetal-elements from the left an dright sides of teh periodic table - have a relatively large changeEN and typically form an ionic compound. Two nonmentals - both from the right side of the table - have a small changeEN and form a covalent compound. When we combine chlorine with each of the other Period 3 elemnts, starting with sodium, we observe a steady decrease in changeEN and gradation in bond type from ionic through polar covalent to nonpolar covalent.
look at figure 9.26 |
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Term
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Definition
Metals can transfer electrons to nonmetals and form ionic solids, such as NaCl. And experiments with metals in the gas phase show that two metal atoms can even share their valence electrons to form gaseous, diatomic molecules, such as Na2. But waht holds the atoms together in a chunk of Na metal? THe electron-sea model of metalic bonding proposes that all the metal atoms in the sample contribute their valence electrons to form a delocalized electron “sea” throughout the piece, with the metal ions (nuclei and core electrons) lying in an orderly array. All the atoms in the sample share the electrons, and the piece is held together by the mutual attraction of the metal cations for the mobile, valence electrons. Thus, bonding in metals is fundamentally different than the other two types:
- In contrast to ionic bonding, the metal ions are not hel in place as rigidly.
- In contrast to covalent bonding, no particular pair of metal atoms is bonded through a localist electron pair.
- Instead of forming compounds, two or more metals typically form alloys*, solid mixtures of variable composition. Alloys appear in car and airplane bodies, bridges, coins, jewelry, dental fillings, and many other familiar objects. |
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Term
| Metals: melting point and boiling points |
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Definition
Nearly all metals are solids with moderate to high melting points and much higer boiling points. These properties are related to the energy of the metallic bonding. Melting points are only moderately high because the cations can move without breaking th eattraction to the surrounding electrons. Boiling points are very high because each cation ant its valence electrons must break away from the others. Gallium provides a striking example: it can melt in your hand but doesnt boil until over 2400 C.
-Down a group, melting points decrease because the larger metal ions have a weaker attraction to the electron sea.
- Across a period, melting point increase. Alkaline earth metals [Group 2A(2)] have higher melting points than alkali metals [Group 1A(1)] because their 2+ cations have a stronger attractions to twice as many valence electrons. |
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Term
| Metals: Mechanical properties |
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Definition
When a piece of metal is deform by a hammer, the meions do not repel each other, but rather slide past each other through the electron sea and end up in new positions; thus metals dent and bent.
All the Group 1B(11) metals - copper, silver, and gold - are soft enough to be mechined into sheets (malleable) and wires (ductile), but gold is in a class of itself. |
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Term
| Metals: Electrical conductivity |
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Definition
| unlike ionic and covalent substances, metals are good conducters of electricity in both the solid and liquid states because of their mobile terminal electrons. When a piece of metal wire is attaches to a bettery, electrons from from one terminal into the wire, replacing the electrons that flow from the wire into the other terminal. Foreign atoms disrupt the array of metal atoms and reduce conductivity. Copper used in electrical wiring is over 99.99% pure because traces of other atoms drastically restrict electron flow. |
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Term
| Metal Thermal conductivity |
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Definition
| mobile electrons also make metals good conductors of heat. Place your hand on a piece of metal and a piece of wood that are both at room temp. THe metal feels colder becaus it conducts heat from your hand much faster than the wood. THe mobile, delocalized electrons in the metal disperse the heat from hand more quickly than the localized electron pairs in the covalent bonds of wood. |
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Term
| Common bonds created by hydrogen, carbon, nitrogen, oxygen, surrounding halogens form one bond, fluorine is always a surrounding atom |
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Definition
- Hydrogen atoms form one bond
- Carbon atoms form four bonds.
- Nitrogen atoms form three bonds
- Oxygen atoms form three bonds
- Surronding halogens form one bond; fluorine is always a surrounding atom |
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Term
| Lewis structure for molecules with single bonds |
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Definition
1) Place the atoms relative to each other. The atom with the least EN is in the middle. if the atoms have the same group number, place the atom with the higher period number (lowest EN) in the center.
2) Determine the total number of valence electrons. For polyatomic ions, add one e- for each negative charge, or subtract one e- for each positive charge
3) Draw a single bond from each surrounging atom to the central atom, and subtract, 2e- form the total for each bond to find the number of e- remaining
4) Distribute the remaining electrons in pairs so that each atom ends up with 8e-. If any electrons remain, place them around the central atom until it reaches eight. |
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Term
| Molecules with Multiple bonds |
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Definition
In most cases, if there are not enough electrons for the central atom(s) to attain an octet, a multiple bod is present, and we add the following step to the procedure for writing a Lewis structure:
5) Cases involving multiple bonds. If a central atom does not end up with an octet, change a lone pair on a surrounding atom into another bonding pair to the central atom, thus forming a multiple bond. |
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Term
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Definition
two lewis structures and a two-headed resonance arrow (<->) between them. Resonance structures have the same relative placement of atoms but different locations of bonding and lone electron pairs. page 367 (pic)
Resonance structures are not real bonding depictions. THe actual molecule is a resonance hybrid, a average of the resonance forms. |
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Term
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Definition
| Our need for more tahn one Lewis structure to depict O3 is due to electron-pair delocalization. In a single, double, or triple bond, each electron pair is localized between the bonded atoms. In a resonance hybrid, two of the electron pairs (one bonding and one lone pair) are delocalized: their density “spread” over a few adjacent atoms. (this delocalization involves just a few e- pairs, so it is much less extensive than the electron delcalization in metal that we saw in 9.6) |
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Term
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Definition
total number of valence electrons minus all of its unshared valence electrons and half of its shared valence electrons.
Formal charges must sum to the actual charge on the species: zero for a molecule or the ionic charge for an ion.
Only when an atom has a zero formal charge does it have its usual number of bons; the same hold for C in CO3(2-), N in NO3-, and so forth. |
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Term
| Choosing the More Important Resonance Form |
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Definition
- smaller formal chages (positive or negative) are preferable to larger ones
- The same nonzero formal charges on adjacent atoms are not preferred.
- A more negative formal charge should reside on a more electronegative atom. |
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Term
| Formal Charge Versus Oxidation Number |
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Definition
Formal charge (used to examine resonance structures) is not the same as oxidation number (used to monitor redo reactions):
- For a formal charge, bonding electrons are shared equally by the atoms (as if the bonding were nonpolar covalent), so each atom has half of them:
Formal charge = valence e- - (lone pair e- + 1/2 bonding e-)
For an oxidation number, bonding electrons are trasferred completely to teh more electronegative atom (as if the bonding were pure ionic):
Oxidation number = valence e- - (lone pair e- + bonding e-) |
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Term
| Molecules with Electron-Deficient Atoms |
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Definition
Gaseous molecules containing either beryllium or boron as the central atom are often electron deficient: they have fewer than eight electrons around the central atom.
There are only four electrons around Be and six around B. Sorrounding halogen atoms don’t form multiple bonds to the central atoms to give them an octet, because the halogens are much more electronegative. Formal charges make the following structures unlikely:
Electron-defiecient atoms of ten attain an octect by forming additional bonds in reactions. When BF3 reacts with ammonia, for instance, a compound forms in which boron attains an octet:
pg. 370 |
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Term
| Molecules with Odd-electrons Atoms |
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Definition
A few molecules contain a central atom with an odd number of valence electrons, so they cannot have all their electrons in pairs. Most have a central atom from an odd-numbered group, such as N (Group 5A) or Cl (Group 7A). These are called free radicals, species taht contain a lone electron, which makes them paramagnetic and extremely reactive. Free radicals are dangerous because they can bond to an H atom in a biomolecule and extract it, which forms a new free radical. This step repeats and can disrupt genes and membranes.
Consider the free radical nitrogen dioxide, NO2, a major contributer to urban smog, that is formed when the NO in auto exhaust is oxidized. NO2 has several resonance forms. Two differ in terms of which O atom is doubly bonded, as in the case of ozone. Two others have the lone electron residing on the N or on an O, so the resonance hybrid has the lone electron delocalized over these two atoms.
Free radicals often react with each other to pair their lone electons. When two NO2 molecules react, the lone electrons pair up to form the N-H bond in dinitrogen tetraoxide (N2O4) and each attain an octet:
formal charge is not very useful for picking the most important resonance form |
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Term
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Definition
| Many molecules (and ions) have more tahn eight valence electrons around the central atom. An atom expands its valence shell to form more bonds, which releases energy. The central atom must be large and have empty orbitals that can hold the additional pairs. Therefore, expanded valence shells occur only with nonmetals form Period 3 or higher because thy have d orbitals available. Such a central atom may be bonded to more than four atoms, or to four or fewer. |
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Term
| Central atom bonded to more than four atoms |
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Definition
| Phosphorus pentachloride, PCl5, is a fuming yellow-white solide used to manufacture lacquers and films. It forms when phosphorus trichloride, PCL3, reacts with chorine gas. The P in PCL5 forms, one Cl-Cl bond bond breaks and two P-Cl bond form |
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Term
| Valence-shell electron-pair repulsion (VSEPR) |
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Definition
| to minimize repulsions, each group of valence electrons around a central atom is located as far as possible from the others. A “group” of electrons is any number that occupies a licalized region around a atom: single bond, double bond, triple bond, lone electron. Th molecular shape is the three-dimensional arrangement of nuclei joined by the bonding groups. |
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Term
| Valence-shell electron-pair repulsion (VSEPR) |
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Definition
| to minimize repulsions, each group of valence electrons around a central atom is located as far as possible from the others. A “group” of electrons is any number that occupies a licalized region around a atom: single bond, double bond, triple bond, lone electron. Th molecular shape is the three-dimensional arrangement of nuclei joined by the bonding groups. |
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