Term
| How do you calculate bond order? |
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Definition
| (# of bonding electrons - # of antibonding electrons)/2 |
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Term
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Definition
| species with unpaired electrons are weakly attracted to a magnetic field. |
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Term
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Definition
| species with paired electrons are repelled out of a magnetic field. |
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Term
| Why is the hybridization diagram different for B, N, and C? |
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Definition
electrons in the σ2s MO will repel those going into the π2p
and so make the, higher in energy
than we previously expected |
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Term
| Atomic orbitals of _____ electronegative elements are _______ in energy. |
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Definition
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Term
| Atomic orbitals of the highest electronegative element will dominate the character of the_______ molecular orbital |
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Definition
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Term
| Atomic orbitals of the lowest electronegative element will dominate the character of the_______ molecular orbital |
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Definition
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Term
| What do you do with nonbonding electrons? |
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Definition
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Term
| What are the properties of gases? |
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Definition
1. Expand to fill available volume
2. Compressable
3. Volume and pressure vary greatly with temperature |
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Term
| To directly measure pressure, you would use a ______. |
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Definition
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Term
| To indirectly measure pressure, you would use a ______. |
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Definition
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Term
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Definition
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Term
| What is the implication of Boyle's Law? |
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Definition
| Pressure is inversely related to volume |
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Term
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Definition
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Term
| What is the implication of Charles' Law? |
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Definition
| Volume is proportional to temperature in kelvin |
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Term
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Definition
Standard Temperature and Pressure.
T=0 degrees C or 273.15K
P= 1 atm or 101.3 kPa |
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Term
| What is the Combined Gas Law? |
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Definition
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Term
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Definition
V1/n1=V2/n2
At the same temperature and pressure, two equal volumes of all gases have the same number of molecules. So, the number of molecules is proportional to volume. |
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Term
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Definition
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Term
| What is standard molar volume? |
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Definition
The volume of 1 mole of an ideal gas at STP
Standard Vm=22.414L/mol |
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Term
| What is the Ideal Gas Equation? |
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Definition
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Term
| What is the universal gas constant? |
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Definition
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Term
| What is Dalton's Law of Partial Pressures? |
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Definition
| For a sample of several gases in a container, each gas exerts pressure as if none of the other gases were present. |
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Term
| What is partial pressure? |
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Definition
| The pressure exerted by a gas in a mixture of gases |
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Term
| How do I find the total pressure of a container of mixed gases? |
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Definition
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Term
| What is another way to write out Dalton's Law? |
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Definition
| By using a molar fraction, Xa=(# of moles of a/total # of moles), we can rewrite the equation to Pa=Xa*Ptotal |
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Term
| What is Kinetic Molecular Theory? |
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Definition
A way of modeling the behavior of ideal gases:
1.Gases are very small molecules that are spaced very far apart
2.These molecules are in constant, random, straight-line motion until they collide with one another or the or the vessel wall
3. Any collision is perfectly elastic, no energy is gained or lost
4. In between collisions, the molecules do not interact with one another and travel at constant speeds. (Individuals travel at different speeds) |
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Term
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Definition
| Energy possessed due to motion. KE=(1/2)mv^2 |
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Term
| What is the equation for Kinetic Energy of a mole of particles? |
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Definition
KEmolar=Na*(1/2)mV(rms)^2
OR
KE=(3RT)/2 |
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Term
| What does Kinetic Molecular Theory have to do with temperature? |
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Definition
| Absolute temperature is directly proportional to the mean KE of gas molecules |
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Term
| What expression would I use to find average velocity of the particles? |
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Definition
v(rms)=(3RT/2)^(1/2)
When M=molar mass in kg, R=8.314J/mol*K |
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Term
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Definition
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Term
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Definition
| The escape of gases through a tiny hole |
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Term
| What does the rate of both diffusion and effusion depend on? |
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Definition
| The average speed of the gas particles |
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Term
| At the same temperature, on average, _____ molar mass gas particles move faster. |
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Definition
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Term
| What is Graham's Law of Diffusion and Effusion of Gases? (1) |
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Definition
The ratio between the effusion or diffusion of two different gases is inversely proportional to the square root ratio of their molecular weights:
(R1/R2)=(M1/M2)^(1/2)
Note that R=Amount (of gas)/time (t) to effuse or diffuse |
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Term
| What is Graham's Law of Diffusion and Effusion of Gases? (2) |
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Definition
The ratio between the effusion or diffusion of the same gas at two different temperatures is directly proportional to the square root ratio of their temperatures:
(R1/R2)=(T1/T2)^(1/2)
Note that R=Amount (of gas)/time (t) to effuse or diffuse |
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Term
| What are the limitations of Kinetic Molecular Theory? |
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Definition
| Gases do in fact interact when they get close to one another (ex- jogger passing pretty lady slows down, as long as he's close enough to get a good look) |
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Term
| What is Van der Waal's equation? |
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Definition
(P+(n^2*a)/V^2)*(V-nb)=nRT
Where:
a=measurement of the strength of attractive forces between molecules (more polar molecules=larger a values)
b=related to the size of the gas molecules
P=measured pressure
V=volume of vessel |
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Term
| What is the first term of Van der Waal's equation? |
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Definition
| At low temp, KE is small: interactions are significant- slows particles so pressure is reduced on vessel walls and thus measured P will be lower than predicted by ideal gas equation |
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Term
| What is the second term of Van der Waal's equation? |
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Definition
| At high pressure, volume is small: volume of particles is significant- volume available to gas is lower than volume of vessel. |
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Term
| What happens in Van der Waal's equation at a low pressure and high temperature? |
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Definition
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Term
| Do polar or non polar gases deviate the most from ideal gas laws? Why? |
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Definition
| Polar, because they have larger a and b terms |
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Term
| Under what conditions do ideal gas laws work fairly well? |
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Definition
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Term
| What hybridization bonds make up a single bond? |
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Definition
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Term
| What hybridization bonds make up a double bond? |
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Definition
| one sigma and one pi bond |
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Term
| What hybridization bonds make up a triple bond? |
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Definition
| one sigma bond and two pi bonds |
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Term
| How are sigma bonds formed? |
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Definition
| By the overlapping of pair of atomic orbitals |
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Term
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Definition
| By two unpaired electrons forming a hybridization bond |
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Term
| What are homonuclear diatomics? |
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Definition
| Molecules created by the bonding of two of the same atoms |
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Term
| As bond order increases, it can be assumed that the bonds are_______, have ______ bond energy, and ______ bond length. |
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Definition
| stronger, higher, shorter |
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Term
| The MO diagram is skewed for _____nuclear diatmoics |
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Definition
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Term
| In a MO diagram of a heteronuclear diatomics, the atomic energies of the ______ electronegative element are lower than the corresponding ones of the other element. |
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Definition
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Term
| How do lop sided MO diagrams affect the MO's of heteronuclear diatomics? |
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Definition
the antibonding orbitals 'look' more like the atomic orbitals of the less electronegative element.
the bonding orbitals 'look' more like the atomic orbitals of the more electronegative element. |
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Term
| If given the heteronuclear diatomic, NO, which MO diagram would you use? |
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Definition
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Term
| What are 'next door' diatomics and how do we handle them? |
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Definition
| CN, OF, etc. are examples of these. We handle them by pretending they are homonuclear diatomics and draw the relevent diagrams and answer questions as above for homonuclear diatomics. |
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Term
| What should you remember about stoichiometric coefficients? |
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Definition
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Term
| What should you do with any chemical equation? |
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Definition
| Check that it's balanced! |
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Term
| What is the Maxwell-Boltzmann Distribution? |
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Definition
| Measures the distributions of speeds as absolute temperature increases. Shows that for two gases at the same temperature, the heavier of the two has a narrower hump. Overall, shows that average KE of gas molecules is directly proportional to absolute temperature. |
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