Term
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Definition
organize and understand underlying principles
of nature
- outgrowth of natural philosophy or the
philosophical speculation of nature
Science involves two facets:
- technological (or factual)
- philosophical (or theoretical) |
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Term
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Definition
- Testable
- Reproducible
- Explanatory
- Predictive
- Tentative
Limitations:
- studying that which is observable
- natural processes in which
variables can be controlled |
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Term
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Definition
Technology - direct application of knowledge to
solve problems |
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Term
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Definition
Chemistry is the study of matter and its
changes. |
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Term
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Definition
Science and technology are interrelated. Their use
involves both risks and benefits.
Does the benefit of use outweigh the risk?
“Risk-benefit analysis” involves an estimation called
the desirability quotient (DQ).
DQ = Benefit
Risks |
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Term
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Definition
- testable explanation of observed data
- tested by designing and performing experiments |
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Term
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Definition
set of tested hypotheses that explain natural phenomena
- best current explanation for (a group of) natural
phenomena
- always tentative; (parts) may change as observations
change |
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Term
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Definition
- description of a natural phenomenon
- generally one “action”
- true for stated conditions
- Law of Gravity, Law of Conservation of Mass/Matter
- can often be stated mathematically
- Boyle’s Law (PV = k) |
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Term
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Definition
- can be measured or described
- “5 senses” |
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Term
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Definition
- found when matter reacts or rearranges its
atoms |
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Term
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Definition
the identity/
composition of the
substance is not
changed. |
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Term
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Definition
the chemical identity
of the substance is
changed. |
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Term
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Definition
anything that has
mass and volume (takes up
space) |
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Term
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Definition
measure of the
amount of matter in an
object |
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Term
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Definition
measure of gravitational
force for the matter in an object |
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Term
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Definition
Substance - “pure matter” that has a fixed composition that
does not vary
Atoms
- the smallest particles of an element to retain properties
Elements
- substance composed of one type of atom
- represented by chemical symbols, e.g., Cl, H, Mg
Compounds
- composed of two or more elements chemically
combined
- fixed proportions
- frequently bond together as groups of atoms in units ⇒
molecules
Mixture
- >2 substances that may be mixed in variable amounts
- varying compositions = varying properties
- heterogenous - very different compositions (solid, liquid)
- homogenous - intimately similar compositions |
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Term
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Definition
the amount of matter
in a given amount of space
d = m/V |
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Term
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Definition
- the ability to do work or transfer heat
- usually accompanies physical and chemical
changes |
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Term
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Definition
flow of energy that is transferred from hotter
objects to cooler objects |
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Term
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Definition
the average kinetic energy of an
object, measure of “hotness” or “coldness” |
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Term
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Definition
K = oC + 273.15
Example: Human body temperature is 37
oC. Convert this to Kelvin.
K = oC + 273.15
= 37 + 273.15
= 310.15 K |
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Term
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Definition
Aristotle
- all matter is composed
of 4 elements, and all
matter is continuous, not
atomistic.
Leucippus and
Democritus
- “atomos”: there is a point at which matter
can no longer be subdivided, leaving
indivisible particles
- atoms - from ‘atomos’ - “cannot be cut”
- each type of atom distinct shape and size
- all substances combo of various
atoms
Robert Boyle - substances (compounds) could be
broken down into simpler substances called elements |
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Term
| Law of Conservation of Mass |
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Definition
Antoine Lavoisier “father of modern chemistry”
- decomposition of mercuric oxide, weighed products
and found they were same mass (he named oxygen)
- did other rxns to find same result
- during a chemical change, matter is neither created nor
destroyed |
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Term
| Law of Definite Proportions |
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Definition
Joseph Proust
A compound always contains the same elements
- in certain definite proportions
- in no other combinations
- pure compounds exhibit consistent properties |
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Term
Law of Multiple
Proportions |
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Definition
John Dalton
elements may combine in more than one set of
proportions
- each set corresponds to a different compound |
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Term
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Definition
1. All matter is composed of extremely small particles
called atoms
(Dalton assumed atoms are indivisible, though this is not the
case)
2. All atoms of a given element are alike and differ
from the atoms of any other element
(Dalton assumed all atom masses were identical, but this is
not always true)
3. Compounds are formed when atoms of different
elements combine in fixed proportions.
4. A chemical reaction involves the rearrangement of
atoms. No atoms are created, destroyed or broken
apart in a chemical reaction.
(except in nuclear reactions) |
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Term
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Definition
Dmitri Mendeleev
- elements arranged in order of increasing
atomic mass
- also by similar properties
- left gaps for undiscovered elements
- predicted the properties of unknown elements
- many of his predictions were found to be
accurate |
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Term
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Definition
Electrolyte - compound that conducts
electricity when molten or dissolved in
water
Electrodes - carbon rods of metallic strips
that carry electrical current
- anode: positive electrode
- cathode: negative electrode
Electrolysis - separation of an ionic substance
in solution using electricity
- ion: atom or group of atoms with a charge
- anion: negative ion
- cation: positive ion |
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Term
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Definition
Crookes Tube
- better vacuum, conducted electricity in air
(whereas Faraday’s failed)
- ZnS coated screen allowed “beam” to be seen
from cathode to anode (“cathode ray”)
Joseph Thomson:
- used magnets perpendicular to cathode rays
- rays moved toward “+” plate - must be “-” in nature
- all electrodes showed same behavior
- “electrons” shown to have mass-to-charge ratio by
varying magnetic field (same for all gases)
Goldstein
- tube filled with gas &
perforated cathode
- electrons moved to
anode
- (+) particles moved to
cathode
-mass “+” depended on
type of gas used
Hydrogen - “+” particle 1837x as
heavy as “-”
Robert Millikan: the oil drop experiment
- discovered the charge of an electron (1.6x10-19 C).
- thus, mass of electron 9.1x10-28 g (much less than
any atom)
Ernest Rutherford (Geiger & Marsden)
- fired alpha particles at thin metal foils
- most went through, small amount deflected
- hitting nucleus (or close)
- nucleus is small, but
- dense
- (+) charge |
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Term
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Definition
Protron mass=1
Electron mass=1/1837
Neutron mass=1 |
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Term
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Definition
alpha particle (2 protons + 2
neutrons, He2+)-mass 4
- beta particle (e-) - like cathode
rays mass=1/1837
- gamma - very energetic
electromagnetic radiation (like Xrays mass=0 |
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Term
| Subatomic particles-Nuclear symbols |
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Definition
Atomic Number - number of protons
in a nucleus (unique)
Mass Number - sum of protons and
neutrons in the nucleus
- 1914, Rutherford suggested “+”
particle be called proton, equal and
opposite charge to electron
- neutron found by James Chadwick
(1932); same size as proton, no
charge
Electrons - negatively charged
particles in cloud around nucleus |
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Term
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Definition
same atomic number (# protons) but
different mass numbers (different # of neutrons)
- every element occurs as mixture of isotopes
- natural abundance: % of atoms occurring as a given
isotope (p+ = 1.6726x10-27 kg, n0 = 1.6749x10-27 kg)
- isotopes may have slightly different physical properties
(related to mass) |
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Term
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Definition
- light from a gaseous substance
passed through a prism
- produces a line spectrum related
to specific energies of light.
Blue/violet light is higher energy than red light
(energy order ROY G BIV). |
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Term
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Definition
Niels Bohr - line spectra exist b/c electrons can have only specific
energies
- energy values called “energy levels”
Quantum - unit of energy (tiny); plural: quanta
- produced or absorbed when an electron transitions between energy levels
- the absorption (and
shortly thereafter)
emission of energy
results in flame
coloration (line spectrum)
- electron cannot make
“intermediate” jumps
Ground state - electrons
are in the lowest energy
state
Excited state - electrons
absorb energy (e.g.,
flame) and are promoted
to a higher energy state
When an excited state ereturns
to a lower energy
state
- emits a photon of energy
- may be observed as light
Ground state -
- (+) & (-) charges attract each other
- energy of e- prevents fatal attraction
into nucleus
If enough energy is supplied, e- dissociates from atom
- energy of free e- is higher than bound e-
- energy of free e- at rest is “0” reference point of energy level
diagram
B/c only certain energies are absorbed, e- levels are well-defined |
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Term
| Heisenberg’s Uncertainty Principle |
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Definition
the position of a specific,
moving e- cannot be precisely defined
- as e- position becomes more accurate, the momentum becomes more
uncertain (vice versa) |
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Term
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Definition
a probability-based model;
- Principle Energy Levels (shells, n): roughly correlate to the
distance of an electron from the nucleus; can hold > 1 orbital
- Sublevels (subshells): each principle energy level (n) is divided
into sublevels = s, p, d, f, etc.
- Orbitals: regions in space representing a high probability of
locating an electron
- each sublevel has one or more orbitals that hold 2e- each
1st shell - 1 orbital - 1s
2nd shell - 4 orbitals - 2s and 3x 2p (x, y, z)
3rd shell - 9 orbitals - 3s, 3x 3p and 5x 3d (xy, xz, yx, etc.)
For each element, an electron configuration can be written for the ground state
(most stable state). |
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Term
| Metals, Nonmetals, and Metalloids |
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Definition
Metals (right)
• metallic luster, conduct heat and electricity, malleable and
ductile. Ex: sodium, copper
Nonmetals (left)
• dull luster, nonconductors, brittle
Ex: sulfur, chlorine
Metalloids (stairstep)
• demonstrate properties of both metals and nonmetals (B, Al,
Si, Ge, As, Sb, Te, Po) |
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Term
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Definition
electrons in the outermost principle energy level of an
atom
- electrons that are gained, lost, or shared in a chemical
reaction
- elements in a group/family have the same number of
valence electrons |
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Term
| Groups in the periodic table have special names |
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Definition
Alkali Metals: Group 1A
– valence electron configuration: ns1 (1 valence e-)
• Alkaline Earth Metals: Group 2A
– valence electron configuration: ns2 (2 valence e-)
• Halogens: Group 7A
– valence electron configuration: ns2np5 (7 valence e-)
• Noble Gases: Group 8A
– valence electron configuration: ns2np6 (8 valence e-) |
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Term
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Definition
- the forces that hold atoms together in molecules
- sharing or transfer of e |
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Term
| Consequences of chemical bonding & structures |
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Definition
- physical state of compound at room temperature
- strength of materials
- “texture” of liquid (light/volatile vs. heavy/viscous)
- taste, odor
- drug activity
- toxicity |
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Term
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Definition
- developed visual representations of the
valence electrons
- dots around the symbol of an atom |
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Term
| Formulas and naming: Binary ionic compounds |
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Definition
Cations: charge is the same as the family (group)
number
- the name is the name of the element
Examples:
Na+ = sodium ion
Mg2+ = magnesium ion
Anions: charge is equal to the family number – 8
- the name is the element root name (1st syllable) plus
the suffix –ide
Examples:
Cl- = chloride ion
O2- = oxide ion
To name the compounds of simple binary ionic
compounds (binary: two elements)
…name the ions.
Examples:
NaCl = sodium chloride
MgO = magnesium oxide
To write formulas: determine ion charges and then
transpose charges for subscripts
Examples:
lithium hydride = Li+ & H- = LiH
calcium fluoride = Ca2+ & F- = CaF2 |
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Term
| Formulas and naming; Transition metals |
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Definition
Transition metals
- many can exhibit more than one ionic charge
- may lose s & d electrons
- Roman numerals denote the charge of the ions
Examples:
Fe2+ = iron (II) ion
Fe3+ = iron (III) ion
Cu2+ = copper (II) ion
Cu+ = copper (I) ion |
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