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Chemistry Finals
Chemistry—study of the composition, structure, and properties of matter and the
changes it undergoes
Element—pure substance made of only one kind of atom
Compound—substance that is made from the atoms of two or more
elements that are chemically bonded
Chemical Change or Reaction—one or more substances are converted into different
substances
Physical Property—characteristic that can be observed or measured
without changing the identity of the substance
Chemical Property—substance’s ability to undergo changes that
transform it into different substances
Intensive Property—does not depend on the amount of matter present
Extensive Property—depend on amount of matter present
Physical
Color
Melting Point
Boiling Point
Chemical
Reacts With…
Intensive
Electricity Conductor
Melting/Boiling Pt.
Density
Extensive
Volume
Mass
Conversion Factor—ratio derived from the equality between two different units that can
be used to convert from one unit to the other
Scientific Method—logical approach to solving problems by observing and collecting
data, formulating hypotheses, testing hypotheses, and formulating theories that are
supported by data
Hypothesis—testable statement, often stated as “if-then”
Significant Figures—consist of all digits known with certainty plus one final digit,
which is somewhat uncertain or estimated
Volume—amount of space occupied by an object
Theory—broad generalization that explains a body of facts or phenomena
Model—(in science) more than a physical object; it is often an explanation of how
phenomena occur and how data or events are related
Mass—measure of quantity of matter (measured in Kilograms (Kg))
Scientific Notation—numbers are written in the form M * 10N, where the factor M is a
number greater than or equal to 1 but less than 10 and N is a whole number
Weight—measure of gravitational pull on matter (measured in Newtons (N))
TABLE 2-5 Rules for Determining Significant Figures
Rule
Examples
1. Zeroes appearing between nonzero digits a. 40.7 L has three significant figures
are significant
b. 87 009 km has five sig figs
2. Zeros appearing in front of nonzero
a. 0.095 897 m has five significant figures
digits are not significant
b. 0.0009 kg has one sig fig
3. Zeros at the end of a number and to the
a. 85.00 g has four sig figs
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right of a decimal are significant
4. Zeros at the end of a number but to the
left of a decimal may or may not be
significant. If such a zero has been
measured or is the first estimated
digit, it is significant. On the other
hand, if the zero has not been
measured or estimated but is just a
placeholder, it is not significant. A
decimal placed after the zeros
indicates that they are significant.
b. 9.000 000 000 mm has 10 sig figs
a. 2000 m may contain from one to four
sig figs, depending on how many
zeros are placeholders
b. 2000. m contains four sig figs,
indicated by the presence of the
decimal point
EXAMPLE TABLE Scientific Notation
6.022 * 1023
2.56 * 103
3.14 * 101
(Avagadro’s #)
7.77 * 1077
1.0 * 10100
(One Googol)
Cathode Ray Tube—glass tube (See pg. 70). Experiment using a cathode ray tube, in
1897, by English physicist Joseph John Thomson, concluded that all cathode rays
(known now as electrons) are composed of identical negatively charged particles
Mole—scientific counting number (like a dozen) that uses Avogadro’s Number.
Looking at the periodic table, elements have their atomic weights listed below
them in Atomic Mass Units (AMU). One mole of that element takes that number
from AMUs into grams. I.e., one mole of carbon-12 (carbon atoms that all have a
mass of 12 AMU) would have a mass of 12 grams
Avogadro’s Number—6.022 * 1023
Atom—smallest particle of an element that retains the chemical properties of that
element
John Dalton—(1808) English schoolteacher who proposed an explanation for the law of
conservation of mass, the law of definite proportions, and the law of multiple
proportions. Here are his rules:
1.
2.
3.
4.
5.
All matter is composed of extremely small particles called atoms.
Atoms of a given element are identical in size, mass, and other properties; atoms
of different elements differ in size, mass, and other properties.
Atoms cannot be subdivided, created, or destroyed.
Atoms of different elements combine in simple whole-number ratios to form
chemical compounds.
In chemical reactions, atoms are combined, separated, or rearranged.
John Dalton—(cont.) turned Democritus’ idea into a scientific theory that could be
tested by experiment. Not all parts of Dalton’s Atomic Theory have proven
correct—one example is that atoms can be broken into quarks
Bohr’s Model—shows how an electron from a hydrogen atom absorbs energy, goes up
one energy level, emits the energy, and returns to the previous (ground) level.
See more on pg. 96
Nuclide—general term for any isotope of any element
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Isotopes—atoms of the same element that have different masses
Parts of the Atom—nucleus containing protons and neutrons and electron cloud
containing (GUESS!) electrons
TABLE 3-1 Properties of Subatomic Particles
Particle
Symbol
Relative Charge
Electron
e-1
+
Proton
p
+1
Neutron
n0
0
Mass Number
0
1
1
Molar Mass—mass of one mole of a pure substance
Mass Number—total number of protons and neutrons in the nucleus of an
isotope
Nuclear Forces—short-range proton-neutron, proton-proton, and neutron-neutron forces
that hold nuclear particles together
Law of Conservation of Mass—mass is neither created nor destroyed during ordinary
chemical or physical reactions
Robert A. Millikan—continued from Thomson’s cathode ray tube experiment. Millikan
discovered that the mass of the electron is in fact about one two-thousandth the
mass of the simplest type of hydrogen atom and that electrons carry a negative
charge. Thus, (1) because atoms are electrically neutral, they must contain a
positive charge to balance the negative electrons, and (2) because electrons have
so much less mass than atoms, atoms must contain other particles that account for
most of their mass
The electrons of the outer shell (valence) are separated from the kernel
(non-valence) electrons by AN ENERGY LEVEL. (Does it make sense now?)
Electromagnetic Radiation—form of energy that exhibits wavelike behavior as it
travels through space. Also travels through matter as a particle (p. 93—Einstein’s
dual wave-particle theory)
Wave Speed— c =  v. C is the speed of light (or any electromagnetic radiation), 3.0 *
108.  (lambda) is the wavelength—distance between corresponding points on
adjacent waves (pg. 91). V is the frequency—the number of waves that pass a
given point in a specific time, usually one second, expressed in waves per second
(a hertz (Hz))
The most stable state of an atom can be one of two things—either the
most stable isotope of an element, which would happen to be the most common
isotope, or (back to Bohr’s Model) when the electrons of an atom are at ground
level (by not absorbing additional photon energy)
Electron Cloud—space surrounding nucleus, containing the electrons of an atom
Orbital—3-d region around the nucleus that indicates the probable location of an
electron
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Quantum—minimum quantity of energy that can be lost or gained by an atom,
formulated as E = h v. E is the energy, in joules, of a quantum of radiation; v is
the frequency of the radiation emitted; h is Planck’s constant, 6.626 * 10-34 J s
(joules per second)
Orbital Notation—uses arrows to represent electrons in an atom
Electron Configuration Notation—uses superscripts, like 1s1 for
hydrogen, to represent electrons. If you don’t know what these notations are, you
shouldn’t have slept in class (or at least not as much)
Group Names—groups are vertical (Families are horizontal)
Alkali Metals—all of group 1 except hydrogen
Alkaline-Earth Metals—group 2
Halogens—group 17
Noble Gases—group 18
S-Block—groups 1 and 2
P-Block—groups 13-18
D-Block—groups 3-12 (transitional elements)
F-Block—lanthanides and actinides
Ion—atom or group of bonded atoms that has a positive or negative charge
Periodic Table Arrangement—everything on the table reoccurs periodically; elements
with similar properties fall in the same column or group (pg. 123-125)
Atomic Radii of Ions—anions, missing electrons, have a larger atomic radius due to less
atomic charges pulling the atom together. Cations, if you haven’t guessed, have a
smaller atomic radius due to increased atomic charges pulling the atom together
Atomic Radii of Neutral Elements of the Table—going down a group, atoms have a
larger atomic radius due to more electron levels. Going across a family, atoms
(usually) have a smaller atomic radius due to more electrons within the same
electron levels and sublevels. The rare exception occurs within new sublevels
created in an already used level, such as a new electron in the P sublevel after its
last F and D sublevels were filled
Electronegativity—measure of the ability of an atom in a chemical compound to attract
electrons. Elements are rated on a scale of 0 to 4, with the perfect 4 going to
Fluorine. Electronegativity determines the type of bond that occurs between
elements—the difference between two elements causes the following bonds:
Ionic—greater than 1.7
Polar Covalent—from 0.3 to 1.7
Non-Polar Covalent—less than 0.3
Electronegativity Table—pg. 151
Ionic Bonding—chemical bonding that results from the electrical attraction between
large numbers of cations and anions
Covalent Bonding—results from the sharing of electron pairs between two atoms
Non-Polar Covalent Bond—covalent bond in which the bonding
electrons are shared equally by the bonded atoms, resulting in a balanced
distribution of electrical charge
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Polar-Covalent Bond—covalent bond in which the bonded atoms have
an unequal attraction for the shared electrons
Ionic Compound—composed of positive and negative ions that are combined so that the
numbers of positive and negative charges are equal
Molecular Compound—chemical compound whose simplest units are molecules—a
neutral group of atoms that are held together by covalent bonds
Formula Unit—simplest collection of atoms from which an ionic
compound’s formula can be established
Chemical Formula—indicates the relative numbers of atoms of each kind
in a chemical compound by using atomic symbols and numerical subscripts
Molecular Formula—shows the types and numbers of atoms combined
in a single molecule of a molecular compound. The chemical formula of a
molecular compound is referred to as this
Common Compounds will be on the final! Table 7-2 will be provided
How many moles are in C6H12O6?—easy as pie or any other sugar-object—6 mol
carbon, 12 mol hydrogen, and 6 mol oxygen
Balance this—___HNO3 + ___Mg(OH)2  ___Mg(NO3)2 + ___H2O
Answer—2HNO3 + Mg(OH)2  Mg(NO3)2 + 2H2O
More on pg. 252
MoleGram Conversion—multiply an element’s number of moles by its atomic mass
GramMole Conversion—divide an element’s number of moles by its atomic mass
Formula Weight—the mass, in AMU, of a formula. Multiply each element’s mass in
the formula by the number of atoms each element has, then take the sum
Molar Mass—the mass, in grams, of one mole of a formula
Percent Composition—for each element, divide its mass over the mass of the entire
formula. Multiply by 100 and label with “% (element)”
Empirical Formula—the symbols for the elements combined in a compound with
subscripts showing the smallest whole-number mole ration of the different atoms
in the compound
Gift of Density—considered a gift due to how easy it is—density = mass / volume
Metric Meter Know-How
1000 meters make up one kilometer
100 centimeters make up one meter
1000 millimeters make up one meter
10 millimeters make up one centimeter
How many millimeters are in a kilometer?
1 km
x
1000 m
1 km
x
100 mm
1m
= 10,000 mm
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