Quick Review of 1st Semester Topics
Unit 1: Measurement, Matter and Change
Quantity
Being
Measured
Units of Measurement
Unit Symbols
mass
kilogram, gram, centigram, milligram
kg, g, cg, mg
volume
kiloliter, liter, milliliter, cubic meter,
cubic decimeter, cubic centimeter, cubic
millimeter
kL, L, mL, m3,
dm3, cm3, mm3,
length
kilometer, meter, centimeter, millimeter
km, m, cm, mm
energy
kilojoules, joules;
chemist unit for heat: calorie
kJ, J;
cal
pressure
millimeters of mercury, Pascals,
kilopascals, atmospheres
mmHg, Pa,
kPa, atm
density
temperature
mass
volume
degree Celsius;
g__;
mL
Kelvin
o
C;
kg__
cm3
K
1 km = 1 000 m
1 m = 100 cm
1 cm = 10 mm
1 m = 1 000 mm
1 kg = 1 000 g
1 g = 100 cg
1 cg = 10 mg
1 g = 1 000 mg
1 kL = 1 000 L
1 L = 1 000 mL
1 dm3 = 1 L
1 mL = 1 cm3
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matter:
has mass, occupies space, has inertia
pure substance:
same composition throughout sample
compound
element
composed of two or
more elements
chemically bonded
together; can be
decomposed only by
chemical means
composed of only
one type of atom;
mixture:
two or more substances physically
mixed together;
composition varies throughout
sample;
can be separated by physical means
can be changed only
by nuclear reactions
ionic compound
molecular compound
composed of two
or more charged
particles called
ions held together
by ionic bonds
composed of two or
more atoms of
different elements held
together by covalent
bonds
Physical means of separating a mixture include: filtration, evaporation, using known freezing
points and boiling points to separate different liquids, distillation (boiling off the liquid to leave
the solid component, and then condensing the vapor back to the liquid state).
Physical states of matter:
solid: particles packed very tightly together, particles are “fixed” in position relative to each
other
liquid: particles still very close together but particles can move around each other
gas: particles very far apart from each other
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Physical changes involve changes in physical state (solid ↔ liquid ↔ gas: melting, boiling,
condensation, freezing), cutting or crushing a large sample into smaller pieces, dissolving in an
appropriate solvent.
Chemical changes involve the rearrangement of the atoms of one or more substances to form
one or more new substances.
Physical properties of matter include: density, color, physical state at a given temperature,
boiling point, freezing point, odor, malleability, brittleness, hardness, solubility in a given
solvent, crystal shape.
Chemical properties involve how the substance behaves in the presence of another substance:
does the substance give up electrons easily to another substance? does the substance take
electrons from another substance? when the substances are mixed together does a new substance
form? does a new substance form when the two substances are heated together? does the
substance combine readily with oxygen gas to form new substances?
Unit II. Phase Changes and Thermochemistry
Thermal Changes: changes in heat energy
endothermic change: substance absorbs heat energy; feels “cool or cold” because heat
radiating from your body is being absorbed by the substance
exothermic change: substance gives off heat energy; feels “warm or hot” because there
is more heat energy outside your body
Calculations involving heat changes:
a) when there are changes in temperature of a substance: Q = m x ∆T x c; where
Q is heat energy in joules or calories being released or absorbed
m is mass of substance in grams
∆T is the change in Celsius temperature; ∆T = (T2 − T1)
c is the specific heat of the substance; unit is
J__ or
cal___
(g)(oC)
(g)(oC)
b) when there are changes in physical state: Q = H x m; where
Q is the amount of heat absorbed or released
H is the: heat of fusion, Hf (heat absorbed to melt 1.00 g of substance at the melting pt)
heat of solidification, Hf (heat released by 1.00 g of substance during freezing)
heat of vaporization, Hv (heat absorbed to vaporize 1.00 g of substance at boiling pt)
heat of condensation, Hc (heat released by 1.00 g of substance during condensation)
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Unit III.
Atomic Structure and Nuclear Chemistry
Atomic Structure
Atom: smallest particle of an element; composed of protons, neutrons, and electrons.
Atomic Particle
Location
Electrical Charge
Mass (atomic mass units)
proton
nucleus of atom
+1
1 a.m.u.
neutron
nucleus of atom
0
1 a.m.u.
electron
outside of nucleus
─1
0 a.m.u.
(mass is too small to be
significant)
In a neutral atom, the number of protons equals the number of electrons.
Number of protons + number of neutrons = mass number of the atom (atomic mass in a.m.u.)
The atomic number of an element is determined by the number of protons in every atom of that
element.
If “X” represents the symbol of the element, the mass number is written as a superscript at the
upper left of the symbol, and the atomic number is written as a subscript at the lower left of the
symbol:
mass number
X
atomic number
Isotopes: atoms of the same element that have the same atomic number (number of protons) but
have different numbers of neutrons and therefore different mass numbers.
and 157N are isotopes of nitrogen. Both have atomic number “7” but because the different
numbers of neutrons, the atoms have different mass numbers, 14 and 15 respectively.
Another way of writing isotopes is to use the element symbol followed by the mass number:
N-14, and N-15.
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7N
Nuclear Chemistry
Nuclear chemistry involves changes that occur within the nucleus of the atom. Elements that
have unstable nuclei due to the ratio of protons to neutrons spontaneously give off radiation from
the nucleus. The major forms of radiation are alpha particles (nuclei of helium atoms), beta
particles (electrons formed within the nucleus – these are NOT the “usual” electrons found
outside the nucleus of an atom), gamma rays (very high energy particles). Other types of
radiation particles include neutrons and positrons.
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There are four types of nuclear reactions:
1) natural radioactive decay: transmutations occur by natural radioactivity (the ability of a
nucleus to emit a nuclear particle and energy without external stimulation);
2) artificial transmutation: transmutations occur by nuclear disintegration caused by external
stimulation as a scientist bombards a nucleus with a particle (the addition of another nuclear
particle makes the nucleus unstable);
3) fission: certain nuclei having a large mass are bombarded with special particles that cause
the nuclei to split into two nuclei each having a smaller mass;
4) fusion: nuclei of light elements are combined to form heavier nuclei.
Nuclear reactions are written using symbols in the maX notation. The symbols for the major
nuclear particles involved in nuclear reactions are given below.
Alpha particle
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Proton
1
1H
Neutron
1
0n
Electron
0
-1e
Positron
0
1e
Gamma ray
0
0
2He
(an alpha particle is a helium nucleus)
(the most common hydrogen nucleus is a proton)
or
0
-1
(also called a beta particle)
0
1
or
(most harmful to living things)
The table below contains information about nuclear reactions involving the emission of
radiation.
Type of Reaction
Alpha emission ()
Beta emission ()
Positron emission (+)
Gamma emission
Radiation
(Particle)
4
0
-1e
2He
or
0
or
1e
0
0
-1
0
1
0
NUCLEAR CHANGES
Effect on the
Effect on the
Atomic Number
Atomic Mass
decrease by 2
decrease by 4
increase by 1
no change
decrease by 1
no change
no change
no change
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In nuclear equations, the total number of positive charges (represented by the atomic numbers)
of the reactants (substances on the left of the reaction arrow) equals the total number of
positive charges of the products (substances on the right of the reaction arrow). The total
mass of the reactants must also equal the total mass of the products.
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4Be
+
4
14
7N
+
4
2
27
13Al
+
45
21Sc
+
238
92U

2
He 
He 

1
1H
1
0n

234
90 Th
12
6C
17
1
+
+
1
26
13Al
+
8O
42
19K
4
2
+
0n
1H
+
2
4
1H
2He
He
Unit IV: Electrons in Atoms and the Periodic Table
Electrons in Atoms: Electrons in atoms are found outside the nucleus of the atom in regions
called principal energy levels. Each electron must have a specific KE to remain in a particular
energy level. The lowest principal energy level is “1”; the 1st principal energy level is closest to
the nucleus. Additional energy levels are farther from the nucleus.
Each principal energy level has certain available sublevels. Each particular sublevel also has a
specific energy requirement. Within each sublevel are orbitals. Orbitals can be thought of as
“pathways” along which electrons travel. Only two electrons can occupy a single orbital.
Table of the First 4 Principal Energy Levels: Sublevels, and Number of Orbitals in
Each Sublevel
Principal
Energy Level
Number of
Orbitals
Maximum Number of
Electrons per Sublevel
1
s
1
2
2
s
p
1
3
2
6
3
s
p
d
1
3
5
2
6
10
s
p
d
f
1
3
5
7
2
6
10
14
4
6
Available
Sublevels
There are currently 7 principal energy levels that have been identified. However, only four (4)
sublevels have been identified as containing electrons: s, p, d, and f.
An electron enters the lowest energy principal energy level available. Within that principal
energy level, the electron will enter the lowest energy sublevel available. Within that sublevel,
the electron will enter the lowest energy orbital available. Once every available orbital in a
given sublevel contains one (1) electron, a second electron may then enter an orbital to form a
pair of electrons in the orbital.
The following Aufbau diagram illustrates the order of filling energy levels, and sublevels.
7s
7p
7d
7f
6s
6p
6d
6f
5s
5p
5d
5f
4s
4p
4d
4f
3s
3p
3d
2s
2p
1s
It is the number and arrangement of the outermost electrons (electrons in the outermost
principal energy level) that determine the chemical properties of an element.
Electrons may “jump” from one sublevel or principal energy level to a higher one ONLY IF the
electron absorbs ENOUGH energy to make the jump AND that the KE is equal to the energy
requirement of the higher level. When an electron absorbs that specific amount of energy and
jumps to a higher level, the electron is said to be “excited.”
When the excited electron returns to its normal energy level, energy is released by the electron as
it descends in “steps;” the descent is not made all at once. Each “packet” of energy that is
released is called a “photon” and has a specific wavelength and frequency.
It is the frequency of the photon that determines the energy of the photon
Electrons may be excited by heat, light, electricity, or any other form of energy. The photons
released by excited electrons returning to their normal energy levels accounts for the colors we
see in flame tests, fireworks, any fire such as that in a fireplace or a lit match, and in the colors of
our clothes (the electrons of the atoms in dye molecules are excited by light energy).
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Periodic Table
Mendeleev proposed an arrangement of the known elements on the basis of increasing atomic
mass and the periodicity of the properties of the elements. Mosley improved the arrangement of
the elements by using the increasing atomic number. By using atomic number instead of atomic
mass, several elements were placed in their correct “families” or groups having similar properties
instead of just switching them to be in the same family as Mendeleev did.
The vertical columns are called GROUPS or “families” and are numbered left to right as 1-18.
The horizontal rows are called PERIODS and are numbered top to bottom as 1-7 on the current
periodic table.
Elements in the same group or family have similar chemical and physical properties.
Elements in the same period will have the same number of outermost principal energy levels
occupied by one or more electrons.
Group 1 is the ALKALI METAL group. The atoms of these metals all have an s1 valence
electron.
Group 2 is the ALKALINE EARTH METAL group. The atoms of these metals all have s2
valence electrons.
Groups 3-12 are the TRANSITION METAL groups. These metals often form compounds that
are colored (red, green, yellow, etc.) The number of valence electrons of these metals varies.
Group 13 is known as the BORON group and each of the atoms of these elements have s2 p1
valence electrons.
Group 14 is called the CARBON group; each atom of these elements has 4 valence electrons.
Group 15 is the NITROGEN group; these elements have s2p3 as the valence electron
configuration.
Group 16 is the OXYGEN group; the valence electron configuration for this group is s2p4.
Group 17 is known as the HALOGEN group; the valence electron configuration is s2p5.
Group 18 is the NOBLE GAS group. These elements have s2p6 as the valence electron
configuration.
Metal atoms tend to give up all of their valence electrons; nonmetals tend to gain enough
valence electrons to have an s2p6 configuration – the same as the noble gases. Nonmetals will
also share valence electrons in order to achieve the s2p6 configuration.
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Trends and the Periodic Table
Electronegativity (attraction for electrons in a covalent bond) increases left to right across a
period; decreases down a group.
Metallic property (tendency to give up valence electrons): decreases left to right across a
period; increases down a group.
Atomic radius: decreases left to right across a period; increases down a group.
Ionic radius: decreases left to right for metals and increases left to right for nonmetals across a
period; increases for both metals and nonmetals down a group.
Ionization energy (energy required to remove one valence electron): increases left to right
across a period and decreases down a group.
Unit V. Chemical Bonding
Ionic Bond: a bond involving the mutual attraction of oppositely charged ions.
An ION is an atom or molecule that has gained or lost one or more electrons.
An ionic bond is an ATTRACTION between ions – there is nothing physical holding the ions
together.
METALS: tend to lose their valence electrons and form positively charged ions.
NONMETALS: tend to gain valence electrons and form negatively charged ions.
Covalent Bond: a bond formed by the mutual sharing of one or more pairs of valence
electrons between two atoms; atoms that are covalently bonded together form molecules.
single covalent bond: one pair of valence electrons is shared
double covalent bond: two pairs of valence electrons are shared
triple covalent bond: three pairs of valence electrons are shared
Coordinate Covalent Bond: a single covalent bond between two atoms BUT the pair of
electrons that is shared is donated by only one of the two atoms.
Polar Covalent Bond: a covalent bond in which the shared electrons are UNEQUALLY shared
by the two atoms; the electrons are held more closely to the atom having the greater
electronegativity. In such a bond, the atom with the shared electrons closer to itself is slightly
negative with respect to the other atom. This sets up “poles” along the bond: a “more positive
pole” and a “more negative pole.”
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Lewis Dot Structures: uses “dots” to represent the valence electrons of each atom.
Structural Diagram: uses lines to represent pairs of shared electrons: one line represents 1 pair
of shared electrons ( a single covalent bond); two lines represents 2 pairs of shared electrons (a
double covalent bond); three lines to represent 3 pairs of shared electrons (a triple covalent
bond).
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