Study Guide for Midterm Exam for Chemistry A
I. The Modern Period Table, pp.154-158
 know how the elements are organized in the periodic table—elements are arranged in
order of increasing atomic number or number of protons
 know elements are classified as metals, nonmetals, and semimetals and where they are
found in the periodic table
 know the periodic trends: atomic radius, ionic radius, ionization energy, and
electronegativity, pp.163-169
II. The Structure of the Atom, Chapter 4
 the atom consists of a nucleus, containing positively charged particles (Protons) and
particles that are neutral in charge (neutrons)
 most of the mass of an atom is in the nucleus of the atom
 the number of protons = the atomic number
 the number of neutrons = atomic mass - atomic number
 orbiting the nucleus are negatively charged particles (electrons); the number of electrons
is equal to the number of protons in an atom
 electrons can occupy different energy levels (principal energy levels, 1-7) and suborbitals
(s, p, d, f); an atom’s electron configuration can be written to show how electrons occupy
principal energy levels and suborbitals (know how to write the electron configuration
for an atom of any element and know the maximum number of electrons that can
occupy each principal energy level and its suborbitals, p.135-140)
 despite having a nucleus and all the different particles, an atom mostly consists of empty
space
III. Chemical Bonding, Chapter 8 and 9
 know the 3 main types of bonding that occurs between atoms
o covalent bonds—valence electrons are shared between atoms; covalent bonds
generally form between atoms with similar electronegativities, such as carbon
and hydrogen;
 Lewis Dot or Electron Dot diagrams can help illustrate how valence
electrons are shared between atoms in covalent bonds
o ionic bonds—electrostatic forces that hold oppositely charged particles (atoms)
together; ionic bonds generally form between atoms with dissimilar
electronegativities
 atoms held together in ionic bonds have opposite charges; these atoms
are called ions (an atom or group of atoms that have a positive or
negative charge); instead of sharing valence electrons, as in covalent
bonding, an atom like sodium transfers its 1 valence electron to an atom
like chlorine (with greater electronegativity), thereby giving sodium a
charge of 1+ and chlorine a charge of 1-

o
ionic compounds form a crystal lattice structure; ionic bonds formed by
elements/atoms with greater electronegativity form compounds that
have greater melting and boiling points than ionic compounds formed
between atoms with lesser electronegativity, p.215-220
metallic bonds—formed by the attraction of particles (atoms) with unlike
charges; “a metallic bond is the attraction of metallic cations (metal atoms with a
positive charge) for delocalized electrons”, p.228-229
 while metals often form lattice structures in the solid state like ionic
compounds, metallic bonds differ in that they do not form because of the
total transfer or giving up of valence electrons by some atoms entirely to
other atoms, as in the case with ionic bonds
 in metallic bonds, the outer energy levels of the metal atoms overlap; the
electrons in these outer energy levels (valence electrons) are “not held by
any specific atom and can easily move from one atom to the next”; these
valence electrons, because they are free to move, are often referred to as
“delocalized electrons”
 the “electron sea model” proposes that “all metal atoms in a metallic solid
contribute their valence electrons to form a “sea” of electrons”
 as valence electrons move freely throughout the solid, metal cations
(metal atoms with a + charge) are formed (hint: think about what the
charge of an atom would be if its valence electrons were free to move to a
neighboring atom, even for a short time); each such cation is bonded
together with neighboring cations by the “sea of electrons”
IV. The Mole, Chapter 11
 The mole is an unit of measure (SI base unit) used to measure the amount of a chemical
substance. Since chemical substances are made of small particles (atoms, molecules,
formula units (f.u.) in compounds, electrons, or ions), there are simply too many
particles to quantify (to give a value or quantity of) using conventional units of measure,
like a dozen or gross (12 dozen).
 The mole is defined as the number of representative particles or carbon atoms in
exactly 12 grams (g) of pure carbon-12 (no isotopes)
 Through experimentation, it has been determined that a mole of any substance contains
6.0221367 x 1023 representative particles; for calculations in class, we will use 6.02 x 1023
 The number, 6.0221367 x 1023, is called Avogadro’s number; it is correct to say, “a mole
of any substance contains Avogadro’s number of particles”
o know how to determine the number of representative particles (atoms,
molecules, f.u., electrons, or ions) when given the number of moles of a
substance
o know how to determine the number of moles of a substance when the number of
representative particles are given
o know how to calculate the molar mass of a substance
o
know how to use the molar mass to determine the mass of a substance when
given the number of moles of the substance, and vice versa
V. Chemical Reactions and Balancing Chemical Equations, Chapter 10
 Chemical reactions occur when the atoms of one or more substances (reactants) are
rearranged to form different substances (products); a chemical reaction is another name
for chemical change
 Chemical reactions involved the transfer of energy; know the difference between an
endothermic reaction (energy absorbed or transferred from surroundings) and an
exothermic reaction (energy released to surroundings)
 Chemical reactions follow important laws:
o The Law of Conservation of Mass—mass is neither created nor destroyed during
a chemical reaction, but is conserved
o The Law of Conservation of Energy—during chemical or physical processes,
energy may be changed from one form to another, but it is neither created nor
destroyed
 Chemical reactions can be represented through statements called chemical equations
o chemical equations show the starting substance(s) (reactants) and the different
substance(s) (products) formed during the chemical reaction
o a general example of a chemical equation, involving 2 reactants that yields 2
products following the reaction, looks like:
reactant 1 + reactant 2  product 1 + product 2
o a specific example is:
CH4(g) + 2O2(g)  CO2(g) + 2H2O(g)
o the equation above is written as a “balanced equation”; a balanced equation
shows that the number of each type of atom is equal on both sides of the
equation after a chemical reaction; this is important to show since mass
cannot be created or destroyed during a chemical reaction according to the
Law of Conservation of Mass
o NOTE: Although we have not covered “balancing equations” in class yet, it is
not a difficult concept. Read pp.278-283, “Representing Chemical Equations”
and “Balancing Chemical Equations” to help you learn how to write chemical
equations and balance them.