Download H3AsO4 + 3 I- + 2 H3O+ H3AsO3 + I3- + H2O

AP Chemistry
Fall Semester Review Worksheet
Name _________________________
Period _____
Unit 1-7 Summary
Measurement in Chemistry
Science knowledge is advanced by observing patterns (laws)
and constructing explanations (theories), which are supported by
repeatable experimental evidence.
Measurements are made using the metric system, where the
standard units are called SI units, which are based on the meter,
kilogram, and second as the basic units of length, mass, and time,
respectively. The SI temperature scale is the Kelvin scale, although
the Celsius scale is frequently used in chemistry. The metric
system employs a set of prefixes to indicate decimal fractions or
multiples of the base units; k (10-3), c (10-2), m (10-3),  (10-6) and n
(10-9).
All measured quantities are inexact to some extent. The
precision of a measurement indicates how closely different
measurements of a quantity agree with one another. The accuracy
of a measurement indicates how well a measurement agrees with
the accepted value. Significant figures indicate the level of certainty
in a measurement. Significant figures in a measured quantity
include one estimated digit; the last digit of the measurement.
Calculations involving measured quantities are reported with the
appropriate number of significant figures. In multiplication and
division, the number of significant figures is used. In addition and
subtraction, the position of the least accurate significant figure is
used. Relative difference between an experiment value (E) and a
true value (T) is % difference: %  = 100|E – T|/T.
Mass and volume measure amount of matter. Density relates
mass to volume, d = m/V.
Chemical processes involve interaction of particles, which are
measured in moles. The number of particles in a mole is 6.02 x 10 23
(Avogadro's number), which is the number of atoms in a sample of
the element that has a mass equal to its atomic mass measured in
grams.
Molar mass (MM) is the sum of atomic masses in the chemical
formula. For example, the mass of one H2O molecule is 18.0 u, so
the molar mass of H2O is 18.0 g.
Dimensional analysis is a calculation strategy, where units are
conserved. The given units are multiplied by a series of conversion
factors, which are ratios of equivalent quantities. After canceling out
units algebraically, what remain are the target units.
Atomic Nature of Matter
Atoms are the basic building blocks of matter; they are the
smallest units of an element that can combine with other elements.
Atoms are composed of even smaller subatomic particles.
Experiments led to the discovery and characterization of subatomic
particles.
Scientist
Discovery
Dalton
Atomic Theory (atoms are conserved in reactions)
Thomson Electron charge/mass ratio (cathode ray bending)
Millikan
Electron mass (charged oil drops in electric field)
Rutherford Planetary atomic model (alpha rays and gold foil)
The atom's nucleus contains protons and neutrons, whereas
electrons move in the space around the nucleus.
Particle
Charge
Mass
Symbol
1 p or 1 H
Proton
+1
1
1
1
1 n
Neutron
0
1
0
0 e
Electron
–1
-1
0
Elements are classified by their atomic number or Z value,
which equals the number of protons. The mass number or A value
is the sum of protons and neutrons. Atoms of the same element
that differ in mass number are called isotopes.
In a neutral atom, the number of protons equals the number of
electrons. Anions are formed when electrons are added to neutral
atoms. Cations are formed when electrons are removed from
neutral atoms.
The unified atomic mass scale (u) is 1/12 the mass of a C-12
atom. The average atomic mass of an element is calculated using
the formula: 100mav = %1m1 + %2m2 ...
The two kinds of pure substances are elements and
compounds. Elements are identified by a chemical symbol.
Compounds are composed of two or more elements joined
chemically and identified by a chemical formula, which shows the
composition. Molecular compounds have a defined size, whereas
crystalline compounds are unbounded, where their formula shows
the ratio of atoms in the compound.
Mixtures are composed of multiple pure substances in an
object or container and have variable compositions. They can be
homogeneous or heterogeneous. Homogeneous mixtures are also
called solutions and are uniform throughout.
Radioactivity
There are four kinds of radioactive decay: emission of alpha
particles ( or 42He), beta particle ( or 0-1e), positron particle (,
0 e), and gamma radiation (0 ).
1
0
In nuclear equations, reactant and product nuclei are
represented by AZX, which is its nuclear symbol. In a balanced
equation the sum of reactant A and Z values equal the sum of
product A and Z values.
Nuclear transmutations, induced conversions of one nucleus
into another, can be brought about by bombarding nuclei with either
charged particles or neutrons.
The decay rate (radioactivity) is proportional to the number of
radioactive atoms, rate = kNt. The time for half of the radioactive
atoms to decay is constant, t½ = (ln2)/k. The time interval t for No
number of radioactive atoms to reduce to Nt is determined by the
formula, kt = ln(No/Nt).
Electron Structure—Bohr Model
The electronic structure of an atom describes the energies and
arrangement of electrons around the atom. Much of what is known
about the electronic structure of atoms was obtained by observing
atomic spectra, which is the radiant energy emitted or absorbed by
matter.
Equations for radiant energy, Ephoton = hf and speed of light, c =
f are combined in Ephoton = hc/ = 2 x 10-25 J•m/.
Bohr analyzed the wavelengths of light emitted by hydrogen
atoms and proposed a model that explains its atomic spectrum. In
this model the energy of the hydrogen atom depends on the value
its quantum number n, where En = -2.18 x 1018 J/n2. The value of n
is a positive integer (1, 2, 3 . . .). As n increases, the energy of the
electron increases until it reaches a value of 0 J, where n equals
infinity and the electron is ionized (leaves the atom). The lowest
energy state where n = 1 is called the ground state. Other values of
n correspond to excited states. Light is emitted when the electron
drops to a lower energy state and light is absorbed when the
electron is excited to a higher one. The energy of light emitted or
absorbed equals the difference in energy between the two states,
Ephoton = En-final – En-initial = 2.00 x 10-25 J•m/.
Quantum Mechanical Model
In the quantum mechanical model each electron has a known
energy, but according to the Heisenberg Uncertainty Principle, the
location of the electron cannot be determined exactly; rather, the
90 % probability of it being at a particular point in space is given by
its orbital. An orbital is described by a combination of four quantum
numbers. The principal quantum number n is indicated by the
integers 1, 2 . . . This quantum number relates to the radius and
energy of the orbital. The sublevel quantum number l is indicated
by the letters s, p, d, and f, which correspond to l = 0, 1, 2, and 3
respectively. The l quantum number defines the shape of the
orbital. For a given value of n, l can have integer values from 0 to
(n – 1). The orbital quantum number ml relates to the orientation of
the orbital in space. For a given value of l, ml can have integral
values ranging from –l to +l. The spin quantum number ms defines
the orientation of the electron's magnetic field and has two
possible values +½ and –½. The Pauli Exclusion Principle states
that no two electrons in an atom can have the same spin in the
same orbital. This principle limits the number of electrons that
occupy any one atomic orbital to two.
Electron Arrangements in Atoms and Ions
Energy increases as n increases (1 < 2 < 3, etc) and within
the same value of n, energy increases as the sublevel progresses
from letters s  p  d  f. Orbitals within the same sublevel are
degenerate, meaning they have the same energy.
The energies of s and p sublevels are less than the energy of
the next higher s sublevel, whereas the energies of d and f
sublevels are greater than the next higher s sublevel. This restricts
the outermost occupied sublevels for any atom to s and p.
Electrons that occupy the outermost sublevels are involved in
chemical bonding and are called valence electrons. Non-valence
electrons are called core electrons.
The periodic table is partitioned into different types of
elements based on their electron arrangement. Elements with the
same valence energy level form a row or period. Elements with the
same number of valence electrons form a column or group. The
elements in which an s or p sublevel is being filled are called the
main-group elements, which include group 1—alkali metals, group
2—alkaline earth metals, group 17—halogens and group 18—
noble gases. Transition metals are where the d-sublevel is filling.
Electron configurations show how electrons are distributed
among the atom's sublevels. In ground state configurations
electrons occupy the lowest sublevel available until its capacity is
reached. Additional electrons fill the next lowest sublevel until its
filled, etc. Excited state configurations have gaps.
Orbital diagrams show how the electrons fill the specific
orbitals, where arrows are used to represent electrons; () for ms =
+½ and () for ms = –½. When electrons occupy a sublevel with
more than one degenerate orbital, Hund's Rule applies. The rule
states that the lowest energy is attained by maximizing the number
of electrons with the same electron spin.
Transition metals in columns 6 and 11 have a half-filled s
sublevel in order to have a Half-filled or fully occupied d sublevel,
which is more stable than other arrangements.
The electron arrangement for monatomic ions is the same as
the element with the same number of electrons. Elements within
three squares on the periodic table of a noble gas form ions with
the same electron arrangement as the noble gas and are
isoelectronic. Transition metals form ions by losing all s level
electrons first.
Elements with unpaired electrons have reinforcing magnetic
fields, which makes the atom paramagnetic. If all of the electrons
are paired, then the atom is diamagnetic.
Periodic Properties—Main Groups
Core electrons are very effective at screening the outer
electrons from the full charge of the nucleus, whereas electrons in
the valence shell do not screen each other very effectively. As a
result, the effective nuclear charge (Zeff) experienced by valence
electrons increases as we move left to right across the main-group
elements because the number of protons in the nucleus increases,
without a corresponding increase in core electrons. The increase
in Zeff is less pronounced with transition metals because the added
electrons enter the core and cancel the added protons.
As a result of the Zeff atomic radius decreases as we go from
left to right across a period of main group elements. Atomic radii
increase as we go down a group because electrons fill the next
higher energy level.
Cations are smaller than their parent atoms; anions are larger
than their parent atoms, but group and period trends are the same
as atomic radius. For an isoelectronic series the radius decreases
with increasing nuclear charge.
Ionization energy is the energy needed to remove an electron
from a gaseous atom; forming a cation. Successive ionization
energies show a sharp increase after all the valence electrons
have been removed, because of the much higher effective nuclear
charge experienced by the core electrons. For the main group
elements, the first ionization energy trend is generally opposite the
atomic radii trend, with smaller atoms having higher first ionization
energies, except for columns 13 (removing p orbital electron) and
column 16 (removing an electron from a full orbital).
Electron affinity measures the energy change when adding an
electron to a gaseous atom; forming an anion. A negative electron
affinity means that the anion is stable; a positive electron affinity
means that the anion is not stable relative to the separate atom
and electron. In general, electron affinities become more negative
from left to right across the main groups except for column 2
(adding an electron to a p orbital), column 15 (adding an electron
to a half-filled orbital) and column 18 (adding an electron to the
next higher energy level).
Bonding
Bonds are classified into three broad groups: ionic bonds are
the result of electrostatic forces between cations and anions;
covalent bonds form when electrons are shared between non-metal
atoms; and metallic bonds, which bind metal cations with mutually
shared valence electrons.
Bonds involve the interaction of valence electrons, which are
represented by electron-dot symbols, called Lewis symbols. The
tendencies of atoms to gain, lose, or share their valence electrons
often follow the octet rule, which can be viewed as an attempt by
atoms to achieve a noble gas electron configuration, which is a
lower energy state.
The strength of the electrostatic attractions between ions is
measured by the lattice energy, which increases with ionic charge
and decreases with distance between ions.
Electronegativity measures the ability of an atom to attract
electrons in a covalent bond. Electronegativity generally increases
from left to right in the periodic table and decreases down a
column. The difference in atoms' electronegativities is used to
determine the polarity of a covalent bond; the greater the
difference, the more polar the bond. A polar molecule has a positive
side (+) and a negative side (–). The separation of charge
produces a dipole, the magnitude of which is given by the dipole
moment. Polar bonds are stronger and shorter than non-polar
bonds.
Bond strength and length is also affected by the number of
shared electrons. Sharing of one pair of electrons produces a single
bond; whereas the sharing of two or three pairs of electrons
produces double or triple bonds, respectively. Multiple bonds are
stronger and shorter than single bonds.
The procedures used for naming two-element, binary,
molecular compounds follow the rules below.
1. The lower electronegative element is written first in the formula
and named as an element.
2. The name of the second element is given an –ide ending.
3. Prefixes are used to indicate the number of atoms of each
element; mono is not used with the first element.
Lewis Structures
Electron distribution in molecules is shown with Lewis
structures, which indicate how many valence electrons are involved
in forming bonds and how many remain as unshared electron pairs.
If we know which atoms are connected to one another, we can
draw Lewis structures for molecules and ions by a simple
procedure, where eight electrons are placed around each atom.
When there are too few valence electrons, then it will be necessary
to add double or triple bonds. When there are too many valence
electrons (and the central atom has at least third energy level
electrons), then it will be necessary to place additional electrons (up
to 10 or 12) around the central atom; forming an expanded octet.
When the total number of valence electrons is an odd number, then
bond and one  bond and a triple bond consists of one  and two 
bonds.
Sometimes a  bond can be placed in more than one location.
In such situations, we describe the molecule by using two or more
resonance structures. The molecule is envisioned as a blend of
these multiple resonance structures and the  bonds are
delocalized; that is, spread among several atoms. The bond order
value represents the actual bond strength and is sum of the  bond
plus a share of the  bond(s).
Hydrocarbons
Carbon molecules (except CO, CO2) are called organic.
Hydrocarbons are organic molecules that contain mostly carbon
and hydrogen. The four groups of hydrocarbons are alkanes,
alkenes, alkynes, and aromatic. The naming of hydrocarbons is
based on the longest continuous chain of carbon atoms in the
structure.
Prefix based on number of carbons in longest chain
1
meth
6
hex
2
eth
7
hept
3
prop
8
oct
4
but
9
non
5
pent
10
dec
Hydrocarbon branches are named according to the number of
carbons in the branch with a "yl" ending and located by the number
of the main chain carbon to which the branch is attached.
Ring structures have the prefix cyclo.
The names of alkenes (double bond) and alkynes (triple bond)
are based on the longest continuous chain of carbon atoms that
contain the multiple bond; with the location of the multiple bond
specified by a numerical prefix.
The chemistry of an organic compound is dominated by the
presence of the functional group.
3
0
no
 planar
Alcohol
Acid
Amine
Trigonal planar 120o
••
2
1
Bent
yes
R–C–R
R–C–O-H
R–N–R
|
||
|
4
0
Tetrahedral
no
O-H
O
R
3
1
Tetrahedral
109o
yes
 pyramidal
Water Soluble
Acid (releases H+)
Base (Absorbs H+)
Isomers
are
substances
that
possess
the
same
molecular
2
2
Bent
yes
formula, but differ in the arrangements of atoms. In structural
5
0
no
 bipyramidal
isomers the bonding arrangements differ. Different isomers are
given different names. Alkenes exhibit not only structural isomerism
o
4
1
Seesaw
yes
Trigonal
90 ,
but geometric isomerism (cis-trans) as well. In geometric isomers
o
bipyramidal
120
3
2
T-shaped
yes
the bonds are the same, but the molecules have different
geometries. Geometric isomerism is possible in alkenes because
2
3
Linear
no
rotation about the C=C bond is restricted.
6
0
Octahedral
no
Gas State
Gases at room temperatures tend to be molecular with low
5
1
Octahedral
90o
 pyramidal
yes
molar mass. Air, a mixture composed mainly of N2 and O2. Some
4
2
 planar
no
liquids and solids can also exist in the gaseous state, where they
Valence-Bond Theory
are known as vapor. Gases' volume can change because they are
Valence-bond theory is an extension of Lewis's notion of
compressible and they mix in all proportions because their
electron-pair bonds. In valence-bond theory, covalent bonds are
component molecules are far apart.
formed when atomic orbitals on neighboring atoms overlap. The
The gas state is characterized by four variables: pressure (P),
bonding electrons occupy the overlap region and are attracted to
volume (V), temperature (T), and quantity (n). Volume is measured
both nuclei simultaneously, which bonds the atoms together.
in liters, temperature in kelvins, and quantity of gas in moles.
To extend valence-bond theory to polyatomic molecules, s, p,
Pressure is the force per unit area. In chemistry, pressure is
and sometimes d orbitals are blended to form hybrid orbitals, which measured in atmospheres (atm), torr (named after Torricelli),
overlap with orbitals on another atom to make a bond. Hybrid
millimeter of mercury (mm Hg) and kilopascals (kPa).
orbitals also hold non-bonding pairs of electrons. A particular mode
1 atm = 101 kPa = 760 torr = 760 mm Hg.
of hybridization can be associated with each of the five common
A barometer is used to measure atmospheric pressure and a
2
electron-domain geometries (linear = sp; trigonal planar = sp ;
manometer is used to measure the pressure of enclosed gases.
tetrahedral = sp3; trigonal bipyramidal = sp3d; and octahedral =
The ideal-gas law equation is PV = nRT, where V is in L, n is
3
2
sp d ).
in moles and T is in K. The term R is the gas constant, which is
Covalent bonds formed between hybridized electrons are
0.0821 when P is in atm or 8.31 when P is in kPa. The conditions
called sigma () bonds, where the electron density lies along the
of 273 K and 1 atm are known as the standard temperature and
line connecting the atoms. Bonds that form between non-hybridized pressure and abbreviated as STP, where the molar volume of all
p orbitals are called pi () bonds. A double bond consists of one 
gases is 22.4 L/mol. Additional equations using molar mass (MM)
are MM = mRT/PV and MM = dRT/P.
it will be necessary to place seven electrons around the atom with
the odd number of valence electrons.
When there are multiple valid Lewis structures for a molecule
or ion, we can determine which is most likely by assigning a formal
charge to each atom, which is the sum of half the bonding electrons
and all the unshared electrons. Most acceptable Lewis structures
will have low formal charges with any negative formal charge on the
more electronegative atom.
VSEPR Model
The valence-shell electron-pair repulsion (VSEPR) model
rationalizes molecular geometries based on the repulsions between
electron domains, which are regions about a central atom where
electrons are likely to be found. Pairs of electrons, bonding and
non-bonding, create domains around an atom, which are as far
apart as possible. Electron domains from non-bonding pairs exert
slightly greater repulsions, which leads to smaller bond angles than
idealized values. The arrangement of electron domains around a
central atom is called the electron domain geometry; the
arrangement of atoms is called the molecular geometry.
Certain molecular shapes have cancelling bond dipoles,
producing a nonpolar molecule, which is one whose dipole moment
is zero. In other shapes the bond dipoles do not cancel and the
molecule is polar (a nonzero dipole moment). In general nonbonding pairs of electrons around the central atom produce polar
molecules.
The table below summarizes the Domain Geometry, Ideal
bond angle, Molecular Geometry and Molecular polarity based on
the number of bonded atoms (–) and non-bonding pairs of
electrons (• •) surrounding the central atom.
Domain
Bond
Molecular
(–)
(• •)
Polar?
Geometry
Angle
Geometry
2
0
Linear
180o
Linear
no
In gas mixtures the total pressure (Ptot) is the sum of the
partial pressures (PA) that gas A would exert if it were present
alone under the same conditions: Ptot = PA + PB ... and the
pressure of gas A is proportional to its mole fraction (XA): PA =
XAPtot. In calculating the quantity of a gas collected over water,
correction must be made for the partial pressure of water vapor.
The kinetic-molecular theory accounts for the properties of an
ideal gas in terms of a set of statements about the nature of gases:
Molecules are in continuous, chaotic motion; the volume of gas
molecules is negligible compared to the volume of their container;
the gas molecules have no attraction for one another; their
collisions are elastic; and the molecules' kinetic energy is
proportional to the absolute temperature: K = 3/2RT.
Molecules of a gas do not all have the same kinetic energy at
a given instant. Their speeds are distributed over a wide range; the
distribution varies with the molar mass of the gas and with
temperature. The root-mean-square speed, u = (3RT/MM)½.
Effusion (rate of escape through a tiny hole into a vacuum)
and diffusion (rate spreading of one gas through another) are
related to molar mass by Graham's law: rA/rB = (MMB/MMA)½.
Departures from ideal behavior increase as pressure
increases (real gases molecules occupy a significant fraction of
container volume resulting in an observed volume that is greater
than predicted) and as temperature decreases (real gases attract
each other resulting in an observed pressure that is less than
predicted).
Phase Change
Substances that are gases or liquids at room temperature are
usually composed of molecules. In gases the intermolecular
attractive forces are negligible compared to the kinetic energies of
the molecules; thus, the molecules are widely separated and
undergo constant, chaotic motion. In liquids the intermolecular
forces are strong enough to keep the molecules in close proximity;
nevertheless, the molecules are free to move with respect to one
another. In solids the inter-particle attractive forces are strong
enough to restrain molecular motion and to force the particles to
occupy specific locations in a three-dimensional arrangement.
Three types of intermolecular forces exist between neutral
molecules: dipole-dipole forces, London dispersion forces, and
hydrogen bonding. London dispersion forces operate between all
molecules and are the result of temporary polarization of the
molecule. Dispersion forces increase in strength with increasing
molecular mass, although molecular shape is also an important
factor. Dipole-dipole forces exist between polar molecules, where
the positive pole of one molecule is attracted to the negative pole
of its neighbor. Hydrogen bonding occurs in compounds containing
N, O or F bonded to H. Hydrogen bonds are stronger than dipoledipole or dispersion forces.
Phase changes are transformations from one state to another.
Changes of solid to liquid (melting), solid to gas (sublimation) and
liquid to gas (vaporization or boiling) absorb energy. The reverse
processes (freezing, deposition, condensation, respectively)
release energy. A gas cannot be liquefied by application of
pressure if the temperature is above its critical temperature.
The vapor pressure is the partial pressure of the vapor when it
is in equilibrium with the liquid. At equilibrium the rate of
evaporation, transfer of molecules from the liquid to the vapor,
equals the rate of condensation, transfer from the vapor to the
liquid. The higher the vapor pressure of a liquid, the more readily it
evaporates and the more volatile it is. Vapor pressure increases
nonlinearly with temperature. Boiling occurs when the vapor
pressure equals the atmospheric pressure. The normal boiling
point occurs at 1 atm pressure.
The equilibria between the solid, liquid and gas phases of a
substance as a function of temperature and pressure are displayed
on a phase diagram. Equilibria between any two phases are
indicated by a line. The line through the melting point usually
slopes slightly to the right as pressure increases, because the solid
is usually more dense than the liquid. The melting point at 1 atm is
the normal melting point. The point on the diagram at which all
three phases coexist in equilibrium is called the triple point.
Crystalline Solids
In a crystalline solid, particles are arranged in a regularly
repeating pattern. An amorphous solid or glass is one whose
particles show no such order.
The properties of solids depend both on the type of particles
and on the attractive forces between them. Molecular solids, which
consist of atoms or molecules held together by intermolecular
forces, are soft and low melting. Covalent network solids, which
consist of atoms held together by covalent bonds that extend
throughout the solid, are hard and high melting. Ionic solids are
hard and brittle and have high melting points. Metallic solids, which
consist of metal cations held together by a sea of electrons, exhibit
a wide range of melting points.
Solubility
Solutions form when one substance disperses uniformly
throughout another. The dissolving medium of the solution (usually
in the greater amount) is called the solvent. The substance
dissolved in a solvent (usually the smaller amount) is called the
solute. The attractive interaction of solvent molecules with solute is
called solvation. When the solvent is water, the interaction is called
hydration. The dissolution of ionic substances in water is promoted
by hydration of the separated ions by the polar water molecules.
The overall change in energy upon solution formation may be
either positive (endothermic) or negative (exothermic), depending
on the relative value of lattice energy (positive) and hydration
energy (negative).
The equilibrium between a saturated solution and undissolved
solute is dynamic; the process of solution and the reverse process,
crystallization, occur simultaneously. For a solution in equilibrium
with undissolved solute, the two processes occur at equal rates,
producing a saturated solution. The amount of solute needed to
form a saturated solution at any particular temperature is the
solubility of that solute at that temperature.
The solubility of one substance in another depends on the
tendency of systems to become more random, by becoming more
dispersed in space, and on the relative intermolecular solute-solute
and solvent-solvent energies compared with solute-solvent
interactions. Polar and ionic solutes tend to dissolve in polar
solvents such as water and alcohol, and nonpolar solutes tend to
dissolve in nonpolar solvents ("like dissolves like"). Liquids that mix
in all proportions are miscible; those that do not dissolve
significantly in one another are immiscible. Hydrogen-bonding
interactions between solute and solvent often play an important
role in determining solubility; for example, ethanol and water,
whose molecules form hydrogen bonds with each other, are
miscible. The solubilities of gases in a liquid are generally
proportional to the pressure of the gas over the solution, as
expressed by Henry's law: Sg = kPg. The solubilities of most solid
solutes in water increase as the temperature increases. In
contrast, the solubilities of gases in water generally decrease with
increasing temperature.
Concentrations of solutions can be expressed quantitatively
by several different measures, including
mass percent: % = 102(msolute/mtotal) or 102(msolvent/mtotal)
mole fraction: Xsolute = molsolute/moltotal or Xsolvent = molsolvent/moltotal
molarity: M = molsolute/Vsolution(L)
molality: m = molsolute/msolvent(kg)
Conversions between concentration units is possible if molar
mass of solute and solvent are known and/or the density of the
solution is known.
Colligative Properties
A physical property of a solution that depends on the
concentration of solute particles present, regardless of the nature
of the solute, is a colligative property. Colligative properties include
vapor-pressure lowering, freezing-point lowering, boiling-point
elevation, and osmotic pressure. The lowering of vapor pressure
follows the equation, Pvap = XsolventPosolvent. A solution containing a
nonvolatile solute possesses a higher boiling point than the pure
solvent. The molal boiling-point constant, Kb, represents the
increase in boiling point for a 1 m solution of solute particles as
compared with the pure solvent. Similarly, the molal freezing-point
constant, Kf, measures the lowering of the freezing point of a
solution for a 1 m solution of solute particles. The temperature
changes are given by the equations Tb = Kbm and Tf = Kfm.
When NaCI dissolves in water, two moles of solute particles are
formed for each mole of dissolved salt. The boiling point or
freezing point is thus elevated or depressed, respectively,
approximately twice as much as that of a nonelectrolyte solution of
the same concentration. The multiplier is called the van't Hoff
factor i. Similar considerations apply to other strong electrolytes.
Osmosis is the movement of solvent molecules through a
semipermeable membrane from a less concentrated to a more
concentrated solution. This net movement of solvent generates an
osmotic pressure  which can be measured in units of gas
pressure, such as atm. The osmotic pressure of a solution as
compared with pure solvent is proportional to the solution molarity:
 = MRT.
Chemical Reactions
One of the important concepts of stoichiometry is the law of
conservation of mass, which states that the total mass of the
products of a chemical reaction is the same as the total mass of
the reactants. Likewise, the same numbers of atoms of each type
are present before and after a chemical reaction. A balanced
chemical equation shows equal numbers of atoms of each element
on each side of the equation. Equations are balanced by placing
coefficients in front of the chemical formulas for the reactants and
products of a reaction, not by changing the subscripts in chemical
formulas.
Among the reaction types described in this unit are (1)
combination reactions, in which two reactants combine to form one
product; (2) decomposition reactions, in which a single reactant
forms two or more products; (3) combustion reactions in oxygen, in
which a hydrocarbon or related compound reacts with O2 to form
CO2 and H2O; (4) electron exchange reactions in which reactants
exchange electrons; (5) ion exchange reactions, in which ions
exchange "partners"; and (6) proton exchange reactions, in which
atoms exchange H+. The last three reactions will be discussed in
more detail later in the year.
Aqueous reactions (electron exchange, ion exchange and
proton exchange) usually involve ions, where only one of the two
ions in an ionic compound or acid react, the other is non-reacting
and is called a spectator ion. In a net ionic equation, those ions
that go through the reaction unchanged are omitted.
The coefficients in a balanced equation give the relative
numbers of moles of reactants and products. To calculate the
grams of a product from the grams of a reactant, first convert
grams of reactant to moles of reactant, then use the coefficients to
convert the number of moles of reactant to moles of product, and
finally convert moles of product to grams of product. Many
reactions occur in water (aqueous). Molarity makes it possible to
convert solution volume to number of moles of solute.
A limiting reactant is completely consumed in a reaction.
When it is used up, the reaction stops, thus limiting the quantities
of products formed. The theoretical yield is the quantity of product
calculated to form when all of the limiting reagent reacts. The
actual yield is always less than the theoretical yield. The percent
yield compares the actual and theoretical yields.
Gravimetric Analysis
The empirical formula can be determined from its percent
composition by calculating the relative number of moles of each
atom in 100 g of the substance. Similarly, the empirical formula
can be determined from the mass of each element in the
compound, or if it is a combustion reaction, from the mass of CO2
and H2O produced. If the substance is molecular in nature, its
molecular formula can be determined from the empirical formula if
the molecular mass is also known.
Volumetric Analysis
Solutions of known molarity can be formed either by adding a
measured mass of solute and diluting it to a known volume or by
the dilution of a more concentrated solution of known
concentration (a stock solution). Adding solvent to the solution (the
process of dilution) decreases the concentration of the solute
without changing the number of moles of solute in the solution
(Mstock)(Vstock) = (Mstandard)(Vstandard).
In titration a solution of known concentration (a standard
solution) is combined with a solution of unknown concentration in
order to determine the moles of unknown. By knowing moles of
unknown, the molar mass of unknown or concentration of unknown
solution can be determined. The point in the titration at which
stoichiometrically equivalent quantities of reactants are brought
together is called the equivalence point. An indicator can be used
to show the end point of the titration, which coincides closely with
the equivalence point.
Change in Enthalpy (H)
Chemical reactions typically involve breaking some bonds
between reactant atoms and forming new bonds. Breaking bonds
absorbs energy, therefore the chemical system gains bond energy
and the surroundings lose energy, typically in the form of heat. In
contrast, forming bonds releases energy; resulting in lose of
energy by the chemical system and a gain in energy by the
surroundings (also in the form of heat).
When energy required to break bonds is greater than the
energy released to form new bonds, then products are at a
higher energy state than reactants (making the product bonds
weaker than the reactant bonds) and energy of the system
increases (+H), which is described as endothermic because the
surroundings typically lose heat energy and cool down.
Alternatively, when energy required to break bonds is less than
the energy released to form new bonds, then products are at a
lower energy state than reactants (making the product bonds
stronger than the reactant bonds) and energy of the system
decreases (–H), which is described as exothermic because the
surroundings typically gain heat energy and warm up. The
change in enthalpy, H, is listed to the right of a balanced
chemical equation. H can be treated in the same way as a
coefficient when using dimensional analysis.
The amount of heat transferred between the system and the
surroundings is measured experimentally by calorimetry. A
calorimeter measures the temperature change accompanying a
process. The temperature change of a calorimeter depends on its
heat capacity, the amount of heat required to raise its temperature
by 1 K. The heat capacity for one mole of a pure substance is
called its molar heat capacity; for one gram of the substance, we
use the term specific heat. Water has a very high specific heat, c =
4.18 J/g•K. The exchange of heat, Q, with the surroundings is the
product of the surrounding medium's specific heat (c), mass (m),
and change in temperature (T), such that Q = mcT. If a Bomb
calorimeter is used, then the bomb constant (C) is in the equation:
Q = (C + mc)T.
Bond energy, B.E., measures the energy needed to break a
covalent bond in a diatomic, gaseous molecule. The bond energy
is approximately the same for any gaseous molecule. Change in
enthalpy is estimated by adding the bond energies of all bonds that
are broken and subtracting the bond energies of all bonds formed:
H = B.E.reactants – B.E.products.
Change in Entropy (S)
All chemical systems have an inherent amount of disorder
because of the complexity of the atomic arrangement within
molecules, the spacing of molecules with respect to each other;
and the overall motion of the system. Increases in complexity,
spacing and overall motion result in increased disorder as
measured by change in entropy, S. A positive S for physical
changes can be predicted based on whether the molecules spread
out. Evaporation, diffusion and effusion have +S values.
Dissolving is more complicated because spreading out solute and
solvent increases disorder, but formation of hydration bonds
between solute and solvent decreases disorder, therefore it is
impossible to predict the sign for S (although most dissolving is
+S. All chemical reactions that result in more moles of gas
products compared to reactants have a +S.
Thermodynamic Data
The standard enthalpy of formation, Hfo, of a substance is the
enthalpy change for the reaction in which one mole of substance is
formed from its constituent elements under standard conditions of
1 atm pressure and 25oC (298 K). For any element in its most
stable state under standard conditions, Hfo = 0 kJ/mol. Most
compounds have negative values of Hfo. Large negative Hfo
indicate a strong bond and stable compound. The standard
entropy So is based on H+ having So = 0 kJ/mol•K (although the AP
exam often lists the values in J/mol•K). The thermodynamic data
chart lists the Hfo and So for common substances.
Hfo applies to situations involving more than one mole, where
o
Hf is multiplied by the number of moles, and involving
decomposition, where H = -Hfo. An important use of Hfo and So
is calculate H and S for a wide variety of reactions under
laboratory conditions, where H  Ho = Hfoproducts – Hforeactants
and S  So = Soproducts – Soreactants.
H depends only on the initial and final states of the system.
Thus, the enthalpy change of a process is the same whether the
process is carried out in one step or in a series of steps. Hess's
law states that if a reaction is carried out in a series of steps, H
for the reaction will be equal to the sum of the enthalpy changes
for the steps. We can therefore calculate H for any process, as
long as we can write the process as a series of steps for which H
is known.
Change in Free Energy (G)
The Gibbs free energy (or just free energy) G combines
enthalpy and entropy. For processes that occur at constant
temperature, G = H – TS. The sign of G relates to the
spontaneity of the process. When G is negative, the process is
spontaneous. When G is positive, the process is
nonspontaneous; the reverse process is spontaneous. At
equilibrium the process is reversible and G = 0 kJ.
The values of H and S generally do not vary much with
temperature. As a consequence, the dependence of G with
temperature is governed mainly by the value of T in the expression
G = H –TS. The threshold temperature, Tthreshold = H/S, is
when a reaction goes from spontaneous  nonspontaneous. This
only occurs when H and S are both positive or both negative.
When H and S are both positive, the reaction is spontaneous at
all temperatures above the threshold. When they are both
negative, the reaction is spontaneous at all temperatures below
the threshold. When H and S have opposite signs, then the
reaction is spontaneous at all temperatures (-H and +S) or
spontaneous at no temperature (+H and –S).
Reaction Rate
Reaction rates are expressed as changes in molarity per unit
time, M/s. For most reactions, a plot of molarity versus time shows
that the rate slows down as the reaction proceeds. The
instantaneous rate is the slope of a line drawn tangent to the
concentration-versus-time curve at a specific time. Rates can be
written in terms of products, which have positive rates, or in terms
of reactants, which have negative rates. The reaction rates for the
each reactant and product are proportional to the coefficients in
the balanced equation.
The quantitative relationship between rate and concentration
is expressed by a rate law, which has the form: rate = k[A]m[B]n,
where A and B are reactants k is called the rate constant, and the
exponents m and n are called reaction orders. The sum of the
reaction orders gives the overall reaction order. Reaction orders
must be determined experimentally. The unit of the rate constant
depends on the overall reaction order. The unit for k is Mxt-1, where
x = 1 – overall order.
The overall order for the reaction is used to determine which
equation is used to determine the reaction rate, time needed to
change from [A]o to [A]t, the half-life (t½) and the linear plot.
Order
0
1
2
k
k[A]
k[A]2
rate =
kt =
[A]o – [A]t
ln([A]o/[A]t)
1/[A]t – 1/[A]o
t½ =
[A]o/2k
ln2/k
1/k[A]o
[A] vs. t
ln[A] vs. t
1/[A] vs. t
linear Plot
(Note: radioactive decay is a first order process)
Collision Model
The collision model, which assumes that reactions occur as a
result of collisions between molecules, helps explain why the rate
constant increases with increasing temperature. At higher
temperature, reactant molecules have more kinetic energy and
their collisions are more energetic. The minimum energy required
for a reaction to occur is called the activation energy Ea. A collision
with energy Ea or greater can cause the atoms of the colliding
molecules to reach the activated complex, which is the highest
energy arrangement in the pathway from reactants to products.
Even if a collision is energetic enough, it may not lead to reaction;
the reactants must also have correct orientation for a collision to
be effective. Because the kinetic energy depends on temperature,
the rate constant is dependent on temperature. The two point
equation is ln(k1/k2) = (Ea/R)(1/T2 – 1/T1), where R = 8.31 J/mol•K. The
slope of lnk versus 1/T equals -Ea/R.
Reaction Mechanism
Many reactions occur by a multistep mechanism, involving two
or more elementary reactions, or steps. A reaction mechanism
details the individual steps that occur in the course of a reaction.
Each of these steps has 1 or 2 reactants and low activation
energy. The rate law for each step corresponds exactly to the
number of reactant molecules, so that reactant coefficients
become exponents in the rate law. An intermediate is produced in
one elementary step and is consumed in a later elementary step
and therefore does not appear in the overall equation for the
reaction. When a mechanism has several elementary steps, the
overall rate is limited by the slowest elementary step, called the
rate-determining step.
A catalyst is a substance that increases the rate of a reaction
without undergoing a net chemical change itself. It does so by
providing a different mechanism for the reaction, one that has
lower activation energy. A homogeneous catalyst is one that is in
the same phase as the reactants. It is consumed in the slow step
and reappears in a later step. As a result, it is not included in the
overall reaction, but is included in the rate law. A heterogeneous
catalyst has a different phase from the reactants and is written
above the reaction arrow.
Unit 1 Practice Multiple Choice
1.
2.
3.
4.
Based on the data, the density of the solid in g/mL is
Mass of metal
19.611 g
Volume of water
12.4 mL
Volume of water + metal
14.9 mL
(A) 7.8444 (B) 7.844 (C) 7.84
(D) 7.8
Which scientist is correctly matched with the discovery?
(A) Millikan discovered the electron charge-to-mass ratio.
(B) Thomson discovered the charge of an electron.
(C) Bohr discovered the four quantum numbers.
(D) Rutherford discovered the nucleus.
Which represents a pair of isotopes?
(A) 146C and 147N
(B) 189F and 3517Cl
(C) 5626Fe2+ and 5626Fe3+ (D) 3517Cl and 3617Cl
Copper has two isotopes, 63Cu and 65Cu. What is the
abundance of 63Cu if the average atomic mass is 63.5?
(A) 90%
(B) 75%
(C) 50%
(D) 20%
5.
6.
7.
8.
9.
10.
11.
12.
13.
Which of the following is correct about beta particles?
I. mass number of 4 and a charge of +2
II. more penetrating than alpha particles
III. electron
(A) I only
(B) III only (C) I and II (D) II and III
For the types of radiation given, which is the correct order of
increasing ability to penetrate a piece of lead?
(A)  <  < 
(B)  <  < 
(C)  <  < 
(D)  <  < 
249 Cm is radioactive and decays by the loss of one beta
96
particle. The other product is
(A) 24594Pu (B) 24997Bk (C) 24896Cm (D) 25096Cm
251 Cf  2 1 n + 131 Xe + . . .
98
0
54
What is the missing product in the nuclear reaction?
(A) 11842Mo (B) 11844Ru (C) 12042Mo (D) 12044Ru
The radioactive decay of C-14 to N-14 occurs by
(A) beta particle emission (B)alpha particle emission
(C) positron emission
(D) electron capture
What is the resulting nucleus after 21484Po emits 2  and 2 ?
(A) 20683Bi (B) 21083Bi (C) 20682Pb (D) 20882Pb
235 U + 1 n  141 Cs + 3 1 n + X
92
0
55
0
Neutron bombardment of uranium can induce the reaction
represented above. Nuclide X is which of the following?
(A) 9235Br
(B) 9435Br
(C) 9137Rb (D) 9237Rb
If 87.5 % of a sample of pure Pb-210 decays in 36 days,
what is the half-life of Pb-210?
(A) 6 days (B) 8 days (C) 12 days (D) 14 days
The half-life of isotope Y is 12 minutes. What mass of Y was
originally present if 1 g is left after 60 minutes?
(A) 8 g
(B) 16 g
(C) 24 g
(D) 32 g
Unit 2 Practice Multiple Choice
Briefly explain why the answer is correct in the space provided.
Questions 1-3
(A) Heisenberg uncertainty principle
(B) Pauli exclusion principle
(C) Hund's rule
(D) Shielding effect
1. Can be used to predict that a gaseous carbon atom in its
ground state is paramagnetic.
2. States that an orbital can hold no more than two electrons.
3. Predicts that it is impossible to determine simultaneously the
exact position and the exact velocity of an electron.
4. Which set of quantum numbers describes the highest energy
valence electron in a ground-state gallium atom (Z = 31)?
(A) 4,0,0,½ (B) 4,0,1,½ (C) 4,1,1,½ (D) 4,1,2,½
5. Which has the outer electronic configuration, s2p3?
(A) Si
(B) Cl
(C) Se
(D) As
Questions 6-9 refer to atoms with the atomic orbitals shown.
(A) 1s 2s
(B) [He] 2s2p
(C) [He] 2s2p (D) [Ar] 4s3d
6. Represents an atom that is chemically unreactive.
7. Represents an atom in an excited state.
8. Represents an atom that has four valence electrons.
9. Represents an atom of a transition metal.
Questions 10-12
(A) 1s2 2s22p5 3s23p5
(B) 1s2 2s22p6 3s23p6
(C) 1s2 2s22p62d10 3s23p6
(D) 1s2 2s22p6 3s23p63d3 4s2
10. An impossible electronic configuration
11. The ground-state configuration for a halogen anion
12. The ground-state configuration for an alkaline earth cation
13. Which represents the ground state for the Mn3+ ion?
(A) 1s2 2s22p6 3s23p63d4
(B) 1s2 2s22p6 3s23p63d5 4s2
(C) 1s2 2s22p6 3s23p63d2 4s2
(D) 1s2 2s22p6 3s23p63d8 4s2
14. In which group are the three species isoelectronic?
(A) S2-, K+, Ca2+
(B) Sc, Ti, V2+
22(C) O , S , CI
(D) Mg2+, Ca2+, Sr2+
15. The ionization energies for element X are listed in the table
below. On the basis of the data, element X is most likely
Ionization Energies for element X (kJ mol-1)
First
Second
Third
Fourth
Fifth
580
1,815
2,740
11,600
14,800
(A) Na
(B) Mg
(C) Al
(D) Si
16. In the periodic table, as the atomic number increases from 11
to 17, what happens to the atomic radius?
(A) It remains constant. (B) It increases only.
(C) It decreases only.
(D) It increases, then decreases.
17. Which elements have most nearly the same atomic radius?
(A) Be, B, C, N
(B) Ne, Ar, Kr, Xe
(C) Mg, Ca, Sr, Ba
(D) Cr, Mn, Fe, Co
Questions 18-19 Use the following options.
(A) O
(B) Rb
(C) N
(D) Mg
18. What element has the most negative electron affinity?
19. Which of the elements above has the smallest ionic radius for
its most commonly found ion?
20.
Ca, V, Co, Zn, As
Which gaseous atoms of the above are paramagnetic?
(A) Ca and As only
(B) Zn and As only
(C) Ca, V, and Co only
(D) V, Co, and As only
21. Which property generally decreases across the periodic table
from sodium to chlorine?
(A) 1st ionization energy (B) Atomic mass
(C) Ionic radius
(D) Atomic radius
Questions 22-23 Consider a ground state atom of each element.
(A) S
(B) Ca
(C) Ga
(D) Sb
22. The atom that contains exactly two unpaired electrons
23. The atom that contains only one electron in the highest
occupied energy sublevel
Unit 3 Practice Multiple Choice
1.
Types of hybridization exhibited by the C atoms in propene,
CH3CHCH2 include which of the following?
I. sp
II. sp2
III. sp3
(A) I only
(B) II only (C) III only (D) II and III
2. Which molecule contains 1 sigma () and 2 pi () bonds?
(A) H2
(B) F2
(C) N2
(D) O2
3. Which molecule has the shortest bond length?
(A) N2
(B) O2
(C) Cl2
(D) Br2
4. Which molecule has only one unshared pair of valence
electrons?
(A) Cl2
(B) NH3
(C) H2O2
(D) N2
5. The electron pairs in a molecule where the central atom
exhibits sp3d2 hybrid orbitals are directed toward the corners of
(A) a tetrahedron
(B) a square pyramid
(C) a trigonal bipyramid (D) an octahedron
6. The SbCl5 molecule has trigonal bipyramid structure.
Therefore, the hybridization of Sb orbitals should be
(A) sp2
(B) sp3
(C) dsp2
(D) dsp3
7. For which molecule are resonance structures necessary to
describe the bonding satisfactorily?
(A) H2S
(B) SO2
(C) CO2
(D) OF2
8. Which molecules have planar configurations?
I. BCl3
II. CHCl3
III. NCl3
(A) I only
(B) II only (C) III only (D) II and III
9.
CCl4, CO2, PCl3, PCl5, SF6
Which does NOT describe any of the molecules above?
(A) Linear
(B) Octahedral
(C) Square planar
(D) Tetrahedral
10. The geometry of the SO3 molecule is best described as
(A) trigonal planar
(B) trigonal pyramidal
(C) square pyramidal
(D) bent
11. Pi () bonding occurs in each of the following EXCEPT
(A) CO2
(B) C2H4
(C) CN(D) CH4
12. According to the VSEPR model, the progressive decrease in
the bond angles in the series of molecules CH4, NH3, and
H2O is best accounted for by the
(A) increasing strength of the bonds
(B) decreasing size of the central atom
(C) increasing electronegativity of the central atom
(D) increasing number of unshared pairs of electrons
Questions 13-15 refer to the following molecules.
(A) PH3
(B) H2O
(C) CH4
(D) C2H4
13. The molecule with only one double bond
14. The molecule with the largest dipole moment
15. The molecule that has trigonal pyramidal geometry
16. Which molecule has a zero dipole moment?
(A) HCN
(B) NH3
(C) SO2
(D) PF5
17. Which molecule has the largest dipole moment?
(A) CO
(B) CO2
(C) O2
(D) HF
18. Which molecule has a dipole moment of zero?
(A) C6H6 (benzene)
(B) NO
(C) SO2
(D) NH3
19. Which pair of atoms should form the most polar bond?
(A) F and B
(B) C and O
(C) F and O
(D) N and F
20. Which pair of ions should have the highest lattice energy?
(A) Na+ and Br(B) Li+ and F(C) Cs+ and F(D) Li+ and O221. Which compound has the greatest lattice energy?
(A) BaO
(B) MgO
(C) CaS
(D) MgS
22. Which molecule has the weakest bond?
(A) CO
(B) O2
(C) Cl2
(D) N2
23. How are the bonding pairs arranged in the best Lewis
structure for ozone, O3?
(A) O–O–O (B) O=O–O (C) OO–O (D) O=O=O
24. Which species has the shortest bond length?
(A) CN(B) O2
(C) SO2
(D) SO3
25. Which species has a valid non-octet Lewis structure?
(A) GeCl4 (B) SiF4
(C) NH4+
(D) SeCl4
26. The Lewis structure for SeS2 with zero formal charge has
(A) 2 bonding pairs and 7 nonbonding pairs of electrons.
(B) 2 bonding pairs and 6 nonbonding pairs of electrons.
(C) 3 bonding pairs and 6 nonbonding pairs of electrons.
(D) 4 bonding pairs and 5 nonbonding pairs of electrons.
27. Which molecular shape cannot exhibit geometric isomerism?
(A) tetrahedron
(B)square planar
(C) trigonal bipyramid
(D)octahedron
28. Which molecule is NOT polar?
(A) H2O
(B) CO2
(C) NO2
(D) SO2
29. Which species has sp2 hybridization for the central atom?
(A) C2H2
(B) SO32(C) O3
(D) BrI3
30. In which species is the F-X-F bond angle the smallest?
(A) NF3
(B) BF3
(C) CF4
(D) BrF3
31. For ClF3, the electron domain geometry of Cl and the
molecular geometry are, respectively,
(A) trigonal planar and trigonal planar.
(B) trigonal planar and trigonal bipyramidal.
(C) trigonal bipyramidal and trigonal planar.
(D) trigonal bipyramidal and T -shaped.
32. The size of the H-N-H bond angles of the following species
increases in which order?
(A) NH3 < NH4+ < NH2(B) NH3 < NH2- < NH4+
(C) NH2- < NH3 < NH4+
(D) NH2- < NH4+ < NH3
33. What is the molecular geometry and polarity of BF3?
(A) trigonal pyramidal and polar
(B) trigonal pyramidal and nonpolar
(C) trigonal planar and polar
(D) trigonal planar and nonpolar
34. Which set does not contain a linear species?
(A) CO2, SO2, NO2
(B) H2O, HCN, BeI2
(C) OCN-, C2H2, OF2
(D) H2S, CIO2-, NH2-
35. The hybrid orbitals of nitrogen in N2O4 are
(A) sp
(B) sp2
(C) sp3
(D) sp3d
36. How many sigma and how many pi bonds are in
CH2=CH–CH2–C–CH3?
||
O
(A) 5 sigma and 2 pi.
(B) 8 sigma and 4 pi.
(C) 11 sigma and 2 pi.
(D) 13 sigma and 2 pi.
37. What is the best estimate of the H-O-H bond angle in H3O+?
(A) 109.5o (B) 107o
(C) 90o
(D) 120o
38. Which of the following pairs of compounds are isomers?
(A) CH3–CH2–CH2–CH3 and CH3–CH(CH3)–CH3
(B) CH3–CH(CH3)–CH3 and CH3–CH3–C=CH2
(C) CH3–O–CH3 and CH3–CO–CH3
(D) CH3–OH and CH3–CH2–OH
39.
CH3–CH2–CH2Br
Which structural formulas represents an isomer of the
compound that has the above structural formula?
(A) CH2Br–CH2–CH2
(B) CH3–CHBr–CH3
(C) CH2Br–CH2–CH2Br (D) CH3–CH2–CH2–CH2Br
Unit 4 Practice Multiple Choice
Questions 1-2 The molecules have the normal boiling points.
Molecule
HF
HCl
HBr
HI
Boiling Point, oC
+19
-85
-67
-35
1. The relatively high boiling point of HF can be correctly
explained by which of the following?
(A) HF gas is more ideal.
(B) HF molecules have a smaller dipole moment.
(C) HF is much less soluble in water.
(D) HF molecules tend to form hydrogen bonds.
2. The increasing boiling points for HCl, HBr and HI can be best
explained because of the increase in
(A) dispersion force
(B) dipole moment
(C) valence electrons
(D) hydrogen bonding
3. A sample of an ideal gas is cooled from 50oC to 25oC in a
sealed container of constant volume. Which of the following
values for the gas will decrease?
I. The average kinetic energy of the molecules
II. The average distance between the molecules
III. The average speed of the molecules
(A) I only
(B) II only (C) III only (D) I and III
Questions 4-7 refer to the phase diagram of a pure substance.
4.
5.
6.
7.
Which phase is most dense?
(A) solid
(B) liquid
(C) gas
(D) can't determine
Which occurs when the temperature increases from 0°C to
40°C at a constant pressure of 0.5 atm?
(A) Sublimation
(B) Condensation
(C) Freezing
(D) Fusion
Which occurs when the pressure increases from 0.5 to 1.5
atm at a constant temperature of 60°C?
(A) Sublimation
(B) Condensation
(C) Freezing
(D) Fusion
The normal boiling point of the substance is closest to
(A) 20oC
(B) 40oC
(C) 70oC
(D) 100oC
Questions 8-9The graph shows the temperature of a pure solid
substance as it is heated at a constant rate to a gas.
8.
Pure liquid exists at time
(A) t1
(B) t2
(C) t3
(D) t4
9. Which of the following best describes what happens to the
substance between t4 and t5?
(A) The molecules are leaving the liquid phase.
(B) The solid and liquid phases coexist in equilibrium.
(C) The vapor pressure of the substance is decreasing.
(D) The average intermolecular distance is decreasing.
10. Which actions would be likely to change the boiling point of a
sample of a pure liquid in an open container?
(A) Placing it in a smaller container
(B) Increasing moles of the liquid in the container
(C) Moving the container to a higher altitude
(D) Increase the setting on the hot plate
11. Gas in a closed rigid container is heated until its absolute
temperature is doubled, which is also doubled?
(A) The density of the gas
(B) The pressure of the gas
(C) The average speed of the gas molecules
(D) The number of molecules per liter
12. A 2.00-L of gas at 27oC is heated until its volume is 5.00 L. If
the pressure is constant, the final temperature is
(A) 68oC
(B) 120oC (C) 477oC (D) 677oC
13. Under the same conditions, which of the following gases
effuse at approximately half the rate of NH3?
(A) O2
(B) He
(C) CO2
(D) Cl2
14. What is the partial pressure (in atm) of N2 in a gaseous
mixture, which contains 7.0 moles N2, 2.5 moles O2, and 0.50
mole He at a total pressure of 0.90 atm.
(A) 0.13
(B) 0.27
(C) 0.63
(D) 0.90
Questions 15-16 refer to the following gases at 0°C and 1 atm.
(A) Ne
(B) Xe
(C) O2
(D) CO
15. Has an average atomic or molecular speed closest to that of
N2 molecules at 0°C and 1 atm
16. Has the greatest density
17. A 2-L container will hold about 4 g of which of the following
gases at 0oC and 1 atm?
(A) SO2
(B) N2
(C) CO2
(D) C4H8
18. Which is the same for the structural isomers C2H5OH and
CH3OCH3? (Assume ideal behavior.)
(A) Gaseous densities at STP
(B) Vapor pressures at the same temperature
(C) Boiling points
(D) Melting points
19. As the temperature is raised from 20oC to 40oC, the average
kinetic energy of Ne atoms changes by a factor of
(A) ½
(B) (313/293)½
(C) 313/293
(D) 2
20. The system shown above is at equilibrium at 28°C.
At this temperature, the vapor pressure of water
is 28 mm Hg. The partial pressure
(in mm Hg) of O2(g) is
(A) 28
(B) 56
(C) 133
(D) 161
21. What is the mole fraction of benzene in a mixture of toluene
(pressure = 22 torr) and benzene (pressure = 75 torr)?
(A) 0.23
(B) 0.29
(C) 0.50
(D) 0.77
22. In which of the processes are covalent bonds broken?
(A) I2(s)  I2(g)
(B) CO2(s)  CO2(g)
(C) NaCl(s)  NaCl(l)
(D) C(diamond)  C(g)
23. Of the following compounds, which is the most ionic?
(A) SiCl4
(B) BrCl
(C) PCl3
(D) CaCl2
24. Which of the following oxides is a gas at 25°C and 1 atm?
(A) Rb2O
(B) N2O
(C) Na2O2 (D) SiO2
25. Which of the following has the highest melting point?
(A) S8
(B) I2
(C) SiO2
(D) SO2
26. Under which conditions is O2(g) the most soluble in H2O?
(A) 5.0 atm, 80oC
(B) 5.0 atm, 20oC
(C) 1.0 atm, 80oC
(D) 1.0 atm, 20oC
27. Which is lower for a solution of a volatile solute compared to
the pure solvent?
(A) normal boiling point (B) vapor pressure
(C) normal freezing point (D) osmotic pressure
Questions 28-30 Refer to 0.20 M solutions of the following salts.
(A) NaBr
(B) KI
(C) MgCl2 (D) C6H12O6
28. Has the lowest freezing point
29. Has the lowest conductivity
30. Has the lowest boiling point
31. The mole fraction of ethanol in a 6 molal aqueous solution is
(A) 0.006
(B) 0.1
(C) 0.08
(D) 0.2
32. What additional information is needed to determine the
molality of a 1.0-M glucose (C6H12O6) solution?
(A) Volume
(B) Temperature
(C) Solubility of glucose (D) Density of the solution
33. The mole fraction of toluene (MM = 90) in a benzene (MM= 80)
solution is 0.2. What is the molality of the solution?
(A) 0.2
(B) 0.5
(C) 2
(D) 3
Unit 5 Practice Multiple Choice
_ Fe2O3 + _ CO  _ Fe + _ CO2
When the equation is balanced and reduced to lowest terms,
the coefficient for CO2 is
(A) 1
(B) 2
(C) 3
(D) 4
2.
1 CH3CH2COOH + _ O2  _ CO2 + _ H2O
How many moles of O2 are required to oxidize 1 mole of
CH3CH2COOH according to the reaction above?
(A) 2
(B) 5/2
(C) 3
(D) 7/2
3. C3H8 burns in excess oxygen gas. What is the coefficient for
O2 when the equation is balanced with lowest terms?
(A) 4
(B) 5
(C) 7
(D) 10
4.
CaCO3 + 2 HCl  CaCl2 + CO2 + H2O
What is the mass percent of CaCO3 in a 1.25-g rock that
generate 0.44 g of CO2 when reacted with HCl?
(A) 35 %
(B) 44 %
(C) 67 %
(D) 80 %
5. 8.0 mol of F2 and 1.7 mol of Xe are mixed. When all of the
Xe reacted, 4.6 mol of F2 remain. What is the formula?
(A) XeF
(B) XeF3
(C) XeF4
(D) XeF6
6. What mass of Ca(NO3)2 contains 24 g of oxygen atoms?
(A) 164 g
(B) 96 g
(C) 62 g
(D) 41 g
7. Compounds contain 38 g, 57 g, 76 g, and 114 g of element Q
per mole compound. A possible atomic mass of Q is
(A) 13
(B) 19
(C) 28
(D) 38
8. What is the percent nitrogen by mass in N2O3?
(A) 18 %
(B) 22 %
(C) 36 %
(D) 45 %
9. Which formula is 54 % water by mass?
(A) CaCO3 • 10 H2O
(B) CaCO3 • 6 H2O
(C) CaCO3 • 2 H2O
(D) CaCO3 • H2O
10. Which formula forms 88 g of carbon dioxide and 27 g of water
when burned in excess oxygen?
(A) CH4
(B) C2H2
(C) C4H3
(D) C4H6
11. How many moles of H2O are produced when 0.56 g of C2H4
(MM = 28 g) is burned in excess oxygen?
(A) 0.04
(B) 0.06
(C) 0.08
(D) 0.4
1.
12.
13.
14.
15.
16.
17.
18.
19.
20.
21.
22.
23.
24.
25.
26.
27.
28.
29.
30.
31.
BrO3- + 5 Br- + 6 H+  3 Br2 + 3 H2O
How many moles of Br2 can be produced when 25 mL of
0.20 M BrO3- is mixed with 30 mL of 0.45 M Br-?
(A) 0.0050 (B) 0.0081 (C) 0.014
(D) 0.015
3 Ag + 4 HNO3  3 AgNO3 + NO + 2 H2O
If 0.10 mole of silver is added to 10 mL of 6.0 M nitric acid,
the number of moles of NO gas that can be formed is
(A) 0.015
(B) 0.020
(C) 0.033
(D) 0.045
What is the simplest formula of a compound that contains
1.10 mol of K, 0.55 mol of Te, and 1.65 mol of O?
(A) KTeO (B) KTe2O (C) K2TeO3 (D) K2TeO6
In which compound is the mass ratio of chromium to oxygen
closest to 1.6 to 1.0?
(A) CrO3
(B) CrO2
(C) CrO
(D) Cr2O
In which is the mass percent of magnesium closest to 60 %.
(A) MgO
(B) MgS
(C) MgF2
(D) Mg3N2
A student obtained a percent water in a hydrate that was too
small. Which is the most likely explanation for this?
(A) Hydrate spattered out of the crucible during heating
(B) The anhydrous absorbed moisture after heating.
(C) The amount of hydrate sample used was too small.
(D) The amount of hydrate sample used was too large.
2 N2H4 + N2O4  3 N2 + 4 H2O
What mass of water can be produced when 8.0 g of N2H4
(MM = 32 g) and 9.2 g of N2O4 (MM = 92 g) react?
(A) 9.0 g
(B) 18 g
(C) 36 g
(D) 7.2 g
A student wishes to prepare 2.00 L of 0.100 M KIO3
(MM = 214 g). The proper procedure is to weigh out
(A) 42.8 g of KIO3 and add 2.00 kg of H2O
(B) 42.8 g of KIO3 and add H2O to a final volume of 2.00 L
(C) 21.4 g of KIO3 and add H2O to a final volume of 2.00 L
(D) 42.8 g of KIO3 and add 2.00 L of H2O
The volume of distilled water that is added to 10 mL of 6.0 M
HCI in order to prepare a 0.50 M HCI solution is
(A) 50 mL (B) 60 mL (C) 100 mL (D) 110 mL
What volume of 12 M HCl is diluted to obtain 1.0 L of 3.0-M?
(A) 4.0 mL (B) 40 mL (C) 250 mL (D) 1,000 mL
400 mL of distilled water is added to 200 mL of 0.6 M MgCI2,
what is the resulting concentration of Mg2+?
(A) 0.2 M (B) 0.3 M (C) 0.4 M (D) 0.6 M
When 70. mL of 3.0 M Na2CO3 is added to 30. mL of 1.0 M
NaHCO3 the resulting concentration of Na+ is
(A) 2.0 M
(B) 2.4 M
(C) 4.0 M
d. 4.5 M
The mass of H2SO4 (MM = 98 g) in 50 mL of 6.0-M solution
(A) 3.10 g (B) 29.4 g (C) 300. g (D) 12.0 g
What mass of CuSO4• 5 H2O (MM = 250 g) is required to
prepare 250 mL of a 0.10 M solution?
(A) 4.0 g
(B) 6.3 g
(C) 34 g
(D) 85 g
2 KOH + SO2  K2SO3 + H2O
What mass of SO2 reacts with 1.0 L of 0.25-M KOH?
(A) 4.0 g
(B) 8.0 g
(C) 16 g
(D) 20. g
How many moles Ba(NO3)2 should be added to 300. mL of
0.20-M Fe(NO3)3 to increase the [NO3-] to 1.0 M?
(A) 0.060
(B) 0.12
(C) 0.24
(D) 0.30
What is the concentration of HC2H3O2 if it takes 32 mL of
0.50-M NaOH solution to neutralize 20. mL of the acid?
(A) 1.6 M (B) 0.80 M (C) 0.64 M (D) 0.60 M
2 HCl + Ba(OH)2  BaCl2 + 2 H2O
What volume of 1.5-M HCI neutralizes 25 mL of 1.2-M
Ba(OH)2?
(A) 20. mL (B) 30. mL (C) 40. mL (D) 60. mL
What is the concentration of OH- in a mixture that contains
40. mL of 0.25 M KOH and 60. mL of 0.15 M Ba(OH)2?
(A) 0.10 M (B) 0.19 M (C) 0.28 M (D) 0.40 M
What mass of Au is produced when 0.0500 mol of Au2S3 is
reduced completely with excess H2?
(A) 9.85 g (B) 19.7 g (C) 24.5 g (D) 39.4 g
Unit 6 Practice Multiple Choice
I2(g) + 3 Cl2(g)  2 ICl3(g)
According to the data in the table below, what is the value of
Ho , in kJ, for the reaction represented above?
Bond
CI–CI
I–I
I–Cl
Bond Energy (kJ/mole)
150
240
210
(A) - 870
(B) - 390
(C) +180
(D) + 450
2.
C2H4(g) + 3 O2(g)  2 CO2(g) + 2 H2O(g)
For the reaction, H is -1,300 kJ. What is the value of H, in
kJ, if the combustion produced liquid water rather than water
vapor? (H for H2O(l)  H2O(g) is 45 kJ/mol)
(A) -1,300 (B) -1,210 (C) -1,345 (D) -1,390
3.
CH4 (g) + 2 O2(g)  CO2(g) + 2 H2O(l) Ho = -900 kJ
What is the standard heat of formation of CH4, in kJ/mol?
(HfoH2O = -300 kJ/mol, HfoCO2 = -400 kJ/mol)
(A) -200
(B) -100
(C) 100
(D) 200
4.
H2(g) + ½ O2(g)  H2O(l)
Ho = x
2 Na(s) + ½ O2(g)  Na2O(s)
Ho = y
Na(s) + ½ O2(g) + ½ H2(g)  NaOH(s)
Ho = z
What is H for the reaction: Na2O(s) + H2O(l)  2 NaOH(s)?
(A) x + y + z
(B) x + y – z
(C) x + y - 2z
(D) 2z - x - y
5. Which is true when ice melts at its normal melting point?
(A) H < 0, S > 0, G = 0 (B) H < 0, S < 0, G > 0
(C) H > 0, S < 0, G < 0 (D) H > 0, S > 0, G = 0
6. Which of the following reactions has the largest positive
value of S per mole of Cl2?
(A) H2(g) + Cl2(g)  2 HCl(g)
(B) Cl2(g) + O2(g)  Cl2O(g)
(C) Mg(s) + Cl2(g)  MgCl2(s)
(D) 2 NH4Cl(s)  4 H2(g) + Cl2(g)
7. Ice is added to hot water in an insulated container, which is
then sealed. What has happened to the total energy and the
total entropy when the system reaches equilibrium?
(A) Energy and entropy remain constant
(B) Energy remains constant, entropy decreases
(C) Energy remains constant, entropy increases
(D) Energy decreases, entropy increases
8.
N2(g) + 3 H2(g)  2 NH3(g)
The above reaction is thermodynamically spontaneous at 298
K, but becomes nonspontaneous at higher temperatures.
Which of the following is true at 298 K?
(A) G, H, and S are all positive.
(B) G, H, and S are all negative.
(C) G and H are negative, but S is positive.
(D) G and S are negative, but H is positive.
9.
3 C2H2(g)  C6H6(g)
What is the standard enthalpy change, Ho, for the reaction
represented above?
(HfoC2H2 is 230 kJ•mol-1; HfoC6H6 is 80 kJ•mol-1)
(A) -610 kJ (B) 150 kJ (C) -770 kJ (D) 610 kJ
10. When solutions of NH4SCN and Ba(OH)2 are mixed in a
closed container, the temperature drops and a gas is
produced. Which of the following indicates the correct signs
for G, H, and S for the process?
(A) –G –H –S (B) –G +H –S
(C) –G +H +S (D) +G –H +S
11.
X(s)  X(l)
Which of the following is true for any substance undergoing
the process represented above at its normal melting point?
(A) S < 0
(B) H = 0
(C) H = TG
(D) H = TS
12. For a reaction, Ho = -150 kg/mol and So = -50 J/mol•K.
Which statement is true about this reaction?
(A) It is spontaneous at high temperature only.
(B) It is spontaneous at low temperature only.
(C) It is spontaneous at all temperatures.
(D) It is non-spontaneous at all temperatures.
1.
Unit 7 Practice Multiple Choice
Questions 1-2 refer to the following types of energy.
(A) Activation energy
(B) Free energy
(C) Ionization energy
(D) Lattice energy
1. The energy in a chemical or physical change that is available
to do useful work
2. The energy required to form the transition state in a chemical
reaction
3. For the reaction whose rate law is, Rate = k[X], a plot of
which of the following vs. time is a straight line?
(A) [X]
(B) 1/[X]
(C) ln[X]
(D) [X]2
4. The initial-rate data in the table were obtained for the
reaction: 2 NO(g) + O2(g)  2 NO2(g).
[NO]o
[O2]o
Initial Rate of Formation
Exp
(mol•L-1)
(mol•L-1)
of NO2 (mol•L-1•s-1)
1
0.10
0.10
2.5
2
0.20
0.10
5.0
3
0.20
0.40
80.
What is the experimental rate law?
(A) Rate = k[NO][O2]
(B) Rate = k[NO][O2]2
2
(C) Rate = k[NO] [O2]
(D) Rate = k[NO]2[O2]2
Questions 5-6 The oxidation of iodide ions by arsenic acid in
acidic aqueous solution occurs according to the reaction.
H3AsO4 + 3 I- + 2 H3O+  H3AsO3 + I3- + H2O
The rate law is: Rate = k[H3AsO4][I-][H3O+]
5. What is the order of the reaction with respect to I-?
(A) 0
(B) 1
(C) 2
(D) 3
6. According to the rate law for the reaction, an increase in the
concentration of H3O+ has what effect on this reaction?
(A) The rate of reaction increases.
(B) The rate of reaction decreases.
(C) The value of the rate constant increases.
(D) The value of the rate constant decreases.
Questions 7-8 The concentrations of X measured over a period of
time for the reaction X + Y  Z is graphed below.
What is the half-life for this reaction?
(A) 2 minutes
(B) 7 minutes
(C) 10 minutes
(D) half-life is not constant
8. What is the order of reaction with respect to X?
(A) 0
(B) 1
(C) 2
(D) can't be determined
9. How long does it take in minutes for the partial pressure of the
reactant in a first order reaction to decrease from 1.0 atm to
0.125 atm, where the half-life is 19 minutes?
(A) 38
(B) 57
(C) 76
(D) 152
10. Which is associated with relatively slow reaction rates?
(A) The presence of a catalyst
(B) High temperature
(C) High concentration of reactants
(D) Strong bonds in reactant molecules
11. Which of the following best describes the role of a match in
setting a fire?
(A) The match decreases Ea for the slow step.
(B) The match increases the concentration of the reactant.
(C) The match supplies Ea for the combustion reaction.
(D) The match provides a more favorable activated complex
for the combustion reaction.
12.
2 A(g) + B(g)  2 C(g)
Which best explains the observation that there is no change in
reaction rate when the concentration of B(g) is doubled?
(A) The order of the reaction with respect to B is 1.
(B) The overall order of the reaction is 0.
(C) B is not involved in the rate-determined step.
(D) substance B is a catalyst.
13.
2 NO + O2  2 NO2
The mechanism for the overall reaction is the following.
(1) NO + NO  N2O2
slow
(2) N2O2 + O2  2 NO2
fast
Which rate law is consistent with this mechanism?
(A) Rate = k[NO]2
(B) Rate = k[NO][O2]-1
(C) Rate = k[NO]2[O2]-1
(D) Rate = k[NO]2[O2]
14. A reaction was observed for 20 days and the percentage of
the reactant remaining after each day was recorded below.
time 0
1
2
3
4
5
6
7
10
20
% 100 79 63 50 41 31 25 20
10
1
Which best describes the order and half-life of the reaction?
(A) First order with a 3 day half-life
(B) First order with a 10 day half-life
(C) Second order with a 3 day half-life
(D) Second order with a 6 day half-life
Questions 15-17 refer to the proposed steps for the catalyzed
reaction between Ce4+ and Tl+.
Step 1: Ce4+ + Mn2+  Ce3+ + Mn3+
Step 2: Ce4+ + Mn3+  Ce3+ + Mn4+
Step 3:
Mn4+ + TI+  Tl3+ + Mn2+
15. The products of the overall catalyzed reaction are
(A) Ce4+ and TI+
(B) Ce3+ and Tl3+
(C) Ce3+ and Mn3+
(D) Ce3+ and Mn4+
16. The catalyst for the reaction is
(A) Ce4+
(B) Mn2+
(C) Ce3+
(D) Mn4+
17. Intermediates in the reaction are
(A) Ce4+ and Ce3+
(B) Ce3+ and Mn3+
(C) Tl+ and Tl3+
(D) Mn3+ and Mn4+
18.
Rate = k[M][N]2
If the concentrations of M and N are doubled, the reaction
rate will increase by a factor of
(A) 2
(B) 4
(C) 6
(D) 8
Questions 19-20 The energy diagram for the reaction
X + Y  Z is shown below.
7.
19. Which of the following is true for this reaction?
(A) The reaction is exothermic where Ea > |H|
(B) The reaction is endothermic where Ea > |H|.
(C) The reaction is exothermic where Ea < |H|.
(D) The reaction is endothermic where Ea < |H|.
20. The addition of a catalyst to this reaction would cause a
change in which of the indicated energy differences?
(A) I only
(B) II only (C) Ill only (D) I and II only
Unit 1 Practice Free Response
7.
15. Hydrogen atoms in the ground state are ionized by UV light.
a. Calculate the energy needed to ionize an electron from n =
1? (En = -2.18 x 10-18/n2 J).
b.
Calculate the wavelength of UV light.
(Ephoton = (2.00 x 10-25 J•m)/)
16. Answer the following questions regarding light and its
interactions with molecules, atoms, and ions.
a. The longest wavelength of light with enough energy to
break the CI-CI bond in CI2(g) is 495 nm.
(1) Calculate the frequency in s-1 of the light.
(2) Calculate the energy in J of a photon of the light.
8.
Use the principles of atomic structure to explain each.
a. The atomic radius of Li is larger than that of Be.
b.
The electron affinity for K is less than Ca.
c.
The first ionization energy of Se is less than As.
Consider the element strontium (Sr). Justify each answer.
a. What is the outer electron configuration of Sr?
(3) Calculate the energy in kJ•mol-1 of the CI-CI bond.
b.
A certain line in the spectrum of atomic hydrogen is
associated with the electronic transition in the H atom
from the sixth energy level (n = 6) to the second energy
level (n = 2).
(1) Indicate whether the H atom emits energy or
whether it absorbs energy during the transition.
Justify your answer.
(2) Calculate the wavelength in nm of the radiation
associated with the spectral line.
Unit 2 Practice Free Response
6.
The table shows the first three ionization energies in kJ/mol
for third period elements, which are numbered randomly.
Use the information to answer the questions.
Element
First
Second
Third
1
1,251
2,300
3,820
2
496
4,562
6,910
3
738
1,451
7,733
4
578
1817
2745
a. Which element is most metallic in character? Explain.
Identify element 3. Explain.
How does the atomic radius of Sr compare to Rb?
c.
How does the atomic radius of Sr compare to Ca?
d.
Compare Sr to Ca and Rb in first ionization energy.
e.
How does the Sr2+ ion compare in size to the Sr atom?
f.
How does the Sr2+ ion compare in size to the Br- ion?
g.
As successive electrons are removed from the Sr atom,
where does the largest jump in ionization energy occur?
h.
Is strontium diamagnetic or paramagnetic?
Unit 3 Practice Free Response
2.
b.
b.
There are several oxides of nitrogen; among the more
common are N2O, NO2 and NO3-.
a. Draw the Lewis structures of these molecules.
c. Complete the following.
Electron configuration of element 3
Common ion charge of element 2
Chemical symbol of element 2
Element with the smallest atomic radius
Chemical symbol for element 4
b.
N2O
NO2
NO3Which of these molecules "violate" the octet rule?
Explain.
c.
Draw resonance structures of N2O.
d.
For each resonance structure from question 2c,
calculate the formal charge and evaluate which structure
is most likely. Explain.
8.
Compounds of Xe and F form molecules where the
hybridization of Xe is sp3d and sp3d2. Write the formula and
draw the Lewis structure for the two molecules.
sp3d
sp3d2
1.
a.
3.327 g of an unknown gas occupies 1.00-L at 25oC and
103 kPa. What is the molar mass of the gas?
b.
What is the density of this gas at STP (standard
temperature—0oC, and pressure—1 atm)?
c.
Which noble gas would have twice the effusion rate?
Unit 4 Practice Free Response
e.
f.
Which side of the N–O bond is +? Explain.
Rank the strength of the N–O bond in order of strongest
(1) to weakest (3). Explain your answer.
N2O
NO2
NO3-
2.
4.
Consider the ion SF3+.
a. Draw a Lewis structure.
b.
5.
Predict whether the F-S-F bond angle is equal to,
greater than or less than 109.5°. Explain
Consider the ion SF5a. Draw a Lewis structure.
b.
Calculate the total moles of gas.
c.
Calculate the partial pressure of each gas.
3.
Explain why methane gas does not behave as an ideal gas at
low temperatures and high pressures.
4.
10 g of water is added to a 10.0-L container filled with dry air at
20oC (PH2O = 20 torr). The container is sealed.
a. How many grams of the water will evaporate?
Identify the type of hybridization exhibited by sulfur.
b.
c. Identify the electron-domain and molecular geometries.
Electron domain geometry
Molecular geometry
6.
b.
Identify the type of hybridization exhibited by sulfur.
c. Identify the electron-domain and molecular geometries.
Electron-domain geometry
Molecular geometry
d.
N2 with V = 200. mL, P = 99.7 kPa, and T = 27.0oC is mixed
with O2 and transferred to a 750.-mL container at 27.0oC. The
total pressure of the mixture is 90.4 kPa, at 27.0oC.
a. Calculate the moles of N2.
Two Lewis structures can be drawn for the OPF3 molecule.
Structure 1
Structure 2
..
:O:
:O:
.. | ..
.. || ..
:F–P–F:
:F–P–F:
.. | ..
.. | ..
:F:
:F:
..
..
Which Lewis structures best represents a molecule of OPF3?
Justify your answer in terms of formal charge.
5.
Would the amount of water that evaporates increase (),
remain the same (=) or decrease () for the following
changes?
=


Use a 5.0 L container
Use humid air
Raise the temperature to 25oC
Add 20.0 g of water
Consider the following solids.
a. Rank the solids from highest melting point (1) to lowest.
CH4
H2O
MgO
Na
NaCl
SiO2
b.
Justify your relative ranking of CH4 and H2O.
c.
Justify your relative ranking of MgO and NaCl.
6.
Explain the following observations.
a. NH3 boils at 240 K, whereas NF3 boils at 144 K.
b.
7.
8.
What is the molality of the solution?
d.
The vapor pressure of pure ethanol at 25oC is 59.0 mm
Hg. What is the vapor pressure of ethanol in the solution
at this temperature?
e.
What is the osmotic pressure of the solution at 25oC?
f.
What is the boiling point of the solution? The normal
boiling point of ethanol is 78.26oC (Kb = 1.22oC/m).
g.
The molar mass of cortisone acetate is determined by
dissolving 2.50 g in 50.0 g camphor (Kf = 40.0oC/m). The
freezing point of the mixture is 173.44oC; that of pure
camphor is 178.40oC. What is the molar mass of
cortisone acetate?
At 25°C and 1 atm, F2 is a gas, whereas I2 is a solid.
Hydrogen gas is produced when aluminum foil is added to a
solution of hydrochloric acid.
a. The hydrogen is collected over water at 25oC and a total
pressure of 756 torr. What is the mole fraction of H2(g)
in the wet gas? (PH2O) at 25oC = 23.8 torr.
b.
If 255 mL of wet gas is collected, what is the yield of
hydrogen in grams?
c.
What is the density of the wet gas?
When CaCl2 is added to water, the temperature of the
solution decreases.
a. Justify which bond is stronger; hydration bonds between
ions and water or ionic bonds between ions?
b.
c.
Justify why you expect CaCl2 to be more or less soluble
in warm water compared to cold water?
11. Explain the following observations. Your responses must
include specific information about all substances.
a. When table salt (NaCI) and sugar (C12H22O11) are
dissolved in water, it is observed that
(1) both solutions have higher boiling points than pure
water
Unit 5 Practice Free Response
6.
Answer the following questions about acetylsalicylic acid.
a. What is mass percent of acetylsalicylic acid in a 2.00-g
tablet that contains 0.325 g acetylsalicylic acid?
b.
Acetylsalicylic acid contains H, C and O. Combustion of
3.00 g yields 1.20 g H2O and 3.72 L of CO2 at 50oC and
1.07 atm, calculate the mass of each element.
c.
Determine the empirical formula of acetylsalicylic acid.
d.
1.625 g of pure acetylsalicylic acid reacts with NaOH.
The reaction requires 88.43 mL of 0.102 M NaOH.
(1) Calculate the molar mass of the acid. (it takes one
mole of acid to react with each mole of NaOH)
(2) the boiling point of 0.10 M NaCI(aq) is higher than
that of 0.10 M C12H22O11(aq).
b.
Ammonia, NH3, is very soluble in water, whereas
phosphine, PH3, is only moderately soluble in water.
12. Consider camphor, C10H16O, a substance obtained from the
Formosa camphor tree. It has considerable use in the polymer
and drug industry. A solution of camphor is prepared by mixing
30.0 g of camphor with 1.25 L of ethanol, C2H5OH (d = 0.789
g/mL). Assume no change in volume when the solution is
prepared.
a. What is the mass percent of camphor in the solution?
b.
What is the molarity of the solution?
(2) What is the molecular formula of the acid?
(3) Suppose the NaOH buret was rinsed with distilled
water resulting in the first few drops of NaOH to be
more dilute, how would this affect the calculated
value for the equivalent mass of the acid?
(4) Suppose some of the solid acid was left in the
weighing cup, how would this affect the calculated
value for the equivalent mass of the acid?
d.

2.
7.
What is the spontaneous temperature range?
Consider the synthesis reaction: N2(g) + 3 F2(g)  2 NF3(g)
(Ho298 = -264 kJ mol-1, So298 = -278 J K-1 mol-1)
a. Calculate Go298 for the reaction.

Bismuth (Bi) reacts with fluorine to form BiF3.
a. Calculate the mass percent of Bi and F in the compound.
b.
For what temperature range is the reaction spontaneous?

b.
Calculate the mass of fluorine required to form 16.5 g of
compound.
c.
Write a balanced equation for the reaction.
d.
How many moles of F2 are required to react with 0.240
mol Bi?
c.
Calculate the heat released when 0.256 mol of NF3(g) is
formed from N2(g) and F2(g) at 1.00 atm and 298 K.

e.
How many grams of F2 are required to react with 1.60 g
Bi?
f.
If 5.00 g of Bi react with 2.00 g F2, what is the limiting
reactant?
g.
What is the theoretical yield of BiF3 when 5.00 g Bi and
2.00 g F2 react?
h.
When BiF3 reacts with water, one of the products is a
compound containing 85.65 % Bi, 6.56 % O, and 7.79 %
F. What is the simplest formula of this compound?
d.

3.
The combustion of carbon monoxide is represented by the
equation: CO(g) + ½O2(g)  CO2(g)
a. Determine Ho for the reaction above using the values.
C(s) + ½ O2(g)  CO(g)
Ho298 = -110.5 kJ•mol-1
C(s) + O2(g)  CO2(g)
Ho298 = -393.5 kJ•mol-1

b. Determine So for the reaction above using the table
CO(g)
CO2(g)
O2(g)
Substance
So (J/mol•K)
197.7
213.7
205.1

c.
Write a balanced equation for the reaction between BiF3
and H2O.
The dissolving of AgNO3(s) in water is represented by the
equation: AgNO3(s)  Ag+(aq) + NO3-(aq)
a. Is G positive, negative, or zero? Justify your answer.

b.
Unit 6 Practice Free Response
1.
Consider the combustion of butanoic acid at 25oC:
HC4H7CO2(l) + 5 O2(g)  4 CO2(g) + 4 H2O(l) Ho= -2,183.5 kJ
So (kJ/mol•K)
Substance
Hfo (kJ/mol)
CO2(g)
-393.5
0.2136
H2O(l)
-285.8
0.0699
5.
O2(g)
0.0
0.2050
C3H7COOH(l)
?
0.2263
a. Calculate Hfo, for butanoic acid.
b.
Calculate So for the combustion reaction at 25oC.

c.

Calculate Go for the combustion reaction at 25oC.
Determine Go for the above reaction at 298 K.

4.
i.
Calculate the F–F bond energy using the information
above and the bond energies
(NN = 946 kJ/mol, N–F = 272 kJ/mol).
The solution cools when AgNO3(s) is dissolved. Is H
positive, negative or zero? Justify your answer.

c.
Is S positive, negative, or zero? Justify your answer.

Consider the thermochemical equation:
C2H6(g) + 7/2 O2(g)  2 CO2(g) + 3 H2O(l) H = -1559.7 kJ
a. Calculate H for the thermochemical equations:
2 C2H6(g) + 7 O2(g)  4 CO2(g) + 6 H2O(l)
2 CO2(g) + 3 H2O(l) C2H6(g) + 7/2 O2(g) 
b. The heat of vaporization of H2O(l) is +44.0 kJ/mol.
Calculate H for the equation:
C2H6(g) + 7/2 O2(g)  2 CO2(g) + 3 H2O(g)
c.
The heat of formation of CO2(g) and H2O(l) are -393.5
kJ/mol and -285.8 kJ/mol. Calculate Hfo of C2H6(g).
c.
4.
d.
How much heat is evolved when 1.00 g of C2H6(g) is
burned to give CO2(g) + H2O(l) in an open container?
e.
What is the bomb constant C if the change in
temperature is 13.13oC when 1.00 g of C2H6 reacts in
the bomb calorimeter that contains 250 g H2O?
Unit 7 Practice Free Response
1.
Consider the reaction: 2 HI(g)  H2(g) + I2(g), where the rate
of the reaction in terms of [HI] = -0.100 M•min-1.
a. What is the rate of the reaction in terms of H2(g)?
b.
2.
What is the rate of the reaction in terms of I2(g)?
Consider the reaction: A + B2  Products. The following
experimental data at 22oC were obtained:
A (M)
B2 (M)
Rate (M•s-1)
0.100
0.100
0.080
0.500
0.100
0.40
0.100
0.500
2.0
a. What is the order of the reaction with respect to each
reactant?
5.
What is the half-life of the reaction?
The redox reaction between Tl+ by Ce4+ occurs by the
following mechanism.
Ce4+(aq) + Mn2+(aq)  Ce3+(aq) + Mn3+(aq)
Ce4+(aq) + Mn3+(aq)  Ce3+(aq) + Mn4+(aq)
Mn4+(aq) + Tl+(aq)  Mn2+(aq) + Tl3+(aq)
a. What is the balanced equation for the overall reaction?
b.
What molecule acts as a catalyst for this reaction?
c.
What molecules is an intermediate for this reaction?
d.
The rate law for this reaction is rate = k[Ce4+][Mn2+].
Which step is the slow step in the mechanism?
I-(aq) + CIO-(aq)  IO-(aq) + Cl-(aq)
Three initial-rate experiments were conducted; the results are
shown in the following table.
Experiment
[ClO-]
[I-]
[IO-]/t
(M)
(M)
(M•s-1)
1
0.017
0.015
0.156
2
0.052
0.015
0.476
3
0.016
0.061
0.596
a. Determine the order of the reaction with respect to each
reactant listed below. Show your work.
(1) I-
(2) CIO-
b.
What is the rate constant for the reaction, including
units?
b.
For the reaction,
(1) write the rate law for the reaction.
(2) calculate rate constant, k, and specify units.
c.
What would cause an increase in the rate constant?
c.
d.
The activation energy for the reaction is 115 kJ/mol.
What is the rate constant of the reaction at 27oC?
The rate constant for this reaction is determined at
various temperatures. The data is graphed in order to
determine the activation energy, Ea.
(1) What variables are graphed?
(2) Explain how to calculate the activation energy from
this graph.
3.
Consider the first-order decomposition of A. The rate
constant is 1.7 x 10-2 min-1.
a. Calculate the rate of the reaction when the initial
concentration of A is 0.400 M.
b.
What percent of A will be used up in two hours?