Answers to the
Honors Chemistry Final Exam Review Questions
Q1. Name the five major divisions of chemistry and give examples
of the types of subjects studied within each division.
A1. Inorganic Chemistry: geology, manufacturing; Organic Chemistry:
biology, medicine, agriculture; Analytical Chemistry: geology, manufacturing;
Physical Chemistry: geology, physics, manufacturing; Biochemistry: biology,
agriculture, medicine.
Q2. Identify five examples of matter and five examples of
non-matter.
A2. The examples should follow this rule: Matter has mass and volume;
non-matter has no mass or volume.
Examples of matter are: air, water, metal, rock, solutions. Examples of non-matter are: energy, vacuum,
force, magnetic or electric or gravitational fields.
Q3. Explain the difference between a qualitative and a
quantitative measurement. Provide
examples to illustrate this difference.
A3. A qualitative measurement is a measurement that gives descriptive,
non-numeric results; a quantitative measurement is a measurement that gives
definite, usually numeric results.
"The rock is heavy" would be a qualitative measurement. "The rock weighs 110 grams" would
be a quantitative measurement.
Q4. Explain what is meant by dimensional analysis. Use an example in your explanation.
A4. In the dimensional analysis approach to problem solving, the units of
the unknown are identified. Then the
relationship between the knowns is established and an appropriate conversion
factor or series of conversion factors is selected. The conversion factor or series of factors is appropriate if,
when it is applied to the knowns, the units cancel so as to yield the units of
the unknown. If the units do not cancel
in this way, it can be inferred that something has gone wrong during the
problem solving process.
For instance,
suppose the problem is to determine an airplane's speed in km/h when you know
that the plane travels 2000 m every 10 seconds. The units of the unknown are km/h. The knowns are 2000 m and 10
seconds, and they are related by the fact that the plane travels 2000 m in each
10 seconds, i.e. the plane's speed is 2000 m/10 s. According to dimensional
analysis, it is necessary to go from the units of the known (m/s) to the units
of the unknown (km/h). The appropriate conversion factors to do
this are 1 km/1000 m and 60 s/h. Units of knowns x conversion factor
x conversion factor = units of unknown m/s x km/m x s/h = km/h By applying
these conversion factors to the related knowns the answer is obtained in the
correct units.
Q5. Explain how the atoms of one element differ from
those of another element.
A5. Differences among the atoms of different elements result from different
numbers of protons in the nuclei of atoms.
All atoms of the same element have the same number of protons. Since atoms are neutral, the number of
electrons in an atom equals the number of protons. The atoms of an element may have different numbers of neutrons in
their nuclei and still be atoms of the same element.
Q6. Explain the terms molecular formula and formula
unit. Give an example of each.
A6. A molecular formula shows the kinds and numbers of atoms present in a
molecule of a compound. CO2
is an example of a molecular formula. A
formula unit is the lowest whole-number ratio of ions in an ionic
compound. BaCl2 is an
example of a formula unit. The ratio of
barium to chloride ions is 1:2. BaCl2
is not a molecule.
Q7. What is the advantage of using the specific term,
gram molecular mass, instead of the general term, gram formula mass?
A7. There may be some confusion when the terms are applied to diatomic
gases such as nitrogen and oxygen. For
instance, the gram molecular mass of nitrogen is 28 g. The gram formula mass
could be either 14 g or 28 g, depending upon whether the subject is nitrogen
atoms or nitrogen molecules.
Q8. What determines whether one metal will replace
another metal from a compound in a single-replacement reaction?
A8. Whether one metal will replace another is determined by the relative
reactivity of the two metals. The
activity series of metals lists metals in order of decreasing reactivity. A reactive metal will replace any metal found
below it in the activity series.
Q9. What is the importance of the coefficients in a
balanced chemical reaction?
A9. The coefficients in a balanced chemical equation indicate the relative
number of moles of reactants and products.
From this information the amounts of reactants and products can be
calculated. The number of moles may be
converted to mass, volume, or number of representative particles.
Q10. Name the three basic assumptions of the kinetic
theory.
A10. The kinetic theory is based upon the assumptions that a gas is
composed of particles, that these particles are in constant random motion, and
that all collisions between particles are elastic.
Q11. Explain the difference between temperature and heat.
Also, state what determines the direction of heat transfer.
A11. Temperature is a measure of the average kinetic energy of the
particles composing an object. Heat is
the energy that is transferred between two objects, of different temperature,
that are in contact with each other.
Temperature determines the direction of heat transfer. Heat always flows from the object of higher
temperature to the object of lower temperature.
Q12. What are some of the differences between a real gas
and an ideal gas?
A12. An ideal gas is one that follows the gas laws at all conditions of
pressure and temperature. The behavior
of a real gas deviates from the behavior of an ideal gas at various
temperatures and pressures. Also,
kinetic theory assumes that the particles of an ideal gas have no volume and
are not attracted to each other. This is not true for real gases. Real gases can be liquefied and sometimes
solidified by cooling and applying pressure, while ideal gases cannot.
Q13. Describe the shapes and relative energies of the s,
p, d, and f atomic orbitals.
A13. An "s" orbital has the shape of a sphere and is the orbital
having the lowest energy. A
"p" orbital is dumbbell-shaped and has the next highest energy. A "d" orbital has a more complex
shape and a higher energy than either an "s" orbital or a
"p" orbital. An "f"
orbital has the highest energy of these four orbital types; this orbital has a
very complex shape.
Q14. Describe the periodic trends in atomic radii that
can be observed in the periodic table.
Provide examples.
A14. Atomic radii increase with increasing atomic number down a column or
family--sodium is bigger than lithium; potassium is bigger than sodium;
rubidium is bigger than potassium, for example. Atomic radii decrease with increasing atomic number across a
period--lithium is bigger than beryllium; beryllium is bigger than boron; and
boron is bigger than carbon, etc., for example.
Q15. Explain the octet rule and give an example of how it
is used.
A15. The electron configuration (filled s and p orbitals s2p6)
of the noble gases is extremely stable.
In this configuration, repulsion between electrons is minimized and the
energy state is, therefore, relatively low.
The octet rule states that, in chemical reactions, elements gain or lose
electrons to achieve the stable electron configuration of a noble gas. This stable configuration is called an octet
because it includes a total of 8 valence electrons (s2p6):
2 from the outermost s orbital and 6 from the outermost p orbital. An example of use of the octet rule is the
following. Oxygen has the electron configuration 1s22s22p4. When oxygen reacts to form ionic compounds,
it completes its octet (2s22p4) by gaining two electrons
from the element it reacts with. These
two elements add to the p orbital of oxygen, giving it the electron
configuration (1s22s22p6) of the noble gas,
neon. Oxygen's incomplete octet is thus
completed by the gain of two additional electrons to give the stable octet of
the noble gas (2s22p6).
Q16. Explain what an unshared pair of electrons is. Give an example.
A16. An unshared pair of electrons is two valence electrons that are not
shared between atoms. Each fluorine
atom in a molecule of fluorine has three unshared electron pairs, for example.
Q17. Describe the structure of the water molecule and
indicate how this structure is responsible for many of the unique properties of
this vital compound.
A17. Water is a simple, triatomic molecule. Each OH covalent bond in the
water molecule is highly polar. Because
of its greater electronegativity, the oxygen atom attracts the electron pair of
the covalent OH bond and acquires a slightly negative charge. The hydrogen
atoms, being less electronegative than the oxygen, acquire a slightly positive
charge. The atoms of the water molecule are joined in a 105º angle. As a
result, the slight charges on the individual atoms do not cancel each other out
and the molecule itself is polar. There is a slight negative charge in the
region around the oxygen and a slight positive charge in the region around the
hydrogens. Because water molecules are polar, they attract each other. The
hydrogen of one molecule is attracted to the oxygen of another molecule. This
attraction is termed hydrogen-bonding and it is stronger than other polar
attractions because of the fact that the hydrogen nucleus is not shielded by an
electron cloud in the way that other nuclei would be (hydrogen atoms have only
1 electron). It is the strong intermolecular attraction associated with hydrogen-bonding
that is responsible for many of the unusual properties of water, including its
high surface tension, low vapor pressure, high specific heat, high heat of
vaporization, and high boiling point.
Q18. Explain what a saturated solution is. Give a specific example.
A18. A saturated solution contains the maximum amount of solute for a given
amount of solvent at a constant temperature. For example, no more than 36.2 g
of sodium chloride will go into solution in 100 g of water. Above this
concentration, there is a dynamic equilibrium between the solid and its
dissolved ions. In this equilibrium,
just as many ions are going out of the solution as are going in per unit time,
and solid will, therefore, always be present.
Q19. Explain the effects of reactant concentration and
particle size on the rate of a reaction.
A19. A small particle size increases the rate of a reaction because there
is more surface area for a given mass of particles and so more collisions are
possible per second. A high concentration of reactants increases the reaction
rate because more molecules are present to collide each second.
Q20. Compare and contrast the properties of acids and
bases.
A20. Both acids and bases are electrolytes, they cause indicators to change
colors, and they react with each other to form water and a salt. Acids taste
sour, while bases taste bitter. Bases feel slippery. Acids react with some
metals to produce hydrogen gas.
Q21. What happens in a neutralization reaction? Use an example.
A21. An acid reacts with a base to produce a salt and water. When
hydrochloric acid reacts with sodium hydroxide, the hydrogen and hydroxide ions
form water and the sodium and chloride ions form sodium chloride.
Q22. What are the steps of the oxidation number change
method of balancing REDOX reactions?
A22. Assign oxidation numbers to all the atoms. Identify the oxidized and reduced atoms. Connect the oxidized and
reduced forms of the atoms with a line.
Use coefficients to make the total change in oxidation number equal for
both the oxidation and die reduction. Check the balance of atoms and charges.
Q23. What are the steps of the half-reaction method of
balancing REDOX reactions?
A23. Write out the equation in ionic form. Write the two half-reactions.
Balance the atoms in each half-reaction. Balance the charges in each
half-reaction. Multiply each
half-reaction by a factor that will make the charges equal in both. Add the
half-reactions, leaving out terms that appear on both sides.
Q24. Explain what occurs in an electrochemical cell.
A24. Electrical energy is converted into chemical energy or vice
versa. Electrons are transferred from
one atom to another, for instance from zinc to copper. One substance changes from oxidized to
reduced form, and the other substance does the opposite. The substance losing electrons more easily
is the one that is oxidized.
Q25. Explain why carbon is able to form such a wide
variety of compounds.
A25. Carbon has four valence electrons and can therefore form four covalent
bonds. Carbon can bind other atoms and it can also bind to itself.
Carbon-carbon bonds are quite stable, and this fact, coupled with its capacity
to form four covalent bonds, enables carbon to form long chains that may be
either straight or branched; carbon atoms can also be joined to each other in
ring structures. Carbon can also form double and triple bonds with itself. This
allows for even greater variety in the types of carbon compounds that can be
formed.
Q26. Given the name of an alkane, indicate how you can reconstruct
its structural formula according to IUPAC rules.
A26. (1) Find the root word (ending in -ane) in the hydrocarbon name.
Then write the longest
carbon chain to create the parent
structure.
(2) Number
the carbons of this parent carbon chain.
(3) Identify the substituent
groups. Attach the substituents to
the
numbered parent chain at
the proper positions.
(4) Add hydrogens as needed.
Q27. Explain how geometric isomers differ from each
other. Describe the difference between
the trans and cis configurations of geometric isomers. Provide an example of each configuration for
molecules that exhibit geometric isomerism.
A27. Geometric isomers differ only in the geometry of their substituted
groups. Geometric isomerism is possible
whenever each carbon of a double bond has at least one substituent. In the trans configuration, the substituted
groups are on opposite sides of the double bond. In the cis configuration, the substituted groups are on the same
side of the double bond. Examples are
trans-2-butene and cis-2-butene.
Q28. Explain the concept of resonance. Give an example of a compound that displays
resonance.
A28. Resonance occurs when two or more equally valid structures can be
drawn for a molecule. The benzene molecule exhibits resonance. In benzene, each carbon atom can participate
in a double bond. A double bond may be
formed on one side of the carbon and a single bond may be formed on the other
side. However, the reverse structure can also be formed. Benzene is said to resonate between these
equally possible structures.
Q29. Describe the process of making sucrose. Give an example.
A29. One glucose molecule and one fructose molecule are joined through an
ether linkage. A water molecule is
removed.
Q30. Give an example of a substitution reaction and
describe what happens in the reaction.
A30. CH4 + Br2 --> CH3Br + HBr
The halogen replaces the
hydrogen because it is more reactive.
Q31. Distinguish between alpha particles, beta particles,
and gamma rays. Indicate how the atomic
number and atomic mass number change when each type of radiation is emitted.
A31. Alpha particles are helium nuclei; they have 2 protons and 2
neutrons. Alpha particles have low
energy. Beta particles are electrons or
positrons. They have medium
energy. Gamma rays are electromagnetic
radiation. They have high energy. Alpha particle emission reduces atomic
number by 2 and atomic mass number by 4. Beta particle emission increases
atomic number by 1 and does not affect atomic mass number. Gamma ray emission does not affect atomic
mass number or atomic number.