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
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
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.
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
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
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
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
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.