AP Semester Exam Objectives & Study Hints
Chapter 1 - Matter and Measurement
Objectives:
1. To distinguish between substances and mixtures, chemical and physical properties, and intensive and extensive properties.
2. To convert lengths, masses, and liquid volumes from one metric (SI) unit to another and in some cases convert from metric to non-metric units. Dimensional Analysis.
3. To understand the relationship between mass, volume and density.
4. To convert temperatures from one to another using three scales, Fahrenheit, Celsius and Kelvin.
5. To apply the rules for significant figures.
6. To distinguish between precision and accuracy.
Study Hints:
1. In the textbook, the problems are worked with a calculator and the answer is rounded to the correct number of significant figures only at the end of the problem. Rounding off intermediate values may produce slightly different results from those in the text.
2. Don’t underestimate the importance of density. It is a simple concept, but one that will be important for many problems later in the course, so be sure to remember it.
3. Most of the unit conversions you will do in this course will involve only metric units. Remember the prefix indicates the magnitude of the conversion factor, and dimensional analysis will help you decide whether to divide or multiply by the conversion factor.
4. When working with percentage, the first step is usually to convert the percent into a fraction (or decimal). Don’t forget the need for this conversion.
Chapter 2 - Atoms and Elements
Objectives:
1. To be familiar with the basic assumptions of Dalton's Atomic Theory.
2. To state the Law of Conservation of Matter.
3. To know the relationships among atomic number, mass number, number of protons, number of neutrons, and number of electrons.
4. To calculate the average atomic mass of an element from the relative abundances of the component isotopes.
5. To use the periodic table to locate the alkali metals, the alkaline earth metals, the transition metals, the halogens, and the rare gases, and to distinguish among metals, nonmetals, and metalloids.
6. To convert among grams, moles, and number of particles.
Study Hints:
1. It is imperative that you learn how to convert moles to grams and grams to moles. This simple conversion is a basic step in a great many problems that will be encountered throughout chemistry.
2. Remember that a single chemical symbol can have several different meanings. For instance, Fe can mean a) one atom of iron, b) one mole of iron atoms, or c) one molar mass of iron. Since all of these interpretations are possible, you must decide which one is most appropriate in a given case.
3. The terms molar mass, gram molecular mass, gram molecular weight, and gram formula mass are different names for the same quantity.
Chapter 2Part 2 - Molecules, Ions, & Their compounds
Objectives:
1. To know the names (including spelling) and symbols of the common elements.
2. To determine the charge for monatomic ions and predict formulas of ionic compounds.
3. To know the names, formulas, and charges of the common polyatomic ions.
4. To name ionic compounds and binary molecular compounds.
5. To distinguish between a molecular formula and an empirical formula.
6. To calculate the molar mass of a compound.
7. To state the law of constant composition.
8. To calculate the percent composition of a compound from the empirical formula and vice versa, and the molecular formula from the empirical formula.
Study Hints:
1. Pay attention to the spelling of the names of elements that frequently cause problems (commonly confused letters are underlined): beryllium, fluorine, silicon, phosphorus, sulfur, chlorine, chromium, nickel, and zinc.
2. Don't confuse the symbols of the elements having names that begin with S (sodium, sulfur,silicon, and scandium) or those that begin with P (potassium and phosphorus).
3. Remember that under normal conditions pure hydrogen, nitrogen, oxygen, & halogens exist as diatomic molecules, whereas rare gas elements exist as uncombined atoms.
4. When working problems related to chemical formulas, always remember that a formula is basically a mole ratio. Therefore, the initial stage of such a problem often involves the determination of the number of moles of the components in the formula.
5. Like chemical symbols, chemical formulas can have several different meanings. For instance, H2O can mean a) one molecule of water, b) two atoms of the element hydrogen combined with one atom of oxygen, c) one mole of water molecules, or d) one molar mass of water molecules.
Chapter 3 - Chemical Equations & Stoichiometry
Objectives:
1. To balance simple chemical equations.
2. To predict the products and balance the resulting equation for some common reactions, such as the combination of a metal with oxygen or a halogen, the combustion of a hydrocarbon, and the decomposition of a metal carbonate due to heating.
3. To use a balanced chemical equation to write the stoichiometric factors for a chemical reaction and use these values to calculate the relationships between the moles or mass of products and reactants.
4. To recognize the limiting reagent in a chemical reaction and use it to perform stoichiometric calculations.
5. To calculate percent yield using actual yield and theoretical yield.
Study Hints:
1. Whenever you begin to work on a stoichiometry problem, your first step should be to make sure that any equations provided are complete and balanced. In problems like this, an equation that hasn't been balanced is a mistake waiting to happen.
2. It is very important to realize that the mole ratio in the balanced equation is the key to doing most of the problems in this chapter. Most of these problems can be thought of in three steps: first, convert whatever you are given into moles; second, use the stoichiometric ratio to convert to moles of another substance in the balanced equation; and third, convert from moles of that new substance to whatever units are requested.
3. Notice that the limiting reagent is not necessarily the reactant present in the smallest amount; you must also consider the stoichiometric factor. The reagent present in the largest amount may be the limiting reagent if there is not enough of it to satisfy the requirements of the balanced equation.
Chapter 4 - Reactions in Aqueous Solution
Objectives:
1. To describe solutions using the terms solute, solvent and solution.
2. To distinguish among strong and weak electrolytes.
3. To distinguish among acids and bases.
4. To name and write formulas for the common acids and bases.
5. To predict the solubility of ionic compounds by using general solubility guidelines.
6. To determine which ions will form when an ionic compound is dissolved in water and to write net ionic equations for chemical reactions involving ionic compounds.
7. To identify the important types of reactions in aqueous solution, including exchange reactions, precipitation reactions, acid-base reactions, and gas-forming reactions and to predict the products in many of these reactions.
8. To write equations for the preparation of compounds using exchange reactions, precipitation reactions, or acid-base reactions.
9. To use the definition of molarity to do solution calculations, including those required to prepare a solution of a given concentration either by dissolving the solute directly in a solvent or by dilution of a previously known solution.
10. To use a balanced equation and the molarity of the reactants, to perform solution stoichiometry calculations for titration and precipitation reactions.
11. To determine the oxidation numbers of the elements in compounds, identify the oxidized and reduced substances in balanced equations, and to balance oxidation-reduction equations in either acidic or basic solutions.
12. To predict products and write net ionic equations for redox reactions.
Study Hints:
1. Don't confuse the weak electrolyte ammonia, NH3, with the ammonium ion, NH4+.
2. Objective 7 suggests that you should be able to predict the products of acid-base exchange reactions. It is also helpful to be able to work backwards, predicting what acid and base would have to be reacted in order to form a given salt.
3. When you learn the definition of molarity, be sure to learn moles of solute per liter of solution, not just moles per liter. Remember that when preparing a solution, liters of solvent added is not always the same as the final volume of the solution. If you learn the complete definition for molarity, you are less likely to confuse it with molality.
4. Dilution problems should be fairly easy once you memorize the formula, but there is one frequent point of confusion. You must read carefully to distinguish between the volume of water added and the volume of the final solution.
5. Remember that the oxidized substance is the reducing agent, and the reduced substance is the oxidizing agent. Also remember that the oxidizing agent and reducing agent in a chemical equation must both be reactants.
6. When balancing oxidation-reduction equations, it is essential that you follow a systematic procedure. Some common errors are forgetting to balance electrons in both half reactions, orforgetting to simplify the final equation. The charges on the ions are very important, so be sure to write the charges clearly and neatly. Otherwise it's very easy to overlook an ion in adding up the charges and the balancing will be incorrect.
Chapter 5 - Energy & Chemical reactions
Objectives:
1. To recognize the various forms of energy and understand how the transformations from one form to another are governed by the conservation of energy principle.
2. To convert energy values from calories to joules and vice versa.
3. To make calculations involving changes in temperature and/or state.
4. To understand the first law of thermodynamics, and terms related to this concept, including endothermic and exothermic reactions, energy change, and enthalpy change.
5. To use Hess's Law to calculate the enthalpy change for a reaction that is a simple combination of the given equations.
6. To understand the meaning of the term standard conditions, be able to predict the standard states for common substances, and be able to calculate the standard enthalpy change for reactions based on values from an appropriate table of standard enthalpies of formation.
7. To describe the basic experimental procedures of calorimetry, the method used to measure heats of reactions in the laboratory.
Study Hints:
1. Thermodynamic symbols often include subscripts and superscripts that represent critical information about the conditions, such as temperature and pressure. For example, the superscript ° on the symbol < H° should tell you the temperature, the pressure, and the physical state of the substance involved. Don't overlook this valuable source of information.
2. Be sure that all of the equations are balanced, before you begin to do a thermochemistry problem. Failure to do this can waste time (if you notice the error later on) or cause your solution to be incorrect.
Chapter 19: The Spontaneity of Chemical Reactions
Objectives:
1. Understand the concept of entropy and how it relates to spontaneity.
2. Predict whether a process is product- or reactant-favored.
3. Use tables of data in thermodynamic calculations.
4. Define and use a new thermodynamic function, free energy.
Chapter 20: Electrochemistry: The Chemistry of Oxidation-Reduction Reactions
Objectives:
1. To explain how an oxidation-reduction reaction in a voltaic cell can be used to produce an electric current and to recognize the various components of a cell such as anode, cathode and salt bridge.
2. To use a table of standard reduction potentials to predict whether or not a specific combination of half-reactions will occur spontaneously under standard conditions.
3. To use the Nernst equation to calculate the potential of an electrochemical cell when conditions are nonstandard.
4. To use the Nernst equation to determine equilibrium constants from standard reduction potential values.
5. To determine the relationship between current flow and the amount of chemical reaction that can occur.
6. To explain the conditions that are most likely to produce corrosion as well as how corrosion can be prevented.
Study Hints:
1. Remember that in the Nernst equation, the value of n is determined by the number of electrons transferred when balancing the net equation. Normally it isn’t possible to simply look at one half-reaction and determine what the n value will be for a net equation. Be sure to balance the two half-reactions in the usual way, and then determined n from the number of electrons that are canceled out when the two half-reactions are added.
2. In many cases, chemists have agreed to always do certain things the same way when writing electrochemistry problems. For example, the anode is normally written on the left on an electrochemical cell, and the table of standard electrode potentials is normally written with the most negative potentials on the top of the list. Don’t depend too much on these conventions. It can be a rude shock for those who memorize that the strongest oxidizing agents are at the top of the reduction potential table, then encounter a table that lists the half reactions in the opposite order. It’s always better to try to understand rather than just memorize isolated facts, but that is especially true in electrochemistry.
3. When using the table of standard reduction potentials, don’t forget that the half-reactions listed include both an oxidizing agent and a reducing agent. When asked to identify the oxidizing agent in a process, don’t respond by giving the entire half-reaction.
Chapter 23 - Nuclear Chemistry
Objectives:
1. To write nuclear reactions for nuclear fission, nuclear fusion, spontaneous radioactive decay, and artificial transmutations.
2. To describe the factors that are important in determining nuclear stability and the importance of binding energy in the determination of nuclear stability.
3. To do calculations involving the relationship between elapsed time and the amount of radioactive material which has been transformed.
4. To explain the basic principles that govern nuclear fission reactors as well as the assets and liabilities that these technologies present.
5. To describe some of the commercial applications of radiation in medicine, chemical research, and chemical analysis.