AP Chemistry Learning Targets

Chapter 1, Measurement and Basics

1. Understand that matter is composed of atoms of various types called elements. An element is a category not a physical substance.
2. Understand that pure substances may either be elements, which are composed of just one kind of atom, or compounds, which are composed of two or more kinds of atoms where there is a specific proportion of the numbers of each kind.
3. Understand that atoms may be joined together in molecules, in which chemical bonds hold them together to make a single particle.
4. Understand that mixtures (physical combinations of two or more pure substances), unlike compounds, have a variable composition in which the component pure substances may vary in their proportions.
5. Understand that heterogeneous mixtures are not uniform in the distribution of the components of the mixture. And that homogeneous mixtures, or solutions, are uniform.
6. Understand the phases of matter (solid, liquid, and gas) based on the behavior of the particles of matter which compose each phase.
7. Understand the difference between a physical change, such as a phase change, and a chemical change, in which substances are transformed into chemically different substances.
8. Understand that intensive properties are based on ratios (such as the amount of mass per unit volume, or density) and so do not depend on the amount of material present. And that extensive properties are specifically measures of the amount of material.
9. Be able to convert between units in the metric system based on the prefixes G, M, k, c, m, µ, n, p, and f.
10. Be able to convert between units of temperature (°C, °F, and K).
11. Be able to show work for conversions between units, or other calculations, using the method of dimensional analysis in which units are explicitly shown to cancel or combine as required to produce the correct unit for the result of the calculation.
12. Understand that accuracy is a measure of how true a measurement is. Understand the precision is a measure of the degree to which multiple measurements of the same quantity agree with one another.
13. Understand that significant figures are a way to express the precision of a measurement. The more digits in a measurement, the more precise it is implied to be. This is because the smaller the final digit is, the closer multiple measurements or the same quantity can be to one another. When making measurements the number of significant figures is determined by the digits of which one can be certain plus one estimated digit.
14. Be able to identify the number of significant figures based just on the written measurement; be able to determine which zeros in a measurement are significant and which are not.
15. Be able to perform simple calculations in such a way as to preserve information about the precision of measurements by following the given rules to produce a calculated result with the correct number of significant figures.

Chapter 2, Part I, Atomic Structure

1. Understand and apply concepts of atomic structure such as atomic number, mass number, isotopes, and ionic charge.
2. Be able to calculate average atomic mass, or the mass or a specific isotope, given appropriate data.
3. Handle and make calculations with number used to describe atomic properties such as size on ångstroms (Å) or charge in coulombs (C).
4. Be conversant with important milestones in the history of the discoveries involved in establishing atomic structure. For example, the Thomson cathode ray tube, Millikan oil drop, and Rutherford scattering experiments.
5. Be able to name groups in the periodic table (metals vs. non-metals and groups such as the alkali metals and the halogens) and also be able to discuss the properties of elements based on group membership.

Chapter 2, Part II, Nomenclature

1. Be able to write formulas and names for molecular compounds, binary ionic compounds, ionic compounds involving polyatomic ions, and acids.
2. Know how and when to use roman numerals in cation names to indicate charge.
3. Memorize common polyatomic ion names and formulas, including oxyanions (-ite, -ate, etc.).
4. Be able to write formulas and names for simple organic compounds.

Chapter 3, Stoichiometry

1. Be able to write chemical equations using correct chemical formulas given names. Also, be able to balance a chemical equation.
2. Know how to recognize synthesis, decomposition, and combustion reactions. Be able to predict the product of a synthesis reaction of a metal with a non-metal. Be able to write an equation for a combustion reaction given the formula of a hydrocarbon (X + O₂ → CO₂ + H₂O).
3. Be able to calculate molar mass given chemical formulas. Also, be able to relate mass in grams to mass in amu using Avogadro’s number: 6.02 × 10²³ atoms = _____ g = 1 mol where the mass in grams of 1 mol is numerically the same as the mass in amu.
4. Be able to find the empirical formula of a compound given percent composition or lab data which can give percent composition. Also be able to find the empirical formula based on data from a combustion analysis.
5. Calculate the molecular formula of a compound given its empirical formula and molar mass.
6. Use data to identify a limiting reactant in a chemical reaction. Know that a limiting reactant is used up entirely but other reactants are not; be able to calculate amount remaining for excess reactants. Know that the limiting reactant determines how much of every product forms.
7. Use data to calculate theoretical yields and understand how to find percent yield given an actual yield or how to find an actual yield given a percent yield.

Chapter 4, Reactions in Aqueous Solution

1. Be able to define, discuss, and apply an understanding of aqueous solutions including solvent, solute, strong electrolyte, non-electrolyte, and weak electrolytes.
2. Be able to recognize, based on names and chemical formulas, which substances are strong electrolytes, non-electrolytes, and weak electrolytes.
3. Memorize the solubility guidelines so that you can correctly predict which product is insoluble in a precipitation reaction.
   

   Always Soluble	Soluble with Exceptions		(almost) Always Insoluble
   NH4+1, Li+1, Na+1, K+1, Rb+1, Cs+1 NO3–1, CH3COO–1	Cl–1, Br–1, I–1	SO4–2	F–1, S–2, CO3–2, OH–1, PO4–3
   	exceptions: Ag+1, Pb+2, Hg2+2	exceptions: Sr+2, Ba+2, Pb+2, Hg2+2
4. Be able to write molecular, complete ionic, and net ionic chemical equations for a given reaction. Molecular chemical equations are the familiar chemical equations which simply relate the complete chemical formulas of the reactants to the products. Complete ionic equations break apart soluble electrolyte compounds into their component ions as they would exist in aqueous solution. Net ionic equations leave out ions which remain unchanged in solution from the reactants to the products.
   For example,
   

   Molecular	Cu(s) + 2AgNO3(aq) → Cu(NO3)2(aq) + 2Ag(s)
   Complete Ionic	Cu(s) + 2Ag+ + 2NO3–(aq) → Cu2+(aq) + 2NO3–(aq) + 2Ag(s)
   Net Ionic	Cu(s) + 2Ag+ → Cu2+(aq) + 2Ag(s)

5. Know the definitions of acids, bases, and neutralization reactions and be able to write chemical equations. Memorize the strong acids (HCl, HBr, HI, HNO₃, H₂SO₄, HClO₃, and HClO₄) and bases (MOH, where M is a metal ion) and be able to recognize weak acids (includes COOH structure) and bases (includes nitrogen atoms). Be able to predict products of reactions between acids and bases which produce gases.
6. Memorize how to assign oxidation numbers to elements based on chemical formulas. Be able to write equations for oxidation-reduction reactions (single displacement). Be able to track electron exchange in such reactions and, given an activity series, predict products of reactions.
   • Pure elements have an oxidation number of 0.
   • Monatomic ions have oxidation numbers equal to their charge.
   • Nonmetals tend to have negative oxidation numbers, although some are positive in certain compounds or ions.
   • In compounds oxygen has an oxidation number of -2, except in the peroxide ion (O₂^2–) in which it has an oxidation number of -1.
   • Hydrogen is -1 in compounds with metals, +1 in compounds with non-metals.
   • In compounds nonmetals usually have negative oxidation numbers. Central atoms in oxyanions, among others, have positive oxidation numbers
   • Fluorine always has an oxidation number of -1 in compounds.
   • The other halogens have an oxidation number of -1 when they are monatomic ions or they are not the central atom in a molecule; in oxyanions they have positive oxidation numbers.
   • The sum of the oxidation numbers in a neutral compound is 0.
   • The sum of the oxidation numbers in a polyatomic ion is the charge on the ion.
7. Be able to calculate concentration of solutions in mol/L (M) and to calculate concentrations in dilutions (M₁V₁ = M₂V ₂).
8. Be able to calculate stoichiometric amounts of materials and corresponding volumes of solutions for any reaction involving aqueous solutions but especially for titrations.

Chapter 10, Gases

1. Know the characteristics of gases and vapors (compressibility, mixing) and the composition of the atmosphere.
2. Know the following units of pressure and how to convert between them: atm, mm Hg (millimeters of mercury), torr, in Hg (inches of mercury), Pa, kPa, bar, and mbar. 1 atm = 760 torr = 760 mm Hg = 101 kPa = 1.0 bar.
3. Be able to describe the physical state of gases using volume (V), pressure (P), number of moles (n), and absolute temperature (T, measured in kelvin, K).
4. Be able to describe barometers and manometers (mercury-filled tubes) and how they work.
5. Be able to use the ideal gas equation (PV = nRT) to calculate gas variables and know Boyle’s Law (PV = k), Charles’s Law (V = kT), and Avogadro’s Law (V = kn).
6. Know that at 1 atm (P) and 273 K (T) (or 0°C) 1 mol of an ideal gas has a volume of 22.41 L. These conditions are known as standard temperature and pressure (STP). Know the value of R, the universal gas constant: R = PV/nT = (1 atm)(22.41 L)/(1 mol)(273 K) = 0.08206 L·atm/K·mol.
7. Be able to use the ideal gas equation to relate the molar mass of a gas to its density. The molar mass of a gas is directly proportional to its density.
   

       m   n·M    P·M           n     P
   D = — = ——— = ———— (because ——— = ————)
       V    V     R·T           V     RT
   (M represents molar mass in g/mol, 
   V is measured in liters, L)   
   

       P·M          DRT
   D = ———  or  M = ————
       R·T           P
   

8. Be able to do stoichiometry with gases by volume. This works because at constant pressure and temperature volume is directly proportional to number of moles. Be able to do stoichiometry with gases by partial pressures. This works because at constant volume and temperature partial pressure is directly proportional to number of moles. (n = PV/RT)
9. Understand that the sum of the partial pressures of a mixture of gases gives the total pressure: P_total = P₁ + P₂ … P_n. The partial pressure of a gas (n) is also given by its mole fraction [(mol n)/(total mol) = χ] times the total pressure: P_n = χ_nP_total. Be able to do relevant calculations.
10. Understand the kinetic-molecular theory (KMT) of gases, which explains the behavior of gases based on the constant, chaotic motion of molecules which have negligible volume and which neither attract nor repel each other.
11. Understand the direct proportion between the absolute temperature of a gas and the average kinetic energy of its molecules (KE_avg = ³/₂RT, where R must have units of J/K·mol, R = 8.314 J/K·mol).
12. Know that the molecules of a gas do not all travel at the same speed at a given temperature but that they have a distribution of possible speeds. At higher temperatures the average (most probable) speed is faster and there are more possible different speeds.
    
13. Know that at the same temperature gases with different molecular weights will have different distributions of speeds. Molecules with smaller molar masses with have faster average (or most probable) speeds and will have more possible different speeds.
    
14. Know that average speeds for molecules are inversely proportional to the square root of the molar mass of the gas.
    u_rms = √(3RT/M), where M = molar mass in kg/mol.
15. Know that effusion is the escape of gas particles through a small hole. Rate of effusion is inversely proportional to molar mass because larger molecules have slower average speeds. Know that diffusion is the mixing of gas particles. The rate at which a gas diffuses through another gas is inversely proportional to its molar mass. This is because larger molecules have slower average speeds.

Chapter 5, Thermochemistry

1. Understand the concept of energy as the capacity to do work or to transfer heat. Energy may take on many forms including gravitational potential energy (E_P = mgh with mass in kg and g = 9.8 m/s² and h is height in meters) and kinetic energy (E_K = 1/2mv² with mass in kg and v is velocity in m/s) and heat (see calorimetry).
2. Know and be able to convert between units of energy 1 cal = 4.184 J, 1 kcal = 1 Calorie, where a capital “C” Calorie is a dietary calorie.
3. Know the definitions of surroundings, open systems (matter and energy can be exchanged with surroundings), closed systems (only energy can be exchanged with the surroundings), and isolated systems (nothing can be exchanged with surroundings).
4. Know that work is energy expended against a resisting force over some distance. And that heat is the energy which is transferred from a hotter object to a colder one.
5. Understand the First Law of Thermodynamics as the idea that energy is neither created nor destroyed and that the change in internal energy (ΔE) is equal to the sum of heat added or taken from a system (q) and the work done on or by the system (w). ΔE = q + w.
6. When heat is transferred out of a system (q < 0) and the process is exothermic. When heat is added into a system (q > 0) and the process is endothermic. Work done on a system: w > 0; Work done by a system: w < 0. Exothermic and endothermic are terms that do not apply to work.
7. Know that state functions (E, H, P, V, and T) depend only on the current conditions of a system but non-state functions (q and w) depend on the specific circumstances of how a change in a state took place.
8. Enthalpy (H) is a state function that is carefully designed (never mind how) to equal the heat absorbed or given off in physical and chemical changes. Change in enthalpy (ΔH) is equal to the heat involved in a process at constant pressure.
9. Though absolute enthalpy (H) cannot be measured, the change in enthalpy (ΔH) is a useful quantity, which can be thought of in this way: ΔH = H_final – H_initial.
10. Know that ΔH is proportional to the amount of material in a reaction and reversing a reaction reverses the sign (+ or –) of ΔH. Note: the state of matter is relevant to ΔH because phase changes may be endothermic (+ΔH) or exothermic (–ΔH).
11. Know that calorimetry is the measure of heat absorbed or given off by chemical and physical changes. Heat capacity is the amount of energy needed to raise the temperature of a material by 1°C (or 1 K) per gram (or per mole) of the material. Units for specific heat capacity: C_s = J/g·K. Units for molar heat capacity: C_m = J/mol·K.
12. Know that calorimetry involves measuring changes in temperature (ΔT = T_f – T_i) and that heat is calculated by multiplying mass times specific heat capacity times the change in temperature: q = mcΔT.
13. Know that heat lost by one object in calorimetry equals the heat gained by another so that the sum of heat gained and lost is zero: q_lost + q_gain = 0.
14. Know Hess’s Law and be able to apply it in order to deduce the ΔH for a target reaction using reactions with known values of ΔH. Hess’s Law concerns the fact that ΔH is a state function: the sum of ΔH values for a series of steps in a process is the same no matter what series of steps actually occur. For example:

    A → B ΔH = +40 kJ so A → C ΔH = +40 + (+30) = +70 kJ
    C → B ΔH = –30 kJ    because:
                        A → B ΔH = +40 kJ
                        B → C ΔH = +30 kJ
                       -------------------
                    A + B → B + C ΔH = +70 kJ
    

15. Know how to use standard enthalpies of formation (ΔH_f°) to calculate standard enthalpies of reaction (ΔH_rxn°) for any reaction. Enthalpies of formation are tabulated values for the formation of one mole of a compound from its elements in their standard states (at 1 atm and 298 K).
    ΔH_rxn° = ΣnΔH_f°_(products) – ΣmΔH_f°_(reactants)

Chapter 6, Electronic Structure of Atoms

1. Learn the definitions of wavelength (λ), frequency (ν), and the speed of light (c = 3.00 × 10⁸ m/s) and how to relate them (c = λν). Know the names and approximate wavelength ranges for the different parts of the electromagnetic spectrum (gamma → radio).
2. Know that although light is described as a wave, all of its interactions with atoms take place as if it were made of particles of pure energy, or photons. Know how to relate the frequency of light to energy per photon (E = hν) using Planck’s constant (h = 6.626 × 10^–34 J·s).
3. Know that the photoelectric effect, in which electrons are ejected from a surface only when light with sufficiently high frequency strikes it, is evidence for the quantum nature of light.
4. Know how the line spectrum of an element arises from the transitions between energy levels within atoms. Energy levels are quantized (they are granular, not smoothly varying) and can be numbered using integers (1, 2, 3, …). Know how to apply the terms ground state, excited state, emission, and absorption.
5. Know that matter, too, is quantized and so can be understood as particles. Still, atoms and molecules can also be modeled as waves. This is a reflection of the modern understanding of the nature of matter: It is granular, its behavior can only be predicted probabilistically, and it is only meaningful to describe interactions. The state of a particle—its exact speed and location—cannot be known between interactions.
6. Know that when treated as if they are waves, particles with mass have a wavelength given by the de Broglie relation: λ = h/mv. In this equation h has its usual value, m is mass in kilograms (kg), and v is velocity in m/s. Particles with very small mass have meaningful wavelengths but ordinary, human-scale objects do not.
7. Know that the Heisenberg Uncertainty Principle (HUP) states that it is not possible, no matter how precise a measurement can be made, to know both the position and the momentum (which depends on velocity) of a particle at the same time.
   Another way to think about this is to say that the position and momentum of a particle between interactions with other particles cannot be known. This principle is the reason why quantum events (which concern atomic and subatomic particles) can only be predicted probabilistically.
8. Know that the Schrödinger wave equation is how the probability density of, for example, an electron’s position in an atom is calculated. Probability density is a measure of the likelihood of finding an electron within a particular volume.
9. Know that the orbitals are the solutions of the wave equation and that they are governed by three quantum numbers. An orbital is a volume centered on an atomic nucleus with a specific shape and within which up to two electrons may be found.
10. Know that a shell is a set of orbitals which all share the same principal quantum number (n, which can equal 1, 2, 3, …). A sub-shell is a set of orbitals which all have the same angular momentum quantum number (l = 0, 1, 2, …, n – 1). The number of orbitals in a sub-shell (and the 3-D shape of each orbital) is determined by the magnetic quantum number (m_l = –l, …, –1, 0, 1, …, +l).

    n	l	ml	Description
    1	0	0	Shell 1, sub-shell 0, a single 1s orbital
    2	0	0	Shell 2, sub-shell 0, a single 2s orbital
    2	1	–1, 0, 1	Shell 2, sub-shell 1, three 2p orbitals
    3	0	0	Shell 3, sub-shell 0, a single 3s orbital
    3	1	–1, 0, 1	Shell 3, sub-shell 1, three 3p orbitals
    3	2	–2, –1, 0, 1, 2	Shell 3, sub-shell 1, five 3d orbitals

    See also Table 6.2, pg. 229
11. Know that electron shells each have a different negative potential energy. Since an electron not attached to an atom has zero potential energy the potential energies of electrons within an atom are negative. Think of this as a well with electrons below the surface; energy is required to pull them out of the well and this is why they have a negative potential energy. A guideline to comparing potential energies is that the lower the value of n, generally the lower the potential energy.
12. Know that electrons are assigned to orbitals in an atom in the ground state starting from the lowest potential energy toward higher energy orbitals. This is known as the Aufbau Principle.
13. Know that atoms with precisely one electron have sub-shells which are all at the same potential energy within each shell. Atoms with two or more electrons have sub-shells with different potential energies. In such atoms the energy of sub-shells increases from s to p to d to f.
14. Know the filling order of orbitals by their relative potential energies. Generally, lower values of n fill first and electrons enter orbitals in sub-shells in order as s then p then d then f. There are some complications, however. Specifically, know that the (n + 1)s subshell is always lower in potential energy than the nd sub-shell (4s is lower than 3d, for example). Also, the (n + 2)s sub-shell is always lower in potential energy than the nf sub-shell (6s is lower in energy than 4f).
15. Know how to interpret graphs of radial probability and how they relate to orbital shapes. Know what nodes are (special distances away from the nucleus within an orbital where the electron probability is zero). Know the shapes of s, p, and d orbitals.
16. Know that each orbital in a sub-shell may hold one or at most two electrons due to the Pauli Exclusion Principle. Electrons have a quantum property called spin. This is notated with a fourth quantum number called the magnetic spin quantum number (m_s). This quantum number has only two possible values +½ or –½. No two electrons within the same atom may have the same values for all four quantum numbers (n, l, m_l, and m_s). Two electrons in the same orbital must have opposite spins.
17. Know that even though a set of orbitals in a sub-shell all have the same potential energy an atom is in its lowest overall potential energy state when as many electrons as possible in the same sub-shell have the same (or parallel) spin. This is called Hund’s Rule. For example, for nitrogen, the electron configuration is 1s²2s²2p³. The three electrons in the 2p sub-shell are in separate orbitals and all have the same spin (which can be either up or down).
18. Know how to write the condensed ground state electron configuration (using noble gas cores) for any element using the periodic table as a guide.
    
19. Know how to distinguish valence and core electrons.
20. Know that elements in the same group have the same valence electron configuration but in the next higher shell. This is why elements in a group have similar characteristics.

Chapter 7, Periodic Trends

1. Know that Mendeleev put the elements in order according to similarities in chemical and physical properties. Know also that Moseley established the concept of atomic number, which equals the number of protons in the nuclei of atoms of an element, and that this provided a better way to put the elements in order than the average atomic mass.
2. Know that the valence orbitals of an atom are responsible for the observed characteristics and behavior of the atoms of an element. Know also that the similarity in the valence shell configuration is the reason why elements in a group have similar characteristics. For example, all the elements in group 16 have electron configurations ending in ns² np⁴.
3. Know that the valence orbitals of an atom are those with the highest value of the principal quantum number. For p-block elements which have (n-1)d sub-shells the ns and np sub-shells are the valence orbitals and contain the electrons that are ionized first, can form bonds, and may have unfilled orbitals which can accept electrons and form negative ions.
4. Be able to explain the relevance of the concept of effective nuclear charge to a wide variety of atomic properties. Effective nuclear charge is the idea that the core electrons shield the valence electrons from the full charge of the nucleus. The core electrons have a negative charge and partially cancel out the positive charge of the protons. For example, sulfur has 16 protons and 16 electrons. Sulfur has 6 valence electrons. The 10 core electrons cancel out the charge of 10 of the protons, leaving an effective nuclear charge of +6 for the valence electrons.
5. Know the general trend for effective nuclear charge as it can be used to explain a variety of atomic properties.
   1. One trend is that effective nuclear charge increases from left to right across a period. This is because although both protons and electrons increase in number the valence electrons, which are the ones being added, do not shield each other very much from the nucleus and the number of core electrons is constant.
   2. Another trend is that effective nuclear charge is approximately constant within a group. This is because the number of valence electrons is constant and the number of core electrons increases, balancing out the increase in positive charge in the nucleus. A slightly more nuanced view is that effective nuclear charge increases slightly from top to bottom within a group because the shielding electrons are spread out over a larger volume in larger atoms and do not shield the valence electrons quite as effectively.
6. Know that the bonding atomic radius is one half of the bond length between atoms. Atomic radii increase from top to bottom within a group because the electrons cannot all occupy the same shell and so must occupy positions in larger shells. Atomic radii decrease from left to right within a period because the effective nuclear charge increases.
7. Know that cations are smaller than their parent atoms and anions are larger. Know that for ions with the same charge ion size increases from top to bottom in a group just as atomic radius.
8. Know that an isoelectronic series is a set of atoms and ions which all have the same number of electrons. Sizes of ions in an isoelectronic series decreases with increasing nuclear charge. For example, in the series Al³⁺, Mg²⁺, Na⁺, F^–, O^2–, N^3– all of the ions have 10 electrons. The largest ion in the series is N^3– and the smallest is Al³⁺. Note that a neutral Ne atom also has 10 electrons but its actual radius does not fit in this series between F^– and Na⁺. This is because Ne does not form bonds and so its tabulated radius is based on its collisional cross-section, rather on bond lengths.
9. Know what the term ionization energy refers to. It is the energy required to remove one electron from a given atom or ion. For example, the first and second ionizations for sodium can written as an equation:
   Na → Na⁺ + e^– (first ionization)
   Na⁺ → Na²⁺ + e^– (second ionization)
10. Know that the successive ionization energies for a given atom always increase in size. Know also that there is a sharp increase in the ionization energy for electrons being removed from the core electrons of an atom as compared to the energies required to remove a valence electron.
11. Know that smaller atoms and atoms with a higher effective nuclear charge have higher first ionization energies. This is for two reasons. First, the closer an electron is to a nucleus the stronger the force of attraction. Second, the higher the charge in a nucleus, the stronger the force of attraction. The combination of these ideas leads to the conclusion that first ionization energies decrease going from top to bottom within a group and increase going from left to right within a period.
12. Know that when atoms form positive ions the first electrons lost are always those in the highest quantum shell. For example, an iron atom will form an ion by losing the two 4s electrons to make Fe²⁺. If this ion loses one more electron it is one of the six 3d electron to make Fe³⁺.
13. Know that electron affinity is a property of monatomic gases and represents the energy released when an electron binds to an atom. The point of learning about it is that it shows that while some atoms release energy when an electron binds to them, other atoms do not. For example, noble gas atoms have completely full valence shells and any additional electron would have to be added to a still higher shell. From the point of view of this higher shell the effective nuclear charge is zero because all of the protons are cancelled out by all of the electrons. Therefore there is no electrical force of attraction for an additional electron and no energy can be released by binding.
14. Know the regions of the periodic table labeled metals, non-metals, and semi-metals (or metalloids). Hydrogen is a non-metal but otherwise all elements to the left of the semi-metals are metals. The semi-metals are B, Si, Ge, As, Sb, Te. Elements to the right of the semi-metals are non-metals.
15. Know that metals have a characteristic luster (or shininess) and are good conductors of heat and electricity. When metals react with non-metals they form cations while the non-metals form anions. Most metal oxides are basic and react with water to form hydroxides or react with acids to form a salt and water.
16. Know that non-metals lack luster and are generally poor conductors of heat and electricity (with the exception of some forms of carbon). Compounds composed only of non-metals are molecular, not ionic. Non-metal oxides are acidic and react with water to form acids. Non-metal oxides react with bases to produce a salt and water.
17. Know the alkali metals (group 1), alkaline earth metals (group 2), halogens (group 17), and noble gases (group 18) and their characteristics. Specifically, what ions they form and some of their chemical reactions (if any).
18. Know the common molecular forms of some non-metals: oxygen (O₂), ozone (O₃), phosphorus (P₄), sulfur (S₈), hydrogen (H₂), nitrogen (N₂), fluorine (F₂), chlorine (Cl₂), bromine (Br₂), and iodine (I₂).

Chapter 8, Chemical Bonds

1. Be able to describe and distinguish ionic, covalent, and metallic bonds. Know the octet rule.
2. Know how ionic compounds are structured—as a 3-D lattice of alternating positive and negative ions.
3. Understand the concept of lattice energy as the energy released by the formation of the lattice from gaseous ions. Be able to compare the relative size of lattice energies for different compounds based on ionic charge and size.
4. Know how to draw Lewis structures using pairs of electrons in lone (non-bonding) pairs, single, double, and triple bonds to give atoms an octet of shared and unshared electrons. Know the number of valence electrons for each element as these are the only ones involved in bonding.
5. Know the periodic trend for values of electronegativity—the tendency to attract electrons preferentially to one end of a bond.
6. Know how to use electronegativity values to compare the degree of polarity (qualitatively) of two covalent bonds and to distinguish ionic, polar covalent, and pure covalent bonds.
7. Know how to calculate formal charge and how to use it to determine the dominant Lewis structure for a molecule. Generally, prefer to follow the octet rule for as many atoms as possible. When this is not possible then keep formal charge values low (±1) and keep negative values on electronegative atoms.
8. Know how to draw resonance structures for molecules which have no single dominant Lewis structure. Resonance structures are all structures with the same skeletal structure but different distributions of electrons. Sometimes all resonance structures are equivalent by symmetry but sometimes they are not. Remember that resonance is not reflection or rotation.
9. Know about exceptions to the octet rule:
   1. Molecules with an odd number of electrons such as NO₂.
   2. Molecules with a central atom which has less than an octet such as BeCl₂ and AlF₃.
   3. Molecules with a central atom which has more than an octet such as SF₆ and SF₄.
10. Know that breaking bonds is endothermic (+ΔH) and forming bonds is exothermic (–ΔH).
11. Know how to use bond enthalpies to estimate ΔH_rxn.
    ΔH_rxn = +ΔH_{(bonds broken)} + (–ΔH_{(bonds formed)}).

Chapter 9, Molecular Structure

1. Learn the five basic shapes of molecules of the general formula AB_n:
   

   General Formula	Shape	Illustration
   AB	Linear
   AB2	Linear
   AB3	Trigonal Planar
   AB4	Tetrahedral
   AB5	Trigonal Bipyramidal
   AB6	Octahedral

2. Learn the VSEPR model for determining the shapes of molecules. VSEPR means Valence Shell Electron Pair Repulsion. Electron domains are either a non-bonding pair (also called a lone pair) or a bond. The bond may be single, double or triple but in all three cases counts as one electron domain. Since electrons repel each other all domains around a central atom separate as far as possible from one another.
3. Learn the molecular geometry names for shapes that result when one or more electron domain is non-bonding. See table 9.2 pg. 350 and table 9.3 pg. 353 and the VSEPR shapes handout.
4. Know the bond angles associated with each VSEPR shape. Also, know how those angles are modified by lone pairs and double or triple bonds— such electron domains squeeze neighboring bond angles to slightly smaller values.
5. Know how to predict whether a molecule is polar or non-polar based on bond dipoles and molecular shapes. Perfectly symmetrical molecules with linear, trigonal planar, tetrahedral, trigonal bipyramidal, and octahedral shapes ar non-polar. Determine polarity as follows:
   1. Draw the Lewis structure
   2. Determine the 3-D shape
   3. Determine bond polarities
   4. Add bond polarities (vector addition)
6. Know that valence bond theory is based on the idea that the overlap of atomic, or hybrid atomic, quantum orbitals is what makes a covalent bond.
7. Know that hybrid orbitals are linear combinations of atomic orbitals that account for the 3-D geometries observed in molecules.
   Here is a table of hybrid orbitals, which shows how they are effectively codes for VSEPR shapes:

   Orbitals on Central Atom	Hybird Orbitals	Shape	Examples
   s, p	2 sp hybrid orbitals	linear	CO2, N2
   s, p, p	3 sp2 hybrid orbitals	trigonal planar	BF3, SO2
   s, p, p, p	4 sp3 hybrid orbitals	tetrahedral	CH4, NH3, H2O
   s, p, p, p, d	5 dsp3 hybrid orbitals	trigonal bipyramidal	PCl5, SF4, ClF3, XeF2
   s, p, p, p, d, d	6 d2sp3 hybrid orbitals	octahedral	SF6,BrF5, XeF4

8. Know the sigma (σ) and pi (π) bonding convention. A single bond is a σ-bond. A double-bond is a σ-bond and one π-bond. A triple-bond is a σ-bond and two π-bonds.
9. Know the limitations placed on the geometry of molecules when they contain π-bonds. For example, there is no rotation around the axis of a bond containing one or more π-bonds. Know that π-bonds in structures that require resonance forms indicate that electrons are delocalized—spread out over 3 or more atoms.

Chapter 11, Intermolecular Forces/Chapter 12, Solids and Modern Materials

Ch. 11

1. Understand phases of matter dynamically as the interplay between attractive intermolecular forces and the kinetic energy of molecules at different temperatures.
2. Understand that substances will be gases when the kinetic energy of molecules exceeds the (weak) bond energy of intermolecular of intermolecular bonds. Liquids form when the two energies are similar in size. Solids exist when bond energy is larger than the kinetic energy of the molecules or atoms.
3. Understand that the strength of intermolecular forces depends on molecular structure.
4. Remember the different types of intermolecular forces and what kinds of materials are affected by each.
   In order from strongest individual forces to weakest:
   ion-ion: attracts ions in ionic compounds
   ion-dipole: attracts polar molecules to ions in solutions
   hydrogen bonding: attracts polar molecules containing O—H, N—H, and F—H bonds.
   dipole-dipole: attracts polar molecules
   dispersion: attracts all molecules based on tiny differences in electric charge across the surface of molecules, even if they are neutral and/or have no dipole. The strength depends on polarizability and larger molecules have stronger dispersion forces.
5. Know that the majority of attractive forces between molecules is due to London dispersion (a.k.a. van der Waals) forces. The other intermolecular forces simply add to the total.
6. Know that viscosity is the degree of resistance to flow for a liquid. The stronger the intermolecular forces (IMF) the higher the viscosity.
7. Know that the surface tension is a measure of a liquid’s resistance to an increase in its surface area. The stronger the IMF, the higher the surface tension.
8. Know that capillary action is due to two forces: cohesion (which holds molecules of a liquid together) and adhesion (which sticks molecules of a liquid to a surface). Liquids rise higher in a capillary when these forces are stronger and/or when the capillary is more narrow.
9. Know that phase changes involve the absorption or release of energy depending on whether intermolecular bonds are broken (melt, sublime, or vaporize) or are formed (condense, deposit, freeze).
10. Know how to calculate energy absorbed or given off in temperature changes for a single phase using the calorimetry equation: q = msΔT.
11. Know how to calculate energy involved in a phase change using the heat of fusion (ΔH_fusion) for melting/freezing and the heat of vaporization (ΔH_vap) for vaporization/condensation.
12. Know how to apply such calculations to heating curves showing temperature as a fuction of heat added to find total heat required for raising temp. and to drive phase changes.
13. Know the terms critical temperature, critical pressure, and supercritical fluid.
14. Know that the vapor pressure of a liquid (or solid) is the partial pressure of the vapor phase when in equilibrium with the liquid or solid phase. Vapor pressure is a function of temperature with higher values for higher temp.
15. Know that vapor pressure is related to a liquid’s evaporation rate—the higher the P_vap, the higher the rate of evaporation. Liquids with high evap. rates are volatile. The stronger the IMF in a liquid, the less volatile it is.
16. Know that a liquid boils when its vapor pressure equals the external pressure. Since vapor pressure is lower at lower temp. a liquid will boil if external pressure is lowered until it is equal to the vapor pressure of the liquid.
17. Know how to label and interpret a phase diagram.

Ch. 12

1. Know the types of solids and the types of forces holding particles together:
   1. Metallic solids: metallic bonding (any metal)
   2. Ionic solids: ion-ion forces (any ionic solid)
   3. Covalent network solids: covalent bonding throughout (C, Si, glass [SiO₂], some polymers [for example bakelite])
   4. Molecular solids: molecular compounds held together by IMF only (H₂O, CO₂, etc.)
2. Know that some solids are crystalline with an orderly, repeating arrangement for atoms or molecules. Some solids are amorphous, such as glass, and have no long-range repeating patterns in their structures.
3. Know that metallic solids have high thermal and electrical conductivity and are malleable and ductile.
4. Know the types of alloy.
   • Substitutional alloys form when the atoms of the mixed elements are similar in size.
   • Interstitial alloys are when one atom is much smaller than the other and fits in the spaces between the larger atoms.
   • Heterogeneous alloys have components with a non-uniform distribution.
   • Intermetallic compounds may also form, which are true compounds rather than mixtures.
5. Know how the electron-sea model accounts for both the thermal and electrical conductivity of metals as well as their deformability.
6. Know how the structure of ionic compounds accounts for their high melting and boiling points and brittleness.
7. Know how the melting and boiling points of molecular substances depends on both the strength of intermolecular forces and the efficiency with which they can be packed.
8. Know some examples of covalent network solids and be able to explain the characterstics of graphite and diamond based on the nature of bonding in those materials.

Chapter 13, Solutions

1. Understand the solution formation process using the concept of energy. In forming a solution a solute becomes uniformly dispersed throughout a solvent. To achieve this state some intermolecular bonds are broken and others are formed. The energy of solution formation has three components:
   1. ΔH₁: bonds holding solute particles together are broken (endothermic)
   2. ΔH₂: bonds holding solvent particles together are broken (endothermic)
   3. ΔH₃: bonds form between solute and solvent particles (exothermic)
2. Know that the formation of bonds between solute and solvent particles is called solvation (or hydration if the solvent is water).
3. Know that the sum of ΔH₁, ΔH₂, and ΔH₃ plays a role in determining whether a solution will form. For example, solutions of vegetable oil and water do not form because ΔH₁ and ΔH₂ are too large compared to ΔH₃. When the sum is positive it makes solution formation unlikely.
4. Know that when solutions form it usually involves an increase in entropy. Entropy is a measure of the number of ways the particles in a system can be arranged. Generally, when entropy increases in some process it means the proces can be expected to happen on its own.
5. Know that a saturated solution is one in which the process of dissolution occurs at the same rate as crystallization, a situation known as dynamic equilibrium.
6. Know that the concentration of a saturated solution is the solubility of a given solute at a specific temperature. Solubility is a function of temperature and varies considerably depending on the substance.
7. Know what the terms unsaturated and supersatuared mean with respect to solutions.
8. Know how to predict, based on the similarity of applicable intermolecular forces, whether a substance is soluble in a given solvent. “Like dissolve like.” Know the terms miscible and immiscible.
9. Know that the solubility of gases follows Henry’s Law S_g = kP_g where solubility in mol/L is directly proportional to the partial pressure of the gas. Know also that, unlike solids, gas solubility decreaes with increasing temperature; also, know why this is so.
10. Know all of the various units of concentration and how to convert between them.
    • molarity (M or mol/L) depends on temperature due to thermal expansion/contraction; density of the solution is required to convert to/from other units
    • molality (m or mol/kg) moles of solutes per kilogram of solvent; independent of temperature
    • mass percent (%) calculated as mass of solute divided by total mass of solution × 100%
    • parts per million (ppm) like mass percent but times 1 × 10⁶ rather than 100
    • parts per billion (ppb) like ppm and mass percent but times 1 × 10⁹
11. Know the colligative properties of solutions and how and why they work. In the following equations m stands for molality (mol/kg) and i stands for the van't Hoff Factor, which gives the numbers of moles of particles per mole of solute. For example, i = 2 for NaCl because it forms two ions per mole when dissolved in water.
    • Vapor pressure reduction—Raoult’s Law
      P_soln = χ_solventP°_solvent
    • Freezing Point Depression
      ΔT = –ik_Fm
    • Boiling Point Elevation
      ΔT = ik_Bm
    • Osmotic Pressure
      the pressure required to just stop spontaneous osmosis is Π (capital π, like P for pressure)
      Π = iMRT
      M = mol/L and R = 0.0821 L atm/K mol

Chapter 14, Kinetics

See also the reference information page for chemical kinetics.

1. Know the notation used for concentrations of reactants and products:
   [X] = concentration in mol/L for substance X
   [X]₀ = initial concentration in mol/L for substance X
   [X]_eq = final (equilibrium) concentration in mol/L for substance X
2. Understand the meaning of a reaction rate and be able to express it mathematically in terms of change in concentration of a reactant or product as a function of time in units of M/s = (mol/L)/s. For the reaction:
   aA + bB → cC + dD
   

   1 Δ[A] rate = – ———·————— a Δt

   1 Δ[B] = – ———·————— b Δt

   1 Δ[C] = + ———·————— c Δt

   1 Δ[D] = + ———·————— d Δt

3. Know the factors which affect reaction rate and why they affect it.
   1. physical state (s, l, g) of reactants
   2. reactant and product concentrations
   3. temperature
   4. catalysts
4. Know the definition of a differential rate law. The differential rate law relates the rate of reaction to reactant concentrations. The instantaneous rate is what is meant here, not an average rate. An instantaneous rate is the slope of a line tangent to the curve of a concentration vs. time graph. The general equation of a differential rate law is:
   rate = k[A]^m[B]ⁿ
   (where in calculus notation rate = –d[A]/dt)
   
5. Understand what is meant by the order of reaction, which is the exponent on the concentration in a rate law equation.
   zero order: rate = k[A]⁰ = k (the rate is constant at all conc.)
   first order: rate = k[A]¹
   second order: rate = k[A]²
6. Know that order of reaction is not taken from the stoichiometric coefficients and must be determined experimentally.
7. Be able to write correct units for the rate constant, k:
   For example:
   

   Order	Units of k
   0	M/s (or M s–1)
   1st	1/s (or s–1)
   2nd	1/M·s (or M–1 s–1)

8. Be able to determine a rate law from measurements of concentration and rate. This means identifying the order of reaction with respect to each reactant and calculating the value of the rate constant with its units.
9. Know and be able to use the integrated rate laws and how they relate to differential rate laws for reactions with one reactant. The kinetics reference page shows these equations in several algebraic arrangements and uses calculus notation, in case that is informative to you.
   

   Order	Differential	Integrated
   0	rate = k	[A]t = –kt + [A]0
   1st	rate = k[A]1	ln[A]t = –kt + ln[A]0
   2nd	rate = k[A]2	1/[A]t = kt + 1/[A]0




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10. Know the characteristic kinetic plots and how to use them to identify order of reaction and to determine the value of k. See also the kinetics reference page.
    

    Appearance of a Concentration vs. Time Graph
    Zero Order	First Order	Second Order
    slope = –k
    Modified Plots to Confirm Order of Reaction
    (no modified plot is needed to confirm whether a reaction is zero order)	Plot the natural log of conc. vs. time slope = -k	Plot the inverse of conc. vs. time slope = k

11. Know the meaning and mathematical forms for reaction half-life and be able to perform relevant calculations, especially for first order reactions.
    

    Order	Half-life Expression
    0	t½ = [A]0/2k
    1st	t½ = ln(2)/k
    2nd	t½ = 1/k[A]0

12. Know the collision model and why temperature affects reaction rate (at higher T there are more frequent collisions and collisions have a larger average energy).
13. Know the meaning of activation energy (E_a), which represents the energy needed to initiate a reaction. Also, know how to interpret a reaction energy diagram.
    (at the top of the curve the materials of the reaction are an ‘activated complex’ or a transition state)
14. Understand the Arrhenius Equation, which relates the value of the rate constant for a reaction to activation energy and temperature.
    k = Ae^(–E_a/RT)
    k = rate constant
    A = frequency factor (related to the frequency of effective collisions)
    E_a = activation energy in kJ/mol
    R = the universal gas constant (8.314 J/K·mol)
    T = absolute temperature (K)
    Be able to predict the effect on the value of the rate constant when either E_a or T are changed.



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15. Know the logarithmic form of the Arrhenius equation and how it can be used to find the value of E_a by using measurements of k at different temperatures.
    ln(k) = –E_a/R·(1/T) + ln(A)
    The slope of a graph of a plot of ln(k) vs. 1/T is equal to –E_a/R.
16. Know how to interpret reaction mechanisms, elementary steps, and how to recognize intermediates and catalysts.
17. Know how to write rate laws for elementary reations and that the slowest (the rate-determining) step is the one whose rate law determines the overall reaction rate law.
18. Be able to understand a mechanism whether the first step is fast or rate-determining.
19. Know the descriptive materials regarding catalysts. Be able to describe heterogeneous and homogeneous catalysts, explain how surface adsorption works, and to discuss enzymes.

Chapter 15, Equilibrium

1. Know that chemical equilibrium is the situation when the forward reaction has the same rate as the reverse reaction. Although both reactions continue, all concentrations reach constant values.
2. Know the generalized law of mass action:
   For a reaction of the form, aA + bB ⇌ cC + dD, the quantitative equilibrium constant in terms of concentration (K_c) is expressed as:
   

        [C]c[D]d
   Kc = —————————
        [A]a[B]b
   

3. Know that the concentrations at equilbrium are not fixed values. Only the ratio of concentrations raised to the appropriate powers is constant. The actual concentrations can take on a wide variety of values.
4. Know how to write the expression for K_c for any reaction.
5. Know that only gases and substances in solution appear in the K_c expression. This is because pure solids and liquids have a constant concentration as long as any amount of them is present.
6. Know that the equilibrium constants can also be written in terms of partial pressures of gases:
   

        (PC)c(PD)d
   Kp = —————————
        (PA)a(PB)b
   

7. Know the formula for the relationship between K_c and K_p for the same reaction: K_p = K_c(RT)^Δn
   where T is in kelvin, R = 0.08206 L atm/K mol, and Δn = (c + d) – (a + b)
8. Know how to interpret the size of the value of K_c:
   1. for K_eq > 1, equilibrium lies to the right; products are favored
   2. for K_eq < 1, equilibrium lies to the left; reactants are favored
9. Know how manipulations of a chemical equation affect the value of K_eq (and the equilibrium constant expression).
   1. For a reaction written in reverse, K_new = (K_old)^–1
   2. For a reaction multiplied by a constant, n, K _new = (K_old)ⁿ
   3. For the sum of two reactions, K₁ · K₂
10. Know how to calculate the numerical value of K_eq from given equilibrium concentrations.
11. Know how to use stoichiometry to find equilibrium concentrations given initial and at least one final (or equilibrium) concentration.
12. Know how to calculate Q, the reaction quotient and how to use it to determine which concentrations increase as the reaction approaches equilibrium.
    1. For Q = K, the reaction is already at equilibrium
    2. For Q > K, the reaction shifts left to reach equilibrium, consuming products and forming reactants
    3. For Q < K, the reaction shifts right to reach equilibrium, consuming reactants and forming products
13. Know how to use given initial concentrations, the value of Q, stoichiometry, and the equilibrium constant expression to calculate equilibrium concentrations.

14. Know Le Châtelier’s Principle, which is that if a reaction at equilibrium is disturbed then the reaction will proceed in the direction which minimizes the impact of the disturbance. That is concentrations will change to establish a new set of concentrations which satisfy the equilibrium constant expression.

    Disturbance	Result
    Reactant concentration is increased	Reaction shifts right
    Product concentration is increased	Reaction shift left
    Volume available to gas-phase reaction is reduced	Reaction shifts to reduce total pressure (no shift occurs if the number of moles of gas products equals the number of moles of gas reactans in the balanced equation)
    Volume available to gas-phase reaction is increased	Reaction shifts to increase total pressure (no shift occurs if the number of moles of gas products equals the number of moles of gas reactans in the balanced equation)
    An inert gas is added	Equilibrium concentrations do not change
    Increase temperature for an endothermic reaction	Reaction shifts right
    Decrease temperature for an endothermic reaction	Reaction shifts left
    Increase temperature for an exothermic reaction	Reaction shifts left
    Decrease temperature for an exothermic reaction	Reaction shifts right
    A catalyst is added	Equilibrium concentrations do not change

Chapter 16, Acid-Base Equilibria

1. Know the Arrhenius definition of acids and bases:
   • Acidic aqueous solutions contain H⁺ (or H₃O⁺) ions
   • Basic aqueous solutions contain OH^– ions
2. Know the Brønsted-Lowry definition of acids and bases:
   • Acids donate H⁺ to another substance
   • Bases accept H⁺ from acids by binding to them
3. Know the conjugate-acid/conjugate-base terminology. Every acid has a conjugate base with one fewer H⁺. Every base has a conjugate acid with one additional H⁺.

   Acid	Conjugate Base		Base	Conjugate Acid
   HA	A–		A–	HA
   HCl	Cl–		B	BH+
   HC2H3O2	C2H3O2 –		OH–	H2O
   H3O+	H2O		NH3	NH4+

4. Know that water is amphiprotic, meaning that it is both a hydrogen ion donor and acceptor.
   

   Water undergoes autoionization: H2O + H2O ⇌ H3O+ + OH–

   Water as acid: H2O + B ⇌ OH– + BH+

   Water as Base: HA + H2O ⇌ H3O+ + A–

5. Know the ion-product constant for the autoionization of water at 25°C:
   H₂O ⇌ H⁺ + OH^–
   K_w = [H⁺][OH^–] = 1 × 10^–14
6. Know that the product of [H⁺] and [OH^–] is constant and that acidic solutions have [H⁺] > [OH^–] and basic solutions have [H⁺] < [OH^–]. Neutral solutions have [H⁺] = [OH^–] = 1 × 10^–7 M.
7. Know the definition of pH = –log₁₀[H⁺].
8. Be able to estimate pH for a given concentration of hydrogen ion by using the exponent for concentrations written in scientific notation.
9. Know that pK_w = –log₁₀(K_w) = 14.00. And that pOH = –log₁₀[OH^–]. Combined with the definition of pH this gives that pK_w = pH + pOH.
10. Know that significant figures for pH are only those after the decimal. The digit before the decimal effectively only gives the power of ten in the scientific notation expression for the concentration of H⁺. For example, if [H⁺] = 1.00 × 10^–3 then pH is properly reported as 3.000.
11. Know that strong acids are strong in that they completely dissociate into ions in aqueous solution:
    HX → H⁺ + X^– so that the concentrations of at H⁺ and X^– equilibrium are equal to the initial concentration of the acid.
    [H⁺]_eq = [X^–]_eq = [HX]₀
12. Memorize the strong acids: HCl, HBr, HI, HNO₃, H₂SO₄, HClO₃, and HClO₄. The anions, or conjugate bases, of these acids have zero basic activity and will not bind to H⁺ in aqueous solution.
13. Know that weak acids are weak electrolytes in that, when dissolved in water the majority of acid molecules do not dissociate into ions.
14. Know that the acid dissociation constant of a weak acid (K_a) is the K_eq for HA ⇌ H⁺ + A^–.
15. Know that K_a for all weak acids is less than 1 and acid strength is higher for acids with larger K_a values.
16. Be able to calculate [H⁺]_eq, [A^–]_eq, and pH given [HA]₀ and the value of K_a.
17. Know that for polyprotic acids such as H₃PO₄, H₂SO₄, and H₂CO₃ the successive K_a values decrease from one to the next by orders of magnitude. For this reason, only the first dissociation constant is important in finding the equilibrium concentration of H⁺. Once that is determined, along with the concentration of conjugate base from the first dissociation, it is possible to calculate the concentration of all other ions.
18. Know that many weak bases contain nitrogen and have structures similar to ammonia (NH₃) such as CH₃NH₂.
19. Know that the base dissociation constant, K_b, governs the equilibrium: B + H₂O ⇌ BH⁺ + OH^–.
    K_b = [BH⁺][OH^–]/[B].
    Be able to use this expression to find pOH and pH given the concentration of a weak base and its K_b value.
20. Know that the K_a for a weak acid is related to K_b for its conjugate base by this expression: K_a·K_b = K_w (or pK_a + pK_b = pK_w). This equation expresses the concept that the stronger a weak acid is, the weaker its conjugate base is and vice versa.
21. Know that salts are composed of a cation, which is the conjugate acid of a base, and an anion, which is the conjugate base of an acid.
22. Know that salts may have acid-base properties, depending on their composition.
    • Salts made by the reaction of a strong acid with a strong base make neutral solutions; this is because their respective conjugates have zero acid-base activity.
    • Salts which include the anion from a weak acid are basic because the conjugate base of a weak acid is a weak base.
    • Salts which include the cation from a weak acid are acidic because the conjugate acid of a weak base is a weak acid.
    • Salts which include an anion from a weak acid and a cation from a weak base may have basic or acidic solutions depending on which ion is stronger.
23. Know the structural features of acids that determine whether they are strong or weak. Strong acids have:
    1. A highly polar H—X bond and/or
    2. A weak H—X bond and/or
    3. a stable anion (X^–)
24. Know that oxyacids are strongest when the central atom is highly electronegative and when there is a larger number of oxygen atoms bound to it.
25. Know about organic acids, which contain a structure called a carboxyl group. They are collectively known as carboxylic acids.
26. Know the Lewis Acid-base definition. Lewis acids accept electron pairs and Lewis bases donate electron pairs.
    In either case, a new bond forms using the electron pair.
    • Lewis acids include the hydrogen ion (or proton, H⁺), metal ions (such as Fe³⁺), and molecules with very polar bonds (such as BF₃ or CH₂O, whether or not the molecule is polar). Lewis acids form bonds with Lewis bases using the electron pair located on the Lewis base. The site of bond formation is the center of positive charge on the Lewis acid.
    • Lewis bases include the hydroxide ion (OH^–), ammonia (NH₃), water (H₂O), the cyanide ion (CN^–), and the thiocyanate ion (SCN^–). In all cases, a lone pair on an atom in the structure is used to make a bond to a Lewis acid.

Chapter 17, Additional Aspects of Aqueous Equilibria

1. Know the meaning of the term the common ion effect. This is when there is a non-zero initial concentration of a product ion in a weak acid or base dissociation equation or in the solubility equilibrium of a sparingly soluble substance.
2. Know what a buffer is: a solution containing significant concentrations of both HA and A^– or both B and BH⁺.
3. Know the mechanism by which buffers are able to resist change to pH with small additions of a strong acid or base.
4. Know that buffer capacity depends on two things: the ratio of the concentrations of A^– and HA ([A^–]/[HA]) and the proximity of the pH of the buffer to the pK_a of the acid.
   1. If the concentration is too small only a little added base or acid destroys the buffer.
   2. If the [A^–]/[HA] ratio is not near 1 then changes to pH are resisted in one direction only.
5. Know how to calculate pH of a solution during a titration based on the relevant equilibrium and stoichiometry:

   Stage	Description	How to calculate pH
   Start	Zero base added	Solve HA ⇌ H+ + A– using the initial concentration of HA and an initial concentration of zero for H+ and A–.
   Prior to Equivalence	Any amount of added base up to but not incl. the stoichiometric equivalence point	Do stoichiometry calculation to determine how much HA is converted to A–. ([HA] = [(HA]0 – x)/(total volume) and [A–] = x/(total volume), where x is the number of moles of added base). Solve HA ⇌ H+ + A– using the new concentration of HA and the new concentration for A– leaving H+ as zero initially.
   At Equivalence	The stoichiometric equivalence point where moles of base added equals the original moles of HA	Now that all of HA has been converted to A– the relevant equilibrium is no longer the dissociation of the acid. Rather, it is the reaction of the conjugate base with water: A– + H2O ⇌ HA + OH–. Determine the number of moles of A– and the total volume; calculate [A–]. Use [A–] and the equilibrium given above to find [OH–], pOH, and pH.
   After Equivalence	An amount of added base that exceeds the number of moles of HA originally in solution	Now that the amount of strong base added exceeds the original amount of the weak acid the pH is determined by the concentration of hydroxide ions. The conjugate base, A–, is a much weaker base and has no effect on the pH of the solution. Simply find the moles of excess OH–, divide by the total volume to find [OH–], calculate pOH, and pH = 14 - pOH.

6. Know the shapes and features of a pH curve graph for various types of titrations:
   
7. Know that indicators are weak acids which absorb light differently depending on whether or not the molecules do or do not include an attached hydrogen ion. As a result their color in solution depends on pH.
8. Know that indicators change color in a pH range from 1 pH unit below the pK_a of the indicator to 1 pH unit above the pK_a of the indicator. This is because at 1 pH unit above the pK_a there is a 10 to 1 ratio of the more basic form and because at 1 pH unit below the pK_a of the indicator there is a 10 to 1 ratio of the more acidic form.
9. Know that the rapid change in pH at the equivalence point allows one to choose any indicator which changes color in the range of pH at equivalence plus-or-minus 1.
10. Know that some ionic compounds are sparingly soluble but not entirely insoluble. The concentration of the ions of these compounds at equilibrium can be described using the solubility product constant, K_sp.
11. Know that K_sp refers to the heterogeneous equilibrium of a pure solid with its dissolved aqueous ions. For example, for MgF₂ the solubility equilibrium is written as:
    MgF₂(s) ⇌ Mg²⁺(aq) + 2F^–(aq)
    and the K_sp expression is:
    K_sp = [Mg²⁺][F^–]²
12. Know that the solubility of a compound is based on the molar concentration of an ion in the compound with a 1:1 ratio with the compound. For example, the solubility of MgF₂(s) is equal to the concentration of the Mg²⁺ion.
13. Know that the solubility-product constant is an equilibrium constant (K_sp) but the molar solubility of the compound is the concentration of the compound at equilibrium.
    For the example above, K_sp = [Mg²⁺][F^–]² and solubility is [MgF₂]_eq, [Mg²⁺]_eq, or [F^–]_eq/2. The concentration of the fluoride ion would be twice as large as the concentration of the magnesium ion.
14. Know the factors that affect solubility and how to work out solubility under different conditions.
    1. Temperature—solubility may increase or decrease as temperature rises.
    2. Common ions—the presence of ions in the compound of interest due to a soluble compound also being in the mixture decreases the solubility of the sparingly soluble ionic compound (this includes the OH^– ion, whose concentration is determined by pH).
    3. Low pH–compounds containing an ion which is the conjugate base of a weak acid (a basic anion) are more soluble at low pH (in acidic solution). Compounds containing the conjugate base of a strong acid (such as Cl^–) are not affected by pH.
    4. High pH–compounds containing an ion which is the conjugate acid of a weak base (an acidic cation) are more soluble at high pH (in basic solution).
15. Know that the presence of certain Lewis bases causes the solubility of some compounds to increase. This is due to the formation of complex ions. This is when a molecule with an available lone pair of electrons uses those electrons to bind to a Lewis acid such as a metal ion. We have seen these in the paint pigments lab with the Prussian blue which contained the [Fe(CN)₆]^4– ion and the equilibrium lab, where we studied the formation of the FeSCN²⁺ ion.
16. Know how to use complex ion formation constants to calculate total metal ion concentration.
17. We will not study precipitation and separation of ions (section 17.6) or qualitative analysis for metallic elements (section 17.7).

Chapter 19, Chemical Thermodynamics

1. Know that all real processes are irreversible in that restoring initial conditions after a change can only be accomplished by making a permanent change to the surroundings.
2. Know that spontaneous processes happen ‘on their own’ and non-spontaneous processes require the expenditure of energy.
3. All real processes are quantitatively spontaneous in one direction and non-spontaneous in reverse. For example, it is spontaneous for a nail to rust but not for a rusty nail to come clean.
4. Know that entropy (S) is a thermodynamic state function with an absolute value defined as the number of microstates, or ways of arranging atoms and energy states, that could give the same macrostate. S = k·ln W, where W is the number of ways a system can be arranged at the molecular scale.
5. Know that a macrostate is the total energy, pressure, temperature, and volume of a system.
6. Know that the entropy of a system (or the surroundings) increases when there is an increase in the number of possible energy states, volume, temperature, or motions of molecules (which include translation, vibration, and rotation).
7. Know that the change in entropy for a system (ΔS_sys) is defined as
   ΔS_sys = q_rev/T
   where q_rev is the hypothetical amount of heat absorbed or given off reversibly. Since q_rev is hypothetical, and there is only one reversible pathway for any process, ΔS is a state function with units J/K.
8. Know that a process is defined as spontaneous when the process results in an increase of the entropy of the universe (ΔS_univ).
9. Know the Second Law of Thermodynamics, which states that in a spontaneous (irreversible) process ΔS_univ > 0, where ΔS_univ = ΔS_sys + ΔS_surr.
   • Note: ΔS_surr is the change in entropy of the surroundings. The special case of a purely reversible process has ΔS_univ = 0; such processes are impossible in a practical sense but useful conceptually.
10. Know that for an isothermal process (one which takes place at a constant temperature) ΔS_surr = –ΔH_sys/T
11. For a given process, ΔS_sys is positive when:
    1. The volume occupied by particles increases as when a gas expands or a solid dissolves in a solvent.
    2. The temperature of the system increases, which involves increasing the speed, rotation, and amount of vibration in molecules.
    3. The phase of matter changes from solid to liquid/gas or liquid to gas.
    4. The number of moles of gas increases in a chemical reaction.
12. Know the Third Law of Thermodynamics which says that the entropy of a perfect crystals at a temperature of absolute zero (0 K) is zero. This is becaause there is only one microstate (or way) to arrange a perfect crystal with no thermal motion. When w = 1, S = k·ln(1) = k·0 = 0.
13. Know that the entropy of a standard substance under standard conditions, symbolized S°, is an absolute value, recorded in tables of thermodynamic functions. (Remember, standard conditions are P = 1 atm, T = 298 K, conc. = 1.0 M).
14. Know how to use standard molar entropies (S°) to find the standard change in entropy (ΔS°):
    ΔS° = ΣS°_products – ΣS°_reactants
15. Know that chemical thermodynamics is concerned, in part, with predicting whether a given reaction is spontaneous or not. Since spontaneity depends on whether or not the entropy of the universe increases, and this is difficult to determine, we use a quantity called the Gibb’s Free Energy (ΔG). When ΔG is negative for a process, then it is spontaneous. When ΔG is positive, it is not spontaneous (or, equivalently, it is spontaneous in reverse).

16. Gibb’s Free Energy is defined as
    
    ΔG = –TΔS_univ
    
    Since ΔS_univ is challenging to measure (to say the least) this definition is re-written in terms of system-based quantities based on the following algebraic manipulations:
    ΔS_univ = ΔS_sys + ΔS_surr
    ΔS_univ = ΔS_sys + (–ΔH_sys/T)
    multiplying through by –T gives:
    –T(ΔS_univ = ΔS_sys + (–ΔH_sys/T))
    –TΔS_univ = –TΔS_sys + ΔH_sys
    And since ΔG = –TΔS_univ we can write:
    ΔG = ΔH_sys – TΔS_sys
17. Know how to calculate ΔG_rxn° (at 298 K, the standard temperature) for a chemical reaction using ΔH° and ΔS°:
    ΔG_rxn° = ΔH° – TΔS°
    Remember that ΔH° and ΔS° can be calculated using the standard values given in tables of thermodynamic quantities, as in the appendix at the end of your textbook.
18. Know that Hess’s Law applies to any state function so that ΔG_rxn° can be calculated using standard free energies of formation in the same way that the standard enthalpy of reaction (ΔH_rxn°) is calculated using standard enthalpies of formation.
    ΔG_rxn° = ΣG°_{f (products)} – ΣG°_{f (reactants)}
19. Since ΔG is a function of temperature it is useful to estimate its value at temperatures other than the standard temperature of 298 K. To do this, enter the values of ΔH° and ΔS° into the equation:
    ΔG_rxn = ΔH – TΔS
    This will give a value for ΔG that is not standard (note, there is no longer a little circle next to the symbol). Note that ΔH and ΔS are also both dependent on temperature but that they are generally assumed not to change value enough to affect the outcome of a calculation of an estimated, non-standard value of ΔG.
20. Know that estimating ΔG (non-standard) at various temperatures is useful for determining the spontaneity of a reaction at temperatures other than 298 K. The value you find for ΔG does not matter as much as the sign: if it’s positive then the reaction is not spontaneous and if it’s negative it is spontaneous. The following table provides a guide:
    

    ΔGrxn = ΔH – TΔS
    ΔH	ΔS		ΔG	Spontaneity
    –	+	=	–	spontaneous at all temperatures
    –	–		+ or –	spontaneous at low temperatures (so that |TΔS| < |ΔH|)
    +	+		+ or –	spontaneous at high temperatures (so that |TΔS| > |ΔH|)
    +	–		+	not spontaneous at any temperature (spontaneous in reverse)

21. Know that ΔG (non-standard; no circle) is not only a function of temperature but also of reactant and product concentrations. Its value can be calculated using the formula:
    ΔG = ΔG° + RTln Q
    (R = 8.314 J/K·mol)
22. Know that ΔG = 0 for a reaction at equilibrium. This allows the establishment of a relationship between the standard value of free energy (ΔG°) to the equilibrium constant:
    0 = ΔG° + RTln K
    ΔG° = –RTln K
23. Know that ΔG is the maximum amount of work (w) that a process can do:
    ΔG = -w_max
    (This work (w) is the same as that in the expression of the First Law of Thermodynamics: ΔE = q + w).
    In practice, this amount of work is not obtained (say, from an internal combustion engine) because the process of putting the energy to work also contributes to an increase in entropy, which reduces the energy available for work.

Chapter 20, Electrochemistry

1. Be able to assign oxidation numbers to chemical species. Review section 4.4 for the rules (pg. 138).
2. Be able to determine what species in a redox reaction is oxidized and which is reduced.
   OIL Oxidation Is Loss of electrons
   RIG Reduction Is Gain of electrons
3. Know that an oxidizing agent causes another species to become oxidized by being reduced.
4. Know that a reducing agent causes another species to become reduced by being oxidized.
5. Know how to balance redox reactions using the half-reaction method in which two reduction half-reactions are combined by reversing one.
6. Know how to construct a voltaic (or galvanic) cell physically—be able to draw and label a diagram with cathod, anode, direction of electron flow, direction of ion migration in a salt bridge, and the identity of all substances.
7. Be able to write the balanced chemical equation for a voltaic cell and calculate the standard cell voltage using standard reduction potentials:
   E°_cell = E°_red     +     (–E°_red)
                cathode          anode
   (cathode refers to the reduction half-reaction and anode refers to the oxidation half-reaction, which is a standard reduction half-reaction run in reverse)
8. Know the applicable standard conditions for E°_cell, which are 1 M for solutions, 298 K, and 1 atm for gases. For elements, it is the standard state of matter at 1 atm and 298 K.
9. Know the definition of voltage as electrical potential energy per unit charge (1 V = 1 J/C). A coulomb (C) is the SI unit of electrical charge. One coulomb is equal to the charge of 6.24 × 10¹⁸ protons. Or, 1 elementary charge is equal to 1.602 × 10^–19 C.
10. Be able to compare substances in the table of standard reduction potentials and identify those that are stronger oxidizing agents (highest E°_red) and those that are stronger reducing agents (the product in a reduction half-reaction with lowest E°_red).
11. Know how free energy is related to cell potential: ΔG = –nFE where n is the number of moles of electrons transferred in the reaction and F is Faraday’s constant, which is the amount of electrical charge per mole of electrons.
    F = 96,485 C/mol = 96,485 J/V·mol.
12. Know how ΔG, E, and K_eq are related:
    

          ΔG°    –RTln(K)    RTln(K) 
    E° = ———— = ————————— = —————————
         –nF      –nF          nF
    

    
    
13. Be able to find K_eq given E° and vice versa. You must know the value of n from the specific reaction. The gas constant in this context is R = 8.314 J/K·mol and temperature must be expressed in kelvin (K).
14. Know the Nernst Equation which allows calculation of E_cell under non-standard conditions given E°_cell and Q, the reaction quotient.
    

    ΔG = ΔG° + RTln(Q)  (from Ch. 19)

    
    

    –nFE = –nFE° + RTln(Q)
    

    
    

               RT
    E = E° – ——————·ln(Q)	
               nF	
    

    
    
15. Know about how batteries work, how corrosion occurs, and how cathodic protection can prevent corrosion.
16. Know about electrolytic cells (voltaic cells run in reverse using a source of electrical power).
17. Be able to calculate the mass of electroplated metal using the balanced redox reaction (so you know n, or moles of electrons per mole of metal) and the electrical current used. Current is measured in amperes which measures coulombs of charge per second (A = C/s).

Organic Chemistry, Ch. 01

1. Structure Communication Examples

   Be able to draw full Lewis structures, line-angle structures (or skeletal structures), and be able to interpret single-line-of-text structures, for example, CH₃CHCHCH₂COCH₂COOH.
2. Be able to draw and interpret resonance structures, which must follow these rules:
   1. All resonance structures must be valid Lewis structures. See examples below.
   2. Nucleuses must remain in the same relative positions. All bond angles remain unchanged—there's only one bond angle. The only allowed differences are in the placement of electrons in radicals, lone pairs, double-, or triple-bonds
   3. The number of unpaired electrons in radical molecules (e.g., neutral OH) must remain the same in all resonance structures.
   4. When comparing resonance structures of a single molecule, when those structures are not equivalent, the major contributor is the structure that has the lowest energy.
      1. All atoms have octets, or as many as possible.
      2. As many bonds as possible.
      3. Any negative formal charges will be on electronegative atoms like N or O.
      4. As little charge separation as possible.
   5. Resonance stabilization is most important when it allows charge to be delocalized over two or more atoms.

   Resonance Examples

3. Be able to recognize organic acid structures including but not limited to carboxylic acids, aldehydes, alcohols, and (not pictured) protonated amines, unprotonated amines, sulfate derivatives, ketones, alkanes, etc.
   
4. Know the determinants of the strength of an organic acid:
   1. Acids are stronger when H⁺ is bonded to a more electronegative atom.
   2. The more stable the anion the results from deprotonation, the stronger the acid. For example, a larger radius with the same charge is more stable.
   3. The more stable the anion, based on the resonance structures that can be drawn, the stronger the acid. For example, the acetate ion (CH₃COO^–) is more stable than the methanoate ion (CH₃O^–) because resonance spreads the negative charge between two oxygen atoms rather than confining it to just one atom.
5. Know how to predict the products of organic acid-base proton exhange reactions.
6. Know how to predict whether products or reactants are favored at equilibrium based on the pK_a of the organic acids (see appendix data included in your handout).
7. Know how to calculate the K_eq value based on the pK_a values of the organic acids involved in a proton-exchange reaction.