College chemistry formula sheet, Cheat Sheet of Chemistry

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College Chemistry First Semester Review Sheet Fall, Dr. Wicks Chapter 1: Chemistry: The Study of Change • I can explain how the subject of chemistry fits into science and into everyday life. • I can explain the scientific method to someone not enrolled in AP Chemistry. • I can understand the language used in the scientific method and I can distinguish between an experiment, data, an independent variable, and a dependent variable. • I can distinguish between qualitative and quantitative results. • I can distinguish between and give examples of a law, a hypothesis, and a theory. • I can use a classification scheme for chemical matter. • I can distinguish between and give examples of homogeneous mixtures and heterogeneous mixtures. • I can distinguish between and give examples of elements, compounds, atoms, and molecules. • I can explain the difference between chemical and physical changes and give examples of chemical and physical properties. • I can distinguish between intensive and extensive properties. • I can use metric-metric and English-metric conversion factors to solve problems. Tera- T trillion 1012 = 1,000,000,000,000 1 inch (in.) = 2.54 cm Giga- G billion 109 = 1,000,000,000 1 pound (lb.) = 454 g Mega- M million 106 = 1,000,000 1 quart (qt.) = 0.946 L Kilo- k thousand 103 = 1,000 one 100 = 1 1 mL = 1 cm3 = 1 cc Deci- d tenth 10-1 = 0.1 Centi- c hundredth 10-2 = 0.01 pph = parts per hundred = % Milli- m thousandth 10-3 = 0.001 ppm = parts per million Micro- µ millionth 10-6 = 0.000001 ppb = parts per billion Nano- n billionth 10-9 = 0.000000001 Pico- p trillionth 10-12 = 0.000000000001 Chemical Matter Pure Substance Element Atoms Molecules (Atoms of same element) Compound Molecules (Atoms of different elements) Mixture Homogeneous Mixture Heterogeneous Mixture College Chemistry First Semester Review Sheet, Page 2 • I can convert between oF, oC, and K. oC = 5 9 ⎛ ⎝⎜ ⎞ ⎠⎟ oF − 32( ) oF = 9 5 ⎛ ⎝⎜ ⎞ ⎠⎟ oC + 32 K = oC + 273.15 • I can explain the difference between precision and accuracy. • I can apply the rules for using significant figures in calculations. • I can use dimensional analysis for problem solving. • I can use densities and percents as conversion factors in problem-solving. Chapter 2: Atoms, Molecules, and Ions • I can describe how Dalton’s atomic theory explained the law of conservation of mass, the law of definite proportions (law of constant composition), and the law of multiple proportions. See Table 1. Table 1: Laws Explained by Dalton’s Atomic Theory Law Meaning Law of Conservation of Mass Matter can be neither created nor destroyed. Law of Definite Proportions (Law of Constant Composition) Different samples of the same compound always contain its constituent elements in the same proportions by mass. Law of Multiple Proportions If two elements can combine to form more than one compound, the masses of one element that combine with a fixed mass of the other element are in ratios of small whole numbers. • I can explain the historical development of atomic theory and identify some of the scientists who have made important contributions. • I can explain the significance of Millikan’s oil drop experiment and Rutherford’s gold foil experiment. • I can describe the structure of the atom using protons, neutrons, and electrons. • Given atomic numbers and mass numbers, I can calculate the number of protons, neutrons, and electrons in atoms of given elements. • I can explain what isotopes are and how isotopic abundance can be used to calculate the atomic mass of an element. • Given a periodic table of the elements, I can identify the location of groups, periods, metals, nonmetals, metalloids (semimetals), alkali metals, alkaline-earth metals, halogens, noble gases, and transition metals. • I can give examples of allotropes. • I can use the periodic table to determine charges for ions of given elements. • I know the names, chemical formulas, and charges for common polyatomic ions. H3O+ Hydronium HCO3 - Hydrogen Carbonate ClO4 - Perchlorate NH4 + Ammonium (also called Bicarbonate) ClO3 - Chlorate OH- Hydroxide CO3 2- Carbonate ClO2 - Chlorite C2H3O2 - Acetate SO4 2- Sulfate ClO- Hypochlorite CN- Cyanide PO4 3- Phosphate MnO4 - Permanganate NO3 - Nitrate HPO4 2- Hydrogen Phosphate CrO4 2- Chromate NO2 - Nitrite H2PO4 - Dihydrogen Phosphate Cr2O7 2- Dichromate College Chemistry First Semester Review Sheet, Page 5 Chapter 4: Reactions in Aqueous Solution • I can explain the difference between a strong electrolyte, a weak electrolyte, and a nonelectrolyte. • I can use the solubility rules to predict the solubility of ionic compounds in water. See Table 4. • I can write the ions formed when an ionic compound dissolves in water. • I can predict products for precipitation reactions (double replacement reactions). • I can write molecular, ionic, and net ionic equations, and I can identify spectator ions. • I can recognize common acids and bases. • I can explain the difference between strong and weak acids and strong and weak bases. • I can write equations for acid ionization and base dissociation. • I can write molecular, ionic, and net ionic equations for acid-base neutralizations (double replacement reactions). • I understand that when acid-base reactions form “salts,” this does NOT mean that they all form table salt, NaCl. The word “salt” in this context refers to an inorganic compound whose cation comes from a base and whose anion comes from an acid. Table 4: Solubility Rules for Common Ionic Compunds in Water at 25oC General Rule Exceptions to the Rule Almost all compounds containing alkali metal ions (Li+, Na+, K+, etc.) and NH4 + are soluble Almost all compounds containing nitrates (NO3 -), bicarbonates (HCO3 -), and chlorates (ClO3 -) are soluble Most compounds containing halides (Cl-, Br-, and I-) are soluble Halides of Ag+, Hg2 2+, and Pb2+ Most sulfates (SO4 2-) are soluble Sulfates of Ag+, Ca2+, Sr2+, Ba2+, Hg2 2+, and Pb2+ Carbonates (CO3 2-), phosphates (PO4 3-), chromates (CrO4 2-), and sulfides (S2-) are usually insoluble Compounds containing alkali metal ions (Li+, Na+, K+, etc.) and NH4 + Hydroxides (OH-) are usually insoluble Hydroxides containing alkali metal ions (Li+, Na+, K+, etc.) and Ba2+ • I can identify Bronsted acids and Bronsted bases. • I can predict the products of simple gas forming reactions. • I can identify and write balanced chemical equations for four common types of reactions in aqueous solution: 1. precipitation reactions - form a solid that does not dissolve in water (double replacement). 2. acid-base reactions - form a “salt” and water (double replacement). 3. gas-forming reactions - form a gas like CO2. 4. oxidation-reduction reactions - transfer electrons (combination, decomposition, combustion, and single replacement). • I can use an activity series to predict whether a single replacement reaction will occur. College Chemistry First Semester Review Sheet, Page 6 • I can calculate oxidation numbers for elements in a compound. • I can recognize “redox” reactions (oxidation-reduction reactions) and write simple half reactions. • I can identify oxidizing agents and reducing agents. (Recall “OIL RIG,” which stands for “oxidation is loss” and “reduction is gain.” • I understand how molarity is used as a unit for concentration. • I can calculate the molar concentration of ionic substances dissolved in water using Molarity = Moles of Solute Liters of Solution ⎛ ⎝⎜ ⎞ ⎠⎟ • I can prepare solutions of various molar concentrations: 1. from solids. 2. by diluting liquids. (Recall C1V1 = C2V2 where C = concentration and V = volume.) • I can solve stoichiometry problems involving concentrations. • I can perform a “titration” and a “standardization” experimentally and determine the unknown concentration for a reactant. Chapter 19: Electrochemistry • I can balance redox equations using the ion-electron method. • Here are the steps for balancing a redox equation occurring in acidic solution: 1. Divide the equation into two half reactions 2. Balance atoms other than H and “O” 3. Balance “O” by adding water 4. Balance H by adding hydrogen ion 5. Balance the net charge by adding electrons 6. Make electron gain = electron loss, and add the half-reactions 7. Cancel anything the same on both sides of the final equation • Here are the steps for balancing a redox equation occurring in basic solution: 1. Divide the equation into two half reactions 2. Balance atoms other than H and “O” 3. Balance “O” and H by adding water and hydroxide ion 4. Balance the net charge by adding electrons 5. Make electron gain = electron loss, and add the half-reactions 6. Cancel anything the same on both sides of the final equation Chapter 5: Gases • I understand how pressure measurements are made. • I can convert between various pressure units: 1atm = 760 mm Hg = 760 torr = 101.3 kPa = 14.7 psi • I know the standard conditions of temperature and pressure (STP) are 0oC (273 K) and 1 atm. • I recall that all temperature-related gas laws use Kelvin rather than Celsius temperature. • I understand that absolute zero on the Kelvin scale is the temperature at which all atomic and molecular motion stops. It can be determined by extrapolating a line to zero volume on a V-T diagram (Charles’ law) or by extrapolating a line to zero pressure on a P-T diagram (Gay-Lussac’s law). • I can distinguish between Boyle’s law, Charles’ law, Gay-Lussac’s law, Avogadro’s law, the combined gas law, the ideal gas law, Dalton’s law, and Graham’s law, which are presented in Table 5 along with values for the gas law constant, R . College Chemistry First Semester Review Sheet, Page 7 • I can calculate gas densities and molar masses by substituting n = m MM into the ideal gas law, PV = nRT , to get d = m V = PMM RT and MM = mRT PV = dRT P . • I understand that the volume occupied by one mole of an ideal gas at standard conditions of temperature and pressure (STP, 0oC and 1 atm) is 22.4 L/mole, and I can use this conversion factor in problem solving. • I can apply the gas laws when solving stoichiometry problems involving gaseous materials. • I understand that the partial pressure of a gas in a mixture, PA , is given by its mole fraction, XA , multiplied by the total pressure of the mixture, Ptotal : PA = XAPtotal . (The mole fraction of a component A in a mixture is defined as the number of moles of A divided by the total moles of all components.) • I can apply the kinetic-molecular theory of gas behavior at the molecular level. 1. KE = 1 2 mu2 where KE = average kinetic energy of the gas molecules, m = mass of the gas, and u2 = the average of the squares of the molecular speeds. 2. Since KE is determined by temperature, heavier molecules move with a slower average speed than lighter molecules at a given temperature. 3. For a given gas, the only way to change its average KE is to change its temperature. • I can distinguish between ideal gases and real gases. Ideal gas molecules are assumed to have no volume and to not interact with one another by intermolecular forces. Deviations from ideal behavior are greatest at high pressure and low temperature. • Ideal gas characteristics: 1. Gases are composed of very tiny molecules with lots of empty space between them. 2. Gas molecules move rapidly, move in straight lines, and travel in random directions. 3. Gas molecules do not attract each other. 4. Gas molecules have elastic collisions. In other words, the total kinetic energy of the molecules before and after a collision is the same. 5. The average kinetic energy of gas molecules is proportional to the absolute (Kelvin) temperature. • In kinetic theory, the molecules of a gas are imagined to be a large number of points bouncing off the walls of a container. I understand that gas pressure is related to the number of collisions that occur with the walls per unit time. • I can calculate the root mean square (rms) speed of molecules in a gas at Kelvin temperature T using urms = 3RT MM where MM is the molar mass in kg/mole. • I can compare and contrast diffusion and effusion and I know how to use Graham's Law. Graham's Law states that the rates of diffusion (or effusion) of two gases are inversely proportional to the square roots of their molar masses at the same temperature and pressure: r1 r2 = MM 2 MM1 • When studying the behavior of real gases, I can use the van der Waals equation: P + an 2 V 2 ⎛ ⎝⎜ ⎞ ⎠⎟ V − nb( ) = nRT where a and b are van der Waals constants which correct for intermolecular forces and molecular volume.