C hapt er 10 – C hem ical Q uant it ies Vocabulary: Representative Particle, Percent Composition, Avogadro's Number, Empirical Formula, Molar Volume, Mole, Molar Mass, STP
Topics: Be able to use the Factor-Label Method to convert between moles, particles, and grams (Avogadro’s Number). Know how to calculate formula mass for a compound. Be able to calculate the percent composition of a compound. Starting with % composition, know how to calculate the empirical formula for a molecule. Be able to determine when to use atoms, formula units, and molecules as representative particles. Be able to count the number of atoms in a formula unit or molecule. Be able to figure out how many conversions to use in a mole problem and solve it. Be able to find the molecular formula of a compound given the molar mass and the percent composition of each element.
Problems 1. Find the molar mass of the following: a. ammonium acetate c. aluminum sulfate
b. tin (II) hydroxide d. CuSO4 · 5 H2O
2. Identify the representative particle for each of the following: a. zinc (metal) d. strontium oxide
b. carbon tetrachloride (CCl4) e. chlorine (gas)
c. xenon hexafluoride (XeF6) f. iron (III) fluoride
3. Mole problems a. b. c. d. e.
How many moles of zinc (II) chloride are in 143 grams? 4.82 moles of sodium acetate contain how many formula units? 24 How many moles are there in 5.00 x 10 atoms of calcium? How many formula units are in 3.45 grams of magnesium chloride? 22 How many grams of dinitrogen trioxide (N2O3) are equivalent to 5.78 x 10 molecules?
4. Find the percent composition of: a. copper in CuS2O3
b. sodium in sodium phosphate c. water in BaSO4 · 4 H2O
5. Empirical and Molecular Formulas a. What is an empirical formula? How is the molecular formula different than the empirical? b. What is the empirical formula of a compound that is 67.6% Hg, 10.8% S, and 21.6% O? c. Methyl butanoate has a percent composition of 58.8% C, 9.8 % H, and 31.4 % O. Its molecular weight is 102 g/mol. Find the molecular formula. d. A compound is found to have 9.09 g C, 1.52 g H, and 14.4 g F. What is its empirical formula? If the compound has a molecular mass of 66 g/mol, what is its molecular formula?
C hapt er 11 – C hem ical R eact ions Vocabulary: Chemical equation, Balance equation, combination reaction, Activity Series, Decomposition Reaction, Single-Replacement Reaction, catalyst, combustion reaction, coefficient, skeleton equation, double-replacement reaction
Topics: What are the four phase indicators? Know the four types of reactions. Be able to write chemical equations from the worded equation using an ion sheet. Know how to balance chemical equations.
Chapter 12 – Stoichiometry Vocabulary: Stoichiometry, actual yield, excess reagent, limiting reagent, mole ratio, percent yield, stoichiometry, theoretical yield.
Topics: Be able to distinguish between excess reactant and limiting reactant. Be able to solve MASSMASS problems. Be able to calculate which reactant in a chemical reaction is the limiting reactant. Know the difference between actual yield and theoretical yield. Be able to use the aforementioned to calculate the percentage yield of a reaction. Be able to define “mole ratio” and explain why it is essential for stoichiometry. Given grams of one substance, find grams of another substance.
Problems 1.
____ CaCl2 + ___ Al2O3 ___ CaO + ___ AlCl3 a. How many moles of calcium chloride would react with 5.99 moles of aluminum oxide? b. If 2.44 moles of calcium oxide are made, how many grams of aluminum chloride are also made? c. If 14.5 grams of CaCl2 react with excess Al2O3, how many grams of CaO are produced?
2. Ammonium chloride reacts with lead (IV) nitrate. a. Write a balanced equation for this reaction. b. If 18.5 grams of each reactant is present: i. Which reactant limits? ii. How many grams of the excess reactant would remain? iii. How much lead (IV) chloride is made? iv. If 9.50 grams of lead (IV) chloride are experimentally made, what is the % yield? 3. Butane (C4H10) is combusted in the air. a. Write the balanced equation for this combustion reaction. b. How many moles of carbon dioxide are made if 43.0 grams of oxygen gas react? c. If 3.22 grams of butane react with 10.4 grams of oxygen gas, what is the maximum mass of water vapor that can be produced?
Chapter 13 – States of Matter Vocabulary: Gas Pressure, amorphous solid, barometer, freezing point, phase diagram, evaporation, sublimation, melting point, vapor pressure, Kinetic energy, Crystal, triple point, kinetic theory, unit cell, glass, Vacuum, Allotrope, Standard atmosphere, pascal, Boiling point, normal boiling point, vaporization, atmospheric pressure
Topics: Know the parts of a “Change of State” diagram. How does the Kinetic-Molecular Theory explain the different states of matter in terms of energy, particle speed, collisions, and temperature? Know the key parts of a “Phase Diagram.” (101.3 kPa = 760 mm).
1. Diagrams: a. Draw a heating/cooling curve for a substance that melts at 48°C and boils at 287°C. In each section of the heating/cooling curve, label the states of matter and what is happening to the molecules. Identify the phase changes that could occur. b. i. Label the states of matter in the diagram. ii. What states of matter exist at each of the three dotted line intersections. iii. Label the possible phase changes that can occur by name. iv. Label the triple point and the critical point.
2. Magnesium reacts with hydrochloric acid. a. Write a balanced equation for this reaction. b. How many grams of magnesium would react completely with 50.0 mL of 2.50 M HCl? c. If 4.68 grams of magnesium is initially placed in the flask, how much excess magnesium theoretically remains after reaction? d. The experiment is run and 2.00 grams of magnesium remain experimentally. What is the percent error for this experiment?
Chapter 14 – The Behavior of Gases Vocabulary: Gay-Lussac's Law, Boyle's Law, Ideal gas constant, Charles’s Law, Diffusion, Effusion, Graham's Law of Effusion, Ideal Gas Law, Partial Pressure, Dalton's law of partial pressure, combined gas law, Compressibility Topics: Be able to solve problems and pressure conversion problems found in chapter (101.3 kPa = 760 mm). Be able to recall the three numbers and units associated with standard pressure. Utilize the factor-label method to convert between units of pressure. Utilize the Boyle’s Law equation to solve for an unknown pressure or volume. Define and explain absolute zero. Recall the values associated with STP. Write a formula and solve for a variable using the Combined Gas Law. Recognize and write the ideal gas law. Use the Ideal Gas Law to solve for an unknown quantity. Solve gas stoichiometry problems using the ideal gas law.
Solve gas stoichiometry problems using the molar volume at STP (1 mole = 22.4 L)
Able to explain the process of diffusion and effusion and relate it to molar masses of molecules (Graham’s Law of Effusion)
Problems 1. For each of the following gas laws: (1) write an equation for the gas law (2) state whether they have a direct or inverse relationship (3) draw what the graph would look like for a, b, c only a. b. c. d. e. f. g.
Boyle’s Law Charles’ Law Gay-Lusacc’s Law Combined Gas Law Ideal Gas Law Dalton’s Law of Partial Pressure Graham’s Law of Effusion
2. Gas Law Problems a. If the volume goes from 4.0 L to 0.50 L, what would be the new temperature if the pressure stays constant? b. Find the volume of hydrogen at 45.0°C at a pressure of 15.0 kPa, knowing that there was 22.4 L of the gas at STP. c. What happens to the pressure inside a balloon if it began at 721 mm Hg and 1.3 L, and it was inflated to 3.8 L? d. What is the temperature if a gas at 834 mm Hg and 21.0°C is compressed to 1.69 atm? e. If 45.0 mL of gas are held at 89.6 kPa and 33.0°C, at what temperature would they have a pressure of 95.6 kPa and 37.5 mL?
3. Ideal Gas Law problems: a. How many grams of methane (CH4) are inside the container if there are 34.4 L of gas at 15.0°C and 0.989 atm? b. What is the pressure inside the container if 56.6 grams of sulfur hexafluoride (SF6) are placed in a 385 mL tank at 79°C?
4. Gas Stoichiometry problems: a.
_1_ Al2(CO3)3 (s) _1_ Al2O3 (s) + _3_ CO2 (g) If 32.2 grams of aluminum carbonate are decomposed, what pressure of CO2 is produced in a 785 mL tank at 153°C?
b.
Hydrogen gas and fluorine gas produce hydrogen fluoride gas. After writing the balanced equation: i. How many Liters of fluorine gas react with 14.7 L of hydrogen gas at STP? ii. How many Liters of hydrogen fluoride gas are produced at STP?
5. Diffusion/Effusion a. Define diffusion. b. Define effusion. c. Rank the following from the slowest rate of diffusion to the fastest: Helium, carbon dioxide, nitrogen gas, sulfur dioxide (SO2), methane (CH4)
Chapter 15 – Water and Aqueous Systems Vocabulary: Deliquescent, suspension, emulsion, solvent, surfactant, surface tension, anhydrous, nonelectolyte, effloresce, solute, Desiccant, Solvation, Brownian motion, water of hydration, strong electrolyte, hydrate, colloid, hygroscopic, aqueous solution, weak electrolyte, electrolyte, Tyndall effect, hydrogen bonds, surface tension
Topics: 15.1 Water and Its Properties I. Liquid Water A. Surface Tension 1. Surface Tension a. A force that tends to pull adjacent parts of a liquid's surface together, thereby decreasing surface area to the smallest possible size b. Hydrogen bonding in water creates stronger than normal surface tension 2. Capillary Action a. The attraction of the surface of a liquid to the surface of a solid B. Vapor Pressure 1. Water has a very low vapor pressure due to the strong hydrogen bonding on the surface II. Water in the Solid State A. Density 1. Water is one of only a few substances that is less dense as a solid than as a liquid B. High melting point 1. No other substance with such small molar mass has so high freezing/melting point 15.2 Homogeneous Aqueous Systems I. Solutions A. Soluble 1. Capable of being dissolved B. Solution 1. A homogeneous mixture of two or more substances in a single phase C. Solvent 1. The dissolving medium in a solution D. Solute 1. The dissolved substance in a solution E. Types of solutions 1. Gaseous mixtures a. Air is a solution 2. Solid solutions a. Metal alloys 3. Liquid solutions a. Liquid dissolved in a liquid (alcohol in water) b. Solid dissolved in a liquid (salt water) II. Solutes: Electrolytes vs. Nonelectrolytes A. Electrolyte 1. A substance that dissolves in water to give a solution that conducts electric current 2. Solutions of acids, bases and salts are electrolytes B. Nonelectrolyte 1. A substance that dissolves
1. Good conductors a. Lamp glows brightly, ammeter registers a substantial current 2. Moderate conductors a. Lamp is dull, ammeter registers a small current 3. Nonconductors a. Lamp does not glow, ammeter may not register a current at all
15.3 Heterogeneous Aqueous Systems I. Suspensions A. A mixture from which particles settle out upon standing II. Colloids A. Colloidal Dispersions (Colloids) 1. Tiny particles suspended in some medium 2. Particles range in size from 1 to 1000 nm. B. Tyndall Effect 1. Scattering of light by particles a. Light passes through a solution
Chapter 16 – Solutions Vocabulary: Concentrated Solution, Supersaturated Solution, Miscible, Boiling point elevation, Henry's Law, Molality, Dilute Solution, mole fraction, concentration, colligative property, unsaturated solution, boiling-point elevation constant, Solubility, freezing-point depression constant, freezing-point depression, saturated solution, immiscible, Molarity,
Topics: Define solute and solvent and how they are used to describe solutions. Identify 3 factors that will increase the rate of dissolving and explain why. Define solubility. Describe the conditions needed for solutions to be saturated, unsaturated, saturated with a precipitate, and supersaturated. Calculate the amount of solute/water given the amount of the other and the temperature. Determine the temperature that a certain amount of solute will dissolve with calculations. Define and calculate molarity. Calculate grams of solute or volume of solvent by using the molarity equation. Calculate a new molarity or volume
using the dilution equation. Explain what colligative properties depend upon. Solve problems for grams, volume, or moles using molarities as a conversion and solution stoichiometry.
Problems 3. Define a solute, a solute, and a solution. 4. What three things can be done to increase the rate of the dissolving process? a. As you __________ the temperature, you can dissolve MORE solid. b. As you __________ the temperature, you can dissolve MORE gas.
5. Molarity: a. What is molarity? b. Find the molarity of a solution if 4.89 moles of sodium nitride are dissolved in 1.86 L of water? c. How many Liters of water are required to make a 0.228 M solution of CuF2 when 0.0388 moles of the solute are added? d. What is the molarity of FeCl3 if 18.2 grams of iron (III) chloride are dissolved in 350. mL of water? e. How many grams of calcium nitrate are needed to make 150.0 mL of a 3.20 M solution?
6. Dilutions: a. What happens to the molarity of a solution if you dilute it? Why? b. How would you prepare 2000.00 mL of a 0.50 M Ca(NO3)2 from a 3.00 M stock solution? c. If 55.0 mL of 1.00 M HCl are diluted to a final volume of 333 mL, what is the new molarity?
7.
___ Ba(NO3)2 (aq) + ___ Li2SO4 (aq) ___ BaSO4 (s) + ___ LiNO3 (aq) a. If 38.0 mL of 0.433 M barium nitrate is added, how many mL of 0.787 M lithium sulfate is needed for complete reaction? b. If 125 mL of 2.50 M lithium sulfate reacts, how many grams of the precipitate are made?
Chapter 19 – Acids, Bases and Salts Vocabulary: acid, Arrhenius theory, base, Bronsted-Lowry theory, conjugate acid, conjugate base, end point, equivalence point, pH, strong acid, strong base, weak acid, weak base, Standard Solution, Basic Solution, Acidic Solution, neutralization reaction, Salt Hydrolysis, Buffer capacity, hydronium ion, Lewis Base, Lewis Acid, Titration, pH, Ampoteric, Monoprotic, diprotic, triprotic, buffer
Topics: Identify properties of acids and bases.Predict the pH of a substance using indicator data. Explain what makes a strong acid/base strong and identify strong acids and strong bases. Calculate between [H+], [OH-], pH, and pOH. Define an Arrhenius acid and base and a Bronsted-Lowry acid and base. Predict the conjugate acid and conjugate base produced in a Bronsted-Lowry reaction. Define neutralization and predict the products of a neutralization reaction. Use Stoichiometry to calculate values in an acid-base reaction. Use titration to find the molarity of an unknown acid or base.
Problems 1. List 5 properties of acids and 5 properties of bases. 2. What is the difference between a strong acid/base and a weak acid/base? 3. pH calculations: a. Find the pH of the following and state whether they are acidic, basic, or neutral: + -5 -13 i. [H ] = 2.5 x 10 M ii. [OH ] = 3.90 x 10 M + b. Find the [H ] for the following: i. pH = 10.5 ii. pOH = 6.30 4. Arrhenius and Bronsted-Lowry acids and bases: a. How did Arrhenius define acids and bases? b. How did Bronsted-Lowry define acids and bases? c. Complete the reaction, and label the acid, base, conjugate acid, and conjugate base: i. H3PO4 + H2O ii. CH3NH2 + H2O 5. For the following reaction: ___ NaOH + ___ H2SO4 a. Predict the products and balance the equation. b. How many moles of base are required to neutralize 2.3 moles of acid? c. How many Liters of 0.345 M H2SO4 are required to neutralize 45.0 mL of 0.750 M NaOH? 6.
In a titration lab, it took 15.0 mL of 3.00 M HCl to neutralize 25.0 mL of sodium hydroxide. What is the concentration of the sodium hydroxide solution?
