Name: _________________________________________ Date: _____________________ Period: _____ AP Chemistry Summer Packet Pre-AP Chemistry Material Covered 2017-2018 ¨ Unit 01: Matter - C.4A Differentiate between physical and chemical properties - C.4C Compare solids, liquids, and gases in terms of compressibility, structure, shape, and volume. - C.4D Classify matter as pure substances or mixtures through investigation of their properties. ¨ Unit 02: Atomic Structure & the Periodic Table - C.5B Identify and explain the properties of chemical families, including alkali metals, alkaline earth metals, halogens, noble gases, and transition metals, using the Periodic Table. - C.5C Interpret periodic trends, including atomic radius, electronegativity, and ionization energy, using the Periodic Table. - C.6B Describe the mathematical relationships between energy, frequency, and wavelength of light using the electromagnetic spectrum. - C.6C Calculate average atomic mass of an element using isotopic composition. ¨ Unit 03: Chemical Bonding - C.6D Express the arrangement of electrons in atoms of representative elements using electron configurations and Lewis valence electron dot structures. - C.7E Classify molecular structure for molecules with linear, trigonal planar, and tetrahedral electron pair geometries as explained by Valence Shell Electron Pair Repulsion (VSEPR) theory. ¨ Unit 04: Chemical Formulas - C.7A Name ionic compounds containing main group or transition metals, covalent compounds, acids, and bases using International Union of Pure and Applied Chemistry (IUPAC) nomenclature rules. - C.7B Write the chemical formulas of ionic compounds containing representative elements, transition metals and common polyatomic ions, covalent compounds, and acids and bases. ¨ Unit 05: Chemical Equations & Reactions - C.8E Write and balance chemical equations using the law of conservation of mass. - C.8F Differentiate among double replacement reactions, including acid-base reactions and precipitation reactions, and oxidation-reduction reactions such as synthesis, decomposition, single replacement, and combustion reactions. ¨ Unit 06: Mole Concept - C.8A Define and use the concept of a mole. - C.8B Calculate the number of atoms or molecules in a sample of material using Avogadro’s number. - C.8C Calculate percent composition of compounds. - C.8D Differentiate between empirical and molecular formulas.
¨ Unit 01: Matter - C.4A Differentiate between physical and chemical properties - C.4C Compare solids, liquids, and gases in terms of compressibility, structure, shape, and volume. - C.4D Classify matter as pure substances or mixtures through investigation of their properties.
¨ Unit 02: Atomic Structure & the Periodic Table - C.5B Identify and explain the properties of chemical families, including alkali metals, alkaline earth
metals, halogens, noble gases, and transition metals, using the Periodic Table. - C.5C Interpret periodic trends, including atomic radius, electronegativity, and ionization energy,
using the Periodic Table. - C.6B Describe the mathematical relationships between energy, frequency, and wavelength of light
using the electromagnetic spectrum. - C.6C Calculate average atomic mass of an element using isotopic composition.
¨ Unit 03: Chemical Bonding - C.6D Express the arrangement of electrons in atoms of representative elements using electron
configurations and Lewis valence electron dot structures. - C.7E Classify molecular structure for molecules with linear, trigonal planar, and tetrahedral electron
pair geometries as explained by Valence Shell Electron Pair Repulsion (VSEPR) theory.
¨ Unit 04: Chemical Formulas - C.7A Name ionic compounds containing main group or transition metals, covalent compounds,
acids, and bases using International Union of Pure and Applied Chemistry (IUPAC) nomenclature rules.
- C.7B Write the chemical formulas of ionic compounds containing representative elements, transition metals and common polyatomic ions, covalent compounds, and acids and bases.
¨ Unit 05: Chemical Equations & Reactions - C.8E Write and balance chemical equations using the law of conservation of mass. - C.8F Differentiate among double replacement reactions, including acid-base reactions and
precipitation reactions, and oxidation-reduction reactions such as synthesis, decomposition, single replacement, and combustion reactions.
¨ Unit 06: Mole Concept - C.8A Define and use the concept of a mole. - C.8B Calculate the number of atoms or molecules in a sample of material using Avogadro’s
number. - C.8C Calculate percent composition of compounds. - C.8D Differentiate between empirical and molecular formulas.
¨ Unit 07: Stoichiometry - C.8G Perform stoichiometric calculations, including determination of mass relationships between
reactants and products and percent yield. - C.8H Describe the concept of limiting reactants in a balanced chemical equation.
¨ Unit 08: Gases - C.8G Perform stoichiometric calculations, including determination of mass and gas volume
relationships between reactants and products and percent yield. - C.9A Describe and calculate the relations between volume, pressure, number of moles, and
temperature for an ideal gas as described by Boyle's law, Charles' law, Avogadro's law, Dalton's law of partial pressure, and the ideal gas law.
- C.9B Describe the postulates of kinetic molecular theory.
¨ Unit 09: Solutions - C.10B Apply the general rules regarding solubility through investigations with aqueous solutions. - C.10C Calculate the concentration of solutions in units of molarity. - C.10D Calculate the dilutions of solutions using molarity. - C.10E Distinguish among types of solutions such as electrolytes and nonelectrolytes; unsaturated,
saturated, and supersaturated solutions. - C.10F Investigate factors that influence solid and gas solubilities and rates of dissolution such as
temperature, agitation, and surface area.
¨ Unit 10: Acids & Bases - C.10G Define acids and bases and distinguish between Arrhenius and Bronsted-Lowry definitions
and predict products in acid-base reactions that form water. - C.10H Define pH and calculate the pH of a solution using the hydrogen ion concentration.
¨ Unit 11: Thermochemistry (Will cover in AP Chemistry) - C.11A Describe energy and its forms, including kinetic, potential, chemical, and thermal energies. - C.11B Describe the law of conservation of energy and the processes of heat transfer in terms of
calorimetry. - C.11C Classify reactions as exothermic or endothermic and represent energy changes that occur
in chemical reactions using thermochemical equations or graphical analysis. - C.11D Perform calculations involving heat, mass, temperature change, and specific heat.
1. Count the total number of valence electrons for all atoms
2. Account for formal charges (oxidative states) for total number of valence electrons ¨ Cation (+) Loss of # e-: Subtract electron(s) from total ¨ Anion (-) Gain of # e-: Add electron(s) to total
3. Identify the central atom (the first atom written unless that atom is hydrogen). Arrange all terminal atoms
around that atom in accord with the normal valances of the atoms. ¨ Hydrogen and Fluorine atoms NEVER have more than one bond. ¨ Carbon is central when present ¨ Almost always the Least electronegative element will be the central atom for the Lewis structure.
Hydrogen and Fluorine will always be outer atoms. ¨ Place Hydrogen last on the structure
4. Complete the octet (8 electrons) for all atoms in the Lewis structure with lone pairs of electrons
¨ Carbon, Nitrogen, Oxygen and Fluorine ALWAYS have 8 electrons around them ¨ Hydrogen will always have a total of two electrons from its one single bond ¨ Boron can have fewer than 8 electrons but never more than eight. ¨ Nonmetallic elements in periods beyond the second period (P, S, Cl, Se, Br, I) usually have 8
electrons around them but can have more.
5. Insert pairs of electrons between all pairs of atoms that are to be bonded together.
6. Place any remaining electrons on peripheral atoms as unshared pairs, starting with the most electronegative such atom.
¨ Fill this atom up to an octet. ¨ Then proceed to the next most electronegative, and so on. ¨ Remember that hydrogens can only have two electrons, and so cannot have any unshared pairs.
7. If electrons remain unused, place them on the central atoms as unshared pairs, again beginning with the most
electronegative atom. Fill that atom to an octet. Then proceed to the next most electronegative, and so on.
v Type I: Ionic Compounds -Cation is always named first and the anion second, or the metal is named first, and the nonmetal is named second. -The cation, or metal takes its name from the name of the element -The anion, or nonmetal takes the first part of the element name (the root) and the ending is changed to -ide.
v Type II: Ionic Compounds w/Transition Metal -The same rules apply as type I compounds -Roman numeral is assigned for the transition metal that indicates its oxidative state
-The first element in the formula is named first as its elemental name -The second element is named as thought it were an anion; therefore, the ending is changed to -ide. -Prefixes are used to denote the number of atoms present.
§ All compounds containing alkali metal cations and the ammonium ion are soluble. § All compounds containing NO3
-, ClO4-, ClO3
-, and C2H3O2- anions are soluble.
§ All chlorides, bromides, and iodides are soluble except those containing Ag+, Pb2+, or Hg22+.
§ All sulfates are soluble except those containing Hg22+, Pb2+, Sr2+, Ca2+, or Ba2+.
§ All hydroxides are insoluble except compounds of the alkali metals, Ca2+, Sr2+, and Ba2+. § All compounds containing PO4
3-, S2-, CO32-, and SO3
2- ions are insoluble except those that also contain alkali metals or NH4
+.
¨ Solubility/Dissociation
• Strong Electrolytes: Strong electrolytes consist of substances that ionize completely (approx. 100%) in water.
• These substances are strong acids, strong bases, and soluble ionic salts. • Weak Electrolytes: Substances that ionize approx. 1-5% in water form weak electrolytes. These
substances are weak acids, weak bases, and slightly soluble ionic salts. • Non-Electrolytes: Substances that dissolve in water but are considered polar covalent, like table
sugar and some alcohols, are considered non- electrolytes.
¨ Rules of Naming Acids - If the anion does not contain oxygen, the acid is named with the prefix hydro- and the suffix -ic
attached to the root name for the element. - When the anion contains oxygen, the acid name is formed from the root name of the central element
of the anion or the anion name with a suffix of -ic, or -ous (refer to polyatomic ions table). - When the anion ends in -ate, the suffix -ic is used. - When the anion ends in -ite, the suffix -ous is used.
¨ Unit 02: Atomic Structure & the Periodic Table § Periodic Trends §
1. Predict the order of increasing electronegativity in each of the following groups of elements. A. C, N,O B. S, Se, Cl C. Si, Ge, Sn D. Tl, S, Ge
2. Predict the order of decreasing electronegativity in each of the following groups of elements. A. Na, K, Rb B. B, O, Ga C. F, Cl, Br D. S, O, F
3. Arrange the following groups of atoms in order of increasing atomic radii. A. Te, S, Se B. K, Br, Ni C. Ba, Si, F
4. Arrange the following groups of atoms in order of decreasing atomic radii. A. Rb, Na, Be B. Sr, Se, Ne C. Fe, P, O
5. In each of the following sets, which atom or ion has the smallest atomic radii? A. H or He B. Cl, In, or Se C. Nb, Zn, or Si D. Na-, Na, Na+
6. In each of the following sets, which atom or ion has the smallest ionization energy? A. Ca, Sr, or Ba B. K, Mn, or Ga C. N, O, or F D. S2-, S, or S2+ E. Cs, Ge, Ar
7. Order the atoms in each of the following sets from the least exothermic electron affinity (smallest electron
affinity) to the most. A. S, Se B. F, Cl, Br, I C. N, O, F D. Al, Si, P
§ Indicate which type of chemical reaction (synthesis, decomposition, single-displacement, double displacement or combustion) is being represented for the following chemical equations.
1. Given 2.98g of copper (II) sulfate, CuSO4, calculate the number of molecules.
2. A sample of barium ferrate, BaFeO4 was crystallized through a process of vacuum filtration. 0.95mg of the compound was extracted. Determine the number of molecules of barium ferrate. (1g = 1000mg)
3. A sample contains 1.054x1021 molecules of cadmium arsenide, Cd3As2. Calculate the mass in grams of the sample.
§ Percent Composition
4. Calculate the percent composition for ammonium phosphate (NH4)3PO4.
5. Calculate the percent composition for magnesium nitrate, Mg(NO3)2.
§ Empirical and Molecular Formula
6. An acid commonly used in the automotive industry is shown to be 31.6% phosphorous, 3.1% hydrogen, and 63.5% oxygen. Determine the empirical formula of this acid.
7. A white powder is analyzed and found to contain 43.64% phosphorus and 56.36% oxygen by mass. The compound has a molar mass of 283.88g/mol. What are the compound’s empirical and molecular formula.
8. Caffeine, a stimulant found in coffee, tea, and chocolate, contains 49.48% carbon, 5.15% hydrogen, 28.87% nitrogen, and 16.49% oxygen by mass and has a molar mass of 194.2g/mol. Determine the molecular formula of caffeine.
- Given percentage = Mass - Convert mass to moles for each element. Contemplate atomic weight - Divide each value by the smallest mole value - Approximate each value to a whole integer (i.e. 1.96 =2, 3.14 = 3) However, if you cannot round
(i.e. 2.5, 3.5, etc.) multiply all the moles by an integer to obtain a whole number. - Write out your empirical formula - The molecular formula can only be solved if the problem gives you molecular weight - To calculate molecular formula, first solve for the empirical formula mass - Divide the given molecular formula weight by the empirical formula weight to obtain a whole number - Multiple the empirical formula by the whole number to obtain the molecular formula.
1. Sodium hydroxide reacts with carbon dioxide as follows:
NaOH(s) + CO2(g) à Na2CO3(s) + H2O(l)
A. Balance the chemical equation B. Which is the limiting reactant when 1.85 mole NaOH and 1.00 mole CO2 are allowed to react to
produce sodium carbonate, Na2CO3 C. How many grams of sodium carbonate can be produced?
2. The fizz produced when an Alka-Seltzer tablet is dissolved in water is due to the reaction between sodium bicarbonate (NaHCO3) and citric acid (H3C6H5O7):
NaHCO3(aq) + H3C6H5O7(aq) à CO2(g) + H2O(l) + Na3C6H5O7(aq)
A. Balance the chemical equation B. In a certain experiment, 1.00g of sodium bicarbonate and 1.00g of citric acid are allowed to react.
Determine how many grams of carbon dioxide can be formed. C. What is the limiting and excessive reagent?
3. Sodium carbonate reacts with nitric acid to produce sodium nitrate, carbon dioxide and water.
Na2CO3(aq) + HNO3(aq) à NaNO3(aq) + H2O(l) + CO2(g)
A. Balance the chemical equation B. If 30g of sodium carbonate react with excess nitric acid, how many grams of sodium nitrate should
be produced? C. If the final mass of sodium nitrate is 45.5 grams, what is the percent yield?
4. If 23.2 grams of butane (C4H10) and 93.7 grams of oxygen (O2) are available in the following reaction
C4H10 + O2 à CO2 + H2O
A. Balance the equation for the reaction B. Determine which reactant is the limiting reagent C. Calculate the theoretical yield of CO2 in grams D. If the actual yield of CO2 is 67.4 g CO2, what is the percent yield
5. Iron reacts with oxygen (diatomic gas) to form iron (III) oxide.
A. Write and balance a chemical reaction B. Identify the type of reaction C. If 100 grams of iron react with excess oxygen, how much rust should be produced? D. If the actual mass if iron (III) oxide at the end of the experiment is 130 grams, what is the percent
yield?
Conversions and Formulas
Given Mass (Grams)
Mass (g) A Given
1 (mol) A Mole B Balanced Equation
Molar Mass (g) B
Molar Mass (g) A Mole A Balanced Equation
1 (mol) B
Given Mass (Kilograms)
Mass (kg) A Given
1000g Conversion
1 (mol) A Mole B Balanced Equation
Molar Mass (g) B
1kg Conversion
Molar Mass (g) A Mole A Balanced Equation
1 (mol) B
Given Mass (Milligrams)
Mass (mg) A Given
1g Conversion
1 (mol) A Mole B Balanced Equation
Molar Mass (g) B
1000mg Conversion
Molar Mass (g) A Mole A Balanced Equation
1 (mol) B
Given Moles
Mol A GIVEN
Mole B Molar Mass (g) B
Mole A 1 (mol) B
Percent Yield: Actual Yield / Theoretical Yield x 100%
1. What is the pressure of a gas if you compressed the gas from its original 500 mL at 3.4 torr to a volume of 302 mL?
2. What pressure is required to compress 196.0 liters of air at 1.00 atmosphere into a cylinder whose volume is 26.0 liters?
¨ Charles’ Law
3. At what temperature will a gas be at if you allow it to expand from an original 456 mL at 65 °C to 3.4 L?
4. A container containing 5.00 L of a gas is collected at 100 K and then allowed to expand to 20.0 L. What must the new temperature be in order to maintain the same pressure?
¨ Gay-Lussac’s Law
5. If at 1.00 atm of pressure water boils at 100 ºC, at what temperature would water boil if the pressure is 600. torr? (This shows why food doesn't cook well at higher elevation)
6. At 1 atm of pressure water boils at 100 ºC, if the sample was placed under 2 atm of pressure, what would be the temperature? (This would be like a pressure cooker).
¨ Avogadro’s Law
7. If 2.00 mol of gas occupies 4.50 L at STP. How much of the same gas will occupy 3.00 L at STP?
8. If 3.25 mol of argon gas occupies a volume of 100. L at a particular temperature and pressure, what volume does 14.15 mol of argon occupy under the same conditions?
¨ Combination Gas Law
9. You have a gas at 453 mm Hg with a volume of 700 mL and a temperature of 25 °C, what will the temperature of the gas be, if you change the pressure to 278 mm Hg and a volume of 1200 mL?
10. A gas balloon has a volume of 106.0 liters when the temperature is 45.0 °C and the pressure is 740.0 mmHg. What will its volume be at 20.0 °C and 780.0mmHg?
11. If you have 0.56 moles of an ideal gas at 87 °C and a pressure of 569 torr, what volume will the gas take up?
12. 2.50 grams of XeF4 is introduced into an evacuated 3.00-liter container at 80. °C. Find the pressure in atmospheres in the container.
¨ Gas Stoichiometry & Ideal Gas Law
13. Calcium carbonate decomposes at high temperatures to form carbon dioxide and calcium oxide:
___CaCO3(s) à ___CO2(g) + ___CaO(s)
If 2.56g of calcium carbonate CaCO3 decomposes under STP conditions, what volume of carbon dioxide is formed at the end of reaction?
14. Ethylene burns in oxygen to form carbon dioxide and water vapor:
___C2H4(g) + ___O2(g) à ___CO2(g) + ___H2O(g) How many liters of water can be formed if 1.25g of ethene C2H4 are consumed in this reaction under STP conditions?
¨ Density Calculations
15. Calculate the density of sodium bicarbonate (NaHCO3) in grams per liter (g/L) at 0.760 atm and 43ºC.
16. What is the molar mass (g/mol) of an unknown compound with a density of 1.84g/L, at 1.284 mmHg and 227ºC?
v Equations
¨ Boyle’s Law P1V1 = P2V2 ¨ Charles’ Law V1/T1 = V2/T2 ¨ Gay-Lussac’s Law P1/T1 = P2/T2 ¨ Avogadro’s Law V1/n1 = V2/n2 ¨ Ideal Gas Law PV=nRT
P (Pressure) V (Volume) n (Number of moles) R (Gas Constant 0.0821L.atm/K.mol) T (Temperature)
Name or give the chemical formula for each of the following compounds. State whether they are soluble (will dissolve) or insoluble (will not dissolve) in solution. Use the solubility rules!
For each of the following combinations of reactants, do the following: - Predict possible products with phase labels - Write the balanced chemical equation. If no visible reaction occurs, write NR. If a
precipitate form, give the formula and name of the precipitate.
1. a. What is the solubility of KCl at 5°C? _______
b. What is the solubility of KCl at 25°C? _______
c. What is the solubility of Ce2(SO4)3 at 10°C? _______
d. What is the solubility of Ce2(SO4)3 at 50°C? _______
2. At 90°C, you dissolved 10 g of KCl in 100. g of water. Is this solution saturated or unsaturated?
3. A mass of 100 g of NaNO3 is dissolved in 100 g of water at 80ºC. a) Is the solution saturated or unsaturated? ______________________________ b) As the solution is cooled, at what temperature should solid first appear in the solution? Explain. 4. Use the graph to answer the following two questions:
Which compound is most soluble at 20 ºC? ________ Which is the least soluble at 40
ºC? ________ 5. Which substance on the graph is least soluble at 10°C? __________
3. Calculate the pH of each of the following solutions of a strong base in water.
A. 0.10M NaOH
B. 0.03M KOH
C. 0.25M Ba(OH)2
4. Fill in the missing information in the following table.
pH pOH [H+] [OH-] Acidic, Basic, or Neutral
Solution A
6.88
Solution B
8.4x10-14M
Solution C
3.11
Solution D
1.0x10-7M
5. What is the hydronium ion [H3O+] concentration in a fruit juice solution that has a pH of 3.3?
6. A commercial window-cleaning liquid has a pH of 11.7. What is the hydroxide ion [OH-] concentration?
7. A solution is prepared by dissolving 22.5g of HNO3 in enough water to make 1.0L of solution. Determine the molar concentration of the acid and calculate the pH.
8. A solution is prepared by dissolving 17.6g of hydrobromic acid in 500mL of water. Determine the molar concentration of the acid and calculate the pH.
9. A solution is prepared by dissolving 5.54g of sodium hydroxide in enough water to make 1.5L of solution. Determine the molar concentration of the base and calculate the pH.
10. Glacial acetic acid, pure HC2H3O2, has a concentration of 17.54M. If 85.5mL of glacial acetic acid is diluted to 250mL, what is the acetic acid concentration? (Use M1V1 = M2V2)
11. Sodium hydroxide, NaOH, has a concentration of 4.5M. If 45.2mL of sodium hydroxide is diluted to 700mL, what is the sodium hydroxide concentration? Calculate the initial and final pH of sodium hydroxide.
§ Chemistry 9th. Steven S. Zumdahl and Susan A. Zumdahl (will use in AP Chemistry) § Chemistry 9th. Raymond Chang § Bozeman Science (YouTube) § Professor Dave (YouTube) § Crash Course Chemistry (YouTube) only recommended as a quick review
All AP Chemistry material will be provided within the first 3week grading period of the 1st 9weeks. Please make a good effort to revisit all course material that was taught in Pre-AP Chemistry and work on the summer packet. The summer packet only covers material that was taught in Pre-AP Chemistry. I will upload an answer key for the summer packet sometime in August. My email is [email protected]. Feel free to reach out to me if you if you need any help or have any questions about AP Chemistry.
