1. Dimensional Analysis: convert the following values a. 47,340 cm to m Unit 1: Chemistry Matters b. 40.64 km to m c. 58,700 ml to L 2. Calculate the number of significant figures a. 0.0210 b. 3.6056 c. 0.0006 3. Solve the following problems with the correct number of significant figures a. 319.15-32.614 = b. 104.630 + 27.08362 + 0.61 = 4. List the five indicators of a chemical change: a. b. c. d. e. 5. Define precipitate. How can you determine if something is a precipitate or not? 6. Label each of the following as a physical or chemical change a. A tire is inflated with air. c. Water is added to red solution, it turns pink. b. Food is digested in the stomach. d. Water is heated and changed into steam.

7. Define the following terms. Give an example of each Term Definition Example Element Compound Mixture Pure Substance 8. Complete the following chart Subatomic Particle Charge Mass (amu) Location Proton Neutron Electron 9. How many protons, neutrons and electrons are in a Cu-64 2+ ion? Write this ion in nuclear notation. 10. Protons + Neutrons = 11. Atomic Mass - = Neutrons 12. What is an ion? What is the difference between a cation and anion? 13. Isotopes differ in the number of. 14. Fill out the chart below for each of the three isotopes Isotope Electrons Protons Mass Number Atomic Number Neutrons 238 U 16 O 2-23 Na +1 15. What is the difference between the isotopes U-240 and U-238?

Unit 2: The Electron 1. Give the symbol for the following a. Alpha b. Beta c. Gamma 2. Which type of decay listed above is the most energetic? The least energetic? 3. Balance the following nuclear reactions and label alpha decay or beta decay b. 220 Rn 4 He + b. 216 Po 0-1e + 4. Tritium has a half-life of 12.3 years. How long would it take for a 40.0g sample to decay down to 1.25g? 5. Fe-61 has a half-life of 6.00min. Of a 100.0g sample, how much will remain after 18.0min? 6. After 20.0 days, a 120kg sample of Bi-210 decays down to just 7.5kg. What is its half-life? 7. Given an initial 500g sample of H-4, how much remains after 3 half-lives? 8. What is the difference between group and period on the PToE? 9. What defines the chemical properties of elements? 10. Where are the following elements located on the periodic table: a. Metals b. Nonmetals c. Metalloids 11. Label the following elements as a metal, nonmetal, or metalloid a. Oxygen b. lead c. silicon d. magnesium e. boron f. neon

12. On the periodic table below, label the following a. Group number b. Period number c. Alkali metals d. Alkaline earth metals e. Halogens f. Noble gases g. Transition elements h. Valance electrons i. Charges j. Sublevels (which blocks)

13. Draw the orbital diagram for the following: a. Sulfur b. Nickel 14. Write the electron configuration (both standard and noble gas notation) for the following elements a. Boron b. Copper c. Aluminum 15. Identify the element represented by the configurations below d. 1s 2 2s 2 2p 4 b. [Ar]4s 2 3d 10 4p 3 16. Determine the number of valence electrons in the following configurations e. 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 5 b. [Kr]5s 1 17. Define atomic radius. 18. Put the following in order by decreasing atomic radius: Al, Na, S, K and explain why. 19. Define ionization energy. 20. What is the general trend for ionization energy? Explain the reasoning. 21. Put the following elements in order of decreasing ionization energy: Rb, Al, S, Mg 22. Define electronegativity. 23. What is the general trend for electronegativity? Explain the reasoning. 24. Put the following elements in order of increasing electronegativity energy: F, B, N, Li

25. Label each diagram with the element (name and symbol) they represent. a. b. 26. Describe the difference between the ground state and an excited state. How do electrons move from one to the other? 27. (Circle the correct answer) Light is emitted when an electron moves from the ground/excited state to the ground/excited state. 28. What is the wavelength of a photon emitted when the electron falls from the third energy level to the second energy level? What type of electromagnetic radiation is it? 29. If an electron in a hydrogen atom moves from n=5 to n=2, what color of light is emitted? 30. If an electron in a hydrogen atom moves from n=1 to n=3, what color of light is emitted? 31. If you have light with a wavelength of 600nm, what color is it? 32. In order to emit red light, what energy level transition must occur? Unit 3: Bonding and Nomenclature 1. Describe metallic bonds. 2. Give the ionic charge for the following groups a. Group 1 b. Group 2 c. Group 13 d. Group 15 e. Group 16 f. Group 17 3. What types of elements form ionic compounds? 4. Give four properties of ionic compounds and explain why they have these properties. 5. Write the formula for the following: a. Magnesium fluoride d. Calcium nitride b. Sodium carbonate e. Ammonium phosphate c. Copper (III) bromide f. Tin (IV) oxide 6. Write the names for the following formulas: a. FeS d. VBr 3 b. NH 4NO 2 e. Ba(OH) 2 c. Al 2S 3 f. (NH 4) 3P

7. Describe covalent bonding. 8. What type of elements are typically used in a covalent bond? 9. Write the formula for the following: a. Carbon tetrachloride c. Dinitrogen pentoxide e. Sulfurous acid b. Carbon monoxide d. heptanitrogen difluoride f. Hydrochloric acid 10. Write the names for the following formulas a. PF 5 c. NO e. CO 4 b. N 2O 3 d. H 2P 6 f. NO 3 11. List 4 properties of covalent bonds and explain why they have these properties. 12. Explain why hydrogen bonds are stronger than dipole-dipole forces which are stronger than dispersion forces. 13. What are the kinds of intermolecular forces? When do they occur? 14. Determine the type of bond of the following: a. N-Br b. Cl 2 c. CaF 2 d. CO

Unit 4: Chemical Reactions 1. List the five types of reactions and their general equations. a. b. c. d. e. 2. For the following equations, identify the type and balance a. BaS + PtF 2 BaF 2 + PtS Type: b. N 2 + H 2 NH 3 Type: c. CH 4 + O 2 CO 2 + H 2O Type: d. K + Cl 2 KCl Type: 3. Use activity series to predict whether the following reactions take place. a. Cu + 2AgNO 3 Cu(NO 3) 2 + 2Ag Yes reaction occurs or No reaction b. 2KBr + I 2 2KI + Br 2 Yes reaction occurs or No reaction 4. Write a complete balanced reaction for each of the following (make sure to use activity series if needed!) a. Magnesium oxide reacts with water b. Magnesium hydroxide reacts with potassium iodide c. Lithium reacts with copper (II) chloride d. Sodium chloride reacts with magnesium carbonate e. Chlorine reacts with potassium fluoride

5. Label the potential energy diagram below. Is this an endothermic or exothermic reaction? Unit 5: The Mole & Stoichiometry 1. What is molar mass? How do we find it? 2. What is the molar mass of one mole of silver? 5. Determine the number of liters in 3.26g of H 2O 6. How many grams are in 3.0 moles of H 2SO 4? 3. What is the mass of 7 moles of strontium? 4. How many particles are in 3.27 moles of Zn? 7. How many moles are in 84.2 grams of CO 2?

8. In the equation 2KClO 3 2KCl + 3O 2, how many moles of oxygen are produced when 3.0 moles of KClO3 decompose completely? 9. Butane, C 4H 10, burns in air to form carbon dioxide and water. If 12 grams of carbon dioxide are formed, how many grams of butane were burned? The equation: C 4H 10 + O 2 CO 2 + H 2O 10. Using the following reaction solve: 2FeCl 3 + 3H 2S Fe 2S 3 + 6HCl. How many particles of HCl is produced when 90.0 g of FeCl 3 reacts with excess H 2S? 11. What is the percent by mass of oxygen in sodium hydroxide? 12. Calculate the percent by mass of water in lithium chromate. 13. Analysis of a chemical indicates that it has a composition of 65.45% C, 5.45% H, and the rest is oxygen. The molar mass is found to be 165.0g/mol. Determine the empirical and molecular formulas.

1. Convert the following quantities Unit 6: Gas Laws a. 1000 torr to kpa b. 518 kpa to atm 2. If I have 4 moles of a gas at a pressure of 5.6 atm and a volume of 12 liters, what is the temperature? 3. A helium balloon contains 125L of gas at a pressure of 0.974atm. What will be the pressure if the volume is changed to 212L? 4. A sample of gas in a closed container at a temperature of 102 o C and a pressure of 3.2atm is heated. The new pressure was measured at 4.6atm. What was the final temperature in o C? 5. Oxygen gas from the decomposition of KClO 3 was collected over water. The total pressure was found to be 750.3torr. The water s pressure was measured to be 23.4torr. What is the pressure of oxygen? 6. An engineer pumps 45.2g of carbon dioxide gas into a cylinder that has a capacity of 23.5L. What is the pressure inside the cylinder at 23.8 o C? 7. A student collects 584L of oxygen at a temperature of 25.8 o C and a pressure of 683mmHg. How many grams of oxygen did the student collect? 8. If 3.2 moles of oxygen and 4.6 moles of helium occupy 15 atm of pressure, what are the partial pressures of oxygen and helium? 9. If you have 15 atm of oxygen, 30 atm of oxygen, and 45 atm of xenon, what is the total pressure of the container?

Unit 7: Heat and Energy/Thermochemistry 1. What two conditions can change to cause a phase change? 2. Explain the difference between heat and temperature. 3. Fill out the chart Phase Change Name of Phase Change Endothermic or Exothermic Solid to Liquid Liquid to Solid Liquid to Gas Gas to Liquid Solid to gas Gas to Solid 4. Answer the following questions using the phase diagram. a. What do A, B and C represent? b. What is the P and T of the triple point? c. What phase is the substance in at 30atm and 200 o C? d. At what T would the substance boil at 30atm? e. Draw an arrow representing sublimation. 5. What is the critical point? 6. Label the heating curve to the left a. include thermal energy, phase energy and the appropriate energy equations 7. Answer the following questions using the heating curve. a. What segment represents a solid being heated? b. What segment represents a liquid being heated? c. What segment represents a gas being heated? d. What segment represents melting? e. What segment represents condensation? 8. During the winter, you forget your 414g cup of tea outside and it freezes. How much energy was lost?

9. A 10g sample of ice heats up from -20C and then melts completely into liquid water. How much energy was transferred? 10. The temperature of a 94.8 grams sample of material increases from 37 o C to 79 o C when it absorbs 3584J of heat. What is the specific heat of this material? 11. What mass of liquid water at room temperature can be raised to its boiling point with addition of 24kJ of heat energy? 1. Define solution. What are the 2 parts of a solution? Unit 8: Solutions, Molarity and ph 2. Define molarity. What is the formula? 3. Solve the following molarity problems: a. What is the molarity of a solution containing 50 moles of NaCl dissolved in 5 L of water? b. How many grams must be dissolved to make 2.5L of a 5.0M solution of HCl? c. How many moles of LiOH are in 56.6L of a 2.56 M solution?

4. What is solvation? What are the three factors that affect the rate of solvation? 5. Use the solubility curve in your reference tables to answer the following questions a. How many grams of sodium nitrate can be dissolved in 100g of water at 40 o C? b. At what temperature can 22 grams of potassium chlorate be dissolved in 100g of water? c. What salt is the most soluble at 70 o C? d. What salt is the least soluble at 30 o C? e. If 100g of potassium iodide is dissolved at 10 o C, is the solution saturated, supersaturated, or unsaturated? 6. Label each of the following as an acid, base, both or neutral a. Can turn litumus paper blue b. Contains more hydronium ions than hydroxide c. Has a ph of 7 d. Feels slippery f. Tastes bitter g. Is an electrolyte h. Turns phenolphthalein pink i. Has a ph of 4 j. Proton donor in solution e. Tastes sour 7. What color would the litmus paper turn if the solution had a ph of 3.4? 8. A solution with a H + concentration of 1.00 x 10-7 M is said to be neutral. Why? 9. Determine the poh of a.0034 M HNO 3 solution. 10. Determine the ph of a 4.3 X 10-4 M NaOH solution