Chemistry FINAL: CONTENT Review Packet Name: Period: Date: Classification of Matter & Chemical/ Physical Changes 1. are substances that are made up of two or more elements which are chemically combined 2. is made from two or more substances that are physically combined 3. are substances that are made up of only one type of atom 4. is anything that has both mass and volume 5. is a solid, produced by a reaction that separates a solution. 6. A consists of a solute in a solvent 7. List 2 lab safety precautions. 8. Identify the following pieces of equipment: 9. 10. Which piece of equipment is used as a cover for the evaporating dish? 11. Where is the hottest part of the Bunsen burner flame? 12. How do you take a meniscus reading? 1
13. Define what a heterogeneous mixture and a homogeneous mixture are and provide an example of each. 14. Suppose that during a reaction a chemistry student touches the beaker and observes that it feels COLD. The student should conclude that the chemical reaction is. a) Why? 15. Suppose that during a reaction a chemistry student touches the beaker and observes that it feels HOT. The student should conclude that the chemical reaction is. a) Why? 16. Classify the following as an element, compound, homogeneous mixture or heterogeneous mixture. a) Table salt b) Aluminum c) Dirt d) Sugar water 17. Define the terms chemical and physical change 18. Classify the following as chemical change or physical change. a) Iron rusting b) Ice melting c) Evaporation d) HCl reacting with Mg to create H 2 gas 19. Another name for a homogeneous mixture is a(n) 20. What are four indicators of a chemical reaction? 2
Measurement & Density 1. What is the difference between a qualitative measurement and a quantitative measurement? 2. ml is a unit that measures 3. cm is a unit that measures 4. g is a unit that measures 5. mm Hg is a unit that measures 6. K is a unit that measures 7. What is the difference between accuracy and precision? 8. Try these conversions: a. 5000 mg = g b. 1L = ml 9. Write the following numbers in scientific notation. a. 74,600 km b. 0.0443 km 10. Write the following numbers in ordinary notation. a. 8.5 x 10 7 km b. 9.67 x 10-4 km 11. Define density. 12. Find the density of an unknown solid given that a 5.62g sample occupies 2.35 cm 3 3
Atomic Structure, Quantum Numbers, and Electron Notation 1. proton a. the total number of protons and neutrons in the nucleus of an atom 2. atom b. the weighted average mass of the atoms in a naturally occurring sample of an element 3. mass number c. 1/12 the mass of a carbon-12 atom 4. atomic mass unit d. the number of protons in the nucleus of an element 5. electron e. atoms with the same number of protons but different number of neutrons 6. mole f. negatively charged subatomic particle 7. average atomic mass g. the smallest particle of an element that retains the properties of that element 8. atomic number h. a counting unit; 6.022 x 10 23 9. neutron i. positively charged subatomic particle 10. isotopes j. subatomic particle with no charge 1. Suppose an atom has 27 protons and 32 neutrons. What is the mass number? 2. Suppose an atom has 29 protons and 35 neutrons. How many electrons does it have? 3. Suppose an isotope has a mass number of 27 and an atomic number of 13. a. Write the hyphen notation for this isotope. b. Write the nuclear symbol for this isotope. 4. What is electromagnetic radiation? Provide examples. 5. What is the BEST way to describe the nature of an electron? 6. Fill in the chart below. What do the four quantum numbers indicate? Quantum Number: Indicates: Principal Quantum Number Angular Momentum Quantum Number Magnetic Quantum Number Spin Quantum Number 4
7. One orbital can hold a MAXIMUM of electrons. 8. What it the correct electron configuration notation and orbital notation for Na and Al Diagonal Rule: 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 6 4d 10 4f 14 5s 2 5p 6 5d 10 5f 14 6s 2 6p 6 6d 10 7s 2 7p 6 Periodic Table 1. are elements that show the properties of both metals and nonmetals. 2. How are elements arranged on the Periodic Table? 3. What is electronegativity? 4. What is the most electronegative element? 5. Write whether an increase or decrease of a trend occurs as you go across the Periodic Table from left to right in a period, and when you go from top to bottom in a group. Trend Across a Period Down a Group 1st Ionization Energy Atomic radius Electronegativity 5
6. What name is given to the following? a. Group #1 Elements: b. Group #2 Elements: c. Groups #3-12 Elements: d. Group #18 Elements: e. Group #17 Elements: 7. Use the elements from part of period three, listed below, to answer the following questions. Write the symbol of the element described in the blank provided. (1pt.each) 11 Na 22.99 12 Mg 24.30 13 Al 26.98 14 Si 28.09 15 P 30.97 16 S 32.07 17 Cl 35.45 18 Ar 39.95 a). an element with 15 electrons b). the most reactive metal c). the most reactive nonmetal d). an element with 12 protons i). an element in the same group as potassium (refer to wall chart to find potassium-k ) j). an element that forms an ion with a charge of 2 k). an element with a atomic number of 14 e). an element with the atomic number of 17 f). the MOST stable element h). an element in the halogen group Bonding 1. A bond where there is a sharing of electrons: a) mutual b) electronegativity c) covalent d) ionic 2. A bond where there is transferring of electrons so an attraction between positive and negative ions is established: a) mutual b) electronegativity c) covalent d) ionic 3. What is the difference between nonpolar covalent and polar covalent bond type? 4. How many valance electrons are found in the following groups? a. Group 1 b. Group 2 c. Group 13 d. Group 14 e. Group 15 f. Group 16 g. Group 17 h. Group 18 6
5. Circle the lone pair electrons in the following dot formula of water. 6. With respect to bonds formed between the following pairs of atoms: Determine the electronegativity difference. SHOW WORK! Determine the probable bond type (ionic, polar covalent, or nonpolar covalent). Assign partial charges to atoms that are part of a polar covalent bond. Pairs of Atoms Electronegativity Difference Probable Bond Type Partial Charge (if polar-covalent) H and H S and O K and Br 7. A molecule is a group of atoms held to together by bonds. 8. What is the difference between a cation and an anion? 9. List the naturally occurring diatomic elements? 10. What is the molecular geometry of BF 3 11. What is the molecular geometry of CH 4 12. What is the molecular geometry of CO 2 7
Writing Formulas & Chemical Equations Name the following: Write formulas for the following: Pb 3 (PO 4 ) 2 Zinc chloride Cr 2 O 3 Aluminum hydroxide 1. What are the four types of chemical reactions? Write a simple equation to represent each type of reaction. 2. The empirical formula is the a) simplest whole number ratio b) actual composition c) actual mole ratio 3. A compound that contains water is called a(n) a) hydrate b) hydroxide c) anhydrate d) none of these 4. A calculation that determines the mass/mass or mass/mole relationship is a) quantitative b) qualitative c) stoichiometry d) molar mass 5. The substance that comes through the funnel during the filtering process is the a) filter residue b) the filter paper c) the filtrate d) the decant 6. Heating until you get a matching weight is referred to as a) being careful b) following directions c) heating to accuracy d) heating to precision 7. A hydrate that has been heated until dry is a a) decomposed b) anhydrous c) hydrated d) oxidized 8. An insoluble product that forms from the reaction of two liquids is called a) solute b) precipitate c)flakes d) residue 9. To balance an equation, one uses a) subscripts b) roman numerals c) coefficients d) stoichiometry 10. The law of conservation of mass states that the mass of the reactants must be the mass of the products a) greater than b) less than c) equal to d) cannot be determined 11. Tarring (re-zeroing) the scale is taking into consideration the a) mass of a substance b)mass of an empty container c) mass of the reactants d)mass of the products 12.When the equation Al 2 (SO 4 ) 3 + Ca(OH) 2 Al(OH) 3 + CaSO 4 is correctly balanced, the coefficient for Ca(OH) 2 is a) 3 b) 2 c) 1 d) 4 13. Using the equation in #12, what is the mole ratio between Al 2 (SO 4 ) 3 and CaSO 4? a) 2:3 b) 1:2 c) 2:3 d) 1:3 8
14. A substance that increases the rate of a reaction without itself being changed is a(n). a) precipitate b) catalyst c) inhibitor d) none of the above 15. What is a reversible reaction? How is a reversible reaction represented? 16. Label the products, reactants and yields arrow. CH 4 + 2O 2 CO 2 + 2H 2 O 17. Fill in the correct symbol that would used when writing a chemical equation based on the meaning provided. Symbol Meaning yields ; indicates result of a reaction Indicates a reversible reaction A reactant or product in the solid state Alternative to (s); used only for a precipitate (solid) falling out of solution A reactant or product in the liquid state A reactant or product in aqueous solution (dissolved in water) A reactant or product in the gaseous state Alternative to (g); used only for a gaseous product Reactants are heated Pressure at which the reaction is carried out, in this case 2 Temperature at which reaction is carried out, in this case O C A catalyst is used to speed up the reaction rate, in this case MnO2 would be used to speed up the reaction rate. 9
Kinetic Theory, Solids, Liquids, Gases, Phase Changes 1. Name the 3 parts of the kinetic molecular theory. 2. How does temperature affect the kinetic energy of a substance? 3. List the values for STP. 4. If the volume of a gas is decreased, then gas pressure will. 5. If the volume of a gas is increased, then gas temperature will. 6. How does evaporation differ from sublimation? 7. Complete the chart to name the four phases of matter and compare their volumes, shapes, and average kinetic energy Phase Shape Volume Average Kinetic Energy Liquid definite definite definite Not definite Very FAST (violent) Not definite Fast 8. What is the difference between a crystalline solid and an amorphous solid? 9. Which of the 4 phases of matter is considered to have the most significant intermolecular forces? a) The least significant intermolecular forces? 10. A liquid will boil when its equilibrium vapor pressure EQUALS. 11. When a system at equilibrium is disturbed by application of a stress, it attains a new equilibrium position that minimizes the stress. This is known as. 10
12. Use the phase diagram below to answer the following questions: a. The area on the graph that represents the gas phase is: b. The area on the graph that represents the liquid phase is: c. The area on the graph that represents the solid phase is: d. When does the triple point occur? e. What happens to a substance as you move from point A to point B? A B C Mixtures: Solutions, Suspensions & Colloids 1. Define the following terms: solution, colloid, suspension 2. Suppose 15 g of sugar is dissolved in 1000 g of water. The sugar is the and the water is the 3. Explain what the phrase like dissolves like means and give an example 4. Explain the difference between an unsaturated, saturated and a supersaturated solution. Define these terms! 5. What 3 things can increase the rate at which a solute will dissolve? 7. How could you identify a colloid? 11
Acids/Bases/Salts 1. Describe acids and bases according to the theory of Arrhenius. 2. Explain how blue and red litmus paper determine if a substance is an acid, base or salt. 3. Explain the process of neutralization using a chemical equation. 4. Define ph and draw the ph scale showing the locations of acids, bases, and neutral substances. 5. What happens during a splint test? How do you know O 2, CO 2 or H 2 is produced? Nuclear Chemistry 1. Name and describe each of the 3 types of radiation. 2. What is a half-life? 3. Write the nuclear symbol and the hyphen notation for an isotope of phosphorus with a mass number of 32. 12
4. What is the difference between fission and fusion? 5. Balance the following nuclear reactions. Things to study: o ALL Notes o Labs o Review Guides (Content & Math) o Mrs. Hostetter s website (review PowerPoints) o www.epasd.org/hostetter The ACCUMULATIVE Final Is Worth 100 points: o 100 multiple choice questions STUDY & GOOD LUCK!! 13
Chemistry FINAL: MATH Review Name: Period: Date: 1. Calculate the molar mass of iron (II) phosphate. 358g 2. Find the percent composition of sodium carbonate. %Na= 43%, %C = 11%, %O= 45% 3. Convert 6.25 g of sodium chloride to moles. 0.108mol 4. Convert 0.500 mol of potassium iodide to grams. 83g 5. What is the empirical and molecular formula of a compound with the following percent composition: P = 26.7 %, N = 12.1 %, Cl = 61.2 %. The molar mass of the compound is 695 g/mol. EF = PNCl2, MF = P6N6Cl12 14
6. How many grams of sodium phosphate are needed to completely react with 25.0 g of barium chloride? 13.205g 7. Suppose 100 ml of solution contains 5.92 g of dissolved calcium chloride. Calculate the molarity. 0.538M 8. Suppose 9.56 g of sodium hydroxide is dissolved in 150. g of water. Calculate the molality. DID NOT COVER!! 9. Write the electron configuration notation, orbital notation, and electron dot symbol for arsenic. 10. Suppose a gas is collected in a 145 ml flask where the temperature of the gas is 23 C and the pressure of the gas is 785 mm Hg. Calculate the volume of the gas at STP. 138.132mL 15
11. Suppose 25.0 ml of a gas is at 24 C. Calculate the temperature at which the gas will occupy 30.0 ml assuming constant pressure. 356.4K 12. Suppose the length of a piece of string is measured to be 33.45 cm. Calculate the percent error if the string is actually 34.50 cm long. 3.04% 13. Suppose 4.25 g of sodium chloride is collected after a chemical reaction. Through stoichiometry calculations it is determined that the theoretical yield of sodium chloride is 5.00 g. Calculate the percent yield. 85% 14. Suppose an unknown radioactive substance with a mass of 120 g has a half-life of 6 months. How much of the substance will remain after 2 years? 7.5g 15. What is the volume of 10.5 g of copper? (The density of copper is 8.94 g/cm 3.) 1.173cm 3 16