Chemistry Midterm Review 2006

Chemistry Midterm Review 2015 (Thematic) Scientists Tools Objective 1.1 – 1.3 1. Define chemistry. 2. What is the difference between an independent variable and a dependent variable? 3. Why is a constant important in some labs? 4. What is the difference between a hypothesis and a prediction? 5. What is the difference between a theory and a law? 6. State the difference between quantitative and qualitative data. Objective 1.4 1. The metric system is based on the power of ______2. What is the basic metric unit of length, mass, and volume? 3. What is the SI unit for volume? 1cm3 = 1 ______4. Define volume. What type of lab equipment measures approximate volume? 5. What are the prefixes for the metric units and what do they equate to? 6. Using dimensional analysis, convert 157 cg into g, convert 8.6 kg into cg, convert 100m into cm , and convert 7068 mm into m Objective 1.5: HONORS ONLY 1. What is the difference between accuracy and precision? 2. Define significant figures. What rule is used to count sig figs? 3. How many significant figures are in each: 5.730 x 108, 3000, 0.01552, 9009, 629.55, 1.777 x 10-3? 4. What is the answer expressed in proper significant figures for the following: 6.54 + 3.053 and 8.95 x .02? Objective 1.6 1. Place the following numbers into scientific notation a. 34,500 b. .098 c. 670 d. .000043 2. Take the numbers of scientific notation and place in ordinary notation a. 4.56 x 103 b. 5.78 x 10-2 c. 1.678 x 10-5

Antacids Objective 2.1 1. Define matter and list several examples 2. Matter can be broken down into ______and ______. 3. Differentiate characteristics of pure substances and mixtures. 4. Pure substances can be broken down into______and ______. 5. Differentiate characteristics of an element and a compound. State examples. 6. Mixtures can be classified as ______and ______7. Differentiate characteristics of a homogeneous and heterogeneous mixture. State examples 8. What is another name for a homogenous mixture? 9. By what means can you separate a compound? A mixture?

10. State whether each is a compound or element: Fe, CO, CaCl2, Hg, Co, argon, sodium chloride, manganese(II) nitrate, I2. 11. What is a chemical symbol? Write the symbols for the following. Mercury, gold, iodine, calcium, barium, tin, magnesium, phosphorus. 12. State whether the following are homogeneous or heterogeneous: a. rocky road ice cream b. orange juice with pulp c. fog d. salt solution d. Koolaid e. air without pollution f. brass g. sandy water 13. Differentiate characteristics of a solution, colloid and suspension. State examples Objectives 2.2-2.3

1. What are the characteristics of a covalent and ionic 4. Which of the following compounds are ionic? H2O,

compound? Na2O, CO2, CaS2, SO2, CaCO3. 2. What is a chemical formula? 5. Know the difference between a formula unit and a 3. ______are used to indicate the number of atoms in a molecule. chemical formula of a compound. 6. Define ion. What is the difference between a cation and an anion? 7. What is the chemical formula for the following; sulfide 13. Ag3N and Na3PO4 ion, sodium ion, fluoride ion, mercury (II) ion? 14. Write formulas for the following: potassium nitrate, 8. What are the names of the following ions; Ba2+, Al3+, lithium oxide, calcium phosphate, ammonium O2-, and Sn4+? carbonate, barium chloride, and copper(I) chloride. 9. Metals form ______ions and nonmetals form 15. What is a polyatomic ion? What are the names of the -3 -2 -2 -1 +1 ______ions. following ions: PO4 , SO3 , CO3 , NO2 , NH4 10. What are binary compounds? Ternary compounds? 16. Prefixes are used to indicate the number of atoms in 11. ______are used to indicate the charge of transition ______compounds. elements in a ternary ionic compound. 17. Name the following covalent compounds: CH4, N2O5,

12. Name the following chemical formulas; KClO3, BaSO4, SO3, CO, H2O, Cl2P3

MgBr2, Li2CO3, CoF2, 18. 19. Objectives 2.4 – 2.5 1. List 3-4 characteristics of both an acid and base 7. Know how to predict the products and balance 2. What is the characteristic ion of an acid? For a Base? neutralization (double replacement) reactions. 3. What is the ph range for an acid? For a base? a) H2CO3 + Fe(OH)2 → a. Which acid is stronger, a pH of 4 or a pH of 1? b) HClO4 + NaOH → b. Which base is stronger, a pH of 11 or a pH of 8? c) HBr + Ba(OH)2 → 4. What color does red litmus paper turn in an acid? In a 8. What is the difference between a strong and weak acid base? and base? What is the difference between a concentrated and dilute acid & base? 5. What color does phenolphthalein turn in an acid? In a base? a. Select the strong acids in the diagram 6. What are the products of neutralization? b. Select the dilute bases in the diagram 9. 10. A 12. B 14. C 16. D 11. 13. 15. 17.

    MOH   X  H OH MOH X H X M   OH M H H H H X H MOH MOH OH     X X  M X H  X M OH MOH X X X OH MOH H MOH H  H  X OH H M OH    H M M  X X MOH X H H

18. E 20. F 22. G 24. H 19. 21. 23. 25.

OH  HX M OH M OH MOH  HX  HX M X MOH X X OH  HX M M  HX OH   OH  MOH  M  H HX H OH  OH HX  M  HX OH OH  HX  M  M HX M H  MOH HX OH HX M 26. 27. Classify the substances as either Arrhenius acids or Arrhenius bases and give their chemical name (STANDARD ONLY: only need to name a-g) a) HCl d) HC2H3O2 g) H2CO3 b) H2SO4 e) HNO3 h) HNO2 c) NaOH f) Fe(OH)2 i) H3PO4 j) Objective 2.6 1. What is a chemical equation? a) Identify the reactants and the products of the following equation:

k) Al + O2 → Al2O3 2. What is the arrow called? Where would a catalyst’s chemical formula be placed? If Δ was above the arrow, what would that mean? 3. What are the 5 types of chemical reactions? 4. Identify the different types of chemical reactions. l) Examples:

b) FeCl3 + NaOH → Fe(OH)3 + NaCl e) Pb(NO3)2 + NaCl → PbCl2 + NaNO3

c) Al + O2 → Al2O3 f) Na + H2O → NaOH + H2

d) C2H2 + O2 → CO2 + H2O g) KClO3 → KCl + O2 5. Know how to predict products of a chemical reaction. Predict and balance the following equations(STANDARD ONLY complete a-g) 6. Example:

a) Al + N2 → e) Cl2 + NaBr → i) Mg(ClO3)2 

b) Li2CO3  f) CaS + FeCl2 → j) Mg + H2O(l) 

c) H2O → g) CH4 + O2 

d) K + HCl → h) SO3 + H2O → k) 7. What is an activity series chart? What type of reaction do you use it for? a) Using the activity chart, why can sodium replace hydrogen? 8. What are 5 indicators of a chemical reaction? 9. List the chemical formulas for the 7 diatomic molecules. 10. Know how to translate chemical equations and balance them appropriately. l) Example: a) ammonium chloride reacts with calcium hydroxide to form calcium chloride & nitrogen trihydride (ammonia) & water b) sodium oxide and water yield sodium hydroxide 11. What is the precipitate for the following reaction

a. CaS + FeCl2 → CaCl2 + FeS

b. AgNO3 + PbCl2  AgCl + Pb(NO3)2 12. HONORS ONLY:Write a complete ionic and net ionic equation for the 2 examples above.(2nd quarter topic, not on Fall 2015 midterm!!!) m) n) Objective 2.7 1. What is the law of conservation of q) 4. On the pathway to the left, label the activated mass? complex, activation 2. Differentiate between a subscript energy, reactants, and a coefficient? Which one is products, and enthalpy released used to balance a chemical equation? or absorbed by the reaction. Is this 3. To satisfy the conservation of mass, we balance reaction Endo or chemical equations. Exothermic? o) Example: r) a) __ FeCl3 + __ NaOH → __ Fe(OH)3 + __ NaCl s) 5. List the 5 b) __ Al + __ O2 → __ Al2O3 factors that affect the rate of a reaction and why the rate changes with each factor? c) __ C2H2 + __ O2 → __ CO2 + __ H2O t) 6. How does rate change if you increase the d) __ Pb(NO3)2 + __ NaCl → __ PbCl2 + __ NaNO3 concentration of the reactants? e) __ Na + __ H2O → __ NaOH + __ H2 u) 7. How does rate change if you increase the f) __ KClO3 → __ KCl + __ O2 surface area? Objective 2.8 p) v) 8. How does rate change if you decrease the 1. Define reaction rate temperature? 2. Explain the three criteria of the collision theory. w) 9. How does rate change if you add a 3. Draw in the activation energy for each line. Label the catalyst? pathways as “with catalyst” & “without the catalyst”. x) 10. Define activation energy? Will a Which line represents the faster reaction. reaction that proceeds faster have a higher or y) lower activation energy? z) 11. Give the following values in kilojoules: aa) a. Potential energy of reactants ab) b. Potential energy of products c. Heat of reaction(∆H) d. Activation energy of forward reaction e. Activation energy of the Reverse reaction f. Potential energy of the activated complex g. Is the overall forward reaction endo or exothermic? 13. What are the 2 reasons why an increase in temperature can increase the reaction rate? ac) Airbags ad) Objective 3.1 1. Define the 6 changes of state. Which ones are a. In what state is compound X at a temperature of 600oC exothermic and which ones are endothermic? and a pressure of 80 atm? 2. Create a chart comparing& contrasting

State the temperature and-30 characteristics of solids, liquids, and gases in b. F regards to density, compressibility, particle -40 pressure at the critical -50 size, shape, volume, kinetic energy, attractive -60

) -70 forces and movement point of compound X?C o (

e -80 r u 3. Explain how a solid melts into a liquid using t ◦ a -90 E What phase change doesr X undergo if a sampleD at 500 C

c. e kinetic energy in your explanation. p -100 m e

T ◦ Define boiling point? and 40 atm has its temperature-110 changed to 100 C? 4. -120 B C 5. What is the difference between normal boiling 12. -130 -140 point and boiling point? 13. -150 A 0 2 4 6 8 10 12 14 16 6. What happens to vapor pressure as temperature 14. 11. Know how to read a heatingTime or (min.) cooling increases? curve 7. What happens to boiling point as altitude a. State the boiling point of the substance: increases? 8. Who has a higher boiling point; Mount b. Which segment represents an increase in the kinetic McKinley or Charlotte, NC? energy of the liquid particles? 9. What 2 temperatures measure the same amount c. Which segment represents condensation? during a phase change of a pure solvent? d. How many minutes does it take from the time the 10. 10. Know how to read phase change graphs. substance begins to vaporize until it is completely 11. gaseous? 15. 16. 17. Objective 3.2 – 3.4 1. What is the difference between a physical and chemical dissolving in water, tarnishing silver, acid neutralizing a property? Give examples base, drying a wet towel 2. What is the difference between an intensive and 7. Define density. What is the equation? extensive property? Give examples. 8. Ice floats because it is more or less dense than water? 3. A chemical change is also known as a chemical 9. Put the 3 states of matter in order of increasing density. ______. 10. A copper penny has a mass of 3.1 g and a volume of .35 4. Name 5 buzz words that signify a physical change and 5 cm3. What is the density? that signify a chemical change. 11. A plastic ball has a volume of 19.7 cm3 and a density 5. Name five indicators/observations of a chemical of .8029 g/cm3. What is the mass? change(reaction). 12. The density of silicon is 2.33 g/cm3. What is the volume 6. Classify each as a physical or chemical change: food if its mass is 62.9g? spoils, water boils, nail rusting, baking bread, sugar 13. Objective 3.5 and 3.7 1. Know the 5 assumptions of the kinetic molecular theory. 2. Know the difference between an ideal gas and a real gas. What conditions does a real gas deviate from an ideal gas. 3. What is the difference between effusion and diffusion of c) As the volume of the gas decreases, the pressure of a gas? the gas will ______. 4. What happens to average kinetic energy when Kelvin d) As the moles of the gas decreases, the pressure of temperature doubles? the gas will ______.

5. Both methane gas(CH4) and hydrogen gas(H2) are in a e) As the volume of the gas decreases, the moles of container at 67 C. the gas will ______. 6. a. Which gas will have the greatest speed? 9. Define pressure. What are some common pressure 7. b. Do they have the same average kinetic units? energy of different average kinetic energy? 10. Know how to convert pressure units: 8. 4. Know the relationship between the variables of a gas a) convert .200 atm to mmHg a) As the temperature of a gas decreases, the volume b) convert 345.8 kPa to atm of a gas will ______. c) convert 760 mmHg to kPa b) As the temperature of a gas decreases, the pressure 11. What is standard temperature and standard pressure? of the gas will ______. 12. Not on Fall 2015 Midterm 13. Know how to solve problems using Boyle’s law, Charles law, Gay-Lussac, Combined, Ideal, Density and Molar mass using the Ideal gas law and Dalton’s law of partial pressure. a) A gas occupies a volume of 200. ml at 100. mmHg. What volume will the gas occupy at 300. mmHg? b) Air has a total pressure of 20.6 atm and contains carbon monoxide, oxygen, and nitrogen. If air is made up of 0.6 atm of carbon monoxide, 12.6 atm of oxygen, what would be the partial pressure of nitrogen? c) If a sample of gas occupies 15.9 L at 34 C, what will its volume be at 27 C if the pressure does not change? d) The volume of a sample of oxygen gas is 300.0 ml when the pressure is 1.00 atm and the temperature is 27.0 C. At what temperature would the volume change to 1.00 L and the pressure change to 0.500 atm? e) A sample of gas at 25.0 C has a volume of 11.0 L and exerts a pressure of 660.0 mmHg. How many moles of gas are in the sample? f) A sample of gas in a closed container at a temperature of 100. C and a pressure of 3.0 atm is heated to 300. C. What pressure does the gas exert at the higher temperature? g) How many grams of N2 are in a flask with a volume of 250. ml at a pressure of 3.0 atm and a temperature of 300 K? h) What is the density of silicon tetrafluoride gas at 72 C and a pressure of 144.5 kPa? 14. Know how to solve for the pressure of a gas if it is collected over water. a) A sample of nitrogen gas is collected over water at 23 C with a vapor pressure of .0278 atm. What is the pressure of the nitrogen gas if the atmospheric pressure is 785 mmHg? 15. 10.Use the law of combining volumes, Avogadro’s law, and molar volume to solve these problems. a) 3O2 → 2O3 Both gases are measured at the same temperature and pressure. How many O3 molecules are formed from the reaction of 24 O2 molecules? b) How many moles of O2 are required to make 24 moles of O3? c) How many liters of O3 are formed from 12 L of O2? 16. 17. Objective 3.6 18. 1. Know the representative particle for an ionic compound, covalent compound, element, and a diatomic molecule? 19. 2. Calculate the molar mass of

20. a) NaCl b) C6H12O6 c) Ca3(PO4)2 21. 3. Be able to convert moles into grams using molar mass, moles into particles using Avogadro’s number a) Convert 100.1 grams of HCl into moles. b) Find the mass in grams of 365.8 moles of SO2. c) How many moles are in 5.43 x 1023 atoms of Ca? d) What is the mass in grams of 1.20 x 108 formula units of CuO? 22. 4. Be able to find percent composition.

23. a) Find the percent composition of NH3 24. b) Find the percent composition for 80.0 g Ba and 32.0 g of Cl.

25. 5. Which of the following is a molecular formula of XY3? X2Y3, XY4, X2Y5, X2Y6 26. 6. What is the empirical formula of the following:

27. a) C4H16 b) P4O10 28. Honors Only: 29. 7. Be able to find the empirical formula and molecular formulas. a) What is the empirical formula of a compound that is 25.9% nitrogen and 74.1% oxygen? b) What is the empirical formula of a compound that has a mass of 10.150 grams and contains 4.433 grams of P and 5.717 g of O? c) What is the molecular formula if the empirical formula is NaO and the gram formula mass is 78 g? 30. 31.