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Home , Dibromine pentoxide, Magnesium sulfide, Sodium iodide

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Honors Chemistry - Final Exam Review Final Exam Dates: Thursday, June 5, 2014 and Tuesday, June 10, 2014. The final exam is worth 100 points and is multiple choice format. Part 1 will assess Sections 1 – 6; Part 2 will assess Sections 7 - 12.

Final Review Practice Directions: Complete on a separate sheet of paper. Show all work. Section 1 and 2 – Introduction to Chemistry and Measurement 1. Is the following set of data accurate, precise, both, or neither if the actual value is 23.4 mL? 20.4 mL 20.3 mL 20.1 mL 20.5 mL The measurements are precise but not accurate. 2. Indicate the number of significant figures in the following: a. 1040 b. 600 c. 12.40 d. 0.00230 e. 0.82 3 1 4 3 2 3. Solve each expression using correct number of significant figures and units. a. 24.4g + 110g 134 g b. 145.30 mL – 12.8 mL 132.5 mL c. 9.8 g / 0.123 L 80. g/L d. 1.4500 cm x 2.55 cm x 1.5 cm 5.5 cm3

4. Solve the following using dimensional analysis (factor label method). a. Convert 23.5 yd to ft. 70.5 ft b. A pound of coffee beans yields 50 cups (this number is obtained by counting) of coffee. How many milliliters of coffee can be obtained from 1.00 g of coffee beans? (1 cup = 240mL) 26.4 mL 2 2 2 c. Convert 23.50 m/s to km/min . 84.60 km/min

5. Convert the following to the indicated temperatures - -29 F -60. C

Gas and air are controlled by knobs on the burner. The gas and air mix 6. Explain how a Bunsen burner works. resulting in a smokeless flame that produces very high temperatures. 7. What is the most practical and accurate tool used to measure volume in a laboratory? Graduated Cylinder 8. Explain how you would set up a laboratory apparatus for a. heating a crucible of b. heating a beaker of copper II oxide and Ring stand with an iron ring supporting a clay triangle Ring stand with wire gauze 9. A student takes an object with an accepted mass of 200.00 grams and masses it on his own balance. He records the mass of the object as 196.50 g. What is his percent error? 1.75%

10. What value does each of the following metric unit represent? a. centi b. nano c. kilo d. milli e. micro 0.01 or 1x10-2 0.000000001 or 1x10-9 1000 or 1x103 0.001 or 1x10-3 0.000001 or 1x10-6 Section 3 – Matter and Energy 1. Solve the following problems using significant figures. a. What is the density (g/mL) of a block of metal 6cm x 10cm x 100cm that weighs 6820 dg? What would be the volume (in L) of 2.8 kg of the metal? 20 L b. A block of metal has a mass of 42.3 g. It is placed in a graduated cylinder containing 56.0 cm3 of water. If the final volume is 75.1 mL, what is the density of the metal block? 2.21 g/mL

2. H w w uld y u sep r e… a. iron filings and pepper b. and sand c. and water magnet Dissolve salt, filter sand, evaporate water distillation 3. What is the specific heat of water? 4.184 J/g°C 4. re u res re e w r r ey e e ly e s e ss s ple lu u r s ple w er pl More energy to heat water than aluminum because water has a higher specific heat.

5. A 0 s ple pper ls r - s e e s e er y e spe e copper is 0 -570 J A physical change results in no change to the components of the substance. A chemical change results 6. Distinguish between a physical and chemical change. in a change to the components of a substance. 7. Classify the following as either a physical or chemical change: a. spoiling of milk chemical b. softening glass to bend it into a new shape physical c. burning a piece of paper chemical d. rusting of a nail chemical

8. Identify the following as a pure substance, heterogeneous mixture, or homogeneous mixture. a. copper b. sweetened tea c. calcium carbonate d. sandy water Pure - element Homogeneous Pure - compound Heterogeneous 9. What is the difference between an exothermic reaction and an endothermic reaction? Endothermic reactions absorb energy from their surroundings. Exothermic reactions release energy into their surroundings.

10. Using the potential energy diagram to the right, answer the following questions.

a. How much energy do the reactants have? 200 kJ b. How much energy is needed to get the reaction started? 250 kJ c. What type of reaction is this, endothermic or exotherimic? Endothermic, product has more energy than reactants

Section 4 – Atoms and the Periodic Table

1. s e ss s l er s s l er’s ss u er 107.87 amu; 108 2. Determine the symbol for the atom, , and or isotope that fits the description below. a. 16 p, 18 e, 16 n b. 26 p, 26 e, 30 n c. 13 p, 10 e, 11 n d. 35 p, 35 e, 45 n -2 +3 Sn Fe Al Br

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3. Determine the identity of an element with a mass number of 75, that contains 42 neutrons. As

4. Determine the number of protons, neutrons, and electrons in the following: a. Ar – 40 b. F-1 c. 29 Al d. Co e. 56 Fe 18p; 22n; 18e 9p; 10n; 10e 13p; 16n; 13e 27p; 32n; 27e 26p; 30n; 26e

5. Distinguish between an alpha and beta particle. alpha particle is 2 protons and 2 neutrons (helium nucleus), beta particle is an electron

6. If 36 grams of water are produced and 4 grams of are used in its production, how many grams of are needed? What law is represented here? 32 grams of oxygen, Law of Conservation of mass

7. s e d ere e e wee e ss u er r r d r ’s er e ss amu? Mass number is the sum of protons and neutrons for one isotope. Average atomic mass is the average mass of all of the isotopes of carbon and their abundance in nature. 8. Calculate the atomic mass of lithium if one isotope has a mass of 6.0151 amu and a percent abundance of 7.59% and a second isotope has a mass of 7.0600 amu and a percent abundance of 92.41%. 6.98 amu 9. Describe the debate between Democritus and Aristotle in terms of the atom. Democritus: atoms are the smallest particle in existence and cannot be further broken down, he called them atomos. He states that all matter consists of a collection of atoms and if there was space between the atoms, it contains nothing. Aristotle: believed that everything could be broken down into smaller and smaller pieces

forever. His ideas were widely accepted because of his popularity and fame. 10. L s e e p s ul es D l ’s e ry pl w re s s why it is no longer accepted. 1.) All matter consists of tiny particles called atoms 2.) Atoms of an element are unchangeable and indestructible. 3.) Elements of the same type have the same mass, elements of different types have different masses. 4.) When elements react, their atoms combine in simple, whole number ratios 5.) When elements react, their atoms sometimes can combine in more than one simple, whole number ratio. *Atoms of the same element can have different masses! *Atoms can be further broken down into protons, neutrons and electrons! 11. p re d r s Ru er rd d T s ’s experiments and atomic models. Thomson – experiments using cathode ray tubes and discovers the electron. His model is referred to as the Plum Pudding Model where atoms are sphere shaped and composed of a space of positive charge. The negative electrons are found within this area of positive charge. Rutherford – experiments using sheets of gold foil and alpha particles and discovers that the alpha particles are deflected back differently than what the plum pudding model would predict. His model is referred to as the nuclear model where the electrons are found in the empty space and the positive charge or protons are found in the center of the atom in what he refers to as the nucleus. 12. s e d es e su p r le e re rds ele e ’s pr per es Protons – used to identify an individual atom Neutrons – affect the average atomic mass of an atom Electrons – determine an atoms chemical behavior

13. Write the symbol of the element that fits each description: a. alkali metal in period 3 Na c. the only metalloid (oxygen group ) Te b. alkaline earth metal with 5 energy levels d. with four energy levels Br Sr

14. How many valence electrons does each of the following possess? What ion does each form? a. calcium b. silver c. phosphorus d. lead e. carbon 2; Ca+2 1; Ag+1 5; P-3 2 or 4; Pb+2 or Pb+4 4; C+4 or -4 15. Which is smaller? a. Ca or Ca+2 Ca+2 c. N or F F b. P or O O d. ion or Aluminum ion aluminum

16. Does each of the following trends increase or decrease down a family or across (left to right) a period? Down a Family Across a Period Electronegativity decreases increases Ionization Energy decreases increases Atomic Radius increases decreases

17. Define electron affinity. Which elements have a negative value for electron affinity? Positive value? Why? Electron affinity is the energy change that occurs when a neutral atom acquires an electron. The more negative the value, the greater the amount of energy that is released. Elements that readily gain electrons like the will have the most negative values because little energy is required to force an atom to acquire the electron.

18. Draw the Lewis dot diagrams for the following elements. a. P b. I c. Al

19. Why are the alkali metals one of the most reactive families of elements? They have low ionization energies because they are very willing to give up their one valence electron

20. Which groups make up the representative or main group elements? 1, 2, 13, 14, 15, 16, 17, 18 21. Give three physical properties that differ between metals and nonmetals. Metals will lose their valence electrons; they are good conductors of heat and electricity; have luster; have high boiling points and melting points. Nonmetals gain electrons; are poor conductors; lack luster; are brittle; and have low boiling points and melting points

22. What is a metalloid? List each of them. Metalloids are elements that share characteristics with both metals and nonmetals. Boron, silicon, germanium,

arsenic, antimony, tellurium, polonium, and astatine. Section 5 – Electrons in Atoms 1. What is the maximum number of electrons that can occupy the following? a. 3d 10 b. 3rd energy level c. 4s 2 d. one orbital e. 4th energy level 18 2 32

2. Draw the orbital diagrams for the following elements: a. b. copper c.

3. Write the electron configuration for the following elements/: 1s22s22p63s23p64s23d104p6 5s14d5 2 2 6 2 6 2 a. calcium 1s 2s 2p 3s 3p 4s c. molybdenum e. sulfur ion 2 2 6 2 2 6 2 2 2 6 2 6 2 10 3 1s 2s 2p 1s 2s 2p 3s 3p b. arsenic 1s 2s 2p 3s 3p 4s 3d 4p d. magnesium ion 6

4. Write the noble gas configuration for the following elements. a. Tin [Kr] 5s24d105p2 b. Gold [Xe] 6s1 4f145d10 c. Chromium [Ar] 4s13d5

5. Answer the following questions about the electron configuration of aluminum: a. What does the 3 in 3s2 mean? Third energy level d. How many electrons will an atom of aluminum b. What does the 2 in 3s2 mean? Two electrons in that sublevel(lose or gain) to form an ion? Lose 3

c. How many valence electrons does this atom e. What orbitals will aluminum lose its electrons have? 3 from? 3s and 3p 6. What is needed for an electron to move from one energy level to a higher energy level?

A quanta of energy must be absorbed 7. e de er ed r ’s e ss spe r The wavelength/frequency/energy of light emitted by an atom when electrons return to ground state.

8. What happens when an electron falls back to its ground state? It emits of a specific wavelength/frequency/color

Section 6 – Ionic and Covalent Bonding 1. What is the difference between an ionic and a covalent bond? Ionic bond – the more electronegative atom will take the electrons from the other atom forming ions Covalent bond – the electrons are shared between two atoms 2. Compare the properties of an and a molecular compound. Ionic compounds – ionize when in water, conduct electricity when dissolved in water, high

Molecular compounds – do not ionize in water, do not conduct electricity, low melting point 3. Would the following pairs most likely form an ionic or molecular compound? a. Mg and Cl ionic c. P and Cl molecular e. Ag and S ionic b. I and F molecular d. Sn and O ionic 4. Draw a Lewis dot diagram for each of the following: a. Be b. N c. Al d. O e. Li

5. Using electronegativity values, write the ranges for a polar covalent, non-polar covalent, and ionic bonds. What is the significance of these electronegativity differences? 1.8 and greater = IONIC 0.4 - 1.7 = POLAR COVALENT 0 - 0.3 = NONPOLAR COVALENT The electronegativity difference determines whether or not the electrons are able to be transferred or shared 6. Describe the three types of intermolecular forces. Give one example of each.  Dispersion force – weak force between nonpolar . Any diatomic like F2  Dipole – dipole force – force between oppositely charged ends of two polar molecules. The more polar the molecule, the stronger the force. SCl2  Hydrogen bond – especially strong attraction between the hydrogen end and a fluorine, nitrogen, or oxygen atom on the other dipole. Water

7. Draw the Lewis structures for each of the following compounds:

a. N2 b. O3 c. H2S d. NCl3 e. SiF4  Identify the total number of valence electrons for each.  What is the molecular shape (linear, bent, trigonal pyramidal, trigonal planar, tetrahedral)?  What is the bond angle for each shape?  Polar or nonpolar covalent bonds?  Polar or nonpolar molecule? a. N2 - 10; linear; 180; nonpolar; nonpolar

b. O3 – 18; Bent; less than 120; nonpolar, nonpolar

RESONANCE STRUCTURES

c. H2S – 8; Bent; less than 109.5; polar, polar

d. NCl3 – 26; trigonal pyramidal; less than 109.5; slightly polar, polar

e. SiF4 – 32; tetrahedral; 109.5; polar, nonpolar

Section 7 – Nomenclature 1. Name each of the following:

 Ba(CN)2 Barium cyanide  SO2  Hg3(PO3)2 Mercury (II) phospate

 CoBr2 Cobalt (II) bromide  BaSO3 Barium sulfite  Na2O oxide

 P2O5 Diphosphorus pentoxide  Zn(ClO4)2 Zinc perchlorate  LiC2H3O2 Lithium acetate

 NaH  Be(BrO)2 Beryllium hypobromite  FeCl3 Iron (III) chloride

 MgI2 Magnesium  HNO3 (aq) Nitric acid  CCl4 Carbon tetrachloride

 Ca3(AsO4) Calcium arsenate  Al(OH)3 Aluminum hydroxide  AgNO3 Silver nitrate

 HCl  HBrO4 (aq) Perbromic acid  N2O3 Dinitrogen trioxide

 Sn(SO4)2 Tin (IV) Sulfate  KMnO4 Potassium permanganate  Pb3N2 Lead (II) nitride

2. Write the formula for each of the following:

Chloride NH4Cl  Copper (I) Phosphide Cu3P  Phosphorus PI3

 Iron (III) Nitrate Fe(NO3)3  Selenium (VI) Fluoride SeF6  Potassium Hydroxide KOH

 Phosphoric Acid H3PO4  NaCN  Mercury (II) Sulfate HgSO4

 Titanium (IV) Bromide TiBr4  Hydroselenic Acid H2Se  Gold (III) Bromide AuBr3  CdO  Silicon Dioxide SiO2  Magnesium MgS

 Calcium Hypochlorite Ca(ClO)2  H2SO4  Lead (II) Acetate Pb(C2H3O2)2

 Lead (II) Phosphate Pb3(PO4)2  Aluminum Chlorate Al(ClO3)3  Manganese (III) Sulfite Mn2(SO3)3

Section 8 – Chemical Reactions 1. For each of the following: a. Indicate the type of the reaction b. Predict the products by name c. Write the equation for the complete reaction. If the reaction does not occur, write NR. d. For DD reactions include the states e. Balance the reaction only if it occurs o silver nitrate + copper (II) chloride  silver chloride + copper (II) nitrate

(double displacement - precipitate) 2AgNO3(aq) + CuCl2(aq) --> 2AgCl(s) + Cu(NO3)2(aq) o copper (II) + aluminum sulfate  NR o + ammonium sulfate  NR o iodic acid  diiodine pentoxide + water

(decomposition) 2HIO3 --> I2O5 + H2O o hydrochloric acid + barium hydroxide  water + barium chloride

(double displacement - acid/base) 2HCl(aq) + Ba(OH)2(aq) --> BaCl2(aq) + H2O(l) o sodium + chromium (II) chlorate  + chromium

(single displacement) 2Na + Cr(ClO3)2 --> 2NaClO3 + Cr o copper (II) oxide  copper (II) + oxygen

(decomposition) 2CuO --> 2Cu + O2

o C4H8 + oxygen  + water

(combustion) C4H8 + O2 --> CO2 + H2O o dibromine pentoxide + water  bromic acid

(synthesis) Br2O5 + H2O  2HBrO3 o iron (II) oxide + carbon dioxide  iron (II) carbonate

(synthesis) FeO + CO2  FeCO3 o chlorous acid +  water + magnesium chlorite

(double displacement - acid/base) 2HClO2(aq) + Mg(OH)2(s)  Mg(ClO)2(aq) + 2H2O(l) o aluminum carbonate  aluminum oxide + carbon dioxide

(decomposition) Al2(CO3)3  Al2O3 + 3CO2 o + hydrochloric acid  +

(double displacement - gas forming) Na2S + 2HCl(aq)  2NaCl (aq) + H2S(g)

o hydrobromic acid + calcium carbonate  calcium bromide + carbon dioxide + water (carbonic acid

breaks down!!) (double displacement - gas forming) HBr(aq) + CaCO3(s)  CaBr2(aq) + CO2(g) + H2O(l)

Oxidation is a loss of electrons by one 2. What is the difference between oxidation and reduction? reactant, reduction is a gain of electrons 3. Write the net ionic equations for the reaction of: a. hydrochloric acid + barium hydroxide  c. plumbic sulfide + hydrochloric acid 

+ - + - H (aq) + OH (aq)  H2O(l) 2H (aq) + S (aq)  H2S(g) b. chlorous acid + calcium hydroxide  d. silver nitrate + cupric chloride 

+ - + - H (aq) + OH (aq)  H2O(l) Ag (aq) + Cl (aq)  AgCl (s)

4. What are the four main signs that a has occurred? Gas formation, temperature change, precipitate formation, color change Section 9 – Chemical Composition 2. What is the of each of the following? (varies depending on the Periodic Table used)

a. Al(C2H3O2)3 204.13 g/mol b. Ammonium chloride 53.50 g/mol c. Potassium bromate 167.00 g/mol 3. I know that 8.401 x 1024 atoms of a certain element has a mass of 391.8 g. What element is this? Silicon 4. How many Sulfide ions are present in 12.24 mg of ? 6.683 x 1019 ions 5. How many carbon atoms are in 1 mole of dicarbon monoxide? 1.20 x 1024 carbon atoms

23 6. How many grams are in 2.34 x 10 molecules of C6H10OS2? 63.1 grams 7. How many molecules can be found in a sample of 0.847 grams of carbon dioxide? 1.16x 1022 molecules 8. Determine the percent composition of all of the elements found in the following compounds: a. Sodium chlorate b. Copper (I) sulfide

NaClO3: 45.10% O; 33.30% Cl; 21.60% Na Cu2S: 20.15% S; 79.15% Cu 9. What is the difference between an empirical and molecular formula? The empirical formula shows the most simplified ratio of elements in a compound. The molecular formula shows the actual ratio of the elements in the compound. 10. Determine the empirical formula for a compound using the following information:

a. 58.00% Rb; 9.50% N; 32.50% O RbNO3 b. 49.62% C; 10.82% H; 39.56% O C5H13O 11. What is the molecular formula of a compound with an empirical formula of CHOCl (chlorine) and a molecular

weight of 129 grams? C2H2O2Cl2

12. What are the 7 diatomic elements? Br2, I2, N2, Cl2, H2, O2, F2 13. Determine the molecular formula for a compound containing 7.70% carbon, 1.30% hydrogen, and 81.90%

. Its molar mass is 281 g/mol. C2H4I2

14. 1.00 mol of Al2S3 consists of how many total mol of ions? 5.00 mol (2.00 mol Al and 3.00 mol S)

Section 10 – Stoichiometry 1. If you react 25.4 g of with an excess of sodium iodide and you produce 2320 mg of iodine, what is the percent yield of iodine? 5.75% 2. How many grams of are produced if 10.0 g of sulfuric acid and 7000.0 mg of potassium hydroxide are mixed together? What is the limiting reactant? LR = KOH; 10.872 g 3. How many moles of chloric acid are required to react completely with 2.05 g of ? 0.0387 mol 4. The reusable booster rockets of the U.S. space shuttle employ a mixture of aluminum and ammonium perchlorate for fuel. The balanced equation is:

3Al(s) + 3NH4ClO4  Al2O3(s) + 3NO(g) + 6H2O(l) NH3(g) a. What mass of ammonium perchlorate should be used in a fuel mixture for every 1.00 kg of aluminum used? 4360 grams or 4.36 kg 5. How many grams of are produced if 420 cm3 of chlorine gas reacts with 14.8 g of

Magnesium at STP? 1.8 g MgCl2

Section 11 – Solutions 1. If 960 g of are used to prepare 16000 mL solution. What is the molarity of the solution? 1.5 M 2. If we add 0.15 L of a 6.02M HCl solution and some water to a beaker, creating a 2.42 M HCl solution, what is the volume of the new solution? 0.37 L 3. A 2.34 M solution contains 10. g of . What is the volume of the solution in mL? 71 mL 4. What mass of is consumed if it reacts with 12.3 mL of a 3.45 M HCl solution? 1.32g

5. How many grams of 0.67 M KClO3 are needed to prepare 1.00 Liters of solution? 82 g 6. Describe how 100 ml of a 0.10 M solution of sodium hydroxide would be made. Fill a flask approximate 2/3 full of water. Add 0.4 grams of sodium hydroxide to the flask. Swirl the flask. Add remaining 1/3 of water to the solution until meniscus is at the line on the flask. Mix. 7. Using the pH scale, define an acid and a base. An acid is a substance with a pH value of 0 – 6.99 and a base is a substance with a pH value of 7.01 - 14 8. Calculate the pH when [H+] = 3.9 x 10 -5. 4.41 Calculate the [H+] when the pH value is 9.57. 2.7 x 10-10

9. Which solution is a poor conductor of electricity? 1M NaCl; 1M HC2H3O2; 1M NaOH; 1M HCl

10. What volume of 3.1 M H2SO4 will neutralize 250 mL of 0.30 M Al(OH)3? 360 mL

Section 12 – , Liquids, and Gases 1. Describe the conditions that allow a liquid to boil. The vapor pressure is equal to the atmospheric pressure allowing the liquid to evaporate. 2. Compare the particles within a , liquid, and a gas. The particles of a solid are close together and fixed in place, liquid the particles can slide past each other, and a gas the particles are spread out to move freely and randomly. 3. Relate the intermolecular forces within a solid, liquid, and gas. The intermolecular forces within solid water are much stronger than liquid water. The intermolecular forces within water vapor are minimal. 4. Identify and define the different phase changes of matter. Sublimation – solid to gas; Deposition – gas to solid; Condensation – gas to liquid; Evaporation – liquid to gas; Freezing – liquid to solid; Melting – solid to liquid 5. What is vapor pressure? The force exerted by a vapor that has been released by a liquid. 6. A balloon at 22 °C holds 2.00L of oxygen. If the gas in the balloon is heated to 90° C, what volume will the balloon be? 2.5 L 7. What is the temperature of 0.70 moles of a gas that occupies 470 mL at a pressure of 150 kPa? 12 K

8. 500.0 mL of O2 gas is collected over water. Atmospheric pressure is 101.0 kPa. The temperature of the water is 22 °C. What is the pressure of the gas? You will need a chart from your notes to complete this problem! 737.8 mmHg 9. A gas is at 1.33 atm of pressure and a volume of 682 mL. What will the pressure be if the volume is reduced to 0.419 L? 2.16 atm 10. Nitrogen gas is being held in a 14.3 m3 tank at a temperature of 62oC. What will the volume be when the temperature drops to 24oC? 12.6 m3 11. A gas storage tank is a 1.72 atm and 35oC. What temperature is the gas at if the pressure increases to 2.00 atm? 358 K 12. A gas with a volume of 1.00 L is at 135oC and 844 mm Hg. What is the volume if the conditions change to 14o C and 748 mm Hg? 0.794 L 13. Calculate the mass of 162 L of chlorine gas, measured at STP. 512 g 14. Find the pressure in mm Hg produced by 2.35 g of carbon dioxide in a 5.00 L flask at 18oC. 194 mmHg 15. How many grams of must be placed into a 40.0 L tank to develop a pressure of 965 mm Hg at 23oC? 58.5 g 16. At what Celsius temperature will argon have a density of 10.3 g/L and a pressure of 6.43 atm? 30.6 °C