FINAL EXAM DATES THURSDAY, JUNE 5, 2014 and TUESDAY, JUNE 10, 2014

Name______

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 Chapters 1 – 6; Part 2 will assess Chapters 7 - 13.

FINAL REVIEW PRACTICE

DIRECTIONS: Complete on a separate sheet of paper. Show all work.

CHAPTER 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 mL20.3 mL20.1 mL20.5 mL

The measurements are precise but not accurate.

2. Indicate the number of significant figures in the following:

a. 1040 3b. 600 1c. 12.40 4d. 0.00230 3e. 0.82 2

3. Solve each expression using correct number of significant figures and units.

a. 24.4g + 110g130 gb. 145.30 mL – 12.8 mL 132.5 mL

c. 9.8 g / .123 L80. g/Ld. 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 yd3 to ft3. 635 ft3

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/g

c. Convert 23.50 m/s2 to km/min2. 84.60 km/min2

5. Convert the following to the indicated temperatures

a. -34 ̊C to ̊F -29 Fb. 213 K to ̊C -60. C

6. Explain how a Bunsen burner works. Gas and air are controlled by knobs on the burner. The gas and air mix 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 magnesium – ring stand with an iron ring supporting a clay triangle

b. heating a beaker of copper II oxide and water – ring stand with an iron ring supporting a wire gauze. The beaker will be placed on the gauze and supported by a second iron ring.

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 196.5 g. What is his percent error? 1.75%

10. What value does each of the following metric unit represent?

a. centi .01 or 1 x 10-2b. nano .000000001 or 1 x 10-9c. kilo 1000 or 1 x 10 3 d. milli .001 or 1 x 10-3 e. micro .000001 or 1 x 10-6

CHAPTER 3 – MATTER AND ENERGY

1. Solve the following density 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? 24.5 L = 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. How would you separate…..

a. iron filings and peppermagnet b. salt and sand dissolve the salt and filter the sand then evaporate the water c. methanol and water distillation

3. What is the specific heat of water? 4.184 J/g C

4. Which requires more heat to warm from 22.0 ̊C to 85.0̊C if they have exactly the same mass, a sample of aluminum (c = .902 J/ g ̊C) or a sample of water. Explain. It takes more energy to heat one gram of water than 1 gram of aluminum because water has a higher specific heat.

5. A .0234 kg sample of copper cools from 53 ̊C to - 9.4 ̊C. What is the change in its energy if the specific heat of copper is .387 J/g ̊C? - 570 J

6. Distinguish between a physical and chemical change. A physical change is a change that results in no change to the components of the substance . A chemical change is a change that results in a change to the components of the substance. The identity of the substance will change due to a physical change.

7. Classify the following as either a physical or chemical change:

a. spoiling of milkCb. softening glass to bend it into a new shapeP

c. burning a piece of paper Cd. rusting of a nail C

8. Identify the following as a pure substance, heterogeneous mixture, or homogeneous mixture.

a. copperPure substance - element

b. sweetened tea Homogeneous mixture - solution

c. calcium carbonate Pure substance - compound

d. sandy water – Heterogeneous mixture - suspension

9. What is the difference between an exothermic reaction and an endothermic reaction? Relate each reaction type to enthalpy. Endothermic reactions absorb energy and the products have more energy than the reactants. The enthalpy change for an endothermic reaction is positive. Exothermic reactions release energy and the products will have less energy than the reactants. Exothermic reactions have a negative enthalpy change.

10. Using the potential energy diagram above, 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

CHAPTER 4 – CHEMICAL FOUNDATIONS

1. What is the atomic mass of silver? What is silver’s mass number? 107.87 amu ; 108

2. Determine the symbol for the atom, ion, and or isotope that fits the description below.

a. 16 p, 18 e, 16 nS-2b. 26 p, 26 e, 30 nFec. 13 p, 10 e, 11 nAl+3

d. 35 p, 35 e, 45 n Br

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 – 4018p; 22n; 18 eb. F-19p; 10n; 10e c. 29 Al13p; 16n; 13e d. Co 27p; 32n; 27e e. 56 Fe 26p; 30n; 26e

26

5. Distinguish between an alpha and beta particle. Alpha decay is the emission of a helium nucleus (2 protons and 2 neutrons = alpha particle). Beta decay is the splitting of a neutron into a proton and an electron = beta particle. The proton is added to the nucleus of the atom, changing the identity of the atom.

6. If 36 grams of water are produced and 4 grams of hydrogen are used in its production, how many grams of oxygen is needed? What law is represented here? 32.0 grams of O2 = Law of the Conversation mass

7. What is the difference between the mass number for Carbon 14 and carbon’s average atomic mass of 12.011 amu? Mass number gives the total number of protons and neutrons for one isotope of carbon. The average atomic mass is the calculated sum 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 – small pieces cannot be broken into smaller pieces forever – sooner or later, a single particle so small that cannot be broken down into anything further will remain. He calls this end particle atomos. He states that all matter consists of a collection of atoms and if there was space between the atoms, it contains nothing. Aristotle – he believes that it makes more sense to believe that everything can be broken into smaller and smaller pieces forever. He questions that if matter is composed of particles and empty space, what holds the particles together? Due to Aristotle’s fame and popularity, he wins the debate.

10. List the five postulates to Dalton’s theory. Explain two reasons why it is no longer accepted.

1. All matter consists of tiny particles.

2. Atoms of an element are unchangeable and indestructible.

3. Elements are characterized by the mass of their elements.

– All atoms of the same element have the same mass.

– Atoms of different elements have different masses.

4. When elements react, their atoms combine in simple, whole number ratios.

– Different atoms combine in simple whole number ratios to form compounds.

5. When elements react, their atoms sometimes combine in more than one simple, whole number ratio.

11. Compare and contrast Rutherford and Thomson’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. What significance does each subatomic particle have in regards to an element’s properties?

Protons – used to identify an individual atom

Neutrons – effect the average atomic mass of an atom

Electrons – determine an atom’s chemical behavior

13. Radioactive strontium (Sr - 90) has a half life of 29 years. If you had a 10.0 gram sample, how many years will it take you to get 2.5 grams of sample?

10.0 grams  5.0 grams = 29 years; 5.0 grams  2.5 grams = 29 years TOTAL = 58 years

14. What element undergoes alpha decay to form lead-208? Explain.

When a radioactive element undergoes alpha decay, it emits an alpha particle that is 2 protons and 2 neutrons (Helium nucleus). The element is Po – 212.

CHAPTER 5 – MODERN ATOMIC THEORY/QUANTUM MECHANICS

1. What is the wavelength (in m) of a photon of light containing 3.19 x 10-13 J of energy?

6.23 x 10 -13 m

1. How much energy in Joules will a photon of light with a wavelength of 345.0 nm possess?

5.762 x 10-19 J

3. What is the maximum number of electrons that can occupy the following?

a. 3d 10b. 3rd energy level18c. 4s 2d. one orbital 2

e. 4th energy level 32

4. Draw the orbital diagrams for the following elements:

a. argonb. copper

c. nitrogen

5. Write the electron configuration for the following elements/ions:

a. calcium 1s22s22p63s23p64s2b. arsenic1s22s22p63s23p64s23d104p3

c. molybdenum1s22s22p63s23p64s23d104p6 5s14d5

d. magnesium ion1s22s22p6e. sulfur ion 1s22s22p63s23p6

6. Write the noble gas configuration for the following elements.

a. Tin[Kr] 5s24d105p2b. Gold [Xe] 6s1 4f14 5d10c. Chromium [Ar] 4s13d5

7. Answer the following questions about the electron configuration of aluminum:

a. What does the 3 in 3s2 mean? Third energy level

b. What does the 2 in 3s2 mean? Two electrons in sublevel

c. How many valence electrons does this atom have? 3

d. How many electrons will an atom of aluminum (lose or gain) to form an ion? Lose 3

e. What orbitals will aluminum lose its electrons from? 3s and 3p

8. What is needed for an electron to move from one energy level to a higher energy level?

A quanta of energy must be absorbed

9. What can be determined from an atom’s atomic emission spectra? The wavelength/frequency/energy of light emitted by an atom when electrons return to ground state.

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

CHAPTER 6 – THE PERIODIC TABLE

1. Write the symbol of the element that fits each description:

a. alkali metal in period 3Sodiumb. alkaline earth metal with 5 energy levels Strontium

c. the only metalloid chalcogen (oxygen group) Tellurium (Polonium is acceptable too)

d. halogen with four energy levels Bromine

2. How many valence electrons does each of the following possess? What ion does each form?

a. calcium 2; Ca+2b. silver 1; Ag+1c. phosphorus 5; P-3d. lead 2/4; Pb+2/Pb+4 e. carbon 4; C-4

3. Which is smaller?

a. Ca or Ca+2 b. P or Oc. N or Fd. Sulfur ion or Aluminum ion

4. Does each of the following trends increase or decrease down a family or across (left to right) a period?

5. Define electron affinity. Which elements have a negative value for electron affinity? 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 halogens will have the most negative values because little energy is required to force an atom to acquire the electron.

Metals such as the alkaline earth metals will have positive values. Metals wish to lose electrons, not gain them. In order to add electrons to a metal, energy must be added forcing the addition of the electron.

6. Draw the Lewis dot diagrams for the following elements.

a. b. c.

7. Why are the alkali metals one of the most reactive families of elements? They have low ionization energy.

8. Which groups make up the representative or main group elements? 1, 2, 13, 14, 15, 16, 17, and 18

9. Give three physical properties that differ between metals and nonmetals. Metals will lose electrons; are good conductors of electricity and heat; have luster; exist as solids; high boiling points and melting points. Nonmetals will gain electrons; are poor conductors of electricity and heat, lack luster, are normally found as gases or solid; are brittle; low boiling points and melting points.

10. What is a metalloid? List each of them. Elements that behave as both a metal and a nonmetal. Boron, Silicon, Germanium, Arsenic, Antimony,Tellurium, Polonium, and Astatine

CHAPTER 7 – CHEMICAL BONDING

1. What is the difference between an ionic and a covalent bond?

An ionic bond is a bond commonly formed between a metal and a nonmetal. The electronegativity difference is normally greater than 1.8 causing a transfer of electrons from one atom to the other. A covalent bond is a bond commonly formed between two nonmetals. The electronegativity difference is less than 1.8 causing the electrons to be shared (and not transferred) from one atom to the other.

1. Would the following pairs most likely form an ionic or molecular compound?
2. Mg and Cl - ionicb. I and F - molecularc. P and Cl- molecular
3. d. Sn and O - ionic e. Ag and S - ionic
4. Draw a Lewis dot diagram for each of the following:
5. b. c. d. e.

1. 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; .4 – 1.7 polar covalent; 0 - .3 nonpolar covalent – the electronegativity difference determines whether or not the electrons are able to be transferred or shared.

1. Describe the three types of intermolecular forces. Give one example of each.

Intermolecular forces are weak attractive forces between MOLECULES in a molecular compound.

• Dispersion force – weak force between nonpolar molecules. Any diatomic molecule 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

1. Draw the Lewis structures for each of the following compounds:
2. Identify the total number of valence electrons for each.
3. What is the molecular shape (linear, bent, trigonal pyramidal, trigonal planar, tetrahedral)?
4. What is the bond angle for each shape?
5. Polar or nonpolar covalent bonds?
6. Polar or nonpolar molecule?

1. 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

CHAPTER 8 – NOMENCLATURE

1. Name each of the following:

1. Ba(CN)2
2. CoBr2
3. P2O5 (STOCK)
4. NaH
5. MgI2
6. Ca3(AsO4)
7. HCl
8. Sn(SO4)2
9. SO2 (CLASSIC)
10. BaSO3
11. Zn(ClO4)2
12. Be(BrO)2
13. HNO3 (aq)
14. Al(OH)3
15. HBrO4 (aq)
16. KMnO4
17. Hg3(PO3)2 (CLASSIC)
18. Na2O2ME
19. LiC2H3O2
20. FeCl3 ∙ 6 H2O (CLASSIC)
21. CCl4
22. AgNO3
23. N2O3 (STOCK)
24. Pb3N2 (STOCK)

1. Barium Cyanide
2. Cobalt (II) Bromide
3. Phosphorus (V) Oxide
4. Sodium Hydride
5. Magnesium Iodide
6. Calcium Arsenate
7. Hydrogen Chloride
8. Tin (IV) Sulfate
9. Sulfur Dioxide
10. Barium Sulfite
11. Zinc Perchlorate
12. Beryllium Hypobromite
13. Nitric Acid
14. Aluminum Hydroxide
15. Perbromic Acid
16. Potassium Permanganate
17. Mercuric Phosphite
18. Sodium Peroxide
19. Lithium Acetate
20. Ferric Chloride Hexahydrate
21. Carbon Tetrachloride
22. Silver Nitrate
23. Nitrogen (III) Oxide
24. Lead (II) Nitride

Write the formula for each of the following:

1. Diboron Hexahydride
2. Ammonium Chloride
3. Iron (III) Nitrate
4. Phosphoric Acid
5. Titanium (IV) Bromide
6. Cuprous Phosphide
7. Selenium (VI) Fluoride
8. Sodium Cyanide
9. Hydroselenic Acid
10. Phosphorus Triiodide
11. Barium Chloride Dihydrate
12. Potassium Hydroxide
13. Mercuric Sulfate
14. Auric Bromide
15. Cadmium Oxide
16. Calcium Hypochlorite
17. Lead (II) Phosphate
18. Magnesium Sulfate Heptahydrate
19. Silicon Dioxide
20. Sulfuric Acid
21. Aluminum Chlorate
22. Magnesium Sulfide
23. Plumbous Acetate
24. Manganese (III) Sulfite

1. B2H6
2. NH4Cl
3. Fe(NO3)3
4. H3PO4
5. TiBr4
6. Cu3P
7. SeF6
8. NaCN
9. H2Se
10. PI3
11. BaCl2 · 2H2O
12. KOH
13. HgSO4
14. AuBr3
15. CdO
16. Ca(ClO)2
17. Pb3(PO4)2
18. MgSO4 · 7H2O
19. SiO2
20. H2SO4
21. Al(ClO3)3
22. MgS
23. Pb(C2H3O2)2
24. Mn2(SO3)3

CHAPTER 9 – THE MOLE

1. What is the molar mass of each of the following?

a. Al(C2H3O2)3? 204.13 g/mol b. Ammonium chloride53.50 g/mol

c. Potassium bromate 167.00 g/mol

2. I know that 8.401 x 1024 atoms of a certain element has a mass of 391.8 g.

What element is this? Silicon

3. How many Sulfide ions are present in 12.24 mg of potassium sulfide? 6.683 x 1019 ions

4. How many carbon atoms are in 1.00 mole of dicarbon hexahydride? 1.20 x 1024 carbon atoms

5. How many grams are 2.34 x 1023 molecules of C6H10OS2? 63.1 grams

6. How many molecules can be found in a sample of .847 grams of carbon dioxide?

1.16x 1022 molecules

7. Determine the percent composition of all of the elements found in the following compounds:

a. Sodium chlorateNaClO3 45.10% O; 33.30% Cl; 21.60% Na b. Cuprous sulfide Cu2S 20.15% S; 79.15% Cu

8. What is the difference between an empirical and molecular formula? The empirical formula shows the most simplified ratio of the elements in a compound. The molecular formula shows the actual ratio of the elements in the compound.

9. Determine the empirical formula for a compound using the following information:

a. 58.00% Rb; 9.50% N; 32.50% ORbNO3

b. 49.62% C; 10.82% H; 39.56% OC5H13O

10. What is the molecular formula of a compound with an empirical formula of CHOCl (chlorine) and a molecular weight of 129 grams? C2H2O2Cl2

11. What are the 7 diatomic elements? Br2, I2, N2, Cl2, H2, O2, F2

12. Determine the molecular formula for a compound containing 7.70% carbon, 1.30% hydrogen, and 81.90% iodine. Its molar mass is 281 g/mol. C2H4I2

13. If you take 20.8 g of hydrated aluminum oxide and heat it to remove all of the water, what is the formula of the original hydrate if the mass of the remaining solid after heating is 15.4 g? Al2O3 · 2H2O

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)

CHAPTER 10 – 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

a. silver nitrate + cupric chloride  DD (precipitate) – silver chloride and cupric nitrate

2AgNO3(aq) + CuCl2(aq)  2AgCl(s) + Cu(NO3)2(aq)

b. copper (II) + aluminum sulfate  NO REACTION

c. sodium nitrite + ammonium sulfate  NO REACTION

d. iodic acid  D4 - diiodine pentoxide and water

2HIO3  I2O5 + H2O

e. hydrochloric acid +barium hydroxide DD (neutralization)-water and barium chloride

2HCl(aq) + Ba(OH)2(aq)  BaCl2(aq) + H2O(l)

f. sodium + chromium (II) chlorate  SD – sodium chlorate and chromium

2Na + Cr(ClO3)2  2 NaClO3 + Cr

g. cupric oxide  D2 – copper (II) and oxygen

2CuO  2Cu + O2

h. C4H8 + oxygen (complete) C – carbon dioxide and water

C4H8 + 6O2  4CO2 + 4H2O

1. dibromine pentoxide + water  S4 – bromic acid

Br2O5 + H2O  2HBrO3

1. iron (II) oxide + carbon dioxide  S5 – iron (II) carbonate

FeO + CO2  FeCO3

k. chlorous acid + magnesium hydroxide 

DD (neutralization) – water and magnesium chlorite

2HClO2(aq) + Mg(OH)2(s)  Mg(ClO)2(aq) + 2H2O(l)

1. aluminum carbonate  D5 – aluminum oxide and carbon dioxide

Al2(CO3)3  Al2O3 + 3CO2

m. Sodium sulfide + hydrochloric acid 

DD (gaseous) – sodium chloride and hydrogen sulfide gas

Na2S + 2HCl(aq)  2NaCl (aq) + H2S(g)