Download Chemistry Final Exam Review 2006-2007

Thematic Final Exam Review 2013
2nd Quarter
Light
Objective 4.1-4.2
1. List the assumptions of Dalton’s atomic theory. What were the revisions?
2. What are the 3 subatomic particles and state their charge and where they are
found in the atom?
3. Who is Rutherford and what did he discover about the atom?
4. a. What is the dense center of an atom called? b. What subatomic particles are
found in the center?
5. Where are the electrons found in an atom?
6. Define atomic number and atomic mass.
7. Define average atomic mass.
Honors Only
8. Boron-10 has a mass of 10.013 amu and a % abundance of 19.8. Boron -11’s
abundance is 80.2% with a mass of 11.009 amu. Calculate the average atomic
mass.
9. What does “electrically neutral” mean in terms of the atom?
10. How many protons and electrons are in a carbon atom?
11. How many neutrons in beryllium?
12. a. Define an isotope. b. Write the nuclear symbol for nitrogen-15. c. Write the
hyphen notation for 38Cl.
13. How many protons, electrons, and neutrons are in oxygen-16?
14. Determine the number of neutrons in 226Ra and 15N.
15. Fill in the chart below.
Atomic #
Mass #
protons
electrons
neutrons
7
7
9
10
39
19
59
27
16. How many protons and electrons are in the following? a. Na+1 b. N3- c. F-1
d. Al3+?
17. Transition elements are in the __________ block and the inner transitions are in
the ________ block.
18. Group 1A and 2A are in the ___________ block and groups 3A to 8A are in the
__________ block.
19. Write the electron configuration for the following: a. boron b. magnesium, c.
vanadium,d. strontium, e. iron f. arsenic
20. a. What is an atomic orbital? b. What shape is the s sublevel? c. The shape of the
p sublevel? d. What are the maximum number of electrons allowed in each
sublevel?
21. What is the difference between the Bohr model and the Quantum mechanical
model?
22. a. What are flame tests? b. What area of the electromagnetic radiation spectrum
allows us to observe flame tests? c. Is energy released or absorbed when an
electron falls from a higher energy level to a lower energy level?
23. What is the difference between a ground state and an excited state?
24. What is the lowest energy level? The lowest sublevel?
25. What is the maximum number of electrons in the 4th energy level?
26. State the three rules for filling atomic orbitals with electrons and describe them.
27. How many unpaired electrons are in the following: a. boron b. fluorine?
28. How many valence electrons are in the highest energy level (valence)? a. barium
b. sodium c. aluminum d. oxygen.
29. What is the symbol of the following configurations?a. 1s2 2s2 2p6 3s1 b. 1s2 2s2
2p6 3s2 3p6 4s2 3d2 c. 1s2 2s2 2p6 3s2 3p2 d. [Kr] 5s2 4d10.
30. Write the shorthand configurations for a. barium b. aluminum c. arsenic.
Objective 4.4-4.7
1. a. Define the periodic table. b. What is the difference between a group and a
period? c. What do columns of elements have in common?
2. Group A elements are called ___________. Group B elements are called
__________.
3. a. What side of the periodic table are the metals? b. The nonmetals? c. Where are
the metalloids?
4. Name 4 characteristics of both metals and nonmetals.
5. Define malleable and ductile.
6. Identify each as a metal, nonmetal, or metalloid: K, B, Mo, iodine, uranium, and
aluminum.
7. a. Who is Mendeleev? b. Who is Moseley? c. How is the modern periodic table
arranged?
8. Name the a. group 1A metals b. 2A metals c. 7A nonmetals d. 8A nonmetals.
9. State 3-4 properties of each of the families above.
10. Which family is the most stable?
11. Which family reacts vigorously with water?
12. Which family is extracted from mineral ores?
13. Which family are the most reactive metals?
14. Which family of nonmetals combines with 1A and 2Ametals to make salts?
15. a. What is electronegativity? b. What is the period and group trend? c. Which one
has a higher electronegativity; C, N, or K?
16. a. Define ionization energy. b. What is the period and group trend? c. Which has
a higher ionization energy; Na, K, Mg, or P?
17. a. Define atomic radius. b. What is the period and group trend? c. Which has a
higher atomic radius; C, N, Mg, P, Na, or K?
18. a. What is the period and group trend for the ionic size of cations? Of anions?
b. How does the size of a neutral atom compare with the cation and the anion?
19. What is a valence electron? b. How many valence electrons are in potassium and
oxygen?
20. Draw the Lewis structure for a. Mg b. Si c. Cl.
21. Define wavelength and frequency. B. What kind of relationship do they have with
each other? c. What kind of relationship does the energy and frequency of a wave
have?
22. What form of electromagnetic radiation is released when an electrons moves from
n=3 to n=2? What is the wavelength of this type of light?
23. Using the electromagnetic spectrum, what form of energy has a wavelength of 6.5
x 10-1 m?
Soap
Objective 5.1-5.6
1.
2.
3.
4.
State the octet rule. The duet rule.
a. Define “ion”. b. What is the difference between a cation and an anion?
What is a chemical formula?
What is the chemical formula for the following; a. sulfide ion b. sodium ion
c. fluoride ion d. mercury (II) ion?
5. What are the names of the following ions; a. Ba2+ b. Al3+ c. O2- d. Sn4+?
6. Metals form _______ions and nonmetals form ________ ions.
7. a. What is the difference between ionic and covalent bonds? b. How does
electronegativity difference determine bond type?
8. Write the electron configurations for a. Al+3 b. O-2 c. Ti+2.
9. Draw Lewis structures for the following; a. K2O b. MgCl2 c. KI d. Na3P
10. Which of the following compounds are ionic? A. H2O b. Na2O c. CO2
d. CaS, e. SO2 f. CaCO3.
11. What is the difference between a nonpolar covalent bond and a polar covalent
bond?
12. What are the properties of a covalent and ionic compound in terms of state of
matter, ability to conduct electricity, hardness & solubility in water?
13. a. Define chemical bond.
b. What is a lone pair of electrons?
14. What is the difference between a single, double, and triple bond?
15. Draw Lewis diagrams for a. PBr3 b. N2 c. CF4 d. HBr e. SO2
16. Predict the shapes of the molecules drawn in question 15.
17. a. What is a dipole? b. What direction does it travel?
18. What is the difference between a polar and nonpolar molecule?
19. Using the formulas from question 15, list ones which are polar and nonpolar.
20. Describe the model for a metallic bond.
21. Using the model described above, explain why a metal is a great conductor of
electricity.
22. Differentiate the difference between intermolecular forces and intramolecular
forces.
23. List in order of strength of the following IMF : dipole-dipole forces, London
Dispersion forces, and hydrogen bonding
24. Identify which ones have dipole-dipole forces? PBr3, N2, CF4, HBr, H2O
25. Identify which ones have London dispersion forces? , N2, CF4, HBr, SO2
26. Identify which ones have hydrogen bonding? HCl,, H2, HBr, H2O, CH4
27. Define the physical properties of Viscosity, Surface Tension, Boiling Point and
state what happens to the property as intermolecular forces increase for each one.
28. Define Like dissolves like. Which molecule can dissolve in water? CH4 or NH3
Sports Drinks
Objective 6.1-6.3
1.
2.
3.
4.
5.
6.
Name 3 factors that increase the rate of dissolution of a substance.
Describe solution equilibrium.
What is the effect of temperature and pressure above a gas on gas solubility?
What is the effect of temperature on the solubility for most ionic solids?
What is the molarity of 4.5 moles of Ba(OH)2 in 10.0 L?
A solution has a molarity of 2.8 M and a volume of 250 ml. How many moles of
solute are in the solution?
7. Differentiate between saturated, unsaturated and supersaturated solutions
8. Using the solubility graph from the notes, how much of NaCl can be dissolved at
45°C
9. Using the solubility graph from the notes, 50 g of KClO3 is dissolved in 100 g of
water at 45°C. Is the solution saturated or unsaturated?
10. How many ml of a 2.0 M NaBr solution are needed to make 200 ml of a 0.50 M
solution?
11. Which types of substances produce electrolytes?
12. Which type of substances produce non electrolytes?
13. List some common properties of an acid.
14. List some common properties of a base.
15. Differentiate between a strong and weak acid or base.
16. Differentiate between a dilute & concentrated acid or base
17. Define self-ionization of water.
18. Know how to predict the products and balance neutralization (double
replacement) reactions.
a) H2CO3 + Fe(OH)3 →
19. Know how to calculate the pH from hydrogen and hydroxide ion concentrations
a) What is the pH of a [OH-] = 1.0 x 10-5 M?
b) What is the pH of a [H+] = 1.0 x 10-5 M?
c) What is the pOH of a [H+] = 1.0 x 10-1 M?
d) What is the pOH of a [OH-] = 1.0 x 10-12 M?
e) What is the [H+] when the pH = 8.0
f) What is the [OH-] when the pH = 8.0
20. a. What is the hydrogen ion concentration of 0.001 M HNO3? b. What is the
[OH-]?
Honors Only:
21. What is the hydrogen ion concentration of [OH-] = 3.0 x 10-2 M? What is the pH?
22. What is the pH of a solution if the [H+] = 3.4 x 10-5 M? What is the hydroxide
concentration?
23. Determine the pH of a 2.0 x 10-2 M Sr(OH)2?
24. The pH of a solution is measured and determined to be 7.52? What is the
hydrogen ion concentration? Is the solution acidic or basic?
Objective 6.5A, B & C
1. What do the coefficients mean in a chemical equation?
2. Calculate the mole ratio between calcium and oxygen in a chemical formula.
2Ca + O2 → 2CaO
3. Solve mole to mole, mole to mass, mass to mole, mass to mass problems.
a) How many moles of lithium hydroxide are required to react with 20.0 mol of
carbon dioxide? CO2 + 2LiOH → Li2CO3 + H2O
b) What mass, in grams, of glucose is produces when 3.00 mol of water react with
carbon dioxide? 6CO2 + 6H2O → C6H12O6 + 6O2
c) How many moles of NO are formed when 824 g of ammonia reacts with an
excess of oxygen? (balance the equation first) NH3 + O2 → NO + H2O
Honors Only
d) How many grams of SnF2 are produced from the reaction of 30.0 g HF with Sn?
Sn + 2HF → SnF2 + H2
4. Know how to solve percent yield problems.
a) When 36.8 g of C6H6 react with an excess of Cl2. The actual yield of C6H5Cl is
38.8 g. What is the percent yield of C6H5Cl?
C6H6 + Cl2 → C6H5Cl + HCl
5. In a titration, how much of .15 M NaOH is needed to neutralize 20.0 ml of .500M
HCl solution? HCl + NaOH  H2O + NaCl
6. In a titration, what is the molarity of HNO3 if
25.0 ml of it neutralized 15.0
ml of .60M Ca(OH)2
2 HNO3 + Ca(OH)2  2 H2O + Ca(NO3)2
7. What is the difference between end point and equivalence point?
Hot & Cold Packs
Objective 7.1-7.2
1. Sketch an endothermic reaction graph, labeling the reactants, products, activation
energy, activated complex, and the heat of reaction.
2. What is the sign of an endothermic reaction and exothermic reaction?
3. Using the specific heat values for water and iron from reference sheet, which one
would have the largest temperature change if they have the same mass?
4. Calculate the heat released or absorbed during a physical change.
a.
Calculate the heat absorbed when 15.0 g of ice melts to liquid. See reference
sheet for Hfus
b. Calculate the heat released when 75.4 g of vapor condenses into liquid. See
reference sheet for Hvap
5. Calculate the heat released or absorbed in a chemical reaction?
a) What is the specific heat of a metal that releases 2500 J of energy. The
metal has a mass of 25.0 g and had a temperature change of 5.0°C.
b) How much heat is released when iron is dropped in a beaker of water.
The mass of the metal was 43.0 g and the initial temperature of the
metal was 78.0°C. The water temperature changed from 25 C to 32 C.
The specific heat of the metal is .45 J/gC.
c) What is the amount of heat absorbed by water if 23.4 g of water is
heated from 34.0°C to 78.0° C. See reference sheet for specific heat of
water.
d) Write the equilibrium expression for the following reaction.
a) H2(g) + Cl2(g) 2HCl(g) + heat
6. In the process of chemical equilibrium, what stays constant at equilibrium?
7. In the process of equilibrium, are the rates equal to each other?
8. Using the reaction above, answer the following questions regarding Le Chatelier’s
principle.
a) Which direction does the reaction shift if temperature increases?
b) Which direction does the reaction shift if hydrogen gas is increased?
c) Which direction does the reaction shift if HCl is removed?
d) Which direction does the reaction shift if the volume is decreased?
e) Which direction doe the reaction shift if temperature is decreased?
9. If K = .00045, what side of the reaction will be favored?
10. Write the particle symbol that represents alpha, beta and gamma decay?
11. Which of the 3 types of decay is the most harmful? Which one is the heaviest?
Which one is a high speed electron? Which one(s) can be blocked by paper?
12. Define radioactive decay. Do all nuclei undergo this process? Why do they go
through this process?
13. Write a nuclear reaction for the alpha decay or radium-226, the beta decay of
carbon-14; the beta decay of 60Co; the alpha decay of 241Am.
14. What is transmutation?
15. Define half life. What is the algebraic equation to represent this?
16. If gallium-68 has a half life of 68.3 minutes, how much of a 10.0mg sample is left
after 1 half life? Two half-lives?
17. If the passing of five half-lives leaves 25.0mg of strontium-90, how much was
present at the beginning?
18. The half life of polonium-214 is 163.7 microseconds. How much of a 1.0g sample
will remain after 818 microseconds?
19. What is the difference between nuclear fission and nuclear fusion?