Chemistry exam review.docx

Chemistry Common Exam Review - Essential Standards

1.1.1 Analyze the structure of atoms, isotopes, and ions.

1. Which best describes the relationship between subatomic particles in any neutral atom?

a. The number of protons equals the number of electrons.

b. The number of protons equals the number of neutrons.

c. The number of neutrons equals the number of electrons.

d. The number of neutrons is greater than the number of protons.

2. What is the nuclear composition of uranium-235?

a. 92 electrons + 143 protons

b. 92 protons + 143 electrons

c. 143 protons + 92 neutrons

d. 92 protons + 143 neutrons

3. Which atomic symbol represents an isotope of sulfur with 17 neutrons?

33/16

4. Draw pictures to represent the isotopes of oxygen, oxygen-16 and oxygen-18.

1.1.2 Analyze an atom in terms of the location of electrons.

1. Which best describes the current atomic theory?

a. Atoms consist of electrons circling in definite orbits around a positive nucleus.

b. Atoms are composed of electrons in a cloud around a positive nucleus.

c. Atoms can easily be split, at which time they become radioactive.

d. An atom’s mass is determined by the mass of its neutrons.

2. Which idea of John Dalton is no longer considered part of the modern view of atoms?

a. Atoms are extremely small.

b. Atoms of the same element have identical masses.

c. Atoms combine in simple whole number ratios to form compounds.

d. Atoms of different elements can combine in different ratios to form different compounds.

3. Which is the electron configuration of calcium?

a. 1s22s22p63s23p8

b. 1s22s22p63s23p64s2

c. 1s22s22p63s23p63d2

d. 1s22s22p63s23p6

4. Predict the electron configuration for the following elements.

1. K

2. Cl

3. Ni

4. Ne

1.1.3 Explain the emission of electromagnetic radiation in spectral form in terms of the Bohr model.

1. Which color of light would a hydrogen atom emit when an electron changes from the n = 5 level to the n = 2 level?

a. red

b. yellow

c. green

d. blue

2. Which statement regarding red and green visible light is correct?

a. The speed of green light is greater than that of red light.

b. The wavelength of green light is longer than that of red light.

c. The energy of green light is lower than that of red light.

d. The frequency of green light is higher than that of red light.

3. What energy level transition is indicated when the light emitted by a hydrogen atom has a wavelength of 103 nm?

a. n = 2 to n = 1

b. n = 3 to n = 1

c. n = 4 to n = 2

d. n = 5 to n = 2

4. Consider the spectrum for the hydrogen atom. In which situation will light be produced?

a. Electrons absorb energy as they move to an excited state.

b. Electrons release energy as they move to an excited state.

c. Electrons absorb energy as they return to the ground state.

d. Electrons release energy as they return to the ground state.

5. An electron in an atom of hydrogen goes from energy level 6 to energy level 2. What is the wavelength of the electromagnetic radiation emitted?

a. 410 nm

b. 434 nm

c. 486 nm

d. 656 nm

6. Use the Bohr model to explain the release of energy in the return of electrons to ground state.

1.1.4 Explain the process of radioactive decay by the use of nuclear equations and half-life.

1. Which will complete this equation?

2. Which particle will complete this reaction?

a. electron

b. neutron

c. nucleus

d. proton

3. In the figure below, what type of nuclear activity is represented?

a. fission

b. fusion

c. alpha emission

d. beta emission

4. The half-life of phosphorus-32 is 14.30 days. How many milligrams of a 20.00 mg sample of phosphorus-32 will remain after 85.80 days?

a. 3.333 mg

b. 0.6250 mg

c. 0.3125 mg

d. 0.1563 mg

5. Consider this diagram:

Which of the three types of radiation will penetrate the paper and wood?

a. alpha, beta and gamma

b. alpha and beta only

c. gamma only

d. beta only

6. The half-life of a radioactive isotope is 20 minutes. What is the total amount of time of 1.00 g of sample of this isotope remaining after 1 hour?

a. 0.500 g

b. 0.333 g

c. 0.250 g

d. 0.125 g

7. Explain how to use M&M-type candies to map a decay plot for a hypothetical decay. Begin with 20 “atoms” and show the plot for your experiment.

1.2.1 Compare (qualitatively) the relative strengths of ionic, covalent, and metallic bonds.

1. Construct an energy diagram that indicates the relative energies of the different types of bonds.

2. Which statement compares the amount of energy needed to break the bonds in CaCl2 (E1) and C12H22O11 (E2)?

a. E1>E2, as CaCl2 is a covalent compound.

b. E1

c. E1>E2, as CaCl2 is an ionic compound.

d. E1

1.2.2 Infer the type of bond and chemical formula formed between atoms.

1. Which pair of elements would most likely bond to form a covalently bonded compound?

a. sodium and fluorine

b. barium and chlorine

c. phosphorus and oxygen

d. magnesium and sulfur

2. For each pair of atoms, predict whether the bond formed between the atoms is either ionic or covalent, and write the formula for the predicted compound.

1. Na and O

2. S and F

3. Ag and N

4. Te and H

3. Which statement describes the compound formed between sodium and oxygen?

a. It is NaO2, which is ionic.

b. It is NaO2, which is covalent.

c. It is Na2O, which is ionic.

d. It is Na2O, which is covalent.

1.2.3 Compare inter- and intra- particle forces.

1. Rank the following substances in the order in which they would evaporate, justifying the order of placement for each (using those inter- and intra- particle forces).

1. Water (H2O)

2. Methane (CH4)

3. Sodium chloride (NaCl)

4. Phosphorus trifluoride (PF3)

2. At STP, fluorine is a gas and iodine is a solid. Why?

a. Fluorine has lower average kinetic energy than iodine.

b. Fluorine has higher average kinetic energy than iodine.

c. Fluorine has weaker intermolecular forces of attraction than iodine.

d. Fluorine has stronger intermolecular forces of attraction than iodine.

1.2.4 Interpret the name and formula of compounds using IUPAC convention.

1. What is the name of the compound PbO2?

a. lead oxide

b. lead (II) oxide

c. lead oxide (II)

d. lead (IV) oxide

2. Which is the correct formula for dinitrogenpentoxide?

a. N4O

b. NO2

c. N2O5

d. NO4

3. What is the name of HCl(aq)?

a. chloric acid

b. hydrochloric acid

c. hydrogen chloride

d. perchloric acid

4. What is the chemical formula for calcium nitrate?

a. CaNO3

b. Ca(NO2)2

c. Ca(NO3)2

d. Ca3N2

5. Given the IUPAC name of a compound, infer its formula, (1-3) and given a formula and write the IUPAC name (4-6), recognizing the differing nomenclature systems for ionic and covalent compounds.

1. iron (III) chloride

2. magnesium oxide

3. carbon tetrachloride

4. N2O5

5. Na2SO4

6. NH4HCO3

6. What is the IUPAC name for the compound represented by the formula Mg(OH)2?

a. magnesium hydroxide

b. magnesiumdihydroxide

c. magnesium (II) hydroxide

d. magnesium (II) dihydroxide

1.2.5 Compare the properties of ionic, covalent, metallic, and network compounds.

1. Based on the VSEPR theory, what is the molecular geometry of a molecule of PI3?

a. linear

b. tetrahedral

c. trigonal planar

d. trigonal pyramidal

2. Which is a unique characteristic of the bonding between metal atoms?

a. Atoms require additional electrons to reach a stable octet.

b. Atoms must give away electrons to reach a stable octet.

c. Atoms share valence electrons only with neighboring atoms to reach a stable octet.

d. Delocalized electrons move among many atoms creating a sea of electrons.

3. What type of bonding is associated with compounds that have the following characteristics?

· high melting points

· conduct electricity in the molten state

· solutions conduct electricity

· normally crystalline solids at room temperature

a. covalent

b. ionic

c. hydrogen

d. metallic

4. Based on your knowledge of the following groups: network solids, covalent compounds (polar and non-polar), ionic solids, metallic solids, classify each substance based on the data given.

Melting Point

(high or low)

Boiling Point

(high or low)

Soluble

in water

(yes or no)

Conducts Electricity in Solid Form

(yes or no)

Conducts Electricity in Water

(yes or no)

Classification of Compound

Brass (an alloy of zinc and copper)

Graphite

Potassium bromide

Carbon tetrachloride

5. An unknown substance is tested in the laboratory. The physical test results are listed below.

· nonconductor of electricity

· insoluble in water

· soluble in oil

· low melting point

Based on these results, what is the unknown substance?

a. ionic and polar

b. ionic and nonpolar

c. covalent and polar

d. covalent and nonpolar

1.3.1 Classify the components of a periodic table (period, group, metal, metalloid, nonmetal, transition).

1. The compound formed between element X and oxygen has the chemical formula X2O. Which element would X most likely represent?

a. Fe

b. Zn

c. Ag

d. Sn

2. Classify each element as a metal (M), nonmetal (NM) or metalloid (MD). Identify which elements are transition metals. Identify each element’s group and period.

M, NM, MD

Transition?

Group

Period

Ca

O

H

W

As

In

Rn

3. The nucleus of an atom is shown.

Which statement describes the element?

a. It is a nonmetal in group 15.

b. It is a nonmetal in group 16.

c. It is a nonmetal in group 2.

d. It is a nonmetal in group 17.

1.3.2 Infer the physical properties (atomic radius, metallic and nonmetallic characteristics) of an atom based on its position on the periodic table.

1. Which best explains why cations are smaller than the atoms from which they are formed?

a. The metallic atom gains electrons, causing a larger effective nuclear pull.

b. The metallic atom loses electrons, resulting in loss of an entire energy level.

c. The nonmetallic atom gains electrons, causing a larger effective nuclear pull.

d. The nonmetallic atom loses electrons, resulting in loss of an entire energy level.

2. What will the properties of element 117 be when it is discovered?

3. Which atom has the largest radius? Justify your answer.

a. Bromine

b. Chlorine

c. Selenium

d. Sulfur

1.3.3 Infer the atomic size, reactivity, electronegativity and ionization energy of an element from its position in the periodic table.

1. Which electron configuration represents a transition element?

a. 1s22s22p3

b. 1s22s22p63s2

c. 1s22s22p63s23p64s23d7

d. 1s22s22p63s23p64s23d104p4

2. Given the electron configuration of 1s22s22p4, how many electrons does this element have in its outer level?

a. 2

b. 4

c. 6

d. 8

3. Which correctly lists four atoms from smallest to largest radii?

a. I, Br, Cl, F

b. F, I, Br, Cl

c. Si, P, S, Cl

d. Cl, S, P, Si

4. Which have the lowest electronegativities?

a. alkali metals

b. halogens

c. rare earth elements

d. transition metals

5. Arrange the following elements in order of increasing electronegativity from lowest to highest: F, K, Si, and S.

a. F < K < S < Si

b. K < SI < S < F

c. Si < F < K < S

d. S < Si < F < K

6. Compare the elements in group 2 of the periodic table. Include a description of the atomic sizes, reactivity, electronegativities, and ionization energies. Then, do the same comparison for period 3.

2.1.1 Explain the energetic nature of phase changes.

1. A piece of metal is heated in a Bunsen burner flame and then immersed in a beaker of cool water.

Which statement best describes the effect of the temperature changes on the kinetic energy of the particles?

a. Kinetic energy of metal atoms decreases in the flame.

b. Kinetic energy of water molecules increases when the heated metal is immersed.

c. Kinetic energy of water molecules decreases when the heated metal is immersed.

d. Kinetic energy of metal atoms increases when immersed in the cooler water.

2. The gases helium, neon, and argon are in separate containers at 55°C. Which is true about the kinetic energy of the gases?

a. Helium has the lowest mass and therefore the greatest kinetic energy.

b. They each have a different kinetic energy.

c. Argon has greatest mass and therefore the greatest kinetic energy.

d. They all have the same average kinetic energy.

3. An open container of water is brought to a boil and heated until all of the water is converted to water vapor. Which describes the changes in the water molecules?

a. The molecules speed up and move farther apart.

b. The molecules speed up and move closer together.

c. The molecules slow down and move farther apart.

d. The molecules slow down and move closer together

4. What happens when energy is removed from liquid water?

a. Molecules slow down, and more hydrogen bonds are formed.

b. Molecules slow down, and more hydrogen bonds are broken.

c. Molecules move faster, and more hydrogen bonds are formed.

d. Molecules move faster, and more hydrogen bonds are broken.

5. What causes the process of perspiration to be cooling for human skin?

a. It involves condensation and is exothermic.

b. It involves evaporation and is exothermic.

c. It involves condensation and is endothermic.

6. When is a liquid considered to be boiling? Explain your reasoning.

A liquid boils at a temp at which its vapor pressure is equal to the pressure of the gas above it.

2.1.2 Explain heating and cooling curves (heat of fusion, heat of vaporization, heat, melting point, and boiling point)

1. This is a heating curve for a substance.

Between points X and Y, which would be observed?

a. Solid and liquid will be present.

b. Only vapor will be present.

c. Liquid and vapor will be present.

d. Only liquid will be present.

2. Given the heating curve below, what is occurring between minutes 6 to 12?

a. There is an increase in kinetic energy and vaporization is occurring.

b. There is an increase in kinetic energy and condensation is occurring.

c. There is an increase in potential energy and freezing is occurring.

d. There is an increase in potential energy and melting is occurring

2.1.3 Interpret the data represented in phase diagrams.

1. Consider this phase diagram:

At what temperature does the normal boiling point occur?

a. 45°C

b. 60°C

c. 100°C

d. 110°C

2. Consider this phase diagram:

What process is occurring when a substance changes from point X (-130°C and 50 kPa) to point Y (30°C and 100 kPa)?

melting

3. Compare these phase diagrams. What can be said about the relationship between the processes of melting for the two substances above? What can you determine about the densities of the solids compared to the liquids of each substance?

4. According to the phase diagram below, what is the boiling point of this substance at a pressure of 80 kPa?

a. 100°C

b. 50°C

c. -75°C

d. -90°C

2.1.4 Infer simple calorimetric calculations based on the concepts of heat lost equals heat gained and specific heat.

1. 6.00g of gold was heated from 20.0°C to 22.0°C. How much heat was applied to the gold?

a. 1.55 J

b. 15.5 J

c. 17.0 J

d. 32.5 J

2. A student has a beaker containing 55 g of water at 100°C. How much heat is needed to convert the water to steam?

a. 120,000 J

b. 18,000 J

c. 2,200 J

d. 330 J

3. How many grams of ice will melt at 0°C if the ice absorbs 420 J of energy?

a. 0.186 g

b. 0.795 g

c. 1.26 g

d. 5.34 x 104 g

4. An 18.0 g piece of an unidentified metal was heated from 21.5°C to 89.0°C. If 292 J of heat energy was absorbed by the metal in the heating process, what was the identity of the metal?

Specific Heat Table

Substance

Specific Heat

Aluminum

0.90 J/g·°C

Calcium

0.65 J/g·°C

Copper

0.39 J/g·°C

Gold

0.13 J/g·°C

Iron

0.46 J/g·°C

Mercury

0.14 J/g·°C

Silver

0.24 J/g·°C

a. calcium

b. copper

c. iron

d. silver

5. 1000J of heat is added to 2g of the following substances. Which one will produce the biggest change in temperature?

a. aluminum

b. copper

c. iron

d. lead

6. An 8.80g sample of metal is heated to 92.0°C and then added to 15.0g of water at 20.0°C in an insulated calorimeter. At thermal equilibrium the temperature of the system was measured as 25.0°C . What is the identity of the metal?

2.1.5 Explain the relationships among pressure, temperature, volume, and quantity of gas, both quantitative and qualitative.

1. What happens to the pressure of a constant mass of gas at constant temperature when the volume is doubled?

a. The pressure is doubled.

b. The pressure remains the same.

c. The pressure is reduced by ½.

d. The pressure is reduced by ¼.

2. The total pressure in a closed vessel containing N2, O2, and CO2 is 30atm. If the partial pressure of N2 is 4 atm and the partial pressure of O2 is 6 atm, what is the partial pressure of of CO2?

a. 20 atm

b. 30 atm

c. 40 atm

d. 50 atm

3. What is the pressure, in atmospheres, exerted by a 0.100-mol sample of oxygen in a 2.00L container at 273°C?

a. 4.48 x 10-1atm

b. 2.24 x 100atm

c. 1.12 x 103atm

d. 2.24 x 103atm

4. What causes an inflated balloon to shrink when it is cooled?

a. because cooling the balloon causes gas to escape from the balloon

b. because cooling the balloon causes the gas molecules to collide more frequently

c. because cooling the balloon makes the gas molecules become smaller.

d. because cooling the balloon causes the average kinetic energy of the gas molecules to decrease

5. The Kelvin temperature and the pressure of a sample of gas are doubled. What will be the effect on the density of the gas?

2.2.1 Explain the energy content of a chemical reaction

1. When a chemical cold pack is activated, it becomes cool to the touch. What is happening in terms of energy?

a. An exothermic reaction is occurring, absorbing heat from its surroundings.

b. An exothermic reaction is occurring, releasing heat to its surroundings.

c. An endothermic reaction is occurring, releasing cold to its surroundings.

d. An endothermic reaction is occurring, absorbing heat from its surroundings.

2. This graph represents the change in energy for the two laboratory trials of the same reaction:

Which factor could explain the energy differences between the trials?

a. Heat was added to trial #2.

b. A catalyst was added to trial #2.

c. Trial #1 was stirred.

d. Trial #1 was cooled.

3. Given the energy diagram below, which statement best describes the forward reaction?

a. It is an exothermic reaction with an energy change of 160kJ.

b. It is an exothermic reaction with an energy change of 80kJ.

c. It is an endothermic reaction with an energy change of 160kJ.

d. It is an endothermic reaction with an energy change of 80kJ.

4. What type of reaction is represented by the energy diagram below?

Label the location of the energy of reactants, energy of products, activation energy, and enthalpy (heat of reaction). If a catalyst were added to this reaction, what quantities would change? Justify your reasoning.

2.2.2 Analyze the evidence of chemical change.

1. Which example indicates that a chemical change has occurred?

a. When aqueous solutions are mixed, a precipitate is formed.

b. As ammonium nitrate dissolves in water, it causes the temperature to decrease.

c. Alcohol evaporates when left in an open container.

d. Water is added to blue copper(II) chloride solution. The resulting mixture is lighter blue in color.

2. A student mixes two chemicals in a test tube. The test tube turns hot and bubbles appear. What indicators of chemical reaction is the student observing?

a. Change in color and formation of precipitate.

b. Change in color and formation of gas.

c. Change in temperature and formation of precipitate.

d. Change in temperature and formation of gas.

3. Select the indicators of chemical reactions that would help you distinguish between these two reactions. Write a balanced chemical equation for each reaction (include phases). Identify each type of reaction.

a. Sodium metal dropped into a beaker of water.

b. Silver nitrate is added to sodium chloride.

2.2.3 Analyze the law of conservation of matter and how it applies to various types of chemical equations (synthesis, decomposition, single replacement, double replacement, and combustion).

1. Consider this reaction:

NH3(g) + HCl (g) → NH4Cl (s)

Which type of reaction does this equation represent?

a. combustion

b. decomposition

c. single replacement

d. synthesis

2. Which equation represents a single replacement reaction that can occur?

a. F2 + 2NaCl → 2NaF + Cl2

b. Cl2 + 2NaF → 2NaCl + F2

c. Cu + 2NaCl → CuCl2 + 2Na

d. Zn + 2NaF → ZnF2 + 2Na

3. What products are formed when the metal potassium is added to water?

a. K and H2O

b. KOH and H2O

c. K2O and H2

d. KOH and H2

4. When Na2O reacts with H2O, what is produced?

a. HNaO2

b. Na + H2O

c. NaO + H2

d. NaOH

5. Which equation is correctly balanced?

a. Cu + H2SO4 → CuSO4 + H2O + SO2

b. 2Na + 2H2O → 2NaOH + H2

c. 2Fe + 3O2 → Fe2O3

d. 4Cu + S8 → 8Cu2S

6. What coefficients are required to balance this reaction?

__Fe2O3 + __CO → __Fe + __ CO2

a. 2, 6, 3, 6

b. 1, 3, 2, 3

c. 1, 1, 2, 2

d. 1, 1, 2, 1

7. An aqueous solution of silver nitrate is added to an aqueous solution of iron(II) chloride. Which is the net ionic equation for the reaction that occurs?

a. AgNO2 (aq) + FeCl (aq) → AgCl(s) + FeNO2 (aq)

b. 2AgNO3 (aq) + FeCl2 (aq) → 2AgCl (s) + Fe(NO3)3 (aq)

c. 2 Ag+1(aq) + NO3-1(aq) + Fe+2(aq) + Cl2 (g) → 2AgCl (s)

d. 2Ag+(aq) + 2Cl-(aq) → 2AgCl (s)

8. Consider this combustion reaction equation:

C4H10 + O2 → CO2 + H2O

When the equation is balanced, what will be the coefficient of O2 ?

a. 1

b. 7

c. 10

d. 13

9. 10.3 grams of sodium hydrogen carbonate reacts with an excess of hydrochloric acid. A white crystalline substance is produced and the mass of the product is 7.59 g.

1. What type of reaction occurred?

2. Write the balanced chemical equation for this reaction.

3. What is the identity of the white crystalline product?

4. Based on the data from the reaction, determine the molar ratio between the given reactant and product.

2.2.4 Analyze the stoichiometric relationships inherent in a chemical reaction.

1. How many moles are in 59.6g of BaSO4?

a. 0.256 mole

b. 3.91 moles

c. 13.9 moles

d. 59.6 moles

2. What is the volume of two moles of hydrogen gas at STP?

a. 44.8 L

b. 22.4 L

c. 11.2 L

d. 2.00 L

3. How many molecules are contained in 55.0g of H2SO4?

a. 0.561 molecule

b. 3.93 molecules

c. 3.38 x 1023 molecules

d. 2.37 x 1024 molecules

4. If a sample of magnesium has a mass of 60. g, how many moles of magnesium does the sample contain?

a. 1.1 moles

b. 1.2 moles

c. 2.0 moles

d. 2.5 moles

5. Consider this reaction:

3Ca(s) + 2H3PO4(aq) → Ca3(PO4)2 (s) + 3H2 (g)

How many moles of calcium are required to produce 60.0g of calcium phosphate?

a. 0.145 mole

b. 0.194 mole

c. 0.387 mole

d. 0.581 mole

6. Metallic sodium reacts violently with water to form hydrogen and sodium hydroxide according to the following balanced equation:

2Na + 2H2O → 2NaOH + H2

How many moles of hydrogen gas are generated when 4.0 moles of sodium react with excess water?

a. 1.0 mole

b. 2.0 moles

c. 3.0 moles

d. 4.0 moles

7. According to the equation 2H2O(l) → 2H2 (g) +O2 (g), what mass of H2O is required to yield 22.4L of O2 at STP?

a. 12g

b. 18g

c. 24g

d. 36g

8. Consider this reaction:

3Mg(s) + 2H3PO4 (aq) → Mg3(PO4)2 (s) + 3H2 (g)

How many grams of magnesium phosphate should be produced if 5.40 grams of magnesium react with excess phosphoric acid?

a. 1.80 grams

b. 19.5 grams

c. 58.4 grams

d. 175 grams

9. Methane is burned in oxygen according to this balanced chemical equation:

CH4 (g) + 2O2 (g) → CO2 (g) + 2H2O (g)

What volume of carbon dioxide is formed when 9.36 liters of methane are burned in excess oxygen at STP?

a. 9.36 L

b. 15.0 L

c. 18.7 L

d. 22.4 L

10. Given the balanced chemical equation for the reaction,

P4 + 5O2 → P4O10

What mass of oxygen is needed to completely react with 7.75 g P4?

a. 2.00 grams

b. 5.00 grams

c. 10.00 grams

d. 40.00 grams

11. A 70.0 sample of limestone consists of a large percentage of calcium carbonate. The sample reacts with an excess of hydrochloric acid and 14.0 L of carbon dioxide is generated at STP. What is the percentage of calcium carbonate in the limestone? Write the balanced chemical equation for this reaction.

2.2.5 Analyze quantitatively the composition of a substance (empirical formula, molecular formula, percent composition, and hydrates).

1. Analysis shows a compound to be, by mass, 43.8% N, 6.2% H, and 50.0% O. Which is the possible molecular formula for the substance?

a. NH4NO2

b. NH4NO3

c. NH3OH

d. N2OH

2. A compound has an empirical formula of CH2O and a molecular mass of 180g. What is the compound’s molecular formula?

a. C3H6O3

b. C6H12O6

c. C6H11O7

d. C12H22O11

3.What is the percent by mass of iron in the compound Fe2O3?

a. 70%

b. 56%

c. 48%

d. 30%

4. A compound consisting of 56.38% phosphorus and 43.62% oxygen has a molecular mass of 220 g/mole. What is the molecular formula of this compound?

a. PO

b. PO2

c. P2O3

d. P4O6

5. A 10.10 g sample of barium chloride hydrate is heated in a crucible. After all of the water is driven off, 8.50 g of the anhydrous barium chloride remains in the crucible. What is the formula of the hydrate?

3.1.1 Explain the factors that affect the rate of a reaction (temperature, concentration, particle size and presence of a catalyst).

1. Which statement explains why the speed of some reactions is increased when the surface area of one or all the reactants is increased?

a. increasing surface area changes the electronegativity of the reactant particles

b. increasing surface area changes the concentration of the reactant particles

c. increasing surface area changes the conductivity of reactant particles

d. increasing surface area enables more reactant particles to collide

2. Consider the balanced chemical equation:

2H2O2(aq) 2H2O (l) + O2(g)

Which will increase the rate of the reaction?

a. increasing pressure on the reaction

b. decreasing concentration of the reactants

c. adding a catalyst to the reaction

d. decreasing the temperature of the reaction

3. For the reaction

A+ (aq) + B- (aq) --> AB (s)

increasing the temperature increases the rate of the reaction. Which is the best explanation for this happening?

a. The pressure increases, which in turn increases the production of products.

b. The concentration of reactants increases with an increase in temperature.

c. The average kinetic energy increases, so the likelihood of more effective collisions between ions increases.

d. Systems are more stable at high temperatures.

4. Write a letter to a fellow scientist who missed the lecture on how these factors affect the rate of a chemical reaction. Give examples of each (particle size, temperature, concentration, and presence of catalyst).

5. When a set amount of marble chips (CaCO3) is added to a small amount of dilute hydrochloric acid, a reaction occurs. What should be done to decrease the rate of reaction the next time the experiment is performed?

a. Use more acid.

b. Stir.

c. Use larger marble chips.

d. Add heat.

3.1.2 Explain the conditions of a system at equilibrium.

1. Make a “How-To” poster for laboratory workers to use in determining if a system is at equilibrium.

2. A scientist observes a chemical reaction as it takes place. What can the scientist do in order to tell if the reaction has achieved equilibrium?

a. Measure concentrations of products and reactants over time.

b. Monitor the temperature of the reaction over time.

c. Measure the pH of the solution over time.

d. Wait for the formation of a precipitate.

3.1.3 Infer the shift in equilibrium when a stress is applied to a chemical system (Le Chatelier’s Principle).

1. For the following reaction, predict the direction the equilibrium will shift for each change indicated.

C2H2 + Br2 C2H2Br2 + heat

1. Add Br2

2. Increase pressure

3. Increase temperature

4. Remove product

2. For the reaction

2SO2(g) + O2(g)2SO3(g) + heat

Which action will increase the concentration of SO3?

a. removing SO2

b. increasing the temperature

c. increasing the pressure

d. adding a catalyst

3.2.1 Classify substances using the hydronium and hydroxide ion concentration.

1. What is the pH of a solution of KOH with a hydroxide concentration of [OH] = 1 x 10-4 M?

a. -10

b. -4

c. 4

d. 10

2. A water sample was found to have a pH of 6 at 25 degrees. What is the hydroxide concentration in the water sample?

a. 1 x 10-8 M

b. 6 x 10-8 M

c. 1 x 10-6 M

d. 6 x 10-6 M

3. Based on hydroxide concentration, which unknown substance would be an acid?

a. Substance A, [OH-] = 1.0 x 10-2 M

b. Substance B, [OH-] = 1.0 x 10-4 M

c. Substance C, [OH-] = 1.0 x 10-6 M

d. Substance D, [OH-] = 1.0 x 10-8 M

4. For each substance and their hydronium ion concentrations, classify them as acidic, basic, or neutral and justify your choice.

1. Substance A, [H3O+] = 1.0 x 10-7 M

2. Substance B, [H3O+] = 1.0 x 10-10 M

3. Substance C, [H3O+] = 1.0 x 10-3 M

4. Substance D, [H3O+] = 1.0 x 10-13 M

3.2.2 Summarize the properties of acids and bases.

1. Phenolphthalein is an indicator that turns pink when added to a basic solution. In which solution would phenolphthalein turn pink?

a. NaOH

b. HCl

c. H2O

d. NaCl

2. Consider this chemical equation:

NH3(aq) + HCl(aq) → NH4+(aq) + Cl-(aq)

In this reaction, why is the ammonia considered a base?

a. NH3 increases the hydronium ion concentration.

b. NH3 decreases the hydroxide ion concentration.

c. NH3 accepts a proton.

d. NH3 donates a proton.

3. Given the data table below, which unknown substance would be an acid?

Substance

W

X

Y

Z

Tastes bitter

?

Yes

Yes

No

Tastes sour

No

No

?

Yes

Feels slippery

No

Yes

Yes

?

Turns litmus blue

Yes

Yes

Yes

?

Turns litmus red

?

No

No

Yes

a. Substance W

b. Substance X

c. Substance Y

d. Substance Z

4. Make trading cards (like baseball cards) for three “people”, Anthony A. Acid, Betty B. Base, and I.M. Neutral. List stats (properties) and other pertinent info (such as interactions, concentration ranges, pH’s, etc.).

3.2.3 Infer the quantitative nature of a solution (molarity, dilution, and titration with a 1:1 molar ratio).

1. In a titration experiment, if 30.0 mL of an HCl solution reacts with 24.6 mL of a 0.50 M NaOH solution, what is the concentration of the HCl solution?

a. 0.41 M

b. 0.61 M

c. 1.5 M

d. 370 M

2. How many grams of KCl are necessary to prepare 1.50 liters of a 0.500 M solution of KCl?

a. 224 g

b. 74.6 g

c. 56.0 g

d. 24.9 g

3. What is the molarity of a solution containing 20.0 g of sodium hydroxide dissolved in 1.00 L of solution?

a. 0.500 M

b. 0.400 M

c. 0.300 M

d. 0.200 M

4. What volume of 0.200M HCl will neutralize 10.0mL of 0.440M KOH?

a. 40.0mL

b. 20.0mL

c. 8.00mL

d. 5.00mL

5. 25.0mL is diluted to a total volume of 1.00L. What is the concentration of the newly diluted solution? Justify your answer.

3.2.4 Summarize the properties of solutions.

1. If the volume of an 18.5 g piece of metal is 2.35 cm3, what is the identity of the metal?

a. iron

b. lead

c. nickel

d. zinc

2. Which substance listed in the table is a liquid at 27oC?

Melting Point (oC)

Boiling Point (oC)

I

28

140

II

-10

25

III

20

140

IV

-90

14

a. I

b. II

c. III

d. IV

3. Heat is added to a solution to

a. increase the solubility of a solid solute

b. increase the solubility of a gas solute

c. increase the miscibility of the solution

d. increase the degree of saturation of the solution

4. Make a PowerPoint Presentation of at least 8 slides to illustrate these general properties of solutions: solubility, miscibility, concentration and degree of saturation.

3.2.5 Interpret solubility diagrams.

1. Why does the solubility of NH3 decrease as the temperature increases? Explain this on a molecular level.

2. How many grams of KCl are required to make a saturated solution in 50.0g of water 80°C?

a. 25.0g

b. 50.0g

c. 100.g

d. 150.g

3.2.6 Explain the solution process.

1. Which will increase the solubility of most solid solutes?

a. decreasing the temperature

b. decreasing the amount of solvent at constant temperature

c. increasing the amount of solute at constant temperature

d. increasing the temperature

2. When considering the energetics of the solution process, which process is always exothermic?

a. Solute particles are separate from one another.

b. Solvent particles separate from one another.

c. Solute and solvent particles form attractions for one another.

d. solution formation as a whole is always endothermic.

3. Develop and make an illustrated storyboard for a three-minute short film that explains how solutions form on a particulate level. Use minimum of 6 vignettes.

Studying Strategies

1 Come to review sessions. Take notes and ask questions about items you may be confused about.

2 Study in small sessions over several nights instead of trying to cram the studying into one or two nights.

3 Make sure that you understand the material well, don't just read through the material and try to memorize everything.

4 You may want to form a study group with your friends so that you can help each other with concepts . If you choose to study in a group, only study with others who are serious about the test.

5 If you get all of the questions correct in a section of the review, don’t spend too much time studying that section. Focus your energy on studying for areas where you missed one or two questions.

Test-Taking Strategies

1 RTDQ- Read the darn question! Read the question and (for multiple choice) all of the answers. Answer each question completely, including any explanation required, or choose the best answer for a multiple choice question.

2 Answer every question.

a Eliminate any wrong answers to narrow down your choices for multiple choice questions.

b Check the reference table for helpful information (remember that the formulas and symbols are in there!)

c Look at other questions for clues to help you answer.

d If necessary, guess for multiple choice. Put SOMETHING down for the constructed response portion - just do your best.

3 Monitor your time.

a When you first receive your test, do a quick survey of the entire test so that you know how to efficiently budget your time.

b Pace yourself, don't rush . Read the entire question and pay attention to the details.

c Don’t worry about the time too much, but don’t spend too much time on one question. If you get stuck on a question, put a star by it and move on, then come back to it after you have answered the remaining questions.

d If you have time left when you are finished, look over your test. Make sure that you have answered all the questions. Only change an answer if you misread or misinterpreted the question because the first answer that you put is usually the correct one. Watch out for careless mistakes and proofread your short answer questions.

4 Write legibly. If the grader can't read what you wrote, they'll most likely mark it wrong.

5 BRING A CALCULATOR AND PENCIL (with a good eraser!) TO THE EXAM!!!!