Part a how important is this resonance structure to the overall

ΔHrxn =

Exercise 9.25

Part A

How do you determine the number of electrons that go into the Lewis structure of a molecule?

ANSWER:

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Part B

How do you determine the number of electrons that go into the Lewis structure of a polyatomic ion?

ANSWER:

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Exercise 9.38

Exercise 8.18

Part A

What are valence electrons?

ANSWER:

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Part B

Why are they important?

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Exercise 8.22

Exercise 8.52

Part A

Name an element in the third period (row) of the periodic table with three valence electrons.

Express your answer as a chemical symbol.

ANSWER:

Exercise 8.44

Exercise 10.82

Part A

Use the drawing of energy diagram for HCl to predict the bond order.

Express your answer using two significant figures.

ANSWER:

bond order =

Exercise 10.54

The valence electron configurations of several atoms are shown. How many bonds can each atom make without hybridization?

Part A

B2s22p1

ANSWER:

Part B

Why is polarity important?

ANSWER:

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Exercise 10.48

Exercise 10.15

In Lewis theory, the two bonds in a double bond look identical. However, valence bond theory shows that they are not.

Part A

Describe a double bond according to valence bond theory.

ANSWER:

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Part B

Explain why rotation is restricted about a double bond, but not about a single bond.

ANSWER:

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Exercise 10.34

The following figures show several molecular geometries.

Part A

Give the number of total electron groups, the number of bonding groups, and the number of lone pairs for (a) geometry.

Express your answers as integers separated by commas.

ANSWER:

Ntotal, Nbonding, Nlone =

Exercise 11.48

Part A

Explain how doping can increase the conductivity of a semiconductor. What is the difference between an n-type semiconductor and a p-type semiconductor?

ANSWER:

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Exercise 11.98

Determine the coordination number for each of the following structures.

Part A

Gold

ANSWER:

Coordination number =

Score Summary:

Exercise 11.90

The high-pressure phase diagram of ice is shown at the top of the next column. Notice that, under high pressure, ice can exist in several different solid forms .

Part A

What three forms of ice are present at the triple point marked O?

ANSWER:

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Part C

Explain your answers.

ANSWER:

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Exercise 11.80

Part A

Calculate the amount of heat required to completely sublime 99.0g of solid dry ice (CO2) at its sublimation temperature. The heat of sublimation for carbon dioxide is 32.3 kJ/mol.

ANSWER:

Exercise 11.72

A 65.0mL sample of water is heated to its boiling point.

Part A

How much heat (in kJ) is required to vaporize it? (Assume a density of 1.00 g/mL.)

ANSWER:

Exercise 11.68

Exercise 11.24

Part A

How is vapor pressure related to temperature? What happens to the vapor pressure of a substance when the temperature is increased? Decreased?

Check all that apply.

ANSWER:

Check all that apply.

Exercise 11.19

Part A

Why is vaporization endothermic? Why is condensation exothermic?

ANSWER:

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Exercise 11.17

Part A

What is capillary action? How does it depend on the relative strengths of adhesive and cohesive forces?

ANSWER:

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Exercise 11.62

Water (a) “wets” some surfaces and beads up on others. Mercury (b), in contrast, beads up on almost all surfaces.

Part A

Explain this difference.

ANSWER:

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Exercise 11.63

Part E

For the following pair of compounds, pick the one with the higher vapor pressure at a given temperature: CH3OH or H2CO.

ANSWER:

Exercise 11.54

Arrange the following in order of increasing boiling point.

Part A

H2S, H2Se, H2O

ANSWER:

Exercise 11.58

Part A

For the following pair of compounds, pick the one with the higher vapor pressure at a given temperature: CH4 or CH3Cl.

ANSWER:

Part C

For the following pair of compounds, pick the one with the higher vapor pressure at a given temperature: CH3CH2CH2OH or CH3OH.

ANSWER:

Part E

Exercise 11.9

Part A

Why are intermolecular forces generally much weaker than bonding forces?

ANSWER:

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Exercise 6.8

Part A

What is a state function?

ANSWER:

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Part B

List some examples of state functions.

Check all that apply.

ANSWER:

Check all that apply.

Exercise 6.42

A system absorbs 198kJ of heat and the surroundings do 110kJ of work on the system.

Part A

What is the change in internal energy of the system?

ANSWER:

ΔE =

Part B

Why does the reaction mixture undergo a decrease in temperature even though energy is absorbed?

ANSWER:

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Exercise 6.74

Instant cold packs, often used to ice athletic injuries on the field, contain ammonium nitrate and water separated by a thin plastic divider. When the divider is broken, the ammonium nitrate dissolves according to the following endothermic reaction:
NH4NO3(s)→NH+4(aq)+NO−3(aq)
In order to measure the enthalpy change for this reaction, 1.25 g of NH4NO3 is dissolved in enough water to make 25.0 mL of solution. The initial temperature is 25.8 ∘C and the final temperature (after the solid dissolves) is 21.9 ∘C.

Part A

Calculate the change in enthalpy for the reaction in kilojoules per mole. (Use 1.0g/mL as the density of the solution and 4.18J/g⋅∘C as the specific heat capacity.)

Express your answer to two significant figures and include the appropriate units.

ANSWER:

ΔHrxn =

Exercise 7.8

Part A

What determines the color of a colored object?

ANSWER:

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Part B

This question will be shown after you complete previous question(s).

Exercise 7.12

Part A

Describe the photoelectric effect.

ANSWER:

Exercise 7.42

Calculate the wavelength of each of the following frequencies of electromagnetic radiation.

Part A

104.8MHz (typical frequency for FM radio broadcasting)

Express your answer using four significant figures.

ANSWER:

λ1 =

Part B

What forces air into your lungs?

ANSWER:

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Exercise 5.30

The pressure on top of Mt. Everest averages about 235 mmHg.

Part A

Convert this pressure to torr.

ANSWER:

P =

Part B

How is the root mean square velocity of a gas related to its molar mass?

ANSWER:

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Exercise 5.22

Part A

Describe how the molecules in a perfume bottle travel from the bottle to your nose.

ANSWER:

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Part B

What is mean free path?

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Exercise 5.23

Part A

Explain the difference between diffusion and effusion.

ANSWER:

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Part B

How is the effusion rate of a gas related to its molar mass?

ANSWER:

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Exercise 5.84

Exercise 5.90

A sample of N2O effuses from a container in 47 seconds.

Part A

How long would it take the same amount of gaseous I2 to effuse from the same container under identical conditions?

Express your answer using two significant figures.

ANSWER:

t =

Exercise 5.94

Part A

Which postulate of the kinetic molecular theory breaks down under conditions of low temperature?

ANSWER:

Exercise 5.146

Part A

Which of the following gases would you expect to deviate most from ideal behavior under conditions of low temperature: F2, Cl2, Br2?

ANSWER:

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