Note: Textbook references pertaining to standards are...

NWHS-Aversa Second Semester Chemistry Points of Interest NAME: _______________________________

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Note: Textbook references pertaining to standards are in blue and red font.

Vocabulary

Conduction

Kinetic molecular model

Kelvin temperature

Order (of atoms/molecules)

Pressure-temperature relationship

Pressure-volume relationship

Rotational motion

Temperature-volume relationship

Translational motion

Vibrational motion

C2.2A Describe conduction in terms of molecules bumping into each other to transfer energy.

Explain why there is better conduction in solids and liquids than gases. Practice: page 199; 1, 2, 3, page 217; 13, 14

Section(s) 9.1,9.4,9. 5,9.7,9.8 Page(s)

197,202,204,207,211

1. C2.2A A student has two identical solid metal samples at 900C. She carefully places one of the metal

samples into a beaker of water at 100C and the other into an empty beaker (only air) at 10

0C. Why would the

metal sample placed into water cool faster than the one placed into air?

A. Water has faster-moving molecules than air, enabling more heat transfer by conduction.

B. Water has more molecules per unit volume than air, enabling more heat transfer by conduction.

C. Air is made mostly of nitrogen which is too heavy an element to conduct heat.

D. Air is a gas and gases do not conduct heat.


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C2.2B Describe the various states of matter in terms of the motion and arrangement of the

molecules (atoms) making up the substance. Practice: page 217; 11 Chapter 9 Page(s) 197

2. C2.2B Choose all that apply:

_________Liquid state a. atoms are close to each other

_________Solid state b. atoms have freedom to move relative to each other

_________Gas state c. atoms are in a very disordered state

d. atoms are in a somewhat disordered state

e. atoms are in a very ordered state

f. atoms are moving very fast with much freedom

3. From the following diagram, identify the liquid state of matter and explain why you chose it.

C2.2c Explain changes in pressure, volume, and temperature for gases using the kinetic

molecular model. Practice: page 241; 17, page 314; 52 Chapter 10 Page(s) 221-228,

229-233

4. C2.2c Explains why damage to aerosol cans can occur if they are exposed to very high temperatures?

C3.3A Describe how heat is conducted in a solid.

Chapter 9

Section 9.8 10

Teacher Supplement

5. C3.3A Explain how heat is conducted in solids. Draw a diagram.


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C3.3B Describe melting on a molecular level.

Practice: Page 217; 23, 24, 33, 35, 37 (a-b)

Section(s) 9.8 Page(s) 210, 213, 214

6. C3.3B The graph below represents the relationship between temperature and time as heat is added to a

sample of water. Explain what happens between each interval using the following terms; Vibrational energy,

potential energy average kinetic energy, rotational energy, increases, decreases, remains the same.

C4.3A Recognize that substances that are solid at room temperature have stronger attractive

forces than liquids at room temperature, which have stronger attractive forces than gases

at room temperature. Practice: Page 217; 22(a-b), Page 218; 54

Chapter 9 Page (s) 204-206, 210

7. C4.3A Using the above language explain why chlorine is a gas, bromine a liquid, and iodine a solid at room

temperature.


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C4.3B Recognize that solids have a more ordered, regular arrangement of their particles than

liquids and that liquids are more ordered than gases. Practice: Page 217; 22 (a-b) Chapter 1 Page 6 Chapter 9 Page 198 fig. 9-2,

Page 204 fig. 9-8,

Page 210 dif. 8-18

8. C4.3B The temperature of an unknown substance was measured as it cooled. The temperature of the

substance over time was graphed and the graph was divided into five different zones, as shown below.

Describe what is happening in terms of particles between the different states of matter this substance goes

through. For each of the segments on the graph identify the state of matter, the arrangement of particles (most

order, somewhat ordered, most disordered).


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C4.5a Provide macroscopic examples, atomic and molecular explanations, and mathematical

representations (graphs and equations) for the pressure-volume relationship in gases. Practice: Page 241; 15, 16, Page 227; 4

Section(s) 10.6 Page(s) 226

9. C4.5a If a balloon is squeezed, what happens to the pressure of the gas inside the balloon? Draw a before

and after picture of it including the gas particles.

C4.5b Provide macroscopic examples, atomic and molecular explanations, and mathematical

representations (graphs and equations) for the pressure-temperature relationship in gases. Practice: Page 241; 17, Page 231; 6


11. C4.5b What happens to the pressure of a gas inside a container if the temperature of the gas decreases?

C4.5c Provide macroscopic examples, atomic and molecular explanations, and mathematical

representations (graphs and equations) for the temperature-volume relationship in gases. Practice: Page 230; 5


12. C4.5c Assuming pressure is held constant, which of the following graphs shows how the volume of an

ideal gas changes with temperature?


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Vocabulary

Atomic weight

Boiling point

Chemical bond

Dipole-dipole bond

Dispersion forces

Endothermic process

Exothermic process

Hydrogen bonding

Ion

Ionic solid (crystal)

Melting point

Metal

Network solid

Relative mass

Release of energy

Temporary dipole

C4.3c Compare the relative strengths of forces between molecules based

on the melting point and boiling point of the substances.

Practice: Page 353; 1

Section(s) 14.14,15.5

Page(s) 343,353

1. C4.3c The hydrocarbons in the table below are all liquids at room temperature. Compare the relative

strengths of the forces between these molecules.

Which molecule has the highest intermolecular forces?


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Claim:

Evidence:

Reasoning:

C4.3d Compare the strength of the forces of attraction between

molecules of different elements. (For example, at room

temperature, chlorine is a gas and iodine is a solid.)

Practice: Page 345; 38

Section(s) 14.3

Page(s) 340

2. C4.3d Observe the following table and explain why at room temperature different compounds can exist in

different phases using the terms strongest and weakest intermolecular forces, molecules, and relative freedom

of movement.


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C4.3f Identify the elements necessary for hydrogen bonding (N, O, F).

Practice: Page 341; 9, Page 345; 36 (a-d), 37

Section(s) 14.3

Page(s) 341, 349-353

3. C4.3f In the diagram below, identify the elements necessary for hydrogen bonding between the DNA bases.

C4.3g Given the structural formula of a compound, indicate all the

intermolecular forces present (dispersion, dipolar, hydrogen

bonding).

Section(s) 14.10,14.13

Page(s) 335,340

4. C4.3g Look at the following structures and identify, which one has all of the following intermolecular

forces present (dispersion, dipolar, hydrogen bonding).

C5.4c Explain why both the melting point and boiling points for water

are significantly higher than other small molecules of comparable

mass (e.g., ammonia and methane). Practice: page 341; 8,

page 345; 57

Section(s) 14.13

Page(s) 341

5. C5.4c Examine the following chart of small molecules:


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In comparison to the other molecules of similar mass, explain why water has a much higher melting and

boiling point using the terms strong, weak, intermolecular forces of attraction, and hydrogen bonds.

Claim:

Evidence:

Reasoning:

C5.4d Explain why freezing is an exothermic change of state. Section(s) 8.7,9.9

Page(s) 185, 186, 213, 214

& Teacher Supplement

6. C5.4d Explain why the freezing of water is an exothermic change.

A. As water freezes heat must be added.

B. As water freezes heat is given off.

C. As the water freezes there is no change in the amount of heat content.

D. As water freezes limited changes of state take place resulting in heat exchange.

C5.4e Compare the melting point of covalent compounds based on the

strength of IMFs (intermolecular forces).

Section(s) 14.13,14.14

Page(s) 340,343

7. C5.4e Which of these lists shows substances in order of increasing melting point?

A. H2O, SO2, O2, Ca(OH)2

B. Ca(OH)2, SO2, H2O, O2

C. H2O, Ca(OH)2, SO2, O2

D. O2, SO2, H2O, Ca(OH)2


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Vocabulary

Boiling point elevation

Calorie

Change of state

Chemical bond

Concentration

Convection current

Convection heating

Crystalline solid

Electrostatic attractions

Enthalpy

Entropy

Equilibrium

Exothermic reaction

Freezing point depression

Hess’s Law

Ionic motion

Joules

Kinetic energy

Mass to energy conversion

Potential energy

Release of energy

Solute

Specific heat

Transforming matter and/or energy

C2.1c Compare qualitatively the energy changes associated with melting various types of

solids in terms of the types of forces between the particles in the solid.

Practice: page 365; 6, page 366; 14, 23

Section(s) 9.8,15.5,16.7

Page(s) 210,353,383

1. C2.1c Arrange the following solid substances in order of increasing strength of attractive forces.


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C3.1c Calculate the ΔH for a chemical reaction using simple coffee cup calorimetry.

Practice: page 365; 12

Section(s) 8.7, 15.3

Page(s) 185, 351

2. C3.1c To find the heat change for an organic reaction, a chemist must use hexane as the solvent in her

calorimetry experiment. If the reaction causes the temperature of 0.0700 kilograms of hexane to decrease by

3.57 0

C, what is the heat change for the reaction? (Hexane has a specific heat capacity of 2.26 kJ/kg • 0 C)

3. How many joules are needed to warm 25.5 grams of water from 14.00 0C to 22.50

0C?

4. Calculate the number of joules released when 75.0 grams of water are cooled from 100.00 0C to 27.50

0C.

5. Calculate the heat, in joules, needed to warm 225 g of water from 88.00 0C to its boiling point 100.00

0C.

6. The specific heat of gold is 0.128 J/g0 C. How much heat would be needed to warm 250.0 grams of gold

from 25.00 0C to 100.00

0C?

7. The specific heat of zinc is 0.386 J/g 0 C. How many joules would be released when 454 grams of zinc at

96.00 0C where cooled to 28.00

0C?

C3.1d Priority

Calculate the amount of heat produced for a given mass of reactant from a

balanced chemical equation.

Practice: page 187; 10, 11, page 193; 30, 31, 32

Section(s) 8.7

Page(s) 185

8. C3.1d Calculate the energy required to produce 8.00 mol of Cl O on the basis of the following balanced

equation.

2Cl (g) + 7O (g) + 130 kcal 2Cl O (g)


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C5.4A Priority

Compare the energy required to raise the temperature of one gram of aluminum

and one gram of water the same number of degrees. Section(s) 7.5

Page(s) 154

9. C5.4A Study the following table of specific heat capacities.

On a sunny day at 25 0 C if the sun is shining on a 100 g of water and a 100 g of aluminum, which one will be

the hottest and why?

10. Worksheet: Introduction to Specific Heat Capacities

C5.4B Priority

Measure, plot, and interpret the graph of the temperature versus time of an ice-

water mixture, under slow heating, through melting and boiling. Section(s) 9.9,9.10

Page(s) 213,214

11. C5.4B Examine the following graph:

Explain what is happening along each point on this curve. At what temperature does this substance freeze and

boil? Under normal conditions at room temperature what is the state of matter of this substance?


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Vocabulary

Acid rain

Acid/base reaction

Acidic

Alkaline

Basic

Bronsted-Lowry

Carboxyl group

Hydrogen ion

Hydronium ion

Hydroxide

Ion

Kw,

Neutral

Neutralize

pH

C5.7A Recognize formulas for common inorganic acids, carboxylic acids, and bases

formed from families I and II.

Practice: page 106; 14 & 14, teacher supplement

Section(s) 5.13, 5.11

Page(s) 105-106,

102-103

1. C5.7A Of the following which one is a carboxylic acid?


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C5.7B Predict products of an acid-based neutralization. Practice: page 441-442; 8,9 & 10

Section(s) 18.5

Page(s) 440-442

2. C5.7B When an acid reacts with a base, what compounds are formed?

C5.7C Describe tests that can be used to distinguish an acid from a base.

Section(s) 18.3, 18.6,

18.8

Page(s) 435-437,

445-446, 448-451

3. C5.7C Use Table 6-1of acid-base indicators to answer the following question.

If you exhale carbon dioxide (CO2) into a solution of bromothymol blue, the solution turns from blue

to yellow. When CO2 dissolves in water, what will the pH of the solution be?

C5.7D Classify various solutions as acidic or basic, given their pH. Section(s) 18.3

Page(s) 435-437

4. C5.7D Observe the following table.

Which substance would be classified as an acid? Which one is a base? Which one is neutral?


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C5.7E Explain why lakes with limestone or calcium carbonate experience less

adverse effects from acid rain than lakes with granite beds.

Section(s) 18.5

Teacher Supplement

5. C5.7E Why do lakes with calcium carbonate beds, such as those in many Michigan lakes, experience fewer

adverse effects from acid rain than lakes with granite beds?

A. Calcium carbonate reacts with the acid rain to neutralize it.

B. Calcium carbonate in water does not let the acid rain dissolve in the lake.

C. Granite reacts with acid rain to produce more acid.

D. Granite-based lakes make the acid rain more acidic than it was when it fell as rain.

C5.7g Calculate the pH from the hydronium ion or hydroxide ion concentration.

Practice: page 436 -438; 2-6

Section(s) 18.3, 18.4

Page(s) 435-440

6. C5.7g What is the pH of a solution whose hydronium ion concentration is 5.03 101

M?

C5.7h Explain why sulfur oxides and nitrogen oxides contribute to acid rain. Teacher Supplement

7. C5.7h Environmental scientists have found that the two major chemical compounds that cause acid rain are

sulfur dioxide (SO2) and nitrogen oxides (NOx); these are released into the air when we burn fossil fuels

mainly from power plants and automobiles. Explain why sulfur oxides and nitrogen oxides contribute to acid

rain.


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Vocabulary

Anode

Cathode

Electrochemical Cell

Equilibrium

Keq

Le Chatelier

Oxidation

Oxidation-reduction reactions

Reduction

C5.3a

Describe equilibrium shifts in a chemical system caused by changing

conditions (Le Chatelier’s Principle).

1. C5.3a Suppose that more H2 is added to the following reaction.

2 H2 (g) + O2 (g) ↔ 2 H2O (g)

Which of the following will not change?

A. the entropy of the system (S)

B. the concentration of O2

C. the concentration of H2O

D. the value of the equilibrium constant (K)

C5.3b Predict shifts in a chemical system caused by changing conditions (Le

Chatelier’s Principle).

2. C5.3b What is the effect of adding more water to the following equilibrium reaction?

CO + H O H CO


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C5.6c Explain oxidation occurring when two different metals are in contact.

3. C5.6c The diagram shows strengths of reducing agents, which are elements that reduce a substance in an

oxidation-reduction reaction.

Many underground pipes are made out of iron. An engineer wants to attach a piece of material to an iron pipe

to prevent rust formation (oxidation). Based on the diagram, which of the following materials should the

engineer use?

A. gold C. copper

B. silver D. zinc

C5.6e Identify the reactions occurring at the anode and cathode in an

electrochemical cell.

4. C5.6e Where does oxidation take place in an electrochemical cell.


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Vocabulary

Activation energy

Disorder

Endothermic reaction

Enthalpy

Entropy

Exothermic reaction

Gibb’s Free Energy

Hess’s Law

Reaction rate

Release of energy

Spontaneous

C3.1a Calculate the ΔH for a given reaction using Hess’s Law.

1. C3.1a Use the information below to calculate H for the following reaction.

2NO (g) N O (g)

N (g) + 2O (g) 2NO (g) H = 67.7 kJ

N (g) + 2O (g) N O (g) H = 9.7 kJ


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C3.1b Draw enthalpy diagrams for exothermic and endothermic reactions.

2. C3.1b From the following identify the enthalpy diagram for an exothermic reaction.

C3.2a Describe the energy changes in photosynthesis and in the combustion of

sugar in terms of bond breaking and bond making.

3. C3.2a During the combustion of glucose, more energy is released when bonds form in the products than is

absorbed to break bonds in the reactants. Which term best describes the energy change for the combustion of

glucose?

C3.4B Explain why chemical reactions will either release or absorb energy.

4. C3.4B When a chemical cold pack is activated, it becomes cool to the touch. What is happening in terms of

energy?


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C3.4d Draw enthalpy diagrams for reactants and products in endothermic and

exothermic reactions.

5. C3.4d Draw and label an exothermic reaction including the energy of activation.

Draw and label an endothermic reaction including the energy of activation.

Constructed Response

C5.6e C5.6c Underground iron pipes in contact with moist soil are likely to corrode. This corrosion can be

prevented by applying the principles of electrochemistry. Connecting an iron pipe to a magnesium block with

a wire creates an electrochemical cell. The magnesium block acts as the anode and the iron pipe acts as the

cathode. A diagram of this system is shown below along with an activity series.

State the direction of the flow of electrons between the electrodes in this cell.

Explain, in terms of reactivity, why magnesium is preferred over zinc to protect underground iron

pipes. Your response should include both magnesium and zinc.

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