SEMESTER TWO REVIEW PACKET

UNIT NINE: STATES OF MATTER

• Intermolecular Forces – attractive forces between molecules

• Surface Tension – tendency of liquids to minimize surface area

• Viscosity – resistance of a liquid to flow

• Volatile – easily evaporates

• Vapor Pressure – pressure exerted by a liquid’s vapor molecules

• Melting – phase change from solid to liquid

• Freezing - phase change from liquid to solid

• Boiling - phase change from liquid to gas

• Condensing - phase change from gas to liquid

• Sublimation - phase change from solid to gas

• Deposition - phase change from gas to solid

2. Gases would have weak intermolecular forces whereas solids and liquids would have strong intermolecular forces.

3. The rate of vaporization increases with:- increase in surface area- increase in temperature- decrease in strength of IMF

4. Vapor pressure increases with:- decreases in strength of IMF- increase in temperature

5. To make a liquid boil:- increase the temperature- decrease the atmospheric pressure

6. To make a gas condense:- decrease the temperature- increase the atmospheric pressure

7. How much energy would be needed to vaporize 43.0 g of water?

8. How much energy would be needed to melt 28.0 g of ice?

9. What quantity of heat would need to be added to 123.4 g of water to change its temperature from 52°C to 86°C?

10. Dispersion is a temporary uneven distribution of electrons that increases with molar mass, symmetry and surface area.

11. Dipole – Dipole is a permanent distribution of electrons

12. Hydrogen Bonding occurs when hydrogen is bonded to nitrogen, oxygen, or fluorine.

13. Molecular Solids have neutral corners, low melting points, are nonmetal with nonmetal or a metalloid with a metal or nonmetal.

14. Covalent Solids have neutral corners, high melting points, are metalloid with metalloid, or a metalloid with a metal or nonmetal.

15. Metallic Solids have positive ions and are metals with metals.

16. Ionic Solids have positive and negative ions at the corners and are metals with nonmetals.

17. Why is water unique?Water has a low molar mass yet is a liquid at room temperature, easily dissolves polar and ionic compounds, expands upon freezing, and its ice is less dense than water.

18. What is the percentage of water, by mass, in FeSO4•7H2O?

UNIT TEN: GAS LAWS

19. Standard pressure is:1 atm, 101.325 kPa, 760 mmHg, and 14.7 psi

20. Temperature must be in units of Kelvin. K = 273 + °C

21. What are the relationships between pressure and volume, pressure and temperature, volume and temperature, and volume and the number of moles?- P increases V decreases- V increases and T increases- V increases and n increases- P increases and T increases

22. A gas mixture contains helium, oxygen, and an unknown gas. Calculate the partial pressure (in atm) of the unknown gas when the partial pressure of helium is 503 mmHg, oxygen is 214 mmHg, and the total pressure is 1.13 atm.

23. If a balloon occupies 1.49 L at room temperature (25.0°C), what temperature in °C would you need to increase the volume to 3.07L?

24. If there is a pressure of 536 mmHg in a container and it is heated from 20.0°C to 100.0°C, what is the new pressure?

25. A hand pump with a moveable piston has an applied pressure of 3.08 atm and a volume of 3.25L. What is the volume if the applied pressure is decreased to 1.18 atm?

26. If you transfer a gas at room temperature (25.0°C ) from a 5.89 L cylinder to a 2.41 L container and the pressure increases from 0.653 atm to 1.36 atm, what is the final temperature in °C?

27. If you blow up a balloon with helium gas and the volume increases from 56.3 mL to 1006 mL and there are initially 0.163 moles, how many moles need to be added?

28. If 28.4 L of Cl2 react with an unlimited supply of H2 at STP, how many moles of HCl would be produced?

29. If helium effuses 3.60 times faster than an unknown gas sample, what is the molar mass of the unknown gas?

UNIT ELEVEN: SOLUTIONS

30. Terms

•Solution – a homogeneous mixture of two or more substances

•Solute – the minority component

•Solvent – the majority component

•Suspension – a mixture from which the particles settle out upon standing

•Colloid – a permanent mixture whose particles are smaller than a suspension and larger than a solution

•Tyndall Effect – a way to differentiate solutions from colloids and suspensions based on the scattering of light

•Electrolyte – aqueous solutions containing a solute that dissociates into ions and conducts electricity

•Solubility – the amount of component that will dissolve in a certain amount of liquid

•Saturated – holds the maximum amount of solute under the solution conditions

• Unsaturated – holds less than the maximum amount of solute under the solution conditions

• Supersaturated – holds more than the normal maximum amount of solute

• Concentration – the amount of solute in a solution

• Mass Percent – the number of grams of solute per 100 grams of solution

• Molarity – the number of moles of solute per liter of solution

• Molality – the number of moles of solute per kilogram of solvent

• Colligative Properties – a property that depends on the number of solute particles and not the type of solute particle (how much you have not what you have)

• Freezing Point Depression – difference in temperature between the freezing point of a solution and the freezing point of the pure solvent

• Boiling Point Depression – difference in temperature between the boiling point of a solution and the boiling point of the pure solvent

31. Which of the following will scatter light?- colloid and suspension

32. What would happen to the following if more solute is added to the solution:

- Saturated: nothing will happen- Unsaturated: solute will dissolve- Supersaturated: extra solute will disrupt the solution and

precipitate out

33. Solubility depends on:- identity of solute and solvent- temperature- pressure (only gases)

34. Factors that increase the rate of solution are:- decreasing the particle size- stirring- increasing temperature (except for gases)

35. Solubility Curve

a. Pb(NO3)2 at 30°C

- 66 g

b. 12 g of KClO3

- 35°C

c. 70 g of CaCl2 at 16°C

- saturated

36. A soft drink contains 85.2 grams of sucrose (C12H22O11) in a 741 mL solution. What is the percent mass if the density of the solution is 1.00 g/mL?

37. What is the molarity of a solution that contains 98.5 grams of MgBr2 dissolved to produce a 1.00 L solution?

38. What volume of 12 M HCl do you need to make 500.0 mL of a 3.0 M solution?

39. Determine the volume in milliliters of 0.225 M KOH solution required to neutralize 185 mL of 0.125 M HCl. The neutralization reaction is:

KOH(aq) + HCl(aq) → H 2O(l) + KCl(aq)

40. Calculate the molality of a solution containing 225 grams of glucose (C6H12O6) dissolved in 1.25 L of water. (Assume the density of 1.00 g/mL for water).

41. What is the freezing point depression and the boiling point elevation of a 2.43 m solution of Al2(SO4)3 in water?

UNIT TWELVE: ACIDS AND BASES

42. Determine whether the following are acids (A), bases(B), or both (X).

a. ____ pH>7 b. ____ Tastes bitterc. ____ Tastes sour d. ____ pH<7e. ____ Corrosive to skin f. ____ Blue litmus redg. ____ Feels slippery h. ____ Electrical Conductori. ____ Red litmus blue

43. Name the following formulas:a. CsOH

- cesium hydroxideb. H2S

- hydrosulfuric acidc. HNO3

- nitric acidd. HNO2

- nitrous acid

A

B

X

A

AA

B

B

B

44. Write the formulas of the following names:a. hydrobromic acid

- HBrb. iron (II) hydroxide

- Fe(OH)2

c. phosphoric acid- H3PO4

d. chlorous acid- HClO2

45. Explain how water is amphoteric.Water can act as an acid and a base. It is able to accept and donate protons.

46. List the six strong acids: HCl, HBr, HI, HClO4, H2SO4, HNO3

47. List the eight strong bases: LiOH, NaOH, KOH, RbOH, CsOH, Ca(OH)2, Ba(OH)2, Sr(OH)2

48. The ion product constant for water is Kw = 1.0 x 10-14

49. Calculate the pH, pOH, [H+], and [OH-]. Identify the type of solution as acidic, basic, or neutral.

50. What are the reactants and products in an acid base reaction?

- reactants: acid and base

- products: water and salt (ionic compound)

51. What is the purpose of titrations?

A titration is an experimental technique used to determine the concentration of an unknown acid/base by comparing it to an acid/base solution of known concentration

52. H2SO4 (aq) + Ba(OH)2 (aq) → BaSO4(aq) + 2 H2O(l)

53. 2 Al(OH)3(aq) + 3 H2SO4(aq) → Al2(SO4)3 (aq) + 6 H2O(l)

54. What is the molarity of a solution of NaOH if 34.2 mL of the solution is used to neutralize 25.3 mL of 0.50 M HCl?

55. How many milliliters of 1.25 M HCl must be added to 50.0 mL of 2.50 M LiOH to make a neutral solution?

56. What is a buffer? What do they consist of?

A buffer is a solution that resists a change in pH. Buffers contain a weak acid and its conjugate base.

57. For the following questions label each part of the following chemical reactions with acid (A), base (B), conjugate acid (CA) and conjugate base (CB).

UNIT THIRTEEN: EQUILIBRIUM

58. Terms• Activation Energy - amount of energy that must be absorbed by

reactants before a reaction can occur

• Catalyst - a substance that increases the rate of a chemical reaction but is not consumed by the reaction

• Equilibrium - when forward and reverse reactions are occurring under the same conditions and at the same rate

• Enthalpy - the energy difference between the reactants and products; another word for energy

• Entropy - randomness, disorder

• Equilibrium constant - product of concentrations of products divided by the product of the concentrations of reactants

• Free Energy - the net balance or difference between energy (enthalpy) and entropy

• Le Châtelier’s Principle - when a chemical system at equilibrium is disturbed, the system shifts in a direction that minimizes the disturbance

• Reversible Reaction - a reaction where the products can re-form reactants

• Solubility Product Constant - is an equilibrium expression for a chemical equation that represents the dissolving of an ionic compound

• Spontaneous - happens on its own, does not have to be forces

• Static - at rest; not changing

• Steady State - constant change but unlike equilibrium there is no reverse reaction

• Thermodynamic - the study of energy relationships

59. Describe each of the following situations as either “equilibrium” or “steady state”.

a. Lake Wissota Dam and the water behind it. The level in the lake is constant. Steady State

b. The liquid in a sealed container and the vapor above it at constant temperature Equilibrium

c. The solid solute remaining in a sealed container of solutionEquilibrium

60. For the reaction below, Keq = 0.39. If the equilibrium concentration of HC2H3O2 is 2.3 M, what are the equilibrium concentrations for the two ions below? HC2H3O2 (aq) + H2O (l) H⇌ 3O+

(aq) + C2H3O2 –

(aq)

61. Calculate the value of Keq for the reaction below, using the given equilibrium concentrations: CH4 (g) + 3 Br2 (l) CHBr⇌ 3 (aq) + 3 HBr (aq)

62. What causes a system at equilibrium to shift? Stress

63. Methyl alcohol (“methanol”) is made according to the following net equation. Predict the effect on the equilibrium system (“shift left”, “shift right”, or “no effect”).

CO (g) + 2 H2 (g) CH⇌ 3OH (g) + heat

a. CO is added to the container: RIGHTb. The volume is decreased: RIGHTc. A catalyst is added: NO EFFECT d. Temperature is increased: LEFTe. Pressure is increased: RIGHTf. Hydrogen gas is removed: LEFT

64. Write the complete equation for Fe(OH)2 and write the Ksp expression.

Fe(OH)2 → Fe2+ + 2OH- Ksp = [Fe2+][OH-]2

UNIT FOURTEEN: REDOX

68. Assign an oxidation number to the specified atom in each of the following examples.

71. Will the following redox reactions be spontaneous?

a. Pb(s) + Sn2+(aq) → Pb2+

(aq) + Sn(s) NO

b. Pb(s) + Ag+(aq) → Pb2+

(aq) + Ag(s) YES

72. How can you pre-determine if a metal will dissolve in an acid?

If the metal is below H2 on the activity series it will not dissolve in the acid. If the metal is above H2 on the activity series it will dissolve in an acid.

73. Calculate the cell voltage and give the cell notation for electrochemical cells made from the following electrodes.

Ecell = 1.05v Ni | Ni2+ || 2Ag+ | 2Ag

UNIT FIFTEEN: ORGANIC

80. Draw isomers of hexane.

UNIT SIXTEEN: BIOCHEMISTRY

93. Dogs may eat homework but do they digest it? Explain why.

Paper is made of cellulose which is indigestible so the dog cannot digest the paper.

94. The insulin protein contains 51 amino acids. How many DNA base pairs are required to code for all the amino acids in insulin?

153

95. When wet hair is put into curlers and allowed to dry, the hair tends to retain the shape of the curler. Why?

When hair dries the hydrogen bonds stay in place from when wet, therefore keeping the shape of the curler.

96. Why is shape so important in the functions of proteins?

The shape determines the function. Heat can denature and change shape therefore change function.

97. Name three lipids of the cell and describe their function.

Answer will vary

98. What are the different structures of proteins? How would you identify these structures?

Primary – amino acid sequenceSecondary – short range repeating pattern along protein chainTertiary – large scale bends and foldsQuaternary – arrangement of chains in protein

99. How are alpha-helix proteins different from beta-pleated sheet proteins?

Alpha helix is a coil held together through hydrogen bonds of NH and CO. The side chains extend outward.

For beta-pleated sheets the peptide backbones of the chain interacts through hydrogen bonds. The side chains extend above and below.

100. The anti-HIV medicine AZT introduces a fake thymine molecule into the cell. When a virus attempts to use the fake nucleotide in replicating its DNA, the fake nucleotide doesn’t work. Below are diagrams of the fake thymine on the left and the real thymine on the right. How does AZT keep the virus from replicating when the virus tries to incorporate the fake thymine into its DNA?

Azidothymidine and thymine have the same shape but slightly different structures. Thymine has an OH group and azidothymidine has an amine group (-N=N=NH). The different structure has a different function.

Final Schedule

Friday:

1st Hour 8:35 – 10:30

2nd Hour 10:40 – 12:35

Lunch 12:35 – 1:30

3rd Hour 1:35 – 3:35

Monday:

4th Hour 8:35 – 10:30

5th Hour 10:40 – 12:35

Lunch 12:35 – 1:30

6th Hour 1:35 – 3:35