1. Chemistry - Final Review Guide

4. REMEMBER YOUR CALCULATOR on the day of the exam
REMEMBER YOUR CALCULATOR on the day of the exam. Also bring a sharpened pencils and eraser.

5. Review topics and readings for the exam:

6. Chemical Reactions

7. Types of reactions

8. Composition

9. Decomposition

10. Single- replacement

11. Double- replacement

12. Neutralization

13. Balancing reactions

14. Solubility Rules:

15. Ions of metals from group IA and ammonium ions are soluble

16. Chloride ions are soluble with the exception of silver, mercury and lead

17. Sulfate are soluble with the exception of metals from group IIA and lead

18. Chlorate are soluble

19. All others are insoluble (ppt)

20. Activity series rules:

21. 1. Metals from the top of the series would displace any metal below it

22. 2. Metals from group IA and IIA (Ca, Sr and Ba) would react with water, Mg reacts with steam

23. 3. Metals above hydrogen would displace it from hydrochloric and sulfuric acid

25. Chemical composition

27. 1. Mole Calculations:

28. Mol = mass/Mw

30. Mol = (# particles) / (6.02 x 10^23)

32. Mol = V in L/ 22.4 at STP

33. Mol = M x V in L

34. 2 Calculations for molecular mass

35. Example: CaCO3 Mw = 40 + 12 + 3(16) =100g/mol

37. Percent composition

38. Example: K2SO4

39. Find % K

40. %K = (2(Aw K)/ Mw) x100

41. Empirical formula and Molecular formula

43. Stoichiometry

44. Balancing Chemical equations

45. Interpreting the coefficient in the balance equation

46. Mole-mole calculations

47. Mole-mass calculations

48. Mass-mass calculations

49. Limiting reactants

50. *Limiting reactant: the amount of reactant that will limit the amount of product

51. *Theoretical Yield: the amount of product calculated from the balance equation

52. *Actual (expected) Yield: the amount of product produce in the experiment

53. * % yield= actual yield/theoretical yield x100

55. IV. Solids, liquids and gases

57. 1. KMT - five postulates

58. a) Gas particles are far apart and can be compressed, because there are a lot of empty spaces

59. b) Gas particles are constantly moving and exerting pressure on the side of the container

60. c) Vgas = Vcontainer (particles occupy only small portion of the volume)

61. d) Gas particles don’t stick to each other when they collide

62. e) KE = 3/2RT

64. 2. Distribution change / temperature graph

65. 3. Physical properties of gases - Expansion, Fluidity, Low density, Diffusion and effusion

66. 4. Gas pressure (Force/ area)

67. 5. Measuring the gas pressure- barometer and manometer

68. 6 Physical equation for water - heating curve

69. 7. Properties of liquids according the KMT: shape, volume, distance between the particles and forces between the molecules (dipol-dipol, London-dispersion and hydrogen bonding)

70. 8. Vapor pressure

71. a) Difference between boiling and evaporation

72. b) factors affecting ability of a liquid to evaporate (increases of surface area, temperature and decrease of pressure)

73. c) Factors affecting the rate of evaporation of two different liquids (particle attraction and mass of particles)

74. V. Gas laws:

75. 1. Boyles’s Law - T = const

77. P1V1 = P2V2

78. 2. Charles’s Law - p = const

80. V1/T1 = V2T2

81. 3. Gay-Lussac’s Law - V = const

82. P1/T1 = P2T2

83. 4. Combined Law: P1V1/T1 = P2V2/T2

84. 5. Avogadro’s Law: V1/n1 = V2/n2

85. 6. Ideal gas Law: PV = nRT (R = 0.08205 atmL/molK)

86. 7. Dalton’s Law: P total = Pgas + Pwater vapor for collecting gas by water displacement

87. 8. Calculations for molecular weight: Mw = DRT/P

88. 9. Real gases - at high p and lower temperature

89. a) Molecules occupy volume

90. b) There are attractive forces between the molecules

91. c) Van der Waal’s equation:

92. (P + n2a /V2) ( V-nb) = nRT

93. 10. Phase diagram: triple point, determining boiling and freezing point

95. VI. Solutions

97. 1. Solution - a homogeneous mixture of two or more substances in a single physical state.

98. a) Solute - substance that is dissolved in a solution (the less quantity).

99. b) Solvent- substance that does the dissolving in a solution (the more quantity).

101. 2. Types of solution - gaseous, liquids and solids (review the table from the notes).

102. * Explore 15-1: Solutions of Gases, liquids and solids - in this activity you investigated how substances of different physical states can be mixed to produce a solution.

103. * The nature of solutions

105. 3. Solubility. Factors affecting the rate of solution.

106. * Lab: Preparing solutions (heating, stirring and grinding)

107. * Factors effecting the process of dissolving of liquids (“Like dissolve like” rule), solids and gases (temperature and pressure)

108. * Henry’s Law - The bends (15-3 activity worksheet)

110. 4. Saturated, unsaturated and supersaturated solutions

111. a) Saturated- contains as much solute as can possibly be dissolved under existing conditions ( T and p).

112. b) Supersaturated - a solution that contains more solute particles then are needed to form a saturated solution.

113. c) Unsaturated - a solution that has less then the maximum amount of solute that can be dissolved.

115. 5. Concentration units:

116. a) Mass % = mass solute/mass solution x 100

117. b) Molarity (M) = mol of solute / volume of solution in L

118. c) Molality (m) = mol of solute / kg of solvent

119. d) Mol fraction = mola/moltotal

120. * See practice problems 15-2

121. * 15-2 Cooking divinity worksheet

123. 6. Colligative properties

124. a) Freezing point depression

125. Tf = kf m

126. b) Boiling point elevation

127. Tb = kb m

128. * See: Practice problems 15-4

129. * 15-4 Explore - Colder then ice water - in this activity you have discovered what happens to the freezing point of water when a substance is dissolved

131. VII. Equilibrium

133. 1. The concept of equilibrium

134. In chemistry - most chemical reactions are reversible processes
In chemistry - most chemical reactions are reversible processes. (Activity with the beads and Equal rate - !6.1 explore - investigate a reversible reaction.)

135. a. When the rate of forward reaction is = to the rate of the reverse reaction - chemical equilibrium is established.

136. b. Under the same condition, at equilibrium the concentration of both reactants and products remain constant.

137. c. At the same temperature the equilibrium constant is constant

138. d. Equilibrium may be approached from different starting points.

139. e. At other temperature the value of the equilibrium constant differs.

141. 2. The law of equilibrium (16.2 Review and reinforcement - worksheet)

142. a) Equilibrium constant Keq and the reaction quotient Qeq.

143. b) Collision theory -the molecules must collide in order to react
b) Collision theory -the molecules must collide in order to react. In successful collision the existing bonds in a molecule are broken and new bonds are formed.

144. c) Factors affecting the rate of a chemical reaction

145. * Collision frequency (depends of concentration of the molecules and

146. temperature).

147. * Collision energy

148. * Orientation of the molecules

149. 3. Reaction profile- shows the energy of reactants and products during reaction

150. * Transition state - is the highest point on the reaction profile where reactant and products have the same potential energy.

151. * Activation energy - the energy require for the reaction to achieve the transition state

152. a) Energy profile for exothermic reaction

153. b) Energy profile for endothermic reaction

154. 4. Types of chemical equilibrium

155. a) Homogeneous

156. b) Heterogeneous

157. 5. Le Chatelier’s principle - when a dynamic equilibrium is upset by disturbance, the equilibrium will shift in a direction to minimize the effect of disturbance

158. (See 16-3 Review and reinforcement).

159. a. Effect of concentration

160. b. Effect of temperature

161. c. Effect of pressure

162. d. Effect of catalyst

163. e. Shifting the equilibrium - the Haber process

164. 6. Solubility equilibrium - Ksp ( see 17.1 Problems)

165. Ksp of a substance is the product of the molar concentration of its ions in a saturated solution, each raised to the power that is the coefficient of that ion in the chemical equation.

167. VIII. Acids and bases

169. 1. Properties of acid and bases

170. * Acid -is any substance that produces hydrogen ions in water

171. - changes blue litmus paper to red

172. - Sour taste

173. - pH  7

174. * Base- is any substance that produces OH- ions in water solution

175. - changes red litmus paper to blue

176. - Bitter taste

177. - pH  7

178. - feel slippery or soapy to the touch

179. Both acids and bases can undergo a neutralization reaction.

180. Neutralization reaction is a reaction between acid and base where the products are salt and water.

181. 2. Arrhenius definition

182. * Acid - is a substance that ionizes in water to produce H+

183. * Base - is a substance that dissociate in water to release OH- ions

184. 3. Bronsted - Lowry definition

185. * Acid - is a substance that donates hydrogen ions to any other substance (proton donor).

186. * Base - is any substance that accept a hydrogen ion (proton acceptor).

187. 4. Acids and bases strength

188. * Strong acid will ionize completely. (HCl, H2SO4, HNO3, HClO4)

190. * Weak acid will ionize slightly. (CH3COOH, HNO2, HCN, H3PO4, H2CO3)

191. HA(aq) + H2O(l)  H3O+ (aq) + A-(aq)

193. Ka = H3O+ A- HA

195. * Strong base will dissociate completely ( MeIAOH ).

197. * Weak base will provide relatively few ions in solution

198. B (aq) + H2O (l)  BH+ (aq) + OH- (aq)

200. Kb = BH+OH-B

202. 5. Acid- base titration

203. a) Principles of titration’s

204. * Titration - is the controlled addition and measurement of the amount of a solution of known concentration that is required to react completely with a measured amount of a solution of unknown concentration.

205. * Standard solution - a solution that contains a precisely known concentration of a solute

206. * Titration curve - is used to represent pH data

207. - Strong acid with the strong base (upward graph - end point at pH =7).

208. - Weak acid with a strong base (upward graph - end point at pH = 8.9)

209. - Weak base with strong acid (downward graph - end point at pH = 5.8)

210. * Equivalent point - is the point at which exactly enough standard solution is added to neutralize the unknown solution.

211. * End point - the point at which the indicator changes color.

212. b) Indicators - are weak acids or bases dyes whose colors are sensitive to pH or hydronium ion concentration.

213. 6. Ionization of water - ionization constant

214. Kw = 1 x 10-14

215. if H3O+OH-- neutral

216. if H3O+OH- - acidic

217. if H3O+OH- - basic

219. pH concept

220. pH =- log H-

221. pOH = - log OH-

222. 8. Buffers - are solutions that resist changes in pH when an acid or a base is added

223. * It is composed of an aqueous solution of a weak acid and one of its salts