Chemistry Regents Review Notes

Unit 10 Solutions:

• Solubility is dependent on temperature for solids (upward curve on Table G).
• Table G: Use when given a temperature and inquiring about a salt dissolved in water.
  • On the line: Saturated.
  • Below the line: Unsaturated.
  • Above the line: Supersaturated.
• "Likes dissolve likes": Polar substances mix with polar substances.
• Table F: Insoluble compounds will not fully dissolve in water (solid phase); soluble compounds will dissolve (aqueous phase).
• The more ions present, the greater the change in boiling point (BP) or freezing point (FP).
• Adding a solute (like salt) to water decreases the freezing point. More salt leads to a lower freezing point.
• Adding a solute (like salt) to water increases the boiling point. More salt leads to a higher boiling point.
• Concentration units: m/L or PPM.
• $PPM = (grams \ of \ solute / grams \ of \ solution) \times 1,000,000$ (Formula on Table T).
• $Molarity = moles/L$ (Formula on Table T).

Unit 11 Kinetics:

• For a chemical reaction to occur, sufficient energy and proper orientation are needed for new bonds to form.
• Chemical reactions need effective collisions.
• Increasing temperature increases the reaction rate (more collisions).
• Increasing pressure increases the reaction rate (more collisions).
• Lower concentration slows the reaction rate (fewer collisions).
• Greater surface area = higher rate of reaction.
• Potential energy diagrams show the energy of a reaction; adding a catalyst does not change the energy of reactants and products.
• Activation energy: Energy from reactants to the top of the curve.
• Reverse activation energy: Energy from the products to the top of the curve.
• Exothermic reaction: Potential energy diagram starts high and ends low.
• $Heat \ of \ reaction = potential \ energy \ of \ products - potential \ energy \ of \ reactants$
• Heat of reaction $(\Delta H)$ is in the middle of a potential energy diagram.
• Table I: Shows heat of reaction (energy absorbed $+\Delta H$, energy released $\Delta H$).
• A catalyst speeds up a reaction by providing a different pathway that lowers the activation energy.
• Entropy (disorder): Gases have the highest entropy, then liquids/aqueous solutions, and solids have the lowest.
• Nature undergoes changes towards higher entropy and lower energy.

Safety, Scientific Method and Graphing:

1. Safety precaution for long hair: Tie it back.
2. If you spill a liquid on your arm: Rinse it off and then tell the teacher.
3. After a lab: Dispose of chemicals properly.
4. Scientific method: Keep all variables the same except the one you are testing.
5. Graphing: Use an even scale of numbers and circle final points.

Unit 1 Atomic:

6. Atoms contain subatomic particles (protons, neutrons, electrons) and are divisible.
7. Protons have a +1 charge.
8. Neutrons have a 0 charge.
9. Mass of neutron = 1 amu (atomic mass unit) = mass of proton.
10. Electrons have a charge of -1 and a mass of 0.
11. Charge of electron is -1, charge of proton is +1 (same magnitude, opposite charge).
12. Table O: Shows symbols, mass, and charges of particles (electrons are represented as beta).
13. Protons and neutrons are located in the nucleus of an atom.
14. Charge of an atom's nucleus = (+) number of protons.
15. Atoms have a positively charged nucleus and negatively charged electrons located in "clouds" (orbitals) around the nucleus.
16. Mass \ number = #protons + #neutrons
17. $Atomic \ number = number \ of \ protons$
    • All atoms of the same element have the same atomic number.
18. $Number \ of \ neutrons = Mass \ number - atomic \ number$
19. Isotopes are atoms with the same number of protons but different numbers of neutrons (different mass number).
20. Isotope notation: the top number is the mass number, and the bottom number is the atomic number.
21. Other notations: C-14 or Carbon-14 (number represents mass number).
22. Atomic mass is the weighted average of all naturally occurring isotopes for that element.
23. $Average \ Atomic \ mass = (isotope1 \ mass) ($
24. Abundance: The whole number the atomic mass is closest to on the Periodic Table indicates the most abundant isotope.
25. Rutherford's gold foil experiment showed that an atom is mostly empty space with a small, dense, positively charged nucleus.
26. Thomson and Bohr's models showed electrons present in an atom.
27. Wave-mechanical model (electron cloud model) shows that an orbital (cloud) is the most probable location of electrons.
28. Neutral Atom: An atom has the same number of protons and electrons (total charge of 0).
29. Total \ (Net) \ Charge \ of \ an \ atom = # \ protons - # \ electrons
30. Ion: A charged element (lost or gained electrons); electron configuration will change if it is an ion (possible charges are found on PT).
31. Electron configuration: Shows the location of electrons in their shells (e.g., 2-8-2).
32. The first shell has less energy than the second shell.
33. The first shell can hold a max of 2 electrons, and the second shell can hold a max of 8 electrons.
34. Valence electrons: Electrons in the outermost shell (last number in electron configuration).
35. Lewis dot diagram for a single atom shows the valence electrons (electrons represented by dots, drawn in pairs).
36. When an excited electron (at a higher energy level) moves to the ground state (lower energy level), a specific amount of energy is emitted (sometimes as light/bright line spectrum).
37. Excited electron configuration is not the same as the configuration on the reference table.
38. Energy emitted from an excited electron can be used to determine the identity of the element.
39. When viewing a bright line spectrum, elements must line up exactly to be part of the mixture in the spectrum.

Unit 2 Nuclear:

40. Table O: A positron and beta particles have the same mass (0) & opposite charge (Beta negative, positron positive).
41. A beta particle has less mass and greater penetrating power than an alpha particle (Gamma radiation has the greatest penetrating power).
42. All nuclear reactions are transmutations (fission, fusion, decays).
43. Any element after Po is naturally unstable and will spontaneously decay.
44. Stable Isotopes are not on Table N (do not spontaneously emit particles).
45. Table N: Shows decay modes and half-lives (alpha, beta, positron decay).
46. Natural Transmutations show spontaneous radioactive decay: 1 reactant → 2 products (elements must change).
47. Nuclear decays release the decay particle.
48. Completing nuclear equations: The sum of mass numbers and atomic numbers must be equal on both sides of the arrow.
49. Fusion: Light nuclei combine to form a heavy nucleus and a lot of energy (energy is sometimes in the form of a neutron).
50. Fusion produces more energy than fission.
51. Nuclear reactions (fission or fusion) release more energy than a chemical reaction (redox, substitution, neutralization).
52. In nuclear reactions, mass is converted into energy.
53. Half-life: The length of time it takes for ½ mass of a sample to decay.
54. 1 half-life = ½ sample remains, 2 half-lives = 1/4 sample remains, 3 half-lives = 1/8 sample.
55. Half-life questions: Use the table or timeline method.
56. Radioisotopes are used for dating geological formations (C-14).
57. I-131 is used to diagnose thyroid disorders.
58. Radioisotopes can be used to detect diseases.
59. Radioisotopes can treat cancer but can also cause mutations in healthy cells (Co-60).

Unit 3 Matter:

60. Substance = compound or element.
61. Elements cannot be broken down by chemical means (it is on the Periodic table/Table T).
62. Compounds can be broken down by chemical means.
63. Same compound = same chemical property; different compound = different chemical properties.
64. 7 diatomics (two of the same atom bonded together): BrINCIHOF (Br2, I2, N2, Cl2, H2, O2, F2).
65. Melting Point (MP) and Boiling Point (BP) of elements are on Table T.
66. *Liquid at a specific temperature: MP < Specific Temperature < BP
67. A mixture can vary in the proportion of its components (e.g., saltwater).
68. Homogeneous mixtures (solutions): even distribution of particles (aq-dissolved in water).
69. Heterogeneous mixtures = not even throughout; contains a substance that will not be soluble in water.
70. When substances are mixed, they retain their properties.
71. Mixtures containing substances with different densities and particle sizes can be separated by physical means.
72. Mixtures can be separated by chromatography, distillation, and filtration.
73. Distillation separates liquids with different boiling points (water and alcohol).
74. Chromatography separates particles by solubility and polarity.
75. Evaporation separates a salt dissolved in water.
76. Chemical property: How substances react.
77. Chemical change: Results in the formation of a different substance (e.g., burning).
78. Physical change: Does not form new compounds, commonly phase changes (change in the distance between molecules).
79. Solids = atoms close together; liquids = atoms in the middle; gases = atoms far apart.
80. Solids have a definite shape and definite volume.
81. Deposition = gas to solid phase change.
82. Sublimation: Solid → Gas phase change (ex: CO2).
83. In a phase change diagram, the flat parts represent the phase changes (Potential Energy [PE] changes, and Kinetic Energy [KE] remains the same).
84. In a phase change diagram, the sloped lines represent heating or cooling (PE remains the same, and KE changes).
85. $Density = mass/volume$ (g/L or g/cm³).
86. Higher density sinks to the bottom of a tank.
87. Density never changes for each element (Found on Table S for elements).

Math:

88. Sig figs: Atlantic Pacific rule:
    • Decimal absent: Count from the first non-zero number on the Atlantic side (right) and count all numbers to the left.
    • Decimal present: Count from the first non-zero number on the Pacific side (left) and count all numbers to the right.
89. When multiplying/dividing: The answer should have the lowest number of sig figs.

Unit 4 Energy:

90. $\% \ error = (measured \ value - accepted \ value)/ accepted \ value * 100$

91. $Kelvin = °C + 273$

92. $1kJ = 1000J$

93. Forms of energy: chemical, thermal, electromagnetic, electrical, nuclear, mechanical.

94. Thermal energy (heat) is measured in joules (J) = random motion of atoms and molecules.

95. Average kinetic energy = temperature.

96. When two substances have the same temperature, the substance with the greater mass has more thermal energy.

97. Heat of vaporization is the amount of heat required to vaporize a substance (Table B for water constants) = 2260J or $2.26 \times 10^3 J$.

98. Heat of fusion (heat it takes to melt a substance) is less than heat of vaporization because it requires less heat to melt a substance than boil a substance.

99. Heat flows from hot to cold.

100. $Q= mC\Delta T$

• (q is heat, m is mass, C is specific heat capacity [found on Table B for water], $\Delta T$ is change in temperature). All info on Table T.

90. Exothermic = energy exits (is released).
91. Endothermic = energy absorbed (heat is shown on the left side of the equation) (examples of endothermic phase changes: s→l, l→g, s→g).

Unit 5 Gas Laws:

103. Pressure only affects gases.
104. Pressure is measured in pascals.
105. STP (Standard Temperature and Pressure) on Table A (273 K and 1 atm or 101.3 kPa and 0°C).
106. Pressure and temperature have a direct relationship (as pressure increases, temperature increases).
107. $\frac{P<em>1V</em>1}{T<em>1} = \frac{P</em>2V<em>2}{T</em>2}$

• If something is constant, it can be crossed out of the formula.

103. Same volume = same number of molecules.
104. Conditions for an ideal gas: P↓↑ (Pressure low, ideal gas behavior), T↑ (high temperature).
105. Ideal gases move in random, constant, straight-line motion.
106. Ideal gases are separated by great distances compared to their size.
107. Ideal gases have no attractive forces.
108. Collisions of gas may result in a transfer of energy.
109. Table H: The dotted line is the normal boiling point of the substance.
110. Gases have weaker intermolecular forces (IMF) than solids.

Unit 6 Periodic Table:

116. Mendeleev organized his periodic table by atomic mass.
117. Modern Periodic Table: Elements are arranged in order of atomic number.
118. Periods are horizontal rows on the Periodic Table.
119. Groups are vertical rows on the Periodic Table.
120. Liquids on the periodic table = bromine and mercury.
121. Metals (left of the staircase) are good conductors of heat and malleable.
122. Metals are malleable because of the nature of the bonds between the atoms.
123. Metals have fewer valence electrons than nonmetals.
124. Commonly, metals react with nonmetals.
125. Nonmetals are to the right of the staircase.
126. Metalloids are on the staircase.
127. Transition metals form colored solutions in aqueous ion form.
128. Group 17 (halogen group) forms halide ions in solution.
129. Noble gases (group 18) are stable because of their stable electron configuration (8 valence electrons); they are not reactive.
130. Elements will gain or lose electrons to be like noble gases.
131. MELPS- Metals Electrons Lost form Positive Smaller ions.
132. Some groups will form the same oxidation numbers when gaining or losing electrons.
133. Same group/family = similar chemical properties because they have the same number of valence electrons.
134. Attraction for electrons in a chemical bond = electronegativity (found on Table T).
135. First ionization energy and electronegativity increase left to right (across a period) and decrease top to bottom (down a group).
136. Atomic radius decreases left to right (across a period) and increases top to bottom (down a group).
137. Atomic radius increases down a group because more energy shells are added.

Unit 7 Naming, Formulas, and Equations:

138. Ionic formulas: metal comes first, nonmetal second. Crisscross charges to get the formula (e.g., Calcium phosphate: $Ca^{2+}$ $PO<em>4^{3-}$ = $Ca</em>3(PO<em>4)</em>2$).
139. Naming ionics with Table E: Say the metal and then the ion from Table E.
140. Roman numerals represent the charge of the metal.
141. Polyatomic ion names are on Table E.
142. Balancing: The number of each atom needs to be the same on both sides of the equation (use tallies). Use coefficients to balance.
143. Mass, charge, and energy are conserved in a chemical reaction.
144. Types of chemical reactions: decomposition, single replacement, double replacement, synthesis.
145. Synthesis: Two or more reactants combine to form one product.
146. Decomposition: One compound is broken down into two or more.
147. Single replacement: One element switches partners.
148. Double Replacement:

Unit 8 Bonding:

149. BARF = bond BROKEN energy ABSORBED (endothermic), energy RELEASED bond FORMED (exothermic).
150. Ionic compounds (metal and nonmetals or Table E ions), molecular compounds are all nonmetals with no Table E ions.
151. Ionic compounds with a Table E ion have covalent and ionic bonds.
152. Ionic compounds: Hard, brittle solids with a high melting point, poor conductors as solids, and good conductors in aq (higher concentration = better conductor).
153. Metallic bonding is metals only!
154. Ionic bonds = bond between metal and nonmetal or ion from Table E.
155. Ionic bonds transfer electrons from the valence shell of one atom to the valence shell of another atom.
156. Lewis dot diagrams have brackets for ionics. Molecular compounds: Make sure all elements have eight electrons (except H), and the total number of electrons in the diagram = total number of valence electrons.
157. Covalent bond = nonmetals (molecular compound).
158. A nonpolar covalent bond is a bond between two of the same elements.
159. Multiple covalent bond: double (four electrons shared/2 pairs), triple (six electrons shared/3 pairs).
160. Diatomic Oxygen molecule has a double bond, and diatomic nitrogen has a triple bond.
161. Electronegativity difference = polarity of bond (higher the difference, the more polar the bond).
162. Molecule Polarity: SNAP = symmetrical nonpolar, asymmetrical polar.
163. Stronger Intermolecular forces (IMF) = High boiling point.
164. Example of IMF: hydrogen bonding (between H and F, O, or N).
165. Water molecules (Oxygen is partially negative, and hydrogen is partially positive).
166. Positive and negative charges attract, and like charges repel.

Unit 9 Stoichiometry:

167. Atomic mass on the Periodic Table is in grams/1 mole.
168. GFM (gram formula mass, formula mass, or molar mass) is found by getting the mass of each element (multiplied by their subscript) and adding all masses together.
169. % Composition by mass = (mass of part / mass of whole) (Table T).
170. % Composition can be calculated given a periodic table and chemical formula.
171. Mole = mass/GFM (Table T).
172. Mole: mole ratio = coefficient: coefficient ratio.
173. Empirical formula: simplified formula; molecular is a normal, non-simplified formula.
174. Conservation of mass: mass of products = mass of reactants.

Equilibrium:

175. Forward reaction → Reverse reaction ←
176. Equilibrium ←→
177. At equilibrium, the concentration of reactants and products remains constant because reaction rates of the forward and reverse reaction are equal.
178. Rate of forward reaction = rate of reverse reaction.
179. A closed system is needed to maintain equilibrium.
180. Shift to the right makes more products; shift to the left makes more reactants.
181. UP and AWAY, DOWN, and TOWARDS.
182. Decrease pressure, shift to the side with the most number of moles.
183. Increase temperature, shift away from heat.
184. Adding a catalyst causes no shift.
185. Solution equilibrium: The rate of dissolving equals the rate of crystallization.
186. Dynamic phase equilibrium = two phase changes going on in a sealed flask.

Unit 12 Acids and Bases:

187. Electrolytes: acids, bases, salts.
188. Electrolytes conduct in aqueous solutions, not in the solid phase.
189. Arrhenius theory describes the behavior of acids and bases.
190. Arrhenius theory: acids yield $H^+$ ions and bases have $OH^-$.
191. Acids and bases are on Tables K/L (know $CH_3COOH$ is an acid; all other $CH$ compounds are not electrolytes).
192. Alternate theory: BAAD: BASES ACCEPT $H^+$, ACIDS DONATE $H^+$.
193. The most acidic compounds have the lowest pH.
194. The lower the pH (more acid) = more hydronium ions.
195. Moles of$OH^-$ and $H^+$ are equal in a neutral solution.
196. Neutralization: Acid + base → salt + water.
197. Metals above $H_2$ on Table J will react spontaneously with an acid.
198. $M<em>aV</em>a= M<em>bV</em>b$

• where M is molarity or concentration.

187. Titration determines the concentration of unknown acids or bases.
188. Indicators: The range on the reference table (Table M) is when the color is in-between two primary colors.

Unit 13 Redox:

201. Oxidation number: Single atom = 0; look on the reference table or 1-30 algebra for atoms with multiple charges (can also use reverse crisscross if all else fails).
202. In redox reactions, electrons are transferred (charge changes).
203. Oxidation: electrons are lost (electrons on the right); Reduction: electrons are gained (electrons on the left).

• LEO the lion says GER (Lose Electrons Oxidation, Gain Electrons Reduction)

201. Half-reactions show charges and electrons lost.
202. A balanced redox reaction has the same number of electrons lost and gained (atoms and charges balanced).
203. Lower on Table J is less reactive.
204. More active metals are higher on Table J.
205. Higher on Table J are more likely to be oxidized for the metal side; lower are more easily reduced.
206. If the oxidized element is higher on Table J, then the reaction will be spontaneous.
207. AN OX RED CAT (anode oxidation, reduction cathode)
208. Electrons flow from anode to cathode.
209. The mass of the anode will decrease because solid metal is going into aqueous solution.
210. Voltaic cell: chemical energy is converted to electrical energy spontaneously.
211. Electrons flow through a wire for voltaic cells.
212. Ions move through the salt bridge for a voltaic cell.
213. Electrolytic cell: electrical energy to chemical energy.
214. A battery (power source) is needed for an electrolytic cell to provide electrical energy (electrolysis or electroplating).
215. The cathode in an electrolytic cell is what is electroplated (key or spoon).

Unit 14 Organic:

219. Organic compounds must contain at least one carbon and one hydrogen.
220. Hydrocarbons contain only carbon and hydrogen.
221. Saturated hydrocarbons have all single bonds (2 electrons shared).
222. Unsaturated compounds have double or triple bonds between carbons.
223. Chemical formulas for organic compounds: empirical, structural, molecular.
224. The structural formula shows bonds.
225. Carbon (atomic number 6) can form chains, rings, and networks.
226. Table P tells the number of carbons (remember carbon always makes 4 bonds!).
227. Homologous series on Table Q (alkanes, alkenes, alkynes).
228. Naming alkenes = (use Table Q): start from the side with the double bond and count the number of carbons; put a number to represent the location of the double bond. Count carbons from the side closest to the double bond or functional group.
229. Isomers have the same number of C's and H's but are in a different structural arrangement (same molecular formula).
230. Different structures = different physical/chemical properties.
231. Functional groups are on Table R.
232. Different Functional groups have different chemical properties.
233. Drawing the structure of an alcohol includes putting an OH.
234. Halides contain group 17 elements.
235. An organic acid structure has C with a double-bonded O and OH group.
236. Amines contain N.
237. Types of organic reactions: esterification, addition, saponification, polymerization, fermentation, etc.
238. Addition reaction: A binary compound (two atoms bonded together) gets added to a double/triple-bonded compound to reduce the number of bonds.
239. Saponification = an organic reaction used to make soap.
240. Polymerization = adding the same compound together to make a chain.
241. Fermentation: sugar + enzyme → ethanol (alcohol) + carbon dioxide.