Things to Know to Pass the Chemistry Regents
1. Protons: charge +1, mass 1 amu, in nucleus, = atomic number
*1 amu = 1/12 a carbon-12 atom
2. Neutrons: charge 0, mass 1 amu, in nucleus, = mass number - atomic number
3. Electrons: charge -1, mass 0 (1/1836) amu, in e- cloud surrounding nucleus, = atomic #
4. Nucleons: subatomic particles in nucleus, protons and neutrons, mass of atom
5. Orbitals: most probable location of electrons in e- cloud, modern (wave-mechanical) model
6. Mass number: protons + neutrons, C-14 has a mass of 14 (6p + 8n = 14)
7. Atomic number: equals # of protons, identifies element/atom
*all atoms with 6p are carbon, all atoms of carbon have 6p
8. Number of neutrons = mass number (p+n) - atomic number(p)
9. Atoms are neutral because # of protons(+) = # of electrons(-)
10. Isotopes are atoms with same # of protons and electrons, but different number of neutrons
*therefore also have different mass numbers
11. Metal atoms lose e- to form (+) ions smaller than atom
12. Nonmetals gain e- to form (-) ions larger than atom
13. Dalton’s model solid, uniform sphere
14. Thompson’s plum-pudding model, (+) sphere with e- throughout (cathode ray tube)
15. Rutherford’s gold foil showed atoms small (+) nucleus & mostly empty space with e-
*few deflections = small (+) nucleus, most through = mostly empty space
16. Bohr’s model e- in orbits like planets around sun (orbit does NOT equal orbital)
17. Modern, wave-mechanical model e- in orbitals (most probable location)
18. Electrons absorb energy and move to higher energy levels, electrons emit light when fall back
*light = spectral lines, energy, bright line spectrum, quanta, photons
19. Elements = pure substances that cannot be chemically decomposed (H, He, Li, Be...)
20. Binary compounds made up of two different types of atoms (HCl)
21. Diatomic molecules = elements that form 2-atom molecules (H2, O2, F2....)
*diatomic gases = H2, N2, O2, F2, Cl2; liquid = Br2; solid = I2
22. Zeros only holding the decimal do not count (0.024 = 2), trailing zeros count (0.0240 = 3)
23. When multiply/divide answer cannot have more sig figs than number with least sig figs
*2.44 x 2.4 = 5.9 (not 5.856)
When add/subtract answer cannot have greater place value than least place value
*2.44 + 2.4 = 4.8 (not 4.84)
24. Solutions are homogeneous (uniform throughout) mixtures, NaCl(aq)
*composition variable
25. Heterogeneous mixtures are not uniform, raisin bran, soil, sandy water
*composition variable
26. Solute is the thing that is dissolved, NaCl in NaCl(aq)
27. Solvent is the thing that does dissolving, water in NaCl(aq)
28. Isotopes are written mass # on top, atomic # on bottom, 6C14
29. Atomic mass = weighted average of naturally occurring isotopes
*75% Cl = 35 amu, 25% Cl = 37 amu
*(.75)(35) + (.25)(37) = 35.5 amu
*(abundance)(mass) + (abundance)(mass) + ...
30. Electron configuration shows how many and where e- are located
*2-8-1 (2 e- in 1st level, 8 in 2nd, 1 in 3rd)
31. 1 mole = gram formula mass = 6 x 1023 things = 22.4 L if gas
32. Empirical formula = reduced (molecular formula C5H10 = empirical formula CH2)
33. Electron-dot diagrams (Lewis structures) show valence e-
34. Element symbol in e-dot diagram = kernel = everything about atom except valence e-
35. Polyatomic ions (Table E) are covalently bonded together, have an overall charge
*metal + (-) poly has ionic and covalent bonds
36. Coefficients are written in front of a formula to indicate how many (3HCl = 3 HCls)
37. Chemical formulas are written so the (+) part cancels the (-) part
38. Binary compounds normally (+) part 1st, (-) part 2nd (NaCl, Na+, Cl-)
39. Compounds with polyatomic ions use poly’s name (NaNO3 = sodium nitrate)
40. Roman numerals tell oxidation # (iron (II) oxide, Fe +2, O -2, FeO)
41. Physical changes do NOT form new substances (H2O(l) –> H2O(s))
42. Chemical changes DO form new substances (2H2O(g) –> 2H2(g) + O2(g))
43. Reactants on left products on right side of arrow
44. Temperature = average kinetic energy
45. Exothermic processes release energy, on right with produces, delta H (-)
*Endothermic processes absorb energy, on left with reactants, delta H (+)
*See Table I
46. Only coefficients can be changed when balancing equations
*HCl –> H2 + Cl2 = 2HCl –> H2 + Cl2
47. Synthesis reactions = A + B –> AB
48. Decomposition reactions = AB –> A + B
49. Single replacement = A + BC –> B + A or D + EF –> F + ED
50. Double replacement = AB + CD –> CB + AD
51. Mass, energy, and charge are always conserved
52. Gram formula mass (gram molecular mass, molar mass, mass of one mole )
*Na = 23, Cl = 35, NaCl = 23 + 35 = 58 grams per mole
53. Percent composition = part/whole x 100 (Table T)
*%Na in NaCl = Na/NaCl x 100 = 23/58 x 100 = 40%
54. Avogadro’s law: = volumes of gases at same temp and pressure have = number of molecules
55. Find molecular formula given molar mass and empirical formula
*molar mass = 92 g, empirical = NO2
*NO2 = 14 + (16)(2) = 46, 92/46 = 2, 2(NO2) = N2O4
56. KMT: Assumptions are 1. motion, 2. collisions, 3. no volume, 4. no attraction
*Deviations are *do have volume, *do have attraction
57. Solid particles vibrate in place, close together, geometric arrangement, crystal lattice
*definite shape and volume
*low energy and entropy (randomness)
58. Liquid particles still close together, but free to move around
*shape of container, definite volume
*moderate energy and entropy
59. Gas particles far apart and zooming around
*shape and volume of container
*high energy and entropy
60. Gases can be compressed
61. Gas particles create pressure (tire pressure)
62. Heating/cooling curve: step like, flat during phase change, diagonal during warming or cooling
*flat = phase change, KE =, PE up if warming, down if cooling, 2 phases present
*diagonal = warming or cooling, PE =, KE up if warming, down if cooling
*mp/fp and bp/cp
63. Sublimation = s –> g (dry ice, solid iodine, mothballs)
*deposition = g –> s
64. Vapor pressure created by liquids turning into gases (Table H)
*as temp goes up, vapor pressure goes up
*high vapor pressure = readily evaporates = weak forces of attraction
*low vapor pressure = does not readily evaporate = strong forces of attraction
*normal boiling point at 101.3 kPA (1 atm)
65. STP = standard temperature and pressure, Table A
66. Kelvin = degrees C + 273 (Table T)
67. Heat always goes from warmer to cooler
68. q = mHf (melting), q = mHv(vaporizing), q = mC(change in temperature) (raising/lowering)
*Tables T and B
69. Combined gas law on Table T
*If given STP, given temp and pressure (Table A)
70. Pressure and volume indirect, P up, V down (PVC pipe)
71. Temperature and pressure direct, T up, P up
72. Temperature and volume direct, T up, V up
73. Gases most ideal at high temp and low pressure (have more energy and free to spread out)
*ideal is summer vacation
74. He and H most ideal because small and weak forces of attraction
75. Mixtures can be separated by physical means
*distillation, different boiling points (evaporation too)
*filtration different solubilities
*chromatography different attraction for separating medium
76. Periodic law states properties of elements are periodic functions of their atomic number
77. Elements arranged according to atomic number on modern periodic table
78. Periods are horizontal rows (across), 1 - 7, if in same period same # of energy levels occupied
79. Groups vertical (up and down), 1-18, in same group same # of valence electrons
*similar chemical behavior because same # of valence e-
80. Alkali metals group 1, Alkaline Earth metals group 2, halogens group 17, noble gases 18
81. Metals on left of periodic table, metalloids along zig-zag line, nonmetals right of table
*H on left but nonmetal
82. Metals malleable, ductile, good conductors, shiny
*Nonmetals brittle, poor conductors, dull
*Metalloids properties of both metals & nonmetals
83. Noble gases, group 18, nonreactive because stable e- configurations
*noble gases exist as monatomic molecules (He, Ne, Ar...)
84. Ionization energy decreases going down group & increases across left to right (Table S)
*ionization energy is amount of energy needed to remove e- (metals low, non high)
85. Atomic radii increase going down group because additional energy levels
86. Atomic radii decrease going across left to right because energy level pulled in more by more positive nucleus
87. Electronegativity is measure of attraction an atom has for electrons (Table S)
*Metals low, nonmetals high
88. Electronegativity decreases going down a group & increases going left to right
89. Use Table S to look up and compare properties of elements
*Density, mp/fp, bp/cp, radii, ionization energy, electronegativity
90. Breaking bonds energy absorbed
*bonds are attractions therefore need energy to pull apart
91. Making bonds energy released
*more energy released, more stable bond
92. Last number of e- configuration = # of valence e-
93. Draw e-dot diagrams of elements
*symbol, # of dots = # of valence e-
94. # of valence e- determines reactivity
95. Metallic bonds between metal atoms
*sea of valence e-
96. Atoms react to obtain a noble gas electron configuration (octet)
*electrons involved in reactions
97. Ionic bonding between metal + nonmetal & (metal + (-) poly, (+) poly + non, (+) poly = (-) poly)
*transfer of valence e-
*electronegativity difference 1.7
97. Covalent bonding between nonmetal atoms
*share valence e-
*molecular substances
98. Nonpolar covalent bonds = sharing
*electronegativity difference 0 - .3
99. Polar covalent bonds unequal sharing
*electronegativity difference .4 - 1.6
100. Coordinate covalent bond one atom supplies shared pair (moocher sharing)
*NH4+ and H3O+
101. Nonpolar molecules are covalent compounds with symmetrical shapes
*H2, CH4, CO2
102. Polar molecules are covalent compounds with asymmetrical shapes
*have a (+) end and a (-) end
*NH3, HCl, H2O
103. Covalent compounds have attractive forces between molecules
*dipole-dipole, between (-) end of one molecule and (+) end of another
*hydrogen bonding = super dipole-dipole between molecules containing H plus F, N, or O (small, highly electronegative atoms) because the (+) end is very (+) and the (-) end is very (-), HF, NH3, H2O, accounts for H2Os high boiling point
*weak-intermolecular forces occur between nonpolar molecules
*molecule-ion between (+) ion and (-) end of water and (-) ion and (+) end of water, NaCl(aq), Na+ attracted to O end of water, Cl- attracted to (+) end of water
104. Metallic solids malleable, ductile, soft to hard, low to high mp, good conductors
105. Ionic solids hard, high mp, poor conductors solid, conduct liquid, gas, (aq)
106. Covalent compounds exist as either molecular solids or network solids
*molecular solids soft, low mp, poor conductors (ice)
*network solids, hard, high mp, poor conductors (graphite, diamond, sand, SiC, asbestos)
107. Likes dissolve like (similar polarities)
*polar dissolve polar
*nonpolar dissolve nonpolar
*nonpolar and polar do not mix (oil and water, water polar, oil non)
108. Solids more soluble as temperature increases
109. Gases less soluble as temperature increases
110. Gases most soluble low temp high pressure (soda pop)
111. Table G
*saturated on line (equilibrium)
*unsaturated below line (dilute)
*supersaturated above line (difference piled up on bottom)
112. Table F
*soluble or insoluble
113. Molarity measure concentration (Table T)
*M = moles/liters of solution
114. Percent by mass or percent composition (Table T)
*% comp = part/whole x 100
115. Parts per million (Table T)
*ppm = grams solute/grams solution x 1,000,000
116. Solutes (dissolved particles) raise the bp and lower the fp
*more particles dissolved, great the effect
*ionic dissociates, more effect than covalent
117. Normal boiling point means at standard pressure (101.3 kPa, 1atm), Table H
118. Particles must collide effectively for a reaction to take place
120. Covalent compounds react slower than ionic aqueous
*ionic (aq) already apart and ready to form new bonds
121. Increased concentration of reactants increases rate of reaction because more collisions
*increased surface area increase rate because more particles exposed
122. Increased temp increased rate of reaction because more collisions and stronger collisions
*particles have more energy
123. Catalysts increase rate of reaction by lowering activation energy (alternate pathway)
124. Increased pressure when gases present increases rate of reaction
125. Interpret potential energy diagrams
*PE of reactants
*PE of products
*Activation energy (and show effect of a catalyst)
*Activated complex
*Heat of reaction
*Read forward and reverse
126. Equilibrium (solution, phase, chemical)
*RATES of forward and reverse processes EQUAL
*CONCENTRATIONS CONSTANT (constant volume, constant mass)
127. Equilibrium shifted by stress (Le Chatelier’s)
*Shifts away from increased temp, concentration
*Shifts toward decreases temp, concentration
*Shifts toward smaller volume if gases present with increased pressure
*Catalyst does NOT shift (just get there faster)
128. Enthalpy = heat ( lower energy favored in nature)
129. Entropy = randomness (greater entropy favored in nature)
130. Reaction will occur spontaneously if toward lower energy and greater entropy
131. Oxidation numbers can be assigned to atoms and ions and elements in compounds
132. Redox (oxidation-reduction) reactions involve a change in oxidation number
*something loses e- something gains e-
*can’t have one without other
*# e- lost = # e- gained
133. Oil Rig
*oxidation is loss (Na –> Na+ + e-)
*reduction is gain (Cl + e- --> Cl-)
134. An Ox
*anode is site of oxidation
135. Red Cat
*cathode is site of reduction
136. Electrons are transferred from what lost to what gained
*anode to cathode
137. Table J
*on top oxidized, on bottom reduced
138. Voltaic cells (typical battery)
*redox
*anode oxidation, cathode reduction
*electrons on wire
*ions on salt bridge
*two cells
*uses chemical energy to produce electrical energy
139. Electrolytic cell (and electroplating apparatus) (when recharging phone)
*redox
*anode oxidation, cathode reduction
*electrons on wire
*one cell
*external power source (battery)
*uses electrical energy to produce chemical energy
*electrolytic ant (anode +) Thank you Margaret Ball!
140. Acids, bases, and salts aqueous are electrolytes because form ions in water (dissociate)
141. Acids give off H+ (hydrogen ion) or H3O+ (hydronium ion) (Tables K and M)
*donate H+
*neutralize bases
*sour, pH less than 7, litmus red, phth colorless
*Metals react with to give off H+ (Table J, ones above H)
141. Bases give off OH- (hydroxide ion) in water (Tables L and M)
*accept H+
*neutralize acids
*bitter, pH greater than 7, litmus blue, phth pink
142. Acid + base –> salt + water
*H+ + OH- –> HOH (H2O)
*neutralization
HCl + NaOH –> NaCl + HOH
143. Titration (Table T)
*used to find concentration of unknown acid or base
MaVa = MbVb
144. pH scale
*7 neutral (water, sugar water, salt water, alcohols, etc.) (H+ = OH-)
*<7 acidic (H+ > OH-), farther from neutral = more acidic
*>7 basic (OH- . H+), farther from neutral = more basic
*each move a 10x change in H+ concentration (1 is 10x stronger than 2, 1 is 100x stronger than 3)
145. All organic compounds contain C, carbon
*and (usually) H, hydrogen
146. Carbon ALWAYS forms 4 covalent bonds
147. Organic info on Tables P, Q, and R
148. Hydrocarbons only contain H and C
*alkanes all single bonds, saturated
*alkenes one double bond, unsaturated
*alkynes one triple bond, unsaturated
149. Isomers same molecular formula only
*different arrangement, structure
*different names
*different properties (physical and chemical)
*must have 4 or more carbons to have isomers
*more carbon atoms, more possible isomers
150. Functional groups specific for classes of compounds (Table R)
*alcohols -OH, acids -COOH, esters - COOC, etc...
151. Organic reactions
1. Combustion, burning
*complete produces water and carbon dioxide
*incomplete produces water and carbon monoxide (insufficient oxygen)
2. Substitution
*alkanes, one atom replaces another
3. Addition
*alkenes and alkynes, double or triple opens up to add on atoms
4. Esterification
*alcohol + organic acid –> ester + water
5. Saponification
*produces soap
6. Fermentation
*sugar broken down to form alcohol and carbon dioxide
7. Polymerization
*monomers + monomers –> polymer
*small units combine to form long molecules
*natural = proteins, cellulose, starches
*synthetic = nylon, plastics, rayon
152. Tables N and O radioactivity (also use periodic table)
153. Half-life unchanging (Table N and T)
154. Mass and charge balance in nuclear equations
155. Nuclear notations, mass on top, charge on bottom
156. Natural radioactivity (alpha, beta, positron, gamma decay)
*spontaneous
*due to unstable neutron : proton ratio (C-14 radioactive, C-12 not)
*atomic number 84 and above ALL radioactive
157. Artificial radioactivity
*bombard a nucleus to make it unstable and therefore radioactive
158. Fission (nuclear reactors)
*splitting large nuclei like uranium
159. Fusion (the sun)
*fusing H + H to form He
160. Both VERY exothermic (think mushroom cloud)
161. Uses of radioactivity
*dating, C-12 to C-14 once living, U to Pb rocks
*tracers, C-14 through biological processes like photosynthesis
*I-131 thyroid problems detected and treated
*Co-60 and Tc-99 cancer diagnosis and or treatment
*Irradiated foods
162. Risks of radioactivity
*biological exposure
*nuclear accidents
*radioactive wastes