1. Atomic MassAM=∑ (fraction of isotopes)∗(massof isotopes )
2. Avogadro’s Number 1mol=6.022∗1023 particles
3. Relations1cm3=1mL1m3=1dL
4. Chemical SymbolXZ
A
A = mass number = number of proton + neutronsZ = atomic number = number of protonsX = chemical symbol
5. Periodic Table and Types of Elements
6. Properties of MetalsGood conductors of heat and electricity, malleable, can be drawn into wires (ductility), shiny, tend to lose electrons when undergo chemical changes.
7. Properties of Metalloids
Several classified as semiconductors because of their intermediate (and highly temperature - dependent) electrical conductivity
8. Properties of NonmetalsSome are solids at room temperature, other are liquid or gases, all poor conductors of heat and electricity, all tend to gain electrons when undergo chemical changes
9. Properties of Alkali MetalsAll are reactive metals
10. Properties of Alkaline Earth MetalsAlso all fairly reactive, although not quite as reactive as the alkali metals.
11. Polyatomic Ions
12. Classifying Matter
13. Classification of Elements and Compounds
14. Nomenclature
15. Mass % of an Element in a Compound
Mass%X=( Mass of X∈1mol of CompoundMolar Massof Compound )∗100%
16. Solubility Rules
17. % Yield
%Yield=( Actual YieldTheoritical Yield )∗100%
18. Types of Chemical Reactions1. Combination(Synthesis) Reaction
A+B→C2. Decomposition Reaction
A∆→B+C+…
3. Single Replacement ReactionI. If A and C are metals
AB+C→CB+AII. If E and F are nonmetals
DE+F→DF+E4. Double Replacement Reaction
AB+CD→CB+ADa. Precipitationb. Acid-Base
5. Hydrocarbon Combustion ReactionC xH y+O2→CO2+H 2OC xH yO z+O2→CO2+H2O
19. Molarity (M)
M= moles SoluteLiters Solution
=molL
20. DilutionsM 1V 1=M 2V 2
21. Oxidation–Reduction(Redox) Reactions
→ Reactions in which electrons transfer from one reactant to the other. Any reaction in which there is a change in the oxidation states of atoms in going from reactants to products
22. Oxidation → Loss of electrons(LEO)/Increase in oxidation number
23. Reduction → Gain of electrons(GER)/Decrease in oxidation number
24. Oxidizing Agent → Oxidizes another substance while itself gets reduced
25. Reducing Agent → Reduces another substance while itself gets oxidized
26. Types of Reaction that ARE Redox1. Combination2. Decomposition3. Single Replacement4. Combustion
27. Types of Reactions That Are NOT Redox1. Double Replacement
a. Precipitation b. Acid-Base
28. Oxidation State Rules
29. Ionic Bonding(cation)Metal - (anion)nonmetal
30. Covalent Bondingnonmetal - nonmetalnonmetal - metalloid
31. Boyle’s Law (Gases)
V ∝ 1P
32. Charles’s LawV ∝T
33. Avogadro’s LawV ∝n
34. Combined Gas LawP1V 1
n1T 1=
P2V 2
n2T 2
35. Universal Gas Constant
R=0.08206 L∗atmmol∗K
36. Ideal Gas LawPV=nRT
37. Standard Temperature and Pressure(STP)P = 1.00 atmT = 0.00 °C + 273.15 = 273.15 K
38. Density of Gas
D= P∗MR∗T
where M is molar mass.
39. Partial Pressure of a
Pa=na∗R∗T
V
40. Dalton’s Law of Partial Pressures
Ptotal=Pa+Pb…=ntotal∗R∗T
Vwhere ntotal=na+nb+…
41. Mole Fraction
X a=na
ntotal=
Pa
Ptotal
42. Final Partial Pressure EquationPa=Xa∗P total
43. Kinetic Molecular Theory1. The size of a particle is negligibly small.2. The average kinetic energy of a particle is proportional to the temperature in kelvins.3. The collision of one particle with another(or with the walls of its container) is completely elastic
44. Equation for Change in Internal Energy ΔE=E final−E intial=E products−E reactants = q(heat) + w(work)
45. Amount of Energy Lost by a SystemΔE system=−ΔEsurroundings
46. Energy Flow Out of System(System ---> Surroundings)ΔE system<0 (negative )→ΔEsurroundings>0( positive)
47. Energy Flow Into System(Surroundings ---> System)ΔE system>0 (positive )→ΔEsurroundings<0(negative)
48. Equation for Heat with Heat Capacity q=C∗ΔTwhere C is heat capacity, the amount of heat required to raise the temperature by 1°C or 1 K
49. Equation for Heat with Specific Heat Capacityq=m∗C s∗ΔTwhere C s is the specific heat capacity, the amount of heat required to raise the temperature of 1
gram of the substance by 1°C or 1 K
50. Equation for Heat with Molar Heat Capacityq=n∗CS∗ΔTwhere n is the molar heat capacity, the amount of heat required to raise the temperature of 1
mole of the substance by 1°C or 1 K
51. Equation for System and Surroundings with Massmsystem∗C s , system∗ΔT system=−msurroundings∗C s , surroundings∗ΔT surroundings
52. Equation for System and Surroundings with Molesnsystem∗C s ,system∗ΔT system=−nsurroundings∗C s , surroundings∗ΔT surroundings
53. Constant Volume CalorimetryMeasures ΔEreaction
qcal=C cal∗ΔT qreaction=−qcalorimetry
ΔEreaction=qreaction
n
54. Constant Pressure CalorimetryMeasuresΔH reaction
qsolution=msolution∗C s , solution∗ΔT solution qreaction=−qsolution
ΔH reaction=qreaction
n
55. Endothermic Reaction – absorbs heat, feels cool to the touch. ΔH >0
56. Exothermic Reaction – releases heat, feels warm to the touch. ΔH <0
57. Hess’s Law1. If a chemical equation is multiplied by some factor, then ΔH reaction is also multiplied by the same factor2. If a chemical equation is reversed, then ΔH reactionchanges sign3. If a chemical equation can be expressed as the sum of a series of steps, then ΔH reaction for the overall equation is the sum of the heats of reaction for each step
58. Calculating the Standard Enthalpy Change for a ReactionΔH °reaction=Σ(np∗ΔH °f (ΔH ° products))−Σ (np∗ΔH °i(ΔH °reactants))
59. Standard Statea. Gas, Solid, Liquid – 1 atm @ 25° Cb. Solution – 1 atm, 25° C, 1 M
60. Frequency Equation
v= cλ ,
where c=3.00∗108ms
61. Electromagnetic Spectrum with 1. Increasing wavelength(λ)2. Decreasing Frequency(v)3. Decreasing Energy(E)
Gamma(γ) Ray --> X-ray --> Ultraviolet(UV) Radiation --> Visible Light(VBGYOR) --> Infrared(IR Radiation) --> Microwave --> Radio(FM --> AM)
62. Energy of a Photon
Ephoton=hcλ
=hv
c= λv h=6.626∗10−34 J∗s
63. Energy of Orbital Equation
En=−2.18∗10−18( 1n2
)
64. Calculating the Change in Energy when an Electron Transitions
∆ E=−2.18∗10−18( 1nf2−1ni2 )
65. Relationship Between Energy Change in Atom and Photon Emitted ∆ Eatom=−Ephoton
66. Periodic Table Trends
67. Types of Chemical Bonds
68. Electron and Molecular Geometries
69. Representing Molecular Geometries on Paper
70. Hybridization Scheme from Electron Geometry
71. Pi(π ) Bond – Forms when orbitals overlap side by side
72. Sigma Bond(σ) – Forms when orbitals overlap end to end
73. Single Bond – Consists of a sigma bond
74. Double Bond – Consists of a sigma bond and a pi bond
75. Triple Bond – Consists of a sigma bond and 2 pi bonds
76. Calculating the Heat Needed to Go from Solid to GasAdd all values for "q" 1. Solid Heatingq=m∗C s , solid∗ΔT in Joules2. Solid melting into liquidq=n∗ΔH fusion in Kilojoules3. Liquid Heatingq=m∗C s , solid∗ΔT in Joules 4. Liquid vaporizing into Gasq=n∗ΔH vaporization in Kilojoules5. Gas warmingq=m∗C s , solid∗ΔT in Joules
77. Types of Intermolecular Forces
78. Phase Diagram
79. Solubility of Gas Sgas=kH∗Pgas
where Sgas is solubility of the gas(usually in M), kH is constant of proportionality(called the Henry's law constant), depends on specific solute and solvent and also on temperature, and Pgas is partial pressure of gas(usually in atm)
80. Raoult’s Law - An equation used to determine vapor pressure of a solutionΔP=P°solvent−P solution=i∗X solute∗P° solvent Psolution is the vapor pressure of solution, Xsolvent is mole fraction of solvent, i is vant hoff’s factor, and P°solvent is vapor pressure of the pure solvent at the same temperature
81. Vapor Pressure Lowering(ΔP) - The difference in vapor pressure between the pure solvent and the solution ΔP=P°solvent−P solution
Psolution is the vapor pressure of solution and P°solvent is vapor pressure of the pure solvent at the same temperature
82. Vapor Pressure for Volatile Solute Psolute=i∗X solute∗P°solutePsolvent=i∗X solvent∗P° solvent Ptotal=P solute∗P solvent
83. Calculating the Amount that the Freezing Point is Lowered ΔT f=FPsolvent−FP solution=i∗mxK f
where ΔTf is change in temperature of freezing point in Celsius degrees, usually reported as a positive number, I is the vant hoff factor, m is molality of the solution in moles solute per kilogram solvent, and K f is freezing point depression constant for solvent
84. Calculating the Amount That the Boiling Point is Raised ΔT b=BP solution−BPsolvent=i∗mx Kb
where ΔTb is change in temperature of boiling point in Celsius degrees(relative to boiling point of pure solvent), I is the vant hoff factor, m is the molality of solution in moles solute per kilogram solvent, and Kb is boiling point elevation constant for solvent
84. Equation for Osmotic Pressureπ=i∗M∗R∗Twhere π is the pressure required to stop the osmotic flow, i is vant hoff’s factor, M is the molarity of the solution, T is the temperature(in Kelvins), and R is the ideal gas constant(0.08206 L*atm/mol*K)
85. Solution Concentration Terms
86. Crystalline Solids
87. Deviations from Raoult’s Laws
88. Freezing and Boiling point Elevation
89. Rate Law Summary Table
90. Activation Energy
91. Activated Complex
91. Chapter 13 Key Equation and Relations
92. Chapter 14 Key Equation and Relations
93. Some Strong Acids
94. Some Weak Acids
95. Some Strong Bases
96. Some Weak Bases
97. pH of Salt Solutions
98. Chapter 15 Key Equations and Relationships
99. Chapter 16 Key Equations and Relations