Web viewAtomic Mass. AM = ∑ fraction of isotopes *(mass of isotopes) Avogadro’s Number . 1 mol=6.022* 10 23 particles . Relations. 1 c m 3 =1 mL . 1 m 3 =1 dL

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