GENERAL CHEMISTRY MAIN POINTS

1. C = F 32 * 5/9 F = C + 32 * 9/52. Melting or freezing of H2O is called fusion. Heat of FUSION = 334 J/g3. Density = g/ mL or g/cm3 for solids g/mL for liquids Density = g/L for gases4. Qsp < Ksp (No precipitate) Qsp > Ksp (Yes Precipiate)5. Enthalpy ( H) using bond energy H = endo exo

Exothermic H < 0 favored Endothermic H > 0 NOT favored6. Entropy (S) = products reactants * S increases from Solids Liquids Gases

S > 0 Favored S < 0 NOT Favored 7. pH = -log[H+] [H+] = 10-pH

8. Ka = pH (1.00 3.00) Ka = more dissociation (Strong acids & bases)9. Looking at a reaction and predicting H & S

a. If reaction involves only gases, entropy is related to moles on either side of reactionb. Less moles on product side = S & vice versa. Entropy = - S

10. Rate Law = k[A]X[B]Y X & Y = reaction order a. To find A keep B constant & vice versab. Divide (Higher/Lower)M = (Higher/Lower) Ex : 4M = 16 M = 2c. X + Y = order of reaction k = rate constantd. To find k pick a trial and solve for k. Bigger k = faster reactione. Rate Law is the 3rd when finding kf. 0th Order = ms-1 1st Order = s-1 2nd Order = m-1s-1 3rd Order = m-2s-1

g. To find 2nd Order Rate Law = k[A]2 m/s = k[m]2 m/s * 1/m2 = K m-1s-1

h. 0th Order Reaction Graph

11. Electrolysis = passing of current through an ionic solution. In H2O, its produces H2 & O2. a. H2 produced at anode & O2 at cathode.

12. Charge carried by one mole of electron is called Faradays Constant (F). One F = 96,500 Columbs13. Ions w/ the most negative charges will have the biggest atomic radii14. Higher Ksp = solubility, Lower Ksp = solubility Disregard solids & liquids 15. Henderson-Hasslebalch Equation is used to estimate pH of a solution in a buffer region

a. pKa = -logKab. pH = pKa + log [Conjugate Base] / [Conjugate Acid]c. pOH = pKb + log [Conjugate Acid] / [Conjugate Base]

16. Quantum #s: n = principal energy level, l =sublevel ml = orbital ms = spin +1/2 or -1/2

a. n = 1,2,3,4 (cant be zero) b. S = 0, p = 1, d = 2, f = 3 deals with sublevels (l)c. If S then 1 orbital, if P then 2 orbitals (-1,0,+1) d. If n = 3, & l = 1, it is 3p because s = 0, p = 1, and so on.e. Shape of p orbital = f. Shape of s orbital = , shape of 2s orbital = , Shape of 3s orbital =

17. Reactants | Product || Reactants | Products E = Cathode + Anode a. Anode (Oxidation)|| Cathode (Reduction)b. +E = spontaneous = Reductionc. Emf of a volatic cell always (+)ive, and an electrolytic cell always (-)ived. Galvanic Cells = Anode (-), Cathode (+)e. Electrolytic Cells = Anode (+), Cathode (-) ** Oxidation takes place at anode in both

18. Phase Change: Q= mcT Q = mHf or mHv f = fusion v = vaporization19. Freezing is a Colligative Property meaning it depends on the # of particles present & not their identities

a. Ex: K+ Cl- i = 2 therefore i* concentration i = Vant Hoff Factorb. C6H12O6 (does not dissociation) i = 1 therefore 1* concentrationc. NaNO3 Na+ + NO3- i = 2 therefore 2 * concentrationd. C2H5OH (does not dissociate) i = 1 therefore 1 * concentratione. Vapor Pressure, Freezing Point, Boiling Point, Osmotic Pressure are all colligative property

20. Boiling Point Elevation: Tb = Kb*m*i m = molality T = in C NaCl Na+ + Cl- i = 2 molality = moles of solute/Kg of solution

21. Freezing Point Elevation: Tb = Kb*m*i m = molality T = in C NaCl Na+ + Cl- i = 2 molality = moles of solute/Kg of solution

22. Spectator Ions are the ones which are soluble. (K, Li, Na, Cl, Br, & NO3-) .a. Net Ionic Reaction is writing all soluble materials as ions (w/ charges), all insoluble as it is, &

cancelling out all insolubles. BaSO4 & CaCO3 are insoluble & are commonly asked. 23. Spontaneous Reactions = -H, +S, -G, K > 1, +E

a. +H, +S spontaneous at high temperaturesb. -H, -S spontaneous at low temperaturesc. -H, +S spontaneous at all temperatures

24. State Function only depends on the initial and final state. H, E, G25. 1 Liter = 1000 cm3 1cm3 = 0.001 L 1L of H2O = 1000g H2O26. -Decay ZXA Z-4YA-2 + 4He2 238U92 234Th90 + 27. Decay 138I53 138Xe54 + 0e-128. Positron Decay (+ decay) 12557La -> 12556Ba + 01e29. 235U92 + 1n0 ? + 139Ba56 + 94Kr36

a. 235+1 = 236 139+94 = 233 236-233 = 3 56+ 32 = 92 Therefore, 3n030. Gamma Radiation 40Ar18 = 40Ar18 + Gamma Radiation31. Temperature favors endothermic reaction. Temperature favors exothermic reaction. But if

Kw (equilibrium constant of H2O) , endothermic is favored32. In a closed container, vapor pressure only depends on the temperature. Surface area does not affect vapor

pressure33. A body centered cubic unit cell has 2 atoms. A face centered unit cell has 4 atoms.34. Osmotic Pressure () = MRTi M = molarity R = 0.0821 T = Temp in K i = Vant Hoff Factor

35. Daltons Law of Partial Pressure PTotal = PA + PB + PC PA = PT (nA)/(nTotal) n = # of moles36. Ionization Energy is the energy needed to remove the most loosely held electron from an atom or ion in

the gas phase. Easy to pull = ionization energy Hard to pull = ionization energy37. Potential energy is in gas phase than in solid or liquid phase38. MM = Surface Tension. Surface Tension is how strong the forces b/w the particles are39. Intensive Property does NOT depend on the size/amount of sample. Extensive Property is proportional to

the amount of materials in a sample40. The SI unit of pressure is PASCALS41. Osmolarity = osmotic concentration

a. Osmolarity = OSmoles of solute / Liters of Solution42. 6 Types of Chemical Reactions:

a. Combustion Reaction when hydrocarbon reacts w/ O2 CO2 + H2O. Always exothermicb. Synthesis Reaction 2 or more simple compounds react to form complex products.

i. 8 Fe + S8 8FeSc. Decomposition Reaction Complex molecules breaks to form simple products 2H2O 2H2 + O2d. Single Replacement Reaction when an element trade places w/ another element in a compounde. Double Replacement Reaction anions & cations of 2 different molecules switch places, forming

two different compoundsf. Acid-Base Reaction a double replacement reaction, when an acid & base react to form ionic salt

& H2O43. Dis-Proportanation Reaction when the same compound gets oxidized & reduced. Therefore it is a

REDOX reaction44. Emf & G G = -nFE0cell n = # of electron exchanged F = 96,500 J/V45. Emf & Concentration If you dont have a concentration of 1 M/L

a. Use Ecell = (RT/n) * lnQ lnQ = [products]/[reactants]b. If given Temp is 250 use Ecell = E0cell (0.0591 V/n) * log Q n = # of electrons exchanged

46. Amphoteric = molecule or ion that can act as an acid or base.

a. Must have H+ ions. Ex NH3, H2O, HSO3-

47. Normality (N) > Molarity (M). a. In HBr, HI, HNO3 etc. N = Mb. For OHs, NaOH, KOH etc. N = Mc. If you see two OHs [Ca(OH)2], N is twice that M.d. Similarly, if there are two H, Normality is twice than Molarity [H2SO4]e. Then use NaVa = NbVb (PROBLEM 135)

48. Gases solubility w/ temperature & solubility w/ pressure. 49. Solids solubility w/ temperature.50. Oxides of metals are basic [Na2O, CaO]. 51. Oxides of NON-METALS are acidic [CO2 H2CO3, P4O10 H3PO4]52. Only temperature can change are Ka, Kb, Kw, Keq, Ksp53. When one atom donates both electrons , a coordinate bond is formed

a. The lone pairs from N bonds w/ H+ forming NH4+

54. Strong Acids HClO4, H2SO4, HNO3, HI, HBr, HCl55. Strong Bases LiOH, KOH, NaOH, RbOH, Ca(OH)2, Ba(OH)256. G = H TS (Goose Hunters Take Shotguns)

a. At boiling point G = 0 because its at equilibrium b. G = fast reactionc. G = slow reactiond. If temperature is , H > 0 (endothermic)e. If temperature is , H < 0 (exothermic)

57. Ligand is a Lewis Base that is part of a complex ion. Ligands are part of co-ordinate covalent bond (52)a. [Ag(NH3)2]+ NH3 is the ligandb. [Zn(CN4)]-2 CN4 si the ligandc. Common ligands are H2O, NH3 (ammonia), Co, Cyanide (CN-), Cl-, OH-

58. More reactants = less change of reaction59. VIBGYOR Red = Longest wavelength, E, F Violet = Shortest Wavelength, E, F 60. Color is associated w/ transition metals. When you see blue, think copper (Cu)61. A catalyst will change the reaction of mechanism, specific rate constant, activation energy of forward, &

backward reaction.62. 3 Gas Laws: (Big Clit Girl)

a. Boyles Law = P1V1 = P2V2b. Charles Law = V1/T1 = V2/T2 *Temp = Kc. Gay Lussacs Law = P1/T1 = P2/T2 *Temp = K

63. Cosmic Rays (High Frequency & Energy) > Gamma Rays > X- Ray > Ultraviolet > Visible > Microwave > Radio-wave (Long Wavelength)

64. Vapor Pressure is the tendency of particles to escape from liquid.a. High Vapor Pressure = Low boing point & vice versab. A substance with a high vapor pressure @ normal temperature is called volatilec. To find Vapor Pressure use Raoults Law = XAPA X = mole fraction Po = vapor

pressure of A65. Exothermic & Endothermic Reactions

a.66. In EOCell, the most positive EOCell is the best reducing agent, and the most negative EOCell is the best

oxidizing agent.67. Diamagnetic has all the subshells filled. Ex He (1s2), Be (1s2, 2s2)68. Paramagnetic - has unfilled subshells Ex Li (1s2, 2s1), N (1s2, 2s2, 2p3)69. Triple point is where solid, liquid, and gases are in equilibrium.

a. Critical Pressure is the pressure which is required to liquefy a gas.b. It occurs between vapor-liquid, & liquid-liquid pointsc.

70. An ALLOY is a solid solution of 2 or more metals71. Molality is temperature independent because its moles of solute/Kg of solvent

a. Mass does not change with temperature72. Molarity is temperature dependent because its moles of solute/Liters of solution

a. Volume (Liters) will change with temperature73. Half-Life:

a. 1st Order Reaction: Kt1/2 = 0.693 t = half-life k = 1st order rate constantb. 2nd Order Reaction: t1/2 = 1/(A0) A0 = initial concentration

74. Half Life y = 1/2n n = given half-life y = fraction of original material 75. H using Bond Energy H = H (bonds broken) - H (bonds formed)

a. Only take into account the bonds which gets broken and which forms76. Small specific heat will give rise to Ts & vice versa because C & T are indirectly proportional

(Q = mCT)77. n = 1 has 2 e-, n = 2 has 8 e-, n = 3 has 18 e-

a. Use 2 (n)2 where n = the shell. Ex: n=2 then 2(2)2 = 878. Dow Process extracting Bromine79. Hall Process Production of elemental aluminum80. Haber Process Production of ammonia (NH3) from N2 & H2. N2 + H2 2NH381. Metallurgical Process process applied to metals & alloys to change shape, structure/composition.

82. Isotopes same # of protons but different mass #. Ex 16O8 17O883. Rate of Diffusion & Effusion rate1/rate2 = Molar Mass of B/Molar Mass of A84. Empirical & Molecular Formula

a. Ex 40% of Oxygen, 20 % of Chlorinei. Find 40/ MM of oxygen & 20/MM of Clii. Divide all with the smallest product which iii. Multiply so that all of the coefficients are whole numbers

b. Ex - 40% of Oxygen, 20 % of Chlorine & total MM given (269 g/mol)i. (.4)(269)/MM of Oxygen (.2)(269)/MM of Clii. That will give you empirical formula, then find molecular formula

85. Formal Charge = Group # - [ (shared electrons) + (unshared electrons)]a. Formal Charge of Xe = 8 - [1/2 (4) + (6)] = 0

86. Actual mass of atomic nucleus is a little smaller than the sum of protons & neutrons called Mass Defecta. Mass defect is due to matter converted to energy. This energy is called BINDING ENERGY. b. Binding Energy hold nucleons together in nucleusc. Intermediate Nuclei are more stable than small & large nucleid. Iron (Fe) has the highest binding energy

87. Types of Covalent Bondsa. Polar Covalent Bond atoms w/ electronegativity difference of to 0.4 to 1.7b. Non-Polar Covalent Bond occurs b/w atoms w/ same electronegativity. H2, Cl2, O2 (diatomic

nucleus)c. Co-Ordinate Covalent Bond

i. Lewis Acid accept electrons from covalent bondii. Lewis Base donate electrons from covalent bond

iii.88. The more positive, the smaller the radius, the more the negative, the larger the radius. K+ > Ca2+ > Al3+

89. Monomer molecules that may bind chemically to other molecules to form a polymer. Ex amino acids90. Dimer Consists of structurally similar monomers joined by bonds that can be either strong/weak,

covalent or intermoleculara. Homo-dimer when two molecules are identical (A-A)b. Hetero-dimer when two molecules are different (A-B)

91. Intermolecular Forces a. Dipole Dipole - in Polar molecules. Its prominent in solid, liquid, & have higher boiling points.

Also called Van Der Waals Forcesb. Dispersion Forces aka LONDON FORCES. They are non-polar. Prominent in large

molecules, but also present in all molecular solids, noble gases, & molecules w/ dipoles.

c. Hydrogen Bonding when hydrogen bonds w/ Fluorine, Oxygen, & Nitrogen. Have boiling point.

92. AX2 Linear (1800) AX3 Trigonal Planar (1200) AX4 (109.50) AX5 Trigonal Bi-pyrimidal (1200, 900) AX6 Octahedral (900, 1800)

93. Drawing Lewis Structurea. Place the least electronegative atom in the middleb. Check for formal chargesc. If resonance form exists, the one w/ charge on the most electronegative at is the most stable

94. % Yield = Actual/Theoretical * 10095. At Half-Equivalence Point pH = pKa96. An intermediate is used up in a reaction. A catalyst is not consumed, but regenerated

a.97. Heisenberg Uncertainty Principle one cannot simultaneously know the position & velocity of an

electron precisely98. Hunds Rule every orbital in a subshell is singly occupied with one electron before any orbital is doubly

occupied. Also, electron in a singly occupied orbitals have the same spin.99. Pauli Exclusion Principle no two electrons in an atom have same four quantum numbers 100. Work = -P V P = Pressure V = Volume

a. E = q + w E = internal energy change q = heat w= workb. If q is absorbed, then +q & if q is released, then qc. +w if work done on system BY surroundings, w if work done by system ON surroundings

101. Beers Law A = bc A = Absorbance @ particular wavelength b = cell path = extinction coefficient constant c= concentration