NOTES ON CHEMISTRY
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İ Minna Pöntinen 2002
(note: "¤" acts as the delta mark (indicating change))

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BASICS\____________________________________________________________________

· Compound = contains more than one element
· Empirical formula: the formula obtained by experiment (ie. CH2)
· Molecular formula: The actual number of atoms of each element in a 
  molecule (ie. C2H4)
· Structural formula: Shows the arrangement of atoms and bonds (ie. H2C=CH2)
· Mole = the number equal to the number of carbon atoms in 12 grams of pure
  12C; Avogadro's number = 6.022 * 10^23 units.
· State symbols: (s) = solid, (l) = liquid, (g) = gas, (aq) = in aqueous
  solution

A  <- Mass number = number of protons+neutrons
 X
Z  <- Atomic number = number of protons&electons

· Isotopes are atoms of the same element that contain a different number
  of neutrons
· Cation = positively charged, anion = negatively charged

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BONDING\___________________________________________________________________

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INTERMOLECULAR FORCES\_____________________________________________________

· Intermolecular forces:
  - force that causes substanes to form a liquid or a solid
  - intermolecular forces occur usually between molecules
· London dispersion forces:
  - due to forces that exist among noble gas atoms and nonpolar molecules
  - the movement of electrons can develope nonsymmetrical electron 
    distribution that produces a temporary arrangement of charge
  - temporary dipole can induce electron distribution of neighboring
    atom(s)
· Dipole-dipole forces:
  - molecules with polar bonds/dipole moments can attract each other
    electrostatically
  - dipole-dipole forces are usually only 1% strong as covalent bonds and
    become rapidly weaker as the distance between dipoles increases
· Hydrogen bonding:
  - strong dipole-dipole forces in molecules in which H is bound to be a
    highly electronegative atom, such as N, O or F
  - polarity is also increased because the very small size of the
    hydrogen atom

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GASES\_____________________________________________________________________

· PV = nRT
· molar volume of an ideal gas at STP (0°C, 1 atm) = 22,42 l

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THERMOCHEMISTRY\___________________________________________________________

· exothermic reaction: heat is released => -¤H
· endothermic reaction: heat is absorbed => +¤H
· Standard condition for thermochemistry calculation: 298K & 1atm
· Hess's law states that the 

· Formulas:
  ¤E = q + w	   //E = energy, q = heat, w = work
  w = -P * ¤V	   //w = work, P = pressure, V = volume
  E = m * c * ¤T   //E = energy, m = mass, c = specific heat capacity,
		     T = temperature

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CHEMICAL KINETICS\_________________________________________________________

· Chemical equilibrium = forward and reverse reaction rates are equal 
· Theactors affecting the rate of reaction are
  temperature, surface area, concentrations and catalyst
· Rate = (conc. of A at time t2 - conc. of A at time t1)/(t2-t1)
  = ¤[A]/¤T (unit (mol/l)/s )
· Rate = k[A]^n; k = rate constant, n = order of reaction
· Rate = k[A]^x[B]^y
· k, x and y can only be found out by experimenting, NOT from equations!
· ln[x] = -kt + ln[Xo]
· ln (k2/k1) = Ea/8.3145 * (1/T1 - 1/T2)

· Units of the rate constant (k):
  1st order: s^-1
  2nd order: dm^3*mol^-1*s^-1 
  3rd order: dm^6*mol^-2*s^-1

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CHEMICAL EQUILIBRIUM\______________________________________________________

· Dynamic equilibrium = concentrations of reactants & productants is const.
· Kp = K(RT)^¤n
· aA + bB -> cC + dD -> k = ([C]^c * [D]^d) / ([A]^a * [B]^b)
· If reaction goes nearly to complation, then k >> 1, and vice versa
· Factors that affect the position of equilibrium:
  temperature, concentrations and pressure

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ACIDS & BASES\_____________________________________________________________

· pH -> values from 1 to 14 (7 = neutral, 1 = very acidic, 14 = very basic)
· Classification of acids & bases:
  - Inorganic: H2SO4 (A), HCl (A) ~CO2- 3 (B), ~OH- (B)
  - Organic: H-COOH (A), CH3COOH (B), aminos (B)
  - Strong: inorganic, ~OH- (B)
  - Weak: organic, others
· In solutions: acids produce H+ -ions and bases produce OH- -ions
· An acid is a proton doner and a base is a proton acceptor
· conjugate base = everything that remains of the acid molecule after the
  proton is lost
· conjugate proton is formed when a proton is transferred to the base

· Acids and bases dissociate in water -> they are protolytes
· Degree of dissociation is a measure of the strenght of the acid
· ex1: HCl + H2O -> Cl- + H3O+
· ex2: NH3 + H2O <-> NH+4 + OH-
· Ka = acid dissociation constant = ([H+]*[A-])/[HA]
· Some strong acids: HNO3, HCl, H2SO4
· Kw = [H+]*[OH-]
· pH = -log[H+]

· Some strong bases (hydroxide salts): NaOH, KOH
· Kb = ([B+][OH-])/[B]
· The Kb of weak bases is usually less than one

· pH = pKa + log([A-]/[HA]) 

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ENTROPY\___________________________________________________________________

· An increase in disorder can result from
  - mixing different types of particles
  - changing the state where the distance between the particles increases
    (ex. water -> steam)
  - the increased movement or particles (eg. heating a liquid or gas)
  - increasing the number of particles
· A reaction is spontaneous if it causes a system move from a less stable
  to a more stable state.

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MISCELLANEOUS NOTES\_______________________________________________________

· Periodicity:
  Atomic radius and ionic radius increase LEFT & DOWN
  Electronegativity increases RIGHT & UP