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Chemistry Flashcards

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1
Q

charge on PROTON or ELECTRON

A

1.6 x 10^-19 C

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2
Q

rutherford

A

dense +ve nucleus

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3
Q

bohr

A

single e- in H atom

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4
Q

heisenberg

A

cannot know both momentum AND position of an electron simultaneously

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5
Q

hund’s rule

A

e- fill empty orbitals to stay UNPAIRED before doubling up

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6
Q

paili exclusion

A

each e- has set of 4 distinct quantum numbers

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7
Q

quantum numbers

A

n = energy level – MAX e- per energy level = 2n^2
l = angular momentum (s=0, p=1, d=2, f=3)
ml = magnetic quantum (-l to +l)
ms = spin

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8
Q

periodic table trends: RIGHT

A
  • decreasing atomic radius
  • increasing EN
  • increasing IE
  • increasing electron affinity
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9
Q

periodic table trends: UP

A
  • decreasing atomic radius
  • increasing EN
  • increasing IE
  • increasing electron affinity
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10
Q

highest - lowest EN atoms:

A

F O N Cl Br I S C H

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11
Q

3 exceptions to octet

A
  1. incomplete octet
  2. expanded octet, period 3+, d orbitals
  3. odd # of electrons, eg. NO has 7, can’t distribute evenly
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12
Q

ionic, polar covalent, nonpolar covalent

A

ionic = EN > 1.7 – crystal lattice, electrostatics
polar covalent = 0.5 < EN < 1.7
nonpolar covalent = EN < 0.5

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13
Q

formal charge

A

FC = (valence e-) - (nonbonding e-) - (0.5 x bonding e-)

vs oxidation assumes more EN atom gets ALL e- in bond

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14
Q

electronic geometry vs molecular geometry

A

electronic = spatial arrangement of all e-s (lone + bonds)
molecular geometry = spatial arrangement based on coordination # only (ignores lone pairs)

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15
Q

basic molecular goemetry for coordination #

A

2 = linear
3 = trigonal planar (bent)
4 = tetrahedral (trigonal pyramidal, bent)
5 = trigonal bipyramidal (sawhorse, T, linear)
6 = octahedral (square pyramidal, square planar, T, linear)

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16
Q

percent yield

A

% yield = actual/theoretical x100%

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17
Q

avogadros number

A

6.02 x 10^23 molecules/mol

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18
Q

normality

A

concentration
refrence to # of protons (3N = 3 protons)
Molarity = Normality/n
1N HCl = 1M. VS. 1N H2SO4 = 0.5M

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19
Q

combination reactions

A

A + B = C

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20
Q

decomposition reactions

A

C = A + B

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21
Q

combustion reactions

A

CxHy + O2 = CO2 + H2O

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22
Q

single displacement reactions

A

REDOX
A + BC = AC + B

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23
Q

double displacement reactions

A

eg. neutralization
AB + CD = AD + BC (or: acid + base = salt + H2O)

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24
Q

how many litres is 1 mol of gas at STP

A

22.4L

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25
Q

collision theory

A

rate of reaction is proportional to # of collisions per second (that are EFFECTIVE)

rate = total # collisions x fraction effective

use Arrhenius equation

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26
Q

transition state theory

A

transition state has HIGHEST eenrgy (peak on graph)
but it is theoretical structure, can’t isolate

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27
Q

rate laws

A
  • rate = mol / Ls = M/s
  • Rate = k[A]^x [B]^y
  • ONLY REACTANTS in the equation
  • orders ≠ coefficients – coefficients go inside the brackets, eg. 2CO2 = [2X]^y
  • coefficients on RDS = orders of rate law
  • value + units of k are specific for each rnx at a specific temperature
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28
Q

0th order reaction

A
  • Rate = k
  • independent of ANY change in [ ] of reactants
  • only TEMP, CATALYST can change rate
  • k = M/s
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29
Q

1st order reaction

A
  • Rate = k[A]
  • rate directly proportional to ONE reactant
  • curved graph, log gives straight line
  • eg. radioactive decay
  • k = s-1
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30
Q

2nd order reaction

A
  • Rate = k[A][B] OR k[A]^2
  • proportional to both reactants or square of one
  • more curved graph
  • INVERSE gives linear graph
  • k = M-1s-1
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31
Q

MIXED order reaction

A
  • beginning of reaction = 1st order
  • end of rxn = 2nd order
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32
Q

S and G at equilibrium:

A
  • MAX entropy
  • MIN free energy (G = 0)
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33
Q

Q < Keq, Q = Keq, Q > Keq

A
  • Q<Keq = rxn proceeds forward, more reactants
  • Q = Keq = rxn in equilibrium
  • Q>Keq = reverse reaction, more products
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34
Q

le chatellier’s principle

A
  • rxn will shift to offset any stress (concentration, P, V, temp)
  • endothermic (heat is reactant) VS exothermic (heat is product)
  • eg. temp increases, exothermic reaction shifts reverse
  • must have at least ONE gaseous species to be affected by P/V changes
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35
Q

kinetic product

A

made FIRST
low temperature, fast, less energy to form
UNSTABLE

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36
Q

thermodynamic product

A

high temp, slower, more energy to transition state
STABLE (low NRG)
made over time in rxn after kinetic product

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37
Q

standard conditions VS. STP (ideal gases)

A

Standard Conditions: 25 degrees / 298K, 1M, 1atm
STP: 0 degrees/273K, 1atm

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38
Q

fusion

A

solid - liquid
MELTING
vs freezing (liquid - solid)

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39
Q

condensation

A

gas - liquid
occurs at LOW temp or HIGH presssure

40
Q

*

A
41
Q

boiling point

A

where vapour pressure of liquid = ambient pressure
LOW Vapour Pressure = HIGH BOILING POINT

42
Q

phase diagram

A
  • triple point = temp & P where all 3 phases in eqilibrium
  • lines = phase boundaries / lines of equilibrum
  • solid/liquid line goes to infinity
  • liquid/gas line ends at CRITICAL POINT
  • critical point = above, no distinction between liquid/gas phases
43
Q

temperature VS enthalpy

A

temp = avg kinetic energy of particles
enthalpy = thermal NRG, considers how much of substance

44
Q

process functions

A

WORK and HEAT (Q)

45
Q

heat transfer

A

q = mcΔT
qhot = -qcold

46
Q

gibbs free energy

A
  • determines if rxn is spontaneous (-ve)
  • ΔH/ΔS: -/+ (spont), +/- (non), +/+ (high T), -/- (low T)
47
Q

boyle’s law

A

constant T
P1V1 = P2V2

48
Q

charle’s law

A

constant P
V1/T1 = V2/T2

49
Q

gay-lussac’s law

A

constant V
P1/T1 = P2/T2

50
Q

combined gas law

A

P1V1/T1 = P2V2/T2

51
Q

vapor pressure

A

pressure exerted by evaporated particles above a liquid
vapour pressure forces some gas back into a liquid
solubility of a gas increases with increasing partial pressures
high partial pressure of O2 in alveoli forces more O2 into blood

52
Q

kinetic molecular theory assumptions

A
  1. gas particles have no volume
  2. gas particles have no IMF
  3. gas particles move randomly
  4. collisions are elastic
  5. avg. KE of gas particles is proportional to remperature of gas
53
Q

real gases

A

as pressure increases, volume is LESS than predicted due to IMF
very high pressures, volume is MORE due to particle volume
as temp decreases, volume is LESS than predicted due to IMF
very low pressures, volume is MORE due to particle volume

54
Q

solubility rules:

A
  • ALL salts of NH4+ and group 1 metals are SOLUBLE
  • ALL salts of NO3- OR CH3COO- are SOLUBLE
  • SO42- salts are SOUBLE (except: Ca, Sr, Ba, Pb)
  • ALL halides are SOLUBLE (except: with Ag, Pb, Hg)
  • metal oxides INSOLUBLE
  • hydroxides INSOILUBLE (except: alkali metals, NH4, Ca, Sr, Ba)
  • CO3-, PO4-, SO3- INSOLUBLE (except: alkali metals or NH4+)
55
Q

chetalation

A

central cation bound to SAME ligand in multiple places
used to detoxify/sequester toxic metals (eg. Fe chetalation)

56
Q

% composition by mass

A

mass solute / mass solution x100%

57
Q

dilutions

A

M1V1 = M1V2

58
Q

Ksp

A
  • HIGH Ksp = MORE soluble (more dissolves)
  • temperature dependent
  • for gases: as pressure increases, solubility increases
  • use IP (ion product) similar to Q
59
Q

common ion effect

A
  • in equation, x (concentration) is molar solubility of solid
  • in Ksp expression, whichever ion is present in solution already will be [M]
  • equation becomes: eg. Ksp = x(0.1)^2. where 0.1M is concentration of OH- common ion, and equation has 2OH
60
Q

colligative properties

A

solution properties that depend only on [ ] of particles, NOT identity
1. vapour pressure depression
2. boiling point eleveation
3. freezing point depression

61
Q

vapour pressure depression

A

solutes reduce vapour pressure of solvent
reduces evaporation rate
LOWER VP = HIGH BOILING POINT (need more NRG to overcome)

62
Q

freezing point depression

A

solutes interfere with frozen lattixce formation
need more energy to be removed to allow for freezing
decrease freezing point

63
Q

arrhenius acid

A

dissociates into H+ (must have H in formula)

64
Q

bronsted acid

A

donates hydrogen ions

65
Q

lewis acid

A

electron acceptor
lewis acid-base resembles CCD formation
lewis acids often catalytsts
nucleophiles and electrophiles

66
Q

strong acids (fully dissociate)

A
  • HCl, HBr, HI
  • H2SO4
  • HNO3
  • HClO4, HClO3
67
Q

strong bases

A
  • LiOH, NaOH, KOH
  • Ca(OH)2, Ba(OH)4
  • … few other random OH’s
68
Q

pH shortcut

A

pH = -log[H+]
pH = -log( n x 10^-m)
pH = m - 0.n

69
Q

polyvalent ion titration curve

A
70
Q

buffer solutions

A

WEAK acid/base and its conjugate cation/anion
HA + A- or B + BH+

if you double both [ ]’s, pH is the same but the buffering capacity and ability to resist change improves

71
Q

acetic acid buffer system

A

CH3COOH + H2O = H3O+ + CH3COO-

  • add base (OH-) reacts with H3O+, shifts right
  • add acid (H+) reacts with Ch3COO-, make more acetic acid
72
Q

bicarbonate buffer system

A

CO2 + H2O = H2CO3 = H+ + HCO3-
* add OH-, reacts with H+, makes more H2O
* add H+, reacts with HCO3-, makes more H2CO3

73
Q

spectator ions

A

oxidation change does NOT change
do NOT appear in net ionic equation

74
Q

dismutation

A

undergoes BOTH oxidation AND reduction
eg. SOD: 2O2~ + 2H+ = H2O2 + O2

75
Q

exceptions to oxidation #s:

A
  • hydrides are -1
  • peroxides (O2)2- are -1
  • compounds with high EN atom eg. OF2, oxygen is +2
76
Q

galvanic (voltaic) cell:

A
  • ANOX REDCAT
  • G(-) & Ecell (+) = spontaneous
  • cations attracted to cathode
  • electrodes = 2 diff metals
  • cathode = HIGH (+) reduction potential
  • no external voltage source
  • Daniel Cell
77
Q

electrolytic cell:

A
  • ANOX REDCAT
  • G(+) & Ecell (-) = NONspontaneous
  • cations attracted to cathode
  • electrodes = any material (eg. Pb)
  • cathode = LOWER (-) reduction potential, doesn’t want to get reduced but is forced to
  • uses battery
78
Q

moles of metal deposited on electrode:

A

mol = It/nF (it not fun)

79
Q

Faraday

A

F = 10^5 C / mole of electrons

80
Q

cell diagram

A

anode | anode solution (M) || cathode solution (M) | cathode
eg: Daniel Cell: Zn(s) | Zn2+ (1M) || Cu2+ (1M) | Cu(s)

81
Q

concentration cell

A

type of galvanic cell
both electrodes are the same, but solutions different in [ ]
current STOPS when the concnetration of solutions is equal (V = 0)

82
Q

rechargeable batteries

A

battery disspates = galvanic cell
to charge battery, use electroyltic circuit to reverse process back

83
Q

Ni-Cd batteries

A
  • Cd is anode (ox), Ni is cathode (red)
  • rechargeable!
84
Q

reduction potential

A

Ered values are the REDUCTION
more positive = more likely to get reduced
means wants to be the cathode (reversed in electrolytic)
if substance getting OXidized, REVERSE reduction potential sign
more positive cell potential = spontaneous
only thing that changes Ered is identity of electrode (not moles, temp, etc.)

85
Q

cell potential equation:

A

Ecell = Ered (cathode) - Ered (anode)
never multiply by moles, Ered is not changed by anything! just flip signs
if Ecell is positive, G is negative (spontaneous)
always opposite signs

86
Q

isoelectric focusing

A
  • separate AA’s based on pI
  • cations (+) AA’s travel to the CATHODE
87
Q

nernst equaton: conceptual understanding

A
  • E cell is directly proportional to ln(Q) or ln (Keq)
  • when Q or Keq = 1 (when cell in equilibrium, [ ] same) then ln(1) = 0
  • so Ecell = 0 (V=0)
  • **if Ecell is +, Keq > 1 (products favoured) = galvanic cells
  • if Ecell is -, Keq < 1 (reactants favoured) = electrolytic cells**
88
Q

positron

A

same mass as electron
but positive charge

89
Q

“ic” vs “ous” endings: metal ions

A

“ous” = ion with lower charge eg. Cuprous (Cu+), Ferrous (Fe2+)
“ic” = ion with higher charge eg. Cupric (Cu2+), Ferric (Fe3+)

90
Q

electron density and acidity

A

highest Ka = most acidic proton
EWG make proton more acidic (easier to lose)

91
Q

transition metals

A
  • Solutions of compounds with transition metal ions often are colored (both in solid and aqueous forms) because the transition metals have unfilled d orbitals.
  • This allows for the metal-to-ligand or d-d orbital transfer of electrons through absorption of photons in the visible light region of the electromagnetic spectrum that we see as color.
  • If metal atom compound has full orbitals, it WONT be coloured
92
Q

what can form H bonds

A

H bonded to N, O, F

93
Q

how to determine if protein is in native form

A

compare its FUNCTIONALITY
eg. compare enzyme binding affinity

94
Q

similarities/differences between enantiomers

A

SAME = density, boiling point, IR spectra, all physical properties

DIFFERENT = smell, rotation of PPL is opposite

95
Q

Ki

A

concentration of inhibitor which reaction rate is HALF

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