Chemistry Flashcards
1
Q
charge on PROTON or ELECTRON
A
1.6 x 10^-19 C
2
Q
rutherford
A
dense +ve nucleus
3
Q
bohr
A
single e- in H atom
4
Q
heisenberg
A
cannot know both momentum AND position of an electron simultaneously
5
Q
hund’s rule
A
e- fill empty orbitals to stay UNPAIRED before doubling up
6
Q
paili exclusion
A
each e- has set of 4 distinct quantum numbers
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
8
Q
periodic table trends: RIGHT
A
- decreasing atomic radius
- increasing EN
- increasing IE
- increasing electron affinity
9
Q
periodic table trends: UP
A
- decreasing atomic radius
- increasing EN
- increasing IE
- increasing electron affinity
10
Q
highest - lowest EN atoms:
A
F O N Cl Br I S C H
11
Q
3 exceptions to octet
A
- incomplete octet
- expanded octet, period 3+, d orbitals
- odd # of electrons, eg. NO has 7, can’t distribute evenly
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
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
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)
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)
16
Q
percent yield
A
% yield = actual/theoretical x100%
17
Q
avogadros number
A
6.02 x 10^23 molecules/mol
18
Q
normality
A
concentration
refrence to # of protons (3N = 3 protons)
Molarity = Normality/n
1N HCl = 1M. VS. 1N H2SO4 = 0.5M
19
Q
combination reactions
A
A + B = C
20
Q
decomposition reactions
A
C = A + B
21
Q
combustion reactions
A
CxHy + O2 = CO2 + H2O
22
Q
single displacement reactions
A
REDOX
A + BC = AC + B
23
Q
double displacement reactions
A
eg. neutralization
AB + CD = AD + BC (or: acid + base = salt + H2O)
24
Q
how many litres is 1 mol of gas at STP
A
22.4L
25
collision theory
rate of reaction is proportional to # of collisions per second (that are EFFECTIVE)
rate = total # collisions x fraction effective
use Arrhenius equation
26
transition state theory
transition state has HIGHEST eenrgy (peak on graph)
but it is theoretical structure, can't isolate
27
rate laws
* 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
28
0th order reaction
* Rate = k
* independent of ANY change in [ ] of reactants
* only TEMP, CATALYST can change rate
* k = M/s
29
1st order reaction
* Rate = k[A]
* rate directly proportional to ONE reactant
* curved graph, log gives straight line
* eg. radioactive decay
* k = s-1
30
2nd order reaction
* 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
31
MIXED order reaction
* beginning of reaction = 1st order
* end of rxn = 2nd order
32
S and G at equilibrium:
* MAX entropy
* MIN free energy (G = 0)
33
Q < Keq, Q = Keq, Q > Keq
* QKeq = reverse reaction, more products
34
le chatellier's principle
* 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
35
kinetic product
made FIRST
low temperature, fast, less energy to form
UNSTABLE
36
thermodynamic product
high temp, slower, more energy to transition state
STABLE (low NRG)
made over time in rxn after kinetic product
37
standard conditions VS. STP (ideal gases)
Standard Conditions: 25 degrees / 298K, 1M, 1atm
STP: 0 degrees/273K, 1atm
38
fusion
solid - liquid
MELTING
vs freezing (liquid - solid)
39
condensation
gas - liquid
occurs at LOW temp or HIGH presssure
40
# *
41
boiling point
where vapour pressure of liquid = ambient pressure
LOW Vapour Pressure = HIGH BOILING POINT
42
phase diagram
* 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
temperature VS enthalpy
temp = avg kinetic energy of particles
enthalpy = thermal NRG, considers how much of substance
44
process functions
WORK and HEAT (Q)
45
heat transfer
q = mcΔT
qhot = -qcold
46
gibbs free energy
* determines if rxn is spontaneous (-ve)
* ΔH/ΔS: -/+ (spont), +/- (non), +/+ (high T), -/- (low T)
47
boyle's law
constant T
P1V1 = P2V2
48
charle's law
constant P
V1/T1 = V2/T2
49
gay-lussac's law
constant V
P1/T1 = P2/T2
50
combined gas law
P1V1/T1 = P2V2/T2
51
vapor pressure
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
kinetic molecular theory assumptions
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
real gases
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
solubility rules:
* 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
chetalation
central cation bound to SAME ligand in multiple places
used to detoxify/sequester toxic metals (eg. Fe chetalation)
56
% composition by mass
mass solute / mass solution x100%
57
dilutions
M1V1 = M1V2
58
Ksp
* HIGH Ksp = MORE soluble (more dissolves)
* temperature dependent
* for gases: as pressure increases, solubility increases
* use IP (ion product) similar to Q
59
common ion effect
* 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
colligative properties
solution properties that depend only on [ ] of particles, NOT identity
1. vapour pressure depression
2. boiling point eleveation
3. freezing point depression
61
vapour pressure depression
solutes reduce vapour pressure of solvent
reduces evaporation rate
LOWER VP = HIGH BOILING POINT (need more NRG to overcome)
62
freezing point depression
solutes interfere with frozen lattixce formation
need more energy to be removed to allow for freezing
decrease freezing point
63
arrhenius acid
dissociates into H+ (must have H in formula)
64
bronsted acid
donates hydrogen ions
65
lewis acid
electron acceptor
lewis acid-base resembles CCD formation
lewis acids often catalytsts
nucleophiles and electrophiles
66
strong acids (fully dissociate)
* HCl, HBr, HI
* H2SO4
* HNO3
* HClO4, HClO3
67
strong bases
* LiOH, NaOH, KOH
* Ca(OH)2, Ba(OH)4
* ... few other random OH's
68
pH shortcut
pH = -log[H+]
pH = -log( n x 10^-m)
pH = m - 0.n
69
polyvalent ion titration curve
70
buffer solutions
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
acetic acid buffer system
CH3COOH + H2O = H3O+ + CH3COO-
* add base (OH-) reacts with H3O+, shifts right
* add acid (H+) reacts with Ch3COO-, make more acetic acid
72
bicarbonate buffer system
CO2 + H2O = H2CO3 = H+ + HCO3-
* add OH-, reacts with H+, makes more H2O
* add H+, reacts with HCO3-, makes more H2CO3
73
spectator ions
oxidation change does NOT change
do NOT appear in net ionic equation
74
dismutation
undergoes BOTH oxidation AND reduction
eg. SOD: 2O2~ + 2H+ = H2O2 + O2
75
exceptions to oxidation #s:
* hydrides are -1
* peroxides (O2)2- are -1
* compounds with high EN atom eg. OF2, oxygen is +2
76
galvanic (voltaic) cell:
* ANOX REDCAT
* G(-) & Ecell (+) = spontaneous
* cations attracted to cathode
* electrodes = 2 diff metals
* **cathode = HIGH (+) reduction potential**
* no external voltage source
* Daniel Cell
77
electrolytic cell:
* 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
moles of metal deposited on electrode:
mol = It/nF (it not fun)
79
Faraday
F = 10^5 C / mole of electrons
80
cell diagram
anode | anode solution (M) || cathode solution (M) | cathode
eg: Daniel Cell: Zn(s) | Zn2+ (1M) || Cu2+ (1M) | Cu(s)
81
concentration cell
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
rechargeable batteries
battery disspates = galvanic cell
to charge battery, use electroyltic circuit to reverse process back
83
Ni-Cd batteries
* Cd is anode (ox), Ni is cathode (red)
* rechargeable!
84
reduction potential
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
cell potential equation:
**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
isoelectric focusing
* separate AA's based on pI
* cations (+) AA's travel to the CATHODE
87
nernst equaton: conceptual understanding
* 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
positron
same mass as electron
but positive charge
89
"ic" vs "ous" endings: metal ions
"ous" = ion with lower charge eg. Cuprous (Cu+), Ferrous (Fe2+)
"ic" = ion with higher charge eg. Cupric (Cu2+), Ferric (Fe3+)
90
electron density and acidity
highest Ka = most acidic proton
EWG make proton more acidic (easier to lose)
91
transition metals
* 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
what can form H bonds
H bonded to N, O, F
93
how to determine if protein is in native form
compare its FUNCTIONALITY
eg. compare enzyme binding affinity
94
similarities/differences between enantiomers
SAME = density, boiling point, IR spectra, all physical properties
DIFFERENT = smell, rotation of PPL is opposite
95
Ki
concentration of inhibitor which reaction rate is HALF
