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Equilibria Flashcards

1
Q

equilibria intro

A

reactions do not go to completion
forwards and backward reactions
reactions are dynamic

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2
Q

how do we understand equilibria?

A
  • concentrations
  • partial pressures in gas phase reactions
  • thermodynamics
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3
Q

at equilibrium

A
  • a reaction has not stopped
  • forwards and back are equal
  • concentration of species is constant
  • partial pressures of gases are constant
  • no net change in mixture composition
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4
Q

activity

A

a measure of concentration that takes into account the interactions in the solution

  • concentration can affect physical and chemical properties
  • extent and nature of solute-solute and solute-solvent interactions
  • structure of solution species
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5
Q

ionic strength

A

summative measure of the charge environment
- a characteristic of solution

I = 1/2 sum mi x zi^2

mi is molality or molarity

zi is the charge on the ion

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6
Q

have happens at high ionic strengths

A

greater than 0.1M

effective and nominal concentrations are not the same

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7
Q

what does activity represent?

A
  • the non-ideal behaviour brought about by intermolecular forces
  • can represent a decrease in ion concentration because of counter-ions in the immediate vicinity

in a dilute solution there are electrostatic forces between ions, in a high concentration of salts, there is a weakened attraction because of interference with other ions

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8
Q

equation for activity

A

ai = (yi mi)/standard mi

it is dimensionless (no units)
yi is the activity coefficient which corrects for non-ideal behaviour
standard as the solution has been indefinitely diluted at standard molality and P = 1 bar where solute interactions are negligible. standard mi = 1

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9
Q

other activity equations

A

ai = yimi (because standard mi is 1)

ai = mi (approx) in dilute solutions as yi = 1

ai = ci (approx) as molarity is a common measure of concentration so activity = concentration (mol/L)

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10
Q

are there some more activity equations to learn?

A

yes

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11
Q

what are the important points of activity?

A
  • activity of a dilute solution is basically equal to the numerical value of its concentration (mol/L)
  • the activity of an ideal gas is basically equal to the value of its partial pressure (1 bar)
  • the activity of any pure substance in its standard state is 1
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12
Q

what is the equilibrium constant?

A

concentrations when equilibrium is established

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13
Q

notes on Kc

A
  • temperature dependent (measure temp)
  • pure solids and liquids are not included (activity = 1)
  • dimensionless
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14
Q

Kc values

A

less than 10^-3 reactants favoured
greater than 10^3 products favoured
between these points, at equilibrium

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15
Q

Kp

A

equilibrium constant in terms of pressure

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16
Q

Ka

A

the acid dissociation constant (Kb for a base)

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17
Q

Kw

A

water self-ionisation constant (involved with Ka and Kb)

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18
Q

Ksp

A

solubility product constant for when salts hardly dissolve

Q > k in supersaturated and precipitation

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19
Q

Kstab

A

or beta, the overall stability constant

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20
Q

Van’t Hoff equation

A

ln k = -(standard enthalpy)/RT + c
this is the linear form

can be modified for two temperatures (exothermic if ln k increases as 1/T increases)

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21
Q

Le Chatelier’s Principle

A

when a change is made to the conditions of a dynamic equilibrium the system moves to counteract the change, causing changes in quantities of reactants and products

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22
Q

concentration

A

add reactant or remove product - moves right with no effect on Kc

add product or remove reactant - moves back (no effect on Kc)

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23
Q

pressure

A

increase pressure or decrease volume - moves towards fewer gas molecules

decrease pressure or increase volume - moves towards more gas

no effect on Kc

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24
Q

temperature

A

increase - endothermic direction (right for endo, left for exo) Kc increases for endo and decreases for exo

decrease - exothermic direction (left for endo, right for exo) Kc decreases for endo and increases for exo

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25
catalyst
add or remove - no effect on direction or KC
26
an acid is
a proton donor
27
a base is
a proton acceptor
28
pH
how acidic or basic a solution is pH = -log[H3O+] [H3O+] is the concentration of an acid
29
pOH
for OH- pOH = -log[OH-] [OH-] is concentration of a base
30
conversions between pH and pOH
[H3O+] x [OH-] = 1 x 10^-14 pH + pOH = 14
31
strong acids
- completely dissociate when dissolved in water - hydrohalic: HCl, HBr (not HF) - oxoacids: HNO3, H2SO4, HClO4 (O atoms outnumber H by at least 2)
32
strong bases
- soluble compounds containing O2- or OH- - cations of usually most active metals - M2O or MOH (M = group 1A) - MO or M(OH)2 (M = 2A metals)
33
weak acids
- do not completely dissociate - hydrohalic acid HF - H is not bound to O or halogen - oxoacids where more O than H - organic acids such as RCOOH
34
weak bases
- electron rich nitrogen (lone pairs) such as NH3 | - amines (RNH2)
35
polyphonic acids
- acids that are capable of donating more than one proton - oxalic and carbonic acid - pKa (2) and (1)
36
pKa is
the pH at half the volume to equilibrium
37
pKa and Ka formula
pKa = -log(subscript 10) Ka
38
conjugate pairs
acid 1 pairs with base 1 - strong acid gives a weak base - weak base gives a strong acid
39
ka and kb formulae
ka x kb = [H3O+][OH-] = kw pKa + pKb = pKw = 14
40
can you use ICE to work out pH?
yes
41
adding a base to an acid | adding acid to base
gets more basic | gets more acidic
42
equivalence point
acid has completely reacted with the base
43
strong base with strong acid
- high pH to low pH | - equivalence at 7
44
strong acid with strong base
- low pH to high pH | - eqivelecne at 7
45
weak acid and strong base
- starts less acidic than strong acid - equivalence is basic - halfway to equivalence point, the concentrations are equal and this is pKa
46
weak base and strong acid
- Strats less basic than strong base - equivalence is acidic - halfway to equivalence is point is base and conjugate acid concentrations equal pKa
47
buffer
solutions that maintain an approximately constant pH even when small amounts of acid or base are added
48
what are bases usually made from? how do they work?
- weak acid and its conjugate base - weak base and its conjugate acid - when small amounts of H3O+ or OH- are added, a small amount of one buffer is converted to the other - ions have little effect on pH
49
Henderson-Hasselbach equation
- rearranges ka equation - designs buffers with certain pH ranges pH = pKa + log ([A-]/[HA]) this helps determine the buffer range pH = pKa + log [10]/[1] to pKa + log [1]/[10] (acid over base) pH range = pKa +/- 1 but this does not tell us about how resistant a buffer is to pH change
50
buffer capacity
beta = (dCb/dpH) = - (dCa/dpH) the amount of acid dCa or base dCb that must be titrated to give a defined pH change - big beta values indicate a resistant buffer max buffering occurs when the pH = pKa
51
ocean acidification
- more acidification in colder water - partial pressure increases in atmosphere and CO2 moves into oceans - CO2 becomes carbonic acid which has acidic H+ ions - shells, coral and algae with calcium break down
52
buffering in natural waters
the process that has been unbalanced: | - buffering of CO2 and dissolved mineral carbonate control pH
53
redox reactions
transfer of electrons from one substance to another
54
oxidation
loss of electrons (OIL) | - electrons are a product
55
reduction
gain of electrons (RIG) | - electrons are a reactant
56
balancing in acidic conditions
- identify reactants and products for half equations - balance all atoms except O and H - balance O with H2O - balance H with H+ - balance net charge of reactions by adding e- - balance e- in half equations so that e- gained = e- lost - add balanced equations and cancel - check stoich
57
balancing in basic conditions
- follow rules for acidic conditions - add OH- to reactants and products based on H+ - combine OH- and H+ to form water - cancel H2O on both sides - check stoich and charge
58
spontaneous redox equations
- potential difference between half cells, energy is gained (ei. chemical to electrical) - batteries and galvanic cells
59
non-spontaneous equations
- potential differences (ie. electrical energy) is required to overcome delta G for ox and red reactions to proceed (ie. electrical to chemical). - electrolytic cells
60
potential difference
difference in electrons between two half cells and is measured in volts 1 V = 1J/C
61
oxidation numbers
hypothetical charge that an individual atom/ion would possess in a molecule if the shared electrons in the covalent bond were assigned to the most electrogentaive element in the bond
62
ox no rules
- zero for neutral atoms or molecules - charge on an ion - charges add up to charge on molecule - F is always -1 - H is +1 except in metal hydrides - O is -2 - oxidation: increases ox. no - reduction: decrease in ox. no
63
anode
negative electrode, electrons produced, oxidation occurs, cations produced
64
cathode
positive, electrons gained, reduction occurs, cations are reduced to solids etc.
65
salt bridge
- inert salts such as KCl - move and balance change - required for current
66
galvanic cell
chemical to electrical energy
67
electrodes
- metal ion/metal where the metal is an electrode and dissolves - metal ion/metal ion where an inert electrode such as graphite or platinum is used and both metals are dissolved - non-metal ion/gaseous nonmetal where glass electrode incorporates platinum electrode and the non-metal ion is dissolved
68
conditions of the electrochemical series
25 degrees, 1x10^5 Pa, solution concentration of 1
69
using the electrochemical series
- predicts spontaneity - looking for a z shape - strong oxidants become weak reductants at the top left to right (reduced) - strong reductants become weak oxidants at bottom right to left (oxidised)
70
cell potential
standard E of cell = E reduction - E oxidation positive for galvanic and negative for electrolytic
71
electroplating
not spontaneous, electrons are supplied to make the reaction go the other way, polarities are switched (cathode is negative) - electroplating - recharging batteries

CHM1051 (9 decks)

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