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CHM1051 > Equilibria > Flashcards

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
Q

catalyst

A

add or remove - no effect on direction or KC

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26
Q

an acid is

A

a proton donor

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27
Q

a base is

A

a proton acceptor

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28
Q

pH

A

how acidic or basic a solution is

pH = -log[H3O+]
[H3O+] is the concentration of an acid

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29
Q

pOH

A

for OH-

pOH = -log[OH-]
[OH-] is concentration of a base

30
Q

conversions between pH and pOH

A

[H3O+] x [OH-] = 1 x 10^-14

pH + pOH = 14

31
Q

strong acids

A
  • completely dissociate when dissolved in water
  • hydrohalic: HCl, HBr (not HF)
  • oxoacids: HNO3, H2SO4, HClO4 (O atoms outnumber H by at least 2)
32
Q

strong bases

A
  • 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
Q

weak acids

A
  • 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
Q

weak bases

A
  • electron rich nitrogen (lone pairs) such as NH3

- amines (RNH2)

35
Q

polyphonic acids

A
  • acids that are capable of donating more than one proton
  • oxalic and carbonic acid
  • pKa (2) and (1)
36
Q

pKa is

A

the pH at half the volume to equilibrium

37
Q

pKa and Ka formula

A

pKa = -log(subscript 10) Ka

38
Q

conjugate pairs

A

acid 1 pairs with base 1

  • strong acid gives a weak base
  • weak base gives a strong acid
39
Q

ka and kb formulae

A

ka x kb = [H3O+][OH-] = kw

pKa + pKb = pKw = 14

40
Q

can you use ICE to work out pH?

A

yes

41
Q

adding a base to an acid

adding acid to base

A

gets more basic

gets more acidic

42
Q

equivalence point

A

acid has completely reacted with the base

43
Q

strong base with strong acid

A
  • high pH to low pH

- equivalence at 7

44
Q

strong acid with strong base

A
  • low pH to high pH

- eqivelecne at 7

45
Q

weak acid and strong base

A
  • starts less acidic than strong acid
  • equivalence is basic
  • halfway to equivalence point, the concentrations are equal and this is pKa
46
Q

weak base and strong acid

A
  • Strats less basic than strong base
  • equivalence is acidic
  • halfway to equivalence is point is base and conjugate acid concentrations equal pKa
47
Q

buffer

A

solutions that maintain an approximately constant pH even when small amounts of acid or base are added

48
Q

what are bases usually made from? how do they work?

A
  • 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
Q

Henderson-Hasselbach equation

A
  • 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
Q

buffer capacity

A

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
Q

ocean acidification

A
  • 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
Q

buffering in natural waters

A

the process that has been unbalanced:

- buffering of CO2 and dissolved mineral carbonate control pH

53
Q

redox reactions

A

transfer of electrons from one substance to another

54
Q

oxidation

A

loss of electrons (OIL)

- electrons are a product

55
Q

reduction

A

gain of electrons (RIG)

- electrons are a reactant

56
Q

balancing in acidic conditions

A
  • 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
Q

balancing in basic conditions

A
  • 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
Q

spontaneous redox equations

A
  • potential difference between half cells, energy is gained (ei. chemical to electrical)
  • batteries and galvanic cells
59
Q

non-spontaneous equations

A
  • 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
Q

potential difference

A

difference in electrons between two half cells and is measured in volts
1 V = 1J/C

61
Q

oxidation numbers

A

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
Q

ox no rules

A
  • 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
Q

anode

A

negative electrode, electrons produced, oxidation occurs, cations produced

64
Q

cathode

A

positive, electrons gained, reduction occurs, cations are reduced to solids etc.

65
Q

salt bridge

A
  • inert salts such as KCl
  • move and balance change
  • required for current
66
Q

galvanic cell

A

chemical to electrical energy

67
Q

electrodes

A
  • 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
Q

conditions of the electrochemical series

A

25 degrees, 1x10^5 Pa, solution concentration of 1

69
Q

using the electrochemical series

A
  • 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
Q

cell potential

A

standard E of cell = E reduction - E oxidation

positive for galvanic and negative for electrolytic

71
Q

electroplating

A

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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