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1

Rate of a reaction

proportional to the concentration of a particular reactant raised to the power of it's order.

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2

K

Proportionality/Rate constant

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3

Finding orders of reaction

Only through experiments

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4

Concentration-time graphs plotting

Gas collection, mass loss and colour change

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5

Monitoring rate with a colorimeter

Measures light absorbed by solution. Select a filter with complimentary colours; zero colorimeter with water; plot a calibration curve with known concentrations; take absorbance readings at timed intervals; compare to curve.

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6

Concentration-time graphs - zero order

Gradient = k

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7

Half-life

time taken for half of a reactant to be used up

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8

Concentration-time graphs - first order

Constant half-life. Exponential decay

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9

Finding K on C-T graphs, first order

substitute values of concentration and rate from tangent into rate equation; K=ln2/half-life

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10

Rate-concentration graph - zero order

y-intercept = k

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11

Rate-concentration graph - first order

Gradient = k

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12

Rate-concentration graph - second order

k = gradient of rate-concentration^2 graph

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13

Clock reaction

Calculates initial rate. Time, t, is measured for a visual change. Initial rate is proportional to 1/t

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14

Iodine clock reaction

Formation of iodine, t=time until first appearance of iodine colour. Starch often added to form a complex - dark blue-black colour appears when Na2S2O3 (Sodium thiosulfate) is used up.

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15

Assumptions on average rate of reaction

Constant, and same as initial rate (accurate providing less than 15% of the reaction has taken place)

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16

Why is it expected that reactions happen in a series of steps

Unlikely that multiple species collide simultaneously

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17

The slowest step

rate determining step

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18

The rate-determining step

Only includes species on the rate equation, orders of which match the number of species in the rds

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19

Increase in temperature

increase in rate and rate constant. Shifts boltzmann distribution to the right, increasing the proportion of particles that exceed the activation energy.

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20

The Arrhenius Equation

k=Ae^(-Ea/RT)

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21

lnk = lnA - Ea/RT

Gradient = -Ex/R ; y-intercept = Ln A

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22

Heterogeneous equilibria

Species of different states. Only gases and solutions included in Kc as solids/liquids have constant concentrations.

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23

Kp

equilibrium constant in terms of partial pressures of gases, which are proportional to it's concentration.

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24

Mole fraction of A

moles of A/total moles ; Volume of A/total volume

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25

Sum of the mole fractions

= 1

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26

Partial pressure of A

mole fraction of A x total pressure

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27

Kp expression

includes ONLY gases

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28

Application Kp

Endurance athletes train at high altitude to produce haemoglobin with extra oxygen containing abilities

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29

K is only affected by

temperature

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30

Exothermic reaction, increase temperature

Kc decreases (decrease in product yield) ; Kp: Partial pressures of reactants increase, shift to left, lower Kp value

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31

Endothermic reaction, increase temperature

Kc increases (increase product yield) ; Partial pressures of products increase, shift to right, higher Kp value

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32

Change in concentration or pressure

does not affect K

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33

If total pressure doubles

so does each partial pressure

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34

Increase pressure, Kc

Kc doesn't change. Ratio of expression now less/more than Kc value. Concentrations of reactants/products increase/decrease to return system to equilibrium and restore ratio of Kc

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35

Conjugate acid-base pair

Two species that can be interconverted by transfer of a proton

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36

pH scale

Measures concentration of H+ ions.

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37

pH =

-log[H+]

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38

[H+] =

10^-pH

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39

Change in 1 pH unit

10 fold change in H+

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40

In Strong acids

[H+] = [HA]

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41

Ka

acid dissociation constant

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42

Ka =

[H+][A-]/[HA]

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43

The stronger the acid

the LARGER the Ka value and the SMALLER the pKa value

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44

The weaker the acid

the SMALLER the Ka value and the LARGER the pKa value

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45

Approximation 1 on weak acids

HA dissociates to produce equal ammounts of H+ and A- ; [H+] = [A-]

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46

Approximation 2 on weak acids

Dissociation is small, so [HA]eqm = [HA]start

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47

Ka for weak acids

[H+]^2[HA]

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48

Ka determined experimentally

Prepare standard of weak acid at known concentration, measure pH with a pH metre

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49

Limitations of approximation 1 of weak acid

If ph>6, then [H+] from dissociation of water will be significant compared to acid. So assumption breaks down with very weak/dilute acids

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50

Limitations of approximation 2 of weak acid

Breaks down when [H+] is significant. , so doesnt work for stronger weak acids with Ka>10^-2 and very dilute solutions.

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51

Kw

The ionic product of water, 1x10^-14

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52

Kw =

[H+][OH-]

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53

Change in temperature effect Kw

Kw varies

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54

Water is neutral so

[H+] = [OH-]

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55

For pH values with whole numbers

indices for [H+] and [OH-] add up to -14

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56

A strong base

an alkali that completely dissociates it OH- in solution.

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57

pH of Strong base

[OH-] = [NaOH], and Kw

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58

Ammonia is a weak base

dissolves in water releasing OH- ions from water molecules. NH3 + H2O

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59

Buffer Solution

A system that minimises pH changes on addition of small amounts of an acid or a base

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60

Preparing buffer solutions from weak acid and it's salt

Solution of a weak acid and it's salt. HA only partially dissociates a small amount - source of HA. Salt completely dissolves, provides A- ions.

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61

Preparing buffer solutions from partial neutralisation

Add an aqueous solution of an alkali to an excess of weak acid. Weak acid is partially neutralised, forming conjugate base and some HA left over.

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62

Conjugate base of buffer solution removes added acid

[H+] increases, reacts with A- ions, equilibrium shifts to the left, removing most H+ ions.

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63

Weak acid of buffer solution removes added alkali

[OH-] increases, reacts with H+ ions, HA dissociates, shifting equilibrium to the right to restore H+ ions.

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64

Buffer most effective when

[HA]=[A-] ; pH of buffer is same as pKa of HA

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65

Fine tune pH of buffer solution

ratio of concentrations of the weak acid and it's base can be adjusted

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66

pH buffer solution

[H+] = Ka x [HA]/[A-]

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67

Ideal pH for blood plasma

7.35-7.45

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68

Buffer system in blood

Carbonic acid-hydrogencarbonate

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69

Carbonic acid-hydrogencarbonate

H2CO3

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70

Body prevents build up of H2CO3

Converting it into CO2 which is then exhaled by lungs

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71

pH titration curve - equivalence point

the volume of one solution that reacts exactly with the volume of the other. Amounts use match stoichiometry of the reaction.

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72

acid-base indicator

a weak acid that has a distinctly different colour from its conjugate base.

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73

Methyl orange

Red in acid, yellow in alkali

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74

Added Alkali

shift right

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75

Added Acid

shift left

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76

At the end point

[HA] = [A-], Ka = [H+], pKa = pH

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77

Indicator must have a colour change

which coincides with vertical section of curve. Ideally, end point = equivalence point

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78

weak acid - weak base reaction

No indicator is suitable

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79

Standard enthalpy change of solution

one mole of a solute dissolves in a water under standard conditions

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80

Lattice Enthalpy

The enthalpy change that accompanies the formation of one mole of an ionic compound from its gaseous ions under standard conditions

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81

standard enthalpy of formation

The change in Enthalpy that accompanies the formation of 1 mole of a compound from its elements with all substances in their standard states.

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82

Standard enthalpy change of atomisation

the enthalpy change when one mole of gaseous atoms is formed from an element under standard conditions

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83

First ionisation energy

The energy required to remove one electron from each atom in one mole of gaseous atoms to form one mole of gaseous 1+ ions

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84

First electron affinity

The enthalpy change that takes place when one electron is added to each atom in one mole of gaseous atoms to form one mole of gaseous 1- ions

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85

Enthalpy change of Hydration

The enthalpy change that accompanies the dissolving of gaseous ions in water to form one mole of aqueous ions

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86

Sol H

q/n

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87

Why are first electron affinities exothermic?

Because electrons are attracted to nucleus

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88

Why are second electron affinities endothermic?

Because the ion is already negative so repels the electron. So energy is needed to force it in

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89

Effect of ionic size on Lattice Enthalpy

ionic radius increases, attraction decreases, less negative, melting point decreases

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90

Effect of ionic charge on Lattice Enthalpy

ionic charge increases, attraction increases, more negative, melting point increases

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91

Effect of ionic size on Hydration

ionic radius increases, attraction between ion and water decreases, less negative

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92

Effect of ionic charge on Hydration

ionic charge increasing, attraction increases, more negative

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93

Soluble

attraction between ions in ionic lattice needs to be broken

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94

If sum of hydration is larger than lattice enthalpy

then enthalpy change of solution is exothermic and should dissolve.

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95

Enthalpy change of solution doesn't take into account

temperature and entropy

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96

Entropy

dispersal of energy within the chemicals making up the chemical system

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97

Increasing entropy

Solids -> Liquids -> Gases

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98

If a system changes to become more disordered

entropy change will be positive

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99

Feasibility

describes whether a reaction is able to happen, energetically

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100

Free energy

Overall change in energy during a chemical reaction.

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