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Physical Flashcards

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1
Q

Enthalpy change

A

The amount of heat released or absorbed by a chemical reaction, carried out at constant pressure

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2
Q

Exothermic reaction

Delta H=

A

A reaction where heat energy is released to the surroundings

Negative

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3
Q

Endothermic reaction

Delta H =

A

A reaction where heat energy is absorbed from The surroundings

Positive

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4
Q

Average bond enthalpy

A

The average enthalpy change for the breaking of one mole of bonds in gaseous molecules

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5
Q

Standard conditions

A

25 degrees C / 298k

1 atmosphere/ 100 kPa

Solution must have conc 1 moldm^-3

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6
Q

Standard enthalpy change of reaction

A

The enthalpy change that accompanies a reaction in the molar quantities expressed in a chemical equation under standard conditions, all reactants and products in their standard states

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7
Q

Standard enthalpy change of neutralisation

A

The enthalpy change that accompanies the formation of one mole of H2O from neutralisation, under standard conditions

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8
Q

Standards enthalpy change of formation

A

The enthalpy change when one mole of a compound is formed from its elements, in their standard states under standard conditions

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9
Q

Standard enthalpy change of combustion

A

The enthalpy change for the complete combustion of one mole of a substance under standard conditions, all reactants and products in their standard states

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10
Q

Hess’s law states that

A

The enthalpy change of a reaction depends only on the initial and final states an is independent of the route taken

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11
Q

First law of thermodynamics

A

Energy can be converted from one form to another and cannot be created or destroyed

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12
Q

Delta H is negative

4

A

System releases heat energy to the surroundings
Enthalpy of the system decreases
Temperature of surroundings increases
Enthalpy change is EXOTHERMIC

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13
Q

Delta H is positive

4

A

System absorbs heat energy from the surroundings
Enthalpy of the system increase
Temperature of the surroundings decreases
Enthalpy change is ENDOTHERMIC

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14
Q

When measuring temperature what are you recording the temperature of

A

Surroundings and NOT the system

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15
Q

Enthalpy change of a reaction calculation

A

Enthalpy of products - enthalpy of reactants

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16
Q

A molecular chemical reaction involves

A

Breaking covalent bonds in the reactant molecules and forming new covalent bonds in the product molecules

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17
Q

Breaking bonds is

A

Endothermic

Absorbs energy

Positive

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18
Q

Making/ forming bonds is

A

Exothermic

Releases energy

Negative

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19
Q

The enthalpy change of a reaction depends only on

A

the initial and final states and is independent of the route taken

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20
Q

The rate of a chemical reaction is

A

The change in concentration of a reactant or a product per unit time

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21
Q

Effect of concentration on rate

Collision theory

A

As the concentration of reactant molecules increases, the rate of rection increases
At a higher concentration there are more molecules in a given volume
More frequent successful collisions

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22
Q

Effect of pressure on rate

Collision theory

A

When the pressure of a gas is increased, the gas molecules are pushed closer together
The number of gas molecules in a given volume increases
More frequent successful collisions occur

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23
Q

The effect of temperature on rate

Collision theory

A

As the temperature of a reaction mixture is increased, the rate of reaction increases
At a higher temeperature, the average energy of the molecules increases
A greater proportion of the molecules have energy greater than or equal to the activation energy
More frequent successful collisions occur

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24
Q

Activation energy

A

The minimum energy required for a reaction to take place, by the breaking of bonds in the reactants

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25
Q

Catalyst

A

Increases the rate of reaction without being used up by the overall reaction
It allows the reaction to proceed via. Different route with lower activation energy

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26
Q

Heterogenous catalysis

A

Catalysis of a reaction in which the catalyst has a different physical state from the reactants

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27
Q

Homogenous catalysis

A

Catalysis of a reaction in which the catalyst and the reactants are in the same physical state

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28
Q

Key advantages of using a catalyst in an industrial reaction

A

Low temperature gives;

  • reduced energy demand
  • less CO2 emissions as less combustion of fossil fuels
  • less cost
  • Increased sustainability

Alternative reaction with higher atom economy and less waste can be used

Using an enzyme generates very specific product, no by products

Enzymes operate at close to room temperature and pressure

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29
Q

Effect of temperature

Boltzman

A

At a higher temperature, the average energy of the molecules increases
A greater proportion of the molecules have energy greater than or equal to the activation energy
More frequent successful collisions occur
Rate of reaction increases
Boltzman flattens and shifts to the right

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30
Q

Effect of a catalyst

Boltzman

A

A catalyst increases the rate of reaction without being consumed by the overall reaction
A catalyst lowers the activation energy for the reaction by providing an alternative pathway
A greater proportion of the molecules have energy greater than or equal. To the activation energy
More frequent successful collisions occur
Rate of reaction increases in the presence of a catalyst

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31
Q

Characteristics of a dynamic equilibrium

4

A
  1. The rate of the forwards reaction is equal to the rate of the reverse reaction
  2. The system is closed (no materials added or taken away)
  3. The concentration of reactants and products do not change
  4. The macroscopic (big) properties (temperature, pressure, concentration) do not change
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32
Q

Le Chateliers principle

A

When a system in dynamic equilibrium is subjected to a change, the equilibrium position will shift to minimise the effects of the change

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33
Q

Effect of temperature on dynamic equilibrium

A

Exothermic
Increase= left
Decrease = right

Endothermic
Increase= right
Decrease=left

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34
Q

Significance of Kc

A

If Kc>1 the equilibrium position lies to the right

If Kc<1 the position lies to the left

If Kc=1 the equilibrium lies halfway between reactants and products

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35
Q

The effect of temperature on Kc/p

Exothermic forward reaction

A

If the temperature is increased:
-Thea equilibrium position shifts to the left
-to minimise the effect of an increase in temperature , by absorbing energy
-because the reverse reaction is endothermic
-the new mixture will contain more reactants and less products
The value. Of Kc/p decreases

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36
Q

The effect of temperature on Kc/p

Endothermic forward reaction

A
  • the equilibrium position shits to the right
  • to minimise the effect of an increase in temperature, by absorbing energy
  • because the forward reaction is exothermic
  • the new mixture will contain more products and less reactants
  • the value of Kc/p increases
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37
Q

The effect of concentration Kc

6

A

If the reactant concentration is increased
This means the expression for Kc no longer matches the original Kc value
The equilibrium position shifts to the right
The concentration of the product increases
The concentration of the reactant decreases
A new equilibrium is reached where Kc is restored to its original value

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38
Q

The effect of pressure KP

A

If total pressure is increased
All the partial pressures increase
There are more partial pressures of products than reactants so the numerator (top) of KP increases more than the denominator (bottom) of KP
The ratio in the KP expression increases
The equilibrium position shifts to the left so the partial pressure of the reactant increases, increasing the denominator (ratio decreases)
A new equilibrium is reached where KP is restored to its original value

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39
Q

Efffect of a catalyst on Kc and Kp

A

A catalyst increases the rat of both the forward and backward reaction by the same amount
The equilibrium position does not change
Value of Kc and Kp remains constant

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40
Q

The rate of a chemical reaction is

A

The change in concentration of a reactant or product per unit time

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41
Q

Rate equation

A

Rate = k [A]^m [B]^n

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42
Q

Rate at time ‘t’

A

Gradient of the tangent to the curve at time ‘t’

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43
Q

Initial rate of reaction

A

Is equal to the gradient of the tangent at time = 0

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44
Q

The half life of a reactant

A

The time taken for the concentration f the reactant to fall to half of its original concentration

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45
Q

The rate constant k for a first order reaction can be found by usin g the expression

A

K=ln2/t1/2

46
Q

The rate determining step

A

The slowest step of a reaction mechanism of a multi-step reaction

47
Q

Arrhenius equation

A

K= Ae^-Ea/RT

48
Q

An acid is

A

A proton donor

49
Q

A base is

A

A proton acceptor

50
Q

Monobasic, dibasic and tribasic acid

A 
Mono= requires one mole of OH- ions to neutralise it
Di= requires two moles
Tri= requires three moles
51
Q

A conjugate acid base pair is

A

A pair of two species that transform into each other by the gain or loss of a proton

52
Q

Significance of Ka

A

If Ka is a large number;
[H+] and [A-] are large
The equilibrium position is far to the right Ka>1
A lot of the HA acid is dissociated into its ions (strong acid)

If Ka is a small number:
[H+] and [A-] are small
The equilibrium position is far to the left Ka<1
A lot of the HA acid is not dissociated (weak acid)

53
Q

Ka equation

A

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

54
Q

pKa

A

-log Ka

55
Q

Ka calculation

A

10^-pKa

56
Q

Ka values only depend on

A

Temperature

57
Q

Kw calculation=

A

[H+][OH-]

58
Q

Kw value

A

1.00x10^-14

59
Q

PH

A

-log [H+]

60
Q

[H+]

A

10^-PH

61
Q

Importance of Kw

A

Acidic solution [H+]>[OH-]

Neutral solution [H+] =[OH-]

Basic solutions [H+] < [OH-]

62
Q

Equivalence point=

A

The point in a titration at which they volume of one solution has reacted exactly with the volume of the second solution

This matches the stoichiometry of the reaction taking place

[H+]= [OH-]

63
Q

Strong acid and strong base

A

Vertical section

PH 3-11

64
Q

Strong acid and weak base

A

PH 3-7

65
Q

Weak acid and strong base

A

PH 7-11

66
Q

Weak acid and weak base

A

No vertical section

67
Q

Indicators change colour with PH

Addition of acid

A

HA -> H+ + A-

Causes an increase in[H+]
The indicator equilibrium position will shift to the left to minimise the increase in [H+] by reacting H+ ions with conjugate base A- to form more HA

68
Q

Indicators change colour with PH

Addition of alkali

A

HA-> H+ + A-

OH- + H+ -> H2O
This causes a decrease in [H+]
The indicator equilibrium position will shift to the right to minimise the decrease in [H+]. More of the weak acid indicator (HA) will dissociate into H+ ions (and A-)

69
Q

The end point of an indicator

A

The Ph at which there are equal concentrations of the weak acid (HA) and its conjugate base (A-) forms of the indicator

The colour of an indicator at its endpoint is midway between the colours of the acid (HA) and conjugate base (A-) forms

70
Q

Choosing an indicator

A

The PH range/ end- point of the indicator must coincide with the vertical section of the PH curve

71
Q

A buffer solution is

A

A mixture that minimise PH changes on addition of small amounts of an acid or base

72
Q

A buffer solution is a mixture of

A

Weak acid

Conjugate base of the weak acid

73
Q

Two ways in which a buffer solution can be prepared

A

Mix a weak acid an one of its salts together

React excess weak acid with a strong alkali

74
Q

When a small amount of acid (H+) is added to the buffer

CH3COOH CH3COO- + H+

A

Conjugate base , CH3COO- , reacts with added H+ ions (and ‘mops’ them up)
CH3COO- + H+ ->
Equilibrium shifts to the left to minimise the effect of an increase in [H+]
Most of the added H+ ions are removed

75
Q

When a small amount of base (OH-) is added to the buffer

CH3COOH CH3COO- + H+

A

H+ ions in the equilibrium react with the added OH- ions
H+ + OH- -> H2O
Ch3COOH dissociates more and the equilibrium shifts to the right t minimise the effect of a decrease in. [H+]
Most of the H+ one in the equilibrium are restored

76
Q

Exothermic reactions are usually spontaneous (feasible)

A

The system goes from higher enthalpy to lower enthalpy and so becomes energetically more stable

Some endothermic reactions are also spontaneous

77
Q

Entropy

A

A measure of the dispersal of energy in a system

The entropy is greater when the system is more disordered

78
Q

If delta S is positive

A

The system becomes more disordered during the reaction

79
Q

A system becomes energetically more stable when

A

It becomes more disordered

80
Q

Chemical reaction involving an increase in number of gas molecules

A

More disorder

Delta S is positive

81
Q

Standard entropy chnage of reaction

A

Entropy change that accompanies a reaction in the molar quantities expressed in a chemical equation under standard conditions, all reactants and products being in their standard states

82
Q

If a chnage makes a system more disordered

A

Delta S is positive

83
Q

If a change makes a system less disordered

A

Delta S is negative

84
Q

A process is spontaneous (feasible) if

A

A chemical system becomes more stable and its overall energy decreases

85
Q

How do endothermic reactions take place spontaneously

A 
Endothermic so delta H is positive 
Spontaneous so delta S must be positive
Spontaneous so delta G must be negative 
Temperature must be high enough so that TdeltaS term is more positive than deltaH term
So deltaH-TdeltaS<0
86
Q

Lattice enthalpy is

A

The enthalpy change when one mole of an ionic compound formed from its gaseous ions

87
Q

Lattice enthalpy facts

4

A

Is always an exothermic enthalpy change as an ionic bond is formed , deltaH i always negative
The more exothermic the value for lattice enthalpy, the stronger the attraction between the oppositely charged ions
Lattice enthalpies cannot be measured directly by experiment, because it is impossible to form one mole of an ionic lattice
Lattices enthalpy is determined from a born haber cycle and applying hess’ law

88
Q

The more exothermic the lattice energy;

A

Stronger the attraction between oppositely charged ions
Stronger the ionic bonds in the giant ionic lattice
Higher the melting point of the ionic compound

89
Q

Ionic size (radius)
Factors affecting the value of lattice enthalpy
1+4

A

Smaller ions attract more strongly than larger ions

Same number of shells so same shielding
But nuclear charge increases
Proton;electron ratio increases- the protons in the nucleus attract fewer electrons as ionic charge increases
Nuclear attraction on the electrons increases, pulling inner shells closer to the nucleus

90
Q

Ionic charge
Factors affecting the value of lattice enthalpy
MgO is stronger than Na2O

A

Ions of greater charge attract more strongly than ions of smaller charge

Ions in MgO are Mg2+ and O2-
Ions in Na2O are Na+ and O2-
Mg2+ ion h as a greater charge than the Na+ ion
And Mg2+ ion is smaller than Na+ ion
Stronger attraction between Mg2+ and O2- ions than between Na+ and O2- ions
Lattice enthalpy of MgO is more exothermic than Na2O

91
Q

Standard enthalpy change of formation

+equation

A

The enthalpy change when one mole of compound is formed from its elements in their standard states under standard conditions

Na(s) + 1/2Cl2(g) —> NaCl(s)

92
Q

Standard enthalpy of atomisation

+equation

A

The enthalpy change when one mole of gaseous atoms is formed from its elements in its standard stats under standard conditions

Na(s)—> Na(g)

1/2Cl2 (g) —> Cl(g)

93
Q

First ionisation energy

+equation

A

The enthalpy change when one electron is removed from each atom in one mole of gaseous atoms to from on mole of gaseous 1+ ions

Na(g) —> Na+ (g) + e-

94
Q

Second ionisation energy

+equations

A

The enthalpy change when one electron is removed from each ion in one mole of gaseous 1+ ions to from one mole of gaseous 2+ions

Ca+(g) —> Ca2+(g) + e-

95
Q

First electron affinity

+equation

A

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

Cl(g) + e- —> Cl-(g)

96
Q

Second electrons affinity

+equation

A

The enthalpy change when one electron is added too each ion in one mole of gaseous 1- ions to form one mole of gaseous 2- ions

O-(g). + e- —> O2-(g)

97
Q

Key features of a born-haber cycle

A

Elements in their standard state have zero enthalpy
All delta.H values. Pointing upwards are endothermic so deltaH positive
AlldeltaH values poing downwards are deltaH negative
The sum of the clockwise enthalpy changes= the sum of the anticlockwise enthalpy changes

98
Q

The standard enthalpy of solution

A

The enthalpy change that takes place when one mole of compound is completely dissolved in water under standard conditions

99
Q

Two processes that take place when an ionic compound dissolves in water

A

The ionic lattice dissociates into its gaseous ions

The gaseous ions are hydrated with H2O molecules

100
Q

Dissociation of the ionic lattice

A

Breaking down the ionic lattice is an endothermic process so deltaH positive
Because energy is needed to overcome the attraction between the oppositely charged ions
Dissociation of the ionic lattice is the opposite to lattice enthalpy

101
Q

Hydrating the gaseous ions

A

This process is exothermic, deltaH is negative because energy is released when the gaseous ion attract and bond with H2O molecules
The positive ions attract the slightly negative O atoms in H2O, the negative ions attract the slightly positive H atoms in H20

102
Q

Standard enthalpy of hydration

+equations

A

The enthalpy change that takes place when one mole of isolated gaseous ions is dissolved in water forming one mole of aqueous ions under standard conditions

Na+(g) + aq —> Na+ (aq)
Cl-(g) + aq —> Cl-(aq)

103
Q

Factors affecting the value of enthalpy change of hydration

A

Affected by ionic size and ionic charge
The stronger the attraction between an ion and H2O molecules, the more exothermic the enthalpy change of hydration
Smaller ions attract and bond with H2O molecules more strongly than larger ions
Higher charged ions attract and bond with H20 molecules more strngly than lower charged ions

104
Q

Enthalpy change of hydration up group 7

A

Ionic radii decrease

More exothermic

105
Q

Enthalpy change of hydration across a period

A

Ionic radii decrease
Ionic charge increases
More exothermic

106
Q

The oxidation number of an element is a measure of

A

The number of electrons that an atom uses to bond with atoms of another element.

107
Q

Manganate equation

A

MnO4-(aq)+ 8H+(aq) + 5e- —>

Mn2+(aq) + 4H2O(l)

108
Q

Thiosulfate equation

A

2S2O3(2-) (aq) + I2(aq) —>

2I- (aq) + S4O6(2-) (aq)

109
Q

Moldm^-3 to gdm^-3

A

X mr

110
Q

gdm^-3 to Moldm^-3

A

/mr

Chemistry A LEVEL (10 decks)

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