Physical Flashcards
1
Q
Enthalpy change
A
The amount of heat released or absorbed by a chemical reaction, carried out at constant pressure
2
Q
Exothermic reaction
Delta H=
A
A reaction where heat energy is released to the surroundings
Negative
3
Q
Endothermic reaction
Delta H =
A
A reaction where heat energy is absorbed from The surroundings
Positive
4
Q
Average bond enthalpy
A
The average enthalpy change for the breaking of one mole of bonds in gaseous molecules
5
Q
Standard conditions
A
25 degrees C / 298k
1 atmosphere/ 100 kPa
Solution must have conc 1 moldm^-3
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
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
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
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
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
11
Q
First law of thermodynamics
A
Energy can be converted from one form to another and cannot be created or destroyed
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
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
14
Q
When measuring temperature what are you recording the temperature of
A
Surroundings and NOT the system
15
Q
Enthalpy change of a reaction calculation
A
Enthalpy of products - enthalpy of reactants
16
Q
A molecular chemical reaction involves
A
Breaking covalent bonds in the reactant molecules and forming new covalent bonds in the product molecules
17
Q
Breaking bonds is
A
Endothermic
Absorbs energy
Positive
18
Q
Making/ forming bonds is
A
Exothermic
Releases energy
Negative
19
Q
The enthalpy change of a reaction depends only on
A
the initial and final states and is independent of the route taken
20
Q
The rate of a chemical reaction is
A
The change in concentration of a reactant or a product per unit time
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
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
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
24
Q
Activation energy
A
The minimum energy required for a reaction to take place, by the breaking of bonds in the reactants
25
Catalyst
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
26
Heterogenous catalysis
Catalysis of a reaction in which the catalyst has a different physical state from the reactants
27
Homogenous catalysis
Catalysis of a reaction in which the catalyst and the reactants are in the same physical state
28
Key advantages of using a catalyst in an industrial reaction
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
29
Effect of temperature
| Boltzman
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
30
Effect of a catalyst
| Boltzman
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
31
Characteristics of a dynamic equilibrium
| 4
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
32
Le Chateliers principle
When a system in dynamic equilibrium is subjected to a change, the equilibrium position will shift to minimise the effects of the change
33
Effect of temperature on dynamic equilibrium
Exothermic
Increase= left
Decrease = right
Endothermic
Increase= right
Decrease=left
34
Significance of Kc
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
35
The effect of temperature on Kc/p
| Exothermic forward reaction
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
36
The effect of temperature on Kc/p
| Endothermic forward reaction
- 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
37
The effect of concentration Kc
| 6
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
38
The effect of pressure KP
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
39
Efffect of a catalyst on Kc and Kp
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
40
The rate of a chemical reaction is
The change in concentration of a reactant or product per unit time
41
Rate equation
Rate = k [A]^m [B]^n
42
Rate at time ‘t’
Gradient of the tangent to the curve at time ‘t’
43
Initial rate of reaction
Is equal to the gradient of the tangent at time = 0
44
The half life of a reactant
The time taken for the concentration f the reactant to fall to half of its original concentration
45
The rate constant k for a first order reaction can be found by usin g the expression
K=ln2/t1/2
46
The rate determining step
The slowest step of a reaction mechanism of a multi-step reaction
47
Arrhenius equation
K= Ae^-Ea/RT
48
An acid is
A proton donor
49
A base is
A proton acceptor
50
Monobasic, dibasic and tribasic acid
```
Mono= requires one mole of OH- ions to neutralise it
Di= requires two moles
Tri= requires three moles
```
51
A conjugate acid base pair is
A pair of two species that transform into each other by the gain or loss of a proton
52
Significance of Ka
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
Ka equation
[H+][A-] / [HA]
54
pKa
-log Ka
55
Ka calculation
10^-pKa
56
Ka values only depend on
Temperature
57
Kw calculation=
[H+][OH-]
58
Kw value
1.00x10^-14
59
PH
-log [H+]
60
[H+]
10^-PH
61
Importance of Kw
Acidic solution [H+]>[OH-]
Neutral solution [H+] =[OH-]
Basic solutions [H+] < [OH-]
62
Equivalence point=
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
Strong acid and strong base
Vertical section
| PH 3-11
64
Strong acid and weak base
PH 3-7
65
Weak acid and strong base
PH 7-11
66
Weak acid and weak base
No vertical section
67
Indicators change colour with PH
Addition of acid
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
Indicators change colour with PH
Addition of alkali
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
The end point of an indicator
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
Choosing an indicator
The PH range/ end- point of the indicator must coincide with the vertical section of the PH curve
71
A buffer solution is
A mixture that minimise PH changes on addition of small amounts of an acid or base
72
A buffer solution is a mixture of
Weak acid
| Conjugate base of the weak acid
73
Two ways in which a buffer solution can be prepared
Mix a weak acid an one of its salts together
React excess weak acid with a strong alkali
74
When a small amount of acid (H+) is added to the buffer
CH3COOH CH3COO- + H+
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
When a small amount of base (OH-) is added to the buffer
CH3COOH CH3COO- + H+
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
Exothermic reactions are usually spontaneous (feasible)
The system goes from higher enthalpy to lower enthalpy and so becomes energetically more stable
Some endothermic reactions are also spontaneous
77
Entropy
A measure of the dispersal of energy in a system
| The entropy is greater when the system is more disordered
78
If delta S is positive
The system becomes more disordered during the reaction
79
A system becomes energetically more stable when
It becomes more disordered
80
Chemical reaction involving an increase in number of gas molecules
More disorder
Delta S is positive
81
Standard entropy chnage of reaction
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
If a chnage makes a system more disordered
Delta S is positive
83
If a change makes a system less disordered
Delta S is negative
84
A process is spontaneous (feasible) if
A chemical system becomes more stable and its overall energy decreases
85
How do endothermic reactions take place spontaneously
```
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
Lattice enthalpy is
The enthalpy change when one mole of an ionic compound formed from its gaseous ions
87
Lattice enthalpy facts
| 4
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
The more exothermic the lattice energy;
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
Ionic size (radius)
Factors affecting the value of lattice enthalpy
1+4
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
Ionic charge
Factors affecting the value of lattice enthalpy
MgO is stronger than Na2O
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
Standard enthalpy change of formation
+equation
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
Standard enthalpy of atomisation
+equation
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
First ionisation energy
+equation
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
Second ionisation energy
+equations
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
First electron affinity
+equation
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
Second electrons affinity
+equation
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
Key features of a born-haber cycle
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
The standard enthalpy of solution
The enthalpy change that takes place when one mole of compound is completely dissolved in water under standard conditions
99
Two processes that take place when an ionic compound dissolves in water
The ionic lattice dissociates into its gaseous ions
The gaseous ions are hydrated with H2O molecules
100
Dissociation of the ionic lattice
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
Hydrating the gaseous ions
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
Standard enthalpy of hydration
+equations
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
Factors affecting the value of enthalpy change of hydration
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
Enthalpy change of hydration up group 7
Ionic radii decrease
| More exothermic
105
Enthalpy change of hydration across a period
Ionic radii decrease
Ionic charge increases
More exothermic
106
The oxidation number of an element is a measure of
The number of electrons that an atom uses to bond with atoms of another element.
107
Manganate equation
MnO4-(aq)+ 8H+(aq) + 5e- —>
Mn2+(aq) + 4H2O(l)
108
Thiosulfate equation
2S2O3(2-) (aq) + I2(aq) —>
2I- (aq) + S4O6(2-) (aq)
109
Moldm^-3 to gdm^-3
X mr
110
gdm^-3 to Moldm^-3
/mr
