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OCR Chemistry A > 3.1 & 2 > Flashcards

3.1 & 2 Flashcards

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1
Q

How is the periodic table arranged

A

By increasing atomic no.
In periods showing repeating trends in physical and chemical properties
In groups having similar chemical properties

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2
Q

First ionisation energy

A

Removal of 1 mole of electrons from each atom in 1 mole of gaseous atoms

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3
Q

Metallic bonding

A

Strong electrostatic attraction between cations and delocalised electrons

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4
Q

Solid giant covalent lattices

A

Networks of atoms bonded by strong covalent bonds

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5
Q

Atomic radii across a period

A

Decreases
More attraction to nucleus
Greater no. of protons

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6
Q

When is the ionisation energy higher

A

With a smaller radius (higher nuclear charge)
Greater attraction
More energy needed

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7
Q

How do inner quantum shells affect the IE

A

Inner quantum shells shield the electrons from the attractive force of the nucleus
The more inner quantum shells
The greater the shielding effect
The lower the IE

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8
Q

When a molecule sublimates what bonds are broken

A

The intermolecular forces not the covalent bonds

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9
Q

Why is Mg harder than Na

A

No. of delocalised electrons
Charge on cations
Cations are smaller, so the charge density is even greater

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10
Q

Why is there an increase in mp and bp of metals from Gp 1 to Gp3

A 
Increased strength of metallic bonding: 
The charge on the positive ion is greater 
More delocalised ions (1 in 1+ vs 3 in 3+)
Smaller ions (greater charge density)
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11
Q

Properties of metals

A

Malleable
Ductile
High mp and bp
Good electrical conductors

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12
Q

Solubility of giant ionic lattices

A

Ionic lattices dissolve in polar solvents like water
Cations attract delta negative charges on the O molecule in H2O
Anions attract delta positive charges on the H molecules in H2O

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13
Q

Solubility of simple molecular lattices

A

London forces can form between molecules and non-polar solvents (hexane), weakening the lattice

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14
Q

Group 2 characteristics

A

Fairly high melting and boiling points

Low densities

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15
Q

Group 2 reactions

A 
2 Mg (s) + O2 (g) —> 2MgO (s)
Ca (s) + 2 H20 (l) —> Ca(OH)2 + H2
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16
Q

Second IE

A

X (g) —> X^+ (g) and e^-

X^+ (g) —> X^2+ and e^-

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17
Q

Why does reactivity increase down Group 2

A

Further from nucleus
More shielding
1st and 2nd IE decreases

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18
Q

Why does solubility and pH increase down a group

A

Mg(OH)2 —> Mg2+ and 2 OH-

Dissociation causes more hydroxide ions to be released —> more alkaline solution

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19
Q

Down Group 7

A

No. of full shells increase
Outermost electrons further from nucleus
Atomic radius increases

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20
Q

Forces in Group 7 molecules

A

London forces

More electrons = stronger London forces

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21
Q

What do halogens react to form

A

Halide ions

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22
Q

Are halogens reduced or oxidised

A

Reduced

Oxidising agent

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23
Q

Why does reactivity decrease down Group 7

A

Nuclear charge increases down a group
Atomic radius increases
Shielding increases (outweighs nuclear atraction)
Attraction between nucleus and electron decreases
Halogen less readily accepts an electron

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24
Q

Do halide ions get oxidised or reduced

A

Oxidised

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25
Q

Why is iodide the most powerful reducing agent

A

It wants to lose electrons

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26
Q

How does chlorine react with water

A

To form an acidic solution containing HCl and HOCl (chloric (I) acid)

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27
Q

Risks of chlorinating water

A

Cl2 gas is toxic

Chlorinated hydrocarbons may form (carcinogens)

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28
Q

Benefits of chlorinating water

A

Effectively kills many harmful microorganisms

Reduction in cholera and typhoid rates

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29
Q

Bleach

A

Formed when Cl2 is added to cold dilute NaOH solution
Active ingredient is sodium chlorate (NaOCl)

Cl2 (aq) + 2 NaOH (aq) —> NaOCl (aq) + Nacl (aq) + H2O (l)

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30
Q

Tests for halide ions

A

AgNO3 used to test

Silver halides are insoluble

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31
Q

Cl- (aq) and Ag+ (aq) —>

A

AgCl (s)

White precipitate

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32
Q

Br- (aq) and Ag+ (aq) —>

A

AgBr

Cream precipitate

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33
Q

I- (aq) and Ag+ (aq) —>

A

AgI

Yellow precipitate

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34
Q

What does AgCl dissolve in

A

Dilute ammonia

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35
Q

AgBr dissolves in

A

Concentrated ammonia

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36
Q

Does AgI dissolve

A

No

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37
Q

Colours with cyclohexane

A

Cl2 - pale green
Br2 - orange
I2 - violet

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38
Q

Order of identifying ions

A

Carbonate
Sultate
Halide

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39
Q

Carbonate test

A

Dilute nitric acid
Bubbles of gas, effervescence if positive
If you bubble CO2 through lime water it goes from clear to cloudy

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40
Q

Sulfate test

A

Add Ba^2+ (Ba(NO3)2)

Forms white precipitate if positive - barium sulfate

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41
Q

Halide test

A

Dissolve halide in H2O and add AgNO3
AgX precipitate colour if positive
Also test if dissolves in NH3

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42
Q

Ammonium test

A

Add warm NaOH and place on damp litmus paper

Turns blue

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43
Q

Why is the order of the ion tests important

A

BaCO3 (s) and Ag2SO4 (s) are both insoluble white precipitates but only BaCO3 would give off bubbles of gas so you’d automatically know

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44
Q

If you have a mixture of ions:

A

Do tests in same order
Keep adding HNO3 until reaction stops
Add excess of barium nitrate to ensure all SO4^2- form a precipitate - which can be filtered
Then test filtrate for halide ions

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45
Q

Why is dilute nitric acid used in the carbonate test

A

It does not form any precipitates whereas HCl and H2SO4 do

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46
Q

Direct measurement of enthalpy change

A 
Q = m * c * delta-t 
Delta-H = -Q/n

Use n of moles that is not in excess

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47
Q

Hess’ law

A

Enthalpy is a state function, depends only on state
If a reaction can take more than one route and the initial and final conditions are the same, the total enthalpy change is the same

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48
Q

Catalyst

A

A substance that increases the rate of a chemical reaction without being used up in the overall reaction.
Does this by providing an alternative route with a lower activation energy

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49
Q

Homogeneous catalysis

A

A reaction in which the catalyst and reactants are in the same physical state usually aq or g

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50
Q

Heterogeneous catalysis

A

A reaction in which the catalyst and the reactants are in diff. states. The reactant is usually a gas and the catalyst, a solid

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51
Q

Dynamic equilibrium

A

The equilibrium that exists in a closed system when the rate of the forward reaction is equal to the rate of the reverse reaction and the conc. don’t change

52
Q

Le Chateliers principle

A

When a system in dynamic equilibrium is subjected to a external change, the system readjusts itself to minimise the effect of the change and to restore equilibrium

53
Q

What is Ca(OH)2 used for

A

Neutralising acid soils

54
Q

What are Mg(OH)2 and CaCO3 used as

A

‘Antacids’ in treating indigestion

55
Q

Delta H

A

Reactants - products

56
Q

-ve delta h

A

Exothermic

57
Q

+ve delta h

A

Endothermic

58
Q

Activation energy

A

Minimum energy required to start a reaction

59
Q

Enthalpy change of formation

A

Formation of 1 mol of a compound from its constituent elements, in standard conditions
Products and reactants in their standard form

60
Q

Standard enthalpy of neutralisation

A

Formation of 1 mol of water from neutralisation

61
Q

Standard conditions

A

298K

100 kPa

62
Q

Collision theory

A

For a reaction to take place, particles must collide in the correct orientation with sufficient energy

63
Q

Why does using a catalyst increase the rate

A

It provides an alternative route with lower Ea so a greater proportion of collisions have an energy > Ea

64
Q

Increasing the temperature increases the rate because

A

A greater proportion of molecules have an energy > Ea and there are more frequent successful collisions

65
Q

Changing the pressure/conc. changes the rate because

A

There are more/less particles per unit volume and therefore there are more/less collisions

66
Q

Enthalpy change of formation

A

Formation of 1 mol of a substance from its constituent elements under standard conditions and states

67
Q

Enthalpy change of combustion

A

Complete combustion of 1 mole of a substance with excess O2 under standard conditions and states

68
Q

Enthalpy change of neutralisation

A

Formation of 1 mol of water from neutralisation under standard conditions and states

69
Q

How to measure rate of reaction from a graph

A

Drawing a tangent at said point and calculating the gradient

70
Q

Factors affecting rate of reaction

A

Temperature
Concentration
Surface area
With or without catalysts

71
Q

Effect of increasing surface area

A

All particles have exposed surfaces to react with each other; more frequent collisions

72
Q

What does the highest point of the Boltzmann curve shows

A

Most probable energy

73
Q

What is the total area under the Boltzmann distribution equal to

A

The no. of molecules

74
Q

If the forward reaction is exothermic, how does increased temp affect the equilibrium

A

Shifts to the LHS

75
Q

If the forward reaction is exothermic, how does decreased temp affect the equilibrium

A

Shifts it to RHS

76
Q

How does increasing pressure affect the equilibrium

A

Shifts to side with least moles of gas

77
Q

How does decreasing pressure affect equilibrium

A

Shifts to side with most moles of gas

78
Q

How does the equilibrium shift if the conc of reactants increases

A

Shifts to the RHS

79
Q

How does the equilibrium shift if the conc. of products increases

A

Shifts to the LHS

80
Q

If the forward reaction is endothermic, how does increased temp affect the equilibrium

A

Shifts to the RHS

81
Q

If the forward reaction is endothermic, how does decreased temp affect the equilibrium

A

Shifts to the LHS

82
Q

What is Kc a measure of

A

The conc. of products and reactants at equilibrium

83
Q

How is Kc calculated

A

Kc = [C]^c [D]^d/ [A]^a [B]^b

84
Q

If Kc&raquo_space; 1

A

More products are formed at equilibrium

85
Q

If Kc &laquo_space;1

A

More reactants formed at equilibrium

86
Q

If Kc = 1

A

Equal conc. of reactants and products

87
Q

Why do solid giant ionic structures not conduct electricity

A

They have ions fixed in position by ionic bonds and so don’t conduct

88
Q

Role of a catalyst

A

Increasing rate of reaction without being used up by the overall reaction

Providing alternative route with lower Ea

89
Q

Homogenous catalysts

A

Has same physical state as the reactants. Reacts with reactants to form intermediate. This then breaks down to give the product and regenerates the catalyst

90
Q

Heterogenous catalyst

A

Has diff. physical state to reactants. Reactant molecules are absorbed onto the surface of the catalyst where the reaction take place, activation energy is lowered and the product molecules leave by desorption

91
Q

Catalyst for Haber process

A

Solid Iron

92
Q

Catalyst for reforming

A

Pt (s) or Rh (s)

93
Q

Catalyst in hydrogenation of alkenes

A

Solid Ni

94
Q

Industrial preparation of ethanol

A

C6H12O6 —> 2 C5H5OH + 2 CO2

95
Q

Benefits of catalysts

A

Increased sustainability:

Lowering temps and reducing energy demand from combustion of fossil fuels with resulting reduction in CO2 emissions

96
Q

Thermal decomposition of Group 2 carbonate

A

CaCO3 —> CaO + CO2

97
Q

Periodicity

A

Repeating patterns of properties across diff. periods

98
Q

Why do solid ionic lattices not conduct electricity but molten ones do

A

The ions are fixed in place by ionic bonds when solid but when molten the ions are mobile

99
Q

Which Group 2 carbonate decomposes at the highest temp

A

BaCO3

100
Q

Are giant metallic structures soluble in water

A

No

101
Q

Bonding in gaseous hydrogen halides

A

Mainly covalent with an increasing tendency towards ionic as you go up the group

102
Q

Which particles are attracted in metallic bonding

A

Cations and delocalised electrons

103
Q

Which halogen most readily forms 1- ions

A

Fluorine

104
Q

Percentage uncertainty

A

No. of readings * uncertainty/ (quantity measured) * 100

105
Q

Limitation in enthalpy pag

A

Fuel evaporates
Heat loss to surroundings
Incomplete combustion.

106
Q

Sources of error in titrations

A

Not accurately weighing mass of solid - use more precise balance
Not all acid gets transferred to volumetric flask - rinse out beaker and add rinsings to volumetric flask
Insufficient mixing of solution - invert several times
Burette not rinsed - rinse with acid solution before use

107
Q

Which halogen most readily forms 1- ions

A

Fluorine

108
Q

NO and CO in a catalytic converters

A

2NO + 2CO —> 2CO2 + N2
CO and NO adsorbed onto surface
CO2 and N2 made, lowered activation energy
Products desorb from the surface

109
Q

General trend in mp and bp from Gp 1 to Gp 4

A

Increases

110
Q

Why is there an increase in mp and bp between Gp 3 - Gp 4

A

From Al to Si; giant metallic lattice to giant covalent lattice
Covalent bonds are stronger than metallic bonds in the same period

111
Q

Why is there a decrease in mp and bp between Gp 4 to Gp 7

A

Go from giant covalent lattice to simple molecular lattice w/ weak IMF (LF) and molecules get smaller e.g. P4 and Cl2

112
Q

Simple molecular lattice

A

Weak IMF (LF) between molecules

113
Q

The bigger the molecules …

A

The more electrons
Stronger induced dipole-dipole forces
Higher mp

114
Q

How does repulsion between electrons in orbitals affect the IE

A

Less repulsion leads to larger IE

115
Q

Why doesn’t a catalyst change the position of eqm

A

Increases the forward and backward reaction by the same amount

116
Q

Enthalpy change of reaction

A

No. of moles of reactants specified in the eqn react together

117
Q

Avg. bond enthalpy

A

One mole of bonds breaking

118
Q

Why may the actual bond enthalpy be diff to the calculated value

A

Bonds have diff strengths in diff environments

119
Q

Why do industrial manufacturing processes use catalysts

A 
Higher atom economy 
Reduce CO2 emission 
Enable reactions to occur w/ more specificity (correct stereoisomer) 
Reactions can occur at lower temp 
Saves energy costs
120
Q

What makes a reaction exothermic

A

Bond breaking absorbs energy
Bond forming releases energy
More energy released than absorbed

121
Q

Endothermic

A

Energy enters system from surrounding

122
Q

Energy profile diagram for endothermic reaction

A

Products above reactants

Ea and delta H going up to products

123
Q

How is a dynamic eqm reached

A

Rate of forward reaction slows down and rate of backwards reaction speeds up

124
Q

Ionic eqn of carbonate test

A

CO3 2- + 2H+ —> CO2 + H2O

125
Q

Directly measuring enthalpy change of solution/ neutralisation

A

Measure out sol using a gradated pippette into a polystyrene cup
Weigh out excess of solid
Place thermometer into cup; stir and record temp every 30s for X mins
At X+1 mins add solid to cup
Gently stir and record temp every 30s until 10mins
Draw cooling curve and extrapolate

126
Q

Why must we extrapolate cooling curves to get an accurate value for delta T

A

Reaction isn’t instantaneous

Heat is lost as mixture heats up

127
Q

Why may enthalpy change of neut/sol be diff from accepted value

A

No lid is used - heat loss
Spp heat capacity may be inaccurate
Energy absorbed by polystyrene

OCR Chemistry A (16 decks)

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