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Chemistry > PAPER 2 GCSE > Flashcards

PAPER 2 GCSE Flashcards

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1
Q

what is a mole (mol)

A

the unit for the amount of a substance

The mass of 1 mole of a substance is the relative formula mass (Mr) of the substance in grams.

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2
Q

mol calculation

A

mol = mass/Mr

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3
Q

what is yield

A

Yield is how much product you get from a chemical reaction.

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4
Q

what is theoretical yield

A

The theoretical yield is the amount of product that you would expect to get. This is calculated using reacting mass calculations.

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5
Q

how to calculate percentage yield

A

actual amount of a product / theoretical amount of a product

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6
Q

what is water of crystallisation

A

when some substances crystallise from solution, water becomes chemically bound up with the salt.

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7
Q

what is the empirical formula

A

The empirical formula shows the simplest whole-number ratio between atoms/ions in a compound.

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8
Q

what is the molecular formula

A

The molecular formula shows the actual number of atoms of each type of element in a molecule.

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9
Q

concentration formula (h)

A

measured in mol/dm3
concentration = mol / volume
1 dm3 = 1000 cm3

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10
Q

1 mole of gas, at room temperature and pressure (rtp), will always occupy (h)

A

24 dm3 or 24,000 cm3.

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11
Q

formula between a volume and a number of moles for a given gas (h)

A

mol = vol / 24

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12
Q

why covalent compounds do not conduct electricity (h)

A

because there are no charged particles that are free to move

charged particles means either delocalised electrons or ions.

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13
Q

why ionic compounds conduct electricity only when molten or in aqueous solution (h)

A

When solid the ions are not free to move.

When molten or in solution the ions are free to move.

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14
Q

A negative ion is called (h)

A

an anion. Examples are the bromide ion (Br⁻) and the oxide ion (O²⁻).

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15
Q

a positive ion is called (h)

A

a cation. Examples are the sodium ion (Na⁺) and the aluminium ion (Al³⁺).

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16
Q

electrolysis is (h)

A

The breaking down of a substance caused by passing an electric current through an ionic compound which is molten or in solution

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17
Q

electrolysis of molten ionic compounds (lead bromide) (h)

A
  1. Solid lead bromide is heated and becomes molten
  2. Electrodes attached to a power source are placed in the molten lead bromide (made of graphite or platinum - unreactive and conductive)
  3. delocalised electrons flow from the anode to the cathode.
  4. Negatively charged bromide ions are attracted to the anode (positive electrode). At the anode, bromide ions lose electrons (oxidation) and become bromine
  5. Positively charged lead ions are attracted to the cathode (negative electrode). At the cathode, lead ions gain electrons (reduction) and become lead atoms.
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18
Q

positively charged electrode (h)

A

anode

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19
Q

negatively charged electrode (h)

A

cathode

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20
Q

at the cathode in solution - rules (h)

A

Hydrogen and metal ions are positively charged
The metal will be produced if it is less reactive than hydrogen (copper, silver, and gold)
Otherwise hydrogen gas is produced

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21
Q

at the anode in solution - rules (h)

A

The product of electrolysis is always oxygen gas (O2) unless the solution contains a high concentration of Cl–, Br­- or I– ions, in which case a halogen gas is produced

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22
Q

electrolysis of ionic solutions (sodium chloride solution) (h)

A
  1. Solid sodium chloride is dissolved in water
  2. The solution also contains hydrogen ions (H+) and hydroxide ions (OH–) because water is a very weak electrolyte. It ionises very slightly to give hydrogen ions and hydroxide ions:
  3. Chloride ions (Cl–) and hydroxide ions (OH–) are attracted to the anode.
  4. Sodium ions (Na+) and hydrogen ions (H+) are attracted to the cathode.
  5. chloride ions lose electrons (oxidation) and form molecules of chlorine.
  6. hydrogen ions gain electrons (reduction) and form molecules of hydrogen. The hydrogen ions react at the cathode
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23
Q

chlorine gas colour (h)

A

green yellow gas

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24
Q

bromine gas colour (h)

A

brown gas

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25
Q

iodine gas colour (h)

A

purple gas

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26
Q

Oxidation is (h)

A

the loss of electrons or the gain of oxygen

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27
Q

Reduction is (h)

A

the gain of electrons or the loss of oxygen

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28
Q

practical for electrolysis of aqueous solutions (h)

A

electrolytic cell
The electrolyte is an aqueous solution. For example it might be concentrated sodium chloride, NaCl (aq).

The test tubes over the electrodes must not completely cover them to make sure the ions are free to move throughout the solution.

In the case of NaCl (aq) bubbles of gas will be seen forming at the electrodes. These float up and collect in the test tubes when each gas can be tested to assess its identity.

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29
Q

chlorine at room temperature

A

green gas

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30
Q

bromine at room temperature

A

red liquid

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31
Q

iodine at room temperature

A

grey solid

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32
Q

fluorine at room temperature

A

yellow gas

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33
Q

astatine at room temperature

A

black solid

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34
Q

As you go up group 7 (decreasing atomic number),

A

the elements become more reactive. For example, fluorine is the most reactive and astatine is the least reactive.

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35
Q

A more reactive halogen will

A

Displace a less reactive halogen
By reacting a halogen solution with a potassium halide solution and making observations, the order of their reactivity can be deduced

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36
Q

explain the trend in reactivity in Group 7 in terms of electronic configurations (h)

A

In fluorine the outer electron shell is very close to the positively charged nucleus, so the attraction between this nucleus and the negatively charged electrons is very strong. This means fluorine is very reactive.

However, for iodine the outer electron shell is much further from the nucleus so the attraction is weaker. This means iodine is less reactive.

how readily these elements form ions, by attracting a passing electron to fill the outer shell.

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37
Q

percentage of nitrogen

A

78

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38
Q

percentage of oxygen

A

21

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39
Q

percentage of argon

A

0.9

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40
Q

percentage of carbon dioxide

A

0.04

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41
Q

how to determine the percentage by volume of oxygen in air with iron

A

The iron reacts with the oxygen in the air (rusting).

As long as the iron and water are in excess, the total volume of air enclosed by the apparatus decreases by about a fifth (20%) over several days.

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42
Q

how to determine the percentage by volume of oxygen in air with phosphorus

A

The phosphorus is lit with a hot wire.

It reacts with the oxygen in the air and causes the water level in the bell jar to rise by about 20%.

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43
Q

Magnesium reacts with oxygen producing a

A

bright white flame leaving behind a white ash of magnesium oxide.

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44
Q

Hydrogen reacts with oxygen

A

in an explosive reaction. This is the basis of the ‘squeak pop’ test for hydrogen in test tube.

45
Q

Sulfur reacts with oxygen producing

A

a blue flame.

46
Q

thermal decomposition of metal carbonates

A

On heating metal carbonates thermal decompose into metal oxides and carbon dioxide.

47
Q

thermal decomposition of copper carbonate

A

green powder (CuCO3) changes to a black powder (CuO) and carbon dioxide

48
Q

carbon dioxide is a greenhouse gas which

A

It absorbs infra-red radiation and therefore warms the atmosphere. This leads to global warming.

This may cause climate change.

49
Q

most metals are extracted from (h)

A

ores found in the Earth’s crust and that unreactive metals are often found as the uncombined element

50
Q

If the ore contains a metal which is below carbon in the reactivity series then the metal is extracted (h)

A

by reaction with carbon in a displacement reaction.

51
Q

If the ore contains a metal which is above carbon in the reactivity series (h)

A

electrolysis (or reaction with a more reactive metal) is used to extract the metal

52
Q

where is carbon in the reactivity series (h)

A 
P
S
L
C
M
A
CARBON
Z
I
C
S
G
53
Q

uses of aluminium (h)

A

aircraft and power cables - low density, conductive

54
Q

uses of copper (h)

A

electrical wires and pots and pans - conductor of heat and electricity, ductile and malleable, unreactive

55
Q

uses of iron (h)

A

buildings and saucepans - strong, high melting points

56
Q

what is steel (h)

A

allow of carbon and iron

57
Q

uses of mild steel (h)

A

nails, car bodies

58
Q

uses of high carbon steel (h)

A

cutting tools, nails

59
Q

uses of stainless steel (h)

A

cutlery, cooking utensils

60
Q

why alloys are harder than pure metals (h)

A

the different elements have slightly different sized atoms. This breaks up the regular lattice arrangement and makes it more difficult for layers of ions to slide over each other.

61
Q

litmus paper indicator

A

blue - alkali

red - acid

62
Q

methyl orange indicator

A

yellow - alkali

red - acid

63
Q

phenolphthalein indicator

A

pink - alkali

colourless - acid

64
Q

Universal indicator is

A

a mixture of different dyes which change colour in a gradual way over a range of pH

65
Q

acid is

A

source of hydrogen ions (H+).

66
Q

alkali is

A

source of hydroxide ions (OH–).

67
Q

bases are

A

Metal oxides, metal hydroxides and ammonia (NH₃)

Bases neutralise acids by combining with the hydrogen ions in them.

68
Q

acid + base

A

salt + water

69
Q

describe how to carry out an acid-alkali titration (h)

A

Titration is used to find out precisely how much acid neutralises a certain volume of alkali

  1. measure out acid using a pipette
  2. transfer to conical flask
  3. fill burette with alkali and record initial reading
  4. add a few drops of phenolphthalein indicator to the acid
  5. add alkali until indicator changes colour
  6. take final reading and repeat
70
Q

test for hydrogen

A

Use a lit splint Gas pops

71
Q

test for oxygen

A

Use a glowing splint Glowing splint relights

72
Q

test for co2

A

Bubble the gas through limewater Limewater turns cloudy

73
Q

test for ammonia

A

Use red litmus paper Turns damp red litmus paper blue

74
Q

test for chlorine

A

Use moist litmus paper Turns moist litmus paper white (bleaches)

75
Q

describe how to carry out a flame test

A

A platinum or nichrome wire is dipped into concentrated hydrochloric acid to remove any impurities.

76
Q

lithium flame colour

A

red

77
Q

sodium flame colour

A

yellow

78
Q

potassium flame colour

A

lilac

79
Q

calcium 2 flame colour

A

orange - red

80
Q

copper 2 flame colour

A

blue - green

81
Q

tests for ammonium NH4

A

add sodium hydroxide and warm

if litmus paper turns blue, ammonia gas is made

82
Q

test for cu2

A

sodium hydroxide solution

blue precipitate forms

83
Q

test for fe2

A

sodium hydroxide solution

green precipitate forms

84
Q

test for fe3

A

sodium hydroxide

brown precipitate forms

85
Q

test for chloride ions

A

nitric acid to remove impurities
siver nitrate solution
white precipitate forms

86
Q

test for bromide ions

A

nitric acid to remove impurities
siver nitrate solution
cream precipitate forms

87
Q

test for iodide ions

A

nitric acid to remove impurities
siver nitrate solution
yellow precipitate forms

88
Q

test for sulfate ions

A

hydrochloric acid to remove impurities
barium chloride solution
white precipitate forms

89
Q

test for carbonate ions

A

add hydrochloric acid

fizzing will occur

90
Q

test for the presence of water

A

Add anhydrous copper (II) sulfate

white to blue

91
Q

test to show whether a sample of water is pure

A

If the sample is pure water it will boil at 100oC

92
Q

exothermic definition

A

chemical reaction in which heat energy is given out.

93
Q

endothermic definition

A

chemical reaction in which heat energy is taken in.

94
Q

calorimetry

A

the measurement of the amount of energy transferred in a chemical reaction to be calculated.

95
Q

molar enthalpy change (ΔH)

A

amount of energy transferred (Q) kj / number of moles

96
Q

how to calculate heat energy change (Q)

A

Q = mcΔT

mass x specific heat capacity x temp change

97
Q

In an exothermic reaction, the reactants have

A

more energy than the products

98
Q

In an endothermic reaction, the reactants have

A

less energy than the products

99
Q

bond making in an exothermic reaction (h)

A

the energy needed to break the bonds is less than the energy released in making new bonds.

100
Q

bond breaking in an endothermic reaction (h)

A

the energy needed to break the bonds is more than the energy released in making new bonds.

101
Q

Carboxylic acids contain the functional group (h)

A

COOH

  • C = O
    I
    O - H
102
Q

reactions of aqueous solutions of carboxylic acids with mg (h)

A

dilute ethanoic acid reacts with magnesium with a lot of fizzing to produce a salt and hydrogen, leaving a colourless solution of magnesium ethanoate:

103
Q

reactions of aqueous solutions of carboxylic acids with metal carbonates (h)

A

dilute ethanoic acid reacts with sodium carbonate with a lot of fizzing to produce a salt, carbon dioxide and water, leaving a colourless solution of sodium ethanoate

104
Q

vinegar is (h)

A

an aqueous solution containing ethanoic acid

105
Q

addition polymer is formed by

A

joining up many small molecules called monomers

One bond in the double bond breaks.
Monomers join together to form a long chain.
Polymer contains only single bonds.

106
Q

how to draw and name a polymer

A

poly(ethene)

break double bond, draw brackets and n

107
Q

problems in the disposal of addition polymers

A

Polymers are inert (unreactive) as they have strong C-C bonds.
This makes them non-biodegradeable.

The production of toxic gases when they are burned

108
Q

condensation polymerisation reaction

A

a dicarboxylic acid reacts with a diol and produces a polyester and water
Polyesters are polymers formed when two types of monomer

109
Q

some polyesters, known as biopolyesters, are

A

biodegradable

Chemistry (15 decks)

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