PAPER 2 GCSE Flashcards
(109 cards)
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.
2
Q
mol calculation
A
mol = mass/Mr
3
Q
what is yield
A
Yield is how much product you get from a chemical reaction.
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.
5
Q
how to calculate percentage yield
A
actual amount of a product / theoretical amount of a product
6
Q
what is water of crystallisation
A
when some substances crystallise from solution, water becomes chemically bound up with the salt.
7
Q
what is the empirical formula
A
The empirical formula shows the simplest whole-number ratio between atoms/ions in a compound.
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.
9
Q
concentration formula (h)
A
measured in mol/dm3
concentration = mol / volume
1 dm3 = 1000 cm3
10
Q
1 mole of gas, at room temperature and pressure (rtp), will always occupy (h)
A
24 dm3 or 24,000 cm3.
11
Q
formula between a volume and a number of moles for a given gas (h)
A
mol = vol / 24
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.
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.
14
Q
A negative ion is called (h)
A
an anion. Examples are the bromide ion (Br⁻) and the oxide ion (O²⁻).
15
Q
a positive ion is called (h)
A
a cation. Examples are the sodium ion (Na⁺) and the aluminium ion (Al³⁺).
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
17
Q
electrolysis of molten ionic compounds (lead bromide) (h)
A
- Solid lead bromide is heated and becomes molten
- Electrodes attached to a power source are placed in the molten lead bromide (made of graphite or platinum - unreactive and conductive)
- delocalised electrons flow from the anode to the cathode.
- Negatively charged bromide ions are attracted to the anode (positive electrode). At the anode, bromide ions lose electrons (oxidation) and become bromine
- Positively charged lead ions are attracted to the cathode (negative electrode). At the cathode, lead ions gain electrons (reduction) and become lead atoms.
18
Q
positively charged electrode (h)
A
anode
19
Q
negatively charged electrode (h)
A
cathode
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
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
22
Q
electrolysis of ionic solutions (sodium chloride solution) (h)
A
- Solid sodium chloride is dissolved in water
- 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:
- Chloride ions (Cl–) and hydroxide ions (OH–) are attracted to the anode.
- Sodium ions (Na+) and hydrogen ions (H+) are attracted to the cathode.
- chloride ions lose electrons (oxidation) and form molecules of chlorine.
- hydrogen ions gain electrons (reduction) and form molecules of hydrogen. The hydrogen ions react at the cathode
23
Q
chlorine gas colour (h)
A
green yellow gas
24
Q
bromine gas colour (h)
A
brown gas
25
iodine gas colour (h)
purple gas
26
Oxidation is (h)
the loss of electrons or the gain of oxygen
27
Reduction is (h)
the gain of electrons or the loss of oxygen
28
practical for electrolysis of aqueous solutions (h)
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.
29
chlorine at room temperature
green gas
30
bromine at room temperature
red liquid
31
iodine at room temperature
grey solid
32
fluorine at room temperature
yellow gas
33
astatine at room temperature
black solid
34
As you go up group 7 (decreasing atomic number),
the elements become more reactive. For example, fluorine is the most reactive and astatine is the least reactive.
35
A more reactive halogen will
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
36
explain the trend in reactivity in Group 7 in terms of electronic configurations (h)
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.
37
percentage of nitrogen
78
38
percentage of oxygen
21
39
percentage of argon
0.9
40
percentage of carbon dioxide
0.04
41
how to determine the percentage by volume of oxygen in air with iron
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.
42
how to determine the percentage by volume of oxygen in air with phosphorus
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%.
43
Magnesium reacts with oxygen producing a
bright white flame leaving behind a white ash of magnesium oxide.
44
Hydrogen reacts with oxygen
in an explosive reaction. This is the basis of the ‘squeak pop’ test for hydrogen in test tube.
45
Sulfur reacts with oxygen producing
a blue flame.
46
thermal decomposition of metal carbonates
On heating metal carbonates thermal decompose into metal oxides and carbon dioxide.
47
thermal decomposition of copper carbonate
green powder (CuCO3) changes to a black powder (CuO) and carbon dioxide
48
carbon dioxide is a greenhouse gas which
It absorbs infra-red radiation and therefore warms the atmosphere. This leads to global warming.
This may cause climate change.
49
most metals are extracted from (h)
ores found in the Earth’s crust and that unreactive metals are often found as the uncombined element
50
If the ore contains a metal which is below carbon in the reactivity series then the metal is extracted (h)
by reaction with carbon in a displacement reaction.
51
If the ore contains a metal which is above carbon in the reactivity series (h)
electrolysis (or reaction with a more reactive metal) is used to extract the metal
52
where is carbon in the reactivity series (h)
```
P
S
L
C
M
A
CARBON
Z
I
C
S
G
```
53
uses of aluminium (h)
aircraft and power cables - low density, conductive
54
uses of copper (h)
electrical wires and pots and pans - conductor of heat and electricity, ductile and malleable, unreactive
55
uses of iron (h)
buildings and saucepans - strong, high melting points
56
what is steel (h)
allow of carbon and iron
57
uses of mild steel (h)
nails, car bodies
58
uses of high carbon steel (h)
cutting tools, nails
59
uses of stainless steel (h)
cutlery, cooking utensils
60
why alloys are harder than pure metals (h)
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
litmus paper indicator
blue - alkali
| red - acid
62
methyl orange indicator
yellow - alkali
| red - acid
63
phenolphthalein indicator
pink - alkali
| colourless - acid
64
Universal indicator is
a mixture of different dyes which change colour in a gradual way over a range of pH
65
acid is
source of hydrogen ions (H+).
66
alkali is
source of hydroxide ions (OH–).
67
bases are
Metal oxides, metal hydroxides and ammonia (NH₃)
| Bases neutralise acids by combining with the hydrogen ions in them.
68
acid + base
salt + water
69
describe how to carry out an acid-alkali titration (h)
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
test for hydrogen
Use a lit splint Gas pops
71
test for oxygen
Use a glowing splint Glowing splint relights
72
test for co2
Bubble the gas through limewater Limewater turns cloudy
73
test for ammonia
Use red litmus paper Turns damp red litmus paper blue
74
test for chlorine
Use moist litmus paper Turns moist litmus paper white (bleaches)
75
describe how to carry out a flame test
A platinum or nichrome wire is dipped into concentrated hydrochloric acid to remove any impurities.
76
lithium flame colour
red
77
sodium flame colour
yellow
78
potassium flame colour
lilac
79
calcium 2 flame colour
orange - red
80
copper 2 flame colour
blue - green
81
tests for ammonium NH4
add sodium hydroxide and warm
| if litmus paper turns blue, ammonia gas is made
82
test for cu2
sodium hydroxide solution
| blue precipitate forms
83
test for fe2
sodium hydroxide solution
| green precipitate forms
84
test for fe3
sodium hydroxide
| brown precipitate forms
85
test for chloride ions
nitric acid to remove impurities
siver nitrate solution
white precipitate forms
86
test for bromide ions
nitric acid to remove impurities
siver nitrate solution
cream precipitate forms
87
test for iodide ions
nitric acid to remove impurities
siver nitrate solution
yellow precipitate forms
88
test for sulfate ions
hydrochloric acid to remove impurities
barium chloride solution
white precipitate forms
89
test for carbonate ions
add hydrochloric acid
| fizzing will occur
90
test for the presence of water
Add anhydrous copper (II) sulfate
| white to blue
91
test to show whether a sample of water is pure
If the sample is pure water it will boil at 100oC
92
exothermic definition
chemical reaction in which heat energy is given out.
93
endothermic definition
chemical reaction in which heat energy is taken in.
94
calorimetry
the measurement of the amount of energy transferred in a chemical reaction to be calculated.
95
molar enthalpy change (ΔH)
amount of energy transferred (Q) kj / number of moles
96
how to calculate heat energy change (Q)
Q = mcΔT
| mass x specific heat capacity x temp change
97
In an exothermic reaction, the reactants have
more energy than the products
98
In an endothermic reaction, the reactants have
less energy than the products
99
bond making in an exothermic reaction (h)
the energy needed to break the bonds is less than the energy released in making new bonds.
100
bond breaking in an endothermic reaction (h)
the energy needed to break the bonds is more than the energy released in making new bonds.
101
Carboxylic acids contain the functional group (h)
COOH
- C = O
I
O - H
102
reactions of aqueous solutions of carboxylic acids with mg (h)
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
reactions of aqueous solutions of carboxylic acids with metal carbonates (h)
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
vinegar is (h)
an aqueous solution containing ethanoic acid
105
addition polymer is formed by
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
how to draw and name a polymer
poly(ethene)
| break double bond, draw brackets and n
107
problems in the disposal of addition polymers
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
condensation polymerisation reaction
a dicarboxylic acid reacts with a diol and produces a polyester and water
Polyesters are polymers formed when two types of monomer
109
some polyesters, known as biopolyesters, are
biodegradable
