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Chemistry > PAPER 1 FULLY ASSESSED TOPICS > Flashcards

PAPER 1 FULLY ASSESSED TOPICS Flashcards

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1
Q

what three things affect the strength of forces of attraction between particles?

A
  • the material
  • the temperature
  • the pressure
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2
Q

SOLIDS

A
  • in solids, there are strong forces of attraction between particles, which hold them close together in fixed positions to form a very regular lattice arrangement
  • the particles don’t move from their positions, so all solids keep a definite shape and volume, and don’t flow like liquids
  • the particles vibrate about their positions - the hotter the solid becomes the more they vibrate (causing solids to vibrate slightly when heated)
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3
Q

LIQUIDS

A
  • in liquids, there is a weak force of attraction between the particles. they’re randomly arranged and free to move past each other, but they tend to stick closely together
  • liquids will have a definite volume but don’t keep a definite shape, and will flow to fit the bottom of a container
  • the particles are constantly moving with random motion. the hotter the liquid gets, the faster they move. this causes liquids to expand slightly when heated.
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4
Q

GASES

A
  • in gases, the forces of attraction between the particles is very weak - they’re free to move and are far apart. the particles in gases travel in straight lines.
  • gases don’t keep a definite shape or volume and will always fill any container
  • the particles move constantly with random motion. the hotter the gas gets, the faster they move. gases either expand when heated, or their pressure increases.
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5
Q

CHANGES OF STATE (solid - liquid - gas)

A

1) when a solid is heated, its particles gain more energy
2) this makes the particles vibrate more, which weakens the forces that hold the solid together. this makes the solid expand.
3) at a certain temperature, the particles have enough energy to break free from their positions. this is MELTING and the solid turns into a liquid.
4) when a liquid is heated, again the particles gain more energy.
5) this energy makes the particles move faster, which weakens and breaks the bonds holding the liquid together.
6) at a certain temperature, the particles have enough energy to break their bonds. this is EVAPORATING and the liquid turns into a gas

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6
Q

DIFFUSION EXPERIMENT: potassium manganate(VII) and water

A
  • take a beaker of water and place some purple potassium manganate(VII) at the bottom, the purple colour slowly spreads out to fill the beaker.
  • the particles of potassium manganate(VII) are diffusing among the particles of water
  • its the random motion of particles in a liquid that causes the purple colour to eventually be evenly spread out throughout the water.
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7
Q

DIFFUSION EXPERIMENT: ammonia and hydrogen chloride

A
  • aqueous ammonia (NH3) gives off ammonia gas. hydrochloric acid (HCl) gives off hydrogen chloride gas.
  • set up the experiment in a glass tube with cotton wool soaked with either solution at each end.
  • a white ring of ammonium chloride will form in the glass tube.
  • the NH3 gas diffuses from one end pf the tube and the HCl gas diffuses from the other. when they meet, they react to form ammonium chloride
  • the ring doesn’t form in the middle of the glass tube, it forms nearest the end of the HCl
  • this is because the particles of ammonia are smaller and lighter so they diffuse more quickly
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8
Q

DIFFUSION EXPERIMENT: bromine gas and air

A
  • bromine gas is a brown strongly smelling gas
  • fill half a gas jar full of bromine gas and the other half full of air - separate the gases with a glass plate
  • when you remove the glass plate, you’ll see the brown bromine gas slowly diffusing through the air
  • the random motion of the particles means that the bromine will eventually diffuse right through the air.
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9
Q

solution

A

a mixture of a solute and solvent that does not separate out.

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10
Q

solute

A

the substance being dissolved

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11
Q

solvent

A

the liquid it’s being dissolved into

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12
Q

saturated solution

A

a solution where the maximum amount of solute has been dissolved, so no more solute will dissolve in the solution

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13
Q

solubility

A

the ability of a substance to dissolve in a solvent. measured in grams of solute per 100 grams of solvent..

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14
Q

solubility curves

A
  • see the solubility of a substance at a specific temperature
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15
Q

SOLUBILITY EXPERIMENT: how the solubility of ammonium chloride is affected by temperature

A

1) make a saturated solution by adding excess ammonium chloride to 10cm3 of water in a boiling tube.
2) stir and place the boiling tube in a water bath at 25C
3) after 5 mins, check that all the excess solid has sunk to the bottom of the tube and use a thermometer to check the solution has reached 25C
4) weigh an empty evaporating basin. pour some pf the solution into the basin (no undissolved solid)
5) reweigh the basin and its contents, then gently heat it using a Bunsen burner to remove all the water
6) once all the water has evaporated reweigh the evaporating basin and it’s contents
7) repeat steps 1-6 twice more but with the water bath at different temperatures
8) use the different masses to work out the solubility at each temperature (and plot a graph)

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16
Q

atoms

A

made up of protons, neutrons and electrons

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17
Q

protons

A
  • heavy and positively charged
  • relative mass = 1
  • relative charge = +1
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18
Q

neutrons

A
  • heavy and neutral
  • relative mass = 1
  • relative charge = 0
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19
Q

electrons

A
  • hardly any mass and negatively charged
  • relative mass = 0.0005
  • relative charge = -1
  • move around the nucleus in energy levels called shells
  • the size of their orbitals cover a lot of space and determines the size of the atom
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20
Q

number of electrons equals number of protons

A

if electrons are added or removed, the atom becomes charged and is then an ion

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21
Q

atomic number

A

tells you how many protons there are

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22
Q

mass number

A

total number of protons and neutrons

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23
Q

molecules

A
  • groups of atoms

- held together by covalent bonds

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24
Q

isotopes

A

same atomic number, different mass number

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25
Q

relative atomic mass

A
  • takes all stable isotopes into account
  • average mass of all the isotopes of an element
  • you can find the relative atomic mass of any element using the periodic table
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26
Q

relative atomic mass CALCULATION METHOD

A
  • multiply the mass of each isotope by its relative abundance
  • add those together
  • divide by the sum of relative abundance

e. g. chlorine
(35. 0 x 3) + (37.0 x 1) / 3 + 1

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27
Q

what is COVALENT BONDING

A
  • formed between atoms by SHARING A PAIR OF ELECTRONS.
  • this way both atoms have a full outer shell, making them very stable
  • each covalent bond provides one extra shared electron for each atom
  • there’s a strong electrostatic attraction between the negatively charged shared electrons (the bonding pair) and the positively charged nuclei of the atoms involved,
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28
Q

describe covalent bonding in terms of electrostatic attractions

A

there’s a strong electrostatic attraction between the negatively charged shared electrons (the bonding pair) and the positively charged nuclei of the atoms involved,

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29
Q

COVALENT SUBSTANCES: simple molecular substances

A
  • the atoms WITHIN A MOLECULE are held together by very strong covalent bonds
  • however, the intermolecular forces of attraction BETWEEN the molecules are very weak
  • this results in very low melting and boiling points, because the molecules are easily separated.
  • the melting and boiling points of simple molecular substances increases as the relative molecular mass increases
  • most simple molecular substances are gases or liquids at room temperature, or a solid with low melting and boiling points
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30
Q

COVALENT SUBSTANCES: giant covalent structures

A
  • all the atoms are bonded to each other by strong covalent bonds
  • there are lots of these bonds meaning it takes a lot of energy to break, so giant covalent structures are solids with VERY HIGH MELTING AND BOILING POINTS
  • they don’t conduct electricity at all (except graphite)
  • usually insoluble in water
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31
Q

COVALENT SUBSTANCES: DIAMOND

A
  • made up of a network of carbon atoms that each form four covalent bonds
  • the strong covalent bonds take lots of energy to break, so diamond has a HIGH MELTING POINT
  • the strong covalent bonds hold the atoms in a very rigid lattice structure, so its really hard
  • it doesn’t conduct electricity because it has no free electrons or ions.
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32
Q

COVALENT SUBSTANCES: GRAPHITE

A
  • in graphite, each carbon atom only forms 3 covalent bonds, creating layers of carbon atoms
  • the layers are only held together weakly by intermolecular forces, so are free to slide over each other. this makes graphite soft and slippery
  • HIGH MELTING POINT - the covalent bonds IN the layers need lots of energy to break
  • only 3 out of carbons four outer electrons are used in bonds, so each carbon atom has one delocalised electron that’s free to carry charge, meaning it can conduct electricity.
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33
Q

COVALENT SUBSTANCES: C60 fullerene

A
  • hollow spheres made up of 60 carbon atoms
  • unlike diamond and graphite, C60 isn’t a giant covalent structure - its just made up of large covalent molecules
  • the C60 molecules are held together by intermolecular forces and so they can slide over each other = soft material
  • like graphite, each carbon in C60 has one delocalised electron. however, the electrons can’t move between the molecules so C60 fullerene is a poor conductor of energy
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34
Q

% of nitrogen in atmosphere

A

78%

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35
Q

% of oxygen in atmosphere

A

21%

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36
Q

% of argon in atmosphere

A

nearly 1%

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37
Q

% of carbon dioxide in atmosphere

A

0.04%

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38
Q

PROPORTION OF OXYGEN IN THE ATMOSPHERE EXPERIMENT: iron

A
  • iron reacts with oxygen in the air to form rust - so will remove oxygen from the air
    1) soak some iron wool in acetic acid (catalyse the reaction). then push the wool into a measuring cylinder and invert it into a beaker of water
    2) record the starting position of the water using the scale on the measuring cylinder - this is the starting volume of air
    3) over time, the level of the water in the measuring cylinder will rise
    4) this is because the iron reacts with the oxygen in the air to form iron oxide. the water rises to fill the space the oxygen took up
    5) leave the measuring cylinder for around a week, or until the water level stops changing
    6) record the finishing position of the water - this is the final volume of air.
    7) to calculate the % of oxygen: start vol- final vol/ start vol x100
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39
Q

PROPORTION OF OXYGEN IN THE ATMOSPHERE EXPERIMENT: phosphorus

A
  • place the phosphorus in a tube and attach a glass syringe at either end. make sure one of the syringes is filled with air and the other is empty
  • heat the phosphorus and use the syringes to pass the air over it - the phosphorus will react with oxygen in the air to make phosphorus oxide
  • as it reacts, the amount of air in the syringes will decrease
  • measure the starting and final volumes of air using the scale on one of the syringes
  • then calculate % of oxygen in air
40
Q

combustion of elements in oxygen: MAGNESIUM

A
  • burns with a bright white flame in air
  • forms a white powder that is magnesium oxide
  • magnesium oxide is slightly alkaline when dissolved in water
41
Q

combustion of elements in oxygen: HYDROGEN

A
  • burns very easily in air and can be explosive
  • has an orangey/yellow flame
  • the only product is water (as water vapour)
  • combustion of hydrogen often used as test for hydrogen (squeaky pop!)
42
Q

combustion of elements in oxygen: SULPHUR

A
  • burns with a pale blue flame and produces sulphur dioxide.

- sulphur dioxide is acidic when dissolved in water

43
Q

thermal decomposition of metal carbonate

A
  • produces CO2 and a metal oxide
  • copper (II) carbonate is a green powder that easily decomposes to form carbon dioxide and copper(II) oxide when you heat it
  • heat copper(II) oxide then collect the gas given off in a test tube. (as carbon dioxide is denser than air, it sinks to the bottom of the tube and can be collected)
  • test for CO2 by bubbling through limewater
44
Q

thermal decomposition

A

when a substance breaks down into simpler substances when heated

45
Q

what is carbon dioxide?

A
  • a greenhouse gas

- increasing CO2 is linked to climate change

46
Q

test for chlorine

A

chlorine bleaches damp blue litmus paper, turning it white. (it may turn red for a moment first though - thats because a solution of chlorine is acidic)

47
Q

test for oxygen

A

oxygen relights a glowing splint

48
Q

test for carbon dioxide

A

carbon dioxide turns limewater cloudy - just bubble the gas through a test tube of limewater.

49
Q

test for hydrogen

A

hydrogen makes a “squeaky pop” noise with a lighted splint. (the noise comes from the hydrogen burning with the oxygen in the air to form H2O)

50
Q

test for ammonia

A

ammonia turns damp red litmus paper blue. (it also has a very strong smell)

51
Q

how do you carry out a flame test?

A
  • clean a platinum wire loop by dipping it in dilute HCl then holding it in a flame. once you hold the loop in the flame and it burns without any colour
  • you can dip it in the sample you want to test and put it in the clear blue part of the bunsen flame (hottest part)
  • then say oooooh and aaahhhhh :)
52
Q

what colour does the cation lithium burn with?

A

RED

53
Q

what colour does the cation SODIUM burn with?

A

YELLOW

54
Q

what colour does the cation POTASSIUM burn with?

A

LILAC

55
Q

what colour does the cation CALCIUM burn with?

A

ORANGE-RED

56
Q

what colour does the cation COPPER burn with?

A

BLUE-GREEN

57
Q

how can you test for the cation NH4+?

A
  • first test for ammonia gas using damp red litmus paper. if theres ammonia present it will turn blue
  • then test whether a substance contains ammonium ions (NH4+).
  • add some sodium hydroxide to a solution of the mystery substance in a test tube. if theres ammonia gas given off (which smells like cat wee) this means there are ammonium ions in the substance
58
Q

how can you test for the cation Cu2+?

A
  • add a few drops of sodium hydroxide solution to a solution of your mystery compound in a test tube.
  • if you get a coloured insoluble hydroxide you can then tell which metal was in the compound
  • copper (II) should be BLUE precipitate
59
Q

how can you test for the cation Fe2+?

A
  • add a few drops of sodium hydroxide solution to a solution of your mystery compound in a test tube.
  • if you get a coloured insoluble hydroxide you can then tell which metal was in the compound
  • iron (II) should be SLUDGY GREEN
60
Q

how can you test for the cation Fe3+?

A
  • add a few drops of sodium hydroxide solution to a solution of your mystery compound in a test tube.
  • if you get a coloured insoluble hydroxide you can then tell which metal was in the compound
  • iron (III) should be REDDISH BROWN
61
Q

how can you test for carbonates?

A
  • add dilute hydrochloric acid to your test sample
  • if carbonates (CO32-) are present then carbon dioxide will be released
  • you can test for carbon dioxide using limewater
62
Q

how can you test for sulphates?

A
  • sulphate ions produce a white precipitate
  • add dilute HCl, followed by barium chloride solution BaCl2
  • a white precipitate of barium sulphate means the original compound was a sulphate
63
Q

how can you test for halides?

A
  • halides = Cl-, Br-, I-
  • add dilute nitric acid (HNO3) followed by silver nitrate solution.
  • a CHLORIDE ion gives a WHITE precipitate of silver chloride
  • a BROMIDE ion gives a CREAM precipitate of silver bromide
  • an IODIDE ion gives a YELLOW precipitate of silver iodide
64
Q

how can you test for the presence of water?

A
  • when copper(II) sulphate is bound to water it forms lovely blue crystals
  • if you heat the blue hydrated copper(II) sulphate crystals it drives the water off
  • this leaves a white anhydrous copper(II) sulphate copper sulphate powder, which doesn’t have any water bound to it.
  • if you add a couple of drops of water to the white powder you get the blue crystals back again
  • if you want to test for water, add water to anhydrous copper(II) sulphate and see if it turns blue.
  • this will tell you if water is present but not if its pure
65
Q

how can you show whether a sample of water is pure?

A
  • pure = one substance, which means SET physical properties like boiling and freezing point
  • water boils at 100 and freezes at 0. if it boils over a range of temps it is impure
66
Q

how do you measure the affect of SURFACE AREA on rates of reaction?

A
  • using marble chips and hydrochloric acid
  • measure the volume of gas produced using a gas syringe. take readings at regular time intervals and record results.
  • repeat the experiment with exactly the same volume and concentration of acid, and exactly the same mass of marble chips, but with the marble chips more crushed up
  • then repeat with powdered marble
67
Q

how do you measure the affect of CONCENTRATION on rates of reaction?

A

the reaction between marble chips is also good for measuring how changing the reactant concentration affects rate.
- this time repeat the experiment with exactly the same volume of acid, but using different concentrations of acid

68
Q

how do you measure the affect of TEMPERATURE on rates of reaction?

A
  • sodium thiosulphate and hydrochloric acid are both clear solutions. they react together to form a yellow precipitate of sulphur
  • measure the rate by timing how long it takes for a black mark to disappear under the cloudy sulphur
  • can be repeated at different temperatures
69
Q

how do you measure the affect of CATALYSTS on rates of reaction?

A
  • the decomposition of hydrogen peroxide is normally quite slow but manganese oxide catalyst speeds it up.
  • oxygen is given off so you can use the gas syringe to measure
  • set up the apparatus
  • add MnO2 powder to the H2O2
  • measure the volume of gas produced at regular time intervals
  • repeat with exactly the same volume and concentration of hydrogen peroxide but using a different catalyst (eg copper oxide) (amount must be kept the same)
70
Q

what are the 4 things that rate of reaction depends upon?

A
  1. temperature
  2. the concentration of a solution/the pressure of a gas
  3. surface area
  4. the presence of a catalyst
71
Q

how does increasing the temperature increase the rate of reaction?

A
  • when the temperature is increase, the particles move faster. if they move faster, they have more collisions
  • higher temperatures also increase the energy of the collisions, since the particles are moving faster. reactions only happen if the particles collide with enough energy
  • this means at higher temperatures there will be more SUCCESSFUL COLLISIONS
  • so increasing temperature increases rate.
72
Q

how does increasing the concentration (or pressure) increase the rate of reaction?

A
  • if a solution is made more concentrated it means there are more particles of reactant in the same volume. this makes collisions more likely, thus increasing reaction rate
  • in a gas, increasing the pressure means the particles are more crowded. this means the frequency of collisions between particles will increase, so rate will also increase
73
Q

how does increasing the surface area increase the rate of reaction?

A
  • if one reactant is a solid, breaking it into smaller pieces will increase its surface area to volume ratio (i.e. more of the solid will be exposed, compared to its overall volume)
  • the particles around it will have more area to work on, so the frequency of collisions will increase
  • this means that the rate of reaction is faster for solids with a larger SA:V ratio
74
Q

how does a catalyst increase the rate of reaction?

A
  • a catalyst is a substance which increases the rate of reaction without being chemically changed or used up in the reaction
  • you can’t use any catalyst for any reaction, they are quite specific
  • catalysts work by DECREASING THE ACTIVATION ENERGY needed for a reaction to occur
  • they do this by providing an alternative reaction pathway that has a lower activation energy
  • as a result more of the particles have at least the minimum amount of energy needed for a reaction to occur when the particles collide.
75
Q

what is particle collision theory?

A

The rate of a chemical reaction depends on:

1) the collision frequency of reacting particles (how often they collide). the more collisions there are, the faster the reaction is.
2) the energy transferred during a collision. particles have to collide with enough energy for the collision to be successful (form products)

76
Q

what is the activation energy and why is it important?

A
  • the minimum amount of energy that particles need to react
  • particles need this much energy to break the bonds in the reactants and start the reaction. the greater the activation energy, the more energy needed to start the reaction- this has to be supplied, e.g. by heating the reaction mixture.
77
Q

what are reaction profiles?

A

diagrams that show the relative energies of the reactants and products in a reaction, and how the energy changes over the course of a reaction.

78
Q

describe an exothermic reaction profile

A
  1. the products are at a lower energy than the reactants. the difference in height represents the energy given out (per mole) in the reaction, ΔH is NEGATIVE here.
  2. the initial rise in energy represents the energy needed to start the reaction. this is the ACTIVATION ENERGY.
  3. the activation energy is the minimum amount of energy the reactants need to collide with each other and react.
79
Q

describe an endothermic reaction profile

A
  1. the products are at a higher energy that the reactants.

2. the difference in height represents the energy taken in (per mole) during the reaction. ΔH is positive here.

80
Q

what is crude oil?

A

a mixture of hydrocarbons.

81
Q

how does the industrial process of fractional distillation separate crude oil into fractions?

A
  1. the oil is heated until ,post of it has turned to gas. the gases enter a fractionating column (and the liquid bitumen is drained off at the bottom.)
  2. in the column theres a temperature gradient (hot at bottom and gets gradually cooler). when the substances that make up crude oil reach a part of the column where the temperature is lower than their boiling point they condense.
  3. the longer hydrocarbons have high boilings. they condense and drain out of the column early on, when they’re near the bottom
  4. the shorter hydrocarbons have lower boiling points. they turn to liquid and drain out much later on, near the top of the column where its cooler.
  5. bubble caps in the fractionating column stops the separated liquids from running back down the column and remixing. you end up with the crude oil mixture separated out into different fractions. each fraction contains a mixture of hydrocarbons with similar boiling points. each fraction may contain saturated or unsaturated hydrocarbons.
82
Q

what is the use for REFINERY GASES?

A

used in domestic heating and cooking. boil at coolest temperature.

83
Q

what is the use for GASOLINE?

A

used as car fuel. second coolest boiling point.

84
Q

what is the use for KEROSENE?

A

aircraft fuel. third coolest boiling point.

85
Q

what is the use for diesel?

A

used as a fuel in some cars and larger vehicles e.g. trains. third hottest boiling point.

86
Q

what is the use for FUEL OIL?

A

fuel for large ships and also in some power stations. second hottest boiling point.

87
Q

what is the use for bitumen?

A

surface roads and roofs. hottest boiling point.

88
Q

what are hydrocarbons?

A

molecules made up of hydrogen and carbon only.

89
Q

what is the trend in colour, boiling point and viscosity of the main fractions?

A

As you go up the fractionating column, the hydrocarbons have:

  • lower boiling points
  • lower viscosity (they flow more easily)
  • higher flammability (they ignite more easily)

This means that in general hydrocarbons with small molecules make better fuels than hydrocarbons with large molecules.

90
Q

what is a fuel?

A
  • a substance that, when burned, release heat energy
91
Q

what are the possible products of complete combustion of hydrocarbons with oxygen?

A
  • when you burn hydrocarbons in plenty of oxygen, the only products are CARBON DIOXIDE and WATER. this is called complete combustion.
  • hydrogen + oxygen –> carbon dioxide + water.
92
Q

what are the possible products of incomplete combustion of hydrocarbons with oxygen?

A
  • if theres not enough oxygen around for complete combustion, you can get incomplete combustion. this can happen in some appliances, eg. boilers, that use carbon compounds as fuels.
  • the products of incomplete combustion contain less oxygen that carbon dioxide.
  • as well as carbon dioxide and water, incomplete combustion produces carbon monoxide CO, a toxic gas, and carbon in the form of soot.
  • binds to haemoglobin etc.
93
Q

how are nitrogen oxides formed?

A
  • nitrogen oxides are created when the temperature is high enough for the nitrogen and oxygen in the air to react.
  • this often happens is car engines.
  • nitrogen oxides include nitrogen monoxide and nitrogen dioxide.
94
Q

how is sulphur dioxide formed?

A
  • a lot of the fractions obtained from crude oil are burnt as fuels.
  • when they’re burnt, sulphur dioxide (and nitrogen oxides) may be produced.
  • the sulphur dioxide comes from sulphur impurities in hydrocarbon fuels.
95
Q

how does nitrogen oxides and sulphur dioxide cause acid rain?

A
  • when sulphur dioxide and nitrogen oxides mix with water vapour in clouds they form dilute sulphuric and nitric acid.
  • the rain that falls is called acid rain
  • acid rain causes lakes to become acidic and many plants and animals die as a result.
96
Q

what is cracking and why is it needed?

A
  • a form of thermal decomposition, which just means breaking down molecules into simpler molecules by heating them.
  • long chain hydrocarbons have high boiling points and are viscous. short chain have low and are thinner and paler.
  • demand for short chain hydrocarbons like octane, which is used in gasoline, is much higher that for longer chain hydrocarbons
  • so, to meet this demand, long chain hydrocarbons are spilt into more useful short chain molecules using cracking.
97
Q

what are the conditions for cracking and why?

A
  • catalyst, e.g. silica (SiO2) or alumina (Al2O3)
  • high high high temperatures like 600-700 C.
  • cracking with a catalyst is called catalytic cracking.

Chemistry (24 decks)

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