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Edexcel GCSE Chemistry Higher > RAG - Red > Flashcards

RAG - Red Flashcards

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1
Q

Acid + Metal >

A

Salt + Hydrogen

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2
Q

Acid + Metal Oxide >

A

Salt + Water

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3
Q

Acid + Carbonate >

A

Salt + Water + Carbon Dioxide

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4
Q

Acid + Metal Hydroxide >

A

Salt + Water

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5
Q

Hydrochloric Acid + Magnesium >

A

Magnesium Chloride + Hydrogen (salt + hydrogen)

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6
Q

Sulfuric Acid + Copper Oxide >

A

Copper Sulfate + Water (salt + water)

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7
Q

Hydrochloric Acid + Copper Carbonate >

A

Copper Chloride + Water + Carbon Dioxide (salt + water + carbon dioxide)

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8
Q

Nitric Acid + Sodium Hydroxide >

A

Sodium Nitrate + Water (salt + water)

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9
Q

What is the method for a neutralisation reaction?

A
  • Use a measuring cylinder to add dilute hydrochloric acid to a beaker.
  • Dip a clean glass rod into the contents of the beaker. Use it to transfer a drop of liquid to a piece of universal indicator paper on a white tile. Wait for 30 seconds then match the colour to a pH colour chart and record the pH.
  • Add a level spatula of calcium hydroxide powder to the beaker.
  • Stir thoroughly, then estimate and record the pH.
  • Repeat this until there are no more changes in pH.
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10
Q

Method for crystallisation:

A
  • Place the sulphuric acid in a conical flask and warm it in a warm water bath.
  • Add a spatula of copper oxide powder to the acid and stir with a glass rod.
  • Continue adding the copper oxide powder until it is in excess.
  • Filter the mixture to remove the excess copper oxide.
  • Pour the filtrate into an evaporating basin.
  • Heat the copper Sulfate solution to evaporate half the water.
  • Leave the solution by a window to allow all the water to evaporate.
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11
Q

Method for an acid-Alkali titration:

A
  • Use the pipette and pipette filler to add a measured volume of alkali to a clean conical flask.
  • Add a few drops of indicator and put the conical flask on a white tile.
  • Fill the burette with acid and note the starting volume.
  • Slowly add the acid from the burette to the alkali in the conical flask, swirling to mix.
  • Stop adding the acid when the
    end-point is reached (when the indicator first permanently changes colour). Note the final volume reading.
  • Repeat steps 1 to 5 until you get
    concordant titres.
  • More accurate results are obtained if acid is added drop by drop near to the end-point.
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12
Q

Purpose of using a pipette in a titration:

A
  • to accurately measure the volume of a reactant before transferring it to a conical flask.
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13
Q

Purpose of using a burette in a titration:

A
  • to add small, measured volumes of one reactant to the other reactant.
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14
Q

Naming salts - hydrochloric acid

A

Chloride salts

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15
Q

Naming salts - nitric acid

A

Nitrate salts

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16
Q

Naming salts - sulfuric acid

A
  • Produces Sulfate salts
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17
Q

To make a soluble salt from an acid and an insoluble reactant:

A
  • Add some dilute hydrochloric acid to a beaker.
  • Add powdered insoluble reactant to some acid in a beaker, one spatula at a time, stirring to mix. The mixture will effervesce. Continue adding powder until some unreacted powder is left over - it is in excess.
  • Filter the mixture in the beaker to remove the excess powder.
  • Allow the water in the solution to evaporate (by heating and/or leaving for a few days) to obtain pure dry crystals of the salt.
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18
Q

Dilute:

A
  • contains a relatively small amount of dissolved solute
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19
Q

Concentrated solution:

A
  • contains a relatively large amount of dissolved solute
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20
Q

Strong acids:

A
  • Strong acids completely dissociate into ions in solution.
  • For example, hydrochloric acid is a strong acid.
  • It completely dissociates to form hydrogen ions and chloride ions.
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21
Q

Weak acids:

A
  • Weak acids only partially dissociate into ions in solution.
  • For example, ethanoic acid is a weak acid.
  • It only partially dissociates to form hydrogen ions and ethanoate ions.
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22
Q

Electrolytes:

A
  • Ionic compounds that are:
  • in the molten state (heated so they become liquids).
    Or
  • dissolved in water.
  • Under these conditions, the
    ions in electrolytes are free to move within the liquid or solution.
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23
Q

Electrolysis:

A
  • a process in which electrical energy, from a direct current (dc) supply, decomposes electrolytes.
  • The free moving ions in electrolytes are attracted to the oppositely charged electrodes, which connect to the dc supply.
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24
Q

Cations:

A
  • Positively charged ions are called
    cations.
  • They move towards the negatively charged electrode, which is called the cathode.
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25
Q

Anions:

A
  • Negatively charged ions are called
    anions.
  • They move towards the positively charged electrode, which is called the anode.
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26
Q

Products of electrolysis - cations

A
  • gain electrons from the negatively charged cathode
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27
Q

Products of electrolysis - anions:

A
  • lose electrons at the positively charged anode
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28
Q

Molten lead bromide, PbBr2(l), is an electrolyte. During electrolysis:

A
  • Pb2+ ions gain electrons at the cathode and become Pb atoms.
  • Br- ions lose electrons at the anode and become Br atoms, which pair up to form Br2 molecules.
  • So lead forms at the negative electrode and bromine forms at the positive electrode.
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29
Q

Why can pure water conduct?

A
  • because a small proportion of its
    molecules dissociate into ions.
  • The two ions formed in water are, hydrogen ions, H+, and hydroxide ions, OH-.
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30
Q

During the electrolysis of water:

A
  • H+ ions are attracted to the
    cathode, gain electrons and form hydrogen gas.
  • OH- ions are attracted to the anode, lose electrons and form oxygen gas.
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31
Q

Hydrogen production at the cathode:

A
  • the metal is produced at the cathode if it is less
    reactive than hydrogen.
  • hydrogen is produced at the cathode if the metal is more reactive than hydrogen.
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32
Q

Production at the anode:

A
  • Either oxygen or a non-metal from the electrolyte can be produced at the anode:
  • for the most common compounds oxygen is produced (from the hydroxide ions)
  • if halide ions (chloride, bromide or iodide ions) are present, then the negatively charged halide ions lose electrons to form the corresponding non-metal halogen (chlorine, bromine or iodine).
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33
Q

Investigate electrolysis of copper Sulfate solution using inert electrodes:

A
  • Pour some copper sulfate solution into a beaker.
  • Place two graphite rods into the copper sulfate solution. Attach one electrode to the negative terminal of a dc supply, and the other electrode to the positive terminal.
  • Completely fill two small test tubes with copper sulfate solution and position a test tube over each electrode as shown in the diagram.
  • Turn on the power supply and observe what happens at each electrode.
  • Test any gas produced with a glowing splint and a burning splint.
  • Record your observations and the results of your tests.
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34
Q

Investigate electrolysis of copper Sulfate solution using copper electrodes:

A
  • Pour some copper sulfate solution into a beaker.
  • Measure and record the mass of a piece of copper foil. Attach it to the negative terminal of a dc supply, and dip the copper foil into the copper sulfate solution.
  • Repeat step 2 with another piece of copper foil, but this time attach it to the positive terminal.
  • Make sure the electrodes do not touch each other, then turn on the power supply. Adjust the power supply to achieve a constant current as directed by your teacher.
  • After 20 minutes, turn off the dc supply.
  • Carefully remove one of the electrodes. Gently wash it with distilled water, then dip it into propanone. Lift the electrode out and allow all the liquid to evaporate. Do not wipe the electrodes clean. Measure and record the mass of the electrode.
  • Repeat step 6 with the other electrode. Make sure you know which is which.
  • Repeat the experiment with fresh electrodes and different currents.
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35
Q

What happens during the process of purification of copper by electrolysis:

A
  • Four Cu ions are attached to the rod on the right, and four Cu²+ ions are floating in the space between the rods.
  • A battery is connected between the rods and the Cu ions are pulled towards the left rod.
  • There are now four Cu ions attached to the left rod, with four Cu²+ ions floating in the middle.
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36
Q

Purifying copper by electrolysis - how does pure copper form on the cathode?

A
  • A beaker with pure and impure copper rods dipped into copper sulfate solution.
  • The pure copper rod is connected to the negative terminal of a battery and the impure rod is connected to the positive terminal.
  • The pure copper rod has increased in size while the impure rod has deteriorated, leaving a pool of anode sludge at the bottom of the beaker.
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37
Q

Oxidation is:

A

The loss of electrons.

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38
Q

Where does oxidation happen?

A

The anode

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39
Q

What is reduction?

A
  • The gain of electrons
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40
Q

Where does reduction happen?

A

The cathode

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41
Q

the more reactive a metal is:

A
  • the more vigorous its reactions are
  • the more easily it loses electrons in reactions to form positive ions
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42
Q

Metal + water >

A

Metal Hydroxide + Hydrogen

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43
Q

Metal + Acid >

A

Salt + Hydrogen

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44
Q

A metal’s relative tendency to form cations and its resistance to oxidation are both related to its position in the reactivity series. In general:

A
  • the higher up a metal, the greater the tendency to form cations
  • the lower down a metal, the greater its resistance to oxidation
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45
Q

Redox Reaction

A

When oxidation and reduction take place at the same time.

46
Q

Where are unreactive metals found?

A

In the Earth’s crust as uncombined elements.

47
Q

What is an ore?

A

a rock that contains enough of a metal or a metal compound to make extracting the metal worthwhile.

48
Q

Low grade ores:

A
  • contain a small percentage of the metal or its compound
49
Q

High-grade ores

A
  • Contain a larger percentage of the metal or its compound.
50
Q

How to extract a metal less reactive than carbon:

A
  • Heat it with carbon
51
Q

Metal oxide + carbon >

A

Metal + Carbon Dioxide

52
Q

How to extract metals more reactive than carbon?

A

Electrolysis of molten compounds

53
Q

Steps of recycling metals

A
  • collecting and transporting the used items to a recycling centre
  • breaking up and sorting the different metals
  • removing impurities from the metals
  • The metals can then be used to manufacture new metal items.
54
Q

Advantages of recycling metals:

A
  • more economic - less
    energy is needed to produce a metal
  • less damage to the environment - fewer quarries and mines, less noise and less heavy traffic.
  • saves valuable raw materials - reserves of metal ores will last longer
55
Q

Disadvantages of recycling:

A
  • the collection and transport of used items needs organisation, workers, vehicles and fuel
  • it can be difficult to sort different materials from one another
56
Q

Stages of Life-Cycle Assessment:

A
  • obtaining the raw materials needed
  • manufacturing the product
  • using the product
  • disposing of the product at the end of its useful life
57
Q

What is a LCA?

A
  • A ‘cradle to grave’ analysis of the impact of a manufactured product on the environment.
58
Q

Impact of obtaining raw materials:

A
  1. Uses up limited resources such as ores and crude oil.
  2. Damages habitats through quarrying, mining or felling trees
59
Q

Impact of manufacture

A
  • Uses up land for factories
  • The use of machines and people
60
Q

Impact of use

A
  • The impact of a product on the environment during its use depends on the type of product.
  • For example, a wooden chair has very little impact, unless it needs cleaning or repair.
  • On the other hand, a car will have a significant impact.
61
Q

Impact of Disposal:

A
  • using up land for landfill sites
  • whether any or all of the product can be recycled or reused
62
Q

Phytoextraction:

A
  • plants are grown on a low-grade
    ore that contains lower amounts of metal.
  • the plants absorb metal ions through their roots and concentrate these ions in their cells
  • the plants are harvested and burnt - the ash left behind contains a higher concentration of the metal than the original ore
  • the ash is processed to obtain the metal
63
Q

Advantages and Disadvantages of Phytoextraction:

A
  • Slow
  • Reduces the need to obtain new ore by mining
  • Conserves limited supplies of more valuable ores with higher metal content
64
Q

Bioleaching:

A
  • Certain bacteria can break down low-grade ores to produce an acidic solution containing metal ions.
  • The solution is called a leachate.
  • Bioleaching can be used to extract copper metal.
65
Q

Examples of reactions that go to completion are:

A
  • Complete Combusion of a fuel
  • Many precipitation reactions
  • Effervescence reactions in which a gas escapes
66
Q

Reversible Reaction of Copper Sulfate:

A
  • Bunsen burner heats a bowl of hydrated copper Sulfate.
  • Water is driven off, leaving anhydrous copper sulfate.
  • The burner is turned off and water is added using a pipette.
  • The bowl now contains hydrated copper sulfate again.
67
Q

In the Haber Process:

A
  • nitrogen (extracted from the air) and hydrogen (obtained from natural gas) are pumped through pipes
  • a compressor increases the gas pressure to 200 atmospheres
  • the pressurised gases are heated to 450°C and passed through a reaction chamber containing an iron catalyst to speed up the reaction
  • the reaction mixture is cooled so that ammonia liquefies and can be removed
  • unreacted nitrogen and hydrogen are recycled
68
Q

Increasing the pressure in reversible reactions

A

Position moves towards the fewest molecules of gas

69
Q

What happens if you increase the concentration of a reactant in reversible reactions?

A

Position moves away from that reactant

70
Q

What happens if you increase the temperature in a reversible reaction?

A
  • Position moves in the direction of the endothermic reaction
71
Q

Properties of transition metals:

A

Compared to other metals, the transition metals have these typical properties:

  • high melting points
  • high densities
  • they form coloured compounds
  • they (and their compounds) can act as catalysts
72
Q

What are catalysts?

A
  • Catalysts are substances that speed up the
    rate of reaction without being used up in the reaction.
73
Q

What is corrosion?

A
  • The destruction of a metal by oxidation or chemical action, e.g rusting.
74
Q

Rusting Experiment:

A
  • One tube with air, water and a nail.
  • One tube with water, no air, oil and a nail.
  • One tube with air, no water, calcium chloride and nail.
  • The nail only rusts in the left-hand test tube. It does not rust:
  • in the middle test tube, where there was water but no oxygen (because there was no air in the water)
  • in the right-hand test tube, where there was oxygen (air) but no water
75
Q

What is rusting?

A
  • A corrosion process in which iron or steel reacts with oxygen and water.
76
Q

Explain whether iron is oxidised or reduced when it forms rust.

A

Iron is oxidised because it gains oxygen during rusting.

77
Q

How can oxygen be excluded?

A
  • by storing the metal in an atmosphere of unreactive nitrogen or argon
78
Q

How can water be excluded?

A
  • Storing the metal with a desiccant (a substance that absorbs water vapour, helping to keep air or objects dry) such as calcium chloride.
79
Q

Desiccant

A

a substance that absorbs water vapour, helping to keep air or objects dry

80
Q

Physical Barriers to Oxygen and Water:

A
  • Painting
  • Oiling and Greasing
  • Coating with plastic
81
Q

Electroplating:

A
  • involves using electrolysis to put a thin layer of a metal on the object:
  • the cathode (negatively charged electrode) is the iron or steel object.
  • the anode (positively charged electrode) is the plating metal.
  • the electrolyte contains ions of the plating metal
82
Q

sacrificial metal

A

A more reactive metal than iron, attached to an iron or steel object to prevent the object rusting.

83
Q

Process of sacrificial protection with iron:

A
  • Iron can be protected from rusting if it is in contact with a more reactive metal, such as zinc.
  • The more reactive metal oxidises more readily than iron, so it ‘sacrifices’ itself while the iron does not rust.
84
Q

Galvanising:

A

Coating iron or steel with a layer of zinc to prevent rusting.

85
Q

What is an alloy?

A

An alloy is a mixture of two or more elements, at least one of which is a metal.

86
Q

Explaining Alloy Strength:

A
  • Solid metals have a regular lattice structure.
  • When a force is applied to a metal, layers of atoms can move past each other.
  • The more difficult it is for the layers to move, the more force is needed and the stronger the metal.
  • Copper and zinc atoms have different sizes.
  • This distorts the regular lattice structure in brass, so layers of atoms cannot slide over each other so easily. - This makes brass stronger than copper or zinc alone.
87
Q

Alloy Steels - Mild Steel

A
  • Made from carbon and iron
  • Malleable, ductile
88
Q

Alloy Steels - Tool Steel

A
  • Made from iron and tungsten
  • Hard, resistant to high temperatures
89
Q

Alloy Steels - Stainless Steel

A
  • Made from iron and chromium
  • Hard, resistant to rusting
90
Q

What is mild steel used for?

A
  • Making car body parts because it is easily pressed into shape.
  • Although mild steel rusts, it can be protected by galvanising and painting.
91
Q

What is tool steel used for?

A
  • Making drill bits.
  • These do not easily become damaged by the heating caused by friction during drilling.
92
Q

Why does aluminium foil work?

A
  • Aluminium does not react with water.
  • Its surface is protected by a natural layer of aluminium oxide that allows the metal to resist
    corrosion.
  • Aluminium foil is used in the home for wrapping and storing food because it does not react to substances in food.
    -It is malleable, so it is easily folded into shape around the food.
  • Aluminium has a low density, so pieces of aluminium are relatively lightweight.
93
Q

Magnalium Use

A
  • Magnalium is stronger than aluminium alone but still has a low density. It is used to make aircraft parts.
94
Q

Suggest an explanation for why aluminium is used to make overhead electricity cables.

A
  • Aluminium is a good electrical conductor.
  • Its low density prevents the wires from sagging too much or breaking under their own weight.
95
Q

Uses of copper and brass:

A
  • Copper and brass resist corrosion and are good electrical conductors.
  • Copper is a better conductor than brass, and it is used in electrical wiring.
  • Brass is stronger than copper, so it is used for the pins in electrical plugs.
96
Q

Uses of gold:

A
  • Gold is a very soft and malleable metal. It is also very
    unreactive, so it resists corrosion and stays shiny.
  • The gold used for jewellery is gold alloyed with other metals, often copper.
  • This makes the jewellery much stronger while keeping its ability to stay shiny.
97
Q

Suggest an explanation for why thin gold foil is used to decorate picture frames.

A

Gold stays shiny and has an attractive appearance.

98
Q

What do fertilisers do?

A
  • Provide mineral ions needed for healthy growth in plants.
  • As plants grow, they absorb mineral ions from the water in the soil through their root hair cells.
  • Over time, the concentration of these ions decreases, so farmers and gardeners add fertilisers to the soil.
99
Q

Why must fertiliser compounds be soluble in water?

A
  • So they can be absorbed by the root hair cells.
100
Q

Ammonium nitrate - formula and essential element

A

NH4NO3 Nitrogen

101
Q

Ammonium sulfate - formula and essential element

A

(NH4)2SO4 Nitrogen

102
Q

Ammonium phosphate - formula and essential element

A

(NH4)3PO4 Nitrogen, phosphorus

103
Q

Potassium nitrate - formula and essential element

A

KNO3 Potassium, nitrogen

104
Q

Urea, (NH2)2CO, is used as a fertiliser. Name the essential element it provides.

A

Nitrogen.

105
Q

Making Ammonia:

A

N2(g) + 3H2(g) ⇌ 2NH3(g)

106
Q

The reaction conditions in the Haber process are:

A
  • a temperature of 450°C
  • a pressure of 200 atmospheres
  • an iron catalyst
107
Q

Explain whether ammonia is oxidised or reduced when it is converted to nitric acid.

A

It is oxidised because it reacts with oxygen.

108
Q

How can ammonium sulfate be made in the lab?

A
  • Using dilute ammonia solution and dilute sulfuric acid:

Ammonia + sulfuric acid → ammonium sulfate

2NH3(aq) + H2SO4(aq) → (NH4)2SO4(aq)

  • Both reactants are soluble, so a titration must be used.
109
Q

Method for making ammonium sulfate in the lab?

A
  • Put some dilute sulfuric acid into a beaker.
  • Add a few drops of methyl orange
    indicator.
  • Add dilute ammonia solution drop by drop, stirring in between.
  • Continue step 3 until the colour permanently changes from red to yellow.
  • Add a few more drops of dilute ammonia solution.
  • Pour the reaction mixture into an evaporating basin, and heat carefully over a boiling water bath.
  • Stop heating before all the water has evaporated
  • Leave aside for crystals to form.
  • Pour away excess water and leave the crystals to dry in a warm oven (or pat dry with filter paper).
110
Q

Explain why a few more drops of dilute ammonia solution are added at step 5, after the end-point.

A
  • This is to make sure that all the dilute sulfuric acid has reacted.
  • During heating at stage 6, any excess ammonia gas evaporates and leaves the evaporating basin, so only ammonium sulfate solution is present.
111
Q

When a reversible reaction happens in a closed container, it reaches a dynamic equilibrium. At equilibrium:

A
  • the forward and backward reactions are still happening.
  • the forward and backward reactions have the same rate of reaction.
  • the concentrations of all the reacting substances remain constant (they do not change).
112
Q

How quickly an equilibrium is reached depends upon:

A
  • the pressure of a reacting gas
  • the concentration of a reacting solution
  • the temperature of the reaction mixture
  • the presence of a catalyst

Edexcel GCSE Chemistry Higher (12 decks)

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