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OCR Gateway A GCSE Chemistry > C3 - Chemical Reactions > Flashcards

C3 - Chemical Reactions Flashcards

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1
Q

why are the formulae for metal elements always written as empirical formulae

A
  • because metals exists as giant metallic lattices
    > you don’t need to include numbers in their formulae because they would be too huge
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2
Q

what are diatomic molecules

A
  • non-metal elements in G7 exist as diatomic molecules
    > diatomic molecule has 2 atoms covalently bonded together
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3
Q

give examples of diatomic elements

A
  • G7: F2, Cl2, Br2, I2
  • H2, O2, N2 (not in g7 but still exist as diatomic)
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4
Q

what is the formula for sulfur

A

S8

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5
Q

what is the formula for phosphorus

A

P4

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6
Q

what charge do group 1 ions produce

A

+1

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7
Q

what ions are produced in metal groups 1,2,3

A
  • positive ions
    > as electrons are lost
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8
Q

what ions are produced in non-metal groups 5,6,7

A
  • negative ions
    > as electrons are gained
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9
Q

what are compound ions

A
  • ions that contain more than one element
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10
Q

what is the formula of ammonium ions

A

(NH4)+

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11
Q

what is the formula of sulphate ions

A

(SO4)2-

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12
Q

what is the formula for carbonate ions

A

(CO3)2-

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13
Q

what is the formula for nitrate ions

A

(NO3)-

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14
Q

what is the formula for hydroxide ions

A

(OH)-

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15
Q

write the chemical formula of magnesium chloride

A
  • ionic compound
  • magnesium = Mg2+
  • chlorine = Cl-
  • chemical formula = MgCl2
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16
Q

write the chemical formula of phosphorus trichloride

A
  • covalent compound
  • phosphorus needs to gains 3 electrons to become stable so forms bonds with 3 chlorine atoms
  • chemical formula = PCl3
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17
Q

what does the law of conservation of mass state

A
  • atoms cannot be created or destroyed by chemical reactions, only rearranged
  • the same atoms are present at the start and end of a reaction
    > so the total mass stays the same in a closed system
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18
Q

what is a closed system

A
  • a container in which no substances can enter or leave during the reaction
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19
Q

why does the mass seem to change during some reactions

A
  • substances can enter / leave the reaction mixture in a non-closed system
    > usually happens when the reaction involves a substance in the gas state
    > the gas often escapes the container, and so the mass appears to go down
    > the mass can go up as well as sometimes the reactants can react with oxygen in the air and produce a higher mass
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20
Q

what is the relative atomic mass (Ar)

A
  • mean mass of an atom of an element compared to 1/12th the mass of a carbon 12 atom
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21
Q

how can you find the relative atomic mass of an element

A
  • it’s usually the top small number of an element in the periodic table
    > e.g. Oxygen = 16
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22
Q

what is the relative formula mass (Mr)

A
  • the mean mass of a unit of a substance compared to 1/12th the mass of a C-12 atom
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23
Q

how can you figure out the relative formula mass of a subtance

A
  • add up all their atomic masses
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24
Q

what does a balanced equation show

A
  • how atoms are rearranged in a reaction
  • the relative amounts of each substance involved
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25
Q

what does it mean if a substance is aqueous (aq)

A
  • dissolved in water
    > water is the solvent
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26
Q

what do the state symbols ; s, l, g, aq mean

A
  • s = solid
  • l = liquid
  • g = gas
  • aq = aqueous solution
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27
Q

what are half equations

A
  • half equations show the change that happens to one reactant in aa chemical reaction
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28
Q

the balanced symbol equation for the reaction between sodium + chlorine is: 2Na (s) + Cl2 (g) —> 2NaCl (s)
construct the half equations

A

Na —> Na+ + e-
> sodium ions form when sodium atom loses electrons

Cl2 + 2e- —> 2Cl-
> chloride ions form when chlorine gains electrons but since it is diatomic, you must balance the charges

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29
Q

how do you construct an ionic equation using the example of:
CuSO4 (aq) + 2NaOH (aq) —> Cu(OH)2 (s) + Na2SO4 (aq)

A

CuSO4 (aq) + 2NaOH (aq) —> Cu(OH)2 (s) + Na2SO4 (aq)

  1. split aqueous substances into ions

Cu 2+(aq) + SO4 2-(aq) + 2Na+(aq) + 2OH- (aq) —> Cu (OH)2 (s) + 2Na+(aq) + SO4 2- (aq)

  1. remove spectator ions (ions that appear in the same form on both sides)

Cu2+ (aq) + 2OH-(aq) —> Cu (OH)2 (s)

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30
Q

what is The Mole

A
  • the unit for amount of substance
    > one mole of anything contains the same number of things
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31
Q

define mole

A
  • the amount of any substance that contains the same number of particles (entities) as there are atoms in 12.0g of carbon-12
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32
Q

what is Avogadro’s constant

A
  • the number of entities in 1 mol
    > 6.02x10*23/mol
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33
Q
A
  • 1 mol of any element is it’s relative atomic mass in grams
    > e.g. oxygen Ar = 16 so 1 mol of oxygen has a mass of 16g
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34
Q

what equation links the Avogadro constant to number of moles

A

number of particle (atoms etc) = moles x Avogadro constant

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35
Q

how can you find the number of moles

A
  • you need to know the mass
    moles (n) = mass (g) / mr
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36
Q

what is a limiting reactant

A
  • the reactant that’s in limiting amount
    > the amount of product formed is determined by the amount of limiting reactant
  • one reactant is usually in excess - more is present than needed to react with the other reactant so some is left behind
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37
Q

what is stoichiometry

A
  • refers to the relative amounts of each substance in a balanced equation
    > sometimes called molar ratios
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38
Q

how do you find the limiting reactant in a reaction

A
  • find moles of both reactants
    > whichever has less moles is limiting
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39
Q

what is the empirical formula

A
  • the simplest whole number ratio of the atoms of each element in a compound
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40
Q

how can you calculate the empirical formula of a substance, using the example of:
a substance is made of 20.2% aluminium and 79.8% chlorine, find the empirical formula

A

Al = 20.2% Cl = 79.8%
20.2/27 = 0.748 79.8/35.5 = 2.25
0.748/0.748 = 1 2.25/0.743 = 3
AlCl3
1. divide the percentage or mass by the Ar
2. divide the answers of the above by the smallest number
3. the numbers you get, you place them before their elements

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41
Q

what does it mean if a reaction is exothermic

A
  • during the reaction, energy is released to the surroundings in the form of heat
  • chemical bonds forming releases energy + so it’s exothermic
  • surroundings warmer as heat released
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42
Q

what are some examples of exothermic reactions

A
  • combustion e.g. fuel burning
  • neutralisation
  • respiration
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43
Q

what does it mean if a reaction is endothermic

A
  • during the reaction energy is taken in from the surroundings in the form of heat
  • chemical bonds breaking required energy + so it’s endothermic
  • surroundings colder as heat taken away
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44
Q

what are some examples of endothermic reactions

A
  • photosynthesis
  • thermal decomposition
  • reaction between sodium hydrogen carbonate + citric acid in sherbet sweets
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45
Q

how can you investigate if a reaction is exo/endothermic

A
  • do an experiment:
    1. measure start temp of solution in an insulated container
    2. add the other reactant + stir
    3. measure end temp of solution
    4. calc the different between start + end temp
    > if temp has gone up its’s an exo
    > if temp has lowered, it’s an endo
46
Q

during a chemical reaction, energy is transferred:

A
  • from the surroundings to break bonds in the reactants to form separate atoms - endothermic
  • to the surroundings from the reacting particles when bonds form between atoms - exothermic
47
Q

what is a reaction profile

A
  • a chart showing the energy involved
48
Q

what is the energy change

A
  • the difference between the energy transferred from the surroundings to break bond in the reactants and the energy transferred to the surroundings when bonds form in products
49
Q

how can you work out the energy change

A
  • energy change = energy to break bonds - energy to form bonds
    > kJ/mol
  • b-f
50
Q

describe / draw how the reaction profile for an exothermic reaction would look like

A
  • reactants higher
  • products lower
    > because energy is released so products have lower energy
  • difference in height = energy given out, negative
51
Q

describe / draw how the reaction profile for an endothermic reaction would look like

A
  • reactants lower
  • products higher
    > because energy is taken in so products have more energy
  • difference in height = energy taken in, positive
52
Q

in terms of bond energies, what happens in a chemical reaction, describe exo/endothermic reaction in terms of bonds breaking/forming

A
  • energy is needed to break bonds + energy is released when making bonds
  • endothermic: energy used to break bonds is greater than the energy released when making bonds
  • exothermic: energy released when forming bonds is greater than the energy used to break bonds
53
Q

what is activation energy

A
  • activation energy is the minimum energy needed for a reaction to start
54
Q

how can the activation energy be provided (often)

A
  • by a flame or heating
    > and it breaks bonds in the reactants
55
Q

what is a bond energy

A
  • the energy needed to break 1 mol of a particular chemical bond
56
Q

what is a redox reaction in terms of oxygen

A
  • a reaction in which reduction + oxidation happen at the same time
57
Q

in terms of oxygen describe reduction + oxidation

A
  • reduction = loss of oxygen from a substance
  • oxidation = gain of oxygen by a substance
58
Q

what are reducing + oxidising agents

A
  • reducing agent: a substance which is oxidised when it reduces another substance
  • oxidising agent: a substance which is reduced when it oxidises another substance
59
Q

what are redox reactions in terms of electrons

A
  • oxidation is loss of electrons
  • reduction is gain of electrons
    OIL RIG
60
Q

a copper ion is formed from copper, is it oxidation or reduction

A
  • oxidation
    > because electrons are lost from Cu to form Cu2+
61
Q

how do half equations present redox reactions in electrons in terms of reduction + oxidation using the following:
- aluminium atoms turn to aluminium ions
- iron(III) ions form ion atoms

A
  • oxidation = aluminium, as electrons are lost to form the aluminium ion
    Al —> Al3+ + 3e-
  • reduction = iron, as electrons are gained to form iron atoms
    Fe3+ + 3e- —> Fe
62
Q

what is an acid

A
  • a substance that releases hydrogen ions, H+ (aq) when dissolved in water to make an aqueous solution
63
Q

what is a base

A
  • a substance (usually metal oxides / metal hydroxides) that can neutralise acids
64
Q

what is an alkali

A
  • a base that can dissolve in water, releasing hydroxide ions, OH- (aq)
65
Q

what does the pH of a solution describe

A
  • its relative acidity or alkalinity
66
Q

what pH values describe what on the pH scale

A

pH < 7 = acid
pH = 7 - neutral
pH > 7 = alkaline

67
Q

how can pH be measured

A
  • universal indicator
  • pH probe
68
Q

what is neutralisation

A
  • the reaction between an acid and a base (or alkali) to form salt + water only
  • acid + base —> salt + water
69
Q

how does the pH change with neutralisation

A
  • increases if base/alkali is added to an acid
  • decreases if acid is added to a base/alkali
70
Q

how do you predict the salt made in a neutralisation reaction

A
  • first part comes from the metal in base/alkali
  • second part of the name comes from the acid used
71
Q

what salt will be produced from the following acids:
- hydrochloric acid HCl (aq)
- nitric acid HNO3 (aq)
- sulfuric acid H2SO4 (aq)
- phosphoric acid H3PO4 (aq)

A
  • hydrochloric acid HCl (aq) = chloride
  • nitric acid HNO3 (aq) = nitrate
  • sulfuric acid H2SO4 (aq) = sulfate
  • phosphoric acid H3PO4 (aq) = phosphate
72
Q

what happens during neutralisation in solution

A
  • acids have hydrogen ions H+ (aq)
  • alkaline have hydroxide ions OH-(aq)
    > the two react in neutralisation producing water
    H+(aq) + OH-(aq) —> H2O(l)
73
Q

what is the ionic equation for neutralisation incl state symbols

A

H+(aq) + OH-(aq) —> H2O(l)

74
Q

what happens when acids react with metal carbonates

A
  • acid + carbonate —> salt + water + carbon dioxide
    > carbonates are ionic compounds containing carbonate ion (CO3)2-
    > carbonates react with acids to form salt + water + CO2
75
Q

what happens to the carbon dioxide when an acid + carbonate react

A
  • the CO2 is releases as bubbles in the acid during the reaction
76
Q

what does effervescence mean

A
  • bubbling or fizzing
77
Q

what happens when acid react with metals

A
  • acid + metal —> salt + hydrogen
    > salt made depends on acid + metal used
78
Q

what does a solution consist of

A
  • solute dissolved in a solvent
79
Q

what do the terms concentrated + dilute mean

A
  • concentration is the amount of substance in a given volume of a solution
  • dilute solution is where there’s a small amount of solute in a give volume of solvent
  • concentrated solution is where there’s a large amount of solute in a given volume of solvent
80
Q

in terms of acids what’s the difference between dilute + concentrated

A
  • dilute acid contains a low ratio of acid to volume of solution
  • concentrated acid contains a high ratio of acid to volume of solution
81
Q

what are weak + strong acids

A
  • weak acids are partially ionised - only small fraction of their molecules release H+ ions
  • strong acids are fully ionised - all of their molecules release H+ ions
82
Q

what arrow symbol is used to show a weak + strong acid

A
  • weak = double arrow to show that the reaction doesn’t go to completion
  • strong = single arrow to show that the reaction goes to completion
83
Q

give examples of strong + weak acid

A
  • strong acids = hydrochloric acid, sulfuric acid, nitric acid
  • weak acid = ethanoic acid
84
Q

in a solution, if the pH decreased by 1 what happens to the concentration off H+ ions

A
  • increases by a factor of 10
85
Q

describe the concentration of H+ ions and pH of a strong acid

A
  • high concentration of H+ ions as they are fully ionised
  • low pH close to 0
86
Q

what is electrolysis

A
  • a process in which electric current is passed through a compound, causing a chemical change
87
Q

what is the electrolyte

A
  • the substance which you are breaking down
    > can only be an ionic substance
88
Q

what 3 components do you need for electrolysis to work

A
  1. electrolyte - a compound in its liquid state or in solution, which contains mobile ions and conducts electricity
  2. two electrodes - made from metal or graphite which conducts electricity to the electrolyte
  3. an electric supply such as a power pack / battery
89
Q

when you break down the electrolyte what do you produce

A
  • ions
  • positive ions are called cations
  • negative ions are called anions
90
Q

what are the 2 electrodes

A
  • cathode = negative electrode
  • anode = positive electrode
91
Q

what happens at the 2 electrodes during electrolysis of molten ionic compounds

A
  • positive ions (cations) gain electrons at the cathode and become atoms (reduction)
  • negative ions (anions) lose electrons at the anode and become atoms (oxidation)
92
Q

when ions become atoms/molecules at an electrode, you say that the ions have been

A
  • discharged
93
Q

describe the process of electrolysis in a molten ionic compound

A
  • the molten ionic compound allows the ions to move freely and carry electric charge
    > electrolyte
  • an electric current is passed through the electrolyte causing the ions to move to the electrodes
  • cations move to cathodes
  • anions move to anodes
  • here at the electrodes the ions lose/gain electrons to become neutrally charged
    > cations gain electrons + anions lose electrons
    > leads to gases + solids being formed
94
Q

what are inert electrodes

A
  • electrodes that aren’t changed during electrolysis
    > usually made from unreactive metals e.g. copper, graphite
95
Q

why are inert electrodes used

A

to avoid any unwanted chemical reactions

96
Q

water is naturally partially ionised, so what ions are contained in water

A
  • hydrogen ions H+ (aq)
  • hydroxide ions OH-(aq)
    H2O(l) double arrow H+(aq) + OH-(aq)
97
Q

what happens during the electrolysis of water

A
  • hydrogen ions are discharged at the cathode as hydrogen in its gas state
  • hydroxide ions are discharged at the anode forming water and oxygen in its gas state

4H+(aq) + 4e- —-> 2H2(g) (reduction)
4OH-(aq) —-> 2H2O(l) + O2(g) + 4e- (oxidation)

98
Q

where does reduction + oxidation take place in the electrodes

A
  • cathode = reduction
  • anode = oxidation
99
Q

what happens during the electrolysis of aqueous solutions

A
  • aqueous solution contains ions from dissolved ionic compound + unlike molten electrolyte it also contains hydrogen + hydroxide ions from water
    > the reaction of each ion compete at the electrodes and only one ion is discharged at each electrode
100
Q

what is produced at the cathode in electrolysis of aqueous solutions

A
  • hydrogen is produced
  • but if ions from a less reactive metal than hydrogen is present then that metal will be produced instead
    > copper
    > silver
    > gold
    > platinum
101
Q

what is produced at the anode in electrolysis of aqueous solutions

A
  • oxygen is produced
  • but if halide ions are present then they are produced instead
102
Q

name all the elements in the reactivity series going from most —> least

A
  • potassium
  • sodium
  • calcium
  • magnesium
  • aluminium
  • carbon
  • zinc
  • iron
  • lead
    (hydrogen)
  • copper
  • silver
  • gold
  • platinum
103
Q

what is electroplating

A
  • a type of electrolysis using non-inert electrodes which change during electrolysis
104
Q

what is the cathode, anode and electrolyte in electropllating

A
  • cathode is the object you want to coat with a metal
  • anode is the piece of metal you want to coat the object (cathode) with
  • electrolyte is a solution containing ions of the coating metal
105
Q

how does electroplating work

A
  • positive metal ions in the electrolyte are attracted to the negative cathode and are discharged on the surface of the object (cathode) to be coated
  • these ions are placed by metal ions leaving the surface of the anode - the anode takes part is the reaction and is non-inert
  • overall, metal leaves the anode and is deposited on the cathode
    > the process continues until the anode is used up
106
Q

why must copper be purified

A
  • copper obtained from copper ores must be purified soo that it can conduct electricity well enough to be used in electrical cables
107
Q

what are the cathode, anode and electrolyte in copper purification

A
  • cathode = very pure copper
  • anode = impure copper
  • electrolyte = copper sulfate solution ( has Cu2+ ions)
108
Q

describe the process of purification of copper

A
  • during electrolysis, the cathode gains copper atoms + increases in mass
  • the anode loses copper atoms + decreases in mass
  • impurities from the anode fall off and collect underneath it
109
Q

write the half equations for the purification of ccopper

A
  • cathode = Cu2+ + 2e- —-> Cu (reduction)
  • anode = Cu —-> Cu2+ + 2e- (oxidation)
110
Q

how can you measure the pH of a solution using universal indicator

A
  • add a few drops of universal indicator to the solution you’re testing
  • compare the colour he solution goes to the pH chart
    > pH 0 (acid) = red all the way to pH 14 (alkali) where it goes purple
    > essentially a rainbow
111
Q

how can you use a pH probe to measure the pH of a solution

A
  • pH probe is attached to pH meter to measure pH electronically
  • pH probe is placed in solution + pH is given on digital display as a numerical value
    > make sure to rinse the probe with deionised water inbetween readings

OCR Gateway A GCSE Chemistry (7 decks)

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