Chemistry Paper 1 Flashcards
1
Q
Atoms
A
Smallest part of an element that can exist.
2
Q
Elements
A
Have chemical symbols (e.g. H, O, Na).
3
Q
Compounds
A
2+ elements chemically combined in fixed ratios.
4
Q
Formulas
A
Represented by formulas (e.g. H₂O, CO₂).
5
Q
Chemical reactions
A
Formed by chemical reactions (new substances + energy change).
6
Q
Mixtures
A
Not chemically combined → can be separated.
7
Q
Separation techniques
A
Methods to separate mixtures.
8
Q
Filtration
A
A separation technique.
9
Q
Crystallisation
A
A separation technique.
10
Q
Simple/fractional distillation
A
A separation technique.
11
Q
Chromatography
A
A separation technique.
12
Q
Dalton’s model
A
Atoms = solid spheres.
13
Q
Thomson’s model
A
Plum pudding model (electrons in positive ‘pudding’).
14
Q
Rutherford’s model
A
Nuclear model (mass & positive charge in nucleus).
15
Q
Bohr’s model
A
Electrons in energy levels.
16
Q
Chadwick’s discovery
A
Discovered neutrons.
17
Q
Proton
A
Charge: +1, Relative Mass: 1.
18
Q
Neutron
A
Charge: 0, Relative Mass: 1.
19
Q
Electron
A
Charge: -1, Relative Mass: ≈0 (very small).
20
Q
Atomic Number
A
Number of Protons.
21
Q
Mass Number
A
Protons + Neutrons.
22
Q
Size of Atoms
A
Atom radius ≈0.1 nm (1 × 10⁻¹⁰ m).
23
Q
Nucleus size
A
Nucleus ≈1 × 10⁻¹⁴ m.
24
Q
Relative Atomic Mass
A
Average of all isotopes (based on abundance).
25
Relative Atomic Mass formula
Relative Atomic Mass = Σ (isotope mass × abundance) ÷ 100.
26
Electronic Structure
Electrons occupy lowest available energy levels (shells).
27
Example of Sodium's structure
Sodium = 2,8,1.
28
1st 20 elements' structures
Know 1st 20 elements' structures.
29
Hydrogen structure
H = 1.
30
Helium structure
He = 2.
31
Lithium structure
Li = 2,1.
32
Periodic Table
Elements arranged by atomic number.
33
Groups
Columns in the periodic table that have the same number of outer electrons, leading to similar properties.
34
Periods
Rows in the periodic table that indicate the number of electron shells.
35
Mendeleev's Periodic Table
Early tables based on atomic weight, where gaps were left and elements were swapped to fit patterns.
36
Isotopes
Discovery of isotopes explained irregularities in the periodic table.
37
Group 0
Elements with a full outer shell, making them very unreactive.
38
Group 1
Elements with 1 outer electron, which are very reactive.
39
Group 7
Elements with 7 outer electrons, which are reactive non-metals.
40
Transition Metals
Elements that have higher melting points, densities, and strength compared to Group 1, and are less reactive.
41
Ionic Bonding
A type of bonding where metal atoms lose electrons to become positive ions and non-metals gain electrons to become negative ions.
42
Covalent Bonding
A type of bonding where atoms share electrons, forming strong covalent bonds.
43
Metallic Bonding
A type of bonding characterized by a giant structure with delocalized electrons.
44
States of Matter
Solid, Liquid, Gas - particle model where change of state depends on bond strength.
45
Ionic Compound Properties
Ionic compounds have high melting/boiling points and conduct electricity only when molten or dissolved.
46
Small Molecule Properties
Small molecules have low melting/boiling points and do not conduct electricity.
47
Polymers
Very large molecules with strong covalent bonds in chains, solid at room temperature.
48
Giant Covalent Structures
Structures where all atoms are bonded with strong covalent bonds, resulting in very high melting points.
49
Properties of Metals
Metals have high melting/boiling points and are soft when pure.
50
Diamond
A form of carbon with 4 covalent bonds per atom, very hard, with a high melting point and does not conduct electricity.
51
Graphite
A form of carbon with 3 covalent bonds per atom, allowing layers to slide and one delocalized electron per atom, which conducts electricity.
52
Graphene
A single layer of graphite that is very strong and conductive.
53
Fullerenes
Hollow structures, such as C₆₀ Buckminsterfullerene.
54
Nanoparticles
Particles sized 1-100 nm with unique properties compared to bulk materials.
55
Law of Conservation of Mass
States that no atoms are lost or made in a reaction, leading to mass of reactants equaling mass of products.
56
Relative Formula Mass (Mr)
The sum of the relative atomic masses of all atoms in a compound.
57
Moles
1 mole equals 6.02 × 10²³ particles (Avogadro constant).
58
Limiting Reactants
The reactant that is used up first in a reaction, limiting the amount of product formed.
59
Percentage Yield
The formula for percentage yield is (actual yield / theoretical yield) × 100.
60
Atom Economy
A measure of how much of the reactants end up as useful product, calculated as (Mr of desired product / Total Mr of reactants) × 100.
61
Solution Concentrations
Concentration in mol/dm³ is calculated as moles divided by volume.
62
Volumes of Gases
At room temperature and pressure, 1 mole of gas occupies 24 dm³.
63
Reactivity Series
A list of metals arranged by reactivity, from most reactive (K) to least reactive (Cu).
64
Oxidisation and Reduction
Oxidisation is loss of electrons, and reduction is gain of electrons.
65
Neutralisation Reactions
Reactions where acids react with alkalis or bases to produce salt and water.
66
Separation Techniques
Methods to separate mixtures: Filtration, Crystallisation, Simple/fractional distillation, Chromatography.
67
Proton
Particle with a charge of +1 and a relative mass of 1.
68
Neutron
Particle with a charge of 0 and a relative mass of 1.
69
Electron
Particle with a charge of -1 and a relative mass of approximately 0 (very small).
70
Atom Size
Atom radius ≈0.1 nm (1 × 10⁻¹⁰ m).
71
Isotopes
Same protons, different neutrons (e.g. Carbon-12 & Carbon-14).
72
Sodium Electronic Structure
Example: Sodium = 2,8,1.
73
Periodic Table Structure
Elements arranged by atomic number.
74
Groups
Columns in the periodic table with the same number of outer electrons → similar properties.
75
Periods
Rows in the periodic table indicating the number of electron shells.
76
Group 0
Full outer shell → very unreactive.
77
Group 1
1 outer electron → very reactive.
78
Group 7
7 outer electrons → reactive non-metals.
79
Transition Metals
Higher melting points, densities, strength compared to Group 1.
80
Ionic Bonding
Transfer of electrons (metal + non-metal).
81
Covalent Bonding
Sharing of electrons (non-metal + non-metal).
82
Metallic Bonding
Delocalized electrons shared in metal atoms.
83
States of Matter
Solid, Liquid, Gas - particle model (spheres).
84
Ionic Compound Properties
High melting/boiling points (strong bonds); conduct electricity only when molten or dissolved.
85
Small Molecule Properties
Low melting/boiling points (weak intermolecular forces); do not conduct electricity.
86
Polymers
Very large molecules with strong covalent bonds in chains.
87
Giant Covalent Structures
All atoms bonded with strong covalent bonds; very high melting points.
88
Properties of Metals
High melting/boiling points; soft when pure (layers slide).
89
Diamond
4 covalent bonds per atom; very hard, high melting point, doesn't conduct.
90
Graphite
3 covalent bonds per atom, layers slide; one delocalized electron per atom = conducts electricity.
91
Graphene
One layer of graphite, very strong and conductive.
92
Nanoparticles
1-100 nm; high surface area:volume ratio.
93
Law of Conservation of Mass
No atoms are lost or made in a reaction → Mass of reactants = mass of products.
94
Relative Formula Mass
Add up Ar (relative atomic masses) of all atoms in a compound.
95
Moles
1 mole = 6.02 × 10²³ particles (Avogadro constant).
96
Limiting Reactants
One reactant is used up first = limits amount of product formed.
97
Percentage Yield
% Yield = (actual yield / theoretical yield) × 100.
98
Atom Economy
Atom Economy = (Mr of desired product / Total Mr of reactants) × 100.
99
Solution Concentrations
Concentration = moles / volume.
100
Gas Volumes
At room temp & pressure: 1 mole of gas = 24 dm³.
101
Metal Oxides
Metals + Oxygen → Metal oxides (oxidation: gain of oxygen).
102
Reactivity Series
Metals arranged by reactivity.
103
Oxidisation and Reduction
Oxidisation = loss of electrons; Reduction = gain of electrons.
104
Reactions of Acids
Acids + Metals → Salt + Hydrogen.
105
Neutralisation Reactions
Acids + Alkalis (soluble bases) → Salt + Water.
106
Insoluble base
A substance that does not dissolve in a solvent, such as copper(II) oxide, used in neutralisation reactions.
107
Filtrate
The liquid that has passed through a filter, containing dissolved substances after filtering out excess solid base.
108
pH Scale
A scale used to measure the acidity or alkalinity of a solution, where pH < 7 indicates acidity, pH > 7 indicates alkalinity, and pH = 7 indicates neutrality.
109
Neutralisation reaction
A chemical reaction between an acid and an alkali that produces water, represented by the equation H⁺ + OH⁻ → H₂O.
110
Universal indicator
A pH indicator that changes color to indicate the pH level of a solution.
111
Titration
A laboratory technique used to determine the concentration of an acid or alkali by reacting it with a solution of known concentration.
112
Indicators
Substances that change color at a specific pH level, such as phenolphthalein or methyl orange.
113
Concordant results
Results that are consistent and agree closely with each other, often obtained by repeating an experiment.
114
Strong acids
Acids that are fully ionised in solution, such as hydrochloric acid (HCl), sulfuric acid (H₂SO₄), and nitric acid (HNO₃).
115
Weak acids
Acids that are partially ionised in solution, such as ethanoic acid and citric acid.
116
pH decrease effect
For every decrease of 1 in pH, the concentration of H⁺ ions increases by a factor of 10.
117
Electrolysis
A process that uses electric current to break down electrolytes, either in molten form or in solution.
118
Cathode
The electrode where reduction occurs, attracting positive ions (cations) during electrolysis.
119
Anode
The electrode where oxidation occurs, attracting negative ions (anions) during electrolysis.
120
Molten ionic compounds
Ionic compounds that are melted to allow ions to move freely, such as lead bromide (PbBr₂).
121
Reactive metals
Metals that are too reactive to be extracted using carbon reduction, often extracted using electrolysis.
122
Aluminium extraction
The process of extracting aluminium from molten Al₂O₃ using cryolite in electrolysis.
123
Aqueous solutions
Solutions where water is the solvent, often used in electrolysis to facilitate the movement of ions.
124
Gas production at cathode
Hydrogen gas is produced at the cathode if the metal is more reactive than hydrogen.
125
Gas production at anode
Oxygen gas is produced at the anode unless a halide is present, in which case a halogen is formed.
126
Inert electrodes
Electrodes that do not react with the electrolyte, such as carbon or platinum, used in electrolysis.
127
Testing gases
The process of identifying gases produced during electrolysis, such as using a lit splint to test for hydrogen.
128
Half Equations
Reduction at cathode: 2H⁺ + 2e⁻ → H₂
129
Half Equations
Oxidation at anode: 4OH⁻ → O₂ + 2H₂O + 4e⁻
130
Exothermic reactions
Transfer energy to the surroundings, increasing the temperature of surroundings.
131
Examples of Exothermic reactions
Combustion, many oxidation reactions, neutralisation.
132
Uses of Exothermic reactions
Hand warmers, self-heating cans.
133
Endothermic reactions
Take in energy from surroundings, decreasing the temperature of surroundings.
134
Examples of Endothermic reactions
Thermal decomposition, citric acid + sodium hydrogen carbonate.
135
Uses of Endothermic reactions
Sports injury packs.
136
Law of conservation of energy
Energy is conserved in chemical reactions.
137
Activation energy
Minimum energy required for particles to react.
138
Exothermic profile
Products have less energy than reactants.
139
Endothermic profile
Products have more energy than reactants.
140
Overall Energy Change
Energy to break bonds - Energy released when bonds form.
141
Bond breaking
Energy in.
142
Bond making
Energy out.
143
Cells
Produce electricity from chemical reactions between electrodes and electrolytes.
144
Voltage
Depends on the type of electrodes and electrolyte.
145
Simple cell
Two different metals in an electrolyte.
146
Battery
Multiple cells in series for higher voltage.
147
Non-rechargeable cells
Stop when reactants used up.
148
Rechargeable cells
Chemical reactions can be reversed with external current.
149
Fuel Cells
Supplied with external fuel (e.g. hydrogen) and oxygen/air.
150
Oxidation of hydrogen in Fuel Cells
Produces electricity + water.
151
Hydrogen fuel cells
A potential alternative to rechargeable batteries.
