Paper 1 Flashcards
(102 cards)
1
Q
Bohr model
A
Describes an atom as a small dense nucleus with electrons orbiting around the nucleus.
2
Q
Relative atomic mass
A
The weighted mean mass of an atom compared with 1/12th mass of an atom of carbon-12.
3
Q
Relative isotopic mass
A
The mass of an atom of an isotope compared with 1/12th mass of an atom of carbon-12.
4
Q
Relative formula mass
A
The mass of the formula unit of a compound with a giant structure.
5
Q
Ionic compound
A
A compound which is made up of oppositely charged ions that are held together by electrostatic forces.
6
Q
Ideal gas
A
A gas which has molecules that occupy negligible space with no interactions between them.
7
Q
Relative molecular mass
A
The average mass of one molecule of an element or compound compared to 1/12th the mass of an atom of carbon-12.
8
Q
Acid
A
Compounds that release H^+ ions in aqueous solution.
9
Q
Alkali
A
Alkalis release OH^- ions into aqueous solution.
10
Q
Base
A
A substance that can accept H^+ ions from another substance.
11
Q
Weak acid
A
An acid that only partially dissociates in solution.
12
Q
Strong acid
A
An acid that completely dissociates in solution.
13
Q
Atomic orbital
A
A region of space around the nucleus that can hold up to 2 electrons with opposite spins.
14
Q
Average bond enthalpy
A
The energy required to break one mole of gaseous bonds. Used as a measurement of the strength of a covalent bond.
15
Q
Covalent bond
A
A strong bond formed between 2 atoms due to the electrostatic attraction between a shared pair of electrons and the atomic nuclei.
16
Q
Dative covalent (coordinate) bond
A
A type of covalent bond in which both of the electrons in the shared pair come from one atom.
17
Q
Electronegativity
A
The ability of an atom to attract bonding electrons in a covalent bond. This is often quantified using Pauling’s electronegativity values.
18
Q
Ionic bond
A
Electrostatic attraction between positive and negative ions.
19
Q
Ionic lattice
A
A giant structure in which oppositely charged ions are strongly attracted in all directions.
20
Q
Preparing standard solution
A
- Weigh the sample bottle containing the solid on a balance.
- Transfer solid to beaker and reweigh sample bottle.
- Record the difference in mass.
- Add distilled water and stir with a glass rod until all the solid has dissolved.
- Transfer to a volumetric flask with washings.
- Make up to the 250 cm^3 mark with distilled water.
- Shake flask.
21
Q
Titration method
A
- Fill the burette with the standard solution of known concentration, ensuring the jet space in the burette is filled and doesn’t contain air bubbles.
- Use a pipette filler and pipette to transfer 25 cm^3 of the solution with unknown concentration into a conical flask.
- Add two to three drops of indicator.
- Record the initial burette reading.
- Titrate the contents of the conical flask by adding the solution to it from the burette until the indicator undergoes a definite, permanent colour change.
- Record the final burette reading and calculate the titre volume.
- Repeat until at least two concordant results are obtained (within 0.1 cm^3 of each other).
22
Q
Amphoteric substances
A
Substances that can act as acids and bases
23
Q
Acid + Carbonate
A
Salt + carbon dioxide + water
24
Q
What is a salt
A
A compound that is formed when H+ of an ion is replaced by a metal ion or positive ion.
25
Acid + metal oxide
Salt + water
26
Acid + alkali
Salt + water
27
Acid + metal
Salt + hydrogen
28
How does a mass spectrometer work?
The sample is made into positive ions.
They pass through the apparatus and are separated according to the mass to charge ratio.
A computer analyses the data and produces mass spectrum.
29
Avogadro’s law
Under the same temperature and pressure, one mole of any gas would occupy the same volume.
30
Ideal gas behaviour
They are in continuous motion.
No intermolecular forces experienced.
Exert pressure when they collide with each other or container.
No kinetic energy lost in the collisions.
When temperature increases, kinetic energy of gases also increase.
31
Ideal gas equation
pV=nRT
32
What does concentrated mean
Large amount of solute per dm^3 of solvent.
33
What does dilute mean
Small amount of solute per dm^3 of solvent.
34
Oxyanions
Negative ions that have an element along with oxygen
35
Intramolecular forces
Forces within a molecule (usually covalent bonds)
36
Intermolecular forces
Forces between the molecules (weaker than intramolecular forces)
37
Why does water have a high melting and boiling point?
Strong intermolecular forces of hydrogen bonding between the molecules.
In ice and water the molecules are tightly held together by hydrogen bonds, a lot of energy is therefore required to break the water molecules apart and melt or boil them.
38
Why is water more dense than ice?
In ice the water molecules are packed in a 3D hydrogen bonded network in a rigid lattice. Each oxygen atom is surrounded by hydrogen atoms. This way of packing the molecules in a solid and the relatively long bond lengths of the hydrogen bonds means that the water molecules are slightly further apart than in the liquid form.
39
How to determine the rate constant for a first order reaction from the constant half life?
k = ln2 / t1/2
40
First ionisation energy
The removal of one mole of electrons from one mole of gaseous atoms.
41
Second ionisation energy
The removal of one mole of electrons from one mole of gaseous 1+ ions to form one mole of 2+ ions.
42
Endothermic
A reaction which takes in energy. More energy is required to break bonds than is released by making bonds.
43
Water treatment
The addition of chlorine to water to kill bacteria. The risks associated with the use of chlorine to treat water are the hazards of toxic chlorine gas and the possible risks from the formation of chlorinated hydrocarbons.
44
Activation energy
The minimum energy required for a reaction to take place.
45
Enthalpy change
The change in the heat content of a system during a reaction.
46
Enthalpy change of combustion
The enthalpy change that takes place when one mole of a substance is completely combusted.
47
Enthalpy change of formation
The enthalpy change that takes place when one mole of a compound is formed from its elements.
48
Enthalpy change of neutralisation
The enthalpy change that takes place when one mole of water is formed from a neutralisation reaction.
49
Exothermic
A reaction which gives out energy. More energy is released by bond making than is used in bond breaking.
50
Dynamic equilibrium
A closed system in which the rates of the forward and reverse reactions are equivalent. The concentrations of reactants and products don’t change.
51
Colorimetry
A technique used to measure the amount of light absorbed by a solution, used to determine the rate of a reaction.
52
Bronsted-Lowry acid
A proton donor
53
Bronsted-Lowry base
A proton accepter
54
Buffer solution
A system that minimises pH change on addition of small amounts of an acid or base. A buffer solution can be formed from a weak acid and a salt of the weak acid or from excess weak acid and a strong alkali.
55
Enthalpy change of atomisation
The enthalpy change that takes place when one mole of gaseous atoms is formed from an element in its standard state.
56
Enthalpy change of hydration
The enthalpy change that takes place when one mole of gaseous ions are dissolved in water (exothermic).
57
Enthalpy change of solution
The enthalpy change that takes place when one mole of solute is dissolved.
58
First election affinity
The amount of energy released when one mole of electrons is added to one mole of gaseous atoms, forming one mole of 1- ions.
59
Lattice enthalpy
The formation of one mole of an ionic lattice from gaseous ions.
60
Entropy
A measure of the dispersal of energy in a system. The greater the entropy the more disordered the system.
61
Free energy change (delta G)
delta G = delta H -T delta S
62
Equimolar solution
A solution of ions in which there is an equal number of rod moles of each ion.
63
Fuel cell
A type of cell that requires a constant supply of fuel and oxygen in order to generate a potential difference.
64
Standard electrode potential
The e.m.f. of a half cell compared with a standard hydrogen half cell. Measured under standard conditions.
65
Storage cell
A type of cell that can store energy. Storage cells convert chemical energy into electrical energy by a reaction and they may be recharged by reversing the chemical reaction.
66
Bidentate ligand
A substance that can form 2 dative covalent bonds with a metal ion.
67
Cis-platin
The cis-isomer of Pt(NH3)2Cl2 used as an anti cancer drug. It binds to DNA, preventing cell replication.
68
Transition elements
d-block elements that can form an ion with an incomplete d-sub shell. They have more than one oxidation state, form coloured ions and can act as catalysts.
69
Limitations of using approximations to aka related calculations for ‘stronger’ weak acids.
we assume the position of equilibrium lies to the left but as the acid gets stronger the position of equilibrium lies more to the right.
70
Limitations of predictions made by delta G about feasibility, in terms of kinetics.
if a low temp is required there is a slow rate so less frequent collisions so the reaction might not start.
71
Importance of Fe in haemoglobin
The Fe^2+ ions from 6 coordinate bonds.
72
Ligand substitution on haemoglobin
On one of the positions there is a water ligand attached to the iron. In the lungs oxygen concentration is high so the water ligand is substituted for an oxygen molecule. When it gets to where oxygen is needed the oxygen molecule is exchanged for a water molecule again. If carbon monoxide is inhaled the haemoglobin swaps its water ligand for a CO ligand. CO is a strong ligand and doesn’t readily exchange with O2 or H2O ligands, meaning the haemoglobin can’t transport oxygen anymore.
73
Interconversion between Fe2+ and Fe3+
Fe2+ is oxidised to Fe3
74
Explain the drop between groups 2 and 3
The outer electron in group 3 elements is in a p orbital rather than an s orbital. A p orbital has a slightly higher energy than an s orbital in the same shell, so the electron is to be found further from the nucleus. The p orbital also has additional shielding provided by the s electrons. These factors override the effect of the increased nuclear charge, resulting in the ionisation energy dropping slightly.
75
Explain the drop between groups 5 and 6
In the group 5 elements the electron is being removed from a singly occupied orbital. In group 6 elements the electron is being removed from an orbital containing two electrons. The repulsion between two electrons in an orbital means that electrons are easier to remove from shared orbitals.
76
Explain the strength of graphite
Strong covalent bonds in the hexagon sheets means it has a very high melting point. Graphite is insoluble in any solvent because the covalent bonds are too strong to break.
77
Explain the conductive properties of graphite
Contains delocalised electrons that are free to move along the sheets so and electric current can flow.
78
What is graphite used for
The weak forces between the layers in graphite are easily broken so the sheets can slide over each other so can feel slippery. This means it is used as a dry lubricant and in pencils.
79
Group 2 redox reaction with oxygen
2Mg (s) + O2 (g) —> 2MgO (s)
80
Group 2 redox reaction with dilute acids
Mg (s) + 2HCl (aq) —> MgCl2 (aq) + H2 (g)
81
Group 2 redox reaction with water
Mg (s) + 2H2O (l) —> Mg(OH)2 (aq) + H2 (g)
82
Graphene is one layer of graphite what is graphene used for
High-speed electronics and aircraft technology useful material for touchscreens
83
What is Ca(OH)2 used for in agriculture?
To neutralise acidic soils.
84
What is Mg(OH)2 and CaCO3 used for?
They are used in indigestion tablets as antacids.
85
cold sodium hydroxide and chlorine:
NaOH + Cl2 —> NaCl + NaClO + H2O
86
Hot sodium hydroxide and chlorine
6NaOH + 3Cl2 —> 5NaOH + NaClO3 + 3H2O
87
What is on the y-axis of a Boltzmann distribution?
Number of molecules
88
What is on the x-axis of a Boltzmann distribution?
Kinetic energy
89
You can use electrode potentials to predict whether a reaction will happen but sometimes the predictions are wrong, why?
The conditions are not standard or the rate of reaction might be slow or if a reactions has a high activation energy.
90
How does a fuel cell produce electricity?
Reacting a fuel, usually hydrogen, with an oxidant, which is most likely to be oxygen.
91
Why do transition elements make good catalysts?
They can change oxidation states by gaining or losing electrons within their d orbitals. This means they can transfer electrons to speed up reactions.
Good at adsorbing substances onto their surfaces to lower the activation energy of reactions.
92
How to test for oxygen
Relight a glowing splint
93
How to test for hydrogen
Squeaky pop
94
How to test for carbon dioxide
Limewater goes cloudy
Ca(OH)2 (aq) + CO2 (g) —> CaCO3 (s) + H2O (l)
95
How to test for ammonia
Moist red litmus turns blue
96
How to test for nitrogen
Will extinguish a lit splint (like CO2) but will not turn limewater cloudy
97
How to test for ethane
Bromine water
Orange —> colourless
98
How to test for CO3^2-
Test with dilute acid and if it fizzes pass the gas produced through limewater which will go cloudy.
99
How to test for SO4^2-
Test with barium chloride
White ppt formed as barium sulphate is insoluble
100
How to test halides
Test with silver nitrate and dilute nitric acid, then add dilute and concentrated ammonia (NH3) and check solubility.
101
How to test NH4^+
Warm NH4^+ (aq) and NaOH (aq) ammonia will be released and damp red litmus will turn blue.
102
Order for sequence of tests
Carbonate
Sulfate
Halides
