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OCR A Level Chemistry > Topic 2 - Foundations in chemistry > Flashcards

Topic 2 - Foundations in chemistry Flashcards

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1
Q

Definition of Relative Atomic Mass (Ar)

A

Weighted mean mass of an atom relative to 1/12th the mass of an atom of carbon 12

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2
Q

Definition of Relative Formula Mass (Mr)

A

The term used when working out the calculation for compounds with giant structures

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3
Q

Name 3 types of giant structures

A

Giant Ionic
Giant Covalent
Giant Metallic

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4
Q

Definition of Relative Molecular Mass

A

The term used when working out the calculation for compounds that are simple molecules

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5
Q

Name a type of simple molecular structure

A

Covalent Compounds

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6
Q

What is the formula for finding out how much of a compound is made up of a particular element?

A

% of element = ((number of atoms of element x relative atomic mass of the element) / relative formula mass of compound) x 100

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7
Q

What is the molar mass?

A

.The mass per mol of a substance in g mol^-1

.Same number as Mr

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8
Q

What is the amount of substance + unit?

A

.A means of counting the number of particles in a substance

.Unit is the mol

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9
Q

What is Avogadro’s Constant?

A

(NA) 6.02x10^23 mol-1

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10
Q

What is everything measured relative to?

A

Carbon-12

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11
Q

Define the mole

A

The amount of any substance containing as many elementary particles as there are carbon atoms in 12 grams of carbon-12 (6.02x10^23) particles

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12
Q

number of mols (mol) =

A

mass (g) / molar mass (g mol-1)

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13
Q

Smallest mass of an atom

A

1.67x10^-27

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14
Q

The largest mass of an atom

A

4.52x10^-25

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15
Q

What do we use the smallest mass of an atom to do?

A

Produce a relative scale called the unified atomic mass unit ‘u’ = 1.67x10^-27

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16
Q

If hydrogen-1 = 1u, what does carbon-12 =?

A

12u

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17
Q

Define relative isotopic mass

A

The mass of an atom of an isotope relative to 1/12th the mass of a carbon 12 atom

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18
Q

What is 1/12th the mass of a carbon-12 atom?

A

1u

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19
Q

Define relative atomic mass

A

The relative atomic mass is the ‘weighted mean’ mass of an atom relative to 1/12th the mass of a carbon-12 atom

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20
Q

formula to find out relative atomic mass/weighted mean?

A

((mass x abundance) + (mass x abundance)) / 100

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21
Q

Define molecular formula

A

The number and type of atoms of each element in a molecule

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22
Q

Define empirical formula

A

Shows the simplest whole number ration of atoms of each element in a compound

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23
Q

State the steps for finding the empirical formula

A
  1. find the mass
  2. calculate the moles
  3. divide all the results by the smallest value to get the ration
  4. adjust the ratio to get whole numbers
  5. workout the empirical formula
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24
Q

What is water of crystallisation?

A

Water molecules that are bonded into a crystalline structure of a compound

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25
Q

What doe solid compounds formed from aqueous solutions have trapped in their crystal structures?

A

The crystals that are formed have water molecules trapped in the crystal structure

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26
Q

What does copper sulphate exist as?

A

Blue crystals

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27
Q

How does copper sulphate lose its water?

A

Heat

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28
Q

What colour is anhydrous copper sulphate?

A

White

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29
Q

What is the formula of copper sulphate?

A

CuSO4.xH2O

x = number of water molecules

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30
Q

Anhydrous definition

A

Contains no waters of crystillisation

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31
Q

Hydrated definition

A

A crystallised compound containing water molecules

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32
Q

Where is the water in copper sulphate?

A

Water forms part of the crystal sturcture

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33
Q

Four steps of the hydrated and anhydrous practical

A
  1. Weigh an empty crucible
  2. Add the hydrated salt into the weighed crucible, weigh the crucible and the hydrated salt
  3. Using a pipe clay triangle, support the crucible containing the hydrated salt on a tripod. Heat the crucible and contents gently for about one minute. Then heat it strongly for a further three minutes.
  4. Leave the crucible to cool. Then weigh the crucible and anhydrous salt
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34
Q

What is a binary compound?

A

Binary compounds contain 2 elements only

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35
Q

How to name a binary compound with one example

A

To name it, take the first element then change second elements name to –ide (for ionic compounds the metal always comes first) for example magnesium oxide (MgO), sodium chloride (NaCl), Calcium sulphide (CaS)

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36
Q

What is a polyatomic ion?

A

Ions comprised of more than one atom

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37
Q

What are brackets in formulas used for?

A

Polyatomic ions need to go into brackets

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38
Q

What is a redox reaction?

A

When reduction and oxidisation happen at the same time

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39
Q

What is OILRIG

A 
Oxidisation
Is
Loss
Reduction
Is
Gain
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40
Q

What assumptions were made during the copper sulphate anhydrous and hydrated salt experiment? Any extra infomation (how to prevent/any problems that could arise)

A

.That all the water has been lost – could heat the mass then weigh and repeat till it stays the same to counteract this
.No further decompositions – copper oxide might have been made if we heated it too much

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41
Q

When are you allowed to use half values for equation balancing?

A

In combustion equations, specifically on the oxygen

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42
Q

In what order do you balance combustion equations (hydrocarbons)?

A

Carbon first then the hydrogen before finally the oxygen

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43
Q

What does aqueous mean?

A

Dissolved in water

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44
Q

Equipment needed for the determination of the formula for magnesium oxide practical

A 
.access to balance accurate to two decimal places
.Crucible and lid
.bunsen burner
.Tripod stand
.heat-proof mat
.clay pipe triangle
.Tongs
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45
Q

Health and Safety for the determination of the formula for magnesium oxide practical

A

.Wear eye protection at all time (safety goggles)
.Take care not to touch any apparatus that is hot
.Take particular care at steps 3 and 4, do not look at any bright light given off by the reacting magnesium while it is being heated
.Do not place the magnesium ribbon directly in the Bunsen flame

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46
Q

Method for the determination of the formula for magnesium oxide practical

A
  1. Measure the mass of crucible and lid
  2. Put the magnesium ribbon to the crucible. You will need to coil the magnesium so that it fits. Reweigh the crucible and lid.
  3. Arrange the equipment with a tripod on a heatproof mat, a clay pipe triangle on the tripod, a crucible in the triangle and a bunsen burner under it. Raise the crucible lid slightly using tongs to control the reaction
  4. When the reaction is nearly complete, place the crucible lid on the heatproof mat and heat the crucible strongly for 5 minutes. During this time, tap the magnesium oxide gently with tongs to break up the residue
  5. Allow the crucible to cool and reweigh the crucible, its contents, and lid
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47
Q

What is molar gas volume?

A

The molar gas volume is the volume per mole of gas molecules at the stated temperature and pressure

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48
Q

What are the conditions at RTP?

A

Conditions: RTP, Room temperature and pressure
20’C
101KPa (1atm) pressure

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49
Q

At RTP what does 1 mole of gas have a volume of?

A

24dm^3 (24000cm^3)

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50
Q

What is the volume and molar gas volume equation?

A

amount n (mol) = Volume V / molar gas volume V

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51
Q

How is a gas a ‘perfect’ or ‘ideal’ gas?

A

If it obeys the ideal gas equation

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52
Q

When is a gas most close to obeying the ideal gas equation and why?

A

real gases obey the equation very closely at low pressure (no more than atmospheric pressure) and high temperature (room temperature). Under these conditions a gas is most like a gas and least like a liquid.

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53
Q

What 5 assumptions does the kinetic theory make about gas molecules?

A
  • The particles are moving in straight lines at random.
  • We can neglect the volume of the particles themselves in comparison with the total volume of the gas (occupy negligible volume).
  • The particles do not attract one another (exert no force on one another).
  • The kinetic energy of the particles is proportional to the temperature of the gas.
  • No energy is lost in collisions between particles.
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54
Q

What is the ideal gas equation?

A

pV = nRT

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55
Q

In the ideal gas equation what is ‘p’ and what is it measured in?

A

Pressure, measured in Pa

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56
Q

In the ideal gas equation what is ‘n’ and what is it measured in?

A

Amount of gas, measured in mol

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57
Q

In the ideal gas equation what is ‘T’ and what is it measured in?

A

Temperature, measured in K

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58
Q

In the ideal gas equation what is ‘V’ and what is it measured in?

A

Volume, measured in m^3

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59
Q

In the ideal gas equation what is ‘R’ and what is it measured in?

A

Ideal gas constant, measured in J/mol/K

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60
Q

What is the ideal gas constant?

A

8.31 J/mol/K

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61
Q

How to convert from cm^3 to m^3

A

x10^-6

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62
Q

How to convert from dm^3 to m^3

A

x10^-3

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63
Q

How to convert from ‘C to K?

A

+273

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64
Q

How to convert from KPa to Pa?

A

x10^3

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65
Q

Max number of electrons in the first shell?

A

2

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66
Q

Max number of electrons in the second shell?

A

8

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67
Q

Max number of electrons in the third shell?

A

18

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68
Q

Max number of electrons in the fourth shell?

A

32

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69
Q

What is an electron shell?

A

A group of atomic orbitals with the same principal quantum number, n

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70
Q

What’s the principal quantum number?

A

Represented as n, a number representing the overall energy level of the orbital. The bigger the number, the further the distance between the energy level and the atomic nucleus

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71
Q

Formula to work out how many electrons in the shell?

A

2(n^2)

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72
Q

What is an orbital?

A

A region of high probability within an atom that can hold 2 electrons with opposite spin

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73
Q

What are shells made up of?

A

Orbitals

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74
Q

What did Pauli discover in 1924?

A

.Orbitals only hold 2 electrons
.Electrons carry a negative charge
.Spin on-axis - generate a magnetic field
.Spin clockwise or anti-clockwise, represented by arrows
.Electrons in the same orbital must spin in different directions

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75
Q

What does an s-orbital look like?

A

A sphere

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76
Q

Which shells have s-orbitals?

A

From n=1 onwards, each shell contains one s-orbital (max two electrons)

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77
Q

What does a p-orbital look like?

A

Dumbbell shaped, like a balloon squashed in the middle

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78
Q

Which shells have p-orbitals?

A

From n=2 onwards, each shell contains three p-orbitals (max 6 electrons)

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79
Q

Which shells have d-orbitals?

A

From n=3 upwards, each shell has 5 d-orbitals (max 10 electrons)

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80
Q

Which shells have f-orbitals?

A

From n=4 onwards, each shell has 7 f-orbitals (max 14 electrons)

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81
Q

What does each part of the notation 1s^2 represent?

A 
1 = energy level
s = type of orbital
2 = number of electrons in orbital
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82
Q

What is Aufbau’s principle?

A

Electrons fill the lowest energy orbitals in sequence

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83
Q

What is the order that the orbitals fill up?

A

1s 2s 2p 3s 3p 4s 3d 4p 5s

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84
Q

What is an electron energy level made up of?

A

An electron energy level is made up of atomic orbitals with the same principal quantum number

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85
Q

What is a sub-shell?

A

Within each shell, orbitals of the same type are grouped together as a sub-shells
Each sub-shell is made up of only one type of orbital only, so there are s, p and d sub-shells

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86
Q

What is Hunds Rule

A

Electrons singly occupy orbitals before pairing up

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87
Q

When ionising, using the orbital model, which electrons are lost?

A

The ones in the outer shell, so even if 3d has electrons in it they will be lost from 4s when ionising

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88
Q

Which two elements don’t follow Aufbau’s principle?

A

Copper and Chromium

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89
Q

How does Copper not follow Aufbau’s principle?

A

copper steals a 4s electron to gain a full 3d orbital

It’s actually 1s2, 2s2, 2p6, 3s2, 3p6, 4s1, 3d10

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90
Q

How does chromium not follow Aufbau’s principle?

A

chromium steals a 4s electron to be able to put an electron in every 3d orbital
It’s actually 1s2, 2s2, 2p6, 3s2, 3p6, 4s1, 3d5

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91
Q

How do copper and chromium gain stable structures?

A

Cr and Cu get stable structures from full and half full 3d sub shells

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92
Q

What do the big numbers in equations show?

A

The molar ratio

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93
Q

What are atoms trying to achieve when they bond?

A

A full outer shell

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94
Q

4 key things about ionic bonding

A

Between metals and non-metals
Its to do with the loss and gain of electrons
Metals form positive charges
Non-metals form negative charges

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95
Q

How does an atom gain a positive charge?

A

It loses electrons, therefore there are more protons than electorns

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96
Q

How does an atom gain a nagative charge?

A

It gains electrons, therefore there are less protons than electrons

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97
Q

What is an ion?

A

A charged particle

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98
Q

How do ionic bonds stay together?

A

The electrostatic attraction between the positive and negative ions

99
Q

What do you use the crosses and dots for on ionic bonding diagrams?

A

use crosses for the metals electrons and dots for the non-metals electrons

100
Q

What is a lattice?

A

A lattice is a regular repeated three-dimensional arrangement of atoms, ions, or molecules in a metal or other crystalline solution.

101
Q

What is the structure of ionic compounds like?

A

A giant ionic lattice, the attraction between the oppositely charged ions acts equally in all directions, which leads to the formation of a giant ionic lattice in three dimensions

102
Q

Moles equation with concentration and volume

A

Mol = (cm^3 x moldm^-3) / 1000

103
Q

Rearrange the moles equation with concentration and volume for concentration =

A

moldm^-3 = (mol x 1000) / cm^3

104
Q

What are almost all ionic compounds at room temperature?

A

solids

105
Q

Why are almost all ionic compounds solid at room temperature?

A

At room temperature there is insufficient energy to overcome the strong electrostatic forces of attraction between the oppositely charged ions in the giant ionic lattice, high temperatures are needed to provide this energy

106
Q

Why is the melting points higher for ionic lattices containing ions with greater ionic charges? What else does ionic attraction depend on?

A

The melting points are higher for lattices containing ions with greater ionic charges, as there is a stronger attraction between ions
.The ionic attraction also depends on the size of the atom

107
Q

Melting point of NaF

A

993’C

108
Q

Melting point of CaF2

A

1423’C

109
Q

Melting point of Na2O

A

1275’C

110
Q

Melting point of CaO

A

2614

111
Q

What do many ionic compounds dissolve in?

A

Polar solvents like water

112
Q

Why might an ionic compound not be soluble?

A

.Polar water molecules break down the lattice and surround each ion in the solution
.But in a compound made of ions with large charges, the ionic attraction may be too strong for water to be able to break down the lattice structure, and the compound will then not be soluble

113
Q

Solubility of NaCL at 20’C

A

6.1 mol dm^-3

114
Q

Solubility of CaCl2 at 20’C

A

0.67 mol dm^-3

115
Q

Solubility of Na2CO3 at 20’C

A

2.0 mol dm^-3

116
Q

Solubility of CaCO3 at 20’C

A

1.3x10^-4 mol dm^-3

117
Q

What 2 processes does solubility require?

A

.The ionic lattice must be broken down

.The water molecules must attract and surround the ions

118
Q

What does the solubility of an ionic compound in water depend on?

A

on the relative strengths of the attractions within the giant ionic lattice and the attractions between ions and water molecules

119
Q

As the solubility decreases in ionic compounds…

A

ionic charge increases

120
Q

When does and when doesn’t an ionic compound conduct electricity?

A

.In the solid state, an ionic compound does not conduct electricity
.But once melted and dissolved in water the ionic compound does conduct electricity

121
Q

Describe an ionic compounds electric properties in solid state

A

.The ions are in a fixed position in the giant ionic lattice
.There are no mobile charge carriers, as the ions cannot move
.An ionic compound is a non-conductor of electricity in the solid state

122
Q

Describe an ionic compounds electric properties in a liquid or dissolved state

A

.The solid ionic lattice breaks down
.The ions are now free to move as mobile charge carriers
.An ionic compound is a conductor of electricity in liquid and aqueous states

123
Q

Summarise the properties of ionic compounds

A

Most ionic compounds –
.Have high melting and boiling points
.Tend to dissolve in polar solvents such as water
.Conduct electricity only in the liquid state or aqueous solutions

124
Q

What is the main component in teeth, bones and tooth enamel

A

hydroxyapatite, Ca5(PO4)3OH

125
Q

What allows tooth decay?

A

Acid conditions, from food, break down enamel and allow tooth decay

126
Q

What does saliva do?

A

Salvia helps to neutralise acidic food and also to replace ions

127
Q

How do fluoride ions help teeth?

A

.Fluoride ions help to replace lost ions by forming fluoropatite, Ca5(PO4)3F, which is stronger than hydroxyapatite and more resistant to acid conditions

128
Q

What do most toothpastes and some water sources contain?

A

Most toothpastes contain fluoride as sodium fluoride, your water may also contain fluoride depending on where you live

129
Q

What is the theoretical yield?

A

The maximum possible amount that can be made

130
Q

Why is the theoretical yield never achievable?

A

.The reactions may have not gone to completion
.Other reactions (side reactions) may have occurred
.Purification of the product may have resulted in loss of product

131
Q

Percentage yield % =

A

(actual yield mol / theoretical yield mol) x 100

132
Q

Describe the use of aqueous barium chloride in qualitative analysis

A

Test for sulphate ions

133
Q

What is atom economy?

A

A measure of the proportion of reactants included in the final useful product

134
Q

What happens in an ideal reaction, in terms of atom economy?

A

All reactant atoms end up within the useful product molecule, no waste produced!

135
Q

What do inefficient reactions have in terms of atom economy?

A

They are wasteful and have a low atom economy

136
Q

What do efficient reactions have in terms of atom economy?

A

.High atom economy
.Important for sustainable development
.Conserve natural resources and create less waste

137
Q

What is a bonding pair?

A

A pair of electrons in a covalent dot n cross diagram that are being used in the reaction

138
Q

What is a lone pair?

A

A pair of electrons in a covalent dot n cross diagram that are not being used in the reaction

139
Q

How is a sigma bond formed?

A

The head-on overlap of orbitals

140
Q

In hydrogen what happens in covalent bonding, in terms of orbitals

A

The two 1s orbitals overlap and they become a molecular orbital

141
Q

How is the octet rule broken in Boron trifluoride?

A

It becomes an electron deficient molecule

142
Q

How can the octet rule be broken in sulphur hexafluoride?

A

3rd shell of S can hold 18 electrons, octet rule can be broken

143
Q

What is NH4 ^+

A

Ammonium ion, a type of molecular ion

144
Q

What is a dative covalent bond?

A

A bond formed when both electrons in the share are donated by one atom

145
Q

For an exam question what should you do?

A

READ IT ALL AND DO EVERYTHING IT TELLS YOU

146
Q

Properties of simple covalent molecules

A

.Low melting and boiling points
.Weak intermolecular forces
.Not soluble in polar solvents – only in other non-polar liquids
.Molecules are not charged so they don’t conduct electricity (no mobile charges)
.Weak and soft when solid

147
Q

What does VSEPR theory stand for?

A 
.Valence
.Shell
.Electron
.Pair
.Repulsion
.Theory
148
Q

What is the shape of a molecule or ion determined by?

A

The shape of a molecule or ion is determined by the number of electron pairs in the outer shell of the central atom, more specifically the number of electron pairs repelling as far away from each other as possible (maximum repulsion).

149
Q

What is the valence shell?

A

The outer shell

150
Q

the bond angle of a linear molecule

A

180’

151
Q

the bond angle of a triangular planar shape

A

120’

152
Q

the bond angle of a tetrahedron

A

109.5’

153
Q

the bond angle of an octahedron

A

90’

154
Q

Molecular formula of hydrochloric acid

A

HCl

155
Q

Molecular formula of sulphuric acid

A

H2SO4

156
Q

Molecular formula of nitric acid

A

HNO3

157
Q

Molecular formula of ethanoic acid

A

CH3COOH

158
Q

Molecular formula of sodium carbonate

A

Na2CO3

159
Q

Molecular formula of sodium hydroxide

A

NaOH

160
Q

Molecular formula of pottasium hydroxide

A

KOH

161
Q

Molecular formula of phosphoric acid

A

H3PO4

162
Q

Define hydrogen ion

A

An ion with no electrons and 1 proton

163
Q

Define acid

A

In water an acid releases hydrogen ions into the solution

164
Q

Define strong acid

A

Releases all of its hydrogen ions into the solution and completely disassociates, this is irreversible

165
Q

Define weak acid

A

Only releases a small number of its hydrogen ions into the solution and partially disassociates, this is reversible

166
Q

How is sulphuric acid both a strong and weak acid?

A

Sulphuric acid is both strong and weak, it can form 2H+ and SO42- or H+ and HSO4-

167
Q

Define a base

A

A compound that neutralises an acid, by accepting a hydrogen ion to form a salt – for example ammonias, hydroxides, carbonates and metal oxides

168
Q

Define an alkali

A

A type of base that dissolves in water to form hydroxide ions

169
Q

Define a salt

A

The product of a reaction in which the H+ ions from the acid are replaced by the metal or ammonium ions

170
Q

Describe and explain the conductivity of sodium Na, chlorine Cl2, and sodium chloride NaCl (6 marks)

A

.Sodium is a metal and conducts as both a solid and liquid
.Free moving delocalised electrons to carry the charge
.Chlorine is a small simple molecule that does not conduct
.No free moving delocalised electrons to carry the charge
.Ionic bond between sodium and chlorine to form sodium chloride, conducts as molten or aqueous but not when solid
.Free moving ions when molten and aqueous so can conduct as they carry the charge

171
Q

Why is the shape of PCl5 special?

A

Trigonal bi-pyramid

Bond angles of 120’ and 90’

172
Q

Describe and explain the shape of a NH3 molecule

A

3 bonding pairs + 1 lone pair = 4 electron pairs
So repels to corners of tetrahedron
Don’t see the lone pair so therefore the shape is pyramidal

173
Q

Bond angle of NH3

A

107’

174
Q

Why do lone pairs of electrons have greater repulsion

A

not attracted by nuclei unlike bonding pairs

175
Q

What is the order of repulsion?

A

LP:LP > LP:BP > BP:BP

176
Q

By how much does one LP in a tetrahedron reduce the bond angle by?

A

2.5’

177
Q

Bond angle of H2O and why?

A

109.5 – 2.5 – 2.5 = 104.5’

2 x LP’s so (– 2.5’) 2 times to the bond angle on a tetrahedron (109.5)

178
Q

Why is CO2 linear?

A

C has 2 electron regions (2 BP’s) so it is linear

179
Q

Define electronegativity

A

a measure of the tendency of an atom to attract a bonding pair of electrons within a covalent bond.

180
Q

Where are the electrons positioned when both atoms in a bond are identical?

A

When the 2 atoms are identical, the electrons are shared equally and stay in the middle

181
Q

What did Linus Pauling invent and when?

A

Linus Pauling – invented the Pauling scale in 1932 which measures electronegativity on an atom

182
Q

Why does it get more reactive as you go down group 1?

A

More reactive as you go down the group
Less nuclear attraction to the outer electron
Easier to lose
The nucleus is being shielded by the electrons in-between
So the electron feels less nuclear attraction

183
Q

Across a period what happens to the electronegativity?

A

increases

184
Q

Down a group what happens to the electronegativity?

A

decreases

185
Q

Why does electonegativity increase across a period?

A

.Size roughly equal across period with same shielding by 2, 8
.Nuclear charge increases across period therefore greater nuclear attraction so electronegativity increases

186
Q

Why does electonegativity decrease down a group?

A

.Down group size of atom increases
.Shielding of outer shell increases
.This reduces the nuclear attraction and electronegativity

187
Q

In a bond, what happens if one atom is more electronegative than the other? Example.

A

.The more electronegative atom will have a greater share of the electrons
.Take H and Cl as an example
.H = 2,1, Cl = 3.0
.Cl more electronegative so the electrons are attracted more towards it

188
Q

What does delta +/- mean?

A

means it has a slight positive or negative charge now

189
Q

When H reacts with Cl what happens to its electron density?

A

.H has lost a small amount of electron density – becomes s+

190
Q

What type of molecule is H-Cl?

A

Polar

191
Q

Why is H2O a polar molecule?

A

the dipoles act in the same direction

192
Q

Why is Carbon Dioxide a non-polar molecule?

A

Symmetrical dipoles will cancel each other out, so CO2 is non-polar

193
Q

What atom gets the delta- charge?

A

The more electronegative atom will take the s- charge

194
Q

Can you predict bonds from eletronegativity?

A

Yes

195
Q

If there is no difference between electronegativity there is what sort of bond?

A

non-polar covalent bond

196
Q

If there is a small difference (0 to 1.8) between electronegativity there is what sort of bond?

A

polar covalent bond

197
Q

If there is a large difference (>1.8) between electronegativity there is what sort of bond?

A

ionic

198
Q

metal + hydrochloric acid –>

A

metal chloride + hydrogen

199
Q

Metal + sulphuric acid –>

A

metal sulphate + hydrogen

200
Q

Metal + nitric acid –>

A

Metal nitrate + hydrogen

201
Q

Metal + phosphoric acid –>

A

Metal phosphate + hydrogen

202
Q

Metal oxide + acid –>

A

Salt + water

203
Q

Metal hydroxide + acid –>

A

salt + water

204
Q

Metal carbonate + acid –>

A

salt + carbon dioxide + water

205
Q

Ammonia + acid –>

A

ammonium salt

206
Q

Where to observe the meniscu

A

eye level

207
Q

What is weighing by difference?

A

With weighting boat – used weighing boat = mass

208
Q

What is an intramolecular force?

A

strong bonds inside molecules

209
Q

What is an intermolecular force?

A

act between different molecules, weak compared to covalent bonds

210
Q

What are intermolecular forces caused by?

A

Intermolecular forces are caused by weak attractive forces between very small dipoles in different molecules

211
Q

3 types of intermolecular forces and their relative strength if ionic and covalent are 1000

A

Hydrogen Bonds 50
Dipole-dipole forces 10
London forces 1

212
Q

What is a dipole dipole force?

A

permanent

A weak attractive force between permanent dipoles in neighbouring polar molecules

213
Q

Explain HCl as an example of dipole dipole intermolecular force

A

Take HCL, H is s+ and Cl s-, these charges weakly attract opposite charges in neighbouring molecules

214
Q

In London forces as the number of electrons increase …

A

so does the strength of the London Forces due to larger instantaneous dipoles created.

215
Q

In terms of intermolecular forces what can and cant dissolve in water

A

If no permanent dipoles it can’t dissolve in water, London forces don’t count as they are temporary dipoles

216
Q

How do London forces work?

A

.Movement of electrons produces a changing dipole in a molecule
.At any instant, an instantaneous dipole will exist, but its position is constantly shifting
.The instantaneous dipole induces a dipole on a neighbouring molecule
.This induced dipole induces further dipoles on neighbouring molecules, which then attract one another

217
Q

How to find uncertainty?

A

Half the smallest possible unit you can make

218
Q

Percentage uncertainty =

A

(uncertainty of equipment/size of measurement) x 100

219
Q

What anomalous properties does water have?

A 
Elevated melting and boiling points 
Ice floats (solid is less dense than the liquid)
relatively high surface tension and viscosity
220
Q

Why does water have an elevated melting and boiling point?

A

due to strong H bonds between molecules

221
Q

Why does ice float?

A

as liquid water freezes the molecules arrange themselves in a more open, regular lattice structure

222
Q

What is a hydrogen bond?

A

A strong dipole-dipole attraction between
.An electron-deficient hydrogen atom on one molecule (O-HS+ or N-HS+)
and
.A lone-pair of electrons on a highly electronegative atom on a different molecule (H-OS- or H-NS-)

223
Q

Why does water have relatively high surface tension and viscosity?

A

The hydrogen bonds

224
Q

In what compounds are hydrogen bonds useful?

A

organic compounds containing O-H and N-H bonds (alcohols, carboxylic acid etc)

225
Q

What are hydrogen bonds responsible for the shape of?

A

They are responsible for shape of proteins and even DNA

226
Q

New concentration equation

A

New conc = (original volume / new volume) * original conc

227
Q

Define oxidation number

A

A measure of the number of electrons that an atom uses t bond with an atom of another element. Oxidation numbers are derived from a set of rules.
Also known as an element’s oxidation state

228
Q

What is the oxidation number of a neutral element?

A

zero

229
Q

What is the oxidation number of a monatomic ion?

A

The same as the charge of the ion

230
Q

What is the sum of all oxidation numbers in a neutral compound?

A

zero

231
Q

What is the sum of all oxidation numbers in a polyatomic ion equal to?

A

the charge on the ion

232
Q

What is the oxidation number of groups 1-3 in compounds

A

In compounds the elements of:
Group 1 have an oxidation number of +1
Group 2 have an oxidation number of +2
Group 3 have an oxidation number of +3

233
Q

What is the oxidation state of hydrogen in a compound? What if it is in a binary metal hydride? What is a binary metal hydride?

A

The oxidation state of hydrogen in a compound is usually +1. If the hydrogen ion is part of a binary metal hydride (compound of metal and some hydrogen), then the oxidation state is -1.

234
Q

What is fluorines oxidation number in compounds?

A

Fluorine has an oxidation number of -1 in compounds.

235
Q

What is the oxidation state of oxygen usually? What is its oxidation state in a peroxide? What is its oxidation state when bonded to fluorine?

A

The oxidation number of oxygen in a compound is usually -2. If, however, the oxygen is in a class of compounds called peroxides (like hydrogen peroxide), then the oxygen has n oxidation number of -1. If the oxygen is bonded to fluorine, the number is +2.

236
Q

What happens with transition metals oxidation numbers?

A

In transition metals, the oxidation number can vary.

237
Q

What are chlorine, bromine, and iodines oxidation numbers? What is the exception?

A

Chlorine, Bromine and iodine usually have numbers of -1, except when it is in a compound with oxygen.

238
Q

In compound names with roman numerals, what does the roman numeral refer to?

A

The oxidation number of the element that is joined in name to the oxygen, but not the oxygen (for example in element(1) element(2)ate (Roman Numeral) then the roman numeral refers to the element (2))

239
Q

In potassium manganate(VII), what does the VII refer to?

A

The oxidation number of manganese

240
Q

In equations, if an elements oxidation state increases what has it undergone?

A

Oxidation

241
Q

In equations, if an elements oxidation state decreases what has it undergone?

A

Reduction

242
Q

What is a disproportionation reaction?

A

when the same element is both oxidised and reduced at the same time in a reaction

243
Q

Rules for writing ionic equations

A

1) Re-write out the equation showing the separate ions present in each substance.
2) Look for ions that change state and/or form non-ionic compounds and underline them
3) Cross out spectator ions
4) Rewrite the equation leaving out the spectator ions.
5) Make sure your equation still balances

244
Q

How does developing chemical processes with high atom economy benefit sustainability?

A

Less environmental waste as a high percentage of atoms are being used in useful products

OCR A Level Chemistry (17 decks)

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