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OCR AS (A-level) Chemistry > Module 2 Foundations in Chemistry > Flashcards

Module 2 Foundations in Chemistry Flashcards

1
Q

The mass of an electron

A

1/1836 the mass of a proton

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2
Q

Isotopes

A

Atoms of the SAME element with the same number of protons but a DIFFERENT number of neutrons

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3
Q

Relative Isotopic Mass

A

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

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4
Q

Relative Atomic Mass

A

The weighted mean mass of an atom of an element relative to 1/12th of the mass of an atom of carbon-12

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5
Q

1 mole (mol)

A

The amount of substance containing 6.02 x 10^23 number of particles

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6
Q

Equation for the amount of substance

A

Moles (n) = mass (m)/Mr (M)

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7
Q

Number of particles (Av Constant)

A

Number of particles = Na (Avogardro’s constant) x moles

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8
Q

The empirical formula

A

The simplest whole number ratio of atoms of each element present in a compound

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9
Q

The complete combustion equation

A

CxHy + O2 —> 1/2yH2O + XCO2

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10
Q

Molecular formula

A

The actual number of atoms of each element in a molecule

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11
Q

Stoichiometry

A

Ratios former by the balancing number to determine the number of moles

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12
Q

Concentration

A

This gives an accurate indication of the amount, in moles, of a solute that is dissolved in a volume of the solvent

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13
Q

Standard solution

A

Solution of known concentration

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14
Q

How do you convert between g dm-3 to mol dm-3

A

Divide by the Mr

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15
Q

Which equation do we use to find the concentration of a solution when the volume contains a dm3 unit

A

Amount (n) = C x V

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16
Q

State the equation used when calculating the moles of a solution when our volume of solution value is in cm3

A

Amount (n) = C x V/1000

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17
Q

State what 1cm3 is equal to

A

1ml

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18
Q

State what 1dm3 is equal to

A

1 litre

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19
Q

How do we convert between cm3 and dm3

A

Divide by 1000

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20
Q

Solution

A

In a solution, a solute is dissolved into a solvent

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21
Q

Concentrated solutions contain…

A

A large amount of solute in 1dm3

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22
Q

Dilute solutions contain…

A

A small amount of solute in 1dm3

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23
Q

Formula for atom economy

A

Mr of desired product/Total Mr of all products x 100

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24
Q

Formula for % yield

A

Actual yield/Theoretical yield x 100

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25
Q

When writing a dissociation equation for a weak acid, which type of arrow is used

A

Reversible

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26
Q

Definition of Acids

A

Proton donors

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27
Q

Strong acids________________dissociate all H+ ions in ________________ _______________

A

Completely
Aqueous solution

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28
Q

Which proton is lost of the carboxyl group in the dissociation of a weak acid

A

The most acidic proton

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29
Q

Formula for ethanoic acid

A

CH3COOH

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30
Q

How do you write out the dissociation of a strong acid?

A

Separate into constituent ions

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31
Q

How do you write an ionic equation?

A
  1. Identify all ionic compounds excluding liquids and gases (acids ionise)
  2. Separate into their constituent ions
  3. Rewrite equation
  4. Remove spectator ions
  5. Rewrite equation
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32
Q

Definition of a base

A

Proton acceptors

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33
Q

4 types of bases?

A

Metal oxides, hydroxides, carbonates and ammonia

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34
Q

Alkalis have ____ in them

A

OH-

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35
Q

In neutralisation, what happens?

A

H+ ions react with a base to form salt and water

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36
Q

When is a salt produced

A

When the H+ ions react of an acid is replaced by a metal ion or ammonia

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37
Q

Acid + metal oxide ——->

A

Salt + water

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38
Q

Acid + metal hydroxide ——->

A

Salt + water

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39
Q

Acid + ammonia ——->

A

Ammonium salt

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40
Q

Acid + metal carbonate ——->

A

Salt + water + carbon dioxide

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41
Q

Definition of a titration

A

A technique to accurately measure the volume of one solution that reacts with another

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42
Q

Describe how to prepare a standard solution

A

See notes

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43
Q

What is a known and unknown solution

A

Known - when we know the concentration, unknown - when we don’t know the concentration

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44
Q

Which acids are weak

A

Carboxyl - defined by ‘oic’

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45
Q

What is an orbital

A

Region within nucleus of an atom which holds 2 electrons with opposite spins

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46
Q

State the 4 subshells

A

S,P,D,F

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47
Q

State how many electrons each sub-shell can hold

A

S=2
P=6
D=10
F=14

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48
Q

State the electron configuration of oxygen

A

1s2 2s2 2p4

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49
Q

Describe how to work out the shorthand electron configuration

A

Take the final electron configuration of atom, travel along periodic table to the noble gases, go up one element and use the selected noble gases symbol, and just place the final configuration next to it

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50
Q

What shape is an S orbital

A

Spherical

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51
Q

What shape is a P orbital

A

Dunbell

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52
Q

Why does 4s2 fill before 3d10

A

Order of increasing energy level

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53
Q

Define an alkali

A

Bases that can dissolve in water and release OH- ions

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54
Q

Give an example of a strong acid

A

Anything that isn’t a carboxyl acid

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55
Q

Why do we rinse the beaker 3 times when preparing a standard solution

A

So all the solution is transferred to the volumetric flask, otherwise, the concentration would diluet

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56
Q

What will happen if the meniscus goes above/below the graduation mark

A

Too much/little water added affecting concentration

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57
Q

What is the ‘mass by difference method’

A

Full weighing boat - empty weighing boat = final accurate mass

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58
Q

Describe how to set out reacting masses calculation

A

Table should include…
1. Molar ratio
2. Mass
3. Mr
4. Moles

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59
Q

Which equipment is involved in titrations

A
  • Burette
  • Conical flask
  • Bulb pippete
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60
Q

When calculating mean titre, which values do we include?

A

Concordant - within 0.1 of each other

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61
Q

Phenolphthalein turns what in alkali?

A

Pink

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62
Q

Phenolphthalein turns what in acid?

A

Clear

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63
Q

Methyl orange turns what in acid?

A

Orange

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64
Q

Methyl orange turns what in alkali?

A

Yellow

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65
Q

True or false

The rough titre is included in the mean tire calculation

A

False

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66
Q

Oxidation number

A

The number of electrons involved in bonding to a different element

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67
Q

What is the oxidation number equal to?

A

The charge

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68
Q

When calculating the oxidation number, which element do we always begin with?

A

The most electronegative

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69
Q

Uncombined elements have an oxidation number of?

A

0

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70
Q

Oxidation number of oxygen?

A

2-

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71
Q

Oxidation number of hydrogen?

A

1+

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72
Q

Oxidation number of hydride?

A

1-

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73
Q

Explain what the roman numerals in a compound mean

A

The positive charge of the ion and its oxidation state

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74
Q

What is an oxyanion?

A

A negative ion containing an oxide

75
Q

Oxyanions always end in which suffix?

A

‘Ate’

76
Q

Using oxidation states, state the formula of chlorate (VII)

A
  • Oxygen and chlorine present, so ClO
  • Cl has a charge of 7+
  • Use the minimum amount of oxidation numbers for oxygen to make it negative
  • In this case, it will be 4
  • So, we used 4 oxygens, meaning a small four will go next to the oxygen and it will have a negative charge

ClO4-

77
Q

Oxidation is when electrons are

A

Lost and the oxidation number increases

78
Q

Reduction is when electrons are

A

Gained, and the oxidation number decreases

79
Q

How can you work out oxidising or reducing agent?

A

Whatever is oxidised in RHS of equation, the same atom on the LHS will be the reducing agent

(Same for oxidising agent)

80
Q

Hydrated

A

When water of crystallisation is present in a compound

81
Q

Anhydrous

A

When all water of crystallisation has been removed from a compound

82
Q

Water of crystallisation

A

When water is present in a compound giving a crystalline appearance (.xH2O)

83
Q

Redox reaction

A

When oxidation and reduction take place simultaneously (OILRIG)

84
Q

Oxidation is when the oxidation number __________, and electrons are _________

A
  1. Increases
  2. Lost
85
Q

Reduction is where the oxidation number _______________ and electrons are ________________

A
  1. Decreases
  2. Lost
86
Q

Explain how to identify the oxidising and reducing agents

A

Using oxidation numbers, whatever is oxidised on the RHS, the same atom on the LHS will be the reducing agents.

Converse for oxidising agent

87
Q

How to tell if a compound is oxidised?

A

Oxidation number increases

88
Q

When writing a half equation for oxidation, the electron number is added on to which side of the equation?

A

Right hand side

89
Q

When writing a half equation for reduction, the electron number is added on to which side of the equation?

A

Left hand side

90
Q

Disproportionation reaction

A

The same element is is oxidised and reduced

91
Q

What constructing difficult half equations, which 4 steps are used?

A
  1. Balance the element being oxidised or reduced (same number of the element on each side)
  2. Write out oxidation numbers, if reduction, electrons go on RHS, if oxidation, LHS
  3. Balance oxidation numbers using electrons e.g., 7+ to 2+, 5e- needed
  4. Balance the charges of each side using H+ or OH-
  5. Use water to balance the H and O on the side where they aren’t present
92
Q

When constructing half equations in acidic conditions, how should the charges be balanced?

A

Using H+ ions

93
Q

When constructing half equations in alkaline conditions, how should the charges be balanced?

A

Using OH- ions

94
Q

All atoms aim to be like….

A

The noble gases

95
Q

The octet rule

A

Atoms of elements chemically combine with each other so each atom has 8 outer electrons. This is achieved by ionic/covalent bonding

96
Q

Ionic bonding is between….

A

Metals and non metals

97
Q

Definition of ionic bonding

A

Electrostatic force of attraction between oppositely charged ions

98
Q

The net charge of any ionic compound is…

A

0

99
Q

Structure of ionic compound

A
  • Giant ionic lattice
  • 6:6 coordination
  • Regular repeated arrangement of ions
100
Q

Melting/boiling point of ionic compound

A

High, strong electrostatic forces of attraction between oppositely charged ions in all directions of the lattice - take lots of thermal energy to break

101
Q

Electrical conductivity of ionic compounds

A
  • Non-electrical conductors when solid, ions fixed in place, can’t move and carry charge
  • Electrical conductors when molten/aqueous, good electrical conductors, ions free to move and carry charge
102
Q

Solubility of ionic compounds

A

Soluble in water and polar solvents
- Giant ionic lattice breaks down
- Water molecules surround each ion in solution

103
Q

Covalent bonding is between….

A

Non metals

104
Q

Define covalent bonding

A

Strong electrostatic forces of attraction between shared pair of electrons and the nuclei of bonded atoms

105
Q

A covalent bond is an ___________ of atomic ___________

A
  1. Overlap
  2. Orbitals
106
Q

What is meant by the term ‘localised’ in covalent bonding?

A

Attraction is localised, acting solely between shared pair of electrons ad nuclei of 2 bonded atoms resulting in a molecule

107
Q

Define ‘lone pair’ in relation to electrons

A

Atoms not bonded covalently

108
Q

Draw out a triple bond

A

3 dots, 3 crosses

109
Q

Which bond is present in ethene?

A

Double bond (alkenes have double bonds)

110
Q

True or false

Sulfur must have a full octet

A

False

111
Q

Dative covalent bonding definition

A

Where only one of the atoms applies both of the electrons shared in covalent bonding

112
Q

When adding a dative covalent bond in a displayed formula, how is this represented?

A

Through an arrow symbol

113
Q

Melting and boiling point of simple covalent compound

A
  • Low
  • Little thermal energy needed to overcome weak intermolecular forces between molecules
114
Q

Electrical conductivity of simple covalent compound

A
  • Non electrical conductors
  • No mobile electrons or ions
115
Q

Solubility of simple covalent compounds

A
  • Soluble in non-polar solvents e.g., hexane
  • Forms induced dipole-dipole forces with these solvents
116
Q

List examples of giant covalent compounds

A
  1. Diamond
  2. Graphite
  3. Buckminster fullerene
  4. Silicon
  5. Silicon dioxide
117
Q

Melting point and boiling point of diamond

A
  • High
  • Strong covalent bonds between carbon atoms
  • Lots of thermal energy needed to break covalent bonds
118
Q

Melting point and boiling point of graphite

A
  • High
  • Strong covalent bonds between carbon atoms
  • Lots of thermal energy needed to break covalent bonds
119
Q

Electrical conductivity of diamond

A
  • Non conductors
  • No mobile electrons or ions
120
Q

Electrical conductivity of graphite

A
  • Electrical conductor
  • Delocalised electrons can carry and move charge
121
Q

Solubility of diamond

A

Insoluble in water

122
Q

When adding a dative covalent bond onto a dot and cross diagram, where should this be added?

A

To a lone pair of electrons

123
Q

Definition of metallic bonding

A

Strong electrostatic force of attraction between positively charged ions and delocalised electrons

124
Q

Structure of metallic compound

A

Giant metallic lattice

125
Q

Draw the structure of a metallic compound

A

Should include positive metals ions with an electron surrounding each one

126
Q

Melting and boiling point of metallic compound

A
  • High
  • Strong electrostatic forces of attraction between positively charged ions and delocalised electrons
  • Lots of thermal energy to break
127
Q

Electrical conductivity of metallic compounds

A
  • Good conductors
  • Delocalised electrons can move and carry charge
128
Q

Solubility of metallic compounds

A

Insoluble in water

129
Q

Different isotopes have similar chemical properties. Explain why (1)

A

Same number of outer electrons
Reject same number of electrons/protons or different number of neutrons

130
Q

Solid aluminium fluoride has a giant ionic lattice structure.

Describe what is meant by the term ionic lattice, in terms of type and arrangement of particles present (2)

A

Repeating pattern (1)

of oppositely charged ions (1)

ALLOW ‘regular’ OR ‘alternating’ OR ‘uniform (arrangement)’ for ‘repeating pattern’
ALLOW positive and negative ions OR aluminium ions and fluoride ions
ALLOW oppositely charged ions from a labelled diagram

131
Q

If a molecule has a larger average bond enthalpy, what does this mean?

A

It has a stronger bond strength

132
Q

Explain the electron pair repulsion theory (2)

A
  1. Bonded pairs repel each other to get as far apart as possible
  2. Lone pairs repel **more strongly/further/greater* than bonding pairs
133
Q

Explain the shape of molecule in a methane atom (3)

A
  1. Tetrahedral shape with a bond angle of 109.5 degrees
  2. The 4 bonding pairs repel each other (as far as possible) equally
134
Q

How much does each lone pairs reduce the bond angle by?

A

2.5 degrees

135
Q

2 bonding pairs, 0 lone pairs

A

Linear, 180 degrees

136
Q

3 bonding pairs, 0 lone pairs

A

Trigonal planar, 120 degrees

137
Q

4 bonding pairs, 0 lone pairs

A

Tetrahedral, 109.5 degrees

138
Q

6 bonding pairs, 0 lone pairs

A

Octahedral, 90 degrees

139
Q

3 bonding pairs, 1 lone pair

A

Trigonal pyramidal, 107 degrees

140
Q

2 bonding pairs, 2 lone pairs

A

Bent, 104.5 degrees

141
Q

State the definition of electronegativity (2)

A
  1. The ability of an atom to attract electrons
  2. In a covalent bond
142
Q

Which elements have the strongest electronegativity?

A

F, O, N, Cl

143
Q

When one atom is more electronegative than the other, it forms a ____________________ ___________

A

Permanent dipole

144
Q

Why do electrons sit closer to the more electronegative atom?

A

The electron cloud is more dense than

OR

The electron density is higher

145
Q

What does the delta symbol (δ) represent

A

A slight charge

146
Q

Definition of a permanent dipole

A

Small difference in charge across a bond as a result of differing electronegativities

147
Q

What is the name of a molecule where a permanent dipole is present?

A

Polar (molecule)

148
Q

Definition of non-polar molecules

A

Two atoms with the same electronegativities

149
Q

Large difference in electronegativity

A

Ionic bonding

150
Q

Medium to small difference in electronegativity

A

Polar covalent bond

151
Q

No difference in electronegativity

A

Non-polar covalent bond

152
Q

Explain Pauling’s electronegativity values (2)

A
  1. (Across periodic table) as nuclear charge
  2. (and) atomic radius increases, electronegativity increases
153
Q

True or false: Symmetrical molecules are non-polar

A

True

154
Q

In symmetrical molecules, all the dipoles __________ ____

A

Cancel out

155
Q

True or false:In unsymmetrical molecules, the dipoles cancel out

A

True

156
Q

Why is a molecule containing lone pairs not symmetrical?

A

Because the dipoles aren’t exactly the same

157
Q

Explain the solubility of the polar molecule NaCl (3)

A
  1. Water molecules attract the Na+ and Cl- ions
  2. Ionic lattice breaks down as it dissolves
  3. (In resulting solution) water molecules surround the Na+ and Cl- ions
158
Q

Where do IM forces exist?

A

Between covalently bonded atoms

159
Q

Explain how permanent dipole-dipole forces arise

A
  • Occurs between molecules containing a permeant dipole
  • A positive ion from one molecule will align with a negative ion from another molecule
  • This forms an IM force
160
Q

Which IM force is strongest?

A

Hydrogen bonds

161
Q

Where do hydrogen bonds exist?

A

Between molecules with H directly bonded to F, O or N and the H atom of —NH, —OH and HF

162
Q

Explain how induced dipole-dipole forces (London forces) arise

A
  • Electrons constantly moving and constantly changing electron density
  • Electrons aren’t always symmetrical around molecule, sometimes imbalanced forming instantaneous dipole
  • This induces a dipole in the neighbouring molecule forming an attraction
163
Q

Where do London forces occur?

A

Between non-polar molecules

164
Q

Which IM force is the weakest?

A

Induced dipole-dipole forces (London forces)

165
Q

Explain the strength of London forces (3)

A
  • The larger the instantaneous dipole and induced dipole
  • The greater the induced dipole-dipole interactions
  • The stringer the attractive forces between the molecule
166
Q

List the 2 factors affecting electronegativity

A
  1. Increases across a period as the number of protons increases and atomic radius decreases (because electrons in the same shell are pulled in more)
  2. Decreases down a group because the distance between the nucleus and outer electrons increases and the shielding of inner shell increases
167
Q

When going down group 7, why do boiling points increase?

A
  • Increasing number of electron (because they’re bigger molecules)
  • This increases the size of the induced dipole-dipole interactions between the molecules
168
Q

Explain why BF3 is a non-polar molecule (3)

A
  1. Symmetrical
  2. All dipoles cancel out
  3. Fluorine is the most electronegative
169
Q

All non-polar molecule will always contain ______________ forces

A

London

170
Q

Anytime a lone pair of electrons is present, the molecule is always _______________

A

Polar

171
Q

Which two atoms are always δ+

A

Carbon and hydrogen

172
Q

Despite having the same type of intermolecular forces of attraction, ammonia has a lower boiling point that water.

Explain why (2)

A
  1. Hydrogen bonds between water molecules are stronger (than ammonia);
     *Allow H-bonds*
  2. Because oxygen is more electronegative (than nitrogen) and has two lone pairs (whereas ammonia only has one);
173
Q

Ammonia has hydrogen bonding, bromine has London forces.

Explain why bromine’s boiling point is higher than ammonias (1)

A

London forces/induced dipole-dipole/dispersion forces/Van der Waals forces between Br molecules are stronger than the hydrogen bonds between the ammonia molecule.

DO NOT NEED TO EXPLAIN WHY

174
Q

When describing the type of IM forces in a molecule, which 2 elements must be included in the statement?

A
  1. Type of molecule e.g., polar
  2. Bonding occurs between molecules
175
Q

Why does water have anomalous properties?

A

Due to hydrogen bonds between molecules

176
Q

List the 2 anomalous properties of water

A
  1. Ice is less dense than water
  2. Water has a relatively high melting and boiling point
177
Q

Explain why ice is less dense than water (2)

A
  1. Water molecules held further apart by hydrogen bonds between
  2. Giving open lattice structure
178
Q

Explain why water has a relatively high melting and boiling point (2)

A
  1. H-bonds relatively strong (stronger than other IM forces)
  2. More thermal energy needed to break
179
Q

Suggest why ice has a higher melting and boiling point than solid ammonia (4)

A
  1. H-bonds between water molecules stronger than H-bonds between ammonia molecules
  2. (Because) O is more electronegative than N
  3. (So) the dipole across the O—H bond is stronger than across N—H bond (because difference in electronegativity is greater in O—H, so, stronger dipole)
  4. (Also) there are 2 hydrogen bonds between water molecules and only 1 between ammonia molecules
180
Q

List the 2 factors that make H-bonds stronger

A
  1. Amount of H-bonds
  2. Greater difference in electronegativity leading to a stronger dipole
181
Q

Why does water have 2 H-bond whereas ammonia only has 1?

A

Water has 2 lone pairs whereas ammonia only has 1 (H-bonds form from lone pairs)

182
Q

Explain the boiling points of the 3 substances listed below.
- H2O = 100
- CHCl = 62
- Cl2 = -34

A
  1. Mentions that all are simple covalent structures
  2. Water is polar, H-bond between water molecules
  3. Trichloromethane is polar, permanent dipole-dipole interactions between molecules
  4. Chlorine is non-polar, London forces between molecules
  5. H-bonds strongest, most thermal energy needed, London forces weakest, least thermal energy to overcome, permenant dipole-dipole have intermediate strength
183
Q

Phosphorus, P4 has a higher melting point than chlorine, Cl2.

Explain this difference (3)

A
  1. Phosphorus has more electrons
  2. Stronger London forces

OR

Stronger induced dipole(-dipole) interactions

  1. More energy required to break the intermolecular forces / bonds OR London forces

Remember - London forces are affected by number of electrons, not electronegativity like H-bonds

184
Q

Explain the boiling points of the molecules below (5)
- Ammonia = -33
- Fluorine = -188
- Bromine = 59

A
  1. Ammonia has H-bonding
  2. Fluorine and bromine have London forces
  3. Forces of attraction are between molecules (or IMF)
    Statement may read: Ammonia has H-bonding between molecules
  4. The London forces in bromine are greater than in fluorine because bromine has more electrons than fluorine
  5. Bromine has a higher boiling point than ammonia because the London forces are stronger than the hydrogen bonds and hydrogen bonding in ammonia is stronger than the London forces in fluorine

OCR AS (A-level) Chemistry (3 decks)

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