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AQA: Chemistry Paper 1 > 3.1.3 Bonding > Flashcards

3.1.3 Bonding Flashcards

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1
Q

Ionic Bonding [definition]:

A

The electrostatic attraction between oppositely charged ions in a lattice

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2
Q

Covalent Bonding [definition]:

A

A shared pair of e⁻s with opposite spins

- one e⁻ donated by each atom

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3
Q

Sulphate formula:

A

SO₄²-

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4
Q

Hydroxide formula:

A

OH−

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5
Q

Nitrate formula:

A

NO₃⁻

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6
Q

Giant Ionic lattice [2]:

A
  • Sodium Chloride

- Magnesium oxide

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7
Q

Covalent simple molecular structures [5]:

A
  • Iodine
  • Ice
  • Carbon dioxide
  • Water
  • Methane
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8
Q

Giant Covalent structures [4]:

A
  • Diamond
  • Graphite
  • Silicon dioxide
  • Silicon
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9
Q

Ionic boiling and melting points [2]:

A
  • high- because of giant lattice of ions with strong
    electrostatic forces between oppositely charged ions
  • requires lot of energy to break
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10
Q

Ionic Solubility in water:

A

Generally good

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11
Q

Ionic compound conductivity (when solid):

[2]

A
  • poor

- ions can’t move/ fixed in lattice

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12
Q

Ionic compound conductivity (when molten):

[2]

A
  • Good

- Ions can move freely

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13
Q

Co-ordinate/ Dative Covalent bond [definition]:

A

A bond that contains a shared pair of electrons with both electrons supplied by one atom.

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14
Q

Simple Covalent compound boiling and melting points:

A

low- because of weak intermolecular forces between molecules (specify type e.g van der
waals/hydrogen bond)

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15
Q

Simple covalent compound solubility in water:

A

Generally poor

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16
Q

Simple covalent compound conductivity (when solid):

[2]

A
  • poor
  • no ions to conduct and electrons are
    localised (fixed in place)
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17
Q

Simple covalent compound conductivity (when molten):

[2]:

A
  • Poor

- No ions

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18
Q

Giant Covalent compound boiling and melting points:

A

High- because of many strong covalent

bonds in macromolecular structure (requires a lot of energy to break many strong bonds)

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19
Q

Giant covalent compound solubility in water:

A

Insoluble

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20
Q

Giant covalent compound conductivity (when solid):

[2]:

A
  • diamond and sand = poor because electrons are localised and can’t move
  • graphite = good, has free delocalised electrons between layers to conduct
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21
Q

Giant covalent compound conductivity (when molten):

A

Poor

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22
Q

Metallic bonding boiling and melting points:

A

high- strong electrostatic forces between positive ions and sea of delocalised electrons

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23
Q

Metallic bonding solubility in water:

A

Insoluble

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24
Q

Metallic bonding conductivity (when solid):

[2]:

A
  • good

- delocalised electrons can move through structure

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25
Q

Metallic bonding conductivity (when molten):

A

Good

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26
Q

Linear Shape [3]:

A
  • BP:2
  • LP: 0
  • Bond angle: 180
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27
Q

Trigonal Planar Shape [3]:

A
  • BP: 3
  • LP: 0
  • Bond angle: 120
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28
Q

Tetrahedral Shape [3]:

A
  • BP: 4
  • LP: 0
  • Bond angle: 109.5
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29
Q

Bent Shape [3]:

A
  • BP: 2
  • LP: 2
  • Bond angle: 104.5
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30
Q

Trigonal Pyramidal Shape [3]:

A
  • BP: 3
  • LP: 1
  • Bond angle: 107
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31
Q

Trigonal Bipyramidal Shape [3]:

A
  • BP: 5
  • LP: 0
  • Bond angle: 120 and 90
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32
Q

Octahedral Shape [3]:

A
  • BP: 6
  • LP: 0
  • Bond angle: 90
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33
Q

Electronegativity [definition]:

A

Electronegativity is the relative tendency of an atom in a covalent bond to attract electrons in a covalent bond to itself

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34
Q

What are the most electronegative atoms? [3]:

A
  • Fluorine
  • Oxygen
  • Nitrogen
    [F, O, N]
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35
Q

Factors affecting electronegativity [2]:

A
  • increases across a period as the number of protons increases and the atomic radius decreases cus electrons in the same shell are pulled in more
  • decreases down a group cus the distance between nucleus and outer electrons increases + shielding of inner shell electrons increases
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36
Q

Purely covalent compound [definition]:

A

A compound containing elements of similar electronegativity and hence a small electronegativity difference will be purely covalent

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37
Q

Ionic compound [definition]:

A

A compound containing elements of very different electronegativity and hence a
very large electronegativity difference (> 1.7) will be ionic

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38
Q

When does a polar covalent bond form?

A

when the elements in the bond have different

electronegativities. (Of around 0.3 to 1.7)

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39
Q

What is a polar covalent bond?

A

When a bond is a polar covalent bond it has an unequal distribution of electrons in the bond and produces a charge separation, (dipole) δ+ δ- ends.

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40
Q

Symmetrical molecule [3]:

A
  • All bonds identical
  • No lone pairs
  • Will not be polar even if individual bonds within the molecule are polar
    (The individual dipoles on the bonds ‘cancel out’ due to the symmetrical shape of the molecule.
    There is no NET dipole moment: the molecule is
    NON POLAR)
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41
Q

Where do Van der Waal forces occur? [2]:

A
  • Occur between all molecular substances and noble gases

- They do not occur in ionic substances

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42
Q

The main factor affecting the size of Van der Waals [2]:

A
  • The more electrons there are in the molecule the higher the chance that temporary dipoles will form
  • This makes the Van der Waals stronger between the molecules and so boiling points will be greater
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43
Q

What is the weakes IMF?

A

Induced dipole/ Van der Waal’s

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44
Q

When do ions form?

A

When e⁻s are transferred between elements that have a large diff in negativity (metal + non-metal)

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45
Q

How are ionic lattices structured?

A

Regular lattice with alternating +ive & -ve ions

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46
Q

What holds an ionic lattice together?

A

Very strong electrostatic forces

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47
Q

What are the properties of an ionic crystal structure? [5]:

A
  • V high melting points
  • Generally soluble in water
  • Electrical insulators when solid
  • Electrical conductors when molten
  • Brittle
48
Q

Why are ionic compounds generally soluble in water? [3]:

A
  • Water is v polar
  • It disrupts electrostatic attractions between ions & breaks them up
  • causes the compound to dissolve
49
Q

Why can ionic compounds NOT conduct electricity when solid? [2]:

A
  • Ions are in a fixed position

- So canny carry a charge

50
Q

Why can ionic compounds conduct electricity when molten? [2]:

A
  • Ions are free to move

- so can carry a charge

51
Q

Why are ionic compounds brittle? [2]:

A
  • When moved ions no longer have alternating arrangement

- so lattice breaks cus same charge ions repel each other

52
Q

When do Covalent bonds occur?

A

Form between elements that have high electronegativity values (non-metals)

53
Q

How many pairs of electrons are in a double bond?

A

2 pairs of electrons

54
Q

How many pairs of electrons are in a triple bond?

A

3 pairs of electrons

55
Q

What are the two covalent structures?

A
  • (simple) molecular

- Macromolecular (giant covalent)

56
Q

What are sum examples of simple covalent molecules? [3]:

A
  • CH₄
  • H₂O
  • H₂
57
Q

What are sum examples of macromolecular covalent structures? [2]:

A
  • Diamond

- Graphite

58
Q

When is a coordinate bond formed?

A

when BOTH electrons are donated by the same atom

59
Q

Coordinate bond =

A

dative covalent

60
Q

What is an example of a coordinate bond?

A

NH₄⁺

61
Q

How do u draw a coordinate bond?

A

with an arrow

62
Q

What type of covalent strucure is ice?

A

(simple) molecular crystal

63
Q

Ice between 0-100 degrees Celsius (water) [2]:

A
  • Hydrogen bonding exists between molecules

- H-bonds hold molecules close together but allow them to move freely (liquid)

64
Q

Ice (at 0 degrees celsius)- properties [4]:

A
  • Molecules have less energy so H-bonds FIX molecules in position (solid)
  • Ice has a 3D hexagonal structure
  • Floats in water
  • Relatively high melting point
65
Q

What causes ice to be able to float in water? [2]:

A
  • Spaces created between molecules, causing it to expand

- So it becomes less dense (can float)

66
Q

What causes the relatively high melting point of ice? [2]:

A
  • Strong hydrogen bonds between molecules

- Requires more energy to break

67
Q

Properties of iodine as a molecular crystal structure [4]:

A
  • Low melting/ boiling point
  • Sublimes into purple gas
  • Slightly soluble in water
  • Shiny grey solid at room temp
68
Q

Molecular crystal structures [2]:

A
  • Ice (H₂O)

- Iodine (I₂)

69
Q

Molecular crystal structure- iodine [3]:

A
  • I₂ is a covalently bonded molecule
  • Covalent bonds very strong
  • Weak Induced dipole (VDW) forces exist between molecules to give its crystal structure
70
Q

What gives I₂ its crystal structure?

A

Weak Induced dipole (VDW) forces exist between molecules to give its crystal structure

71
Q

What causes I₂’s (crystal) low melting & boiling points [2]:

A
  • Weak VDW’s forces between molecules are v easily broken

- I₂ molecule stays intact cus diatomic

72
Q

Why is I₂ only slightly soluble in water? [2]:

A

I₂ is non-polar

- KI or I⁻ dissolve better/ helps it dissolve

73
Q

Diamond- Structure [3]:

A
  • Each C atom forms 4 single covalent bonds with 4 other atoms
  • Forms 3D solid lattice of Carbon
  • Tetrahedral arrangement (109.5)
74
Q

Properties of diamond [4]:

A
  • Very Hard
  • Very high melting/ boiling point
  • Electrical insulator
  • Insoluble
75
Q

Why is diamond hard?

A

Due to strong Covalent bonds

76
Q

Why does diamond have a high melting/ boiling point? [2]:

A
  • Strong covalent bonds between molecules

- Require lots of energy to break

77
Q

Why is diamond a poor conductor? [2]:

A
  • All e⁻s involved in bonding

- so no delocalised electrons to carry a charge

78
Q

Why’s diamond insoluble? [2]:

A
  • Due to strong c-c bonds

- water can’t break through them

79
Q

Graphite- structure [5]:

A
  • Each C atom forms covalent bond with 3 C atoms
  • 4th C is delocalised
  • Trigonal planar (120 degrees)
  • Forms 2D hexagonal structure (1 atom thicc)
  • Weak VDW forces between layers
80
Q

Properties of Graphite [4]:

A
  • Soft
  • Very high melting/ boiling point
  • Electrical conductor
  • Insoluble
81
Q

Why is Graphite soft? [2]:

A
  • Weak VDW forces between layers

- allows layers to slide over each other when force is applied to it

82
Q

Why does graphite have a high melting point? [2]:

A
  • Strong c-c covalent bonds

- Require a lot of energy to break

83
Q

Whys is graphite a good electrical conductor?

A

Delocalised electrons are free able to carry a charge

84
Q

Why is graphite insoluble?

A

Due to very strong c-c covalent bonds

85
Q

Where does metallic bonding happen?

A

Occurs between metals and alloys

86
Q

Metallic crystal structure =

A

A regular lattice of cations surrounded by a sea of delocalised electrons

87
Q

Properties of metallic bonding [5]:

A
  • High melting points
  • Conduct electricity
  • Conduct heat
  • Malleable
  • Ductile
88
Q

Why do metallic structures have high melting points?

A

Due to strong electrostatic attractions

89
Q

Why are metallic compounds good conductors of electricity?

A

Delocalised electrons free to carry a charge

90
Q

Why are metallic compounds good conductors of heat?

A

Delocalised electrons can quickly transfer energy

91
Q

Why are metallic compounds malleable & ductile? [2]:

A
  • Delocalised electrons able to move with the ion when force is applied
  • They maintain the non-directional electrostatic attractions thus, maintain their shape
92
Q

Repulsion in atoms =

A

Electron pairs around central atom repel each other & settle in a position to minimise repulsion

93
Q

__=

A

Flat to paper

94
Q
        • =
A

Into paper (away from u)

95
Q

|\ =

A

Coming out of paper (towards u)

96
Q

what are the combinations of electron pairs [3]:

A
  • LP-LP
  • LP-BP
  • BP-BP
97
Q

LP-LP =

A

Most repulsive

98
Q

BP-BP =

A

Least repulsive

99
Q

Linear shape examples [2]:

A
  • BeCl₂

- CO₂

100
Q

Bent shape [example]:

A

H₂O

101
Q

Trigonal planar example:

A

BF₃

102
Q

Trigonal Pyramidal example:

A

NH₃

103
Q

Tetrahedral shape examples [2]:

A
  • CH₄

- NH₄⁺

104
Q

Trigonal Bipyramidal examples:

A

PCl₅

105
Q

Octahedral shape example:

A

SF₆

106
Q

Square Planar example:

A

ClF₄⁻

107
Q

How to find the shape of an unknown molecule [5]:

A
  1. Find formula
  2. Identify central atom
  3. Draw the molecule
  4. Add all bonding electrons
  5. Add ‘satellite’/ valence electrons
108
Q

electronegativity in a covalent bond =

A

A diff in electronegativity causes pairs of electrons in the bond to be shared unequally

109
Q

What is the strongest intermolecular force?

A

Hydrogen bonding

110
Q

when does a hydrogen bond? lol

A

Occurs btwn molecules that contain a H atom DIRECTLY bonded to N, O, or F

111
Q

What are the simple molecules with hydrogen bonding? [3]:

A
  • Hydrogen Flouride
  • Water
  • Ammonia
112
Q

Groups in organic chem with hydrogen bonding [3]:

A
  • Alcohols
  • Carboxylic acids
  • Amines
113
Q

What is a hydrogen bond?

A

essentially its a VERY STRONG dipole

114
Q

What is the weakest IMF force?

A

Induced dipole (VDW)

115
Q

Induced dipole forces [2]:

A
  • Occurs btwn ALL molecules

- Strength varies ( Increases with increasing Mr cus more electrons)

116
Q

Increased vdw = [3]:

A
  • Increased melting/ boiling point
  • this cus it would be a bigger molecule
    e. g longer alkanes have higher melting/boiling point
117
Q

Permanent dipole forces [3]:

A
  • Occur btwn polar molecules (they have diff in electronegativity)
  • Stronger than induced/vdw but weaker than h bonding
  • Dipole-dipole IMF causes compounds to have higher than expected melting/ boiling points

AQA: Chemistry Paper 1 (10 decks)

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