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A LEVEL - Chemistry (Both) > Topic 2 - Bonding and Structure > Flashcards

Topic 2 - Bonding and Structure Flashcards

1
Q

What is an ion?

A

A positively or negatively charged atom (or group of atoms).

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2
Q

How are ions formed?

A

When electrons are transferred from one atom to another.

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3
Q

What are positive ions called?

A

Cations

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4
Q

What are negative ions called?

A

Anions

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5
Q

How do single atoms form ions?

A

They gain or lose 1, 2 or 3 electrons so that they have a full outer shell.

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6
Q

Is the charge on a metal ion positive or negative?

A

Positive

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7
Q

Is the charge on a non-metal ion positive or negative?

A

Negative

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8
Q

What is an ionic bond?

A

The strong electrostatic attraction between two oppositely charged ions.

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9
Q

What is the force responsible for ionic bonding?

A

Electrostatic attraction

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10
Q

What does the fomula of an ionic compound tell you?

A

The ratio of the ions in that compound.

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11
Q

How does the electrostatic attraction relate to the strength of an ionic bond?

A

The stronger the attraction, the strongere the bond.

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12
Q

What two factors affect the strength of an ionic bond?

A
  • Ionic charges

* Ionic radius

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13
Q

How does ionic charge affect the strength of an ionic bond?

A

The greater the charge, the stronger the bond.

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14
Q

How does ionic radius affect the strength of an ionic bond and why?

A
  • The smaller the radii, the stronger the bond.

* This is because smaller ions can pack more closely together and electrostatic attraction gets weaker with distance.

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15
Q

What is charge density of an ion?

A

The amount of charge per unit area or volume.

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16
Q

How does charge density affect the strength of an ionic bond?

A

Ions with high charge density form stronger bonds than those with low charge densities.

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17
Q

What is an ion with high charge density?

A

An ion with a large charge spread over a small area.

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18
Q

What is an ion with low charge density?

A

An ion with a small charge spread over a large area.

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19
Q

NaF and CaO.

Which has the higher melting point and why?

A
  • CaO
  • It is made up of Ca2+ and O2- ions, while NaF is made up of Na+ and F- ions.
  • So the charges in the CaO are greater, resulting in stronger ionic bonds.
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20
Q

NaF and CsF.

Which has the higher melting point and why?

A
  • NaF
  • The ionic radius of Ca+ is greater than that of Na+.
  • So Na+ and F- ions can pack more tightly than Ca+ and F- ions, resulting in stronger ionic bonds.
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21
Q

Describe how the size of an ion changes as you go down a group. Why?

A

It is increased, because extra electron shells are added.

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22
Q

What are isoelectronic ions?

A

Ions of different atoms with the same number of electrons.

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23
Q

Describe how the radius of isoelectronic ions changes as the atomic number increases.

A

It decreases, because there is greater attraction per electron from the protons, so they are pulled in closer.

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24
Q

What type of diagram is used to show ionic bonding?

A

Dot-and-cross diagrams

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25
Q

Remember to practise drawing out dot-and-cross diagrams for ionic bonding.

A

See pg 20 of revision guide

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26
Q

Describe the structure that ionic compounds form.

A

Giant ionic lattice

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27
Q

In an ionic compound, why does a giant ionic lattice form?

A

Each Ion is electrostatically attracted in all directions to ions of the opposite charge.

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28
Q

What are the physical properties of ionic compounds?

A
  • High melting point
  • Soluble in water, but not in non-polar solvents
  • Non-conductors when solid, but do conduct when molten or dissolved
  • Brittle
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29
Q

Describe how the physical properties of ionic compounds is evidence for the giant ionic lattice model.

A
  • High melting point -> Tells you that the ions are held together by a strong attraction, like that between positive and negative ions
  • Soluble in water, but not in non-polar solvents -> Particles must be charged
  • Non-conductors when solid, but do conduct when molten or dissolved -> Ions are present, which are only free to move when molten or dissolved
  • Brittle -> Layers can’t be pulled over each other, or the electrostatic repulsion would be too great
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30
Q

Aside from the physical properties of ionic compounds, what else is evidence for the theory of ionic bonding? Explain how this works.

A

Migration of ions in electrolysis:
• When you electrolyse a green solution of copper(II) chromate(VI) ions in some wet filter paper, the filter paper turns blue at the cathode and yellow at the anode.
• This is due to blue copper(II) ions in solution and yellow chromate(VI) ions in solution, which move to the electrodes.
• The solution is yellow to start off with because the ions mix.

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31
Q

What is the colour change observed in the migration of ions experiment observed? Why?

A
  • Green colour (copper(II) chromate(VI) ions) splits into blue and yellow at the electrodes
  • Blue is due to copper(II) ions
  • Yellow is due to dichromate(VI) ions
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32
Q

What colour are dichromate(VI) ions?

A

Yellow

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33
Q

What types of bonds hold a molecule together?

A

Covalent

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34
Q

Describe how a covalent bond forms.

A
  • Two atoms share electrons so they’ve both got full outer shells of electrons.
  • A covalent bond is the strong electrostatic attraction between the two positive nuclei and the shared electrons in the bond
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35
Q

What diagrams can be used to represent covalent bonding?

A

Dot-and-cross diagrams

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36
Q

Remember to practise drawing out dot-and-cross diagrams for covalent bonding.

A

Pg 22 of revision guide, including examples

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37
Q

Describe the dot-and-cross diagram for the bonding in CO.

A
  • C has two non-bonding electrons
  • O had two non-bonding electrons
  • In the covalent bond -> 2 electrons from the C + 4 electrons from the O (this includes a dative covalent bond)
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38
Q

What things balance at the given bond length for a bond?

A
  • Repulsion between positively charged nuclei

* Attraction between the positive nuclei and the area of electron density

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39
Q

What determines the enthalpy of a bond?

A

Its length.

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40
Q

How does the length of a bond affect its enthalpy?

A

The shorter the bond, the higher the bond enthalpy.

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41
Q

Compare the length and bond enthalpy of C-C, C=C and C triplet bond?

A
  • C-C is the longest and has the smallest enthalpy

* C triplet bond is the shortest and has the highest enthalpy

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42
Q

What is a dative covalent bond?

A

A covalent bond where bond of the electrons are donated by one atom.

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43
Q

Describe how an ammonium Ion involves a dative covalent bond.

A
  • Ammonia has a lone pair of electrons on the N

* This lone pair is donated to form a dative covalent bond with a H⁺ ion to give an ammonium ion

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44
Q

When a dative covalent bond is being drawn, which way does the arrow point?

A

The direction in which the electrons are being donated.

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45
Q

Describe the structure of AlCl₃.

A
  • Each Al forms single covalent bonds with 3 Cl’s
  • However, in certain conditions, two AlCl₃ molecules can combine to form Al₂Cl₆.
  • This happens when one Cl in each molecule donates a lone pair to the Al in the other molecule, forming two dative covalent bonds
  • This gives aluminium a full outer shell

(See diagram pg 23 of revision guide)

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46
Q

Remember to practise drawing out the structure of Al₂Cl₆.

A

See diagram pg 23 of revision guide

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47
Q

Give some examples of molecules where dative covalent bonding is seen.

A
  • CO
  • NH₄⁺
  • Al₂Cl₆
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48
Q

Order these in terms of the angle size:
• Bonding pair/bonding pair
• Lone pair/bonding pair
• Lone pair/Lone pair

A
  • Smallest angle: Bonding pair/bonding pair
  • Lone pair/bonding pair
  • Largest angle: Lone pair/lone pair
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49
Q

Do lone pair electrons or bonding pair electrons repel more?

A

Lone pair electrons

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50
Q

What is the name for the way of predicting molecular shapes according to how much electron pairs repel?

A

Electron pair repulsion theory

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51
Q

Describe how much an extra lone pair decreases the bonding angles by in a molecule with 4 electron pairs around the central atom.

A

2.5°

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52
Q

Describe a tetrahedral molecule’s bond angle and how they change with extra lone pairs.

A
  • 109.5°
  • Decreases by 2.5° for every extra lone pair, so:
  • 109.5° -> 107° -> 104.5°
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53
Q

Describe how you can predict the shape of a molecule around a central atom.

A

1) Find the central atom.
2) Work out the number of bonding pairs and lone pairs of electrons around it.
3) Use this information to predict the shape of the molecule.

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54
Q

For a molecule with 2 bonding pairs around the central atom, give the:
• Bond angles
• Name

A
  • Bond angles = 180°

* Name = Linear

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55
Q

For a molecule with 3 bonding pairs around the central atom, give the:
• Bond angles
• Name

A
  • Bond angles = 120°

* Name = Trigonal planar

56
Q

For a molecule with 2 bonding pairs and 1 lone pair around the central atom, give the:
• Bond angles
• Name

A
  • Bond angles = 119°

* Name = Bent

57
Q

For a molecule with 4 bonding pairs around the central atom, give the:
• Bond angles
• Name

A
  • Bond angles = 109.5°

* Name = Tetrahedral

58
Q

For a molecule with 3 bonding pairs and 1 lone pair around the central atom, give the:
• Bond angles
• Name

A
  • Bond angles = 107°

* Name = Trigonal pyramidal

59
Q

For a molecule with 2 bonding pairs and 2 lone pairs around the central atom, give the:
• Bond angles
• Name

A
  • Bond angles = 104.5°

* Name = Bent

60
Q

For a molecule with 5 bonding pairs around the central atom, give the:
• Bond angles
• Name

A
  • Bond angles = 90° and 120°

* Name = Trigonal bipyramidal

61
Q

For a molecule with 4 bonding pairs and 1 lone pair around the central atom, give the:
• Bond angles
• Name

A
  • Bond angles = 87° and 102°

* Name = Seesaw

62
Q

For a molecule with 3 bonding pairs and 2 lone pair around the central atom, give the:
• Bond angles
• Name

A
  • Bond angles = 87.5°

* Name = Distorted T

63
Q

For a molecule with 6 bonding pairs around the central atom, give the:
• Bond angles
• Name

A
  • Bond angles = 90°

* Name = Octahedral

64
Q

For a molecule with 5 bonding pairs and 1 lone pair around the central atom, give the:
• Bond angles
• Name

A
  • Bond angles = 81.9°

* Name -> Square pyramidal

65
Q

For a molecule with 4 bonding pairs and 2 lone pair around the central atom, give the:
• Bond angles
• Name

A
  • Bond angles = 90°

* Name = Square planar

66
Q

Remember to practise drawing out and naming different shapes of molecules.

A

Pg 25 of revision guide

67
Q

What type of structure do covalently bonded atoms form?

A
  • Simple covalent molecule

* Giant covalent structure

68
Q

Give three examples of giant covalent structures.

A
  • Carbon
  • Silicon
  • Silicon dioxide
69
Q

Describe the structure of diamond.

A

Tetrahedral, with each carbon bonding with another four carbons.

(See diagram pg 26 of revision guide)

70
Q

Describe the structure of silicon dioxide.

A
  • Silicon atoms form a tetrahedral shape, like in diamond
  • Oxygen atoms are between each of the two silicon atoms

(See diagram pg 26 of revision guide)

71
Q

Describe and explain the properties of giant covalent structures.

A
  • High melting points -> Need to break a lot of very strong bonds
  • Hard -> Due to very strong bonds
  • Insoluble -> Atoms are more attracted to their neighbours in the lattice than to solvent molecules. Also, no ions.
  • Can’t conduct electricity -> No charged ions or free electrons.
72
Q

What is the only exception to giant covalent structures not being able to conduct electricity?

A

Graphite

73
Q

Describe the structure of graphite.

A
  • Each carbon forms 3 single bonds with other carbons -> This forms a hexagonal structure
  • There are multiple sheets held together by weak forces
  • Each carbon has a free electron (to conduct electricity)
74
Q

What allows graphite to conduct electricity?

A

Each carbon has a single free electron that is free to move around and conduct electricity.

75
Q

What is one layer of graphite called?

A

Graphene

76
Q

Describe the structure of graphene.

A
  • It is a single sheet

* Each carbon forms 3 single bonds with other carbons -> This forms a hexagonal structure

77
Q

What are the properties of graphene?

A
  • Conductor of electricity
  • Transparent
  • Strong
  • Light
78
Q

Describe the structure of a metal.

A

Exist as giant metallic lattices:
• The outermost shell of the metal atoms are delocalised -> This leaves a positive metal ion
• Positive metal ions are electrostatically attracted to the sea of delocalised electrons
• The overall lattice structure is made up of layers of positive metal ions, separated by layers of electrons

79
Q

Describe and explain the properties of metals.

A
  • High melting point -> Strong attraction between the positive ions and sea of delocalised electrons
  • Malleable + Ductile -> Layers of metal ions are free to slide over each other, without disrupting the bonding
  • Good thermal conductors -> Electrons can pass kinetic energy to each other
  • Good electrical conductors -> Delocalised electrons are free too move and carry a current
  • Insoluble (except in liquid metals) -> Metallic bonds are very strong
80
Q

What factors affect the melting point of a metal?

A
  • No. of delocalised electrons per atom
  • Size of the metal ion
  • Lattice structure
81
Q

Define electronegativity.

A

The ability of an atom to attract the binding electrons in a covalent bond.

82
Q

On what scale is electronegativity measured?

A

Pauling

83
Q

What value are very electronegative elements given on the Pauling scale?

A

High

84
Q

What are the properties of very electronegative elements?

A
  • High nuclear charges

* Small atomic radii

85
Q

How does electronegativity change across periods and groups?

A
  • Increases across periods

* Increases up groups

86
Q

What causes a polar bond?

A

When the two elements on either side of the bond have very different electronegativities.

87
Q

What is the name for a difference in charge between the atoms at each end of a polar bond?

A

Dipole

88
Q

What are the only truly covalent bonds?

A

Those between two atoms of the same element.

89
Q

How does the difference in electronegativity between two elements tell you about the bond between them?

A

The greater the difference, the more ionic the bond.

90
Q

Do polar bonds always make polar molecules?

A

No, often the polar bonds may point in opposite directions and cancel out.

91
Q

What are the 3 types of intermolecular force?

A

1) London forces
2) Permanent dipole-permanent dipole bonds
3) Hydrogen bonds

92
Q

What is another name for London forces?

A

Instantaneous dipole-induced dipole bonds

93
Q

What type of molecules form London forces?

A

All atoms and molecules do

94
Q

Describe how London forces form.

A
  • Electrons in charge clouds are always moving quickly
  • At any instant, the electrons in an atom are likely to be made to one side than another -> This is an instantaneous dipole
  • This dipole can induce another temporary dipole in the opposite direction on a neighbouring atom -> This is an induced dipole
  • This can induce another dipole in another atom
  • These dipoles are constantly being created and destroyed, but the overall effect is for the atoms to be attracted
95
Q

What types of forces hold molecules in a lattice?

A

London forces

96
Q

Describe the bonding in iodine.

A
  • Iodine atoms within I₂ are held together by strong covalent bond.
  • The whole molecules are then held together in a molecular lattice arrangement by weak London forces
97
Q

What is the effect of stronger intermolecular forces?

A

The melting and boiling posits are higher.

98
Q

What factors affect the size of the intermolecular forces within an organic compound?

A
  • Size

* Surface Area

99
Q

Compare the strength of intermolecular forces in straight and branched-chain organic compounds.

A

Straight molecules have stronger intermolecular forces due to their increased surface area and better ability to pack closer together.

100
Q

Compare the melting and boiling points of butane and methylpropane (both C₄H₁₀).

A

Butane has a higher melting and boiling point because it is straight, not branched-chain.

101
Q

Describe how permanent dipole-permanent dipole bonds form.

A

δ⁺ and δ⁻ charges on polar molecules cause weak electrostatic forces of attraction between molecules.

102
Q

Can a molecule have both London forces and permanent dipole-permanent dipole bonds?

A

Yes

103
Q

Compare the melting and boiling point of molecules with just London forces and those with both London force and permanent dipole-permanent dipole bonds.

A

The ones with both tend to have higher melting and boiling points.

104
Q

What is the strongest intermolecular force?

A

Hydrogen bonding

105
Q

Describe the formation of a hydrogen bond.

A

• H atom bonded to N, O or F can form a hydrogen bond with the lone pair of electrons on N, O or F of another molecule
This is because:
• The H is given a strong positive by the very electronegative N, O or F
• It is also very small, so it has a high charge density

106
Q

How can you tell whether something is a hydrogen bond?

A

There is a H atom with an N, O or F either side of it.

107
Q

What is the effect of hydrogen bonding?

A

It raises the melting and boiling point of a molecule hugely.

108
Q

Describe and explain the graph of boiling points of hydrogen halides.

A
  • HF has a very high boiling point -> Due to hydrogen bonding
  • Large drop to HCl -> Since it can’t form hydrogen bond
  • Steady increases to HBr and HI -> Increasing size of the molecule means that there are stronger London forces (which overrides the effect of the weaker dipoles)
109
Q

Why does HBr have a higher boiling point than HCl, even though the strength of the permanent dipole-permanent dipole increases?

A

The increased strength of the London forces overcomes this decrease.

110
Q

Remember to revise the graph of the boiling points of Group 6 hydrides (like H₂O).

A

Pg 32 of revision guide

111
Q

What is the effect of hydrogen bonds on solubility?

A

It makes the substance soluble, because it can form hydrogen bonds with the water molecules, allowing them to mix and dissolve.

112
Q

Describe the structure of ice.

A
  • 6 H₂O molecules line up to form a hexagon

* There are hydrogen bonds holding the structure together

113
Q

Remember to practise drawing out the structure of ice.

A

Pg 33 of revision guide

114
Q

Why does ice float?

A

In ice:
• 6 H₂O molecules line up to form a hexagon
• There are hydrogen bonds holding the structure together
• This is a structure that wastes a lot of space
When the ice melts:
• The water molecules lose their lattice structure and become more closely packed
• This increases the density, so water has a lower density and floats

115
Q

What is the effect of hydrogen bonding on volatility?

A

It reduces the volatility (i.e. it increases the boiling point)

116
Q

Compare and explain the volatility of alcohols and equivalent alkanes.

A
  • The alcohols have lower volatility (i.e. higher boiling points)
  • Because they can form hydrogen bonds due to the -OH group
  • So it takes more energy to overcome the intermolecular forces and boil them
117
Q

What things have to happen in order for a substance to dissolve?

A
  • Bonds in the substance have to break
  • Bonds in the solvent have to break
  • Bonds have to form between the substance and the solvent

The strength of the new bonds has to be about the same as, or greater than, the bonds broken.

118
Q

What are the two types of solvent?

A
  • Polar

* Non-polar

119
Q

Do all polar solvents form hydrogen bonds?

A

Most, but not all. For example, propanone only forms London forces and permanent dipole-permanent dipole bonds.

120
Q

What intermolecular forces do non-polar solvents have?

A

London forces

121
Q

What is the process of water molecules surrounding ions in dissolving called?

A

Hydration

122
Q

In what solvents do ionic substances dissolve?

A

Polar solvents

123
Q

Why don’t some ionic substances dissolve in water?

A

The bonding between their ions is too strong - so it is stronger than the bonds they would form with the water molecules.

124
Q

How does the solubility of alcohols change with their size?

A

The larger the alcohol, the less soluble it is, because:
• Small alcohols -> The OH groups can form H-bonds with the water molecules
• Large alcohols -> The hydrocarbon chain can’t form H-bonds with the water molecules

125
Q

Do all polar molecules dissolve in water?

A

No, usually it is just those that can form hydrogen bonds with the water.

126
Q

Do halogenoalkanes dissolve in water and non-polar solvents?

A
  • Water -> No

* Non-polar solvents -> Yes, those that form permanent dipole-permanent dipole bonds

127
Q

Explain the solubility of halogenoalkanes in water.

A

They are not soluble because the hydrogen bonding between the water molecules is stronger than the bonds that would be formed with halogenoalkanes.

128
Q

What do non-polar substances dissolve in best?

A

Non-polar solvents

129
Q

What are the general rules for solubility?

A

Substances usually dissolve best in solvents with similar intermolecular forces.

130
Q 
Describe ionic compounds in terms of:
• Melting/Boiling point
• State at RTP
• Electrical conductivity when solid
• Electrical conductivity when liquid
• Solubility in water
A
  • Melting/Boiling point -> High
  • State at RTP -> Solid
  • Electrical conductivity when solid -> No
  • Electrical conductivity when liquid -> Yes
  • Solubility in water -> Yes
131
Q 
Describe simple molecular compounds in terms of:
• Melting/Boiling point
• State at RTP
• Electrical conductivity when solid
• Electrical conductivity when liquid
• Solubility in water
A
  • Melting/Boiling point -> Low
  • State at RTP -> Usually liquid or gas (but may be solid)
  • Electrical conductivity when solid -> No
  • Electrical conductivity when liquid -> No
  • Solubility in water -> Depends on whether it can hydrogen bond
132
Q 
Describe giant covalent compounds in terms of:
• Melting/Boiling point
• State at RTP
• Electrical conductivity when solid
• Electrical conductivity when liquid
• Solubility in water
A
  • Melting/Boiling point -> High
  • State at RTP -> Solid
  • Electrical conductivity when solid -> No
  • Electrical conductivity when liquid -> N/A since they usually sublime
  • Solubility in water -> No
133
Q 
Describe metallic substances in terms of:
• Melting/Boiling point
• State at RTP
• Electrical conductivity when solid
• Electrical conductivity when liquid
• Solubility in water
A
  • Melting/Boiling point -> High
  • State at RTP -> Solid
  • Electrical conductivity when solid -> No
  • Electrical conductivity when liquid -> N/A since they usually sublime
  • Solubility in water -> No
134
Q

Substance X has a melting point of 1045 K. When solid, it is an insulator, but once melted it conducts electricity. Identify the type of structure present in substance X.

A

Ionic

135
Q

Remember to practise predicting the properties of a material from its structure and vice versa.

A

Pgs 36 and 37 of revision guide

A LEVEL - Chemistry (Both) (21 decks)

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