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AQA A Level Chemistry > Bonding, Shapes of molecules and Types of intermolecular forces > Flashcards

Bonding, Shapes of molecules and Types of intermolecular forces Flashcards

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1
Q

Ionic bond

A

Electrostatic force of attraction between oppositely charged ions in a lattice as a result of electron transfer

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2
Q

Why is a sodium ion smaller than a sodium atom

A

Na+ is smaller than Na since it has one less shell of electrons
The proton to electron ratio is greater in Na+
Na+ has a greater effective nuclear charge than Na so the remaining electrons are more strongly attracted and are closer to the nucleus

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3
Q

Structure of an ionic bond

A

Giant ionic lattice

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4
Q

Isoelectronic

A

Same number and arrangement of electrons

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5
Q

Suggest 3 forces of repulsion which exist in an ionic lattice

A

Between ions of the same charge
Between electron clouds of the ions
Between the nuclei of the ions

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6
Q

What are the properties of ionic compounds

A

High melting temperatures
Solids at room temperature
Soluble in water (most ionic compounds)
Conduct electricity when molten or dissolved in water

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7
Q

Why do ionic compounds have a high temperature

A

They have a giant ionic lattice.
There are strong electrostatic forces of attraction between the ions.
A large amount of energy is needed to break the strong ionic bonds

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8
Q

Why can ionic compounds conduct electricity when molten or dissolved in water

A

The ions are free to move and carry the charge

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9
Q

Explain why sodium iodide has a lower melting temperature than that of sodium chloride. The charge on both anions is the same, -1, but the I- is bigger than Cl-

A

I- has a smaller charge density than Cl-

The electrostatic forces of attraction between Na+ and I- ions are weaker therefore less energy is needed to break the weaker ionic bond in NaI

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10
Q

What are major factors that affect the strength of an ionic bond

A

Radius/size of the ion
Charge of the ion

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11
Q

What happens to ionic radius as you go down a group

A

It increases

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12
Q

What happens to ionic radius as you go across a period

A

It decreases

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13
Q

Why does ionic radius increase as you go down a group

A

Outer electrons are further away from the nucleus.

Although the nuclear charge increases down a group, the outer electrons are more shielded by the inner electron shells.

There is a weaker attraction between the nucleus and the outer electrons as you go down a group so ionic radius increases

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14
Q

Why does ionic radius decrease as you go across a period

A

The ratio of protons to electrons increases across a period so nuclear charge increases

There is no change in shielding of the outer shell of electrons.

There is a stronger attraction between the nucleus and the outer electron as you go across a period so ion become smaller

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15
Q

What is a covalent bond

A

The attraction between a shared pair of electrons and the nuclei of the bonded atoms. Each atom gives one electron to the shared pair

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16
Q

What is average bond enthalpy a measurement of

A

Covalent bond strength

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17
Q

What happens to the strength of a covalent bond as average bond enthalpy increases

A

Covalent bond becomes stronger

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18
Q

Properties of simple molecular structure (covalent)

A

They are liquids or gases at room temperature

They have low melting points and boiling points

They do not conduct electricity

Many do not dissolve in water

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19
Q

Why do covalent compounds have low melting and boiling points

A

They have a simple molecular structure and are made up of small molecules with weak Van der Waals forces between the molecules. Not much energy is needed to break these weak intermolecular forces

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20
Q

Why do covalent compounds not conduct electricity

A

There are no free electrons since the outer electrons are all fixed in covalent bonds

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21
Q

Why does iodine, I2, have a low melting point

A

Iodine is made up of small molecules with weak Van der Waals forces between the molecules. Not much energy is needed to break these weak intermolecular forces.

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22
Q

Dative covalent bond

A

The attraction between a shared pair of electrons and the nuclei of the bonded atoms. The shared pair of electrons is given by the SAME atom

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23
Q

What is a co-ordinate bond/ dative bond represented by

A

An arrow showing the direction of donation

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24
Q

An aluminium chloride molecule reacts with a chloride ion to form the AlCl4- ion.
Name the type of bond formed in this reaction and explain how this bond is formed in AlCl4-.

A

type of bond : Dative covalent bond
Explanation : The electron pair on Cl- is donated to the Al atom in AlCl3

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25
Q

Why can element in periods 1 and 2 not expand their octet

A

They have no 2d orbitals

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26
Q

Why can elements in period 3 and above expand their octet

A

They have an empty 3d orbital

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27
Q

Explain why the melting point of BCl3 is much lower than that of NaCl

A

BCl3 is made up of small molecules with weak Van der Waals forces between the molecules as it is covalently bonded. Not much energy is needed to break these weak intermolecular forces.

NaCl has a giant ionic lattice with strong electrostatic forces of attraction between the ions. A large amount of energy is needed to break these strong ionic bonds

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28
Q

What are the allotropes of carbon

A

Diamond, graphite and graphene

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29
Q

Allotropes

A

Different forms of the same element in the same physical state

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30
Q

Structure of diamond

A

Giant covalent lattice

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31
Q

Why does diamond have a high MP

A

it has a giant covalent lattice with many strong covalent bonds between the atoms.
A large amount of energy is needed to break so many bonds

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32
Q

Why is diamond a poor conductor of electricity

A

There are no free electrons. The outer electrons in diamond are all fixed in covalent bonds

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33
Q

Structure of graphite

A

giant covalent lattice

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34
Q

Why is graphite soft

A

There are weak Van der Waals forces between the layers of carbon atoms so the layers can slide over each other

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35
Q

Explain how solid graphite conducts electricity

A

In graphite each carbon has one free electron. Graphite conducts electricity because it has delocalised electrons that move along the layers and carry the charge

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36
Q

What are the 4 types of crystal structure

A

ionic
metallic - magnesium and sodium
macromolecular (giant covalent) - diamond, graphite
molecular - iodine, ice

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37
Q

Metallic bond

A

Electrostatic force of attraction between the delocalised electrons and positive metal ions arranged in a lattice

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38
Q

Structure of metals

A

Positive metal ions are arranged in a regular lattice and held together by electrostatic attraction to delocalised electrons. Metals have a giant metallic lattice

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39
Q

Why are metals malleable/ductile

A

There are no rigid, directed bonds in the metal so the layers of ions can slide over each other when a force is applied to the metal

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40
Q

How do pairs of electrons in the outer shell of atoms arrange themselves

A

as far apart as possible to a position of minimal repulsion

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41
Q

What is the order of pair repulsion

A

lone pair-lone pair repulsion > lone pair-bond pair repulsion > bond pair-bond pair repulsion

42
Q

Lone electron pair

A

2 electrons that do not form a covalent bond

43
Q

How do you work out the total number of bonding pairs and lone pairs for molecules and ions with single bonds only

A
  1. Work out the number of outer shell electrons on the central atom by looking at what group it is in e.g group 5 has 5 electrons. DON’T FORGET ABOUT THE CHARGES.
  2. Work out how many atoms are bonded to the central atom (number of bonding pairs) e.g. NH3 would have 3 as there are 3 hydrogens
  3. Work out how many lone pairs there are by doing 1/2( outer electrons on central atom - number of atoms bonded to central atom)
  4. Add the bonding pairs and lone pairs
44
Q

What is the shape of the molecule and its bond angles if it has 2 bonding pairs and 0 lone pairs

A

Linear and 180 degrees

45
Q

What is the shape of the molecule and its bond angles if it has 3 bonding pairs and 0 lone pairs

A

Trigonal planar and 120 degrees

46
Q

What is the shape of the molecule and its bond angles if it has 4 bonding pairs and 0 lone pairs

A

tetrahedral and 109.5 degrees

47
Q

What is the shape of the molecule and its bond angles if it has 3 bonding pairs and 1 lone pair

A

Trigonal pyramidal and 107 degrees

48
Q

What is the shape of the molecule and its bond angles if it has 2 bonding pairs and 2 lone pairs

A

Bent ( V shaped) and 104.5 degrees

49
Q

What is the shape of the molecule and its bond angles if it has 5 bonding pairs and 0 lone pairs

A

Trigonal bipyramidal and 120, 90 degrees

50
Q

What is the shape of the molecule and its bond angles if it has 4 bonding pairs and 1 lone pair

A

Trigonal bipyramidal and 119,89 degrees

51
Q

What is the shape of the molecule and its bond angles if it has 3 bonding pairs and 2 lone pairs

A

Trigonal bipyramidal and 89 degrees

52
Q

What is the shape of the molecule and its bond angles if it has 6 bonding pairs and 0 lone pairs

A

Octahedral and 90 degrees

53
Q

Why is the beryllium chloride molecule linear

A

There are 2 bonding pairs and no lone pairs of electrons around the central beryllium atom. These electron pairs repel each other as far apart as possible to a position of minimum repulsion. The BeCl2 molecule is linear with bond angles of 180 degrees

54
Q

What is the shapes and bond angles of CO2 and HCN

A

Linear with bond angles of 180 degrees

55
Q

Why is the CO2 molecules linear

A

There are two regions of electron density around the central C atom. These electron regions repel each other as far apart as possible to a position of minimum repulsion. The CO2 molecule is linear with bond angles of 180 degrees. ( USE AS AN EXAMPLE FOR ALL EXCEPTIONS E.G NO3-)

56
Q

How are double and triple bonds treated when predicting the shape of a molecule

A

As a single bond

57
Q

Why is the AlCl3 molecule trigonal planar

A

There are 3 bonding pairs and no lone pairs of electrons around the central Al atom. These electron pairs repel each other as far apart as possible to a position of minimum repulsion. The AlCl3 molecule is therefore trigonal planar with bond angles of 120 degrees

58
Q

What are the shapes and bond angles of NO3- and CO3 2-

A

Trigonal planar with bond angles of 120 degrees

59
Q

What does a straight line represent when drawing the shape of a molecule

A

A bond in the plane of the paper

60
Q

What does a triangle represent when drawing the shape of a molecule

A

A bond in front of the plane of the paper

61
Q

What does a dashed line represent when drawing the shape of a molecule

A

A bond behind the plane of the paper

62
Q

What is the shape and bond angle of SO4 2-

A

Tetrahedral with bond angles of 109.5

63
Q

State and explain the shape and bond angle(s) of the ammonia molecule and

A

There are 3 bonding pairs and one lone pair of electrons around the central N atom. These electron pairs repel each other as far apart as possible to a position of minimum repulsion. The NH3 molecule is pyramidal.

Since lone pair bonding-pair repulsion is greater than bonding pair-bonding pair repulsion, the H-N-H bond angle is lesser than the tetrahedral angle and so is 107 degrees.

64
Q

What is the shape and angle of SO3 2-

A

Pyramidal with bond angles of 107 degrees

65
Q

Electronegativity

A

The power/ability of an atom to attract the pair of electrons in a covalent bond

66
Q

What is the most electronegative element

A

Fluorine

66
Q

What are the trends in electronegativity in the periodic table

A

Non-metals have a higher electronegativity than metals

Electronegativity increases across a period of the periodic table as atom gets smaller across a period

Electronegativity decreases as you go down a group as atom gets bigger down a group

66
Q

In terms of atomic structure, give 2 reasons why oxygen is more electronegative than carbon

A

Oxygen atom has a greater nuclear charge than carbon atom because it has more protons.

Oxygen atom has a smaller atomic radius than the carbon atom

67
Q

How do partial charges arise when 2 different atoms are joined by a covalent bond

A

The more electronegative atom attracts the bonding pair of electrons more and as a result has a delta negative charge whilst the other atom has a delta positive charge

68
Q

When is a bond non-polar

A

The bonded atoms are the same or the bonded atoms have the same/similar electronegativity

69
Q

Dipole

A

Separation of opposite charges

70
Q

Permanent dipole

A

A dipole in a polar covalent bond which does not change

71
Q

Why is CCl4 not polar

A

The CCl4 molecule is symmetrical and the dipoles cancel.

72
Q

What happens in terms of permanent dipoles, if the molecule is symmetrical

A

Any dipoles will cancel and molecule will NOT have a permanent dipole

73
Q

What happens in terms of permanent dipoles, if the molecule is not symmetrical

A

Dipoles do not cancel and molecule has a permanent dipole

74
Q

Why might a molecule have polar bonds but not necessarily be a polar molecule

A

The molecule is symmetrical and dipoles cancel

75
Q

What shaped molecules form polar molecules

A

Bent, pyramidal and trigonal bipyramidal

76
Q

What are the 3 types of intermolecular forces

A

Van der Waals forces
Permanent dipole-dipole interactions
Hydrogen bonding

77
Q

Which type of intermolecular forces is the strongest

A

Hydrogen bonding > permanent dipole-dipole interactions > Van der Waals forces

78
Q

Where do Van der Waals forces exist

A

Between non-polar molecules / all molecular substances and noble gases OR exist between all atoms or molecules

79
Q

Explain how Van der Waals forces arise between non-polar molecules

A

Due to the movement of electrons there is an uneven distribution of electrons in a molecule. This gives rise to an instantaneous dipole.

This instantaneous dipole induces a dipole on the neighbouring molecule.

The induced dipole induces further dipoles on neighbouring molecules which then attracts one another

80
Q

What is the main factor affecting the size of Van der Waals forces

A

The number of electrons.

81
Q

How does the number of electrons affect the size of Van der Waals forces

A

As the number of electrons per molecule increases, the strength of the forces increases. LArger molecules have larger electrons clouds meaning stronger Van der Waals forces

82
Q

Where do permanent dipole-dipole interactions occur

A

Between polar molecules

83
Q

Explain what happens in permanent dipole-dipole interactions

A

Polar molecules have a permanent dipole.

Polar molecules are attracted to each other by permanent dipole-dipole interactions.

The negative end of one molecule is attracted the positive end of another.

84
Q

How is a hydrogen bond formed

A

By the attraction between the hydrogen atom of one molecule and the lone pair of an N, O or F atom of a different molecule

85
Q

What are the conditions for a hydrogen bond to form

A

A hydrogen atom is attached to a small electronegative atom such as N, O or F

An electronegative atom with a lone pair of electrons

86
Q

Explain why a hydrogen bond forms between 2 water molecules

A

Oxygen atom is more electronegative than the hydrogen atom.

Oxygen atom attracts the bonding pair of electrons more and has a delta negative charge hydrogen atom has a delta positive charge

A hydrogen bond is formed by the attraction of the delta positive hydrogen atom of one water molecule and the lone pair on the small delta negative oxygen atom on a different water molecule

87
Q

Why does H2O have a higher BP than NH3

A

Each water molecule form 2 hydrogen bonds but each ammonia molecule can only form one. More energy is needed to break the hydrogen bonds between water molecules.

88
Q

Why does HF have a higher BP than NH3

A

Fluorine is more electronegative than nitrogen so the H-F bond is more polar than the N-H bond.

The hydrogen bonds between hydrogen fluoride molecules are stronger.

More energy is needed to break the hydrogen bonds between HF molecules

89
Q

Polar

A

The electrons are shared unequally in a bond. One atom has a partial positive charge and the other atom has a partial negative charge

90
Q

Explain the structure of ice

A

With 2 hydrogen atoms and 2 lone pairs on the oxygen atom, each water molecule can form 4 hydrogen bonds. The hydrogen bonds extend outwards, holding the water molecules slightly in a 3 dimensional tetrahedral lattice.

This arrangement creates an open lattice structure which decreases its density, hence ice is less dense than water

91
Q

Why is the boiling point of HF greater than HCl, HBr and HI

A

There are hydrogen bonds between molecules of hydrogen fluroride
There are Van der Waals forces between molecules of the other hydrogen halides.
Hydrogen bonds are stronger than VdW forces therefore more energy is needed to separate molecules of HF

92
Q

Why is the BP of HI greater than HCl

A

HI has more electrons per molecule than HCl, therefore the VdW forces between HI molecules are stronger and more energy is needed to separate it.

93
Q

What is the trend in BP for the noble gases

A

It increases as you go down the group since there are more electrons per atom therefore the strength of the VdW forces increases.

94
Q

What happens to the boiling point as the molecule becomes more branched and why

A

BP decreases because there is less contact with molecules with a smaller SA .There are more VdW forces between linear molecules compared to branched or spherical molecules

95
Q

What is the trend in boiling points of the hydrogen halides from HF to HI

A

HF has the greatest boiling point and then it HI, HBr and finally HCl

HF>HI>HBr>HCl

96
Q

Why is the boiling temperature of HF greater than that of HCl, HBr or HI

A

There are hydrogen bonds between molecules of hydrogen fluoride.
There are Van der Waals forces between molecules of the other hydrogen halides.
Hydrogen bonds are stronger than Van der Waals forces therefore more energy is needed to separate molecules of HF

97
Q

Why is the boiling point of HI greater than HCl

A

Hydrogen iodide has more electrons per molecule than hydrogen chloride, therefore the Van der Waals forces between hydrogen iodide molecules are stronger.

More energy is needed to separate molecules of hydrogen iodide.

98
Q

State and explain whether the electronegativity of fluorine is greater than, similar to or less than that of bromine.

Hence explain why HF can form hydrogen bonds but HBr cannot

A

Fluorine is more electronegative than bromine because it is smaller.
The bonding electrons in HF are closer to the nucleus of the fluorine atom.
The H-F bond is polar, but the H-Br bond is not polar enough, so hydrogen bonds cannot form between HBr bonds

99
Q

What is the trend in boiling points of the noble gases and why

A

The boiling points increase as you go down the group.
The number of electrons per molecule increases down the group so VdW forces get stronger .
More energy is therefore needed to separate the atoms.

100
Q

What is the shape and bond angles of a molecules with 2 bonding pairs and 1 lone pair

A

Bent, 118 degrees

AQA A Level Chemistry (27 decks)

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