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Chemistry A - Level > Topic 2 - Bonding and Structure > Flashcards

Topic 2 - Bonding and Structure Flashcards

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1
Q

What is the definition of an Orbital?

A

The area around an atom where there is a high probability of finding an electron

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2
Q

What is the definition of Bonding?

A

How atoms or ions form attractive forces

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3
Q

What is the definition of Structure?

A

How atoms or ions are arranged

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4
Q

What is the definition of Isolelectronic?

A

When two atoms, ions or molecules have the same electronic configuration and same number of valence electrons e.g Na(+), Mg(2+) and Al(3+)

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5
Q

Why do atoms form bonds?

A

In order to become more stable (they do not have to have a full outer shell of this to be the case)

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6
Q

Definition of Ionic Bonding

A

The electrostatic forces of attraction between oppositely charged ions

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7
Q

Definition of Covalent bonding

A

The electrostatic forces of attraction between the negative shared pair of electrons and their adjacent, positively charged nuclei

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8
Q

Definition of Metallic Bonding

A

The electrostatic forces of attraction between the delocalised (sea of) electrons and the positively charged ions

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9
Q

What factors affect the strength of the ionic bond

A
  • The charge of the ions involved
  • the radius of the ions (ionic radius)
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10
Q

What is the trend in ionic radius down a group?

A

The ionic radius increases due to more filled shells of electrons meaning shielding increases so there is less attraction between the electrons and the nucleus

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11
Q

Ionic Charge density formula

A

= charge/ionic radius

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12
Q

Which two atoms form ions that are electron deficient?

A

Be is in group 2 and only has 4 electrons in its outer shell when it forms BeCl2
B is in group 3 and only has 6 electrons in its outer shell when it forms BF3

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13
Q

Which two atoms from ions that are expanded octets and break the octet rule?

A

P has 5 electrons in its outer shell and forms PCl5 and therefore has 10 electrons in its outer shell
S has 6 electrons in its outer shell and forms SF6 and therefore has 12 electrons in its outer shell

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14
Q

What is a dative covalent bond (coordinate bond)

A

Like a normal covalent bond it involves a shared pair of electrons, however, in a dative covalent bond, both of the electrons originate from one atom.

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15
Q

Give an example of an ion that has a dative covalent

A

NH4 (+) when NH3 and H+ are bonded together

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16
Q

Why do dative covalent bonds occur?

A

When an atom, such as nitrogen in NH3, has a non-bonded pair of electrons

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17
Q

How do you show a dative covalent bond?

A
  • In a dot and cross diagram they are from the same species so are drawn as the same, dots or crosses
  • In a stick diagram you draw an arrow from the species with the pair of electrons to the atoms that is accepting them
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18
Q

What two factors does bond length depend on?

A
  • The size of the atoms
  • The number of shared electron pairs (As this increases the strength of the covalent/ionic bond gets greater)
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19
Q

As bond strength decreases, bond length…

A

Increases

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20
Q

What is pure ionic bonding?

A

Pure ionic bonding is when there is only a very little polarising effect of the negative ion by the positive ion so the ions are perfectly spherical. The electrons are shared equally between the two atoms (however, this rarely happens)

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21
Q

What does pure ionic bonding require?

A
  • A large positive ion with a low charge
  • A small negative ion
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22
Q

What is polarisation?

A

When a cation distorts the electron cloud of an anion

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23
Q

Why do certain ions have a higher polarising effect?

A

As they have a high charge density e.g Al (3+), which will attract an anion with a large ionic radius

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24
Q

Why are larger negative ions more easily polarised than smaller negative ions?

A

As their outer electrons are less attracted to their nucleus and so more easily distorted

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25
Q

What is the definition of electronegativity?

A

A measure of the tendency of an atom to attract a bonding pair of electrons

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26
Q

Electronegativity increases as…

A

You move right and up on the periodic table

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27
Q

Where does pure covalent only occur?

A

If the two atoms are identical e.g N2, O2 and H2

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28
Q

How do you draw an electron density map?

A

Draw circles around the separate atoms. In ionic bonding the circles don’t overlap where as, in covalent bonding they join in the middle and are symmetrical.

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29
Q

Why does melting and boiling point increase across a period e.g Na, Mg, Al

A
  • The ionic charge increases and the number of delocalised electrons increases
  • Therefore there is more attraction between the electrons and the nucleus so the ionic radius decreases
  • Therefore, there is a higher charge density and so the forces of attraction within the metallic bonds are stronger so require more energy to overcome.
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30
Q

Why do the melting and boiling points of group 1 metals decrease as you go down the group?

A
  • ionic radius increases
  • the ratio of atoms to electrons in the metallic structure stays the same
  • Charge density decreases, therefore the forces of attraction between positive ions and electrons decrease
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31
Q

What is the order of repulsion strength between pairs of electrons?

A

Lone pair - Lone pair > Lone pair - Bonding pair > Bonding pair - Bonding pair

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32
Q

What is the bond angle in linear molecules? Give examples.

A

180° as they have 0lp’s and 2bp’s that repel to the point of minimum repulsion e.g BeCl2, CO2, HCN and C2H2

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33
Q

What is the bond angle in Triagonal planar molecules? Give examples.

A

120° as they have 0lp’s and 3bp’s that repel to the position of minimum repulsion e.g BCl3

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34
Q

What is the bond angle in Tetrahedral molecules? Give examples.

A

109.5° as they have 0lp’s and 4bp’s that repel to the position of minimum repulsion.
The central atom lies at the centre of the tetrahedron.
The bond coming out of the is represented by a filled triangle and the bond into the page, by a dashed triangle.

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35
Q

What is the bond angle in Trigonal bipyramidal molecules? Give examples.

A

The equatorial atoms that lie on the plane have a bond angle of 120°. The axial atoms that upand down make a bond angle of 90° with the equatorial atoms. The molecule has 5bp’s and 0 lp’s.

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36
Q

What is the bond angle in Octahedral Molecules? Give Examples.

A

90° as they have 6bp’s and 0lp’s that repel to the point of minimum repulsion. Two bonds go into the page and two come out of the page.

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37
Q

What is the bond angle in Triangular Pyramidal Molecules? Give Examples.

A

107° as they have 1lp and 3bp’s which repel to the position of minimum repulsion. One bond is on the plane, one into the page and one out of the page e.g NH3

38
Q

What is the bond angle in Non -linear molecules? Give examples.

A

104.5° as they have 2bp’s and 2lp’s that repel to the position of minimum repulsion. Both bonds are on the plane e.g H2O

39
Q

How can molecules that are electron deficient gain a full outer shell?

A

By sharing a pair of electrons with an atom which has a non-bonding or lone pair of electrons

40
Q

What is the bond angle in Square Planar Molecules? Give examples.

A

90° as they have 4bp’s and 2lp’s. Two bonds go into the plane and two bonds go out of the plane e.g XeF4

41
Q

Why are all the bond angles in Methane 109.5°?

A
  • Methane contains 4 bonding pairs and 0 lone pairs of electrons.
  • Electron pairs repel to the position of minimum repulsion
  • The 4 bonding pairs of electrons will repel each other equally
  • Giving methane it’s tetrahedral shape with bond angle 109.5°
42
Q

What is the trend in ionic radius across period 3, from Na to Cl?

A

As you go from Na (+) → Al (3+) the ions begin already small and get smaller. As you go from P (3-) → Cl (-) the ions begin big and get smaller

43
Q

Why do the negative ions all have a greater ionic radius than the positive ones?

A

This is because the negative ions are gaining electrons, therefore they are gaining a whole extra layer making their radius larger.

44
Q

Why do the metallic ions decrease in radius from group 1 to group 3?

A

Whilst the ions are isoelectronic, as you go across a period the number of protons in the nucleus is increasing giving it a more positive charge. This in turn attracts the electrons more strongly causing a decrease in the ionic radius.

45
Q

Why do the non-metallic elements decrease in radius from group 5 to group 7?

A

The number of electrons is decreasing so the overall charge on the ion gets smaller. Therefore, there is less repulsion between the electrons so a smaller ionic radius.

46
Q

What is the definition of ionic radius?

A

Ionic radius is the distance from the nucleus to the outer edge of the electron cloud of an ion.

47
Q

What is the definition of ionic radius?

A

Ionic radius is the distance from the nucleus to the outer edge of the electron cloud of an ion.

48
Q

What is the definition of a cation?

A

A positively charged ion

49
Q

What is the definition of an anion?

A

A negatively charged ion

50
Q

When will a cation form a stronger ionic bond with an example?

A

When the ion is smaller with a higher charge the ionic bonds are stronger. for example, the size of the ions decreases in radius and increases in charge as you go from Na(+) to Al(3+)

51
Q

When will an anion form a stronger ionic bond?

A

When the ion has a larger ionic radius the electrons are more easily attracted by the positive nucleus of the other ion. For example, as you go from F(-) to O(2-) the strength of the ionic bond that forms increases.

52
Q

How do you draw a dot and cross diagram for the compound ion sulphate?

A
  • The bottom 2 oxygen atoms are bonded to the sulphur by a double bond
  • The top 2 oxygen atoms are bonded to the sulphur by a single bond
  • The extra 2 electrons gained to form the 2- ion are shown by an asterisk
  • The dot and cross diagram is surrounded by brackets and a 2- charge in the top right
53
Q

How do you draw a dot and cross diagram for the compound ion hydroxide?

A
  • The dot and cross diagram is surrounded by brackets and a - charge in the top right
  • The extra electron gained to form the - ion are shown by an asterisk
54
Q

How do you draw a dot and cross diagram for the compound ion carbonate?

A
  • The top 2 oxygen atoms are bonded to the carbon by a single bond
  • The extra 2 electrons gained to form the 2- ion are shown by an asterisk
  • The bottom oxygen atom is bonded to the carbon by a double bond
55
Q

Why does H+ form dative covalent bonds?

A

As it only has 1 proton and no electrons, therefore needing 2 electrons for a full outer shell

56
Q

Why does the bond strength decrease from HF to HI, using the idea of larger atoms to explain your answer?

A

As you move from HF to HI the atomic radius of the halogen increases as it gains more shells. There is also increased shielding between the shared pair of electrons and the adjacent halogen nucleus. Therefore the forces of attraction between the nucleus and the electron cloud decrease.

57
Q

Why is C-C bond weaker than C=C bond?

A

This is because the C-C bond only has one shared pair of electron and has a larger bond length in comparison to C=C which has two shared pairs of electrons and a smaller bond length, so the electrostatic forces of attraction between the electrons and nuclei are stronger

58
Q

What does the statement ‘the ionic compound has covalent character’ mean?

A

When a cation distorts an anion there is an increase in electron density in the space between the ions, some sharing of electrons and partial covalency

59
Q

Explain why negative ions are polarised by positive ions?

A

Positive cations attract the negative electron cloud of the anion

60
Q

Explain why smaller, more highly charged positive ions have a greater polarising effect?

A

If the positive ion is smaller with a large charge they have a greater charge density so a greater attractive force on the electron cloud

61
Q

What happens to the shared pair of electrons when atoms of different electronegativities bond covalently?

A

The shared pair of electrons is attracted to the more electronegative element. This causes the bond to become polarised and have a delta(+) and delta(-) charge

62
Q

What is the definition of a dipole?

A

A bond or molecule who’s ends have opposite charges/ a separation of positive and negative charge between two covalently bonded atoms.

63
Q

Explain the change in bond angle from 109.5° to 107° to 104.5° in methane, ammonia and water.

A
  • There is an increasing number of lone pairs
  • lone pair-lone pair repulsion is greater than lone pair-bonding pair repulsion
  • Therefore the bond angle is decreasing
64
Q

What is a polar bond?

A

A polar bond forms when the Electronegativities of the two atoms in the bond are different so a partial positive and a partial negative charge is gained

65
Q

Can a molecule have polar bonds but not be polar? If yes, why?

A

Yes, this is because if the polar bonds are symmetrical and act in opposite directions then the charges cancel out.

66
Q

How do you draw a dipole onto a diagram of the shape of molecule?

A

An arrow with a base from the partially positive to the partially negative atom

67
Q

What are London Forces?

A

They are the weakest type of intermolecular force that occurs between all molecules

68
Q

How do London Forces form?

A
  • In all molecules the electrons are moving constantly and randomly
  • This leads to fluctuations in the electron density in side the molecule
  • This then can cause parts of the molecule to become more or less negative to from an instantaneous/temporary dipole
  • These temporary dipoles can cause dipoles in neighbouring molecules called induced dipoles
  • The London Force is the force between the two dipoles
69
Q

What are the two factors that
effect the strength of London Forces?

A
  • The number of electrons in a molecule
  • The shape of a molecule
70
Q

What is the definition of polarisability?

A

Polarisability is an indication of the extent to which the electron cloud in a molecule (ion) can be distorted by a nearby electric charge.

71
Q

How does an increasing number of electrons result in stronger London Forces?

A

Bigger molecules that have a large number of electrons have a higher polarisability. Therefore, the possibility for temporary induced dipoles is greater so there are stronger London Forces present.

72
Q

Why do the melting and boiling points of halogens increase down the group?

A
  • As you go down the group the halogen molecules increase in size, therefore their are more electrons contributing to a larger electron cloud
  • The outer electrons become further from the nucleus so the electron cloud is more easily distorted forming more instantaneous dipoles
  • This leads to stronger London Forces being formed so more energy is required to overcome them
73
Q

Why do straight chain molecules have high melting and boiling points in comparison to branched chain molecules?

A

Straight chain molecules have a larger surface area in comparison to branched molecules which are more compact. Therefore, the straight chain molecules have a larger area over which London Forces can form so more energy is required to overcome them.

74
Q

Do London Forces act between all molecules?

A

Yes

75
Q

How do permanent dipole-dipole forces form?

A

These occur when polar molecules, ones that have permanent dipoles, attract each other. This is a more strong intermolecular force in comparison to London Forces.

76
Q

What are hydrogen bonds?

A

A strong type of permanent dipole-dipole interaction and the strongest type of intermolecular force.

77
Q

What 2 features must be present for hydrogen bonds to form between molecules?

A
  • A hydrogen atom must be covalently bonded to a small, highly electronegative atom that is N, O or F
  • There must be a lone pair on a second electronegative atom in a neighbouring molecule
78
Q

What are the 2 reasons for why Hydrogen Bonding is particularly strong in comparison to other permanent dipole-dipole forces?

A
  • There is a very large dipole, due to the difference in electronegativities, between the Hydrogen atom and N, O or F atom. This means that the covalent bond is extremely polarised.
  • The small size of the Hydrogen atom gives a high charge density from the partial positive charge that it has
79
Q

Why is the bond angle around the hydrogen, that is both covalently bonded to an atom and has a hydrogen bond with another molecule, 180°?

A

There are 2 pairs of electrons around the Hydrogen atom that are involved in the Hydrogen bond. These pairs of electrons repel to the position of minimum repulsion, as far apart as possible

80
Q

What are the elements you need to remember when drawing a diagram of hydrogen bonding?

A
  • Partially positive and negative charges on the H atom and O/N/F atom
  • Lone pair(s) of electrons on the atom involved in the hydrogen bond
  • Bond angle of 180° around the hydrogen involved in the Hydrogen bond
  • Dashed lines to represent the Hydrogen bond
81
Q

What gives ammonia, water and hydrogen fluoride their relatively high boiling temperatures?

A

Hydrogen Bonding

82
Q

Why does the density of water increase as it goes from a solid to a liquid?

A

Ice has a lower density than water. This is because, in a solid state, the Hydrogen bonding holds the water molecules in an open structure. As the ice melts this structure collapses, therefore the molecules get closer together causing the density to increase.

83
Q

Why do the boiling points of the noble gases increase as you go down the group?

A

As you go down the group the number of electrons increases, therefore there is a large electron cloud that distorts more easily to form instantaneous dipoles. Therefore, stronger London Forces are created between molecules so the energy required to break these forces increases.

84
Q

Why does hydrogen fluoride have a higher boiling point than the other hydrogen halides?

A

As it is the only hydrogen Halide containing Hydrogen bonds

85
Q

Why does the boiling point increase gradually as you move from HCl to HI?

A

Whilst the difference in electronegativity decreases so the strength of the permanent dipole-dipole forces decreases, this is smaller than the increase in the strength of the London Forces due to an increased number of electrons. Therefore, more energy is required to overcome these attractive forces.

86
Q

Why are the boiling points of all the hydrogen halides much higher than the Noble Gases?

A
  • Noble gases only contain London Forces
  • Hydrogen halides contain both London Forces and permanent dipole-dipole forces or hydrogen bonds. These types of intermolecular force and significantly stronger than London Forces so the addition of them means more energy is required to overcome the forces between molecules.
87
Q

When describing differences in boiling points, what points should you make?

A

London forces - affected by number of electrons and the shape of the molecule (branched or straight chain)
Permanent dipole-dipole - If the molecule is permanently polar (differing electronegativities)
Hydrogen bonds - strongest off all intermolecular forces

Some molecules will contain a more than one type of intermolecular force

88
Q

What is the rule for the solubility of solutes?

A

“Like dissolves Like”

This means that polar solvents dissolve polar solutes and vice versa

89
Q

Why are many ionic solids able to dissolve in water?

A
  • The ions in ionic solids are strongly hydrated by polar water molecules
  • The water molecules cluster around the ions and bind to them
  • In ionic solids, the energy released when the water molecules bind to the ions is enough to overcome the ionic bonding between the ions
90
Q

Why are other salts, not ionic solid, insoluble in water?

A

As the energy that is released when the water molecules attach to the ion, is not large enough to overcome the existing forces between the ions.

Chemistry A - Level (26 decks)

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