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A Level Chemistry > Bonding and Structure > Flashcards

Bonding and Structure Flashcards

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1
Q

Types of bonding

A

Ionic (or electrovalent), colvalent, intermolecular forces, hydrogen, metallic

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2
Q

How are ionic compounds arranged?

A

In a lattice structure, and the ions are arranged in such a way that maximises the attractive forces between the oppositely charged ions and minimises repulsive forces between similarly charged ions.

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3
Q

In what direction do forces act in ionic compounds?

A

In all directions

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4
Q

In what direction do forces act in covalent compounds?

A

One direction only

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5
Q

What is the definition of ionic bonding?

A

Ionic bonding is the electrostatic force of attraction between oppositely charged ions

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6
Q

What experiment gives evidence for ions?

A

Electrolysis of copper (II) chromate
CuCrO4 is green, but when electrolysed it produces copper (II) ions which are blue (at the cathode) and chromate ions which are yellow (at the anode)

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7
Q

What are the physical properties of ionic compounds?

A

Solid at room temperature, high melting and boiling points.

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8
Q

Why do ionic compounds have high melting and boiling points?

A

The electrostatic attractions between the oppositely charged ions are strong and require a lot of energy to break down

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9
Q

What is the solubility of ionic compounds?

A

Many ionic solids are soluble in water, but insoluble in non-polar solvents

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10
Q

What is the conductivity of ionic substances?

A

They do not conduct electricity when solid, but do when melted or dissolved in water. This is because they are split up, and the cations travel to the cathode, and the anions travel to the anode, and are discharged at the electrodes
In a solid, the ions are trapped in the lattice and are not free to move

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11
Q

What is a covalent bond?

A

A covalent bond is the electrostatic force of attraction between 2 nuclei and a shared pair of electrons between them

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12
Q

Facts about covalent giant structures

A

There are no molecules and no intermolecular forces

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13
Q

Melting and boiling points of covalent simple molecular substances

A

Low mp and bp, because the only forces that have to be overcome are the weak intermolecular forces which require very little energy to overcome, as none of the strong covalent bonds between the atoms need to be broken

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14
Q

Melting and boiling points of covalent giant lattice structures

A

Strong covalent bonds have to be broken, which require lots of energy to overcome, and so covalent giant lattice structures often have extremely high melting and boiling points

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15
Q

Mechanical strength of covalent simple molecular structures

A

Force of attraction between neutral molecules is usually very weak, and so when temperature is low enough for a molecular lattice to form then it will be weak and brittle

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16
Q

Mechanical strength of covalent giant lattice structures

A

Very hard if the covalent bonds extend in 3 dimensions (e.g diamond, silica etc), but if it only extends in 2 dimensions (e.g graphite, talc, slate) they are soft and flaky

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17
Q

Solubility of covalent simple molecular substances

A

Generally, molecular substances dissolve better in non-polar solvents (e.g petrol, paraffin) than in water

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18
Q

Solubility of covalent giant lattice structures

A

Insoluble in all solvents

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19
Q

Conductivity of covalent simple molecular substances

A

Cannot conduct electricity, as there are no ions that are free to move, and all electrons are held tightly in covalent bonds and cannot travel through the substance

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20
Q

Conductivity of covalent giant lattice structures

A

Electrons in these substances are localised within covalent bonds and so cannot move through the lattice so do not conduct electricity, however, graphite can conduct electricity because of its delocalised electrons

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21
Q

Bonding in diamond

A

Each carbon atom is covalently bonded to 4 other carbon atoms in a tetrahedral structure (bond angle is 109.4°)

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22
Q

Density of diamond vs graphite

A

Diamond is more dense than graphite

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23
Q

Bonding of graphite

A

Each carbon atom is covalently bonded to 3 other atoms - interlocking hexagons in layers
Layers are held together by weak intermolecular forces

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24
Q

Hardness of diamond

A

Very hard - often used for drillbits

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25
Q

Hardness of graphite

A

Soft - layers can slide over each other

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26
Q

Electrical conductivity of diamond

A

None

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27
Q

Electrical conductivity of graphite

A

Good - the delocalised electrons between the layers are mobile and can move through the structure

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28
Q

What are allotropes?

A

Different forms of the same element that exist in the same physical state

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29
Q

What are carbon’s allotropes?

A

Diamond, graphite and buckminsterfullerene/fullerenes

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30
Q

What is one use of fullerenes?

A

Carbon nanotubes

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31
Q

Properties of graphene

A

Good conductor of electricity (as good as copper), good conductor of heat (better than any other material), strong, transparent, flexible. Thinnest material known. Composites allow for the possibility of plastics that can conduct.

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32
Q

What is graphene?

A

Effectively a two-dimensional material - a one atom thick layer of carbon atoms

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33
Q

What atoms have pure covalent bonding?

A

Diatomic molecules with identical atoms

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34
Q

What causes a polar molecule?

A

If 2 atoms forming a covalent bond are dissimilaar then the attractions each pair of electrons have is probably different, and so the electrons are not shared equally, but will be closer to one atom, making the molecule polar

35
Q

What is electronegativity?

A

A numerical measure of the power of an atom to attract a pair of electrons in a covalent bond in a molecule

36
Q

What are the most electronegative elements?

A

1) Fluorine
2) Oxygen
3) Nitrogen
4) Chlorine

37
Q

How does electronegativity affect bonding?

A

No difference in electronegativity = completely covalent
Small difference in electronegativity = polar covalent bonding
Large difference in electronegativity = ionic bonding

38
Q

What are the trends of electronegativity on the periodic table?

A

Electronegativities increase from left to right across the periodic table and decrease down the group
Noble gases are not included

39
Q

HBr (shape and polarity)

A

Linear molecule

Has a dipole -> polar

40
Q

BCl3 (shape and polarity)

A

Trigonal planar

Has dipoles, but dipoles cancel out because the molecules are symmetrical so it is non polar

41
Q

NH3 (shape and polarity)

A

Trigonal pyramidal

Has dipoles which do not cancel out because the molecule is not symmetrical, so the molecule is polar

42
Q

H2O (shape and polarity)

A

V shaped

Has dipoles which do not cancel out, so molecule is polar

43
Q

CCl4 (shape and polarity)

A

Tetrahedral

Has diploes which cancel, so the molecule is non-polar

44
Q

CHCl3 (shape and polarity)

A

Tetrahedral

Dipoles are different sizes and so do not cancel, so molecule is polar

45
Q

What is polarisation influenced by?

A

Size of the anion - the larger the anion, the less control it has over its outer electrons and the more polarisable it is

46
Q

What is polarising power influenced by?

A

Charge of the cation - the smaller and more highly charged the cation, the more it will distort the outer electron cloud of the anion

47
Q

What gives ionic compounds covalent character?

A

Where the anion is polarised by the cation. The greater the degree of polarisation, the more the covalent character

48
Q

What are the 4 types of intermolecular forces?

A
  • Permanent dipole - permanent dipole forces
  • Permanent dipole - induced dipole forces
  • Instanteous dipole - induced dipole / London dispersion forces
  • Hydrogen bonds
49
Q

How do permanent dipole - permanent dipole forces of attraction work?

A

Molecules with a permanent dipole have a slightly positive and a slightly negative end. The slightly positive end attracts the slightly negative end, causing a permanent dipole - permanent dipole force of attraction

50
Q

How do induced dipole - permanent dipole forces of attraction work?

A

The molecule with a permanent dipole attracts a molecule without a dipole, inducing a dipole in the other molecule, as it repels elections to the other side of the molecule
It only happens in mixtures

51
Q

How do London forces work?

A

In molecules, the distribution of electrons can be uneven in the molecule at any given moment, and, by chance, there may be a momentarily greater electron density at one end of the molecule than the other. The instantaneous dipole has the effect of inducing a second dipole on a neighbouring molecule, resulting in attraction between the molecules. The more electrons in a molecule, the more likely it is that the distribution will be uneven, so the stronger the London forces

52
Q

How does the shape of a molecule influence the boiling point?

A

The surface area of contact between the molecules will determine the boiling point of the molecule if they are isoelectronic. “Sausage” shaped molecules have a greater surface area of contact than branched or spherical molecules, and so will have a higher boiling point

53
Q

What is a hydrogen bond?

A

A hydrogen bond is the electrostatic force of attraction between the poorly shielded proton of a hydrogen atom bonded to a small, highly electronegative atom such as nitrogen, oxygen, or fluorine, and a lone pair of electrons on a N, O or F atom on a neighbouring molecule

54
Q

Why does ice float on water?

A

The presence of 2 hydrogen atom and 2 lone pairs in each water molecule results in a 3D tetrahedral structure. Each O atom in ice is surrounded tetrahedrally by 4 others; 2 covalently bonded and 2 hydrogen bonded.
This creates a very ‘open’ structure, maximising the hydrogen bonding, giving it a lower density than water, explaining why it floats
When ice melts, the lattice breaks up, and water molecules are able to pack more closely, and in liquid water hydrogen bonds are constantly being broken and reformed

55
Q

What determines the differences in boiling points (e.g in pentane and dimethyl propane)?

A

They have the same number of electrons, so difference only depend on the shape of the molecules (as intermolecular forces are the same strength). Pentane is unbranched (sausage shaped), and has a greater surface area of contact than the branched (spherical) dimethyl propane molecule, giving pentane a higher bp

56
Q

What determines the differences in meltingpoints (e.g in pentane and dimethyl propane)?

A

Branched (spherical) molecules pack together better in the solid, which is more important than the surface area of contact, so dimethyl propane has a higher melting point

57
Q

What is boiling point determined by?

A

Strength of intermolecular forces and surface area of contact between molecules

58
Q

What is melting point determined by?

A

Strength of intermolecular forces and how closely packed the molecules are in the solid state

59
Q

What is needed for an ionic substance to dissolve?

A

A lot of energy needs to be put in to separate the ions in the solid lattice (reverse lattice energy) and some to break intermolecular forces between the molecules. Energy is given out when the water molecules form bonds with individual ions, which is called the solvation energy, or hydration energy specifically for water. If more energy is given when bonds are formed than is put in when bonds are broken, then the process is likely to be favourable and it will dissolve (a decrease in enthalpy/exothermic processes are often favoured)

60
Q

Why is water a good solvent for ionic compounds?

A

Water has a dipole, so it can solvate the positive and negative ions and can form strong ion-dipole bonds, releasing lots of energy, meaning they are usually soluble.

61
Q

Why are non-polar solvents not good to dissolve ionic compounds?

A

They are not strongly attracted to ions and so there is no highly exothermic hydration energy to compensate for the reverse lattice energy, so the process is unfavourable as it is overall endothermic

62
Q

Why do covalent simple molecular substances not dissolve in water?

A

The bonds broken are the ones in the simple molecular substance, which requires London forces to be overcome, and the hydrogen bonds in the water. The bonds that would need to be made are dipole - dipole forces OR London forces in the new mixture.
Hydrogen bonds are stronger than the new intermolecular forces and so process is likely to be endothermic overall and therefore unfavourable.

63
Q

When are covalent simple molecular structures able to dissolve in water?

A

1) If the covalent substance reacts with water to form ions, then the substance will dissolve. This happens with ammonia (which reacts to form an ammonium ion and a hydroxide ion), and hydrogen chloride (which reacts to form a hydroxonium ion and a chloride ion). Both of these are soluble in water.
2) When the water molecules can form hydrogen bonds with the covalent substance. This happens with substances such as methanol, ethanol and propanol. The O atom hydrogen bonds to the H atom of a water molecule

64
Q

Why do covalent simple molecular substances dissolve in non-polar solvents?

A

Energy needs to be put in to overcome the intermolecular forces between molecules of the covalent substance and the intermolecular forces between the solvent molecules
Energy is given out when bonds are formed between molecules of the covalent substance and the solvent molecules

These bonds have similar strengths, so enthalpy change is approximately 0, but there is an increase in entropy (disorder), and so it is favoured

65
Q

Why is propanone such a good solvent?

A

Propanone has dipole - dipole and London forces and therefore can interact with a variety of substances

66
Q

What is the octet rule?

A

“Every atom combines in such a way that its outer shell achieves the stable configuration of 8 electrons (or 2 electrons in H, Li, Be); that is, it tends to adopt the electrnic configuration of the nearest noble gas”

67
Q

What are the exceptions to the octet rule?

A

Some stable cations have 18 electrons in their outer shell (Cu+, Zn2+, Hg2+)
Some stable cations have 2 electrons in their outer shell, preceded by 18 (Pb2+, Bi3+)
Transition metal cations have between 9 and 17 electrons in their outer shell (Fe2+ (14), Fe3+ (13), Cr2+ (12))
Covalently, some substances have less than an octet and are electron deficient (BeCl2, BCl3)
Some covalently bonded substances have more than an octet as a central atom can ‘expand its octet’ and use empty, low energy d orbitals for bonding (PF5, SF6, SO3, SO2)

It is also broken when odd electron molecules are formed (NO, ClO2, NO2)

68
Q

What is a dative covalent bond?

A

A dative covalent bond is the electrostatic attraction of 2 nuclei for a shared pair of electrons, both electrons in the pair coming from ONE of the atoms forming the bond.

69
Q

How can the shape of molecules be arranged / how do electrons organise themselves?

A

Electron pairs around a central atom maximise their distance apart to minimise repulsion (and therefore to reach their lowest potential energy state)

70
Q

Spacing of bonding pairs and lone pairs of electrons?

A

The space taken up by a lone pair of electrons is greater than that taken up by a bonding pair. Lone pairs repel other lone pairs more than lone pairs repel bonding pairs more than bonding pairs repel other bonding pairs

71
Q

Changes to bond angles

A
  • Each lone pair reduces bond angle by 2°
  • The space taken up by a double bonding pair of electrons is only slightly greater than the space taken up by a bonding pair of electrons
  • The space taken up by an electron pair also depends on the electronegativity (more electronegative atom pulls electrons from the covalent bond towards the centre which increases the bp-bp repulsion)
72
Q

What is metallic bonding?

A

The attraction of a lattice of positive ions and a sea of delocalised electrons

73
Q

How can the strength of metallic bonding be worked out?

A

The ratio: number of valency electrons / metallic radius

74
Q

Why are metals malleable and ductile?

A

When a stress is applied to the metal lattice, the layers of cations are able to slide over each other, and on a large scale, these atomic slips account for the ductility and malleability of metal

75
Q

Why are metals able to conduct electricity?

A

They have delocalised electrons that are free to move throughout the lattice

76
Q

Why are metals good thermal conductors?

A

The atoms are packed closely together so the kinetic energy of the vibrating atoms can be passed through the lattice easily

77
Q

What are the trends in mp and bp in metals?

A

Group 1 - 1 valency electron, large radius, so weak metallic bonding and low melting point. Metallic radius increases down the group and melting point decreases down the group
Group 2 - 2 valency electrons, smaller radius than group 1, so stronger bonding. Metallic radius increases down the group, so they should decrease, but this is not completely regular due to the difference in packing of the solid
Transition metals - have 3d and 4s electrons involved in bonding, therefore a much higher mp than group 1 or 2

78
Q

How can an electron density map show that a compound is completely ionically bonded?

A

The electron density will drop to 0, and there will be areas with no lines
If a substance is not 100% ionically bonded, then the electron density will not drop to 0 anywhere

79
Q

How can electron density maps show covalent bonding?

A

The electron density between the 2 nuclei does not drop to 0, and electron density between the nuclei will be quite high as the attraction of the 2 nuclei for this high concentration of negative charge that holds the 2 atoms together.

80
Q

What type of overlaps create sigma orbitals?

A
  • 2 1s orbitals (e.g. in hydrogen)
  • an overlap of s and p orbitals (e.g. hydrogen fluoride)
  • a pair of p orbitals (e.g. in fluorine)
    A sigma orbital is always formed when 2 p orbitals approach along the line joining the nuclei of their atoms

Sigma orbitals show a significant electron density between the nuclei.

81
Q

What type of overlaps create pi orbitals?

A
  • A pair of p orbitals at right angles to the line joining the nuclei overlap to form a pi orbital
    They produce regions of electron density above and below the axis joining centres
82
Q

What has to already be present for a pi orbital to form?

A

A sigma orbital.

83
Q

How effective are pi orbitals at holding atoms together?

A

Less effective than sigma orbitals. The sigma + pi combination is less than twice as strong as the sigma alone

84
Q

How are triple bonds formed?

A

A sigma orbital and 2 pi orbitals

A Level Chemistry (25 decks)

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