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Chemistry > Structure > Flashcards

Structure Flashcards

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1
Q

What is electronegativity?

A

The power of an atom to attract the pair of electrons in a covalent bond towards itself

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2
Q

What is the electron distribution in a covalent bond between elements with different electronegativities like?

A

Unsymmetrical

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3
Q

Where does this phenomenon arise from? - (What is the electron distribution in a covalent bond between elements with different electronegativities like?)

A

From the positive nucleus’s ability to attract the negatively charged electrons, in the outer shells, towards itself

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4
Q

What scale is used to assign a value of electronegativity to an atom?

A

The pauling scale

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5
Q

Which is the most electronegative atom?

A

Fluorine with a value of 4.0 on the pauling scale

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6
Q

When is fluorine best at attracting electron density towards itself?

A

When covalently bonded to another atom

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7
Q

What is the nuclear charge?

A

Attraction exists between the positively charged protons in nucleus and negatively charged electrons in energy levels of atom

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8
Q

What does an increase in the number of protons lead to?

A
  • An increase in nuclear attraction for electrons in outer shells
  • So, an increased nuclear charge results in an increased electronegativity
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9
Q

What is the atomic radius?

A

The distance between the nucleus and electrons in the outermost shell

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10
Q

What happens when electrons are closer to the nucleus in terms of atomic radius?

A
  • They are more strongly attracted towards its positive nucleus and vice versa
  • So an increased atomic radius results in a decreased electronegativity
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11
Q

What effect do filled energy levels have on electronegativity?

A
  • They shield the effect of the nuclear charge so outer electrons are less attracted to the nucleus and will experience less of the attractive force
  • So it will result in a decreased electronegativity
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12
Q

What is the trend of electronegativity down a group?

A
  • It decreases
  • Nuclear charge increases but there is increased shielding so atomic radii decrease in size
  • So, there is a decrease in attraction between the nucleus and outer bonding electrons
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13
Q

What is the trend of electronegativity across a period?

A
  • It increases
  • Nuclear charge increases and shielding remains constant so greater attraction between bonding electrons, resulting in smaller atomic radii
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14
Q

What happens when 2 atoms in a covalent bond have the same electronegativity?

A

The covalent bond is non-polar as the electrons are shared equally between the 2 atoms

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15
Q

What will dictate what type of polar bond is formed?

A

The difference in electronegativites

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16
Q

When is an ionic bond formed?

A

When the values are very different (more than 1.7), ions will be formed

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17
Q

When is a covalent bond formed?

A
  • When the value is between 0.3 and 1.7. The bond will be polar
  • The electrons will be drawn towards the more electronegative atom
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18
Q

What happens when a covalent bond is polar?

A
  • The negative charge centre and positive charge centre do not coincide with each other
  • So electron distribution is asymmetric
  • The less electronegative atom gets a partial charge of delta positive
  • The more electronegative atom gets a partial charge of delta negative
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19
Q

What does a greater difference in electronegativity mean?

A

The more polar the bond becomes

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20
Q

How do we determine whether a molecule with more than 2 atoms is polar?

A

We need to consider:
- The polarity of each bond
- How the bonds are arranged in the molecule

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21
Q

Why do some molecules that have polar bonds end up becoming non - polar?

A

The polar bonds in the molecule are arranged in such a way that the individual dipole moments cancel each other out

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22
Q

What are intramolecular forces?

A

Forces within a molecule and are usually covalent bonds
(e.g. single, double, triple and co-ordinate bonds)

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23
Q

What are intermolecular forces?

A

Forces between molecules

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24
Q

What are the 3 types of intermolecular forces?

A
  • Induced dipole-dipole forces (also known as van der walls or london dispersion forces)
  • Permanent dipole-dipole forces
    (are the attractive forces between 2 neighbouring molecules with a permanent dipole)
  • Hydrogen bonding (special type of permanent dipole-dipole force)
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25
Q

Which type of force is stronger?

A

Intramolecular forces

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26
Q

What is the order of the strength of the forces from strongest to weakest?

A

Covalent bonding
Hydrogen bonding
Permanent dipole-dipole
Instantaneous dipole - induced dipole

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27
Q

How do instantaneous dipole-induced dipole forces occur?

A
  • The electron charge cloud in non-polar molecules or atoms are constantly moving
  • During this movement, the electron charge cloud can be more on one side than the other
  • This causes a temporary dipole to arise
  • This temporary dipole can induce a dipole in neighbouring molecules
  • The, the positive end of the dipole in 1 molecule and the negative end in another are attracted towards each other
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28
Q

Why are the dipoles in instantaneous dipole-induced dipole only temporary?

A

because the electron clouds are constantly moving

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29
Q

For small molecules with the same number of electrons, which force is stronger (permanent or induced)?

A

Permanent dipoles
e.g. Butane and propanone have the same number of electrons but butane has induced dipole and propanone has permanent dipole so has a higher b.p. as more energy is required to break the intermolecular forces

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30
Q

What kind of intermolecular forces will polar molecules have?

A

Permanent dipoles - it will always have a negatively and positively charged end

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31
Q

What is needed for hydrogen bonding to take place?

A

A species which has an O, N, or F atom bonded to a hydrogen

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32
Q

What happens when hydrogen is covalently bonded to an O, F or N?

A
  • The bond becomes highly polarised
  • The H becomes so delta positively charged that it can form a bond with the lone pair of an O, N or F atom in another molecule
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33
Q

Give an example of hydrogen bonding

A

Water can form 2 hydrogen bonds, because the O has 2 lone pairs

34
Q

What are the examples of compounds that can form hydrogen bonds?

A

Alcohols
Ammonia
Amines
Carboxylic acids
Hydrogen fluoride
Proteins

35
Q

Why does water have high melting and boiling points?

A
  • Due to the strong intermolecular forces of hydrogen bonding between the molecules in both ice and water
  • A lot of energy is therefore required to separate the water molecules and melt or boil them
36
Q

What is the enthalpy of vaporisation?

A

The energy required to boil a substance

37
Q

What does the high enthalpy change of vaporisation of water suggest?

A

Instantaneous dipole-induced dipole forces are not the only forces present in the molecule - there are also strong hydrogen bonds which cause the high b.p.

38
Q

What kind of surface tension does water have?

A

A high one

39
Q

What is surface tension?

A

The ability of a liquid surface to resist any external forces (i.e. to stay unaffected by forces acting on the surface)

40
Q

What happens to water molecules at the surface of a liquid?

A
  • The molecules are bonded to other water molecules through hydrogen bonds
  • These molecules pull the surface molecules downwards, causing the surface of them to become compressed and more tightly together at the surface
  • This increases water’s surface tension
41
Q

Why are solids denser than liquids?

A

The particles are more closely packed together

42
Q

What are water molecules packed into as a solid?

A

A lattice

43
Q

Why are the water molecules in a solid slightly further apart than in the liquid form?

A

This way of packing the molecules (into a lattice) and the relatively long bond lengths of the hydrogen bonds means the molecules are slightly further apart

44
Q

Does ice have a lower density than liquid water?

A

Yes, by about 9%
The ‘more open’ structure of molecules in ice causes it to have a lower density than water

45
Q

What does a larger surface area mean in a molecule?

A

It will have more contact with adjacent molecules

46
Q

How is the surface area of a molecule reduced?

A

By branching

47
Q

What does a greater ability to induce a dipole in an adjacent molecule result in?

A

Greater London forces and higher melting and boiling points

48
Q

Give an example of how surface area affects the boiling point

A

Both pentane and 2-dimethylpropane have the same number of electrons but pentane has a larger surface area and so has a higher b.p. as more London forces are induced

49
Q

What does a greater number of electrons or molecular mass in a molecule result in?

A

A greater likelihood of a distortion and thus the greater the frequency and magnitude of the temporary dipoles

50
Q

How does the number of electrons affect the melting and boiling points?

A

The greater the number, the stronger the dispersion forces between molecules and also the enthalpy of vaporisation, making the melting and boiling points higher

51
Q

What is the correlation between the number of electrons and energy needed to break forces of attraction?

A

As the number of electrons increases, more energy is needed to overcome the forces of attraction

52
Q

Do alcohols have hydrogen bonds?

A

Yes, there are O-H bonds present and so hydrogen bonds set up between the slightly positive H atoms and lone pairs on oxygens in other molecules

53
Q

Why are the H atoms slightly positive in an alcohol?

A

The bonding electrons are pulled away from them towards the very electronegative oxygen atoms

54
Q

What are the only intermolecular forces in alkanes?

A

temporary induced dipole-dipole forces

55
Q

Why is the b.p. of an alkane lower than its respective alcohol?

A

Hydrogen bonds in an alcohol are stronger than the induced dipole-dipole forces in an alkane and so it takes more energy to separate alcohol molecules than alkane molecules

56
Q

What are the boiling points of the first 4 hydrogen halides?

A

HF - 293 K
HCl - 188 K
HBr - 207 K
HI - 238 K

57
Q

What trend do the rest of the boiling points of the rest of the hydrogen halides follow?

A

They increase

58
Q

Why do the boiling points of hydrogen halides increase as you go down the group?

A
  • The molecules become larger
  • The extra electrons allow greater temporary dipoles and so increase the amount of London dispersion forces between the molecules
59
Q

Why does HF have a higher boiling point?

A

It has hydrogen bonding between the H-F molecules. F molecules have small intense lone pairs and H is very positive so bonds form

60
Q

What is the general principle for solubility?

A

‘like dissolves like’

61
Q

What do non-polar substances dissolve in?

A

Non-polar substances such as hydrocarbons and they form dispersion forces between the solvent and the solute

62
Q

What do polar covalent substances generally dissolve in?

A

Polar solvents as a result of dipole-dipole interactions or the formation of hydrogen bonds between the solute and the solvent

63
Q

Give an example of when polar substances are dissolved

A

Between alcohols and water, hydrogen bonds form between the hydrogen on the water and the oxygen on the alcohol

64
Q

What happens when covalent molecules become larger?

A

Their solubility can decrease as the polar part of the molecule is only a smaller part of the overall structure
(this effect is seen in alcohols e.g. C2H5OH is readily soluble whereas C6H13OH is not)

65
Q

Why are polar covalent substances not able to dissolve well in non-polar solvents?

A

Their dipole-dipole interactions are unable to interact well with the solvent

66
Q

Why do giant covalent substances generally not dissolve in any solvents?

A

The energy needed to overcome the strong covalent bonds in the lattice structures is too great

67
Q

What can covalent compounds be arranged in?

A

Simple molecular lattices
Giant molecular

68
Q

What do covalent substances tend to have?

A

Small molecular structures known as simple molecules e.g. Cl2, H2, CO2

69
Q

What does iodine exist as?

A

It is a simple molecule but it exists as a crystalline structure involving a regular structure held together by weak london dispersion forces

70
Q

Why do giant covalent lattices have very high melting and boiling points?

A

They have a large number of covalent bonds linking the whole structure
A lot of energy is required to break the lattice

71
Q

Why is graphite soft?

A

Forces between the carbon layers are weak

72
Q

Why is diamond and silicon oxide hard?

A

It is difficult to break their 3D network of strong covalent bonds

73
Q

Are covalent compounds soluble in water?

A

No

74
Q

Why does graphite conduct electricity?

A

It has delocalised electrons between the carbon layers which can move along the layers when a voltage is applied

75
Q

Why can’t diamond and silicon oxide conduct electricity?

A

All 4 outer electrons on every carbon atom are involved in a covalent bond so no freely moving electrons available

76
Q

What is the structure of diamond called?

A

A tetrahedron

77
Q

Give details about graphite?

A
  • Forms layers of hexagons
  • Free electrons migrate between the layers
  • Layers can slide over each other
78
Q

Are all covalent bonds in diamond identical?

A

Yes and there are no intermolecular forces

79
Q

What kind of dimensional molecule is graphene?

A

It is essentially a 2D molecule as it is one atom thick

80
Q

Why are the very unusual properties of graphene useful?

A

They make it useful in fabricating composite materials and in electronics

81
Q

When are simple covalent molecules soluble?

A

When they are polar

82
Q

Are metals soluble?

A

No, but some may react

Chemistry (6 decks)

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