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A-Level Edexcel Chemistry > Chapter 2 - Steven > Flashcards

Chapter 2 - Steven Flashcards

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1
Q

What are ions?

A

Positvely or negatively charged atoms

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2
Q

If sodium forms Na+, what has it done?

A

Lost 1 electron to form a full outer shell.

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3
Q

What is an ionic bond?

A

The strong electrostatic attraction between 2 oppositely charged ions.

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4
Q

How do ionic charges affect ionic bonding?

A

The greater the charge on an ion, the stronger the ionic bond

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5
Q

What is the relationship between the strength of the ionic bond and the melting/boiling points?

A

The stronger the ionic bond, the higher the melting/boiling point

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6
Q

What happens to the electrostatic attraction as the ionic radii?

A

It decreases

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7
Q

Why does electrostatic attraction get weaker with distance?

A

The ions sit further apart, so there is less electrostatic attraction.

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8
Q

What happens to the ionic radius as you go down a group?

A

It increases

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9
Q

What happens to the ionic radius as you go across a period?

A

It decreases

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10
Q

What do dot and cross diagrams show?

A

Where the electrons in a bond came from

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11
Q

What do ionic compounds form?

A

Giant Ionic lattice structures

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12
Q

What is a lattice?

A

A regular structure

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13
Q

What is a giant lattice?

A

A structure made up of the same basic unit repeated over and over again

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14
Q

Why do ionic compounds form giant lattice structures?

A

Because each ion is electrostatically attracted in all directions to ions of the opposite charge

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15
Q

As ionic compounds have high melting/boiling points, what does this provide evidence for?

A

Their strong forces of attraction between positive and negative ions.

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16
Q

What shows that the particles in ionic compounds are charged?

A

They are often soluble in water but not in non-polar solvents. The ions are pulled apart by polar molecules like water, but not by non-polar molecles

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17
Q

As ionic compounds can’t conduct when solid, but can when molten, what does this tell you?

A

That they are ions, which are in fixed positions by strong ionic bonds when solid, but are free to move when molten.

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18
Q

Why are ionic compounds brittle?

A

If you ever tried to pull layers over each other, you’d get negative chlorine ions directly over negative chlorine ions so there would be a very high repulsion, so they are brittle

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19
Q

When you electrolyse a green solution of copper chromate on a piece of wet filter paper, what happens?

A

The filter turns blue at the cathode (negative) and yellow at the anode (positve).
Copper fromed at the cathode and chromate formed at the anode

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20
Q

What is isoelectronic?

A

Having the same number of electrons

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21
Q

What is a covalent bond shared by?

A

2 or more atoms bonded together

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22
Q

If the atoms ends up with 8 electrons in its outer shell, is it stable or unstable?

A

It’s very stable

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23
Q

What is a bond containing 2 electron pairs called?

A

A double bond

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24
Q

In covalent molecules, what is the positive nuclei attracted to?

A

The area of electron density between the 2 nuclei.

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25
Q

In covalent bonds, what is the repulsive force?

A

The 2 positively charged nuclei repelling, and the sets of electrons repelling

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26
Q

What is the bond length?

A

The distance between the 2 nuclei in a covalent bond

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27
Q

What is the relationship between bond enthalpy and bond length?

A

The shorter the bond length, the higher the bond enthalpy

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28
Q

Why does the bond enthalpy increase as the bond length decreases?

A

The electron density between the nuclei increases, due to the region being smaller.

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29
Q

What is dative covalent bonding?

A

When both electrons come from 1 atom

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30
Q

What gives Al a full outer shell when bonded with chlorine?

A

When 2 AlCl3 molecules combine to form Al2Cl6. Once Cl in each of the molecules donates a lone pair to the Al and the other molecule, forming 2 dative bonds

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31
Q

How many dative bonds are formed in Al2Cl6?

A

2 are formed

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32
Q

What does molecular shape depend on?

A

The electron pairs around the centre atom

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33
Q

As electrons are negatively charged, what do electron pairs do?

A

Repel each other

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34
Q

What repel more, lone pairs or bonding pairs?

A

Lone pairs

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35
Q

When are bond angles between bonding pairs reduced?

A

When they are pushed together by lone pair repulsion

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36
Q

What is the way of predicting molecular shape known as?

A

Electron pair repulsion theory

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37
Q

What reduces the bond angle more, 2 lone pairs of electrons or 1 lone pair of electrons?

A

2 lone pairs of electrons

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38
Q

If there are 2 electron pairs around the central atom, what will the bond angles be?

A

180*

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39
Q

If there are 2 electron pairs around the central atom, what is it known as?

A

A linear molecule

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40
Q

If there is 3 electron pairs around an atom, with no lone pairs, what is it known as? What are the bond angles?

A

Trigonal Planar

180*

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41
Q

If there is 3 electron pairs around an atom, with 1 lone pair of electrons, what is it known as? What is the bond angle?

A

Non-linear or bent

119*

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42
Q

If there are 4 electron pairs around an atom, what is the bond angle?

A

109.5*

43
Q

As the number of lone pairs increases by 1 in an atom with 4 electron pairs, how much does the bond angle decrease by?

A

2.5*

44
Q

What is it called when ther are 4 electron pairs around an atom with no lone pairs?

A

Tetrahedral

45
Q

What is it called when there are 4 electron pairs around an atom with 1 lone pair?

A

Trigonal pyramidal

46
Q

What is it called when there are 4 electron pairs around an atom with 2 lone pairs?

A

Non-linear or bent

47
Q

What is it called when there are 5 electron pairs around an atom with no lone pairs?

A

Trigonal bipyramidal

48
Q

What is it called when there are 6 electron pairs around an atom with no lone pairs?

A

Octahedral

49
Q

When do covalent bonds form?

A

When atoms share electrons with other atoms

50
Q

Give 2 examples of giant structures of covalently bonded atoms:

A

Carbon and silicon

51
Q

In diamond, how many covalent bonds does each carbon form?

A

4

52
Q

What compound forms a similar lattice arrangement to diamond?

A

Silicon dioxide

53
Q

How does silicon dioxide form a similar lattice arrangement to diamond?

A

Oxygen atoms between each silicon atom

54
Q

As giant structures have high melting/ boiling points, how does this provide evidence for covalent bonding?

A

You need to break a lot of strong bonds before the substance melts, which requires a lot of energy.

55
Q

As giant structures can’t conduct when solid, what does this show?

A

All bonding electrons are held in covalent bonds

56
Q

Why can graphite conduct electricity?

A

Each carbon atom forms only 3 covalent bonds.

57
Q

What is graphene?

A

One layer of graphite

58
Q

How thick is the sheet of graphene?

A

1 atom thick

59
Q

Why can graphene conduct?

A

The delocalised electrons are free to move along the sheet

60
Q

What are delocalised electrons?

A

Electrons that are free to move

61
Q

In metals, what are the positive metal ions attracted to?

A

The negative delocalised electrons

62
Q

What is electronegativity?

A

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

63
Q

What is the most electronegative element on the periodic table?

A

Fluorine

64
Q

What happens to electronegatvity as you go across a period?

A

It increases

65
Q

What happens to electronegativity as you go down a group?

A

It decreases

66
Q

In a covalent bond, if the elements have similar or identical electronegativities, where will the electrons sit?

A

Midway between the 2 nuclei

67
Q

If 2 elements have similar electronegativities, will the bond be polar or non-polar?

A

Non-polar

68
Q

If a covalent bond is between elements with very different electronegatvities, where will the electrons sit?

A

They will be pulled more towards the more electronegatvie element, so there will be a charge across the bond

69
Q

In a polar bond, what does the difference in electronegativity cause?

A

A dipole

70
Q

What is a dipole?

A

A difference in charge between the 2 atoms caused by a shift in electron density in the bond

71
Q

What is a polar molecule?

A

A molecule that has an overall dipole, which is a dipole caused by the presence of a permanent charge across the molecule

72
Q

If the polar bonds are arranged so they point in opposite directions, ios the molecule polar?

A

No

73
Q

If the polar bonds roughly all point in the same direction, will the molecule be polar?

A

Yes

74
Q

Is H-Cl a polar molecule?

A

Yes

75
Q

Is a molecule that has polar bonds always a polar molecule?

A

No, for example 0=C=O

76
Q

What are the 3 main types of intermolecular forces?

A

London forces
Permanent dipole-permanent dipole bonds
Hydrogen bonding

77
Q

What do all atoms and molecules form?

A

London forces

78
Q

Describe a london force:

A
  • At any moment, the electrons in an atom are more likely to be on one side than the other (temporary dipole)
  • This dipole can induce a temporary dipole on a neighbourring atom. The 2 dipoles are then attracted to each other
79
Q

What is a temporary dipole?

A

When the electrons in an atom are more likely to be on one side than the other

80
Q

What is a simple molecular structure?

A

When atoms are held together in pairs and held together in a molecular lattice by weak London forces

81
Q

What happens to the strength of the London forces as the molecular size increases?

A

They increase, as larger molecules have larger electron clouds

82
Q

Why does the strength of London forces increase as the length of the carbon chain increases?

A

There’s more molecular surface contact and more electrons to interact

83
Q

What type of molecules have permanent dipole-permanent dipole bonds?

A

Polar molecules

84
Q

Give an example of a permanent dipole-permanent dipole bond:

A

H-Cl——H-Cl——–H-Cl

85
Q

Why do permanent-dipole-permanent dipole bonds generally have higher melting points than those with just London forces?

A

Because molecules that have permanent dipole-permanent dipole attraction also have London attraction aswell!

86
Q

What is the strongest intermolecular force?

A

London bonding

87
Q

What is the only case where hydrogen bonding occurs?

A

When hydrogen is covalently bonded to fluorine, nitrogen or oxygen.

88
Q

What happens in hydrogen bonding?

A
  • Fluorine, nitrogen and oxygen are very electronegative, so draw electrons from the hydrogen atom
  • The bond is polarised, and hydrogen has such a high charge density that the hydrogen atoms form weak bonds with lone pairs of electrons on the fluorine, nitrogen or oxygen
89
Q

Give an example of a compound where hydrogen bonding is present?

A

Water

90
Q

How does the boiling points of group 7 hydrides show the strength of hydrogen bonds?

A

Hydrogen fluoride has the highest boiling points compared to the other group 7 hydrides

91
Q

How does the boiling points of group 6 hydrides show the strength of hydrogen bonds?

A

Water has the highest boiling points, which is the only molecule that contains hydrogen bonds

92
Q

What type of structure does ice involve?

A

Simple molecular structure

93
Q

In ice, how are the water molecules arranged?

A

So that there is a maximum number of hydrogen bonds, which wastes space

94
Q

As ice melts, what happens?

A

Some of the hydrogen bonds are broken and the lattice breaks down, allowing molecules to ‘fill the spaces’

95
Q

Why are alcohols less volatile than similar alkanes?

A

Hydrogen bonding gives alcohols low volatilties (high boiling points) compared to non-polar compounds

96
Q

For one substance to dissolve in another, what 3 things have to happen?

A
  • Bonds in substance have to break
  • Bonds in solvent have to break
  • New bonds have to form between substance and solvent
97
Q

What are polar solvents made up of?

A

Polar molecules, such as water

98
Q

Give an example of a non-polar solvent:

A

Hexane

99
Q

What do ionic substances dissolve in?

A

Polar solvents

100
Q

How do ionic substance dissolve in polar solvents?

A
  • The ions in the ionic substance are attracted to the oppositily charged ends
  • The ions are pulled away by the solvent, which surround the ions
  • This is called hydration
101
Q

Why do alcohols dissolve in polar solvents, such as water?

A

The polar O-H bond in an alcohol is attracted to the 0-H bond in water

102
Q

Why can’t halogenoalkanes dissolve in water?

A

Their dipoles aren’t strong enough to form hydrogen bonds with water
The hydrogen bonding between water molecules is stronger than the bonds that would have been formed with halogenoalkanes, so halogenoalkanes don’t dissolve

103
Q

What do non-polar substances dissolve best in?

A

Non-polar solvents, because they form similar bonds

A-Level Edexcel Chemistry (24 decks)

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