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A Level Chemistry > Module 2 - chapters 5 and 6 > Flashcards

Module 2 - chapters 5 and 6 Flashcards

(not including electrons)

1
Q

Give a general overview about ionic bonding

A
  • Between metals and non-metals
  • Ions have opposite charges and so are attractive
  • OILRIG
  • Ions exchange electrons to achieve a full outer shell and a stable electron configuration
  • the ions form a giant ionic lattice which has high melting and boiling points due to strong intramolecular forces and electrostatic forces
  • The ions are conductive if molten or aqueous because the ions are free to move
  • the ions can be separated using electrolysis
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2
Q

What is a key feature of strong ionic bonding?

A

A large difference in charge/a larger charge.

E.g. Al³⁺ and O²⁻ has a larger difference in charge than Na⁺ and Cl⁻ so aluminium oxide is stronger than sodium chloride.

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3
Q

Why are ionic substances typically brittle?

A

As you try to move the ions, the ions of the same charge are brought side-by-side and so the crystal repels itself apart

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4
Q

What does giant mean in terms of a giant ionic structure?

A

When you can’t state how many ions there actually are in the structure

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5
Q

What does it mean if an ion is 6:6 co-ordinated?

A

The ion is the centre is being touched by 6 other ions and those ions touched by 6 other ions to form a giant structure.

For example: a sodium ion in the centre of a giant ionic lattice is being touched by 6 chloride ions and each chloride ion by 6 sodium ions.

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6
Q

Do all ionic substances have high or low melting and boiling points?

A

High

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7
Q

How do the sizes of ions effect the attractions of ions?

A

If the ions are smaller they get closer together and so the electrostatic attractions are greater. This is because the charge isn’t spread over a large density.

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8
Q

What defines how soluble an ionic substance is?

A
  • the nature of the ionic lattice that needs breaking down
  • the attraction of the water molecules (or other polar solvents) to the ions
  • ionic lattice strength dominates, solubility decreases with increasing ionic charge
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9
Q

What is said when an ion has a full outer shell?

A

It is isoelectronic to a noble gas

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10
Q

What is a hydrogen ion also known as?

A

A proton

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11
Q

Why do covalently bonded atoms stay together?

A

The positive nuclei are attracted to the shared negative electrons in between the nuclei, this forms strong intramolecular forces which hold atoms together

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12
Q

What prevents the positive nuclei repelling between covalently bonded atoms?

A

High electron density between nuclei balances the repulsion of the nuclei

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13
Q

How is the degree of strength a covalent bond has measured?

A

As average bond enthalpy, the higher this value the stronger the covalent bond

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14
Q

What is a daitive bond also known as?

A

A co-ordinate bond

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15
Q

What is a daitive bond?

A

A covalent bond in which both electrons come from the same atom.

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16
Q

How do you show the daitive bond in a displayed formula?

A

Draw an arrow rather than just a line to represent the bond. The arrow should point to the atom (previously an ion) which didn’t bring any electrons to the bond.

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17
Q

How can you tell the difference between a regular covalent bond and a daitive covalent bond?

A

It is impossible to tell the difference once the bond is formed, there is no difference in reality.

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18
Q

Why do ions remain together after exchanging electrons?

A

The ions are held together by strong electrostatic attractions between the negative charges.

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19
Q

What are some properties of giant ionic lattices?

A
  • High melting point and boiling point
  • Brittle
  • Conductivity
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20
Q

Is the cathode positively or negatively charged?

A

negatively (attracts positive cations)

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21
Q

Is the anode positively or negatively charged?

A

positively (attracts negative anions)

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22
Q

What is formed when ammonia and hydrogen chloride are reacted? What do the reactants look like and what does the product look like?

A

Ammonium chloride,

reactants = colourless gases, product = a thick white smoke

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23
Q

what is the definition for a covalent bond?

A

when two non-metals share a pair of electrons to form structures isoelectronic to noble gases.

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24
Q

Is hydrochloric acid ionically or covalently bonded?

A

Covalently

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25
Q

How do you show a daitive bond in a displayed formula?

A

Draw an arrow pointing to the atom which didn’t bring anything to the molecule

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26
Q

How are ammonium ions formed?

A

by the transfer of a hydrogen ion from the hydrogen chloride (when ammonia and HCl are reacted) to the lone pair of electrons on the ammonia molecule

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27
Q

What do acids form with water?

A

Hydroxonium ions (H₃O⁺)

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28
Q

How is boron trifluoride different from other molecules?

A

it is electron deficient, the boron only has 3 pairs of electrons in its bonding level whereas there would be room for 4 pairs.

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29
Q

What is the valence shell electron pair repulsion theory (VSEPR)?

A

that bonded electron pairs repel bonded electron pairs, but lone pairs repel more. This is to reduce electrostatic attraction.

“Molecules take up the shape that minimises the repulsion between electron pairs within the molecule.”

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30
Q

Is Beryllium chloride electron deficient?

A

yes

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31
Q

How would you draw a double bond in a 3D displayed formula?

A

A double bond only counts as one bond/one area of electron density so you draw it as a singular bond

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32
Q

What are lone pairs also known as?

A

non-bonded pairs

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33
Q

why do lone pairs have a more concentrated negative charge?

A

because they are not shared between two atoms

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34
Q

What is the name of the structure of a molecule with 2 bonded electron pairs?

A

linear

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35
Q

What is the bond angle in a linear molecule?

A

180°

36
Q

What is the name of the structure of a molecule with 2 bonded electron pairs and 2 lone pairs?

A

non-linear

37
Q

What is the bond angle in a non-linear molecule?

A

104.5°

38
Q

What is the name of the structure of a molecule with 3 bonded electron pairs?

A

trigonal planar

39
Q

What is the bond angle in a trigonal planar molecule?

A

120°

40
Q

What is the name of the structure of a molecule with 3 bonded electron pairs and 1 lone pair?

A

trigonal pyramidal

41
Q

What is the bond angle in a trigonal pyramidal molecule?

A

107°

42
Q

What is the name of the structure of a molecule with 4 bonded electron pairs?

A

tetrahedral

43
Q

What is the bond angle in a tetrahedral molecule?

A

109.5°

44
Q

What is the name of the structure of a molecule with 5 bonded electron pairs?

A

trigonal bipyramidal

45
Q

what is the bond angle in a trigonal bipyramidal molecule?

A

120° between the base and 90° between the vertical

46
Q

what is the name of the structure of a molecule with 6 bonded electron pairs?

A

octahedral

47
Q

what is the bond angle in an octahedral molecule?

A

90°

48
Q

What is electronegativity?

A

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

49
Q

Within a group, why is there a decrease in electronegativity as you go down the group?

A

Because of the increasing size of the atoms (increased shielding and increasing atomic radius) as you go down the group

50
Q

Across a period, why is there an increase in electronegativity?

A

the increasing nuclear charge decreases the atomic radius

51
Q

What is a “pure” covalent bond?

A

when the electrons are shared evenly between the two atoms

52
Q

What is a polar bond?

A

A covalent bond in which there is a separation of charge between one end and the other (when one end is slightly positive and the other end slightly negative.)

53
Q

What is a permanent dipole?

A

when the opposite charges within a molecule act as poles on either end of the molecule (2 poles = permanent dipole)

54
Q

If one atom is a lot more electronegative than the other, what happens?

A

The more electronegative atom takes the electrons completely and ions are formed

55
Q

What makes the atoms ionic/completely charged rather than slightly charged?

A

a difference of more than 0.8 in electronegativity

56
Q

when can ionic bonds showing covalency happen?

A
  • If the positive cation is small and highly charged
  • if the negative anion is large (so highly polarisable)
  • dependent on the charge density
57
Q

How can you calculate the charge density?

A

charge of the ion ÷ ionic radius

58
Q

What does it mean by ionic bonds showing covalency?

A

If an atom has a high charge density (e.g. aluminium compared to chlorine) then the electron pairs are attracted back to the atom to such an extent that the electrons are partly shared and so the bonds become covalent

59
Q

when does polarising ability increase?

A

as the positive ion decreases and the number of charges increases

60
Q

What electronegativity difference makes a bond pure covalent?

A

a difference of 0

61
Q

What electronegativity difference makes a bond polar covalent?

A

a difference of 0 > 1.8

62
Q

What electronegativity difference makes a bond ionic?

A

a difference of >1.8

63
Q

What are the types of intermolecular forces?

A

permanent dipole, london forces, hydrogen bonding

64
Q

What are intermolecular forces?

A

attractions between one molecule and a neighbouring molecule

65
Q

What is dipole-dipole attraction/bonding?

A

when molecules line up so the partial positive charges are adjacent to partial negative charges - a net attractive force/weak electrostatic attraction

66
Q

What are the conditions for hydrogen bonding to occur?

A
  • the hydrogen atom must be attached to an atom with high electronegativity (O, N or F)
  • there must be a lone pair of electrons on the electronegative atom
67
Q

how would the properties of water be different without hydrogen bonds?

A
  • decreased surface tension
  • Lower boiling point/melting point
  • ice would be denser than liquid water
  • water wouldn’t stick to itself
68
Q

what are london forces?

A

the weakest of the intermolecular forces

69
Q

what are london forces also known as?

A

temporary dipoles/induced dipoles

70
Q

What are monatomic molecules?

A

molecules of noble gases which consist of a single atom

71
Q

What happens to the strength of london forces as the molecule increases in size?

A

the number of electrons also increases making the molecule more polarisable and producing greater induced dipole attractions

72
Q

Between molecules with long chains of atoms, do london forces get stronger or weaker? why?

A

stronger because there are more points of contact.

73
Q

what is the bond strength of a covalent bond?

A

200-500 kj/mol to break

74
Q

what is the bond strength of a hydrogen bond?

A

5-40kj/mol to break

75
Q

what is the bond strength of london forces?

A

2kj/mol to break

76
Q

what happens to iodine crystals with GENTLE heating?

A

sublimation

77
Q

What are physical properties governed by?

A

the intermolecular forces

78
Q

why do molecular substances tend to be gases, liquids or have low mpt/bpts?

A

because the intermolecular forces of attraction are comparatively weak - no covalent bonds need to be broken to melt/boil a molecular substance

79
Q

If a molecule is soluble in water without reacting, what does this suggest?

A

that the molecule is small and capable of hydrogen bonding or at least is a small very polar molecule

80
Q

if a molecule is a high melting point solid, what does this suggest?

A

that it is a giant structure (ionic, metallic or giant covalent)

81
Q

if a high melting point solid is soluble in water, what is suggested about the bonding of the solid?

A

it is ionic

82
Q

if a substance conducts electricity when solid what is suggested about the type of substance it is?

A

it’s either a metal or graphite because delocalised electrons must be present to conduct electricity

83
Q

what is the strength of the intermolecular forces related to?

A

the number of electrons present and so the size of the molecule (or isolated atom)

84
Q

Which substance will have the higher boiling point out of one polar and one non-polar molecule with the same masses (and so the same number of electrons)?

A

the polar molecule will have the higher boiling point

85
Q

As you go down a group, how do the strength of the London forces and permanent dipoles compare?

A

Strength of London forces increases but the strength of permanent dipole decreases

86
Q

When drawing in 3D, how do you express the arrangement of the bonds using different types of lines?

A
  • An ordinary line represents a bond in the plane of the paper (flat)
  • a dotted line represents a bond going away from you, into the paper
  • a wedge represents a bond coming towards you, out of the paper
87
Q

What is a key point to remember when drawing ammonium ions in 3D?

A

The lone pair in ammonium ion is part of a Daitive covalent bond so electrons are still shared, this only counts as one area of electron density

A Level Chemistry (21 decks)

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