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AQA A-Level 3.1 Physical Chemistry > 3.1.2 BONDING > Flashcards

3.1.2 BONDING Flashcards

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1
Q

What is ionic bonding?

A

The electrostatic attraction between oppositely charged ions in a lattice. Takes place between a non-metal (-) and a metal (+). Electrons are transferred to achieve a full outer shell.

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2
Q

What are the common compound ions and their formulas?

A
  • Sulphate = s04^2-
  • Hydroxide = OH-
  • Nitrate = NO3^-
  • Carbonate = CO3^2-
  • Ammonium = NH4+
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3
Q

What are covalent bonds?

A

Bonding between non-metals.

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4
Q

What does a single covalent bond contain?

A

Shared pairs of electrons.

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5
Q

What do multiple covalent bonds contain?

A

multiple shared pairs of electrons.

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6
Q

What is co-ordinate (dative covalent) bonding?

A

Shared pairs of electrons where both electrons are supplied by one atom.

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7
Q

How do we represent covalent bonds when drawing?

A

Using a line.

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8
Q

How do we represent a co-ordinate bond?

A

Using an arrow.

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9
Q

What is metallic bonding?

A

The attraction between a sea of delocalised electrons and positive ions is arranged in a lattice.

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10
Q

What are the four types of crystal structure?

A

Ionic, Covalent, macromolecular, molecular.

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11
Q

What is the structure and properties of an ionic crystal?

A
  • Giant ionic lattice with a regular arrangement of - and + ions.
  • strong electrostatic forces of attraction so high melting point so more energy required to break these.
  • cannot conduct electricity when solid, but can when molten as ions carry charge freely.
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12
Q

Ionic bonding is stronger and melting points are higher in ions that are…

A

smaller and have higher charges.

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13
Q

What is the structure and properties of metallic bonding?

A
  • giant metallic lattice of positive ions and sea of delocalised electrons.
  • high melting point due to strong attractions between the positive metal ions and electrons and strong metallic bonds so requires a lot of energy to overcome this
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14
Q

Why are metals good thermal and electrical conductors?

A
  • Delocalised electrons can pass kinetic energy to each other
  • Delocalised electrons can move freely and carry current
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15
Q

What are 3 factors affecting the strength of metallic bonding?

A
  1. number of protons
  2. number of electrons
  3. size of the ion
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16
Q

Why does Mg ion have a stronger metallic bonding and higher melting point than Na ion?

A
  • There are more electrons in the outer shell of the Mg ion.
  • It has more protons than Na ion.
  • It is a smaller ion than a Na ion.
  • so there is stronger attraction and higher energy required to break bonds.
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17
Q

What are some examples of macromolecular structures?

A

Diamond and Graphite

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18
Q

Describe the structure of diamond:

A
  • each carbon atom is covalently bonded to 4 other carbon atoms
  • tetrahedral shape
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19
Q

What are the 5 properties of diamond?

A
  • very hard
  • high melting points
  • does not conduct electricity
  • does not dissolve in solvents.
  • good thermal conductors as vibrations travel easily
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20
Q

Describe the structure of graphite:

A
  • Carbon atoms are covalently bonded to 3 other carbon atoms
  • 4th outer electron of each C atom is delocalised.
  • sheets of hexagons (graphene) is bonded by weak van der Waal forces.
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21
Q

What are the 4 properties of graphite?

A
  • slippery and soft
  • electrical conductor
  • insoluble in any solvent
  • high melting points.
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22
Q

Why is graphite slippery and soft?

A

There are weak van der Waal forces between the layers of graphite which can be easily broken so the sheets can slide past each other.

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23
Q

Why do diamond and graphite have high melting points?

A

They have many strong covalent bonds which require a lot of energy to break.

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24
Q

What are 2 examples of molecular crystal structures?

A

Iodine and ice.

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25
Q

What is the structure and properties of molecular crystals?

A

Molecules are arranged in regular arrays by intermolecular forces.
There are strong van der Waal forces because there are many electrons, but they are weaker than covalent bonds with a low melting point.
- does not conduct electricity
- iodine sublime easily to its gaseous form.

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26
Q

What is the order of repulsion for electron pairs?

A

lone pair - lone pair> lone pair to bond pair> bond pair-bond pair.

27
Q

What is the angle degree rule for lone pairs?

A

1 lone pair = 2.5°

2 lone pairs = 5°

28
Q

What is the angle in a linear shape?

A

180°

29
Q

What is the angle in a trigonal planar shape?

A

120°

30
Q

What is the angle in a bent shape with 2 BP and 1 LP?

A

roughly 118°

31
Q

What is the angle in a tetrahedral shape?

A

109.5°

32
Q

What is the angle in a trigonal pyramidal shape with 3 BP and 1 LP?

A

107°

33
Q

What is the angle in a bent shape with 2 BP and 2 LP

A

104.5°

34
Q

What are the 2 angles in a trigonal bipyramidal shape?

A

120° and 90°

35
Q

What are the 2 angles in trigonal pyramidal with 4 BP and 1 LP?

A

119° and 89°

36
Q

What is the angle in a t-shape?

A

120°

37
Q

What is the angle in an octahedral?

A

90°

38
Q

What is the angle in a square pyramidal?

A

89°

39
Q

What is the angle in a square planar?

A

90°

40
Q

What is electronegativity?

A

the power of an atom to attract the pair of electrons/electron density to itself in a covalent bond.

41
Q

What are polar molecules?

A

when there is unequal sharing of electrons in a covalent bond.

42
Q

Which element is the most electronegative?

A

Fluorine with 4.0

43
Q

what does electronegativity depend on?

A
  • nuclear charge of an atom
  • the distance between the nucleus and the outer electrons
  • the shielding of the nuclear charge by electrons in the inner shells.
44
Q

How is the electronegativity of atoms measured?

A

0-4 on the Pauling scale.

45
Q

What are non-polar molecules?

A

when the electron sharing is equal in a bond,

e.g: F₂

46
Q

How do we represent partial charges?

A

δ⁺ and δ⁻

47
Q

Molecules with polar covalent bonds will have what force?

A

permanent dipole force.

48
Q

Describe the trend of electronegativity as you go across a period and why:

A
  • EN increases
  • nuclear charge increases/proton number increases
  • number of shells remain the same
  • so atom size gets smaller as there is strong attraction between nucleus and electrons in shells.
49
Q

Describe the trend of electronegativity as you go down a group:

A
  • EN decreases
  • distance between the nucleus and outer electrons increases
  • shielding of inner shells increases.
50
Q

What are the 3 types of intermolecular forces?

A
  • Van der Waals
  • Permanent dipole-dipole forces
  • Hydrogen bonding
51
Q

What is the difference between INTER and INTRA molecular forces?

A 
INTER = are between molecules in a substance and weaker.
INTRA = are between the atoms in a molecule and stronger.
52
Q

What are Van der Waal forces and how are they formed?

A
  • Occur in all molecules.
  • temporary dipoles created by constant and random movement of electrons in the cloud.
  • as molecule size increases, electrons increase, stronger VdW forces as more energy needed to overcome it.
53
Q

What are permanent dipole-dipole forces?

A
  • Occur between polar bonds due to the larger difference between electronegativity
  • are stronger than VdW forces
54
Q

What is needed for hydrogen bonding?

A
  • a highly electronegative atom with lone pairs

- attracts with H atom.

55
Q

What are the three molecules that have hydrogen bonding? (from highest boiling to lowest)

A

H2O, HF, NH3

56
Q

Why do NH3, H20 and HF not fit the trend in the hydrides graph of boiling points

A

They have higher boiling points than expected for their molecular mass because it has an extra intermolecular force called hydrogen bonding.

57
Q

Why is hydrogen bonding important?

A
  • Ice is less dense than water because the particles are further apart/less closely packed.
  • they enable organisms to thrive in ponds as it would reflect and absorb sunlight to keep the bottom of the pond warm.
58
Q

Identify one molecule with the same number of atoms, the same number of electrons and the same shape as H3O^+ (1)

A

NH3.

59
Q

Deduce the molecule of a compound that has the same number of atoms and electrons and the same shape as AlCl4^- (1)

A

SiCl4.

60
Q

What does Electron Pair Repulsion Theory state?

A

Electron pairs repel each other equally if there are all bond pairs to get as far away as possible to minimise repulsive forces between them.

61
Q

What’s the difference between saturated and unsaturated fats in terms of intermolecular forces?

A
  • Saturated fats have regular chains/branches so stronger IMFs so higher M.P.
  • Unsaturated fats have random branched chains so weaker IMFs so lower M.P.
62
Q

How does hydrogen bonding arise (3)

A
  • There is a large difference in electronegativity
  • so forms delta + and delta - molecule
  • where lone pair will be attracted by/donated to
63
Q

Explain why the O-H bond in methanol is polar (2)

A
  • there is a large difference in EN between the O and H atom

- So O will attract higher e- density around it to form δ⁺and δ⁻ molecule.

64
Q

What are the steps in working out the shape?

A
  1. number of outer electrons in central atom
  2. add one electron from each bonding atom
  3. Add or subtract if the molecular is charged
  4. Total electrons
  5. Divide by two to work out total electron pairs
  6. How many bonding pairs are there? How many lone pairs?

AQA A-Level 3.1 Physical Chemistry (2 decks)

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