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Bonding Flashcards

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1
Q

Ionic bonding

A

The strong electrostatic attraction between oppositely charged ions.

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2
Q

How does ionic radius and charge affect bond strength?

A

Increased charge density and ions can get closer together - increase the strength of the ionic bond.

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3
Q

Explain reason in trend in ionic radii down a group

A

Despite nuclear charge increasing, the number of shells increases (include example here in exam) so the inner shells effectively shield the outer shells from the nuclear charge (effective nuclear charge decreases), weaker forces of attraction between nucleus and outer shell, ionic radii increases.

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4
Q

Explain reason in trend in ionic radii across a period (isoelectronic ions)

A

Isoelectronic structure (state the structure in exam), nuclear charge increasing but number of shells stays constant, similar shielding, effective nuclear charge increases, forces of attraction increases, reducing ionic radius.

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5
Q

Transition metals becoming ions

A

Remove 4s electrons before 3d.

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6
Q

Proof of existence of ions

A

Electrolysis of copper(II) chromate(VI). Blue Cu2+ ions migrate to cathode and yellow chromate ions migrate to anode.
The physical properties of high boiling point, solubility in polar solvents and conductivity of solutions.

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7
Q

What does the ionic model assume and for which ions is this true for?

A

Ions are hard, spherical, non-compressible, non-polarisable.

Li+ and F-

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8
Q

What makes cation better at polarising?

A

Smaller and higher charge, increase charge density

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9
Q

What makes an anion better at being polarised?

A

Larger and higher charge.

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10
Q

How does polarisation affect the bond strength?

A

The covalent character makes the bond more stable than a pure ionic bond.

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11
Q

What is lattice energy?

A

Energy change when gaseous ions combine to form one mole of a solid ionic lattice.

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12
Q

Why does magnitude of lattice energy increase from LiF to MgCl2?

A

Despite increasing ionic radius, significantly more stable because Mg2+ is much more charge dense than Li+ due to the charge doubling.

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13
Q

Why is the experimental lattice energy in LiF same as theoretical?

A

Ions obey ionic model. Fl- is too difficult to polarise and Li+ is not effective enough at polarising.

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14
Q

Why is the magnitude of experimental lattice energy in MgCl2 higher than theoretical?

A

It does not obey ionic model. Polarisation of chloride distorts bond because Mg2+ has high charge density. This strengthens bond and makes lattice more stable.

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15
Q

Definition of covalent bonding

A

Strong electrostatic attraction between two nuclei and the shared pair of electrons between them.

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16
Q

Dative covalent bond

A

When two atoms form a covalent bond with both paired electrons coming from one of the atoms - it donates a lone pair

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17
Q

How is Al2Cl6 formed?

A

Dimerisation of AlCl3 with 2 dative covalent bonds (check OneNote).

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18
Q

Relationship between bond length and bond strength and why

A

Shorter = stronger; longer = weaker because shared pair is less electron dense as it is spread out over large volume, therefore electrostatic attraction between it and atomic nuclei is weaker and therefore less energy to break.

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19
Q

What is electronegativity?

A

The ability to attract the bonding electrons in covalent bond.

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20
Q

What factors increase electronegativity?

A

Higher nuclear charge and smaller atomic radius (more important factor).

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21
Q

Which element has the highest electronegativity?

A

Fluorine

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22
Q

Is CO2 a polar molecule?

A

CO2 is linear so the polar bonds are symmetrical therefore the direction and magnitude of bonds cancel out and the molecule loses polarity.

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23
Q

IS CCl4 polar molecule?

A

Tetrahedral shape therefore symmetrical and charge distribution cancels out (calculated by trig).

24
Q

How to draw VSEPR

A

1) Draw dot-cross Lewis diagram
2) Identify central atom.
3) Count number of areas of electron density.
4) Determine electron pair distribution shape.
5) Determine shape ignoring lone pairs.
6) Apply adjustments to shape/bond angles due to lone pairs.

25
Q

Order of magnitude of repulsion between different electron densities (high to low)

A

lp - lp, lp - bp, bp - bp (multiple bonds, single bonds).

26
Q

2 bonding pairs, 0 lone pairs

A

Linear, 180

27
Q

2 bonding pairs, 1 or 2 lone pair

A

V-shaped, 119.5 or 104.5

28
Q

3 bonding pairs, 2 lone pairs

A

T-shaped, 90

29
Q

3 bonding pairs, 0 lone pairs

A

Trigonal planar, 0

30
Q

3 bonding pairs, 1 lone pair

A

Trigonal pyramidal, 107

31
Q

4 bonding pairs, 0 lone pairs

A

Tetrahedral, 109.5 or square planar, 90

32
Q

5 bonding pairs, 0 lone pairs

A

Trigonal bipyramidal, 90 and 120

33
Q

6 bonding pairs, 0 lone pairs

A

Octahedral, 90

34
Q

Intermolecular forces

A

Weak forces of attraction between different molecules

35
Q

London forces

A

Instantaneous dipole - induced dipole attractions.

All molecules have London forces.

36
Q

Instantaneous dipole

A

Temporary fluctuation in electron density of a molecule due to kinetic effects that causes a temporary dipole.

37
Q

Induced dipole

A

Temporary dipole caused by neighbouring molecules with instantaneous dipole. This causes weak attraction between the two molecules.

38
Q

What causes stronger London Forces?

A

More electrons

39
Q

Permanent dipole

A

Most common in halogens and nitrogen and oxygen.
Molecule with a dipole can attract another molecule with a dipole.
Weaker than London forces.
London + perm dipole forces add up to be stronger than just London alone.

40
Q

Hydrogen bonds

A

Strongest intermolecular force.
Has to be hydrogen directly bonded with N, O or F.
They have extremely high electronegativity compared to H.
In H-X bonds, antibonding empty orbitals become exposed so it is very delta positive. This means electrons from other molecule can be donated to it and bond.
Has to be 180 degrees.

41
Q

Link between H-bonds and viscosity

A

More H-bonds per molecule mean that interconnected chains and networks of molecules can form, increasing viscosity.

42
Q

Why is ice less dense than water?

A

The hydrogen bonds between water molecules in ice are at 180 from each other, creating a structure with many gaps. In liquid water, molecules roll over each other with hydrogen bonds being formed and made, meaning it does not conform to a structure and is denser.

43
Q

Why do molecules with H bonds have higher boiling point?

A

The electronegativity between H and N / O / F is great enough create H bonds. The other molecules do not have great enough of a difference of electronegativity to form H-bonds therefore they only form London forces or permanent dipoles. These require much less energy to break.

44
Q

Why is water a poor solvent for compounds without H-bonds?

A

Water cannot form H-bonds even with polar molecules (such as halogenoalkanes) as it is not energetically favourable to break the H-bonds in water and the London forces and dipole-dipole forces in halogenoalkanes.

45
Q

Why are non-polar substances soluble in non-polar solvents?

A

London forces form between both as the forces are energetically comparable.

46
Q

Why are polar substances insoluble in non-polar substances?

A

There is no dipole be attracted to the polar substance in the solvent and the London forces do not release enough energy in order to form between 2 substances.

47
Q

When are ionic substances soluble in water

A

If the energy to break the ionic bonds (lattice energy) is less than the energy released by the hydration of the ions (sum of hydration enthalpies). These are dipole-ion attractions.

48
Q

Why can water and simple alcohol mix?

A

They have similar strengths of IMF. Water has 2 H bonds and short chain alcohols have 1 but also London forces in addition. Energy released by water-alcohol bonds are similar to energy needed to break IMFs of water and alcohol.

49
Q

Trend in boiling temps of hydrogen halides and why

A

HF is high due to H bonding needing much energy to break. HCl to HI have dipole-dipole and London. Despite dipole-dipole forces decreasing, the London forces significantly increase therefore b.p increases.

50
Q

Effect of branching of carbon chains on b.p of alkanes

A

More branching means b.p decreases as molecules are not able to stack as effectively and London forces act of longer distances and the strength of the attractions decrease.

51
Q

Effect of length of carbon chains on b.p of alkanes.

A

Longer = more electrons = increased London forces

52
Q

Why are alcohols less volatile than alkanes?

A

They have H bonds as opposed to only London bonds so the intermolecular forces for the alcohols are stronger.

53
Q

Bond angles in ethane

A

Co-joined tetrahedral 109.5

54
Q

Bond angles in ethene

A

Co-joined trigonal planar

H-C-H 117

55
Q

Bond angles in ethyne

A

Linear

56
Q

Methanol VSEPR

A

Tetrahedral with angular C-O-H

Chem (19 decks)

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