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Chemistry > Topic 2 - Bonding & Structure > Flashcards

Topic 2 - Bonding & Structure Flashcards

1
Q

What is ionic bonding?

A

The strong electrostatic attraction between oppositely charged ions.

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2
Q

What effects does the ionic radius have of the strength of ionic bonding?

A

As ionic radius increases, charge density decreases ∴ weaker bonds are made. (enthalpy = less exo)

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3
Q

What effects does the ionic charge have on the strength of ionic bonding?

A

The bigger the charge difference the stronger the bond will be.

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4
Q

What are the trends in ionic radii down a group for a set of isoelectronic ions?

A

(+ve) ions are smaller than their atoms as they have less electrons and a greater ration of protons to remaining electrons. (-ve) ions are larger as there are more elctrons fro the same amount of protons, so the pull is less.

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5
Q

How does the migration of ions explain the existence of ions?

A

The (-ve) ions move towards the (+ve) electrode and the (+ve) ions move towards the (-ve) electrode.

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6
Q

How do the physical properties of ionic compounds explain the existence of ions?

A
  • high mp (strong attractive forces between ions)
  • non conductor as solid (ions held together tightly)
  • conductor when molten/in aq (ions free to move)
  • brittle
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7
Q

What is a covalent bond?

A

The strong electrostatic attraction between 2 nuclei & the shared pair of electrons between them.

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8
Q

What are diative covalent bonds?

A

When the shared pair of electrons in the covalent bond come from only 1 bonding ato. (co-ordinate bonding, shown by an arrow from sharing atom)

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9
Q

What is the relationship between bond lengths & strengths in covalent bonding?

A

If their is a greater force of attraction (x2/x3 bonds) the bond will be shorter but stronger.

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10
Q

How do lone pairs and bonding pairs arrange themselves around a central atom & determine the molecules shape?

A

Lone pairs repel more than bonding pairs, which reduces bond angles.

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11
Q

What is a linear molecule?

A
  • 2 bonding pairs, no lone pairs
  • 180° bond angle
  • CO₂, BeF₂
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12
Q

What is a trigonal planar molecule?

A
  • 3 bonding pairs, no lone pairs
  • 120° bond angle
  • SO₃, NO₃⁻, CO₃²⁻
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13
Q

What is a tetrahedral molecule?

A
  • 4 bond pairs, no lone pairs
  • 109.5° bond angle
  • SiCl₄, SO₄²⁻, NH₄⁺
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14
Q

What is a trigonal pyramidal molecule?

A
  • 3 bond pairs, 1 lone pair
  • 107° bond angle
  • PF₃,H₃O⁺
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15
Q

What is a bent/v-shaped molecule?

A
  • 2 bond pairs, 2 lone pairs
  • 104.5° bond angle
  • H₂S, SCl₂
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16
Q

What is a trigonal bipyramidal molecule?

A
  • 5 bond pairs, no lone pairs
  • 120° & 90° bond angles
  • PCl₅
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17
Q

What is an octahedral molecule?

A
  • 6 bond pairs, no lone pairs
  • 90° bond angle
  • SF⁶
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18
Q

What is a square planar molecule?

A
  • 4 bond pairs, 2 lone pairs (variation or octahedral)
  • 90° bond angle w/ lone above & below
  • XeF₄
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19
Q

What is the electron-pair repulsion theory?

A

VSEPR, used to predict the shapes of molecules as they are arranged to produce the minimum amount of repulsion/maximum separation.

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20
Q

What is electronegativity?

A

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

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21
Q

What are the factors affecting electronegativity?

A
  • ↑ across a period as protons ↑ ∴ radius ↓

- ↓ down a period as distance & shielding ↑

22
Q

How does electronegativity affect bonding?

A
  • ↓ elec.neg. difference = covalent

- ↑ elec.eg. difference = ionic (>1.7)

23
Q

What is a polar covalent bond?

A

Not purely ionic or covalent. The least elec.neg. element = δ⁺ & the most elec.neg. = δ⁻

24
Q

How do you know if a bond is polar or non-polar?

A
  • if symmetrical (all bonds same & no lone pairs) = non-polar even if individual bonds are (cancel out)
  • CO₂ = non, CH₃Cl = polar
25
Q

What are London forces?

A

Instantaneous, induced dipole-dipole interactions, occur when fluctuations in e⁻ density causes temporary dipoles to form, which then = induced dipoles.

26
Q

What affects the strength of London forces?

A

↑ electrons = ↑ chance of temporary dipoles = stronger L forces which need ↑ energy to overcome. Seen in ↑ bp down group 7. (I₂ = solid & Cl₂ = g)

27
Q

How does the chain length of alkanes affect the size of L-forces?

A

↑ chain length = ↑ surface area of contact ∴ stronger L-forces. Branches in c-chain = ↓ points of contacts ∴ ↓ bps.

28
Q

What are permanent dipole-dipole forces?

A
  • occur between polar molecules, stronger than L-forces
  • bond has a significant difference in elec.neg.
  • can occur in addition to L forces. (δ⁻ Cl in H-Cl attracted to δ⁺ H in another H-Cl)
29
Q

What are hydrogen bonds?

A

Bonds that occur in compounds that have a H atom attached to F, N or O, and an available lone pair.
-occurs in addition to L-forces

30
Q

What is the reason for the shape of H-bonds?

A

They form a 180° angle around the H atom as the H gives 2 electrons which repel to min repulsion/max separation.

31
Q

How are H-bonds formed in H₂O?

A

Water forms 2 H-bonds per molecule as δ⁻ oxygen has 2 lone pairs ∴ = stronger H-bonds & ↑ pb.

32
Q

Why does ice have a lower density than water?

A

The H-bonds create interlocking rings of 6 water molecule when water freezes, creating distance between the molecules which decreases its density.

33
Q

Why do H₂O, NH₃ & HF have abnormally high bps?

A

They form H-bonds as well as L-forces which needs ↑ energy to break.

34
Q

Why do alcohols have lower volatility compared to alkanes?

A

They can form H-bonds ∴ have ↑ bps so need ↑ energy to overcome.

35
Q

What are the trends in bp of the hydrogen halides from HF to HI?

A
  • HF = ↑due to H-bonds

- ↑ from HCl to HI as L-forces ↑ due to ↑ electrons

36
Q

What does ‘like dissolves like’ mean?

A

Compounds with similar intermolecular forces to those in the solvent will generally dissolve. (Iodine = only L-forces dissolves hexane = only L-forces.)

37
Q

Why is propane a useful solvent?

A

It has polar & non-polar characteristics, so can dissolve polar and non-polar substances.

38
Q

How is water used to dissolve some ionic compounds?

A

Water can hydrate the ions, the δ⁺ H bonds to the (-ve) ion in the bond & the δ⁻ O bonds to the (+ve) ion.

39
Q

How is hydration affected by charge density?

A

↑ charge density = ↑ hydration enthalpy, as ions attract water molecules more strongly.

40
Q

How does water dissolve simple alcohols?

A

↓ alcohols = soluble in water as can form H-bonds, as hydrocarbon chain ↑ the less soluble it becomes.

41
Q

When is water not useful as a solvent?

A

If substances cant form H-bonds (halogenoalkanes) or are non-polar (hexane) they are insoluble in water.

42
Q

What is metallic bonding?

A

The strong electrostatic attraction between metal ions & delocalised electrons.

43
Q

What are the main factors that affect the strength metallic bonding?

A
  • no. of protons (↑ protons = stronger bond)
  • no. of delocalised electrons per atom (↑ = stronger)
  • size of ion (↓ ion = stronger bond)
44
Q

What are giant lattices present in?

A
  • ionic solids (giant ionic lattices)
  • covalently bonded solids (giant covalent lattices)
  • solid metals (giant metallic lattices)
45
Q

What is the structure of covalently bonded substances like I₂ & H₂0?

A

simple molecular

46
Q

What are the allotropes of carbon?

A
  • diamond, tetrahedral, 4 bonds per atom
  • graphite, planar, 3 bonds per atom, layer of delocalised electrons
  • graphene, 1 layer of graphite, ↑ tensile strength & good conductor of electricity
47
Q

What are the physical properties of ionic structures?

A
  • ↑ m/bp due to giant lattice w/ strong electrostatic forces between oppositely charged ions
  • soluble in water
  • poor conductor as solid, good when molten (ions free)
  • crystalline solids
48
Q

What are the physical properties of simple molecular structures?

A
  • ↓ m/bp due to weak intermolecular forces
  • generally poor solubility in water
  • poor conductor as solid & molten (no ions)
  • mostly gases & liquids
49
Q

What are the physical properties of macromolecular structures?

A
  • ↑ m/bp due to many strong covalent bonds
  • insoluble in water
  • diamond poor conductor, graphite good (delocalised)
  • solids
50
Q

What are the physical properties of metallic structures?

A
  • ↑ m/bp due to strong electrostatic forces
  • insoluble in water
  • good conductor as solid & molten
  • shiny & malleable

Chemistry (19 decks)

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