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AQA - A level Chemistry > Bonding > Flashcards

Bonding Flashcards

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1
Q

The melting point of XeF4 is higher than the melting point of PF3 Explain why the melting points of these two compounds are different.
In your answer you should give the shape of each molecule, explain why each molecule has that shape and how the shape influences the forces that affect the
melting point.

A
  • XeF4 4BP and 2LP around Xe
  • PF3 3BP and 1LP around P
  • XeF4 square planar
  • PF3 pyramidal
  • XeF4 has vdw forces and PF3 has dipole- dipole forces (and vdw)
  • Stronger/more intermolecular forces in XeF4
  • Due to larger Mr
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2
Q

Ionic bonding structure and bonding

A

strucutre = giant lattice
bonding = strong ionic

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3
Q

ionic bonding definition

A
  • strong electrostatic forces of attraction
  • between oppositely charged ions
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4
Q

physical properties of ionic compounds

A
  • high MP/BP
  • electrical conductivity when aqueous
  • brittle
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5
Q

covalent bonding structure and bonding

A

structure = simple covalent/macromolecular

bonding = strong covalent

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6
Q

covalent bonding definition

A
  • shared pair of electrons
  • between two atoms
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7
Q

properties of diamond

A
  • tetrahedral
  • each C atom forms 4 covalent bonds
  • high MP
  • hard
  • non-conductor of electricity (no delocalised electrons)
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8
Q

properties of graphite

A
  • each carbon forms 3 covalent bonds
  • each carbon has one free delocalised electron
  • layers held by weak VDW
  • soft (layers can slide over each other)
  • high MP (strong covalent bonds)
  • conducts electricity due to delocalised electrons
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9
Q

define co-ordinate bonding

A
  • shared electron pair
  • from same atom
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10
Q

metallic bonding structure and bonding

A

structure = giant lattice
bonding = strong metallic bonds

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11
Q

why does Mg have stronger bonding than Na

A
  • Mg2+ ions smaller than Na+ ions
  • Mg2+ greater charge
  • more delocalised electrons
  • attraction between Mg2+ and delocalised electrons stronger
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12
Q

properties of metals

A
  • electrical and thermal conductors (delocalised electrons)
  • strong (strong forces of attraction)
  • malleable and ductile
  • high MP/BP
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13
Q

calcium carbonate formula

A

CaCO3

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14
Q

Magnesium nitrate formula

A

Mg(NO3)2

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15
Q

Zinc phosphate formula

A

Zn3(PO4)2

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16
Q

Sodium sulphide formula

A

Na2S

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17
Q

Copper (I) hydroxide formula

A

CuOH

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18
Q

Aluminium dichromate formula

A

Al2(Cr2O7)3

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19
Q

Lead (II) sulphate formula

A

PbSO4

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20
Q

Silver (I) iodide formula

A

AgI

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21
Q

Ammonium carbonate formula

A

(NH4)2 CO3

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22
Q

Potassium oxide formula

A

K2O

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23
Q

state structure for iodine and graphite

A

iodine = simple molecular
graphite = giant covalent/macromolecular

24
Q

explain why graphite has high MP

A

lots of strong covalent bonds need to be broken

25
Q

explain why iodine has low MP

A

weak VDW forces are broken

26
Q

state why iodine is a poor electrical conductor

A

no delocalised electrons

27
Q

Zinc(II) bromide (ZnBr2) is a crystalline solid with a melting point of 394°C.

Suggest the type of bonding present in Zinc(II) bromide and state why the melting point is high.

A
  • ionic bonding
  • strong electrostatic forces of attraction between oppositely charged ions
  • large amount of energy needed to break bonds
28
Q

explain how ions held together in sodium metal

A
  • metallic bonding
  • explain metallic bonding
29
Q

explain how ions are held together in NaCl

A
  • ionic bonding
  • explain ionic bonding
30
Q

In terms of structure and bonding explain why Magnesium has a higher melting point than ethane (C2H6).

A

magnesium = metallic bonding
giant metallic lattice
- explain metallic bonding

C2H6 = simple molecular
- weak VDW forces

  • metallic bonds stronger than VDW
  • more energy needed to break metallic bonds
31
Q

The melting point of sodium chloride is much higher than that of sodium metal. What can be deduced from this information?

A

ionic bond in NaCl is stronger than metallic bond in Na

32
Q

Compare the electrical conductivity of solid sodium metal with that of solid sodium chloride. Explain your answer.

A

comparison - metal conducts electricity but NaCl does not

explanation - Na has metallic bonding and delocalised electrons
no delocalised electrons in NaCl

33
Q

explain why sodium ion is malleable

A

layers of ions can slide over each other

34
Q

explain why Mg has higher MP than Na

A
  • Mg2+ smaller ion
  • stronger attraction between Mg2+ and delocalised electrons
35
Q

. In terms of structure and bonding explain why Sodium Chloride has a higher melting point than sulphur dioxide (SO2).

A

NaCl
- ionic bonding
- giant ionic lattice
- strong electrostatic forces of attraction between oppositely charged ions

SO2
- covalent bonding
- simple molecular
- weak IMF

  • ionic bonding stronger
  • so requires more energy to break ionic bonds
36
Q

Name the shape of HOCl and bond angle

A

v shaped
104.5 degrees

37
Q

Explain why the shape of water (H2O) changes when it becomes a hydronium ion (H3O+).

A

H2O = 2BP 2LP
H3O+ = 3BP 1LP

  • LP repel more than BP
  • bond angle changes from 107 to 104.5
38
Q

Predict the bond angle in CF2 and explain why this angle is different from the 120 bond angle in BF3:

A
  • CF2 117.5
  • CF2 has 2BP 1LP
  • LP repel more than BP
  • bond angle reduced by 2.5
39
Q

The shape of the molecule BF3 is trigonal planar.
Why is each bond angle 120 in BF3?

A
  • BF3 has 3BP 0LP
  • BP repel equally
40
Q

AlCl3 and NH3 are molecules with different bond angles. Predict and explain the bond angles in each molecule.

A

angles
AlCl3 = 120
NH3 = 107

AlCl3
3BP 0LP
All BP repel equally

NH3
3BP 1LP
LP repel more than BP

41
Q

State how co-ordinate bond is formed between ammonia and boron trichloride

A
  • LP on N donates to B atom
  • forming co-ordinate bond
42
Q

Explain, in terms of electronegativity, why the boiling point of H2S2 is lower than that of H2O2.

A
  • S is less electronegative than O
  • no H bonding between H2 S2 molecules
43
Q

Which type of bond is formed between N and B when a molecule of NH3 reacts with
a molecule of BF3?

A

Co-ordinate bonding

44
Q

For each of the following state the type of intermolecular force of attraction and use
this to explain their difference in boiling point.

H2O
boiling point 100 degrees

CHCl3
BP 62

Ar
BP -186

A

H2O
- H bonding between molecules
-

CHCl3
- permanent-dipole-dipole betweeen molecules

Ar
- VDW between atoms

  • H2O > CHCl3 > Ar
  • most energy required to break bonds on H2O then CHCl3 then Ar
45
Q

Sodium fluoride (NaF) and Chlorine monofluoride (ClF) have similar atomic masses but very different boiling points. Predict which compound will have the highest boiling point and which will have the lowest. Explain your prediction in terms of the types of bonding and the forces of interaction

A
  • NaF has higher BP
  • NaF has ionic bonds between positive and negative ions
  • ClF has permanent dipole-dipole forces between molecules
  • Permanent dipole-dipole forces are weaker than ionic bonds and require less energy to overcome
46
Q

Van der Waals’ forces exist between all molecules. Explain how these forces arise.

A
  • Electron movement in first molecule causes temporary dipole
  • This induces a dipole in another molecule
  • (Induced-temporary) attraction or δ+ attracts δ- in different/adjacent molecules
47
Q

Suggest why methaneselenol (CH3SeH) has a higher boiling point than methanethiol (CH3SH)

A
  • CH3SeH bigger molecule (larger Mr)
  • so more/stronger VDW forces between molecules
48
Q

Explain the different boiling points of NH3, F2, and Br2.

Include the names of any relevant forces and particles.

BP NH3 = -33
BP F2 = -188
BP Br2 = 59

A
  • NH3 has hydrogen-bonding forces between molecules
  • F2 has induced dipole-dipole forces between molecules
  • Br2 has induced dipole-dipole forces between molecules
  • Order or strength of attractions: Br2 > NH3 > F2
  • The induced dipole-dipole attractions in Br2 are stronger than the attractions in F2 and NH3 because Br2 is bigger and has more electrons
  • Energy required to break the attractions is greatest for Br2, then NH3, then F2
49
Q

Define electronegativity

A

ability of atom to attract pair of electrons in a covalent bond

50
Q

state and explain trend in electronegativity from Li to F

A
  • increases
  • increase in nuclear charge
  • same shielding
51
Q

suggest why graphene is a good conductor of electricity

A
  • delocalised electrons
  • free to move around to carry charge
52
Q

Explain in terms of structure and bonding why graphene has a high melting point

A
  • macromolecular with strong covalent bonds between atoms
  • requires lots of energy to break bonds
53
Q

The graph shows the melting points of the elements in period three.
In terms of their structure and bonding explain the trend in melting point from sodium to
argon. Your answer should consist of three or four parts to enable a full description

A

Na - Al
- all giant metallic lattices
- with metallic bonds between delocalised electrons and positive ions
- increases as ion charge increases
- so metallic bond strength increases

Si
- macromolecular with strong covalent bonds between atoms
- highest MP as most energy needed to break bonds

P - S - Cl
- all simple molecular structures
- VDW between molecules
- S8 = highest as biggest molecule so most electrons
- Cl2 = lowest as smallest molecule with least electrons

Ar
- Weak VDW between atoms
- smallest particle
- least number of electrons

54
Q

Name and describe the bonding in crystals of Na, MgCl2, SiO2 and Cl2

A
  • Na – Metallic crystal with a giant metallic lattice
  • Strong electrostatic forces (metallic bonds) between positive ions and
    delocalised electrons.
  • MgCl2 – Ionic crystal with a giant ionic lattice
  • Strong electrostatic forces (ionic bonds) between positive and negative
    ions
  • SiO2 – Macromolecular
  • Strong covalent bonds between atoms
  • Cl2 Molecular crystal
  • Strong covalent bonds between atoms and weak induced dipole-dipole
    attractions / forces between molecules
55
Q

working out polar molecules

A
  • trigonal planar with 3 same bonds = non-polar
  • pyramidal = always polar
  • symmetrical = non-polar
56
Q

By a consideration of the intermolecular forces involved, explain why the product HCl is a gas but the product CH3COOH is a liquid at room temperature

A

CH3COOH = strong hydrogen bonds between molecules

HCl = weak dipole-dipole forces between molecules

57
Q

 Suggest why sodium benzoate is soluble in cold water but benzoic acid is insoluble in cold water.

A

Sodium benzoate soluble because it is ionic
Benzoic acid insoluble because: despite the polarity of the COOH group / ability of COOH to form H-bonds, the benzene ring is non-polar.

AQA - A level Chemistry (25 decks)

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