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AQA A-Level Physical Chemistry > [3.1.3] Bonding > Flashcards

[3.1.3] Bonding Flashcards

Ionic Bonding, Nature of Covalent & Dative Covalent Bonds, Metallic Bonding, Bonding & Physical Properties,Shapes of Molecules & Ions, Bond Polarity and Forces Between Molecules

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1
Q

What is ionic bonding?

A

Ionic bonding is the electrostatic force of attraction between oppositely charged ions formed by electron transfer.

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2
Q

What structure do ionic crystals have?

A

Giant lattices of ions.

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3
Q

What factors affect the strength of ionic bonds thus the melting points of ionic compounds?

A
  1. When ions are smaller.
  2. When ions have a higher charge
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4
Q

Describe the radii of positive and negative ions in ionic compounds.

A

POSITIVE IONS

  • Smaller compared to their atoms because it has one less shell of electrons.
  • The ratio of protons to electrons has increased so there is greater net force on remaining electrons holding them more closely.

NEGATIVE IONS

  • Negative ions formed from groups five to seven are larger than the corresponding atom, but have the same number of protons.
  • So the pull of the nucleus is shared over more electrons and the attraction per electron is less, making the ion bigger.
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5
Q

Describe the trend in ionic radii when going down a group.

A
  • Within a group, the size of the ionic radii increases going down the group.
  • This is because as one goes down the group, ions have more shells of electrons.
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6
Q

What’s a covalent bond?

A

A covalent bond is a shared pair of electrons.

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7
Q

What’s a dative covalent bond?

A
  • A dative covalent bond forms when the shared pair of electrons in the covalent bond come from only one of the bonding atoms.
  • A dative covalent bond is also called co-ordinate bonding.
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8
Q

Draw the dative covalent bond in NH₃BF₃.

A

The direction of the arrow goes from the atom that is providing the lone pair to the atom that is deficient.

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9
Q

Draw the dative covalent bond in NH₄⁺.

A

The direction of the arrow goes from the atom that is providing the lone pair to the atom that is deficient.

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10
Q

What’s metallic bonding?

A

Metallic bonding is the electrostatic force of attraction between the positive metal ions and the delocalised electrons.

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11
Q

What is the arrangement of atoms in metallic bonding?

A

Giant metallic lattice.

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12
Q

What three main factors affect the strength of metallic bonding?

A
  1. Number of protons/strength of nuclear attraction.
    • The more protons, the stronger the bond.
  2. Number of delocalised electrons per atom (the outer shell electrons are delocalised)
    • The more delocalised electrons, the stronger the bond.
  3. Size of ion.
    • The smaller the ion, the stronger the bond.
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13
Q

Explain why magnesium has a higher melting point than sodium.

A
  • Mg has stronger metallic bonding than Na and hence a higher melting point.
    • The metallic bonding gets stronger in Mg because there are more electrons in the outer shell that are released to the sea of delocalised electons.
    • The Mg ion is also smaller because it has one more proton so the nucleus pulls its electrons more closely.
      • There is therefore a stronger electrostatic attraction between the positive metal ions and the delocalised electrons and higher energy is required to break bonds.
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14
Q

What are the four types of crystal structures?

A
  • Ionic.
  • Metallic.
  • Macromolecular (giant covalent).
  • Molecular.
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15
Q

Describe and explain the boiling & melting points, solubility, state of matter and conductivity of ionic structures.

A

BOILING & MELTING POINTS

  • High because of giant lattice of ions with strong electrostatic forces between oppositely charged ions.

SOLUBILITY

  • Soluble.

STATE OF MATTER

  • Mostly crystaline solids.

CONDUCTIVITY

  • When solid, conductivity is poor as ions cannot move/fixed in lattice.
  • When molten, conductivity is good as ions can move.
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16
Q

Describe and explain the boiling & melting points, solubility, state of matter and conductivity of molecular structures.

A

BOILING & MELTING POINTS

  • Low because weak intermolecular forces between molecules.
    • For example, van der Waals, dipole-dipole, hydrogen bonds…

SOLUBILITY

  • Generally poor.

STATE OF MATTER

  • Mostly gases and liquids.

CONDUCTIVITY

  • When solid or molten, there are no ions to conduct and electrons are localised (fixed in place) so a current cannot be carried.
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17
Q

Describe and explain the boiling & melting points, solubility, state of matter and conductivity of macromolecular (giant covalent) structures.

A

BOILING & MELTING POINTS

  • High because of many strong covalent bonds that take ** a lot of energy** to break the strong bonds.

SOLUBILITY

  • Insoluble.

STATE OF MATTER

  • Solid.

CONDUCTIVITY (USING DIAMOND & GRAPHITE AS EXAMPLES)

  • When solid, conductivity of diamond is poor as electrons can’t move as they’re localised. However, graphite’s conductivity is good as there are free delocalised electrons between its layers.
  • When molten, conductivity is poor.
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18
Q

Describe and explain the boiling & melting points, solubility, state of matter and conductivity of metallic structures.

What are some other physical properties that are also present in metallic structures?

A

BOILING & MELTING POINTS

  • High because of strong electrostatic forces between positive ions and sea of delocalised electrons in giant metallic structure.

SOLUBILITY

  • Insoluble.

STATE OF MATTER

  • Solid (except mercury).

CONDUCTIVITY

  • When solid and molten, conductivity is good as delocalised electrons can move through structure.

OTHER PHYSICAL PROPERTIES

  • Shiny.
  • Malleable.
    • The positive ions in the lattice are all identical so the planes of ions can slide easily over one another.
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19
Q

Describe and explain the structure of diamond.

Draw a diagram to show the structure of diamond.

A
  • Macromolecular - giant molecular structure.
  • Each carbon atom is bonded to four other carbon atoms by strong covalent bonds.
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20
Q

Describe and explain the structure of graphite.

Draw a diagram to show the structure of graphite.

A
  • Macromolecular - giant molecular structure.
  • Each carbon atom is bonded to three other carbon atoms by strong covalent bonds.
    • This means graphite has one delocalised electron in its structure which can carry a current.
  • Graphite also has weak van der Waals forces between its layers.
    • This means its layers can slide over each other.
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21
Q

Describe and explain the structure of ice.

Draw a diagram to show the structure of ice.

A
  • Molecular structure.
  • Covalent bonds within water molecules between oxygen and hydrogen atoms.
  • Intermolecular forces (hydrogen bonds) between water molecules.
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22
Q

Describe and explain the structure of iodine.

Draw a diagram to show the structure of iodine.

A
  • Molecular structure.
  • Covalent bonds between iodine atoms.
  • Intermolecular forces (van der Waals) between iodine molecules.
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23
Q

Describe and explain the structure of magnesium.

Draw a diagram to show the structure of magnesium.

A
  • Metallic structure.
  • Strong electrostatic forces between positive magnesium ions and sea of delocalised electrons arranged in a giant metallic lattice.
24
Q

Describe and explain the structure of sodium chloride.

Draw a diagram to show the structure of sodium chloride.

A
  • Ionic structure.
  • Giant ionic lattice of sodium and chloride ions with strong electrostatic forces between the positive sodium ions and the negative chloride ions.
25
Q

Describe the repulsion between lone pairs and bonding pairs.

A

Lone pair-lone pair repulsion is greater than lone pair-bond pair repulsion which is greater than bond pair-bond pair repulsion.

26
Q

How do you explain the shape of a molecule?

A
  1. State the number of bonding pairs and lone pairs of electrons.
  2. State that electron pairs repel and try to get as far apart as possible to minimise repulsion.
  3. If there are no lone pairs, state that electron pairs repel equally.
  4. If there are lone pairs of electrons, then state that lone pairs repel more than bonding pairs.
  5. State the actual shape and bond angle.

(Remember that lone pairs repel more than bonding pairs and so reduce bond angles by about 2.5° per lone pair)

27
Q

What equation can you use to calculate the number of lone pairs in a molecule?

A

Lone pairs = (no. of electrons in central atom - no. of bonding pairs) ÷ 2

28
Q

Draw the shape of BeCl₂. State the name of the shape, the bond angle, the number of bond pairs and the number of lone pairs.

A
  • SHAPE = linear
  • BOND ANGLE = 180°
  • BOND PAIRS = 2
  • LONE PAIRS = 0
29
Q

Draw the shape of BeCl₃ including any lone pairs. State the name of the shape, the bond angle, the number of bond pairs and the number of lone pairs.

A
  • SHAPE = trigonal planar
  • BOND ANGLE = 120°
  • BOND PAIRS = 3
  • LONE PAIRS = 0
30
Q

Draw the shape of CH₄ including any lone pairs. State the name of the shape, the bond angle, the number of bond pairs and the number of lone pairs.

A
  • SHAPE = tetrahedral
  • BOND ANGLE = 109.5°
  • BONE PAIRS = 4
  • LONE PAIRS = 0
31
Q

Draw the shape of NH₃ including any lone pairs. State the name of the shape, the bond angle, the number of bond pairs and the number of lone pairs.

A
  • SHAPE = trigonal pyramidal
  • BOND ANGLE = 107°
  • BOND PAIRS = 3
  • LONE PAIRS = 1
32
Q

Draw the shape of H₂O including any lone pairs. State the name of the shape, the bond angle, the number of bond pairs and the number of lone pairs.

A
  • SHAPE = bent
  • BOND ANGLE = 104.5°
  • BOND PAIRS = 2
  • LONE PAIRS = 2
33
Q

Draw the shape of PF₅ including any lone pairs. State the name of the shape, the bond angle, the number of bond pairs and the number of lone pairs.

A
  • SHAPE = trigonal bipyramidal
  • BOND ANGLE = 120° & 90°
  • BOND PAIRS = 5
  • LONE PAIRS = 0
34
Q

Draw the shape of SF₆ including any lone pairs. State the name of the shape, the bond angle, the number of bond pairs and the number of lone pairs.

A
  • SHAPE = octahedral
  • BOND ANGLE = 90°
  • BOND PAIRS = 6
  • LONE PAIRS = 0
35
Q

Draw the shape of XeF₄ including any lone pairs. State the name of the shape, the bond angle, the number of bond pairs and the number of lone pairs.

A
  • SHAPE = square planar
    • (Variation of the octahedral shape)
  • BOND ANGLE = 90°
  • BOND PAIRS = 4
  • LONE PAIRS = 2
36
Q

Draw the shape of BrF₅ including any lone pairs. State the bond angle, the number of bond pairs and the number of lone pairs.

A
  • BOND ANGLE = ~89°
    • (Reduced by lone pair).
  • BOND PAIRS = 5
  • LONE PAIRS = 1
37
Q

Draw the shape of I₃⁻ including any lone pairs. State the name of the shape, the bond angle, the number of bond pairs and the number of lone pairs.

A
  • SHAPE = linear
  • BOND ANGLE = 180°
  • BOND PAIRS = 2
  • LONE PAIRS = 3
38
Q

Draw the shape of ClF₃ including any lone pairs. State the bond angle, the number of bond pairs and the number of lone pairs.

A
  • BOND ANGLE = ~89°
    • (Reduced by lone pair).
  • BOND PAIRS = 3
  • LONE PAIRS = 2

(Shape is a variation of the trigonal bipyramidal)

39
Q

Draw the shape of IF₄⁺ including any lone pairs. State the bond angle, the number of bond pairs and the number of lone pairs.

A
  • BOND ANGLE = ~119° & ~89°
    • (Reduced by lone pair).
  • BOND PAIRS = 4
  • LONE PAIRS = 1

(Shape is a variation of the trigonal bipyramidal)

40
Q

Define electronegativity.

A

Electronegativity is the relative tendency of an atom in a covalent bond in a molecule to attract electrons in a covalent bond to itself.

41
Q

What is the trend in electronegativity across a period? Explain the trend.

A
  • Electronegativity increases across a period as the number of protons increases and the atomic radius decreases because electrons in the same shell are pulled in more.
  • So there is a greater force of attraction between the nucleus and electrons.
42
Q

What is the trend in electronegativity down a group? Explain the trend.

A

Electronegativity decreases down a group because the distance between the nucleus and the outer electrons increases and the shielding of inner shell electrons increases

  • So there is a weaker force of attraction between the nucleus and the electrons.
43
Q

Ionic and covalent bonding are the extremes of a continuum of bonding type. Differences in electronegativity between elements can determine where a compound lies on this scale.

If a compound contains elements of similar electronegativity, what bonding is present?

If a compound contains elements of very different electronegativity, what bonding is present?

A
  • A compound containing elements of similar electronegativity an hence a small electronegativity difference will be purely covalent.
  • A compound containing elements of very different electronegativity and hence a very large electronegativity difference will be ionic
44
Q

What are the four most electronegative elements? Out of the four, which is the most electronegative?

A
  • Fluorine - MOST ELECTRONEGATIVE
  • Oxygen.
  • Nitrogen.
  • Chlorine.
45
Q

Describe and explain how a polar covalent bond forms.

A
  • A polar covalent bond forms when the elements in the bond have different electronegativities.
  • When a bond is a polar covalent bond, it has an unequal distribution of electrons in the bond and produces a charge separation.
46
Q

How would you use partial charges to show that a bond is polar?

A
  • Polar bonds have a partially positive (δ+) end and a partially negative (δ−) end i.e. a dipole.
  • The element with the larger electronegativity has the partially negative charge.
  • The element with the smaller electronegativity has the partially positive charge.
47
Q

What’s the difference between polar and non-polar molecules?

A

NON-POLAR MOLECULES

  • Non-polar molecules are symmetric (all bonds identical and no lone pairs) and will not be polar even if individual bonds within the molecule are polar.
  • The individual dipoles on the bonds cancel out due to the symmetrical shape of the molecule.
  • There is no net dipole moment: the molecule is non-polar.

POLAR MOLECULES

  • Polar molecules are asymmetrical.
  • There is a net dipole moment: polar molecules have a permanent dipole.
48
Q

Are these molecules polar or non-polar?

  • CCl₄
  • CH₃Cl
  • CO₂
A
  • CCl₄ = non-polar
  • CH₃Cl = polar
  • CO₂ = non-polar
49
Q

Where do van der Waals forces occur? How do they form?

A

WHERE DO THEY OCCUR?

  • Van der Waals forces occur between all molecular substances and noble gases.
  • They do not occur in ionic substances.

FORMATION

  • In any molecule, the electrons are moving constantly and randomly.
  • As this happens the electron density can fluctuate and parts of the molecule become more or less negative i.e. small temporary or transient dipoles form.
  • These *instantaneous dipoles** can cause dipoles to form in neighbouring molecules.
  • These are called induced dipoles. The induced dipole is always the opposite sign to the original one.
50
Q

What factors affect the size of van der Waals forces?

A
  • The more electrons there are in the molecules, the higher chance that temporary dipoles will form.
  • This makes the van der Waals stronger between the molecules and so the boiling points of the molecules will be greater.
51
Q

Why do the boiling points of the halogens increase down the group?

A
  • The number of electrons increases in the bigger molecules as you go down the group.
  • So the size of the van der Waals forces increases between the molecules.
  • This explains why I₂ is a solid, whereas Cl₂ is a gas.
52
Q

As you go through the homologous series of alkanes, why does the boiling point increase?

A
  • The number of electrons increases in the bigger molecules.
  • So the size of the van der Waals increases between molecules.
53
Q

Why do long-chain alkanes have a higher boiling point than spherical-shaped branched alkanes?

A
  • Long-chain alkanes have a larger surface area of contact between molecules for van der Waals to form compared to spherical-shaped branched alkanes.
  • So long chain alkanes have stronger van der Waals forces.
54
Q

Where do permanent dipole-dipole forces occur? How do they form?

A

WHERE DO THEY OCCUR?

  • Permanent dipole-dipole forces occur between polar molecules.

FORMATION

  • Polar molecules have permanent dipoles.
  • They are asymmetrical and have a bond where there is a significant difference in electronegativity between the atoms.
55
Q

Where do hydrogen bonds occur? How do they form?

Draw a diagram to represent hydrogen bonding in hydrogen fluoride, ammonia and water.

A

WHERE DO THEY OCCUR?

  • Occurs in compounds that have a hydrogen atom attached to one of the three most electronegative atoms - nitrogen, oxygen and fluorine.

FORMATION

  • Nitrogen, oxygen and fluorine must have an available lone pair of electrons.
  • Which leads to a large electronegativity difference between the H and the O, N, F.
56
Q

Order the intermolecular forces in terms of strength from weakest to strongest.

A

Van der Waals -> Permanent Dipole-Dipolie -> Hydrogen Bonds

  • Molecules with hydrogen bonds will have a higher boiling and melting point than those without.

AQA A-Level Physical Chemistry (11 decks)

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