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AQA Chemistry > Physical 1 Section 1 - Unit 3: Bonding > Flashcards

Physical 1 Section 1 - Unit 3: Bonding Flashcards

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1
Q

Explain how hydrogen bonding arises (3 marks)

A
  • The large difference in electronegativity between Hydrogen and the other element
  • Causes a dipole to be formed
  • So the lone pair of electrons on the other element are attracted towards the δ+H
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2
Q

Define the term electronegativity (2 marks)

A
  • The power of an element to attract electrons towards itself
  • In a covalent bond
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3
Q

Explain why CCl4 is not polar (3 marks)

A
  • CCl4 is a tetrahedral so there is rotational symmetry
  • This means that all the bonding pairs are facing in opposing directions
  • The bonds cancel out so the overall molecule is not polar
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4
Q

Explain why lone pairs have larger bond angles than bonding pairs (2 marks)

A
  • Electron pairs repel each other and try to get to a position of maximum separation
  • Lone pairs repel more strongly than bonded pairs so the bond angle is higher
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5
Q

Compare the molecular geometries of CCl3− and CCl4 (6 marks)

A
  • Any electron pairs repel each other in molecules.
  • Electrons that are in a lone pair will repel other electrons more than bonding electrons.
  • Therefore, CCl3− will have a trigonal pyramidal shape.
  • And the bond angles in CCl3− will be 107°.
  • The shape of CCl4 will be tetrahedral.
  • And the bond angles in CCl4 will be 109.5°.
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6
Q

Explain why the covalent radius decreases across a group (3 marks)

A
  • There is increased nuclear charge
  • Electrons are added to the same shell
  • Therefore the outer electrons are pulled in more strongly
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7
Q

Explain why sodium bromide has a melting point that is higher than that of sodium, and higher than that of sodium iodide (6 marks)

A
  • Na has metallic bonding so there is attraction between the positive ions and the delocalised electrons
  • Na has a giant/lattice structure
  • There is ionic bonding in NaBr and NaI
  • Where there is attraction/ bonding between the oppositely charged ions in a giant lattice structure
  • The ionic bonds are stronger than the metallic bonds
  • And there is stronger attraction between the oppositely charged ions in NaBr than in NaI since the Br– ion is smaller than the I– ion
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8
Q

Deduce the type of intermolecular forces in SiF4 and explain how this type of intermolecular force arises and why no other type of intermolecular force exists in a sample of SiF4 (3 marks)

A
  • Van der Waals
  • One molecule induces dipole in neighbouring molecule
  • Symmetrical molecule so dipoles cancel out
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9
Q

Explain why F- is bigger than Na+ in Sodium Fluoride (2 marks)

A
  • Both Na+ and F have the same electron arrangement

- But sodium (ion) has more protons so it attracts (outer) electrons closer

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10
Q

Explain how permanent dipole-dipole forces arise between hydrogen chloride molecules.

A
  • Difference in electronegativity leads to bond polarity

- And there is an attraction between ∂+ on one molecule and ∂- on another

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11
Q

In terms of structure and bonding explain why the boiling point of bromine is different from that of magnesium. Suggest why magnesium is a liquid over a much greater temperature range compared to bromine (5 marks)

A
  • Bromine is (simple) molecular
  • Magnesium is metallic
  • Br2 has weak (van der Waals) forces between the molecules
  • So more energy is needed to overcome the stronger metallic bonds
  • Mg has a much greater liquid range because forces of attraction in liquid metal are strong(er)
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12
Q

Define ionic bonding (2)

A
  • Electrostatic force of attraction between oppositely charged ions
  • In a lattice
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13
Q

Define covalent bonding (1)

A
  • Shared pair of electrons
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14
Q

Define dative covalent bonding (1)

A
  • When the shared pair of electrons in the covalent bond come from only one of the bonding atoms
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15
Q

Define metallic bonding (1)

A
  • Electrostatic force of attraction between the positive metal ions and the delocalised electrons
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16
Q

State three factors that affect metallic bonding (3)

A
  • Number of protons (more protons, stronger bond)
  • Number of delocalised electrons (more electrons, stronger bond)
  • Size of ion (smaller ion, stronger bond)
17
Q

Properties of ionic bonding (5)

  • Melting/boiling points (2)
  • Solubility in water (1)
  • Conductivity of heat/electricity (2)
A
  • High melting/boiling point
  • Giant lattice of ions with strong electrostatic forces between oppositely charged ions
  • Soluble in water
  • Poor conductor of heat/electricity
  • Ions are fixed in lattice
18
Q

Properties of simple molecular bonding (5)

  • Melting/boiling points (2)
  • Solubility in water (1)
  • Conductivity of heat/electricity (2)
A
  • Low melting/boiling point
  • Weak intermolecular forces
  • Not soluble in water
  • Poor conductor of heat/electricity
  • Since there are no ions and electrons are fixed
19
Q

Properties of macromolecular bonding (5)

  • Melting/boiling points (2)
  • Solubility in water (1)
  • Conductivity of heat/electricity - diamond, sand and graphite (2)
A
  • High melting/boiling point
  • Many strong covalent bonds which require a lot of energy to overcome
  • Insoluble in water
  • Diamond and sand - poor conductors since electrons can’t move
  • Graphite - good conductors as free delocalised electrons between layers
20
Q

Properties of metallic bonding (5)

  • Melting/boiling points (2)
  • Solubility in water (1)
  • Conductivity of heat/electricity (2)
A
  • High melting/boiling point
  • Strong electrostatic forces between positive ions and sea of delocalised electrons
  • Insoluble in water
  • Good conductors of heat/electricity
  • As delocalised electrons can move through structure
21
Q

Linear (3)

  • Bonding pairs
  • Lone pairs
  • Angle
A
  • 2 bp
  • 0 lp
  • 180°
22
Q

Trigonal Planar (3)

  • Bonding pairs
  • Lone pairs
  • Angle
A
  • 3 bp
  • 0 lp
  • 120°
23
Q

Tetrahedral (3)

  • Bonding pairs
  • Lone pairs
  • Angle
A
  • 4 bp
  • 0 lp
  • 109.5°
24
Q

Trigonal Pyramidal (3)

  • Bonding pairs
  • Lone pairs
  • Angle
A
  • 3 bp
  • 1 lp
  • 107°
25
Q

Bent (3)

  • Bonding pairs
  • Lone pairs
  • Angle
A
  • 2 bp
  • 2 lp
  • 104.5°
26
Q

Trigonal Bipyramidal (3)

  • Bonding pairs
  • Lone pairs
  • Angle
A
  • 5 bp
  • 0 lp
  • 120 & 90°
27
Q

Octahedral (3)

  • Bonding pairs
  • Lone pairs
  • Angle
A
  • 6 bp
  • 0 lp
  • 90°
28
Q

Electronegativity trend across a period (3)

A
  • Increases
  • Number of protons increase
  • And atomic radius decreases
29
Q

Electronegativity trend down a group (3)

A
  • Decreases
  • Distance between nucleus and outer electrons increases
  • Shielding increases
30
Q

State and explain the trend in melting point of the Group II elements Ca–Ba (3 marks)

A
  • Decreases
  • Increase in size of ion/atom / more shells
  • Weaker attraction for delocalised electrons / weaker
    metallic bonding
31
Q

Iodine and diamond are both crystalline solids at room temperature. Identify one similarity in the bonding, and one difference in the structures, of these two solids. Explain why these two solids have very different melting points (6 marks)

A
  • Both covalent
  • Iodine = molecular /I2 (stated or in diagram)
  • Diamond = giant macromolecular/giant covalent
  • Iodine = Weak van der Waals’ forces / induced dipole-dipole
  • Diamond = Covalent bonds would need to be broken
  • Many / strong covalent bonds OR much energy needed
32
Q

Explain why the O–H bond in a methanol molecule is polar (2 marks)

A
  • Oxygen more electronegative than hydrogen
  • Higher e– density round oxygen atom / causes Hδ+
    Oδ–
33
Q

Explain why methanol is easy to liquefy (1 mark)

A
  • As hydrogen bonding is a strong enough force (to hold methanol molecules together in a liquid)
34
Q

Explain why molecules with 4bp and 0lp have a tetrahedral shape (2 marks)

A
  • Equal repulsion

- Between four bonding pairs

35
Q

Diamond and graphite are both forms of carbon. Diamond is able to scratch almost all other substances, whereas graphite may be used as a lubricant. Diamond and graphite both have high melting points. Explain each of these properties of diamond and graphite in terms of structure and bonding. Give one other difference in the properties of diamond and graphite (9 marks)

A

Similarity:
Both macromolecular/giant covalent

Explanation of differences:
C atoms in diamond joined to 4 other C atoms
C atoms in graphite joined to 3 other C atoms
Diamond hard / crystal strong
Because of 3-D structure / rigid structure / not layered
Graphite (soft) as layer can slide over each other

High melting points:
Covalent bonds must be broken / overcome
Which are strong / many / hard to break

Another difference:
Diamond is non-conductor of electricity, graphite is conductor

AQA Chemistry (32 decks)

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