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A Level Chemistry > 3.13 - Bonding > Flashcards

3.13 - Bonding Flashcards

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1
Q

How does electronegativity increase along the period table?

A

Electronegativity increases as you go right and up the periodic table

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2
Q

How does charge effect electronegativity?

A

Higher nuclear charge = higher electronegativity

  • number of protons increases
  • attraction between nucleus & pair of electrons in the covalent bond increases
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3
Q

How does atomic radius effect electronegativity?

A

the smaller the atomic radius the higher the electronegativity

  • bonding electrons are closer to the nucleus
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4
Q

How does shielding affect electronegativity?

A

The higher the shielding, the lower the electronegativity

  • greater number of inner shells
  • weaker attraction between nucleus & bonding pair of electrons
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5
Q

How does the difference in electronegativity affect bond polarity?

A
  • The greater the difference in electronegativity the more polar the covalent bond
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6
Q

What are the 3 types of intermolecular forces ? (Strongest to weakest)

A
  • hydrogen bonding
  • dipole-dipole
  • van der waals
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7
Q

Describe hydrogen bonds

A
  • strongest force
  • Only occur between Hydrogen AND - N,O,F
  • there is an attraction between the delta + Hydrogen and the lone pairs of other molecules
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8
Q

Explain how dipole dipole forces occur

A

Occurs in Polar molecules
Has a permanent dipole
- when one element is more electronegative than the other,
- the delta positive on one is attracted to the delta negative on the other

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9
Q

What’s the definition of electronegativity?

A

the power of an atom to attract the pair of electrons in a covalent bond

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10
Q

Explain how Van der Waals forces occur?

A
  • random movement of electrons in one molecule creates a temporary dipole
  • a dipole in a neighbouring molecule is induced
    there is a temporary attraction between delta + and -
  • the bigger the molecule the stronger the force (as there are more electrons) so take more energy to overcome
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11
Q

Describe a Non-polar molecule

A
  • no dipole
  • only has Van der Waals
  • symmetrical because individual dipoles cancel out so no net dipole moment
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12
Q

Describe a polar molecule

A
  • has a lone pair
  • has a permanent dipole
  • may have all 3 intermolecular forces
  • non symmetrical
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13
Q

Why can hydrogen bonding only occur between Hydrogen and N,O,F and not another molecule?

A
  • Molecules other than N,O,F are not very electronegative
  • so not a big enough difference In electronegativity between hydrogen and the molecule
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14
Q

Describe linear

A

2 electron pairs
180° bond angle

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15
Q

Describe triganal planar

A

3 bond pairs
120° bond angle

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16
Q

Describe tetrahedral

A

4 bond pairs
109.5°

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17
Q

Describe triganal bipyramidal

A

5 bond pairs
120/90°

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18
Q

Describe octahedral

A

6 bond pairs
90°

19
Q

Describe bent

A

2 bond pair
2 lone pair
104.5°

20
Q

Describe triganal pyramidal

A

3bp, 1 lp
107°

21
Q

Describe square pyramid

A

5bp, 1lp
89°

22
Q

Describe square planar

A

4bp, 2lp
90°

23
Q

What does each lone pair equal

A

-2.5°

24
Q

What is a dative covalent bond

A

When the shared pair of electrons in the covalent bond come from only one of the bonding atoms

25
Q

Why do some molecules with polar bonds not have a permanent dipole

A
  • the shape of the molecule is symmetrical
  • so there is an even distribution of electrons so no permanent dipole
26
Q

Describe melting points and conductivity of ionic compounds

A
  • High melting point because it is a giant lattice with strong electrostatic forces of attraction between oppositely charged ions.
  • can conduct electricity when molten/in solution because the ions can move
27
Q

Describe properties simple molecular structures (2nd type of covalent bonding)

A

low melting point due to weak intermolecular forces between molecules

28
Q

Why do metallic compounds have high melting and boiling points?

A

Strong electrostatic forces between positive ions and sea of delocalised electrons

29
Q

How to explain shape of a molecule

A
  1. State number of bond pairs and lone pairs of electrons
  2. State that electron pairs repel and try to get as far apart as possible
  3. If there are no lone pairs state that the electron pairs repel equally
  4. If there are lone pairs of electrons then state that lone pairs repel more than bonding pairs
  5. State actual shape and bond angle
30
Q

What’s the definition of ionic bonding

A

The electrostatic force of attraction between oppositely charged ions formed by electron transfer
METAL + NON METAL

31
Q

What makes ionic bonding stronger and the melting points higher

A

When Ions are smaller and have higher charges

32
Q

What happens to ionic radius down a group

A

It increases because the ions have more shells of electrons as you go down a group

33
Q

What 3 factors effect the strength of metallic bonding

A
  1. Number of protons. The more protons the stronger the bond
  2. The number of outer shell/delocalised electrons
  3. Size of ion. The smaller the ion the stronger the bond
34
Q

What are the 3 types of bonding

A
  1. IONIC
  2. COVALENT (2 types) - macromolecular & simple molecular
  3. METALLIC
35
Q

What is covalent bonding between

A

Non metals

36
Q

Describe properties macro molecular structures (1st type of covalent bonding)

A

Macromolecular: eg diamond & graphite
- high melting point due to lots of strong covalent bonds
- slippery due to weak VDW between layers

37
Q

How does electronegativity change across a period

A

Increases because:
the number of protons increases and atomic radius decreases because the ELECTRONS IN THE SAME SHELL are pulled in more

38
Q

How does electronegativity change down a group

A

Decreases because the distance between the nucleus and the outer shell electrons increases and the shielding of inner shell electrons increases

39
Q

What does difference in electronegativity mean for a compound’s bonding type

A

Small difference - covalent
Big difference - ionic

40
Q

Why does a polar covalent bond form

A

When there is an unequal distribution of electrons in the bond which produces a charge separation/dipole

41
Q

How does symmetry effect polar and non polar molecules and why

A

Symmetrical molecule (no lone pairs) - NOT POLAR
- because the individual dipoles on the bonds cancel out so there is no net dipole moment

42
Q

What effects strength of van der waals forces

A

More electrons = bigger molecule = higher chance temporary dipoles will form = stronger VDW

43
Q

Why is a compound brittle

A
  • applying force disrupts the arrangement of ions
  • ions with like charges repel

A Level Chemistry (13 decks)

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