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OCR A LEVEL Chemistry > 2.2.2 Bonding and Structure > Flashcards

2.2.2 Bonding and Structure Flashcards

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1
Q

What is Ionic Bonding?

A

The electrostatic force of attraction between 2 oppositely charged ions; formed by the transfer of electrons.

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2
Q

What does ionic bonding happen between?

A

Metals and Non-metals

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3
Q

At which two points is ionic bonding its strongest?

A
  1. When the ion is small (small ionic radius)
  2. When the ion has a larger charge
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4
Q

Example of an ionic compound?

A

NaCl

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5
Q

What is the arrangement of ionic compounds?

A

Giant Ionic Lattice
3D structure

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6
Q

In a giant ionic lattice, what do the ions do?

A

Each ion is surrounded by oppositely charged ions therefore the ions in the lattice attract in all directions.

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7
Q

What are the properties of ionic compounds?

A

Soluble in Water

High boiling point and melting point

Can conduct Electricity (only molten or aqueous)

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8
Q

High Boiling and melting point (ionic):

A

Strong electrostatic forces between oppositely charged ions.
A great amount of energy is required to overcome these forces, giving it a high MP and BP

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9
Q

Electrical conductivity (ionic):

A

When in their solid state, ionic compounds cannot carry charge or electricity because the ions are not free to move and are in a fixed position.

When dissolved or molten, ionic compounds can carry charge and electricity through the structure because the ions become free to move.

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10
Q

Solubility (ionic):

A

An ionic substances can dissolve in polar solvents (e.g. water)
The polar solvent disrupts the lattice as the slight charges (dipoles) in the polar substance attract the ions in the lattice, creating a solution.

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11
Q

What is covalent bonding?

A

The electrostatic force of attraction between a shared pair of electrons and the nuclei of the bonded atoms

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12
Q

Single covalent bonds consist of only…

A

one shared pair of electrons

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13
Q

Multiple covalent bonds consist of …

A

2 or more shared pairs of electrons

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14
Q

Example of single covalently bonded molecule

A

Hydrogen H-H

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15
Q

Example of multiple covalently bonded molecule

A

Oxygen (O=O)
Nitrogen (triple)
Carbon dioxide (O=C=O)

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16
Q

What are lone pairs?

A

Electron pairs that haven’t been used up in covalent bonding

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17
Q

What can lone pairs do?

A

Give a concentrated region of negative charge around the atom.

decreases bond angle by 2.5°

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18
Q

What is a dative covalent bond?

A

When only one atom provides both electrons which will be shared as a pair.
(one atom provides an electron pair –> these are usually its lone pairs)

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19
Q

Examples of Molecules with dative covalent bonds

A

Ammonium ion (3 covalent and 1 dative)
Oxonium ion (2 covalent and 1 dative)

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20
Q

What is bond enthalpy?

A

The measure of covalent bond strength

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21
Q

Bond enthalpy shows you …

A

The average amount of energy required to break a covalent bond.

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22
Q

The larger the enthalpy value…

A

the stronger the covalent bond

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23
Q

If the covalent bond is strong, the bond enthalpy should have a…

A

Positive value (endothermic)

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24
Q

Endothermic reactions:

A

Breaking bonds (+)

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25
Q

Exothermic reactions:

A

Forming bonds (-)

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26
Q

What are two covalent structures

A

Simple covalent molecules
Giant covalent molecules

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27
Q

Name the properties of a simple covalent molecule:

A

Low melting point and boiling point

Low solubility

Poor electrical conductivity

Usually found as a liquid or gas

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28
Q

Low melting point and boiling point (simple covalent molecule):

A

Weak intermolecular forces of attraction (London forces) between the molecules
A minimal amount of energy is required to overcome these forces in the molecule

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29
Q

Electrical conductivity (simple covalent molecule):

A

Low because, whether in a molten or aqueous state, no delocalised electrons or ions can carry charge through the structure.

30
Q

Examples of simple covalent molecules:

A

Water
Nitrogen
Hydrogen
Ammonia
Carbon dioxide
Oxygen
Iodine
Ice (Solid)

31
Q

What defines the shape of a molecule or ion?

A

The number of electron pairs it has
(Either bonded or lone pairs)

32
Q

Define the electron pair repulsion theory:

A

All electrons are negatively charged
Electron pairs automatically repel each other
The shape the molecule makes allows the electron pairs to be as far as possible to minimise repulsion

33
Q

What happens if the central atom is surrounded by bonded pairs of electrons

A

The bonded pairs repel each other equally

34
Q

What are the 7 shapes of molecules?

A

Linear
Trigonal planar
Tetrahedral
Trigonal pyramidal
Trigonal Bipyramidal
Bent (Non-linear)
Octahedral

35
Q

Linear:

A

2 bonded pairs (double)

0 lone pairs

180 degrees

36
Q

Linear example:

A

Carbon dioxide (O=C=O)

37
Q

Trigonal Planar:

A

3 bonded pairs

0 lone pairs

120 degrees

38
Q

Trigonal planar example:

A

BF3

39
Q

Tetrahedral :

A

4 bonded pairs

0 lone pairs

109.5 degrees

40
Q

Tetrahedral example:

A

Methane (CH4)

41
Q

Trigonal pyramidal:

A

3 Bonded pairs

1 lone pair

107 degrees

42
Q

Trigonal pyramidal example:

A

Ammonia (NH3)

43
Q

Trigonal Bipyramidal:

A

3 Bonded pairs

2 Lone pairs

120 and 90 Degrees

44
Q

Bent (Non-linear):

A

2 Bonded pairs

2 lone pairs

104.5 degrees

45
Q

Bent example:

A

Water (H2O)

46
Q

Octahedral:

A

6 bonded pairs

0 lone pairs

90 degrees

47
Q

Octahedral example:

A

SF6

48
Q

Lone pairs…

A

Reduce bond angles by 2.5 degrees

49
Q

Lone pairs repel…

A

More than bonded pairs

50
Q

What is electronegativity?

A

The tendency for an atom to attract an electron in a covalent bond to itself

51
Q

What are the most electronegative elements?

A

Nitrogen
Oxygen
Fluorine
Chlorine

52
Q

What happens to electronegativity as you go across a period?

A

Increases
The nuclear charge increases (proton number increases)
Atomic radius decreases because of the strong electrostatic force between the positive nucleus and negative electrons.
Therefore the electrons get pulled towards the nucleus
Electrons get easier to gain and harder to lose

53
Q

What happens to electronegativity as you go down a group?

A

Decreases
Electron shielding increases due to more shells and atomic radius increases
The distance between the nucleus and the outer electron increases so the attraction between them weakens
Electrons become easier to lose and harder to gain

54
Q

Compounds with slight electronegative differences will be …

A

Covalent

55
Q

Compounds with large electronegative differences will be …

A

Ionic

56
Q

What side of the periodic table has the most electronegative elements

A

Top right

57
Q

What is a permanent dipole-dipole interaction (Polar covalent bond)?

A

When the elements in a molecule have different electronegativities, it causes an unequal distribution of charge/electrons, producing a separation of charge (A dipole)

This creates a polar molecule.

58
Q

What element will have the negative dipole?

A

The element with the larger electronegativity

59
Q

Symmetrical molecules have…

A

Bonded pairs and no lone pairs

60
Q

Example of a symmetrical molecule:

A

CO2

61
Q

Symmetrical molecules are…

A

Non-polar

62
Q

Why are symmetrical molecules non-polar?

A

The dipoles cancel out, even if there are individual polar bonds within the molecule since there is an equal distribution of electrons (bonded electron pairs) so the dipoles cancel out.

63
Q

Properties of permanent dipole-dipole interactions:

A

Stronger than London forces
Found in asymmetrical molecules
There is always a significant difference in electronegativity.

64
Q

Induced dipole-dipole interaction only occurs in what?

A

Simple covalent molecules and noble gases

65
Q

How are London forces (induced dipole-dipole interactions) created?

A

electrons are constantly moving in a molecule.
Electron density constantly fluctuates, so parts of the molecule can become more negative than others, inducing temporary dipoles (transient dipoles).
These induced dipoles induced more dipoles in neighbouring molecules

66
Q

London forces =

A

Induced dipole-dipole interactions

67
Q

What can affect London forces?

A

The more electrons there are, the higher the chance for temporary dipole to be produced

This increases the strength of the interaction, increasing Boiling points.

68
Q

Hydrogen bonds only occur?

A

When one of the hydrogen atoms in a molecule is attached to one of the most electronegative atoms of either N,O,F or Cl

69
Q

Hydrogen bonding requires

A

A large electronegativity difference and there must be a lone pair

70
Q

Hydrogen bonding is stronger than both…

A

Permanent dipole dipole interactions
London forces

71
Q

Why does water have a high BP?

A

Because of the surplus amount of strong hydrogen bonds , therefore a high amount of energy is needed to break these hydrogen bonds

72
Q
A

OCR A LEVEL Chemistry (5 decks)

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