Brainscape's Knowledge GenomeTM


Top Flashcards (Certified)
All Flashcards

OCR A LEVEL Chemistry > 2.2.2 Bonding and Structure > Flashcards

2.2.2 Bonding and Structure Flashcards

(72 cards)

Start Studying
1
Q

What is Ionic Bonding?

A

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

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
2
Q

What does ionic bonding happen between?

A

Metals and Non-metals

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
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
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
4
Q

Example of an ionic compound?

A

NaCl

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
5
Q

What is the arrangement of ionic compounds?

A

Giant Ionic Lattice
3D structure

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
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.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
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)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
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

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
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.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
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.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
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

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
12
Q

Single covalent bonds consist of only…

A

one shared pair of electrons

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
13
Q

Multiple covalent bonds consist of …

A

2 or more shared pairs of electrons

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
14
Q

Example of single covalently bonded molecule

A

Hydrogen H-H

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
15
Q

Example of multiple covalently bonded molecule

A

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

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
16
Q

What are lone pairs?

A

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

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
17
Q

What can lone pairs do?

A

Give a concentrated region of negative charge around the atom.

decreases bond angle by 2.5°

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
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)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
19
Q

Examples of Molecules with dative covalent bonds

A

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

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
20
Q

What is bond enthalpy?

A

The measure of covalent bond strength

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
21
Q

Bond enthalpy shows you …

A

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

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
22
Q

The larger the enthalpy value…

A

the stronger the covalent bond

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
23
Q

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

A

Positive value (endothermic)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
24
Q

Endothermic reactions:

A

Breaking bonds (+)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
25
Exothermic reactions:
Forming bonds (-)
26
What are two covalent structures
Simple covalent molecules Giant covalent molecules
27
Name the properties of a simple covalent molecule:
Low melting point and boiling point Low solubility Poor electrical conductivity Usually found as a liquid or gas
28
Low melting point and boiling point (simple covalent molecule):
Weak intermolecular forces of attraction (London forces) between the molecules A minimal amount of energy is required to overcome these forces in the molecule
29
Electrical conductivity (simple covalent molecule):
Low because, whether in a molten or aqueous state, no delocalised electrons or ions can carry charge through the structure.
30
Examples of simple covalent molecules:
Water Nitrogen Hydrogen Ammonia Carbon dioxide Oxygen Iodine Ice (Solid)
31
What defines the shape of a molecule or ion?
The number of electron pairs it has (Either bonded or lone pairs)
32
Define the electron pair repulsion theory:
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
What happens if the central atom is surrounded by bonded pairs of electrons
The bonded pairs repel each other equally
34
What are the 7 shapes of molecules?
Linear Trigonal planar Tetrahedral Trigonal pyramidal Trigonal Bipyramidal Bent (Non-linear) Octahedral
35
Linear:
2 bonded pairs (double) 0 lone pairs 180 degrees
36
Linear example:
Carbon dioxide (O=C=O)
37
Trigonal Planar:
3 bonded pairs 0 lone pairs 120 degrees
38
Trigonal planar example:
BF3
39
Tetrahedral :
4 bonded pairs 0 lone pairs 109.5 degrees
40
Tetrahedral example:
Methane (CH4)
41
Trigonal pyramidal:
3 Bonded pairs 1 lone pair 107 degrees
42
Trigonal pyramidal example:
Ammonia (NH3)
43
Trigonal Bipyramidal:
3 Bonded pairs 2 Lone pairs 120 and 90 Degrees
44
Bent (Non-linear):
2 Bonded pairs 2 lone pairs 104.5 degrees
45
Bent example:
Water (H2O)
46
Octahedral:
6 bonded pairs 0 lone pairs 90 degrees
47
Octahedral example:
SF6
48
Lone pairs...
Reduce bond angles by 2.5 degrees
49
Lone pairs repel...
More than bonded pairs
50
What is electronegativity?
The tendency for an atom to attract an electron in a covalent bond to itself
51
What are the most electronegative elements?
Nitrogen Oxygen Fluorine Chlorine
52
What happens to electronegativity as you go across a period?
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
What happens to electronegativity as you go down a group?
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
Compounds with slight electronegative differences will be ...
Covalent
55
Compounds with large electronegative differences will be ...
Ionic
56
What side of the periodic table has the most electronegative elements
Top right
57
What is a permanent dipole-dipole interaction (Polar covalent bond)?
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
What element will have the negative dipole?
The element with the larger electronegativity
59
Symmetrical molecules have...
Bonded pairs and no lone pairs
60
Example of a symmetrical molecule:
CO2
61
Symmetrical molecules are...
Non-polar
62
Why are symmetrical molecules non-polar?
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
Properties of permanent dipole-dipole interactions:
Stronger than London forces Found in asymmetrical molecules There is always a significant difference in electronegativity.
64
Induced dipole-dipole interaction only occurs in what?
Simple covalent molecules and noble gases
65
How are London forces (induced dipole-dipole interactions) created?
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
London forces =
Induced dipole-dipole interactions
67
What can affect London forces?
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
Hydrogen bonds only occur?
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
Hydrogen bonding requires
A large electronegativity difference and there must be a lone pair
70
Hydrogen bonding is stronger than both...
Permanent dipole dipole interactions London forces
71
Why does water have a high BP?
Because of the surplus amount of strong hydrogen bonds , therefore a high amount of energy is needed to break these hydrogen bonds
72

OCR A LEVEL Chemistry (6 decks)

Brainscape helps you reach your goals faster, through stronger study habits.
© 2025 Bold Learning Solutions. Terms and Conditions