Brainscape's Knowledge GenomeTM


Top Flashcards (Certified)
All Flashcards

A-Level Chemistry > Chapter 6 Shapes and intermolecular forces > Flashcards

Chapter 6 Shapes and intermolecular forces Flashcards

1
Q

Electron-pair repulsion theory

A
  • Electron pairs surrounding a central atom determine the shape of the molecule
  • The electron pairs repel one another so that they are arranged as far apart from each other as possible
  • The arrangement of electron pairs minimises repulsion -> holds the bonds in a definite shape
  • Different numbers of electrons result in different shapes
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
2
Q

Wedges

A
  • Solid line
  • Solid wedge
  • Dotted wedge
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
3
Q

Solid line

A

Bond in the plane of the paper

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

Solid wedge

A

Bond comes out of the plane of the paper

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

Dotted wedge

A

Bond goes into the plane of the paper

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
6
Q

Bonding pair and lone pair repulsion comparison

A
  • Lone pairs repel more than bonding pairs
  • for each lone pair, bonding angle decreases by 2.5 degrees
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
7
Q

Tetrahedral

A
  • 4 bonding pairs
  • 0 lone pairs
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
8
Q

Pyramidal

A
  • 3 bonding pairs
  • 1 lone pair
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
9
Q

Non linear

A
  • 2 bonding pairs
  • 2 lone pairs
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
10
Q

4 electron pair shapes

A
  • Tetrahedral
  • Pyramidal
  • Non-linear
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
11
Q

3 electron pairs shapes

A
  • Trigonal planar
  • Bent
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
12
Q

Trigonal planar

A
  • 3 bonding pairs
  • 0 lone pairs
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
13
Q

Bent

A
  • 2 bonding pairs
  • 1 lone pair
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
14
Q

2 electron pair shapes

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

linear

A

2 bonding pairs + 0 lone pairs

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

octahedral

A
  • 6 bonding pairs
  • 0 lone pairs
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
17
Q

Shapes of ions

A
  • Ammonium (NH4+)
  • Sulfate (SO4 2-)
  • Carbonate (CO3 2-)
  • Nitrate (NO3 - )
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
18
Q

Ammonium ion

A
  • 4 bonding pairs + 0 lone pairs
  • Tetrahedral
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
19
Q

Sulfate ion

A

Tetrahedral

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

Carbonate ion

A

Trigonal planar

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

Nitrate ion

A

Trigonal planar

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

Electronegativity

A

The ability of an atom to attract a shared pair of electrons in a covalent bond

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

Pauling electronegativity scale

A

In periodic table left to right:
* Nuclear charge increases
* Atomic radius decreases

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

What happens to electronegativity as Pauling value increases?

A

Electronegativity of the atoms of that element increases

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
25
The most electronegative atoms
N O F Cl
26
Least electronegative atoms
Group 1
27
What happens when electronegativity difference of atoms is large
The more electronegative bonding atoms will gain control of the electrons, and the bond becomes ionic (not covalent)
28
Diatomic molecules
* Atoms are of the same element * Bonded electron pair is shared evenly
29
When bonded atoms are different:
* Atoms are different sizes * Nuclear charges are different * Shared pair of electrons maybe closer to one nucleus than the other
30
Non-polar bonds definition
Bonded electron pair is shared equally between the bonded atoms
31
Characteristics of non-polar bond
* Pure covalent bond
32
Example of non-polar bond
Hydrocarbon liquids (e.g. hexane)
33
When do non-polar bonds occur
* Bonded atoms are the same OR * Bonded atoms have same or similar electronegativity
34
Shape of non-polar molecules
Symmetrical
35
Polar bonds
bonded electron pair is shared unequally between the bonded atoms
36
Characteristics of polar bonds
Polar covalent bond
37
Example of polar bond
e.g. hydrogen chloride * Cl is more electronegative than H so Cl has ,ore attraction to bonded pair of electrons = polar covalent bond
38
δ signs
* δ+ : Atoms with smaller electronegativity * δ- : Atoms with larger electronegativity
39
Dipole
Seperation of opposite charges
40
When do polar bonds occur
* Bonded atoms are different OR * Bonded atoms have different electronegativity
41
Shape of polar bonds
Assymetrical
42
Solubility of ionic compounds
* Positive ions are attracted to oxygen (δ-) in water molecules * Negative ions are attracted to hydrogen (δ+) in water molecules
43
Intermolecular forces
Weak attraction between dipoles of different molecules
44
What do intermolecular forces determine?
physical properties
45
What do covalent bonds determine?
identity + chemical reactions of molecules
46
London forces
Induced dipole-dipole interactions
47
Strength of intermolecular forces
Weakest -> strongest 1. Induced dipole-dipole interactions 2. Permanent dipole-dipole interactions 3. Hydrogen bonding 4. Single covalent bonds
48
Process of formation of induced dipole-dipole interactions
1. Movement of electrons produces a changing dipole in a molecule 2. At any instant, an instantaneous dipole will exist, but its position is constantly shifting 3. The instantaneous dipole induces a dipole on neighbouring molecules 4. The induced dipole induces further dipoles on neighbouring molecules, which then attract one another
49
State of induced dipole-dipole interactions
* Temporary * Can disappear at any moment
50
What happens as the number of electrons in each molecule increases?
1.The larger the instantaneous-induced dipole 2. The larger the induced-induced dipole 3. The stronger the attractive force between molecules (higher b.p.)
51
What molecules do London forces occur in?
Both polar and non-polar molecules
52
What molecules do permanent dipole-dipole interactions occur in?
polar molecules ONLY
53
Simple molecular substance
* Made up of simple molecules * Small units containing a definite number of atoms with a definite molecular formula
54
Structure of simple molecular substances as solids
Regular structure of simple molecular lattice
55
Simple molecular lattice
* Weak intermolecular forces hold the molecules together * Atoms within each molecule are held together by strong covalent bonds
56
Properties of simple molecular substances
* Low m.p. and b.p. -> weak intermolecular forces are easily overcome by small amount of heat energy * The strong covalent bonds do not break
57
Solubility of simple molecular substances
Different for * Polar simple molecular substances * Non-polar simple molecular substances
58
Solubility of non-polar simple molecular substances: Non-polar molecule + non-polar solvent
SOLUBLE intermolecular forces form between the molecules and the solvent These interactions weaken the intermolecular forces in the simple molecular lattice -> and break
59
Solubility of non-polar simple molecular substances: Non-polar molecule + polar solvent
INSOLUBLE Little interaction between molecule and solvent Intermolecular forces in solvent are too strong to be broken
60
Solubility of polar simple molecular substances: Polar molecule + polar solvent
SOLUBLE Attract one another (similar to dissolving of ionic compounds)
61
Electrical conductivity of simple molecular substances
Do not conduct electricity due to: * No mobile charged particles
62
Hydrogen bond
Special type of permanent dipole-dipole interaction between molecules
63
What does a hydrogen bond contain?
* An electronegative atom with a lone pair of electrons (O, F, N) * A H atom attached to an electronegative atom (H-O, H-N, H-F)
64
What is a hydrogen bond shown by?
Dashed line
65
What gives water unique properties?
Hydrogen bonding
66
Unusual properties of water
* Ice (solid) is less dense than water (liquid) * Water has relatively high m.p. and b.p.
67
Why is ice less dense than water
* Hydrogen bonds hold water molecules apart in an open lattice structure * The water molecules in ice are further apart than in water * Solid ice is less dense than liquid water and floats
68
Structure of ice
* Water has 2 lone pairs so can bond to 4 hydrogen atoms * These hydrogen bonds extend outwards, forming an open tetrahedral lattice, full of holes * Bond angle is close to 180 degrees
69
What happens when ice melts
The ice lattice collapses and the molecules move close together
70
Why does water have a relatively high melting point and boiling point?
* Hydrogen bonds are extra forces * Energy is needed to break the hydrogen bonds * When the ice lattice breaks, the rigid arrangement of bonds in ice is broken. When water boils, the hydrogen bonds break completely
71
Ionic lattice
Repeated pattern of oppositely charged ions
72
Covalent bond
shared pair of electrons

A-Level Chemistry (20 decks)

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