M2 - Chapter 6 Flashcards
(114 cards)
1
Q
why do scientists use the electron repulsion theory
A
for explaining and predicting shapes of molecules and polyatomic ions.
2
Q
what is the electron repulsion theory
A
- An electron has negative charge
- Electron pairs repel one another so they occupy a position that will minimise this repulsion – as far apart as possible
3
Q
different number of electrons…
A
different shapes
4
Q
what does a solid line represent in a 3D diagram
A
represents a bond in the plane of the paper
5
Q
what does a solid wedge represent in a 3D diagram
A
comes out of plane of paper
6
Q
what does a dotted wedge represent in a 3D diagram
A
goes into plane of paper
7
Q
pattern in bond angles
A
Bond angle is reduced by about 2.5 degrees for each lone pair
8
Q
difference between bonding pair and bonding region
A
can be used interchangeably
bonding pair is a single covalent
bonding region is a double covalent
9
Q
greater number the electron pairs… (relate to bond angle)
A
smaller the bond angle
10
Q
define lone pair
A
any pair of electrons that is NOT involved in a covalent bond - real definition
11
Q
define bonded pair
A
one bonded pair of electrons in a covalent molecule
12
Q
why repels more strongly, a bonded pair or a lone pair
A
a lone pair
13
Q
why does a lone pair bond more strongly
A
- Lone pair of electrons is slightly closer to central atom and
- occupies more space than a bonded pair.
= Lone pair repels MORE strongly than bonding pair.
14
Q
bond pairs + lone pairs of carbon dioxide
A
2 bonding pairs + 0 lone pairs
15
Q
describe the shape of carbon dioxide
A
2 bonding pairs + 0 lone pairs • Linear shape
• 180 degrees angle
• SYMMETRICAL
16
Q
bond pairs + lone pairs of boron trifluoride
A
3 bonding pairs + 0 lone pairs
17
Q
shape of boron trifluoride
A
3 bonding pairs + 0 lone pairs • Triagonal planar shape
• 120 degrees angle
18
Q
bond pairs + lone pairs of methane
A
4 bonding pairs + 0 lone pairs
19
Q
shape of methane
A
4 bonding pairs + 0 lone pairs • - Tetrahedral shape
• 109.5 degrees angle
• SYMMETRICAL
20
Q
bond pairs + lone pairs of ammonia
A
3 bonding pairs + 1 lone pair
21
Q
shape of ammonia
A
3 bonding pairs + 1 lone pair • Pyramidal shape
• 107 degrees
22
Q
bond pairs + lone pairs of water
A
2 bonding pairs + 2 lone pairs
23
Q
shape of water
A
2 bonding pairs + 2 lone pairs • Non-linear shape
• 104.5 degree angle
24
Q
bond pairs + lone pairs of sulphur hexaflouride
A
6 bonding pairs + 0 lone electrons
25
shape of sulphur hexafluoride
6 bonding pairs + 0 lone electrons
- Octahedral shape
- 90 degrees
26
shape of 2 bonding pairs + 0 lone pairs
• Linear shape
• 180 degrees angle
• SYMMETRICAL
eg. CO2
27
shape of 3 bonding pairs + 0 lone pairs
• Triagonal planar shape
• 120 degrees angle
eg. BF3
28
shape of 4 bonding pairs + 0 lone pairs
• Tetrahedral shape
• 109.5 degrees angle
• SYMMETRICAL
eg. CH4
29
shape of 3 bonding pairs + 1 lone pair
• Pyramidal shape
• 107 degrees
eg. NH3
30
shape of 2 bonding pairs + 2 lone pairs
• Non-linear shape
• 104.5 degree angle
eg. H20
31
shape of 6 bonding pairs + 0 lone electrons
- Octahedral shape
- 90 degrees
eg. SF6
32
angles of linear shape
180
33
angles of trigonal planar
120
34
angles of tetrahedral shape
109.5
35
angles of pyramidal shape
107
36
angles of non-linear
104.5
37
angles of octahedral shape
90 degrees
38
Table summary of all bonding pairs
number of BP + LP // diagram // shape name // angle degree
39
important note when counting bond pairs/lone pairs
Around the central atom!
40
define covalent bond
* The strong electrostatic attraction between a shared pair of electrons and the nuclei of bonded atoms
41
define electronegativity
* A measure of the tendency of an atom to attract a bonding pair of electrons in a covalent bond
42
what does it mean to have a greater electronegativity
* the more it attracts electrons towards it
43
if the atoms in a bonded molecule are the same, what is the polarity / electronegativity like
the electrons shared equally - non-polar
44
will the more electronegative atom have a smaller to larger share of the electrons
larger
45
relationship between electronegativity and dipoles
* The more electronegative – more electrons – slightly negative charge – negative dipole
* The less electronegative – less electrons – slightly positive charge – positive dipole
46
Electronegativity of group 0
There is none - 0
47
# 3
Factors affecting electronegativity
* Nuclear charge
* Atomic radius
* Electron shielding
48
what is the Pauling scale
* Compares electronegativity
49
How does electronegativity change across a period
Increases
50
Why does electronegativity increase across a period
* Because nuclear charge increases
* (One more proton on each element)
* And atomic radius decreases
* (Increased charge – pulls in shells
* – decreasing atomic radius)
* Attracts a bonding pair more strongly
51
How does electronegativity change down a group
Decreases
52
Why does electronegativity decrease down a group
* Atomic radius increases
* More shells = more shielding
* The bonding pair of electrons are attracted less strongly to the nuclei of an atom
53
Electronegativity difference in covalent
0
54
Electronegativity difference in Polar covalent
0 - 1.8
55
Electronegativity difference in Ionic
\>1.8
56
Spectrum of bonds
57
What is a non-polar bond
* The electron pair is shared equally between atoms
58
When is a bond non-polar
* The bonded atoms are the same
* The bonded atoms have similar electronegativities
59
What is a polar bond
* The electron pair is shared unequally between atoms
* This forms a permanent dipole
60
What does a polar bond form
A permanent dipole
61
What is a dipole
A separation of opposite charges
62
When is a bond polar
The bonded atoms have different electronegativities
63
What is a polar molecule
* A molecule is polar if the individual dipoles do NOT oppose each other and cancel each other out
64
When would there be a polar molecule
* the shape of the molecule is NOT symmetrical - dipoles don’t cancel out
65
What is a non-polar molecule
* A molecule is non-polar if the individual dipoles oppose each other and cancel each other out
66
When would a molecule be non polar
If the shape is symmetrical - dipoles cancel each other out
67
Example of a non-polar molecule
CO2
68
Example of a polar molecule
H2O
69
Solubility of polar and non polar substances
* ‘Like’ dissolves ‘like’
* Polar solutes dissolve in polar solvents
* Non-polar solutes dissolve in non-polar solvents
70
What is an intermolecular force
* Weak interactions between dipoles of different molecules
71
What are intermolecular forces responsible for
* Responsible for physical properties
* E.g. melting/boiling point
72
# 3
Types of intermolecular forces
* Induced dipole-dipole interactions – London forces
* Permanent dipole-dipole interactions
* Hydrogen bonding
73
What can some intermolecular forces be referred to as
* Both London forces and induced dipole-dipole interactions can be referred to as Van der Waal’s forces
74
Table of strengths for intermolecular forces - don’t need to memorise exact numbers
75
Strongest type of intermolecular force
Hydrogen bonds
76
Weakest type of intermolecular force
London forces
77
How do induced dipole-dipole interactions (London forces) occur
* The constant movement of electrons produces an instantaneous dipole – the position of this is constantly changing
The instantaneous dipole induces a dipole on a neighbouring molecule
* The induced dipole induces further dipoles on neighbouring molecules, which then attract one another
* The more electrons in a molecule, the stronger the London forces
* Because more electrons can be on one side/atom at once
78
How can you increase the strength of a London force
* more electrons = more electrons can be at one side at once = larger charge difference
79
How should we think of electrons
* Think of electrons as a cloud and is not in one shell but constantly moving in an area, so at one point it will be closer to one atom then another
80
What are London forces - indused dipole-dipole
Weak intermolecular forces that exist between ALL molecules
81
What are permanent dipole-dipole interactions
* Act between the permanent dipoles in different polar molecules
82
If something has a permanent dipole-dipole interaction, what else will it have
London force - acts between ALL molecules
83
More specifically what do London forces act between
ALL SIMPLE COVALENT
84
What are hydrogen bonds
* Special type of permanent dipole-dipole interaction found between molecules containing:
* An electronegative atom with a lone pair (O/N/F)
* i.e. H – O, H – N, H – F
* A hydrogen atom attached to an electronegative atom
85
How do you draw hydrogen bonds
* Dipoles must be shown
* Lone pairs must be shown
* The hydrogen bond is represented by a dashed line
86
What is a simple molecule
* Small units containing a definite number of atoms with a definite molecular formula
87
What do simple molecules form in the solid state
* a regular structure called a simple molecular lattice
88
How are molecules and atoms held together in a simple molecular lattice
* The molecules ae held together by weak intermolecular forces
* The atoms within each molecule are bonded together strongly by covalent bonds
89
# 3
Properties of simple molecular substances
* Low Melting and boiling point
* Solubility
* Electrical conductivity
90
How are simple molecular substances bonded
Covalently
91
why do simple molecular substances have a low melting and boiling point
* The weak intermolecular forces can be broken even by the energy present at low temps
* Therefore they have low boiling/melting points
92
What attraction is broken when a simple molecular lattice / substance melts/boils
* Only the weak intermolecular forces break
* The covalent bonds do NOT break
93
Example of a non-polar solvent
Hexane
94
Why are non-polar simple molecular substances soluble in non-polar solvents
* When a simple molecular compound is added to a non-polar solvent (e.g. Hexane), intermolecular forces form between the molecules and the solvent
* The interactions weaken the intermolecular forces in the simple molecular lattice.
* The intermolecular forces break and the compound dissolve
95
why is a non-polar simple molecular substance not soluble in a polar solvent
* When a simple molecular substance is added to a polar solvent, there is little interaction between the molecules in the lattice and the solvent molecules
* The intermolecular bonding within the polar solvent is too strong to be broken
96
Why may polar substances dissolve in polar solvents
* as the polar solute molecules and the polar solvent molecules can attract each other
* Similar to dissolving an ionic compound
97
Example of dissolving polar in polar
* E.g. sugar dissolves in water (a polar solvent)
* Sugar = polar covalent compound with many O-H bonds
* These attract and bond with polar water molecules
98
Why is solubility of a polar solute difficult to predict
depends on strength of the dipole
99
How can a substance dissolve in both polar and non-polar solvents and examole
If they have both polar and non-polar bonds
eg. ethanol, has both O-H polar bonds and non-polar C-C bonds = dissolve in both polar and non-polar substances eg.
100
define hydrophilic
* polar and contain electronegative atoms (usually O2) that can interact with water
101
Define hydrophobic
* non-polar and comprised of a carbon chain
102
What gives water anomalous properties
Hydrogen bonds
103
Anomalous properties of water
Density
high mp/bp
surface tension - cohesion
104
What is anomalous about the density of water
ice is less dense than water
105
Why is ice less dense than water
* Hydrogen bond hold water molecules apart in open lattice structure
* Water molecules in ice are further apart than in water
* Solid ice is less dense than liquid water and floats
* With two lone pairs on oxygen atom and two hydrogen atoms, each water molecule can form 4 hydrogen bonds. The hydrogen bonds extend outward, holding water molecules slightly apart, forming an open tetrahedral lattice full of holes. Bond angle about hydrogen atom in hydrogen bon, is close to 180 degrees.
* Holes in open lattice structure decrease density of water on freezing, when ice melts the ice lattice collapses and molecules move closer together so liquid water is denser than ice.
106
Consequences of ice being less dense than water
* Forms an insulating layer in ponds and lakes, preventing water from freezing solid, so fish can survive.
107
Why does water have a relatively high mp/bp
* Water has London forces between molecules
* Hydrogen bonds are extra forces over and above London ones
* Fair amount of energy need to break hydrogen bonds in water, so water has much higher mp/bp than expectedWater has London forces between molecules
* Hydrogen bonds are extra forces over and above London ones
* Fair amount of energy need to break hydrogen bonds in water, so water has much higher mp/bp than expected
108
What’s the difference between ice melts and water boils
* When ice lattice breaks down, rigid arrangement of hydrogen bonds in ice is broken.
When water boils, hydrogen bonds break completely.
109
What reduces surface tension
Detergents
110
Does water have a high or low viscosity
High
111
How do you increase the strength of London forces
More electrons = more electrons are on any given side at one point = instantaneous dipole is larger
112
More branching
Weaker London forces….less points of contact
113
Trend in group 5 hydrides b.p
Decreases then it increases
114
draw a dot and cross for ozone
