Structure Flashcards
1
Q
What is electronegativity?
A
The power of an atom to attract the pair of electrons in a covalent bond towards itself
2
Q
What is the electron distribution in a covalent bond between elements with different electronegativities like?
A
Unsymmetrical
3
Q
Where does this phenomenon arise from? - (What is the electron distribution in a covalent bond between elements with different electronegativities like?)
A
From the positive nucleus’s ability to attract the negatively charged electrons, in the outer shells, towards itself
4
Q
What scale is used to assign a value of electronegativity to an atom?
A
The pauling scale
5
Q
Which is the most electronegative atom?
A
Fluorine with a value of 4.0 on the pauling scale
6
Q
When is fluorine best at attracting electron density towards itself?
A
When covalently bonded to another atom
7
Q
What is the nuclear charge?
A
Attraction exists between the positively charged protons in nucleus and negatively charged electrons in energy levels of atom
8
Q
What does an increase in the number of protons lead to?
A
- An increase in nuclear attraction for electrons in outer shells
- So, an increased nuclear charge results in an increased electronegativity
9
Q
What is the atomic radius?
A
The distance between the nucleus and electrons in the outermost shell
10
Q
What happens when electrons are closer to the nucleus in terms of atomic radius?
A
- They are more strongly attracted towards its positive nucleus and vice versa
- So an increased atomic radius results in a decreased electronegativity
11
Q
What effect do filled energy levels have on electronegativity?
A
- They shield the effect of the nuclear charge so outer electrons are less attracted to the nucleus and will experience less of the attractive force
- So it will result in a decreased electronegativity
12
Q
What is the trend of electronegativity down a group?
A
- It decreases
- Nuclear charge increases but there is increased shielding so atomic radii decrease in size
- So, there is a decrease in attraction between the nucleus and outer bonding electrons
13
Q
What is the trend of electronegativity across a period?
A
- It increases
- Nuclear charge increases and shielding remains constant so greater attraction between bonding electrons, resulting in smaller atomic radii
14
Q
What happens when 2 atoms in a covalent bond have the same electronegativity?
A
The covalent bond is non-polar as the electrons are shared equally between the 2 atoms
15
Q
What will dictate what type of polar bond is formed?
A
The difference in electronegativites
16
Q
When is an ionic bond formed?
A
When the values are very different (more than 1.7), ions will be formed
17
Q
When is a covalent bond formed?
A
- When the value is between 0.3 and 1.7. The bond will be polar
- The electrons will be drawn towards the more electronegative atom
18
Q
What happens when a covalent bond is polar?
A
- The negative charge centre and positive charge centre do not coincide with each other
- So electron distribution is asymmetric
- The less electronegative atom gets a partial charge of delta positive
- The more electronegative atom gets a partial charge of delta negative
19
Q
What does a greater difference in electronegativity mean?
A
The more polar the bond becomes
20
Q
How do we determine whether a molecule with more than 2 atoms is polar?
A
We need to consider:
- The polarity of each bond
- How the bonds are arranged in the molecule
21
Q
Why do some molecules that have polar bonds end up becoming non - polar?
A
The polar bonds in the molecule are arranged in such a way that the individual dipole moments cancel each other out
22
Q
What are intramolecular forces?
A
Forces within a molecule and are usually covalent bonds
(e.g. single, double, triple and co-ordinate bonds)
23
Q
What are intermolecular forces?
A
Forces between molecules
24
Q
What are the 3 types of intermolecular forces?
A
- Induced dipole-dipole forces (also known as van der walls or london dispersion forces)
- Permanent dipole-dipole forces
(are the attractive forces between 2 neighbouring molecules with a permanent dipole) - Hydrogen bonding (special type of permanent dipole-dipole force)
25
Which type of force is stronger?
Intramolecular forces
26
What is the order of the strength of the forces from strongest to weakest?
Covalent bonding
Hydrogen bonding
Permanent dipole-dipole
Instantaneous dipole - induced dipole
27
How do instantaneous dipole-induced dipole forces occur?
- The electron charge cloud in non-polar molecules or atoms are constantly moving
- During this movement, the electron charge cloud can be more on one side than the other
- This causes a temporary dipole to arise
- This temporary dipole can induce a dipole in neighbouring molecules
- The, the positive end of the dipole in 1 molecule and the negative end in another are attracted towards each other
28
Why are the dipoles in instantaneous dipole-induced dipole only temporary?
because the electron clouds are constantly moving
29
For small molecules with the same number of electrons, which force is stronger (permanent or induced)?
Permanent dipoles
e.g. Butane and propanone have the same number of electrons but butane has induced dipole and propanone has permanent dipole so has a higher b.p. as more energy is required to break the intermolecular forces
30
What kind of intermolecular forces will polar molecules have?
Permanent dipoles - it will always have a negatively and positively charged end
31
What is needed for hydrogen bonding to take place?
A species which has an O, N, or F atom bonded to a hydrogen
32
What happens when hydrogen is covalently bonded to an O, F or N?
- The bond becomes highly polarised
- The H becomes so delta positively charged that it can form a bond with the lone pair of an O, N or F atom in another molecule
33
Give an example of hydrogen bonding
Water can form 2 hydrogen bonds, because the O has 2 lone pairs
34
What are the examples of compounds that can form hydrogen bonds?
Alcohols
Ammonia
Amines
Carboxylic acids
Hydrogen fluoride
Proteins
35
Why does water have high melting and boiling points?
- Due to the strong intermolecular forces of hydrogen bonding between the molecules in both ice and water
- A lot of energy is therefore required to separate the water molecules and melt or boil them
36
What is the enthalpy of vaporisation?
The energy required to boil a substance
37
What does the high enthalpy change of vaporisation of water suggest?
Instantaneous dipole-induced dipole forces are not the only forces present in the molecule - there are also strong hydrogen bonds which cause the high b.p.
38
What kind of surface tension does water have?
A high one
39
What is surface tension?
The ability of a liquid surface to resist any external forces (i.e. to stay unaffected by forces acting on the surface)
40
What happens to water molecules at the surface of a liquid?
- The molecules are bonded to other water molecules through hydrogen bonds
- These molecules pull the surface molecules downwards, causing the surface of them to become compressed and more tightly together at the surface
- This increases water's surface tension
41
Why are solids denser than liquids?
The particles are more closely packed together
42
What are water molecules packed into as a solid?
A lattice
43
Why are the water molecules in a solid slightly further apart than in the liquid form?
This way of packing the molecules (into a lattice) and the relatively long bond lengths of the hydrogen bonds means the molecules are slightly further apart
44
Does ice have a lower density than liquid water?
Yes, by about 9%
The 'more open' structure of molecules in ice causes it to have a lower density than water
45
What does a larger surface area mean in a molecule?
It will have more contact with adjacent molecules
46
How is the surface area of a molecule reduced?
By branching
47
What does a greater ability to induce a dipole in an adjacent molecule result in?
Greater London forces and higher melting and boiling points
48
Give an example of how surface area affects the boiling point
Both pentane and 2-dimethylpropane have the same number of electrons but pentane has a larger surface area and so has a higher b.p. as more London forces are induced
49
What does a greater number of electrons or molecular mass in a molecule result in?
A greater likelihood of a distortion and thus the greater the frequency and magnitude of the temporary dipoles
50
How does the number of electrons affect the melting and boiling points?
The greater the number, the stronger the dispersion forces between molecules and also the enthalpy of vaporisation, making the melting and boiling points higher
51
What is the correlation between the number of electrons and energy needed to break forces of attraction?
As the number of electrons increases, more energy is needed to overcome the forces of attraction
52
Do alcohols have hydrogen bonds?
Yes, there are O-H bonds present and so hydrogen bonds set up between the slightly positive H atoms and lone pairs on oxygens in other molecules
53
Why are the H atoms slightly positive in an alcohol?
The bonding electrons are pulled away from them towards the very electronegative oxygen atoms
54
What are the only intermolecular forces in alkanes?
temporary induced dipole-dipole forces
55
Why is the b.p. of an alkane lower than its respective alcohol?
Hydrogen bonds in an alcohol are stronger than the induced dipole-dipole forces in an alkane and so it takes more energy to separate alcohol molecules than alkane molecules
56
What are the boiling points of the first 4 hydrogen halides?
HF - 293 K
HCl - 188 K
HBr - 207 K
HI - 238 K
57
What trend do the rest of the boiling points of the rest of the hydrogen halides follow?
They increase
58
Why do the boiling points of hydrogen halides increase as you go down the group?
- The molecules become larger
- The extra electrons allow greater temporary dipoles and so increase the amount of London dispersion forces between the molecules
59
Why does HF have a higher boiling point?
It has hydrogen bonding between the H-F molecules. F molecules have small intense lone pairs and H is very positive so bonds form
60
What is the general principle for solubility?
'like dissolves like'
61
What do non-polar substances dissolve in?
Non-polar substances such as hydrocarbons and they form dispersion forces between the solvent and the solute
62
What do polar covalent substances generally dissolve in?
Polar solvents as a result of dipole-dipole interactions or the formation of hydrogen bonds between the solute and the solvent
63
Give an example of when polar substances are dissolved
Between alcohols and water, hydrogen bonds form between the hydrogen on the water and the oxygen on the alcohol
64
What happens when covalent molecules become larger?
Their solubility can decrease as the polar part of the molecule is only a smaller part of the overall structure
(this effect is seen in alcohols e.g. C2H5OH is readily soluble whereas C6H13OH is not)
65
Why are polar covalent substances not able to dissolve well in non-polar solvents?
Their dipole-dipole interactions are unable to interact well with the solvent
66
Why do giant covalent substances generally not dissolve in any solvents?
The energy needed to overcome the strong covalent bonds in the lattice structures is too great
67
What can covalent compounds be arranged in?
Simple molecular lattices
Giant molecular
68
What do covalent substances tend to have?
Small molecular structures known as simple molecules e.g. Cl2, H2, CO2
69
What does iodine exist as?
It is a simple molecule but it exists as a crystalline structure involving a regular structure held together by weak london dispersion forces
70
Why do giant covalent lattices have very high melting and boiling points?
They have a large number of covalent bonds linking the whole structure
A lot of energy is required to break the lattice
71
Why is graphite soft?
Forces between the carbon layers are weak
72
Why is diamond and silicon oxide hard?
It is difficult to break their 3D network of strong covalent bonds
73
Are covalent compounds soluble in water?
No
74
Why does graphite conduct electricity?
It has delocalised electrons between the carbon layers which can move along the layers when a voltage is applied
75
Why can't diamond and silicon oxide conduct electricity?
All 4 outer electrons on every carbon atom are involved in a covalent bond so no freely moving electrons available
76
What is the structure of diamond called?
A tetrahedron
77
Give details about graphite?
- Forms layers of hexagons
- Free electrons migrate between the layers
- Layers can slide over each other
78
Are all covalent bonds in diamond identical?
Yes and there are no intermolecular forces
79
What kind of dimensional molecule is graphene?
It is essentially a 2D molecule as it is one atom thick
80
Why are the very unusual properties of graphene useful?
They make it useful in fabricating composite materials and in electronics
81
When are simple covalent molecules soluble?
When they are polar
82
Are metals soluble?
No, but some may react
