Physical - Bonding Flashcards
1
Q
What is ionic bonding?
A
The electrostatic attraction between oppositely charged ions in a lattice.
2
Q
What is an ionic bond?
A
The electrostatic attraction. The transfer of electrons has to take place first to make the ions that attract to each other, but ionic bonding doesn’t refer to this transfer.
3
Q
Why can a solution of an ionic compound conduct electricity?
A
Because the ions are free to move.
4
Q
Why can metals conduct electricity?
A
Because the delocalised electrons can move.
5
Q
What is charge density?
A
charge density = charge / size
6
Q
What atoms are involved in ionic bonding?
A
Metals and non-metals. The metal atoms lose electrons to form positive ions (cations) while the non-metal atoms gain electrons to form negative ions (anions).
7
Q
What is the strength of an ionic bond?
A
The smaller the ions and the greater the charge on the ions, the stronger the attraction between the positive and negative ions (usually). However, once distorted, if ions of the same charge are next to each other, they repel and the structure breaks.
8
Q
What is the structure of ionic compounds?
A
Giant, closely packed, ionic lattice structure of positive and negatively charged ions. The ions are held together by the very strong electrostatic attraction between the + and - ions. This +/- attraction is known as an ionic bond, though it is just an electrostatic attractive force.
9
Q
What is the melting and boiling points of ionic compounds?
A
High because a lot of energy is required to overcome the strong attraction. They are solids at room temperature.
10
Q
What is the conductivity of ionic compounds?
A
Can conduct when liquid (molten) and aqueous only because that’s when the ions are free to move. As a solid, they are fixed in place so are unable to move and carry the electric current.
11
Q
Are ionic compounds soluble in water?
A
Generally soluble and dissolve in water.
12
Q
What is covalent bonding?
A
The shared pair of electrons between atoms.
13
Q
What is the strength of a covalent bond?
A
The shorter the bond, the stronger the bond (usually).
Double bonds are stronger than single bonds, while triple bonds are stronger than double bonds.
14
Q
What atoms are involved in covalent bonding?
A
Non-metals. The atoms share electrons to obtain stable electron structures.
15
Q
What structures use covalent bonding?
A
Simple molecular and giant covalent (macromolecular).
16
Q
What are polymers?
A
Molecules made up of long chains of covalently bonded carbon atoms. They’re formed when lots of small molecules called monomers join together. Monomers are weaker forces and so have a lower melting and boiling point.
The more covalent bonds there are, the higher the melting point because the stronger they are and the longer it will take to break the bonds. This is because the longer chains, the more energy is required to break the bonds and cause the material to change state.
17
Q
What is a lone pair?
A
Two non-bonded outer shell electrons.
18
Q
What is a dative covalent bond? (co-ordinate)
A
When both the electrons come from the same species. The atom that accepts the electron pair is an atom that does not have a filled outer main level of electrons - the atom is electron-deficient. The atom that is donating the electrons has a pair of electrons that is not being used in a bond (a lone pair).
Coordinate bonds have exactly the same strength and length as ordinary covalent bonds between the same pair of atoms, but they are charged particles.
19
Q
What does the arrow show in dative covalent bonds?
A
Shows the direction in which the electrons are donated. Once formed, dative covalent bonds are identical to other covalent bonds.
20
Q
What is the structure of simple molecular elements/compounds?
A
Individual molecules with weak intermolecular (van der waals) forces between them. Atoms within molecules are joined by strong covalent bonds.
21
Q
What is the melting and boiling points of simple molecular elements/compounds?
A
Low because only the weak intermolecular forces need to be broken in order to change state. Breaking the strong covalent bonds (intramolecular forces) would be a chemical rather than physical change. Most molecular substances are gases or liquids at room temperature.
22
Q
What is the conductivity of simple molecular elements/compounds?
A
Cannot conduct because there are no charged particles that can move - simple molecular elements/compounds are neutral.
23
Q
Are simple molecular elements/compounds soluble in water?
A
Infrequently soluble, but usually insoluble.
24
Q
What is the structure of giant covalent elements/compounds?
A
Giant, continuous, lattice structures in which all atoms are joined to others by covalent bonds in a tetrahedral arrangement.
25
What is the melting and boiling points of giant covalent elements/compounds?
Very high due to the large amounts of energy that is required to overcome the strong covalent bonds. All are solids at room temperature.
26
What is the conductivity of giant covalent elements/compounds?
Cannot conduct because there are no charged particles to carry the current (except graphite and graphene due to the delocalised electron).
27
Are giant covalent elements/compounds soluble in water?
Insoluble as water molecules cannot break down the covalent bonds to pull the carbon atoms apart.
28
What is metallic bonding?
The electrostatic attraction between the positive metal ion and the delocalised electrons (from the outer shell of the metal atom).
29
What is the strength of a metallic bond?
The stronger the attraction between the positive metal ions and the delocalised electrons, the stronger the metallic bonding. The smaller the metal ions, the stronger the electrostatic attraction between the positive ions and electrons, the closer the electrons are to the positive nucleus, the greater the charge on the ions (more protons) and the more delocalised outer shell electrons there are, the stronger the metallic bonding. Metals also tend to be strong because the delocalised electrons extend throughout the solid so there are no individual bonds to break.
30
What is the structure of metallic elements?
Giant, fixed lattice structure of metal ions with outer shell electrons free to move through the structure. There are strong forces of electrostatic attraction between the positive metal ions and the shared negative electrons. These forces of attraction hold the atoms together in a regular, fixed lattic structure. Include metals and alloys.
31
What is the melting and boiling points of metallic elements?
High because of the strong electrostatic forces of attraction, and their giant structures. However this can vary depending on the metal. All but mercury are solid at room temperature. Alloys have lower melting and boiling points because of the different shaped atoms, making the forces between them weaker.
32
What is the conductivity of metallic elements?
The delocalised electrons that can move throughout the structure explain why metals are such good conductors of electricity as a solid or liquid.
Metals are also good conductors of heat - they have high thermal conductivity. The sea of electrons is partly responsible for this property, with energy also spread by increasingly vigorous vibrations of the closely packed ions.
33
Are metallic elements soluble in water?
Insoluble as the polarity of the water molecules is not large enough to overcome the strong metallic bonding. However, some react with water.
34
What are monatomic substances?
Group 0 elements (noble gases).
35
What is the structure of monatomic substances?
Individual atoms with very weak forces between them.
36
What is the melting and boiling points of monatomic substances?
Very low as a result of the weak attractions between the atoms. All are gases at room temperature.
37
What is the conductivity of monatomic substances?
Cannot conduct because there are no charged particles that can move. The atoms are neutral.
38
Are monatomic substances soluble in water?
Insoluble
39
Where can dative covalent bonds be found? In what structures?
Between non-metal atoms, in simple molecular structures only. This means they have weak intermolecular forces between the molecules, and a low melting/boiling point.
40
What are the melting and boiling points for graphite?
High because there are strong covalent bonds between the majority of carbon atoms in the structure. These take a lot of energy to break, making the melting and boiling points very high.
41
What is the electrical conductivity of graphite?
Conducts electricity as each carbon atom only forms covalent bonds to three other carbon atoms. The fourth outer shell electron from each carbon atom is said to be delocalised and is free to move along the planes of the structure and allow it to conduct electricity.
42
What is the strength of graphite?
It is a very soft material. The weak Van der Waals forces between the hexagonal layers of carbon atoms are easy to break, allowing the layers to slide over each other.
43
What is the solubility of graphite?
Does not dissolve in water and water molecules cannot beak down the covalent bonds to pull the carbon atoms apart.
44
What is the structure of graphene?
Graphene is a single-atom thick layer of graphite with strong covalent bonds between each carbon atom. The atoms are arranged in hexagons.
45
What are the melting and boiling points for graphene?
High as graphene's many covalent bonds are strong and substantial energy is needed to break them.
46
What is the electrical conductivity of graphene?
High as each carbon atom has an unbonded electron. The unbonded electrons are delocalised electrons that are free to move and carry charge.
47
What is the strength of graphene?
Very strong. Graphene's strong covalent bonds makes it 100 times stronger than steel. It is also the thinnest material possible - one atom thick - and very lightweight and transparent.
48
What is the strength of simple molecular substances?
They are soft and break easily. Van der Waals forces between the molecules are much weaker than covalent bonds.
49
What is the melting and boiling points of ice?
Higher than would be expected. Hydrogen bonds between molecules are the strongest type of intermolecular force and therefore more energy is needed than expected to break these to let the molecules move apart, compared to other simple molecular substances.
50
What is the electrical conductivity of ice?
Low because there are no charged particles to carry the current.
51
What is the strength of ice?
The hydrogen bonds between the layers are quite hard to break and the arrangement of the molecules in the solid is like the structure of diamond, so ice is strong.
52
What is absolute zero?
-273 degrees centrigrade = 0 Kelvin
53
What particles possess kinetic energy?
The particles in any substance at any temperature above absolute zero are either vibrating about a fixed position (solid) or moving around (liquid or gas). Therefore they possess kinetic energy. The particles have different amounts of kinetic energy (some are moving faster than others). At absolute zero, the particles do not vibrate and so the particles have no kinetic energy.
54
What is the relationship between temperature and kinetic energy?
The temperature of any substance is directly proportional to the mean kinetic energy of the particles.
55
What happens as a solid is being heated?
As the solid is heated, the particles vibrate faster which increases the mean kinetic energy of the particles. As the mean kinetic energy of the particles increases, temperature increases.
56
What happens as a solid is melting?
The heat energy is used to partially overcome the forces (or bonds) between the particles rather than to increase the kinetic energy of the particles. The solid melts but the mean kinetic energy of the particles remains constant and so temperature remains constant.
57
What happens as a liquid is being heated?
As the liquid is heated, the particles move faster, increasing the mean kinetic energy of the particles. As the mean kinetic energy of the particles increases, temperature increases.
58
What happens as a liquid is boiling?
The heat energy is used to overcome the forces (or bonds) between the particles rather than to increase the kinetic energy of the particles. The liquid boils but the mean kinetic energy of the particles remains constant and so temperature remains constant.
59
What happens as a gas is being heated?
As the gas is being heated, the particles move faster, increasing the mean kinetic energy of the particles. As the mean kinetic energy of the particles increases, temperature increases.
60
What is latent heat?
The energy required to change state without any change in temperature.
61
What happens when a gas is cooled to a liquid and then a solid?
During condensation and solidifying/freezing, energy from forming forces (or bonds) is released as heat energy. This stops the temperature from falling further as the mean kinetic energy of the particles remains constant as it changes state.
62
What happens when a liquid evaporates? (different from boiling)
When a liquid evaporates (which occurs below the boiling point), some kinetic energy is used to overcome forces between particles to allow the particle to escape. As a consequence, the mean kinetic energy of the remaining particles is lower and so the temperature is lower. This explains why liquids cool as they evaporate.
63
What do the shapes of molecules and ions depend on?
- The total number of electron pairs around the central atom which repel each other as far as possible.
- The nature of these pairs (bonding pair or lone pair).
64
What two types of electron pairs do molecules and ions possess? Do they attract or repel?
Bonding pairs
- The two shared electrons in a covalent bond.
- These pairs repel each other equally.
Lone pairs
- The two electrons in a pair not involved in bonding; also known as non-bonding pairs.
- Lone pairs repel other pairs more than bonding pairs because they are more electron dense.
- Each lone pair reduces the bond angle by about 2.5 degrees.
65
What does the strength of the repulsion depend on?
The type of electron pair interaction.
66
What is the strongest repulsion between?
lone pair to lone pair
67
What is the middle repulsion between?
lone pair to bond pair
68
What is the weakest repulsion between?
bond pair to bond pair
69
What dimension, name of shape, and bond angle do 2 electron pairs have?
2 PAIRS
- dimension = 2D
- name of shape = linear
- bond angle = 180 degrees
70
What dimension, name of shape, and bond angle do 3 electron pairs have?
3 PAIRS
- dimension = 2D
- name of shape = trigonal planar
- bond angle = 120 degrees
71
What dimension, name of shape, and bond angle do 4 electron pairs have?
4 PAIRS
- dimension = 3D
- name of shape = tetrahedral
- bond angle = 109.5 degrees
72
What dimension, name of shape, and bond angle do 5 electron pairs have?
5 PAIRS
- dimension = 3D
- name of shape = trigonal bipyramidal
- bond angle = 120 degrees, 90 degrees
73
What dimension, name of shape, and bond angle do 6 electron pairs have?
6 PAIRS
- dimension = 3D
- name of shape = octahedral
- bond angle = 90 degrees
74
Are lone pairs or bonding pairs more compact?
Lone pairs are more compact than bonding pairs because they aren't bonding to anything. This means that lone pairs repel more than bonding pairs. This reduces the bond angles to a small extent.
75
How do you work out the shapes for molecules and ions with single bonds only?
- work out how many outer shell electrons are on the central atom (accounting for any charge on the species)
- work out the number of atoms bonded to the central atom (= number of bonding pairs)
- work out how many lone pairs there are
- add the bonding pairs and lone pairs
76
What shape do you call a molecule with 2 electron pairs with 2 bonding pairs but no lone pairs? What is the angle?
- linear
| - 180 degrees
77
What shape do you call a molecule with 3 electron pairs with 3 bonding pairs but no lone pairs? What is the angle?
- trigonal planar
| - 120 degrees
78
What shape do you call a molecule with 3 electron pairs with 2 bonding pairs and 1 lone pair? What is the angle?
- bent (v-shape)
| - 118 degrees
79
What shape do you call a molecule with 4 electron pairs with 4 bonding pairs but no lone pairs? What is the angle?
- tetrahedral
| - 109.5 degrees
80
What shape do you call a molecule with 4 electron pairs with 3 bonding pairs and 1 lone pair? What is the angle?
- trigonal pyramidal
| - 107 degrees
81
What shape do you call a molecule with 4 electron pairs with 2 bonding pairs and 2 lone pairs? What is the angle?
- bent (v-shape)
| - 104.5 degrees
82
What shape do you call a molecule with 5 electron pairs with 5 bonding pairs but no lone pairs? What are the angles?
- trigonal bipyramidal
| - 120 degrees, 90 degrees
83
What shape do you call a molecule with 5 electron pairs with 4 bonding pairs and 1 lone pair? What are the angles?
- trigonal pyramidal
| - 119 degrees, 89 degrees
84
What shape do you call a molecule with 5 electron pairs with 3 bonding pairs and 2 lone pairs? What are the angles?
- trigonal planar
- 120 degrees, 90 degrees OR 89 degrees
(top and bottom provide equal repulsion)
85
What shape do you call a molecule with 6 electron pairs with 6 bonding pairs but no lone pairs? What is the angle?
- octahedral
| - 90 degrees
86
What shape do you call a molecule with 6 electron pairs with 5 bonding pairs and 1 lone pair? What is the angle?
- square pyramid
| - 89 degrees
87
What shape do you call a molecule with 6 electron pairs with 4 bonding pairs and 2 lone pairs? What is the angle?
- square planar
- 90 degrees
(top and bottom provide equal repulsion?
88
What is the average bond enthalpy?
The energy required to break a covalent bond. The bigger the value, the stronger the bond, the higher the melting and boiling point.
89
What are allotropes?
Different forms of the same element.
90
Why do smaller atoms create a stronger attraction?
Atomic radius is smaller, so there's a stronger attractive force between the negatively charged electrons and positively charged nucleus.
91
Why are ionic substances brittle?
Ionic substances are often brittle materials. This is because they form a lattice of alternating positive and negative ions. When the layers of alternating charges are distorted, like charges repel, breaking apart the lattice into fragments.
92
Why are pure metals malleable?
The layers of atoms that are the same size can easily slide past each other.
93
For every lone pair, how much does the bond angle change by?
2-2.5 degrees
94
Give examples of compounds with a total number of 2 electron pairs and a linear shape.
BeCl2
| CO2
95
Give examples of compounds with a total number of 3 electron pairs and a trigonal planar shape.
BF3
NO3 -
CO3 2-
96
Give examples of compounds with a total number of 3 electron pairs and a bent (V-shape) shape.
SO2
97
Give examples of compounds with a total number of 4 electron pairs and a tetrahedral shape.
```
CH4
SO4 2-
AlCl4 -
NH4 +
BeCl4 2-
SiCl4
```
98
Give examples of compounds with a total number of 4 electron pairs and a trigonal pyramidal shape.
NH3
H3O +
PF3
NF3
99
Give examples of compounds with a total number of 4 electron pairs and a bent (V-shape) shape.
H20
100
Give examples of compounds with a total number of 5 electron pairs and a trigonal bipyramidal shape.
PF5
101
Give examples of compounds with a total number of 5 electron pairs and a trigonal pyramidal or see-saw shape.
SF4
102
Give examples of compounds with a total number of 5 electron pairs and a trigonal planar or T-shape shape.
ICl3
103
Give examples of compounds with a total number of 6 electron pairs and a octahedral shape.
SF6
104
Give examples of compounds with a total number of 6 electron pairs and a square pyramid shape.
IF5
105
Give examples of compounds with a total number of 6 electron pairs and a square planar shape.
XeF4
| XeCl4
106
What is electronegativity?
The power of a bonded atom to attract the electron density (pair of electrons) in a covalent bond towards itself. The more electronegative element in a covalent bond attracts the bonding electrons towards itself.
107
What are the factors the affect electronegativity?
1. nuclear charge
2. atomic radius (distance between the nucleus and the outer shell electrons)
3. shielding of the nuclear charge by electrons in inner shells
108
How does nuclear charge affect electronegativity?
The more protons there are, the stronger the attraction between the nucleus and the bonding pairs of electrons, and the greater the electronegativity.
109
How does atomic radius affect electronegativity?
The smaller the atomic radius, the stronger the attraction between the nucleus and the bonding pairs of electrons, and the greater the electronegativity.
110
How does shielding affect electronegativity?
The fewer shells of electrons between the nucleus and the electrons there are, the less shielding (repulsion) there is and the stronger the attraction between the nucleus and the bonding pairs of electrons.
111
Describe the trend in electronegativity down a group and across a period.
down a group - electronegativity decreases
| across a period - electronegativity increases
112
Why does electronegativity decrease down a group?
- atomic radius increases
- more shielding
- less attraction between the nucleus and the bonding pair of electrons
113
Why does electronegativity increase across a period?
- atomic radius decreases
- higher nuclear charge
- similar shielding
- stronger attraction between the nucleus and the bonding pair of electrons
114
When does a non-polar covalent bond form?
When the two atoms in a covalent bond have the same electronegativity.
- If the two bonding atoms are identical, their attraction for the shared pair of electrons is symmetrical.
- The electrons are equally distributed between the bonding atoms.
- This bond is perfectly covalent.
115
What does a non-polar covalent bond mean?
It's a covalent bond where the two electrons are shared equally.
116
Give an example of a non-polar covalent bond.
Cl-Cl bond in Cl2
117
When does a polar covalent bond form?
When the two atoms in a covalent bond have a different electronegativity.
- If the two bonding atoms are different, their attraction for the shared pair of electrons is unsymmetrical.
- The bonding atom with a greater attraction for the shared pair of electrons is more electronegative.
- The bond is polarised.
118
What does a polar covalent bond mean?
It's a covalent bond where the two electrons are not shared equally. The more electronegative atom has a greater share of the two electrons and is beta negative while the less electronegtive atom has a lower share and is beta positive.
119
Give an example of a polar covalent bond.
HCl bond in HCl (hydrogen is beta positive and chlorine is beta negative)
120
What is the strongest to weakest intermolecular force?
1. Hydrogen bonding
2. Permanent dipole-dipole bonding
3. van der Waals' forces (aka induced dipole-dipole or London dispersion forces)
121
What are van der Waals' forces?
- These are present in all molecular substances.
- All atoms and molecules are made up of positive and negative charges even though they are neutral overall. These charges produce very weak electrostatic attractions between all atoms and molecules.
- They occur because the electrons are constantly moving around and there will be an uneven (unsymmetrical) electron distribution at any given moment in time. This causes a temporary dipole within a molecule.
- This temporary dipole induces a temporary dipole in a neighboring molecule. There is then an attraction between these molecules - this is a temporary induced dipole-dipole attraction.
- The bigger the molecule (i.e. the more electrons), the greater the van der Waal's forces because the dipole is caused by the changing position of the electron cloud.
- Breaking van der Waal's forces does not mean breaking covalent bonds, as this would be a chemical rather than physical change.
122
What is permanent dipole-dipole attraction?
- There are permanent dipole-dipole attractions between polar molecules (e.g. between H-Cl molecules).
- There are only permanent dipole-dipole attractions between polar molecules, if correctly aligned. Some molecules are non-polar but contain polar bonds (e.g. CCl4 and CO2) - these do not have permanent dipole-dipole attractions.
- This attraction is not induced and is permanent as the polarity is always there.
- Two molecules which both have dipoles will attract one another. Whatever their starting positions, the molecules with dipoles will 'flip' to give an arrangement where the two molecules attract.
123
What is hydrogen bonding?
- This is a special case of permanent dipole-dipole attractions - where a hydrogen atom is bonded to a very electronegative atom (i.e. F, O, N).
- Common examples of molecules where they occur are HF, H2O, NH3, alcohols, carboxylic acids, amines, amino acids.
- The polar bond between the H and N/O/F leaves the positive H nucleus exposed as H only has one electron to shield it.
- Therefore there is a strong attraction from the lone pair on the N/O/F of one molecule to the exposed H nucleus of another molecule.
- This is simply a strong intermolecular force - it is not a bond.
- When drawing the hydrogen bonds between two molecules, always show all lone pairs, all beta positive and beta negative charges, and a dotted line between the lone pair on one molecule and the beta positive hydrogen on another.
124
Compare covalent bonds to the forces between molecules.
Covalent bonds are very strong (values in hundred of kJ mol-1). The forces between molecules are much weaker, with van der Waal's forces being in units of kJ mol-1 and hydrogen bonds in tens of kJ mol-1.
125
Bonds that are polar have a bond dipole moment. What does this mean?
This is a measure of the strength and direction of the polarity in the bond. In simple terms, the bigger the difference in electronegativity, the bigger the bond dipole moment.
126
What is an induced dipole?
An induced dipole can form when the electron orbitals around a molecule are influenced by another charged particle.
127
What do London forces occur between?
induced dipole-induced dipole
128
What do permanent-induced dipoles occur between? (dipole-dipole interactions)
permanent dipole-induced dipole
129
What do permanent dipoles occur between? (dipole-dipole interactions)
permanent dipole-permanent dipole
130
What is the Pauling scale?
The Pauling scale is used as a measure of electronegativity. It runs from 0 to 4. The greater the number, the more electronegative the atom. The noble gases have no number because they do not, in general, form covalent bonds.
131
What is electron density?
The electron density relates to the probability of finding electrons at a particular position in space. It can be imagined as a cloud of electrons around the nucleus.
132
What is the spectrum of bonds?
Rather than bonds existing as discretely ionic and covalent, they exist on a spectrum.
Ionic Bonding
- The difference in electronegativity is so great that one atom effectively takes the electron from the other.
Polar-Covalent Bonding
- The difference in electronegativity is small.
- The atoms share the electrons unequally.
- The bond is polarised.
Covalent Bonding
- There is no difference in electronegativity.
- The molecule is electronically symmetrical.
133
Are molecules containing polar bonds always polar?
Molecules containing polar bonds are not always polar. The symmetry of polar bonds can cancel the effect of any permanent dipole.
134
What is a non-symmetrical molecule that contains polar bonds?
- A difference in charge exists across the molecule.
- Electrons are pulled towards the more electronegative atom.
- There is an overall dipole.
- The molecule is polar.
- E.g. water (non-linear shape).
135
What is a symmetrical molecule that contains polar bonds?
- The symmetry of the molecule means that the effect of any permanent dipoles is cancelled out.
- Have a linear, trigonal planar or tetrahedral shape.
- All atoms attached to the central atom are identical.
- No difference in charge exists across the molecule.
- The molecule is non-polar, despite containing polar bonds.
- E.g. carbon dioxide (linear shape).
136
What lines do we use when drawing 3D shapes of molecules?
- Normal Line -> Bond in the plane of the paper.
- Dotted Wedge -> Bond is going into the paper away from you.
- Bold Wedge -> Bond is coming out of the paper towards you.
137
What is an octet?
An octet is 8 electrons on the outer shell with 4 of them used in covalent bonding.
138
Some elements can expand their octet. Where can this occur?
Group 15-17, from Period 3 downwards.
- Group 15, Non-Metals (3/5 covalent bonds)
- Group 16, Non-Metals (2/4/6 covalent bonds)
- Group 17, Non-Metals (1/3/5/7 covalent bonds)
139
How does sharing electrons hold atoms together?
Atoms with covalent bonds are held together by the electrostatic attraction between the nuclei and the shared electrons. This takes place within the molecule. The electrostatic forces balance when the nuclei are a particular distance apart.
140
What forms when some covalent compounds react with water?
ions
141
What are delocalised electrons?
The atoms in a metal element cannot transfer electrons (as happens in ionic bonding) unless there is a non-metal atom present to receive them. In a metal element, the outer main levels of the atoms merge. The outer electrons are no longer associated with any one particular atom. Delocalised electrons are not tied to a particular atom.
142
Why don't the positive ions in a metal repel each other?
The positive ions tend to repel one another, but this is balanced by the electrostatic attraction of these positive ions for the negatively charged 'sea' of delocalised electrons.
143
What does the number of delocalised electrons depend on?
How many electrons have been lost by each metal atom.
144
Why do metals have giant structures?
Because the metallic bonding spreads throughout.
145
Why are metals malleable and ductile?
Metals are malleable (they can be beaten into shape) and ductile (they can be pulled into thin wire) as the metal ions can slide over each other and are still held together by delocalised electrons. After a small distortion, each metal ion is still in exactly the same environment as before so the new shape is retained.
146
Where are the most electronegative atoms are found?
The most electronegative atoms are found at the top right-hand corner of the Periodic Table (ignoring the noble gases which form few compounds). The most electronegative atoms are fluorine, oxygen, and nitrogen followed by chlorine.
147
Why could you say that polar covalent bonds have some ionic character?
It is going some way towards the separation of the atoms into charged ions. It is also possible to have ionic bonds with some covalent character.
148
What do van der Waals forces explain?
- why the boiling points of the noble gases increase as the atomic numbers of the noble gases increase
- why the boiling points of hydrocarbons increase with increased chain length
149
What conditions have to be present for hydrogen bonding to occur?
- A hydrogen atom that is bonded to a very electronegative atom. This will produce a strong partial positive charge on the hydrogen atom.
- A very electronegative atom with a lone pair of electrons. These will be attracted to the partially charged hydrogen atom in another molecule and form the bond.
- The only atoms that are electronegative enough to form hydrogen bonds are oxygen, nitrogen and fluorine.
150
What is the importance of hydrogen bonding?
Although hydrogen bonds are only about 10% of the strength of covalent bonds, their effect can be significant - especially when there are a lot of them. The very fact that they are weaker than covalent bonds, and can break or make under conditions where covalent bonds are unaffected, is very significant.
151
Why is the nitrogen-hydrogen-oxygen system linear?
Because the pair of electrons in the N-H covalent bond repels those in the hydrogen bond between oxygen and hydrogen. This linearity is always the case with hydrogen bonds.
152
Describe the structure of ice.
In water in its liquid state, the hydrogen bonds break and reform easily as the molecules are moving about. When water freezes, the water molecules are no longer free to move about and the hydrogen bonds hold the molecules in fixed positions. The resulting three-dimensional structure resembles the structure of diamond.
153
Describe the density of ice.
In order to fit into the three-dimensional structure, the molecules in ice are slightly less closely packed than in liquid water. This means that ice is less dense than water and forms on top of ponds rather than at the bottom.
154
What are the benefits of ice forming on top of water?
This insulates the ponds and enables fish to survive through the winter. This must have helped life to continue, in the relative warmth of the water under the ice, during the Ice Ages.
155
What is the electron pair repulsion theory?
Each pair of electrons around an atom will repel all other electron pairs. The pairs of electrons will therefore take up positions as far apart as possible to minimise repulsion.
156
What is enthalpy?
Enthalpy is the heat energy change measured under constant pressure whilst temperature depends on the average kinetic energy of the particles and is therefore related to their speed - the greater the energy, the faster they go.
157
What is the enthalpy change of melting?
The energy needed to weaken the forces that act between the particles, holding them in the solid state.
158
What is the enthalpy change of vaporisation?
The energy needed to break all the intermolecular forces between the particles in a liquid.
159
How does the electron pair repulsion theory affect the structure of diamond?
The four electron pairs around the carbon atoms repel each other. In three dimensions, the bonds actually point to the corners of a tetrahedron.
160
Which direction does graphite conduct electricity?
The delocalised electrons can travel freely through the material, though graphite will only conduct along the hexagonal planes, not at right angles to them.
161
Why are intermolecular forces stronger if the molecule is bigger?
Because there are more electrons (we're talking about induced dipole-dipole attractions). This isn't covalent or ionic bonding, where the bigger the atom, the weaker the attraction.
162
What kind of chained molecules experience stronger van der Waals forces?
Straight chain molecules experience stronger van der Waals forces than branched chain molecules as they can line up and pack closer together. This reduces the distance over which the force acts, therefore they are stronger.
163
Name compounds with coordinate covalent bonds.
NH4 +
SO3
H3O +
164
How does van der Waal's forces change down group 7?
increases because more electrons, so more likely that there'll be an uneven distribution
165
If molecules only have van der Waal's forces, which one will have a higher melting or boiling point?
the one with more electrons
166
If molecules have different intermolecular forces, which one will have a higher melting or boiling point?
hydrogen bonding, then permanent dipole dipole forces, then van der Waal's forces
167
If molecules only have dipole dipole forces, which one will have a higher melting or boiling point?
the one with the greater difference in electronegativity
168
Why do polar molecules dissolve in polar substances?
Because hydrogen bonds or dipole dipole bonds will break between the molecules and form new ones with water. If the strongest force is van der Waal's, then it can only form van der Waal's forces with water, which aren't very strong and therefore nonpolar molecules aren't soluble in water.
