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AQA A Level Physical Chemistry > 3. Bonding > Flashcards

3. Bonding Flashcards

1
Q

How do ions in ionic compounds achieve stability?

A

Their electron configuration is the same as a noble gas after they have gained/lost an electron

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2
Q

What is an ionic bond?

A

The electrostatic force of attraction between the two oppositely charged ions

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3
Q

Will ionic bonds exist in isolation?

A

No, they will always form part of a giant ionic lattice

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4
Q

When two ions are more charged, what will this mean for the forces of attraction between them?

A

The more charged the ions are, the greater force of attraction between them

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5
Q

Are ionic compounds ever molecules?

A

No

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6
Q

What does the formula ‘NaCl’ represent?

A

That there is a one to one ratio of sodium to chloride ions

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7
Q

What is the formula for an ammonium ion?

A

NH4+

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8
Q

What is the formula for a chlorate ion?

A

ClO3-

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9
Q

What is the formula for a cyanide ion?

A

CN-

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10
Q

What is the formula for a carbonate ion?

A

CO3 2-

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11
Q

What is the formula for a chromate ion?

A

CrO4 2-

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12
Q

What is the formula for a dichromate ion?

A

Cr2O7 2-

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13
Q

What is the formula for a hydrogen carbonate ion?

A

HCO3 -

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14
Q

What is the formula for a hydrogen phosphate ion?

A

HPO4 2-

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15
Q

What is the formula for a hydrogen sulfate ion?

A

HSO4 -

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16
Q

What is the formula for a hydroxide ion?

A

OH-

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17
Q

What is the formula for a nitrate ion?

A

NO3 -

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18
Q

What is the formula for a nitrite ion?

A

NO2-

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19
Q

What is the formula for a manganate ion?

A

MnO4-

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20
Q

What is the formula for a peroxide ion?

A

O2 2-

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21
Q

What is the formula for a phosphate ion?

A

PO4 3-

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22
Q

What is the formula for a sulfate ion?

A

SO4 2-

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23
Q

What is the formula for a sulfite ion?

A

SO3 2-

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24
Q

What is the formula for a thiosulfate ion?

A

S2O3 2-

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25
Q

What is a covalent bond?

A

A shared pair of electrons, where each electron come from a different atom

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26
Q

What happens when a pair of electrons is shared?

A

A single bond is formed

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27
Q

How is a double or triple covalent bond formed?

A

If two or three pairs of electrons are shared

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28
Q

What type of covalent bond does a H2 molecule have?

A

Single covalent, H-H

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29
Q

What type of covalent bond does an O2 molecule have?

A

Double covalent, O=O

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30
Q

What type of covalent bond does an N2 molecule have?

A

Triple covalent, N≡N

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31
Q

What is a coordinate bond?

A

When both electrons come from the same atom

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32
Q

What is a coordinate bond also known as?

A

A dative covalent bond

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33
Q

What are ammonium ions formed from?

A

NH3 and H+

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34
Q

How is an ammonium ion formed?

A
  • NH3 molecule has a lone pair of electrons, and H+ ion has no electrons
  • the lone pair in NH3 donates electrons to vacant orbital in the hydrogen (COORDINATE BOND FORMED)
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35
Q

How is a coordinate bond represented in displayed formula?

A

As an arrow, showing where electrons have been donated

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36
Q

What are the types of covalent bonding?

A
  • simple molecular

* giant covalent (macromolecular)

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37
Q

Where does metallic bonding exist?

A

Within a metal, as there is no bonding between metals

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38
Q

Why do electrons in metallic bonding become delocalised?

A

There are no non-metal atoms to transfer electrons to

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39
Q

What does the number of delocalised electrons in metallic bonding depend on?

A

The number of outer electrons the metal atom has

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40
Q

What happens in metallic bonding when there are more delocalised electrons?

A

Increased attracted between ions and electrons as ions more strongly charged

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41
Q

What is a metallic bond?

A

In a metal, when positive metal ions and delocalised electrons attract each other and hold the structure together

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42
Q

Do metallic bonds exist in isolation?

A

No - they form part of a giant metallic lattice

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43
Q

What is the strength of ionic, covalent and metallic bonds?

A

Very strong

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44
Q

What are the four types of crystal structure?

A
  • ionic
  • metallic
  • macromolecular (giant covalent)
  • (simple) molecular
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45
Q

Why can ionic compounds dissolve in water?

A

Positive ions in lattice are attracted to opposite charged atom in water and vice versa (as water is a polar molecule)

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46
Q

Why can ionic compounds conduct electricity and heat when molten or in aqueous solutions?

A

Positive and negative ions are free to move and therefore carry charge

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47
Q

Why do ionic compounds have high melting points?

A

Lots of electrostatic forces of attraction between positive and negatively charged ions, which take a lot of energy to break

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48
Q

Why are ionic compounds brittle?

A

If it is bent then layers slide over each other so that like charges are next to each other, then repel and break apart

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49
Q

Why don’t simple covalent compounds conduct electricity or heat?

A

There are no free charged particles to carry charge - only ends of molecules are slightly charged so they are neutral

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50
Q

Why do simple covalent compounds have low melting points?

A

There are weak intermolecular forces between molecules which need little energy to break

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51
Q

Why don’t giant covalent compounds conduct electricity? Is there an exception to this?

A

No charged particles that are free to move

graphite conducts - 1 free electron per atom

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52
Q

Why do giant covalent compounds have high melting points?

A

Lots of strong covalent bonds which take a lot of energy to overcome

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53
Q

Why are giant covalent compounds not soluble in water?

A

No charged particles - all carbon are neutral so don’t attract to H2O molecules

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54
Q

Why do metals conduct electricity and heat?

A

Have delocalised electrons that can carry a charge

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55
Q

Why do metals have high melting points?

A

Electrostatic forces - ‘metallic bonds’ - need a lot of energy to be broken

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56
Q

Why are metals malleable and ductile?

A

Layers of ions can slide past each other easily - bonds flexible although strong

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57
Q

What substances are typical of monatomic structures?

A

Elements: group 0

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58
Q

What substances are typical of simple molecular structures?

A
  • Elements: H₂, O₂, N₂, F₂, Cl₂, Br₂, I₂, At₂, S₈, P₄

* Compounds: different non-metals bonded

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59
Q

What substances are typical of giant covalent structures?

A
  • Elements: Silicon, carbon (diamond and graphite)

* Compounds: SiO₂

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60
Q

What is the structure in monatomic substances?

A

Individual atoms with very weak forces between them

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61
Q

What substances are typical of ionic structures?

A

Compounds: metal and non-metal bonded

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62
Q

What substances are typical of metallic structures?

A

Elements: one metal

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63
Q

What is the structure in simple molecular substances?

A

Individual molecules with weak forces between them

atoms within molecules are joined by covalent bonds

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64
Q

What is the structure in giant covalent substances?

A

Lattice structure in which all atoms are joined to other atoms by covalent bonds

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65
Q

What is the structure in ionic substances?

A

Lattice structure of +ve and -vely charged ions

held together by ionic bonds

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66
Q

What is the structure in metallic substances?

A

Lattice structure of metal ions with outer shell electrons free to move through the structure

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67
Q

What are the particles in monatomic substances?

A

Atoms

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68
Q

What are the particles in simple molecular substances?

A

Molecules

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69
Q

What are the particles in giant covalent substances?

A

Atoms

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70
Q

What are the particles in ionic substances?

A

Ions

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71
Q

What are the particles in metallic substances?

A

Ions and delocalised electrons

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72
Q

What is the formula for monatomic substances?

A

Just the symbol e.g. Ar

from periodic table

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73
Q

What is the formula for simple molecular substances?

A

Need to learn common examples e.g. H₂O, CO₂, CH₄

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74
Q

What is the formula for giant covalent substances?

A
  • Elements: just the symbol

* Compounds: ratio of atoms e.g. SiO₂, C, Si (need to learn common examples)

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75
Q

What is the formula for ionic substances?

A

Need to learn common ions and work out formula

e.g. MgCl₂ - ratio of Mg:Cl ions is 1:2

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76
Q

What is the formula for metallic substances?

A

Just the symbol e.g. Fe

from periodic table

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77
Q

What is the type of structure and bonding in sodium chloride?

A

Ionic

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78
Q

What are the particles present in sodium chloride?

A

Ions - Na+ and Cl-

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79
Q

Explanation of bonding in sodium chloride?

A

+vely charged Na ions attracted to -vely charged Cl ions to create ionic bond

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80
Q

Explanation for melting/boiling point of sodium chloride?

A

High - strong ionic bonds take a lot of energy to overcome

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81
Q

Explanation for conductivity of sodium chloride?

A

Conducts when molten or dissolved in water - ions need to be free to move

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82
Q

Why is sodium chloride brittle?

A

When shifted, ions will be next to same charged ion and repel naturally

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83
Q

What is the type of structure and bonding in magnesium?

A

Giant metallic lattice of Mg ions - metallic bonding within the metal

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84
Q

What particles are present in magnesium?

A

Ions and delocalised electrons

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85
Q

Explanation of bonding in magnesium?

A

Attraction between +vely charged ions and -vely charged electrons - electrostatic forces of attraction

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86
Q

Explanation for melting/boiling point of magnesium?

A

High - electrostatic forces of attraction take a lot of energy to overcome

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87
Q

Explanation for conductivity of magnesium?

A

Can conduct as delocalised electrons can move and carry a charge

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88
Q

What is the type of structure and bonding in diamond?

A

Macromolecular - carbon atoms

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89
Q

What are the particles present in diamond?

A

Carbon atoms

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90
Q

Explanation of bonding in diamond?

A

Each carbon atom is bonded to four other carbon atoms covalently

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91
Q

Explanation for melting/boiling point of diamond?

A

High - many strong covalent bonds - lots of energy to break - and no intermolecular forces

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92
Q

Explanation for conductivity of diamond?

A

Doesn’t conduct - neutral as not charged, and no free electons - no moving particles

(however good thermal conductor)

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93
Q

What is the type of structure and bonding in graphite?

A

Macromolecular - giant covalent structure

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94
Q

What are the particles present in graphite?

A

Atoms

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95
Q

Explanation of bonding in graphite?

A
  • each carbon atom joined to three others covalently
  • carbon atoms form layers with a hexagonal arrangement (layers have weak forces in between)
  • each atom has one non-bonded, delocalised electron
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96
Q

Explanation for melting/boiling point of graphite?

A

High - covalent bonds

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97
Q

Explanation for conductivity of graphite?

A

Conducts - delocalised electron per atom is free to move

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98
Q

What is the type of structure and bonding in iodine?

A

Covalent, simple molecular, with weak van de waal forces in between

99
Q

What are the particles present in iodine?

A

Two iodine atoms - diatomic iodine molecules

100
Q

Explanation of bonding in iodine?

A

Iodine molecules joined by weak van de waal’s (intermolecular) forces

101
Q

Explanation for melting/boiling point of iodine?

A
  • m.p. - 113.7 and b.p. - 183.35

* low - van de waals forces are weak so don’t need as much energy to break

102
Q

Explanation for conductivity of iodine?

A

Doesn’t conduct - covalently bonded so no delocalised electrons

103
Q

Can iodine undergo sublimation?

104
Q

What is the type of structure and bonding in ice?

A
  • Covalent (simple molecular)

* With intermolecular force - hydrogen bond

105
Q

What are the particles present in ice?

A

H₂O molecules - hydrogen and oxygen covalently bonded

106
Q

Explanation of bonding in ice?

A

Slightly +ve regions of one molecule attracted to slightly -ve regions of another to form hydrogen and vice versa

107
Q

Explanation for melting/boiling point of ice?

A
  • b.p. - 100 and m.p. - 0

* relatively high for a simple molecular substance due to hydrogen bonds

108
Q

Explanation for conductivity of ice?

A
  • using distilled water doesn’t conduct - no charged particles that are free to move
  • tap water contains ions
109
Q

Why is ice less dense than water?

A

Molecules are arranged further away in ice

110
Q

What is the shape of a simple molecule and ion determined by?

A

The number and type of electron pairs that surround the central atom

111
Q

How do charge clouds try and arrange themselves?

A

In a way that minimises the repulsion between charge clouds - by moving as far apart as possible

112
Q

What is the shape name and therefore bond angle of CO₂?

A

Linear - 180°

113
Q

Why is CO₂ a linear molecule?

A

It has double bonds

114
Q

What are the two types of electron pairs in a molecule?

A
  • bonding pairs

* non bonding / lone pairs

115
Q

What are bonding pairs of electrons?

A

Those that are involved in a chemical bond

116
Q

What are non-bonding/lone pairs of electrons?

A

Those that are not involved in a chemical bond

117
Q

Which type of electron pair have a greater repulsion in a molecule?

A

Lone pairs

118
Q

How does repulsion increase, in terms of electron pairs?

A

bonding- bonding pair < lone - bonding pair < lone - lone pair

119
Q

What are the steps taken to predict the shape of a molecule?

A
  1. determine no. of electron pairs
  2. identify number of lone pairs
  3. use these to determine shape of molecule and bond angle
120
Q

How do you calculate the number of electron pairs surrounding the central atom?

A

(number of outer electrons of central atom (group no.) + no. of atoms bonded) ÷ 2

121
Q

How do you calculate the number of lone pairs?

A

Number of electron pairs - number of atoms bonded

122
Q

How much do lone pairs reduce bond angle by?

A

Approximately 2.5°

123
Q

Why do lone pairs reduce bond angles?

A

Due to increased repulsion

124
Q

Work out the shape and bond angle of CH₄.

A

Tetrahedral, 109.5°

125
Q

Work out the shape and bond angle of H₂O.

A

V-shaped, 104.5°

126
Q

What is the shape and bond angle of a molecule with two bonding, and no non-bonding electron pairs?

A

Linear, 180°

127
Q

What is the shape and bond angle of a molecule with three bonding, and no non-bonding electron pairs?

A

Trigonal planar, 120°

128
Q

What is the shape and bond angle of a molecule with two bonding, and one non-bonding electron pair?

A

Bent (v-shaped), 117.5°

129
Q

What is the shape and bond angle of a molecule with four bonding, and no non-bonding electron pairs?

A

Tetrahedral, 109.5°

130
Q

What is the shape and bond angle of a molecule with three bonding, and one non-bonding electron pair?

A

Trigonal pyramidal, 107°

131
Q

What is the shape and bond angle of a molecule with two bonding, and two non-bonding electron pairs?

A

Bent (v-shaped), 104.5°

132
Q

What is the shape and bond angle of a molecule with five bonding, and no non-bonding electron pairs?

A

Trigonal bipyramidal, 120° and 90°

133
Q

What is the shape and bond angle of a molecule with four bonding, and one non-bonding electron pair?

A

Trigonal pyramidal or see-saw, 120° and 90°

134
Q

What is the shape and bond angle of a molecule with three bonding, and two non-bonding electron pairs?

A

Trigonal planar or t-shape, 120° and 90°

135
Q

What is the shape and bond angle of a molecule with two bonding, and three non-bonding electron pairs?

A

Linear, 180°

136
Q

What is the shape and bond angle of a molecule with six bonding, and no non-bonding electron pairs?

A

Octahedral, 90°

137
Q

What is the shape and bond angle of a molecule with five bonding, and one non-bonding electron pair?

A

Square based pyramid, 90°

138
Q

What is the shape and bond angle of a molecule with four bonding, and two non-bonding electron pairs?

A

Square planar, 90°

139
Q

What is the shape and bond angle of a molecule with three bonding, and three non-bonding electron pairs?

A

T-shaped, 90°

140
Q

What is the shape and bond angle of a molecule with two bonding, and four non-bonding electron pairs?

A

Linear, 180°

141
Q

What is affected when determining molecule shape in the case of an ion?

A

The number of electrons of the central atom is affected by its charge

142
Q

When determining the shape of a positive ion, what has to be done to the number of outer electrons of the central atom?

A

Subtract number of electrons lost from number of outer electrons

143
Q

When determining the shape of a negative ion, what has to be done to the number of outer electrons of the central atom?

A

Add number of electrons gained to number of outer electrons

144
Q

Work out the shape and bond angle of an NH4+ ion.

A

Tetrahedral, 109.5°

145
Q

Work out the shape and bond angle of a CIF4- ion.

A

Square planar, 90°

146
Q

What is affected when working out the molecule shape of an ion instead of an atom?

A

The number of electrons of the central atom is affected by its charge

147
Q

How can the molecule shape of a positive ion be worked out?

A
  • the central atom has lost that no. of electrons

* so the charge needs to be subtracted from the no. of outer electrons of the central atom

148
Q

How can the molecule shape of a negative ion be worked out?

A
  • the central atom has gained that no. of electrons

* so the charge needs to be added to the no. of outer electrons of the central atom

149
Q

Why is the bond angle in an amide ion smaller than that in an ammonia molecule?

A
  • amide ion has 2 lone pairs; ammonia has one

* more repulsion between two lone pairs and pushes them closer together

150
Q

What is electronegativity?

A

The power of an atom to attract the pair of electrons in a covalent bond

151
Q

What does electron density describe?

A

How the negative charge in an atom is distributed

152
Q

What does electronegativity depend on?

A
  • nuclear charge (no, of protons)
  • atomic radius
  • shielding
153
Q

What is atomic radius?

A

The distance between the nucleus and outer shell electrons

154
Q

What happens to electronegativity as nuclear charge increases?

A

It increases

155
Q

What happens to electronegativity as atomic radius increases? Why?

A

It decreases - bonding pair of electrons is further from the nucleus so less attracted to charge of the nucleus

156
Q

What happens to electronegativity as shielding increases?

A

It decreases

157
Q

What happens to electronegativity across a period?

A

It increases

158
Q

Why does electronegativity increase across a period?

A
  • proton no. ↑ and size ↑ so nuclear charge ↑

* shielding doesn’t ↑ as electron shells are the same

159
Q

What happens to electronegativity down a group?

A

It decreases

160
Q

Why does electronegativity decrease down a group?

A
  • shielding ↑ (as no. of shells ↑)

* and distance from nucleus (atomic radius) ↑

161
Q

What is the most electronegative element?

162
Q

Is there a difference in electronegativity in a covalent bond with two atoms that are the same?

163
Q

Why is electron density evenly distributed in a covalent bond between two atoms that are the same?

A

Because there is no difference in electronegativity

164
Q

What results in a non-polar bond?

A

When there is an equal sharing of electrons due to having no difference in electronegativity

165
Q

What happens in a covalent bond between two atoms that are the same, in terms of electronegativity?

A

A non-polar bond is formed due to no difference in electronegativity

166
Q

What happens in a covalent bond between two atoms that are the different, in terms of electronegativity?

A

A polar bond is formed, due to a difference in electronegativity and therefore the shared electrons not being evenly distributed

167
Q

What does the difference in electronegativity between atoms in a covalent bond result in?

A

The shared electrons not being evenly distributed, and therefore a polar bond

168
Q

When there is unequal sharing of electrons in a covalent bond, which atom will the electron pair be drawn towards?

A

The atom that is more electronegative

169
Q

When does a polar bond form?

A
  • In a covalent bond between two atoms that are different
  • difference in electronegativity therefore unequal sharing of electrons
  • electron pair drawn towards more electronegative atom = partial charges
170
Q

What does a larger difference in electronegativity mean for a polar bond?

A

The bond is more polar

171
Q

How is a polar bond thought to be, in terms of character?

A

Covalent with some ionic character

172
Q

Can ionic bonds show covalent character?

173
Q

How can ionic bonds show covalent character?

A

The ionic nature of the bond is reduced as the electron cloud is distorted by strong charges of the ion

174
Q

What happens if a cation has a high positive charge or small size?

A

It will tend to attract electrons towards itself

175
Q

What happens if an anion has a high negative charge or large size?

A

It will have an electron cloud that is easily distorted

176
Q

What will it mean for an ionic bond if ions have high charges?

A

The polarity of the bond increases, giving the bond covalent character

177
Q

What does the degree of ionic or covalent character depend on?

A

The atoms involved

178
Q

What would happen if there were no forces holding water molecules together?

A

They would move apart and become a gas

179
Q

What are the forces that exist between covalent molecules called?

A

Intermolecular forces

180
Q

How strong are intermolecular forces compared to covalent bond?

A

Between 1/10 and 1/100 of the strength

181
Q

What are the three types of intermolecular forces?

A
  1. Van der Waal’s
  2. permanent dipole-dipole forces
  3. hydrogen bonds
182
Q

Where do Van der Waal’s forces exist?

A

Between all molecules where they’re liquid or solid

183
Q

Why is there an imbalance in charge on a molecule, even though there may be no difference in electronegativity?

A

At any instant the electron distribution in a non-polar covalent bond can be asymmetrical due to constant movement of electrons

184
Q

What is a temporary dipole?

A

When there is an imbalance of charge on a molecule due to the electron distribution in a non-polar bond being asymmetrical because of constant electron movement

185
Q

What does a temporary dipole do to the adjacent molecule?

A

Induces an opposite dipole - and this continues through the structure

(this is known as the van der waal’s force)

186
Q

How are van der waal’s forces formed?

A

When there is an asymmetrical distribution of electrons in a molecule which leads to an imbalance in charge - and this induces an opposite charge on the next molecule

187
Q

What are van der waal’s forces also known as?

A

Temporary dipole-dipole forces

188
Q

Can van der waal’s form in monoatomic substances?

189
Q

What is the only intermolecular force in diatomic and monoatomic molecules?

A

Van der waal’s - but they occur between all molecules

190
Q

Are van der waal’s forces in addition to any other intermolecular force?

191
Q

What is the strength of van der waal’s forces dependant on?

A
  • no. of electrons present/size of molecule

* shape of the molecule

192
Q

What does a molecule having more electrons mean for the strength of its van der waal’s forces?

A

more electrons = stronger van der waal’s

193
Q

Why does having more electrons increase the strength of van der waal’s forces?

A
  • the outer electrons are further from the nucleus and so attracted less strongly by the nucleus
  • so temporary dipoles are easier to induce
194
Q

What does a branching mean for the strength of its van der waal’s forces?

A

When molecules have branches the van der waal’s forces are weaker

195
Q

Why does branching decrease the strength of van der waal’s forces?

A

The molecules are further away when branched, so straight chain alkanes pack closer together

196
Q

What is the weakest type of attraction that can exist between molecules?

A

Van der Waal’s

197
Q

What are polar molecules?

A

Molecules where there is a difference in electronegativity between the atoms involved

198
Q

Which type of molecules have permanent dipoles?

A

Polar molecules

199
Q

What type of dipole do polar molecules have?

200
Q

How do polar molecules attract each other?

A

Each molecules has opposite charges on each end; these opposite charges on the different molecules attract eachother

201
Q

Are permanent dipoles induced?

A

No, they already exist

202
Q

Why are permanent dipoles not constantly changing?

A

Because their dipoles already exist due to differences in electronegativity

203
Q

How will molecules with permanent dipoles arrange themselves?

A

So that oppositely charged ends are closest

204
Q

Do all polar molecules have polarity overall?

205
Q

Which polar molecules do not have overall polarity?

A

Symmetrical molecules

206
Q

Why do symmetrical molecules not have polarity overall?

A

The dipoles can cancel each other out to leave the molecule with no overall polarity

207
Q

What is meant by a symmetrical molecule?

A

A molecule where the dipoles are the same around a central atom or bond

208
Q

What is the strongest intermolecular force that can exist between symmetrical polar molecules?

A

Van der Waal’s

209
Q

Which are stronger: Van der Waal’s or permanent dipole-dipole?

A

Permanent dipole-dipole

210
Q

What is the strength of a permanent dipole dependant on?

A

The difference in electronegativity of the atoms in the polar bond

211
Q

What will happen to a permanent dipole-dipole attraction when there is a larger difference in electronegativity?

A

There will be greater partial charges and therefore a stronger permanent dipole-dipole attraction

212
Q

Which atoms have a large difference in electronegativity to hydrogen atoms?

A

Nitrogen, oxygen and fluorine

213
Q

What are the three most electronegative atoms?

A

Nitrogen, oxygen and fluorine

214
Q

What does the large difference in electronegativity between hydrogen and either nitrogen, oxygen or fluorine mean?

A

Very polar bonds are formed, and therefore a strong permanent dipole

215
Q

What is the strongest type of permanent dipole?

A

Hydrogen bonding

216
Q

What is a stronger; a hydrogen bond or a covalent bond?

217
Q

What is the least electronegative atom?

218
Q

When are strong permanent dipoles formed?

A

When there is a large difference in electronegativity between hydrogen and either fluorine, oxygen or nitrogen

219
Q

How is a hydrogen bond different to a dipole-dipole force?

A

The nitrogen, oxygen or fluorine atom can draw hydrogen’s single electron towards itself

220
Q

What is the result of nitrogen, oxygen or fluorine drawing hydrogen’s single electron towards them?

A

Hydrogen exposes its single proton and therefore a strong δ+ hydrogen end and a strong δ- end

221
Q

What is required for a hydrogen bond to form?

A

A lone pair of electrons on the N, O or F

222
Q

Why is it important that the non-hydrogen atom has a lone pair to form a hydrogen bond?

A

So this lone pair can strongly attract the δ+ end of hydrogen - which hydrogen can’t repel without any non-bonding electrons

223
Q

Why can’t hydrogen repel the non-hydrogen atom in a hydrogen bond?

A

It has no non-bonding electrons

224
Q

What is a hydrogen bond?

A

The attraction between the strong δ+ hydrogen end and the lone pair of electrons on a N, O or F atom

225
Q

What does a hydrogen bond have characteristics of?

A

A dipole-dipole force as well as a covalent bond

226
Q

How strong is a hydrogen bond in comparison to a covalent bond?

A

1/10 the strength of a covalent bond

227
Q

What is the strongest type of intermolecular force?

A

Hydrogen bonds

228
Q

What are the three intermolecular forces responsible for keeping the structure of hair in place?

A
  • hydrogen bonds
  • salt linkages
  • disulfide bonds
229
Q

How do hydrogen bonds affect the structure of molecules in hair?

A

Hydrogen bonds attract the long keratin molecules to each other and them into a spiral shape

230
Q

How can the corresponding strength of intermolecular forces be represented as?

A

Van der waal’s < permanent dipole-dipole < hydrogen

231
Q

When will a substance turn from a solid to a liquid?

A

When there is enough heat energy to break enough intermolecular forces

232
Q

What is melting temperature?

A

The heat energy required to break enough intermolecular forces to turn a substance into a lquid

233
Q

What is boiling temperature?

A

The energy required to break enough intermolecular forces to turn that substance into a gas

234
Q

What do substances with hydrogen bonds tend to have higher boiling points than?

A

Substances with permanent dipole-dipole bonds, and those with van der waal’s

235
Q

What is the pattern of boiling points of the hydrides in periods 2-5?

A

Apart from the 1st in the period, there is a general increase across the period

236
Q

Why does the boiling point of hydrides in groups 2-5 increase across the period?

A
  • molecules are increasing in size
  • due to more electrons
  • greater polarization and dipole increases in strength
  • (van der waal’s also stronger due to bigger molecule)
237
Q

In a graph showing boiling point of hydrides between periods 2-5, how and why do H2O, HF and NH3 not follow the trend?

A
  • how? they are much higher than the other hydrides in their group
  • why? there are hydrogen bonds between molecules
238
Q

Why does H2O have a greater boiling point than HF and NH3?

A
  • H2O - equal no. of hydrogens and oxygens so more bonds - also 2 lone pairs on oxygen so quite -ve attraction between H&O
  • HF - strong bonds (F is most electronegative) but unequal ratio of H to lone pairs
  • NH3 - lower b.p. as less electronegative & unequal ratio of N to H
239
Q

Why are the boiling points of group 4 hydrides lower than those in 5,6, and 7?

A

Molecules in group 4 have symmetrical tetrahedral shape - so electronegativity cancels out = no permanent dipole - only have van der waal’s

240
Q

What happens to molecules when water is heated?

A

They gain energy, moving faster and further apart as a result of hydrogen bonds being broken

241
Q

What happens to water’s volume and density as it is heated?

A

Water becomes less dense and volume increases

242
Q

What happens to water’s volume and density as it is cooled?

A

Water becomes more dense UNTIL it reaches 4°C

243
Q

When is water at its most dense?

244
Q

Why does water float and expand when it is frozen?

A

Hydrogen bonds are quite long and therefore hold the water molecules slightly further apart

AQA A Level Physical Chemistry (12 decks)

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