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A Level Chemistry > Bonding and Structure - Topic 2 > Flashcards

Bonding and Structure - Topic 2 Flashcards

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1
Q

What are London forces

A

The attractive forces between all atoms and molecules

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2
Q

Why does an increase in pi bonds increase the London forces

A

As the bigger the electron cloud the more uneven electron distribution there is so there would be a much stronger imbalance in charge causing there to be a stronger dipole, so a higher polarizability so stronger LDN forces

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3
Q

covalent bonding

A

Strong electrostatic attraction between the nuclei and two atoms and the bonding pair of electrons, a single covalent bond is formed between two atoms when an atomic orbital containing a single electron from one atom overlaps with an atomic orbital that contains one electron of another atom

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4
Q

What is expansion of the octet

A

when an atom is able to have more than 8 valence electrons

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5
Q

What are molecular orbitals

A

The overlap of atomic orbitals from separate atoms

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6
Q

How many electrons can molecular orbitals hold

A

2

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7
Q

Types of molecular orbitals

A

Sigma
Pi

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8
Q

Where are sigma bonds found

A

Between atoms

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9
Q

Where are pi bonds found

A

Below and above atoms

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10
Q

How are sigma bonds formed

A

When orbitals from two separate atoms overlap

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11
Q

How are pi bonds formed

A

When p orbitals in separate atoms overlap sidewards forming pi bonds above and below the atoms

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12
Q

Properties of sigma bonds

A
  • all single bonds are sigma bonds
  • sigma bonds are free to rotate as rotating each nuclei has no impact on the bonding orbital
  • have a high level of attraction between nuclei and shared electrons so are very strong
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13
Q

How many pi and sigma bonds are there in double bonds

A

x1 sigma
X1 pi

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14
Q

How many pi and sigma bonds are there in triple bonds

A

x1 sigma
x2 pi

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15
Q

Why are pi bonds weaker than sigma bonds

A

As the electrons in the pi bond are further away from the nuclei compared to sigma bonds so has there is a weaker level of attraction making the bonds easier to break

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16
Q

What is hybridisation

A

The concept of mixing w atomic orbitals to create a new type of hybridised orbital which has a different shape and energy level

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17
Q

Number of hybridised and unhybridised orbitals and what is the hybridised orbital called in single bonds

A

0 unhybridised orbitals
4 hybridised orbitals
Name: sp3 orbitals

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18
Q

Number of hybridised and unhybridised orbitals and what is the hybridised orbital called in double bonds

A

1 unhybridised p orbital
3 hybridised orbital
Name: sp2

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19
Q

Number of hybridised and unhybridised orbitals and what is the hybridised orbital called in triple bonds

A

2 unhybridised p orbital
2 hybridised orbitals
Name: sp

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20
Q

In terms of hybridised orbitals explain why carbon has a tetrahedral shape

A

Each sp3 orbital has the same energy so the electrons will have equal repulsion between all 4 orbitals causing the orbitals to point away from each other at maximum distances at the angle of 109.5 degrees between then

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21
Q

No lone pairs 2 electron pairs geometric shape name

A

Linear

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22
Q

Bond angle of linear molecule

A

180 degrees

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23
Q

3 electron pairs no lone pairs name of molecule

A

Triagonal planar

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24
Q

What are electron domains

A

Things attached to the central atoms e.g other atoms or lone pairs

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25
Q

What makes up a linear geometric shape

A

2 electron pairs no lone pairs

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26
Q

What makes up a triagonal planar shape

A

3 electron pairs no lone pairs

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27
Q

2 bonding pairs 1 lone pair geometrical shape

A

V shaped

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28
Q

2 bonding pairs 2 lone pairs shape

A

V shaped

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29
Q

Bond angle for triangular planar

A

120 degrees

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30
Q

Bond angle of v shape

A

104.5 degrees

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31
Q

What makes up a v shape

A

2 bonding pairs of electrons and 2 lone pairs

OR

2 bonding pairs and 1 lone pairs of electrons

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32
Q

4 bonding pairs of electrons and no lone pairs geometric shape

A

Tetrahedral

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33
Q

Tetrahedral bond angle

A

109.5 degrees

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34
Q

What makes up a tetrahedral shape

A

4 bonding pairs no lone pairs

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35
Q

3 bonding pairs 1 lone pair geometric shape

A

Tetrahedral pyramidal

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36
Q

pyramidal bond angle

A

107 degrees

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37
Q

3 bonding pairs 1 lone pairs geometric shape

A

pyramidal shape

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38
Q

6 bonding pairs no lone pairs

A

Octahedral

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39
Q

What makes up an octahedral

A

6 bonding pairs

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40
Q

What makes up a triangular bipyramidal shape

A

5 bonding pairs no lone pairs

41
Q

Triagonal bipyramidal bond angles

A

90 and 120 and 180

42
Q

Octahedral bond angles

A

90 and 180

43
Q

why can polar substances
i) molecules containing electronegative atoms
ii) NH3 HF and alcohols
dissolve in water

A

i) as the positive part of water attracts the electronegative part of the molecule and the negative part of the water attracts to the positive part of the molecule
2) as they all form hydrogen bonds between water molecules

44
Q

trends of electronegativity

A
  • increases across a period
  • decreases down a group
44
Q

why does electronegativity increase across a period

A

as the nuclear charge increases causing the atomic radius to decrease but the number of shielding remains constant so there is a stronger attraction between the nucleus and bonding pairs of electrons

45
Q

why does electronegativity decrease down a group

A

as the number of shielding electrons increases the atomic radius increases so this means that the shared electron pair is further from the nucleus so less attraction between nucleus and bonding pairs of electrons

46
Q

why cant non polar substances dissolve in water

A

These molecules don’t have regions of partial positive or partial negative charge, so they aren’t electrostatically attracted to water molecules and their weak intermolecular forces cannot disturb the strong hydrogen bonding between water molecules

47
Q

why does solubility decrease as HC chain length increases in alcohols

A
  • As when the length of the hydrocarbon chain increases, the London forces predominate
  • the non-polar carbon chain cannot form hydrogen bonds, so as we increase the length a greater part of the molecule is unable to hydrogen bond with water, causing long chain alcohols to be less soluble than short-chain
48
Q

what is required for a substance to dissolve

A
  • soluble particles must be separated from each other and become surrounded by solvent molecules
  • the forces of attraction between solute and solvent molecule must be strong enough to overcome the solvent-to-solvent forces and the solute-to-solute forces
49
Q

definition for electronegativity

A

the ability of an atom to attract the bonding pairs of electrons in a covalent bond

50
Q

what is a saturated solution

A

a solution that contains as much solute as possible at a particular temperature

51
Q

definition for solubility

A

a measure of the concentration of a saturated solution of a solute at a specific temperature

52
Q

factors affecting electronegativity

A
  • nuclear charge
  • atomic radius
  • shielding
53
Q

how does nuclear charge affect electronegativity

A

more protons stronger attraction between nucleus and bonding pairs of electrons

54
Q

how does atomic radius affect electronegativity

A

closer to the nucleus the stronger the attraction between nucleus and bonding pairs of electrons

55
Q

how does shielding affect electronegativity

A

the less shielding so the less repulsion and a stronger attraction between the nucleus and bonding pairs of electrons

56
Q

what are the electronegative atoms

A

chlorine bromine and iodine

57
Q

what atoms form hydrogen bonds

A

O F N

58
Q

3 types of intermolecular forces

A
  • induced dipole dipole interactions (London forces)
  • permanent dipole dipole interactions
  • hydrogen bonding
59
Q

What causes induced dipoles to form

A

By the random movement of electrons causing an uneven distribution of electrons on one side of the atom. This forms an instantaneous dipole. The negative charge of the instantaneous dipoles repels the electrons in the other atom causing the electrons to move away from the atoms forming an induced dipole. Each atom them experiences a force of attraction between the positive and negative end of the atom, called a London force

60
Q

How does an increase in electrons affect London forces and the boiling point of an atom

A

The more electrons there are the stronger the London forces so more energy is required to over come them, so the boiling point increases

61
Q

how does the shape of a molecule affect the london forces

A
  • the interactions between long thin molecules are stronger than those between short fat molecules because the interactions between long thin molecules can take effect over a larger surface area
  • so the more branched the molecule the lower the surface area of contact and therefore the weaker London forces
62
Q

What causes permanent dipole dipole interactions

A

This forms when there is an atom with a greater electronegativity so attracts pair of electrons in the covalent bond more strongly causing the electronegative atom to be slightly negative forming a permanent dipole. When two permanent dipoles attract each other they form forming a permanent dipole interaction

63
Q

Why does tetrachloromethane have no overall permanent dipole

A

As the molecule is completely symmetrical so the bond polarities cancel out so it has no overall permanent dipole

64
Q

Why does tetrachloromethane have a higher boiling point than trichloromethane even though it has no overall permanent dipole

A

As it has more electrons so stronger London forces so more energy is required to overcome the London forces

65
Q

What causes hydrogen bonding

A

The electronegative element (nitrogen oxygen or fluorine) strongly attracts the pairs of electrons in the covalent bond causing it to be slightly negative and the hydrogen atom to have a slightly positive charge. The positive hydrogen atom of the other polar molecule (that forms hydrogen bonds) attracts to the lone pair of electrons in oxygen ,fluorine or nitrogen molecule forming a hydrogen bond

66
Q

How many lone pairs of electrons in fluoride nitrogen and oxygen

A

3 - fluoride
2 - oxygen
1 - nitrogen

67
Q

Conditions for hydrogen bonding

A
  • we need a hydrogen atom bonded to a strongly electronegative element
  • the electronegative atom has to have at least one lone pair of electrons
68
Q

Water properties due to hydrogen bonding

A
  • it has a high boiling and melting point due to the hydrogen bonds so requires a lot of energy to overcome hydrogen bonds
  • solid form of water is less dense than liquid form
69
Q

Why is ice less dense than water

A

As when water reaches its freezing point (0 degrees Celsius) the water molecules arrange themselves into an ordered structure which is stabilised by a network of hydrogen bonds so the water molecules are further apart, which makes ice less dense than liquid water

70
Q

How does ice having a low density help organisms

A

Due to its low density it floats on water, insulating the water below preventing the water from freezing completely

71
Q

definition of hydrogen bonding

A

a strong intermolecular force between delta-positive hydrogen atom covalently bonded to a F N or O and a lone pair of electrons on the delta negative of the O F or N atom of a nearby molecule

72
Q

What is dative bonding

A

When an atom uses a lone pair of electrons to form a covalent bond

73
Q

How is a dative bond shown in a displayed formula

A

Using an arrow

74
Q

Why isn’t a carbon to carbon double bond stronger than a carbon to carbon single bond

A

As the double bond contains 1x sigma and 1x pi whilst the single bond contains 1x sigma bonds, and its easier to break the pi bond rather than a sigma bond as sigma bonds are stronger

75
Q

what is ionic bonding

A

the strong electrostatic attraction between two oppositely charged ions

76
Q

How are ionic bonds formed

A

When a metal transfers its electrons to a non metal so that they both have the same electron configuration of a noble gas causing them to be oppositely charged. And the electrostatic attraction between oppositely charged ions is called an ionic bond

77
Q

What do the square brackets represent in ionic bonding drawings

A

That the charge is spread over the whole ion

78
Q

How do ionic compounds/ionic crystals form

A

When a metal and non metal react the ions produced arrange them self into a giant ionic lattice forming an ionic compound

79
Q

Properties of ionic compounds

A
  • they have high mp and bp - due to the strong electrostatic forces of attraction so requires a lot of energy to overcome electrostatic forces of attraction
  • they are soluble in polar substances such as water - as the water cluster around the ions and bind to them and the energy released when water molecules bind to the ions is enough to overcome the electrostatic forces of attractions holding the ions lattice together
  • they don’t conduct electricity when solid - as the ions are in a fixed in place by the electrostatic forces of attraction so they cannot carry a charge
  • hard brittle crystalline substances
80
Q

Why does magnesium oxide have a higher mp than sodium chloride

A
  • magnesium oxide is in group 2 so has a higher charge and a higher charge density so has stronger electrostatic forces of attraction so more energy is needed to overcome so has a higher mp
81
Q

Why does solubility decrease as the charge of ionic compounds increases

A

As the hydration energy released when the ions are hydrated (water molecules binding to them) is not large enough to overcome the electrostatic forces of attraction holding the lattice together

82
Q

isoelectronic definition

A

molecules and ions with the exact same number and arrangement of electrons

83
Q

Properties for simple molecular substances

A
  • they have low melting and boiling points - due to the weak London forces between molecules so require little energy to overcome
  • do not conduct electricity as they don’t have any delocalised electrons or freely moving ions
  • only polar simple substances are soluble in water
84
Q

Why do giant covalent structures have high mp and bp

A
  • have high melting and boiling points due to the strong covalent bonds between atoms so requires a lot of energy to overcome
85
Q

How are the atoms arranged in diamond and what is its bond angle

A

Tetrahedral structure
Bond angle: 109.5

86
Q

Properties of diamond

A
  • high mp and bp
  • does not conduct any electricity as every electron in its covalent bond are fixed between pairs of atoms so no delocalised electrons to act as charge carriers
  • insoluble as solvents cannot disrupt the larger number of covalent bonds
  • good thermal conductor due to the fixed covalent bonds meaning that the atoms close to the heat get hotter and move faster so the vibrations move quickly throughout the structure heating the whole structure
  • very hard due to the strong covalent bonds operating in 3D dimensions
  • sublimes at ordinary pressures due to strong covalent bonds
87
Q

How are the carbon atoms in graphite arrange and what’s the bond angle

A

Planar hexagonal structure
Bond angle: 120

88
Q

Properties of graphite

A
  • conducts electricity due to the delocalised electrons from each carbon atom being able to freely move THROUGHOUT the sheets
  • high mp and bp (same as diamond)
  • can act as a lubricant due to ADSORBED gases on the surface of carbon atoms (gas molecules sticking to the surface) - but lubricating properties are not there in a vacuum
  • insoluble in water (same reason as diamond)
  • lower density as diamond due to the relatively amount of large space between layers
89
Q

What is graphene

A

A single layer of graphite

90
Q

Suggest why boron nitrate can act as a lubricant in a vacuum whilst graphite cannot

A

As they don’t need gaseous atoms to be ADSORBED (not absorbed, ADSORBED) onto the surface of layers

91
Q

Properties of graphene

A
  • conducts electricity
  • can be rolled into a carbon nanotube
  • one atom thick
92
Q

Buckminister fullerene formula

A

C60

93
Q

Properties of silicone dioxide

A
  • high mp and bp due to strong silicone - carbon covalent bonds
  • hard
  • doesn’t conduct electricity
  • insoluble - as there are no attractions which could occur between solvent molecules and the silicone and oxygen atoms which could overcome the covalent bonds in the structure
94
Q

what is metallic bonding

A

The strong electrostatic attraction between positive metal ions and the sea of delocalised electrons

95
Q

Properties if metallic bonding

A
  • high mp and bp - due to the strong electrostatic forces of attraction between cations and delocalised electrons so requires a lot of energy to overcome these forces and move the metal ions away from their positions in the lattice
  • high density - as the atoms are closely packed with little space between them
  • can conduct electricity - due to the freely moving delocalised electrons so are attracted to the positive electrode and flow through the metal and this flow is an electrical current
  • can conduct heat - when the metal is heated the delocalised electrons move around and conduct heat to the other parts of the metal
  • malleable due to metal ions being able to slide over each other when a force is applied (which is known as a slip)
  • insoluble
96
Q

What is the electron pair repulsion theory

A

That the shape of a molecule is determined by the electron pairs surrounding the central atom as the electron pairs repel other electron pairs so move as far as possible to minimise repulsion

97
Q

Why do lone pairs reduce the bond angle by 2.5 degrees

A

As they repel more strongly

98
Q

why do branched molecules have a lower bp than unbranched molecules

A

as the molecules dont pack as close together so London forces are reduced as there are fewer contact points

A Level Chemistry (15 decks)

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