Brainscape's Knowledge GenomeTM

Browse over 1 million classes created by top students, professors, publishers, and experts.


See full index

A Level Chemistry > Bonding > Flashcards

Bonding Flashcards

You may prefer our related Brainscape-certified flashcards:
AP® Chemistry flashcards Periodic Table flashcards Chemistry 101 flashcards Organic Chemistry 101 flashcards
1
Q

Define ionic bonding

A

Ionic bonding occurs between metals and non-metals. The oppositely charged ions are held together by strong forces of electrostatic attraction within a giant ionic lattice.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
2
Q

The greater the ionic charge …

A

The greater the ionic charge, the stronger the electrostatic attraction and therefore the stronger the ionic bond.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
3
Q

The smaller the ionic radius …

A

The smaller the ionic radius, the more tightly packed the atoms can be and therefore the stronger the ionic bond

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
4
Q

The higher the charge density

A

the stronger the ionic bond

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
5
Q

When are cations formed?

A

when elements lose electrons to become stable.
Generally, this happens to metals in groups 1, 2 and 3, each losing their respective number of electrons to become positively charged

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
6
Q

When are anions formed?

A

when elements gain electrons to become stable.

Generally, this happens to non-metals in groups 5, 6 and 7, each gaining 3, 2 or 1 electron(s) respectively to become negatively charged

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
7
Q

Why does ionic radius increase down a group?

A

Because extra electron shells are added

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
8
Q

What are isoelectronic ions?

A

Those that have the same electronic configuration

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
9
Q

Explain why for a set of isoelectronic ions, the atomic number increases, the ionic radius then decreases

A

due to the positive nucleus attracting the electrons more strongly hence drawing them nearer

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
10
Q

What is the melting and boiling points of ionic substances like?

A

High

Oppositely charged ions are held together with strong forces of electrostatic attraction within a giant ionic lattice

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
11
Q

Do ionic substances conduct electricity?

A

When solid, the ions are held in fixed positions within the lattice, but when molten/dissolved they are free to move and hence can carry a charge to form a current.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
12
Q

Ionic substances- brittle

A

When put under strain, the lattice is disrupted and cations could be put over cations, and anions over anions. They would hence repel each other and the lattice would break.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
13
Q

Ionic substances- solubility

A

Only soluble in polar substances like water, because the ions are charged so can be pulled apart by polar molecules but not by non-polar molecules

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
14
Q

What is the migration of ions?

A

Copper Chromate solution is dropped onto damp filter paper
Crocodile clips are connected to a power source and can then be connected to both sides of the paper
Electrolysis occurs causing the negative chromate anions to move to the positive anode, whilst the positive copper cations move to the negative cathode
The solution is green: blue colouring shows the copper ions, yellow colouring shows the chromate ions.
Ions must therefore be charged and able to move/’migrate’ whilst in aqueous solution

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
15
Q

Define covalent bonding

A

Covalent bonding occurs between non-metals only (usually). The positively charged nuclei and negatively charged shared pair of electrons between them are held together by strong forces of electrostatic attraction. This can be within a giant or simple covalent structure

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
16
Q

The shorter a bond…

A

the higher its bond enthalpy

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
17
Q

Explain electron density on bond length

A

In covalent bonding, there are attractive forces between the nuclei and shared electrons. There are also repulsive forces between the two nuclei and shared electrons in themselves

Hence, the distance between the two nuclei is a balance of all of these forces acting opposingly and this is what is defined as the bond length.

So therefore the bond will be shorter if there is more electron density of if the atom’s involved are smaller.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
18
Q

Explain Valence Shell Electron Pair Repulsion Theory

A

All electrons are negatively charged. Therefore, their repulsion about a central atom determines the shape of simple molecules and ions because repulsive forces are so that they are kept to a minimum.

Lone pair-Lone pair > Lone pair-Bonding pair > Bonding pair-Bonding pair

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
19
Q

What do lone pairs cause bond angles to do?

A

Decrease by 2.5

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
20
Q

Describe the linear shape?

A

2 bonding pairs & 0 lone pairs 180˚ angles

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
21
Q

Describe the trigonal planar

A

3 bonding pairs & 0 lone pairs 120˚ angles

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
22
Q

Describe the angular shape?

Could also be described as bent or v shaped

A

2 bonding pairs & 1 lone pair

<120˚ *(119˚) angles

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
23
Q

Describe the tetrahedral shape

A

4 electrons and 0 lone pairs

109.5 angles

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
24
Q

Describe the trigonal pyramidal shape (tetrahedral shape)

A

3 bonding pairs & 1 lone pair

<109.5˚ *(107˚) angles

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
25
Q

Tetrahedral angular shape

A

2 bonding pairs & 2 lone pairs

<109.5˚ *(104.5˚) angles

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
26
Q

Trigonal Bipyramidal shape

A

5 bonding pairs & 0 lone pairs

120˚ and 90˚ angles

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
27
Q

Seesaw shape

A

4 bonding pairs & 1 lone pair

*87˚ and 102˚ angles

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
28
Q

Distorted t shape

A

3 bonding pairs & 2 lone pairs

*87.5 angles

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
29
Q

Linear trigonal Bipyramidal

A

2 bonding pairs & 3 lone pairs

180˚ angles

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
30
Q

Octahedral

A

6 bonding pairs & 0 lone pairs

90˚ angles

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
31
Q

Square pyramidal

A

5 bonding pairs & 1 lone pair

*90˚ and 81.9˚ angles

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
32
Q

Square planar

A

4 bonding pairs & 2 lone pairs

90˚ angles

33
Q

(Distorted) T-shaped

Octahedral

A

3 bonding pairs
& 3 lone pairs
No stable molecules known to exist

34
Q

Linear

Octahedral

A

2 bonding pairs
& 4 lone pairs
No stable molecules known to exist

35
Q

How to predict the shapes and bond angles of simple molecules and ions

A
  1. First draw the Lewis structure of the ion/molecule.
  2. Next the steric number (number of lone pairs add the number of atoms bonded to a central atom) can easily be identified from this.
  3. This can be used to work out the electron distribution (2 = linear, 3 = trigonal planar, 4 = tetrahedral, 5 = trigonal bipyramidal, and 6 = octahedral)
  4. Now the number of lone pairs should be taken into account to work out the final geometry of the ion/molecule.
36
Q

What is electronegativity?

A

Electronegativity is the ability of an atom to attract the bonding pair of electrons in a covalent bond. Informally, it can be described as ‘electron-pulling power’.

37
Q

How is electronegativity measured

A

It is measured on the Pauling scale with fluorine being the most electronegative element, followed by oxygen, chlorine and then nitrogen.

38
Q

Trends of electronegativity across a period and up a group

A

Increase

39
Q

What happens in purely covalent substances VS polar covalent substances

A

In purely covalent substances, the bonding electrons are shared completely equally between the two atoms

In polar covalent substances, the bonding electrons are shared unequally between the atoms, which therefore have partial charges on them

40
Q

What happens in purely ionic substances?

A

In purely ionic substances, one or more electrons are transferred completely from one atom to another, and therefore the atoms have full charges

41
Q

What is the general rule for electronegativity

A

a difference in electronegativity less than 0.4 results in pure covalent bonds; between 0.4 and 1.7 results in polar covalent bonds, and above 1.7 results in ionic bonds

42
Q

What is a dipole

A

A dipole is a difference in charge between the two atoms in a polar bond caused by a shift in electron density in the bond towards the more electronegative element.

43
Q

What does being a polar molecule depend on

A

The presence of polar bonds

The geometry/symmetry of the molecule

44
Q

If polar bonds are present and the molecule is symmetrical is the molecule present?

A

NO

45
Q

If polar bonds are present and the molecule is asymmetrical, is the molecule polar?

A

YES

46
Q

What are instantaneous dipole-induced dipole bonds called?

A

London forces

47
Q

Explain the existence of instantaneous dipole-induced dipole bonds

A

Electrons move around quickly and randomly and can be thought of as a cloud of charge about the nucleus.

At any particular point in time, it is unlikely that the electrons will be evenly distributed and rather that they are more likely to be found on one side

This is an instantaneous dipole where the area concentrated with electrons is partially negative and the area with an absence is partially positive

This dipole is able to induce a dipole in a neighbouring atom/molecule

48
Q

How can dipoles be created/destroyed instantaneously?

A

because the particles are moving constantly as well as the electrons, causes an overall electrostatic force of attraction

49
Q

London forces are the …

A

weakest type of intermolecular force and affects all atoms/ molecules

50
Q

Where do permanent dipole-permanent dipole bonds occur

A

in molecules that contain polar bonds but are asymmetrical
There is electrostatic attraction between the poles of neighbouring molecules. This occurs in addition to instantaneous dipoles

51
Q

Are permanent dipole-dipole bonds stronger than London forces?

A

Yes but not stronger than hydrogen bonding

52
Q

What does Hydrogen bonding affect?

A

Hydrogen bonding affects molecules in which hydrogen is covalently bonded to either fluorine, oxygen, or nitrogen – all highly electronegative elements

53
Q

How does hydrogen bonding occur?

A

The highly electronegative element attracts electrons away from the hydrogen and forms a polar bond.
There is electrostatic attraction between hydrogen and the electronegative elements.
The hydrogen has an extremely high charge density and forms weak bonds with the lone pair(s) of electrons on the other element

54
Q

Hydrogen bonding is the …

A

strongest type of intermolecular force

55
Q

Explain the hydrogen bonding in oxygen

A

The oxygen has a partially negative charge as it is highly electronegative, and hence attracts the
electron pair away from the hydrogen, which is therefore partially positive.

Hydrogen bonds form between the oxygen atoms
(and its two lone pairs) and hydrogen atoms of neighbouring molecules.

56
Q

Explain the bonding in ammonia

A

The nitrogen has a partially negative charge as it is highly electronegative, and hence attracts the electron pair away from the hydrogen, which is therefore partially positive.

Hydrogen bonds form between the nitrogen atoms (and its lone pair) and hydrogen atoms of neighbouring molecules.

57
Q

Explain the bonding in hydrogen fluoride

A

The fluorine has a partially negative charge as it is highly electronegative, and hence attracts the electron pair away from the hydrogen, which is therefore partially positive.

Hydrogen bonds form between the fluorine atoms (and its three lone pairs) and hydrogen atoms of neighbouring molecules.

58
Q

Why is water able to do hydrogen bonding while the other group 6 hydrides can not?

A

Oxygen is electronegative
However the other group 6 hydrides do permanent dipole-permanent dipole bonding instead, however the trend from H2S to H2Te is due to London forces - the dipole becomes greater due to more electrons

59
Q

Hydrogen bonding in ice vs liquid

A

In ice, each water molecule is bonded to 4 other water molecules in a tetrahedral geometry through covalent bonding and hydrogen bonding. This leads to a relatively open hexagonal 3D structure with several holes between molecules, which therefore lead to a lower density.

In liquid water however, the molecules are able to slide over each other because the particles have more kinetic energy. This leads to their being closer together, which is why ice floats in water

60
Q

Predicting the presence of hydrogen bonding in molecules

A

Hydrogen must be bonded to a small highly electronegative element
There must be at least one lone pair of electrons on the highly electronegative element

61
Q

In terms of intermolecular forces describe the trends in boiling temperatures of alkanes with increasing chain length

A

chain length increases –> more total electrons and therefore a bigger
‘electron cloud’ and increase in volume around the molecule.

More more likely to be a more uneven distribution of (more) electrons leading to a greater instantaneous dipole (with more partial charge).

This will induce greater dipoles and therefore there will be more electrostatic attraction between molecules, leading to increasingly high boiling points for greater chain lengths.

62
Q

Explain, in terms of intermolecular forces, the effect of branching in the carbon chain on the boiling temperatures of alkanes

A

Linear molecules- higher boiling points than branched chain molecules because they are able to align themselves closer together by lying parallel to each other. Which makes the electrostatic attraction between dipoles stronger.

Larger surface areas allows for more opportunity for dipoles to exist and therefore more places where molecules can be attracted to each other. Hence giving linear molecules higher boiling points

63
Q

How is volatility determined?

A

By boiling point where high boiling points denote a low volatility.
Alkanes are non-polar molecules and as such can only do instantaneous dipole-induced dipole bonding.
Alcohols, in addition, are able to do hydrogen bonding due to their hydroxyl group.

64
Q

Alcohols will have a higher… / lower… than alkanes of the same size

A

boiling point

viscosity

65
Q

Why is hydrogen bonding less significant amongst longer chain alcohols

A

due to the main intermolecular force acting is London dispersion due to the non-polar hydrocarbon chain

66
Q

Why does HF have the highest boiling point

A

it is able to form hydrogen bonds between molecules

67
Q

What characterises the boiling points frim HCl to HI

A

London dispersion forces.
Due to the increase in electrons, greater dipoles form and there is a stronger electrostatic attraction between molecules

68
Q

What has to happen in order for a substance to be able to dissolve?

A

The solute particles must become surrounded by the solvent particles

The intermolecular forces formed between the solute and solvent particles must be strong enough to overcome the forces between the solute particles and solvent particles respectively.

If the enthalpy change is about 0 then the process is neither energetically favourable nor unfavourable so entropy will facilitate the process.

69
Q

What happens when an ionic compound is dissolved in water?

A

the positive ions are attracted to the partially negative oxygen atoms and the negative ions are attracted to the partially positive hydrogen atoms.

70
Q

What happens during the hydration stage when an ionic compound dissolved in water

A

the ions are pulled out of their giant ionic lattice and surrounded by water molecules

The ionic bonds in the lattice are therefore broken as well as the hydrogen bonds between water molecules. This is an endothermic process and requires energy.

o New hydrogen bonds are formed between the ions and water molecules. This is an exothermic process and gives out energy.

71
Q

What happens to NaCl so dissolution can occur

A

The total enthalpy change is 0, however is still positive (exothermic/ unfavourable), so dissolution occurs because of entropy.

72
Q

Why is Al2O3 insoluble

A

It has strong ionic bonds due to Al3+’s high charge density.

Entropy is not strong enough to drive the process

73
Q

What are the factors that influence the choice of solvents in terms of hydrogen bonding

A

Alcohols are polar substances due to their hydroxyl groups
Hydrogen bonds between the water molecules and the alcohol molecules break and new hydrogen bonds form between the water and alcohol molecules collectively.
For longer alcohol molecules, the hydrogen bonding is ‘diluted’ by the instantaneous dipole-induced dipole bonding performed by the carbon chain, hence the longer the carbon chain, the less soluble an alcohol will be

74
Q

Why are halogenoalkanes insoluble

A

The strongest intermolecular force between halogenoalkane molecules is permanent dipole-permanent dipole bonding. Therefore hydrogen bonding can’t be formed.

The hydrogen bonding in water is stronger than the permanent dipole-permanent dipole bonding that would form between water and halogenoalkanes, hence they are insoluble.
And entropy isn’t enough to make this process favourable because it’s too endothermic.

75
Q

What does non aqueous mean

A

refers to any solvent that is not water

76
Q

Why are cyclohexanes and heptanes often used?

A

‘like dissolves like’
Non polar molecules- the strongest intermolecular forces present are instantaneous dipole-induced dipole bonding so non-polar substances will usually dissolve readily in each other.

77
Q

What is metallic bonding?

A

Metallic bonding occurs between metals only. The metal atoms lose their outer electrons, which become delocalised and they themselves become positive metal ions.

78
Q

How are giant metallic lattices held together?

A

The giant metallic lattice is held together by the strong electrostatic attraction between the positive metal ions and the delocalised
(negative) sea of electrons.

A Level Chemistry (12 decks)

Brainscape helps you reach your goals faster, through stronger study habits.
© 2024 Bold Learning Solutions. Terms and Conditions