Brainscape's Knowledge GenomeTM

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


See full index

A Level Chemistry - AQA > Bonding > Flashcards

Bonding Flashcards

1
Q

Define Ionic Bonding

A

The electrostatic forces of attraction between oppositely charged ions formed by the transfer of electrons.

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

What are simple ions?

A

Single atoms which have lost all gained electrons to complete their outer shells.

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

What are compound ions ?

A

Ions made up of groups of atoms with an overall charge .

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

State the formula for a carbonate ion.

A

C03(2-)

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

State the formula a hydroxide ion.

A

OH(-)

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

State the formula a nitrate ion.

A

NO3(-)

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

State the formula a sulphate ion.

A

S04(2-)

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

State the formula an ammonium ion.

A

NH4(+)

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

State Three physical properties of ionic compounds.

A
  1. they conduct electricity when molten or dissolved
  2. they have high melting points
  3. they are soluble in water
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
10
Q

Why can ionic compounds conduct electricity when molten or dissolved in liquid?

A

Because the ions are free to move and carry the charge. However, in a solid they are all fixed in position by the strong ionic bonds so they are unable to move and carry the current.

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

Why do ionic compounds have high melting points?

A

The giant ionic lattices of held together by strong electrostatic forces of attraction Between the oppositely charged ions which require a lot of energy to be overcome.

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

Why do ionic compounds dissolve in water ?

A

Water molecules are polar the positive ends of the water molecule pulled the negative irons away from their lattice and the negative ends of the water molecule pulled the positive ions away from the lattice causing it to break apart and dissolve .

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

What structure do ionic compounds form?

A

Gian ionic lattices

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

What is a giant ionic lattice?

A

Regular structures made up of the same basic repeating unit.

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

State the type of bonding present in a molecule

A

Covalent bonding.

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

Define covalent bonding.

A

When two nonmetal atoms share a pair of electrons giving them full outer shells.

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

How are the shared electrons held in place?

A

By strong Electrostatic forces of attraction between the positive nuclei and negative shared electrons.

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

What are simple covalent compounds and what holds them together?

A

Compounds made up of lots of individual molecules. The atoms in the molecule are held together by strong covalent bonds but the molecules are held together by weaker intermolecular forces.

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

State three physical properties of simple covalent compounds.

A
  1. they cannot carry a current
  2. they have low melting points
  3. some are soluble whereas others are not
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
20
Q

Why can’t simple covalent compounds carry a current?

A

Simple molecules have no overall charge. All the electrons are held in place by strong covalent bonds so they are not free to move and carry a current.

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

Why do simple covalent compounds have low melting points?

A

The molecules in simple covalent compounds are held together by weak intermolecular forces which are easily overcome so they do not require much energy to break.

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

Explain the differences in the solubility of different covalent compounds .

A

The solubility of these simple covalent compounds depends on how polarised each molecule is.

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

What are giant covalent structures?

A

Huge networks of covalently bonded atoms. For example: diamond and graphite

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

Describe the structure of graphite (4).

A
  1. Each carbon atom is covalently bonded to 3 others
  2. The 4th outer electron Of each carbon Atom is delocalised
  3. Carbon atoms arranged in sheets of flat hexagons
  4. Sheets of hexagons are bonded together by week van der waals forces
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
25
Q

List 5 properties of graphite.

A
  1. low density
  2. lubricant
  3. electrical conductor
  4. insoluble
  5. very high mp
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
26
Q

Explain why graphite has a low density

A

The layers of carbon atoms are far apart in comparison to the covalent bond length. So it has low density and is used to make light strong sports equipment.

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

Explain why graphite can be used as a lubricant

A

The weak IMF between the layers in graphite are easily broken allowing the sheets to slide over each other Making the graphite quite slippery

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

Explain why graphite Is an electrical conductor

A

The delocalised electrons are free to move through the structure and carry the current

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

Explain why graphite is insoluble

A

The many covalent bonds in the sheets are too strong to break

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

Explain why graphite has a very high melting point

A

Because it takes a lot of energy to overcome the many strong covalent bonds

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

Describe the structure of diamond

A

Each carbon Atom is covalently bonded for other carbon atoms arranged in a tetrahedral shape

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

List 5 properties of diamond

A 
1 Very high melting point
2 Very hard
3 Good thermal conductor
4 Can’t conduct electricity
5 Insoluble
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
33
Q

Explain why diamond has a very high melting point

A

It takes a lot of energy to overcome the many strong covalent bond

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
34
Q

Explain why diamond is very hard

A

Because of the rigid network of carbon atoms joined by strong covalent bonds . Making it useful to be used in tools to cut things

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
35
Q

Explain why diamond is a good thermal conductor

A

Because vibrations can travel easily through the rigid lattice

36
Q

Explain why diamond can’t conduct electricity

A

All electrons are held in place by strong covalent bonds so none are free to move and carry the current

37
Q

Explain why diamond is insoluble

A

The covalent bonds in the sheets were too strong to break

38
Q

What is dative covalent bonding?

A

When does shared pair of electrons in the covalent bond come from only one of the bonding atoms.

39
Q

Give an example of dative covalent bonding and explain how the bond forms

A 
Ammonium ion (NH4+)
Forms when the nitrogen iron in ammonia donates a pair of electrons to the hydrogen ion
40
Q

Describe how a dative covalent bond is represented

A

The dative covalent bond is represented by an arrow going from the atom donating the electrons to the atom receiving them

41
Q

Define metallic bonding

A

Electrostatic forces of attraction between positive metal ions and negative delocalised electrons

42
Q

Metal elements exist as …

A

Giant metallic lattices

43
Q

Describe metallic bonding

A

The electrons in the outer most shell of a metal atom are delocalised leaving a positive metal ion. These two oppositely charged particles are attracted to each other, Forming a lattice of closely packed positive ions in a sea of delocalised electrons

44
Q

List 5 properties of metals

A 
1 High melting points
2 good electrical conductors
3 good thermal conductors
4 malleable
5 insoluble
45
Q

Explain why metals have high melting points

A

It takes a lot of energy to overcome the strong electrostatic forces of attraction between the positive metal ions and negative delocalised electrons

46
Q

Why are metals good electrical conductors?

A

The delocalised electrons are free to move on carrier current

47
Q

Explain why metals are good thermal conductors

A

The delocalised electrons are able to pass kinetic energy onto each other

48
Q

Explain why metals are malleable

A

The metal ions are able to slide over each other because nothing is holding them in fixed positions so metals can change shape

49
Q

Explain why metals are insoluble

A

This is due to the strength of the metallic bonds which require a lot of force to be overcome

50
Q

List 3 factors affecting the strength of metallic bonds

A

1 the proton number
2 the number of delocalised electrons per Atom
3 the size of the ion

51
Q

Explain how proton number affects the strength of the metallic bond

A

the greater number of protons the stronger the attraction between the nucleus of the metal ion and the delocalised electrons

52
Q

Explain how the number of delocalised electrons per atom reflects the strength of the metallic bond

A

The more delocalised electrons per atom the greater the charge of the metal ion and so the stronger the electrostatic forces of attraction between the positive metal ions and negative delocalised electrons

53
Q

Explain how the size of the iron affects the strength of the metallic bond

A

The smaller the ion the stronger the bond because the smaller the distance between the positive nucleus of the metal ion and the negative delocalised electrons

54
Q

Explain why magnesium has a higher melting point than sodium

A

Magnesium has 2 electrons in its outer shell, each magnesium atom loses 2 electrons leaving the magnesium ion with a positive charge of 2+. Sodium, however, has one electron in it’s outer shell which it loses giving it a positive charge of 1. Therefore, there are stronger electrostatic forces of attraction between the positive magnesium ions and negative delocalised electrons which require more energy to be overcome, giving magnesium a higher melting point.

55
Q

Bonding pairs and lone pairs of electrons exist in …

A

… charge clouds

56
Q

What is a charge cloud?

A

An area where you have a high chance of finding a pair of electrons

57
Q

Explain why charge clouds repel each other

A

electrons are all negatively charged so charge clouds repel each other until they are as far apart as possible

58
Q

Why are bonding angles reduced when lone pairs of electrons are added?

A

This is because lone pairs of electrons repel more than bonding pairs so the bonding pair angles are often reduced because they are pushed together by lone pair repulsion

59
Q

List the three types of charge cloud angles in order of descending angle size

A

Lone pair-Lone pair, Lone pair-Bonding pair, Bonding pair-Bonding pair

60
Q

Describe how you can find out how many lone pairs and bonding pairs of electrons there are on the central atom of a molecule

A
  1. Identify the central Atom
  2. workout the number of electrons in the outer shell of the central Atom
  3. add one electron for every Atom the central Atom is bonded to
  4. if the molecule is an ion you will need to add 1 electron for each negative charge or subtract 1 for each positive charge
  5. add all the electrons and divide by two to find the number of electron pairs
  6. Compare the number of electron pairs to the number of bonds to find out the number of lone players on the number of bonding pairs
61
Q

Describe the structure of an answer to a question asking you to explain the shape of a molecule:

A
  1. state the number of bonding pairs and lone pairs of electrons
  2. state electron pairs repel and try to get far as possible
  3. If there are no lone pairs state the bonding electron pairs repel equally
  4. if there are lone pairs state lone pairs repel more than bonding pairs
  5. states the actual shape and bond angle
62
Q

How would you draw a molecule containing a double bond?

A

Treat the double bond like a single bond.y

63
Q

Define electronegativity

A

And atoms ability to attract the pair of electrons in a covalent bond

64
Q

List the four most electronegative elements

A

F > O > N > CL

65
Q

List 3 factors affecting electronegativity

A

1 Nuclear Charge
2 Atomic radius
3 Shielding

66
Q

State the trend in electronegativity as we move across a period

A

It increases

67
Q

State the trend in electronegativity as we move down a group

A

It decreases

68
Q

How is a polar bond created?

A

In a covalent bond between two atoms with different electronegativities, the pair of electrons are pulled more strongly towards the more electronegative atom, making the bond polar.

69
Q

When is a covalent bond considered non-polar?

A

When two atoms of the same element have bonded together or when two atoms of different elements with very similar electronegativities (Eg: hydrogen and carbon) have bonded together.

70
Q

Why are some bonds non-polar?

A

This is due to the atoms having equal electronegativities, so the pair of electrons are equally attracted to both nuclei.

71
Q

What is a dipole?

A

A difference in charge between the two atoms in a covalent bond caused by a shift in electron density.

72
Q

How is a permanent dipole formed?

A

When two atoms with different electronegativities are bonded together, the electrons shift towards the more electronegative atom causing a charge is distributed unevenly across the whole molecule.

73
Q

When are molecules with polar bonds not polar and why is this?

A

When the polar bonds are arranged symmetrically, cancelling each other out.

74
Q

Name 3 types of intermolecular forces

A
  1. Permanent dipole-dipole forces
  2. Van der Waals forces
  3. Hydrogen bonding
75
Q

Describe how permanent dipole-dipole forces form

A

In a substance made up of molecule with permanent dipoles, there are weak electrostatic forces of attraction between the slightly positive and slightly negative charges on neighbouring molecules.

76
Q

Where are permanent dipole-dipole forces found?

A

Between polar molecules

77
Q

Describe how Van der Waals forces form.

A

Electrons in charge clouds are always moving really quickly. At any moment, the electrons are more likely to be to one side of the atom than the other. This gives the atom a temporary dipole. This dipole causes another temporary dipole in the opposite direction on a neighbouring atom. The two dipoles are then attracted to each other. The second dipole induces yet another dipole in a third atom. This happens over and over again. However, because the electrons are constantly moving the dye poles are created and destroyed all the time. Despite this, the overall effect is for the items to be attracted to each other.

78
Q

Van der Waal forces hold molecules together in a…

A

Lattice

79
Q

List 3 factors affecting the strength of Van der Waal forces

A
  1. the size of the molecule
  2. the shape of the molecule
  3. number of electrons in the molecule
80
Q

Explain how the size of the molecule affects the strength of Van der Waal forces

A

The larger the molecule the larger the electron clouds resulting in stronger Van der Waal forces

81
Q

Explain how the shape of the molecule affect strength of Van der Waal forces

A

Long straight molecules like closer than branched ones . The closer two molecules oh the stronger the forces between them.

82
Q

How are hydrogen bonds formed?

A

Hydrogen bonding occurs when a molecule contains a hydrogen atom covalently bonded to a F, O or N atom. F, O or N Are very electronegative so they draw the bonding electrons away from the hydrogen Atom. The bond is polarised and hydrogen has a high charge density. This results in the hydrogen atoms forming weak bonds with the lone pairs of electrons on the F. O or N atoms of other molecules.

83
Q

Which is the strongest type of Intermolecular force?

A

Hydrogen bonding

84
Q

Why do substances with hydrogen bonds have higher boiling/melting points than similar molecules?

A

This is because of the extra energy needed to break the hydrogen bonds which are not present in the other molecules.

85
Q

Explain why ice is less dense than liquid water

A

As Liquid water cools to form ice, the molecules make more hydrogen bonds and arrange themselves into a regular lattice structure. The hydrogen bonds between the water molecules are longer in ice, so the molecules are further apart than in liquid water. So, ice is less dense than liquid water.

A Level Chemistry - AQA (6 decks)

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