Brainscape's Knowledge GenomeTM

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


See full index

A-level Chemistry > 3.1 - The periodic table and periodicity > Flashcards

3.1 - The periodic table and periodicity 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 the term ‘periodicity’ ?

A

Periodicity - The study of properties of elements across each period that show a repeating pattern

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

Explain how elements are arranged in the periodic table ?

A
  • They are arranged in increasing atomic number
  • They are arranged in periods
  • They are arranged in groups
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
3
Q

What does the arrangement of elements in periods represent ?

A
  • It represents the number of shells atom contains
  • It represents repeating trends in physical and chemical properties
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
4
Q

What does the arrangement of elements in groups represent ?

A
  • It represents the number of out shell electrons an atom contains
  • It shows elements that have similar chemical properties
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
5
Q

Define the term ‘atomic radius’ ?

A

Atomic radius - Half the distance between the nuclei of two adjacent atoms

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

What is the trend in atomic radius across a period ?

A

As you go across the period, the atomic radius decreases

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

Explain the trend in atomic radius across a period ?

A
  • As you go across the period, the number of protons increases but the shielding remains the same
  • This means that the effective nuclear charge increases which means the outer shell electrons are more attracted to the positive nucleus
  • This means they are pulled more towards the nucleus causing the atomic radius to decrease
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
8
Q

Identify the s, p, d and f blocks on the periodic table ?

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

How do you determine what block is assigned to an element ?

A
  • An element can be assigned to the s, p, d or f block according to where it highest energy electrons are
  • Eg. Magnesium is in the s-block since its outer electrons / highest energy electron is in the 3s orbital ( 1s2 2s2 2p6 3s2 )
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
10
Q

Explain the increasing trend in melting points for the first three elements across periods 2 and 3 ?

A
  • They have a giant metallic lattice structure
  • As you go across the period, the number of outer shell electrons increases
  • This means metals form positive metal ions with a greater positive charge
  • This means the positive metal ions are more attracted to delocalised electrons which means metallic bonding becomes stronger
  • This means more energy is required to overcome strong metallic bonds which causes melting/ boiling point to increase
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
11
Q

Explain the trend in melting points for the fourth element across periods 2 and 3 ?

A
  • They have a giant covalent lattice structure
  • Therefore, atoms are only held together by many strong covalent bonds which require a lot of energy to overcome
  • This means they have a very high melting/ boiling point
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
12
Q

Explain sudden decrease between the fifth and seventh element across periods 2 and 3 ?

A
  • They have a simple molecular lattice structure
  • Although atoms are held together by strong covalent bonds, molecules are held together by weak London forces which don’t require a lot of energy to overcome
  • This means they have a very low melting/ boiling point
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
13
Q

Explain the increasing trend in melting points between the fifth and seventh element across period 2 ?

A
  • As you go across the period, the number of electrons increases
  • This results in larger dipoles and therefore stronger London forces which require more energy to overcome
  • Therefore, the melting/ boiling point increases across the period
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
14
Q

Explain the increasing trend in melting points between the fifth and seventh element across period 3 ?

A
  • As the size of the molecule increases, the number of electrons present increases
  • This results in larger dipoles and therefore stronger London forces that require more energy to overcome
  • Therefore larger molecules have higher melting/ boiling points
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
15
Q

Explain the trend in melting points for the eight element across periods 2 and 3 ?

A
  • These elements are monoatomic
  • This means they have a low number of electrons which results in very small dipoles
  • This means they have very weak intermolecular forces between atoms which don’t require a lot of energy to overcome resulting in very low melting/ boiling points
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
16
Q

What is the melting point determined by ?

A
  • It is dependant on the size of the molecule
  • As the size of the molecule increases, the number of electrons present increases
  • This results in greater dipoles and therefore stronger London forces which require more energy to overcome
  • This causes the melting/ boiling point to increase
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
17
Q

What is the structure of metals ?

A
  • A Giant metallic lattice structure
  • A 3D arrangement of positive metal ions surrounded by a sea of delocalised electrons held together by strong metallic bonds
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
18
Q

Define the term ‘metallic bond’ ?

A

Metallic bond - The strong electrostatic attraction between positive metal ions and delocalised electrons

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

What are the features of metals ?

A
  • High melting/ boiling point
  • Electrical and thermal conductors
  • Insoluble
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
20
Q

Explain why metals have high melting/ boiling points ?

A
  • Metals contain strong metallic bonds that require a lot of energy to break/ overcome
  • This means they have a large melting/ boiling point
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
21
Q

Explain why metals are electrical conductors ?

A
  • Metals contain delocalised electrons which can move
  • This means metals can conduct electricity in the liquid and solid state
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
22
Q

Explain why metals are insoluble ?

A

Metals do not dissolve

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

What is the structure of a simple molecule ?

A
  • They have a simple molecular lattice structure
  • Atoms are held together by strong covalent bonds while molecules are held together by weak London forces/ induced dipole-dipole interactions
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
24
Q

What are the features of simple molecular substances ?

A
  • Relatively low melting/ boiling point
  • Cannot conduct electricity
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
25
Q

Explain why simple molecular substances have a relatively low melting/ boiling point ?

A
  • There are only strong covalent bonds between atoms while there are weak London forces between molecules
  • London forces are less strong/ weaker than strong covalent bonds so less energy is required to overcome them
  • This results in a low melting/ boiling point
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
26
Q

Explain why simple molecular substances cannot conduct electricity ?

A
  • Simple molecular substances do no contain moving electrons or mobile ions
  • Therefore, they cannot conduct electricity
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
27
Q

What is the structure of an ionic compound ?

A
  • Ionic compounds have a giant ionic lattice structure
  • It is a 3D arrangement of oppositely charged ions held together by strong covalent bonds
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
28
Q

Define the term ‘ionic bond’ ?

A

Ionic bond - The strong electrostatic attraction between oppositely charged ions

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

What are the features of an ionic compound ?

A
  • They have a high melting/ boiling point
  • They are soluble
  • They cannot conduct electricity in the solid state
  • They can conduct electricity in the molten/ aqueous state
30
Q

Explain why ionic compounds have a high melting/ boiling point ?

A
  • Oppositely charged ions are held together by strong ionic bonds that require a lot of energy to overcome
  • This means they have high melting / boiling point
31
Q

Explain why ionic compounds are soluble ?

A

Ionic compounds dissolve in most polar solvents

32
Q

Explain why ionic compounds are both poor and good electrical conductors ?

A
  • In the solid state, ions are fixed in the giant ionic lattice structure therefore no mobile charge carriers and cannot conduct electricity
  • In the molten/ aqueous state, ions are mobile so can conduct electricity
33
Q

Define the term ‘covalent bond’ ?

A

Covalent bond - The shared pair of electrons

34
Q

What is the structure of Giant covalent substances ?

A
  • They have a giant covalent lattice structure
  • A 3D network of atoms held together by strong covalent bonds
35
Q

Name some Giant covalent substances ?

A
  • Graphene
  • Graphite
  • Diamond
  • Silicon
36
Q

What is the structure of diamond ?

A
  • Each carbon atom is covalently bonded to 4 other carbon atoms
  • It is very hard, very strong and has a high melting point due to carbon atoms being held together by strong covalent bonds that take a lot of energy to overcome
37
Q

What is the structure of graphite ?

A
  • Each carbon is covalently bonded to 3 other carbon atoms
  • One outer shell electron from each carbon atom becomes delocalised
  • Each layer of graphite is held to the next by weak intermolecular forces ( slippery )
38
Q

What is the structure of graphene ?

A
  • Each carbon atom is covalently bonded to 3 other carbon atoms
  • One outer shell electron from each carbon atom becomes delocalised, so it conducts electricity
  • It is only one atom thick ( transparent + flexible )
39
Q

What are the features of Giant covalent substances ?

A
  • High melting/ boiling point
  • Insoluble
  • Not electrical conductors ( apart from graphene/ graphite )
40
Q

Explain why Giant covalent substances have a high melting/ boiling point ?

A
  • Atoms are held together by many strong covalent bonds that require a lot of energy to overcome
  • This means they have a high melting/ boiling point
41
Q

Explain why Giant covalent substances insoluble ?

A
  • They do not dissolve in almost all solvents
  • Arms are held together by strong covalent bonds which require too much energy to overcome to be broken by interacting with solvents
42
Q

Explain why graphene and graphite electrical conductors ?

A
  • In graphene/ graphite, each carbon atom donates one delocalised electron to the ‘sea’ of delocalised electrons
  • Graphene/ graphite contain delocalised electrons that can move so can conduct electricity
43
Q

What are the factors that affect electronegativity ?

A
  • Nuclear charge
  • Atomic radius
  • Shielding
44
Q

Define the term ‘electronegativity’ ?

A

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

45
Q

Explain how nuclear charge affects electronegativity ?

A
  • The more protons in the nucleus, the stronger the electrostatic attraction between the nucleus and bonding pairs of electrons
46
Q

Explain how atomic radius affects electronegativity ?

A
  • The smaller the atomic radius, the smaller the distance between the outer shell electrons and nucleus
  • This means the electrostatic attraction between the nucleus and bonding pairs of electrons is stronger
47
Q

Explain how shielding affects electronegativity ?

A
  • As the number of shields increases there are more electrons between the nucleus and the bonding pairs of electrons
  • This means the electrostatic attraction between the nucleus and bonding pairs of electrons in weaker electrons
48
Q

What is the trend in electronegativity down the group ?

A

As you go down the group, the electronegativity decreases

49
Q

Explain the trend in electronegativity down the group ?

A
  • As you go down the group the atomic radius and shielding increases
  • This means there is more distance and a larger number of electrons between the nucleus and the bonding pairs of electrons
  • This weak rand the electrostatic force between the nucleus and bonding pairs of electrons
50
Q

What is the trend in electronegativity across the period ?

A

As you go across the period, the electronegativity increases

51
Q

Explain the trend in electronegativity across the period ?

A
  • As you go across the period the atomic radius decreases and the number of protons increases and shielding remains the same meaning the effective nuclear charge increases
  • This means the atomic radius decreases
  • This means there is shorter distance between the nucleus and bonding pairs of electrons
  • This means that the electrostatic attraction between the nucleus and bonding pairs of electrons is stronger
52
Q

Define the term ‘first ionisation energy’ ?

A

First ionisation energy - The amount of energy required to remove one mole of electrons from one mole of atoms in the gaseous state at STP ( Standard conditions of temperature and pressure )

53
Q

Write an equation for the first ionisation energy of Sodium/ Na ?

A

Na (g) –> Na+ (g) + e-

54
Q

Do you know how to write an equation for the ionisation energies of different elements ?

A

Yes

55
Q

What factors affect first ionisation energy ?

A
  • Atomic radius
  • Nuclear charge
  • Shielding
56
Q

How does atomic radius affect ionisation energy ?

A
  • As the atomic radius increases, the greater the distance between the outer shell electrons and positive nucleus
  • This means the outer shell electrons are less targeted to the nucleus so less energy is required to remove them
  • This causes the ionisation energy to decrease
57
Q

How does effective nuclear charge affect ionisation energy ?

A
  • As the number of protons increase, the nuclear charge increases
  • This can cause the effective nuclear charge to increase which means the outer shell electron is more attracted to the nucleus
  • Therefore, more energy is required to remove the outer shell electron causing the ionisation energy to increase
58
Q

How does shielding affect ionisation energy ?

A
  • As the shielding increases, the number of shells increases
  • This means the outer shell electrons are less attracted to the nucleus, meaning less energy is required to remove the outer shell electron
  • This causes the ionisation energy to decrease
59
Q

What is the trend in ionisation energy down the group ?

A

As you go down the group, the ionisation energy decreases

60
Q

Explain the trend in ionisation energy down the group ?

A
  • As you go down the group, the atomic size increases and the shielding increases
  • This means that the outer shell electrons and further away from the nucleus and therefore less attracted to the nucleus
  • Therefore, less energy is required to remove them causing the ionisation energy to decrease
61
Q

What is the trend of ionisation energies across a period ?

A

As you go across a period, the ionisation energy generally increases

62
Q

Explain the trend in ionisation energy across a period ?

A
  • As you go across a period, the number of protons increases but the shielding remains the same
  • This means the effective nuclear charge increases which means the outer shell electrons are more attracted to the positive nucleus
  • this means more energy is required to remove the outer shell electrons which causes the ionsaition energy to increase
63
Q

Draw a graph showing the successive ionisation energies of an element ( ie. Al ) ?

A
64
Q

Can you draw a graph showing the successive ionisation energy of an element ?

A

Yes

65
Q

Interpret this graph, explain what element in the period two this graph represents ?

A
  • There is a ‘jump’ in ionisation energy between the third and fourth electrons removed
  • This means that the fourth electron is removed from an inner shell which means that the element has three outer shell electrons
  • This means that the element is Boron since it is in group 3, period 3
66
Q

Can you interpret values of successive ionisation energy and a successive ionisation energy graph ?

A

Yes

67
Q

Draw a diagram representing the ionisation energies across a period ( eg. period 2 ) ?

A
68
Q

Can you draw a graph representing ionisation energy across a period ?

A

Yes

69
Q

Explain the anomaly in ionisation energies across the period for ie. Boron ?

A
  • Boron has a lower ionisation energy than beryllium since it contains a 2p1 ( different shell for different periods ) electron which is slightly higher in energy
  • This means less energy is required to remove it meaning it has a lower ionisation energy
70
Q

Explain the anomaly in ionisation energies across the period for ie. Oxygen ?

A
  • Oxygen has a lower ionisation energy than nitrogen since it contains four electrons in the p sub-shell
  • This means that two electrons occupy on of the p orbital and repel due to like charges
  • This means less energy is required to remove the 2p4 electron meaning it has a lower ionisation energy

A-level Chemistry (29 decks)

Brainscape helps you reach your goals faster, through stronger study habits.
© 2025 Bold Learning Solutions. Terms and Conditions