Periodicity

Question 1

Who created the modern periodic table?

Answer 1

Dmitri Mendeleev

Question 2

How were the elements ordered by Mendeleev?

Answer 2

Increasing atomic number

Question 3

What is the name for the vertical columns of the periodic table and what can you tell me about them?

Answer 3

Groups - Each element in a group has atoms with the same number of valence electrons and similar properties.

Question 4

What is the name for the horizontal rows of the periodic table and what can you tell me about them?

Answer 4

Periods - The number of the period gives the number of the highest energy electron shell in an elements atom.

Question 5

Why are hydrogen and helium not really apart of any group?

Answer 5

Since the electronic configurations of H and He are unusual, they don’t fit comfortably into any group. They are thus allocated a group based on similarities in physical and chemical properties with other members of the group.

He is placed in group 0 on this basis, but hydrogen doesn’t behave like any other element and so is placed in a group of its own.

Question 6

What is periodicity?

Answer 6

A repeating trend in properties of the elements across each period of the PT.

Question 7

What are the 4 periodic trends we can look at across the PT?

Answer 7

1.Electron configuration
2.Ionisation energy
3.Structure
4.Melting points

Question 8

What’s the chemistry of each element determined by?

Answer 8

Its electron configuration, particularly the valence electron/s in its highest energy shell

Question 9

What’s the trend in electron configuration across period 2?

Answer 9

Across period 2, the 2s subshell fills with 2 electrons, followed by the 2p subshell with 6 electrons

Question 10

What’s the trend in electron configuration across period 3?

Answer 10

Across period 3, the same pattern of filling is repeated for the 3s and 3p subshell ( 2 electrons the 6).

Question 11

What’s the trend in electron configuration across period 4?

Answer 11

Across period 4, although the 3d subshell is involved, the highest energy shell number is n=4 . From the n=4 shell, only the 4s and 4p subshells are occupied

Question 12

What the key to understanding electron configuration across a group?

Answer 12

Each period starts with an electron in a new highest energy shell.

For each period, the s and p subshells are filled in the same way - a periodic pattern.

Question 13

What the trend in electron configuration down a group?

Answer 13

Remember elements in each group have atoms with the same number of valence electrons in the outer energy level.

Elements in each group also have atoms with the same number of electrons in each subshell. This similarity in electron configuration gives elements in the same group their similar chemistry.

Question 14

How are elements in the periodic table divided?

Answer 14

They’re divided into blocks corresponding to their highest energy subshell, giving 4 distinct blocks (s,p,d,f).

Question 15

What are s- block elements and where are they located?

Answer 15

The s-block elements are all those with only s electrons in the outer shell (the valence electrons occupy an s atomic orbital).Located on left of PT.

Question 16

What are P- block elements and where are they located?

Answer 16

The p-block elements are all those with at least one p-electron in the outer shell (one of the valence electrons occupies a p atomic orbital).Located on right of PT.

Question 17

What are D- block elements and where are they located?

Answer 17

The d-block elements are all those with at least one d-electron and at least one s-electron but no f or p electrons in the outer shell (up to 5d).Located in centre of PT.

Question 18

What are F- block elements and where are they located?

Answer 18

The f-block elements are all those with at least one f-electron and at least one s-electron but no d or p electrons in the outer shell.Loctaed at bottom of PT.

Question 19

What is ionisation?

Answer 19

The removal of one or more electrons from an atom.

Question 20

What’s ionisation energy?

Answer 20

A measure of how easily an atom loses electrons to form positive ions?

Question 21

Define first ionisation energy

Answer 21

The energy required to remove one electron from each atom in one mole of gaseous atoms of an element to form one mole of gaseous 1+ ions.

Question 22

What are the 3 factors that affect IE?

Answer 22

1.Atomic radius
2.Nuclear charge
3.Electron shielding

Question 23

How does atomic radius affect ionisation energy?

Answer 23

The greater the distance between the nucleus and valence electrons, the less nuclear attraction. The force of attraction falls off sharply with increasing distance, so atomic radius has a large effect.

Question 24

How does nuclear charge affect ionisation energy?

Answer 24

The more protons there are in the nucleus of an atom, the greater the attraction between the nucleus and the valence electrons. Remember opposite charges attract.

Question 25

How does electron shielding affect ionisation energy?

Answer 25

Electrons are negatively charged and so inner-shell electrons repel outer shell electrons. This repulsion, called the shielding effect, reduces the attraction between the nucleus and the outer electrons.

Question 26

How many ionisation energies does an atom have?

Answer 26

As many ionisation energies as it has electrons.

Question 27

Why is the second ionisation of helium greater than the first?

Answer 27

In a helium atom, there are two protons attracting two electrons in the 1s subshell. After the first electron is lost, the single electron is pulled closer to the helium nucleus .The nuclear attraction of the remaining electron increased and more ionisation energy will be needed to remove the second electron.

Question 28

Define second ionisation energy

Answer 28

The energy required to remove one electron from each ion in one mole of gaseous 1+ ions of an element to form one mole of gaseous 2+ ions.

Question 29

What do successive ionisation energies provide evidence for?

Answer 29

The different energy levels in an atom

Question 30

What could an extremely large difference in ionisation energy suggest?

Answer 30

That the next electron must be being removed from a different shell closer to the nucleus, where there’s a greater nuclear attraction and less shielding.

Question 31

What do successive ionisation energies allow us to make predictions about?

Answer 31

1.The number of electrons in the outer shell
2.The group of the element in the PT
3.The identity of the element

Question 32

What do periodic trends ion first ionisation energy provide evidence for?

Answer 32

The existence of shells and subshells

Question 33

What causes the large jumps in successive ionisation energies?

Answer 33

Moving down to a closer shell, as these electrons are closer so experience a greater nuclear attraction.

Question 34

What predictions can be made from a graph of successive ionisation energies?

Answer 34

The number of electrons in the outer shell, the group of the element in the periodic table and thus the identity of the element.

Question 35

Explain the trend of first ionisation energy down a group

Answer 35

First ionisation energy decreases down every group in the PT, for the same reasons. Although the nuclear charge increases, its effect is outweighed by the increased atomic radius and to a lesser extent the increased shielding.

Atomic radius increases,
More inner shells so shielding increases,
Nuclear attraction on outer electrons decreases,
First ionisation energy decreases.

Question 36

Explain the general trend of first ionisation energy across a period

Answer 36

First ionisation energy increases across a period.

Nuclear charge increases - greater nuclear attraction between protons in nucleus and valence electrons
Same shell/ similar shielding (doesn’t have a great impact)
Atomic radius decreases - closer so greater attraction

Question 37

What are the subshell trends in first ionisation energy?

Answer 37

Although the first ionisation energy shows a general increase across period 2 and 3, it does fall in two places in each period. The drop occurs at the same positions in each period, suggesting a periodic cause. The reason is linked to the exitance of subshells, their energies and how atomic orbitals fill with electrons.

Question 38

In period 2, explain the fall from beryllium to boron of first ionisation energies

Answer 38

The fall, in the first ionisation energy from beryllium to boron marks the start of the filling of the 2p subshell.

The 2p subshell in boron has a higher energy than the 2s subshell in beryllium. Therefore, in boron the 2p electron is easier to remove than one of the 2s electrons in beryllium. The first ionisation energy of boron is less than the first ionisation energy of beryllium.

There is a slight decrease in first ionisation energybetween beryllium and boron as the fifth
electron in boron is in the 2psubshell,which is further awayfrom the nucleus than the 2s subshell
of beryllium
Beryllium has a first ionisation energyof 900 kJ mol as its electron configuration is 1s 2s
Boron has a first ionisation energy of 801 kJ mol as its electron configuration is 1s 2s 2p

Question 39

In period 2, explain the fall from nitrogen to oxygen

Answer 39

The fall in first ionisation energy from nitrogen to oxygen marks the start of electron pairing in the p orbitals of the 2p subshell/

In nitrogen and oxygen the highest energy electrons are in a 2p subshell.
In oxygen, the paired electrons in one of the 2p orbitals repel one another, making it easier to remove an electron from an oxygen atom than a nitrogen atom. Therefore the first ionisation energy of oxygen is less than the first ionisation energy of nitrogen. check notes to make more sense image how the orbitals are filled.

There is a slight decrease in first ionisation energybetween nitrogen and oxygen as the paired
electrons in the 2psubshell of oxygen repel each other,making it easierto remove an electron in
oxygen than nitrogen.
Nitrogen has a first ionisation energyof 1402 kJ mol as its electron configuration is
1s 2s 2p
Oxygen has a first ionisation energy of 1314 kJ mol as its electron configuration is
1s 2s 2p

Question 40

Making predictions from successive ionisation energies, graph analysis, and a general explanation of IE.

Answer 40

The successive ionisation energies of an element increase as removing an electron from a positive ion is more difficult than from a neutral atom.
As more electrons are removed the attractive forces increase due to decreasing shielding and an increase in the proton to electron ratio.
The increase in ionisation energy, however, is not constant and is dependent on the atom’s electronic configuration
The values become very large and difficult to represent meaningfully, so it is more convenient to show the logarithm of the ionisation energies.
This helps us to see significant jumps in ionisation energies.

The first electron removed has a low ionisation energy as it is easily removed from the atom due to the repulsion of the paired electrons in the 4s orbital.
The second electron is a little more difficult to remove than the first electron as you are removing an electron from a positively charged ion.
The third electron is much more difficult to remove than the second one corresponding to the fact that the third electron is in a shell that is closer to the nucleus (3p).
The graph shows there is a large increase in successive ionisation energy as the electrons are being removed from a increasingly positive ion.
The big jumps on the graph show the change of shell and the small jumps are the change of subshell

Question 41

IONISATION ENERGY GRAPH EXAM TIP:

Answer 41

Be careful with how you interpret successive ionisation energy graphs as it is common for students to read them the wrong way around and count outer electrons from right to left instead
of left to right so they get the jumps in the wrong place.
This happens particularly when you are given only a partial successive ionisation energy graph and
have to deduce which group the element comes from.
It’s a good idea if you see an ionisation energy graph in an exam question to label the shells and
subshells so you are less likely to make this mistake!

Question 42

What is metallic bonding?

Answer 42

The electrostatic attraction between the positive metal cations and the negatively charged seas of delocalised electrons

Question 43

Describe to me the process of metallic bonding?

Answer 43

In a giant metallic structure, each atom donates its valence electrons ton a sea of delocalised electrons, which are spread throughout the giant metal structure.

The positive metal cations consists of the nucleus and the inner shells of the metal atoms.

The cations are fixed in position by the strong electrostatic forces of attraction, maintaining the structure and shape of the metal. Whereas the electrons are delocalised (mobile) and can move throughout the structure.

Question 44

What are the 3 key properties of metals?

Answer 44

Strong metallic bonds
good electrical conductivity
high melting and boiling points

Question 45

When and how do metals conduct electricity?

Answer 45

Metals conduct electricity in solid and liquid states. Because the electrons are mobile they can carry the electrical charge throughout the structure.

Question 46

Why do metals have high melting and boiling points?

Why does the melting point and boiling point increase across the metals of a period?

Answer 46

The melting point depends on the strength of the metallic bonds, which depends how many electrons are lost and contributed to the sea of delocalised electrons, obviously the more electrons the stronger the nuclear charge between the electrons and the nucleus.

For most metals a great deal of energy is required to overcome the strong electrostatic forces of attraction between the metal cations and the sea of delocalised electrons, resulting in high melting and boiling points.

Question 47

Are metals soluble?

Answer 47

Metals don’t dissolve. Many interactions would lead to reaction as opposed to dissolving.

Question 48

What are the common properties of metals?

Answer 48

Metals are strong, hard, malleable, ductile, good conductors of heat and electricity.

Question 49

Name the 3 non metals that form giant covalent structures?

Answer 49

BORON,CARBON,SILICON

Question 50

Why do giant covalent compounds have high melting and boiling points?

Answer 50

As they have million of strong covalent bonds which, have a large bond enthalpy and require a great deal of energy to overcome.

Question 51

Are giant covalent compounds soluble?

Answer 51

They’re insoluble in almost all solvents. The covalent bonds holding together the atoms in the lattice are far too strong to be broken by interactions with solvents.

Question 52

Do giant covalent compounds conduct electricity?

Answer 52

They’re non-conductors of electricity Only exceptions are graphite and graphene.

In carbon (diamond) and silicon, all four valence electrons are involved in covalent bonding, so none are available for conducting electricity.

Question 53

What are the properties of diamond?

Answer 53

Diamond is a giant covalent lattice (or macromolecule) of carbon atoms
Each carbon is covalently bonded to four others in a tetrahedral arrangement with a bond angle of109.5
The result is a giant lattice structure with strong bonds in all directions
Diamond is the hardest substance known
For this reason, it is used in drills and glass-cutting tools

Question 53

What’re the properties of giant covalent compounds?

Answer 53

Giant covalent lattices have very high melting and boiling points.
These compounds have a large number of covalent bonds linking the whole structure.
A lot of energy is required to break the lattice. The compounds can be hard or soft.
Graphite is soft as the intermolecular forces between the carbon layers are weak.
Diamond and silicon(IV) oxide are hard as it is difficult to break their 3D network of strong covalent bonds.
Graphene is strong, flexible and transparent, which makes it potentially a very useful material.
Most compounds are insoluble with water
Most covalent substances do not conduct electricity
For example, diamond and silicon(IV) oxide do not conduct electricity as all four outer electrons on every carbon atom is involved in a covalent bond , so there are no free electrons available.
There are some covalent substances that are exceptions because they do conduct electricity.
Graphite has delocalised electrons between the carbon layers, which can move along the layers when a voltage is applied
Graphene is an excellent conductor of electricity due to the delocalised electrons.

Question 54

What are the properties of silicon dioxide?

Answer 54

Silicon(IV) oxide is also known as silicon dioxide,but you will be more familiar with it as the white
stuff on beaches!
Silicon(IV) oxide adopts the same structure as diamond - a giant covalent lattice
macromolecular structure made of tetrahedral units all bonded by strong covalent bonds
Each silicon is shared by four oxygens and each oxygen is shared byt wo silicons
This gives an empirical formula of SiO2

Question 55

What are the properties of graphite?

Answer 55

In graphite,each carbon atom is bonded to three others in a layered structure
The layers are made of hexagons with a bond angle of120
The spare electrons are delocalised and occupy the space between the layers
All atoms in the same layer are held together by strong covalent bonds
However, the layers are held together by weak intermolecular forces
These weak intermolecular forces allow the layers to slide over each other

Answer 56

Some substances contain an infinite lattice of covalentlybonded atoms in two dimensions only to form layers.
Graphene is an example
Graphene is made of a single layer of carbon atoms that are bonded together in a repeating pattern of hexagons
Graphene is one million times thinner than paper ;so thin that it is actually considered two dimensional