The periodic table and periodicity

Question 1

How are the elements ordered in the periodic table

Answer 1

The elements are ordered by increasing atomic number, in other words number of protons.

Question 2

What is a group in the periodic table

Answer 2

A vertical row which consist of a number of elements with the same number of electron in the outer shell.

All the elements in the similar group have similar properties.

Question 3

What is a period in the periodic table

Answer 3

The period is a number that tells us the number of the highest energy electron shell for the elements in that period

Question 4

Key ideas of the periodic table

Answer 4

There is a repeating trend in the property of the elements across a period. (Metals on the left and non-metal on the right)

Between the metal and non-metal are metalloids which have properties of both metals and non-metals

Question 5

What is a pattern of reacting trends called

Answer 5

Periodicity

Question 6

Key idea of periodic trends in electron configuration

Answer 6

The filling of electron subshells follows a periodic pattern. Within the highest energy shell, the s sub shell fills before the p subshell.

Each block in the periodic table is named after the highest block containing the highest energy electron for the elements in that block

The 4s subshell first before the 3d subshell

Question 7

Key idea of atomic radius

Answer 7

We cannot measure the radius of an atom directly as electron clouds do not have a clearer cut-off point

However, one way of calculating atomic radius is to look at pairs of identical atoms that have formed a bond. We take the atomic radius as half the distance between the nuclei of the two atoms.

Question 8

Explain the trends of the atomic radius as we move across and down the period

Answer 8

The atomic radius decreases as we move across a period from left to right. (Each elements gains one more proton in its nucleus than the element before, positive charge increases thus there is an increased attraction between the nucleus and the electrons. This draws the electron causing the atomic Addis to decrease across the period.)

Outer shell electron are partially shielded from the attraction of the nucleus by electrons in the inner shells. However,all elements in period 2 have one inner electron shell. Thus shielding due to the inner electron shell is the same across the period.

Atomic radius increases moving down the group. The number of electron shell increases as we go down the group so the outer electron shell is further from the nucleus. Each element has one more full inner electron shell. This increases the amount of shielding between the nucleus and the outer electron. There is less attraction between the outer electron and the nucleus.

Question 9

Explain the trend in first ionisation energy across a period

Answer 9

It tends to increases as we move across a period. The positive charges in the nucleus increases as the number of protons increases. This increases the attraction between the nucleus and the electrons. The atomic radius deceases across a period

Both the increases nuclear charge and decreased nucleus radius means that the outer electrons are more attracted to the nucleus. This causes the first ionisation energy to increases across a period. (In all these elements, one electron is removed from a shell depending on the group therefore te shielding effect due to the inner electron shell is the same for each element)

However, some elements such as Boron and Oxygen may not fit the pattern of increasing first ionisation energy in group 2. The element Boron has an outer electron on the 2p subshell which has more energy. This means it takes less energy to remove the outer electron of Boron compared to Beryllium. Thus has a lower first ionisation energy.

It will then increase but then fall again. In this case in group 2 the element is oxygen. The 2p subshell will have two electron in one orbital and one in the other two. These electron will repel each other thus will take less energy to remove one of these electrons than if they wear separated. Therefore, the first ionisation energy of oxygen is less than nitrogen in group 2.

The same concept applies to other group. However, depending on the group different electron shells are affected (group 2 and second electron shell, group 3 and third electron shell)

Question 10

Define first ionisation energy

Answer 10

The first ionisation energy is the energy needed to remove one mole of electron from one mole of atoms in their gaseous state to form one mole of 1+ ions (also in their gaseous state)

Question 11

Explain the trend in first ionisation energy moving down a group

Answer 11

The first ionisation in energies decreases moving down a group.This is due to two factors:

The atomic radius increases as we move down the group. The outer electron shell is further away from the nucleus.

Going down a group the number of internal electron shells also increases. This means that there is more shielding between the nucleus and the outer electrons.

Both these factors means that the attraction between the nucleus and the outer electrons decreases. This causes the first ionisation energy to fall.

The nuclear charge increases a we move down a group. However, the effect is offset by two factors that we’ve just looked at.

Question 12

Explain the successive ionisation energies of oxygen

Answer 12

There is an increase of ionisation energy as we remove the first six electrons. As we remove an outer electron the remaining electrons in the outer shell are pulled slightly closer to the nucleus. This means that there is a greater attraction between the outer electrons and the nucleus and this causes the ionisation energy to gradually increase.

There is a massive increase in ionisation energy as we remove the seventh electron. This is because the seventh electron shell is on the inner electron shell a closer to the nucleus. Additionally, electrons in the first shell experience much less shielding. This means that electron in the first shell have a greater attraction to the nucleus compared to the electrons in the second shell.

Question 13

How can you find an element using successive ionisation energies of a graph

Answer 13

Observe the graph and compare the gradual and massive increase in ionisation energy.

You can identify the number of electron in the outer shell by counting the successive amount of ionisation which goes up in energy gradually. This can be used to determine the group

Therefore, the other ionisation must be in an inner shell. This can be used to determine the period

Question 14

How is a metallic lattice formed

Answer 14

In metals the electron in the outer shell are delocalised.The metals in group 1 would donate an electron to become a +1 cation. the negatively delocalised electron are attracted to the positive cations by electrostatic forces of attraction. This is called metallic boding (metallic lattice)

Question 15

Key features of a giant metallic lattice

Answer 15

The cations are fixed in place and cannot move.

The delocalised electrons are free to move. Therefore, are able to be good conductors of electricity when they are both solid or liquid. The delocalised electron are attracted to the positive pole and move towards it. The delocalised electrons act as mobile charge carriers enabling metals to conduct electricity.

Most metals have relatively high melting and boiling points. This is due to the strength of the metallic bond (delocalised electron have strong electrostatic forces of attraction to the cations) . Lots of energy is required to overcome this reaction. This means that metal generally have a high m.p/b.p

Metals do not dissolve. When added to water the metal usually reacts to water rather than dissolving

Question 16

How is a giant covalent structure formed

Answer 16

Billions of atoms are joined by strong covalent bonds. Together they from a giant covalent lattice.

Question 17

Key properties of giant covalent lattices

Answer 17

It takes a great deal of energy to break all of the covalent bonds in a giant covalent lattice. Therefore, all giant covalent structures have a high melting and boiling point.

Giant covalent lattices do not conduct electricity. This is because every electron is in the covalent bond.There are no delocalised electrons to act as charge carriers enabling metals

Giant covalent lattices are insoluble. This is because solvents cannot disrupt the large number of strong covalent bonds.

Question 18

Explain the structure of diamond

Explain the structure of graphite

Answer 18

Diamond:
Carbon atom covalently bond to four other carbon atoms. This forms a giant covalent lattice.

The atoms are arranged in a tetrahedral structure with bond angles of 109.5

graphite:
Carbon atoms covalently bond to three other carbon atoms. This forms graphite.

The carbon atoms from layers of planar hexagonal structures with a bond angle of 120

Question 19

Explain the properties of graphite

Answer 19

One electron from each carbon atom is delocalised and acts as a mobile charge carrier. Therefore, because of these delocalised electrons, graphite is a good conductor of electricity.

Question 20

Explain the properties of graphene

Answer 20

A single layer of graphite (graphene) because of its delocalised electron it is also a good conductor of electricity.

Question 21

Explain what would happen if a simple molecular structure was cooled

Answer 21

Diatomic molecules and noble gases when cooled to a solid would form a simple molecular lattice which have weak intermolecular forces between the molecules. These intermolecular forces do not require a lot of energy to break which is why these elements all have a low m.p/b.p

Question 22

Describe the trends in reactivity of group 2 elements down the group

Answer 22

All elements in group 2 have two electrons in the s sub-shell of their outer shell. When group 2 elements react , these two outer electron are lost,. By transferring these two electrons to other species group 2 metals act as reducing agents

The reactivity of metals increases as we move down group 2. When group 2 elements react they lose their outer two electrons. This requires ionisation energy.

For e.g: Both the second and first ionisation energies for calcium are less than for magnesium. This is because calcium has a greater atomic radius than magnesium so the two outer electrons in calcium are further from the nucleus compared to magnesium. Calcium atoms have one more full inner electron shell than magnesium. The full inner electron shells partially screen the outer electrons from the positive charge of the nucleus. Therefore , the outer electrons in calcium are less attracted to the nucleus than the outer electron in magnesium. Calcium takes less energy to remove outer electrons. Therefore, calcium is more reactive than magnesium.

Question 23

Key idea of elements in group 2

Answer 23

Alkaline earth metals are reactive