periodic table

Question 1

What is a group?

Answer 1

The vertical columns in the Periodic Table

Question 2

What are periods?

Answer 2

The horizontal columns of the Periodic Table

Question 3

What does the s-block consist of?

Answer 3

Elements in Groups 1 and 2. S-block elements have highest energy electron in an s orbital

Question 4

What does the p-block consist of?

Answer 4

Elements in Groups 3 to 8. P-block elements have highest energy electron in a p orbital

Question 5

What does the d-block consist of?

Answer 5

Elements Sc to Zn in period 4 and Y to Cd in period 5. D-block element has its d subshell filled

Question 6

What is periodicity?

Answer 6

A regularly repeating pattern of atomic, physical and chemical properties with increasing atomic number

Question 7

What are examples of periodicity?

Answer 7

Regular repeating pattern of electronic configurations, trends in atomic radii, trends in melting and boiling temperatures, trends in first ionisation energies

Question 8

What is atomic radii of an element?

Answer 8

The distance from the centre of the nucleus to the boundary of the electron cloud. We can find the atomic radius by determining the distance between the two nuclei and dividing by two

Question 9

What are the types of radius?

Answer 9

Covalent radius
Van der Waals radius (only radius that can be determined for neon and argon)
Metallic radius (used for metals)

Question 10

What is the trend of atomic radii across periods? Why?

Answer 10

Radius decreases across each period. As the number of protons in the nucleus increases, so does the nuclear charge, resulting in an increase in attractive force between the nucleus and the outer electrons. The increase in attractive force offsets the increase in electron-electron repulsion as the number of electrons in the outer quantum shell increases

Question 11

Why is the first ionisation energy of helium greater than hydrogen?

Answer 11

There is an increase in nuclear charge from hydrogen (1) to helium (2) as an extra proton is added. This increase in nuclear charge offsets the increase in electron-electron repulsion in the 1s orbital as a second electron is added

Question 12

Why does boron (a group 3 element) have a lower ionisation energy than beryllium (a group 2 element)?

Answer 12

Although the nuclear charge of boron is greater than beryllium, the outer electron in boron is in the 2p orbital rather than the 2s orbital (like beryllium) so the energy required to remove this electron is less. Furthermore, the 2p electron in boron experiences greater electron-electron repulsion (greater shielding) because there are two inner electron sub-shells as opposed to only one in beryllium

Question 13

Why does aluminium (a group 3 element) have a lower ionisation than magnesium (a group 2 element)?

Answer 13

Although the nuclear charge of aluminium is greater than magnesium, the outer electron in aluminium is in the 3p orbital rather than the 3s orbital (like magnesium) so the energy required to remove this electron is less. Furthermore, the 3p electron in aluminium experiences greater electron-electron repulsion (greater shielding) because there are 4 inner electron sub-shells as opposed to only 3 in magnesium

Question 14

Why does oxygen (a group 6 element) have a lower ionisation energy than nitrogen (a group 5 element)?

Answer 14

N: 1s2, 2s2, 2p1, 2p1, 2p1
O: 1s2, 2s2, 2p2, 2p1, 2p1
The first electron removed from oxygen is one of the two paired electrons in the 2p orbital. The presence of two electrons in a single orbital increases the electron-electron repulsion in the orbital so less energy is required to remove oxygen’s 2p2 electron than nitrogen’s 2p1 electron, despite the larger nuclear charge of the oxygen atom