Chapter 1: Atomic Structure

Question 1

Define isotopes

Answer 1

Isotopes are atoms of the same element which have the same proton number but different nucleon number.

Question 2

Define an isotope

Answer 2

An isotope is an atom of an element which has the same proton number as other atoms of the same element but it has a different nucleon number.

Question 3

Describe the behaviour of beams of protons, neutrons and electrons moving at the same velocity in an electric field.

Answer 3

Neutrons continue straight.
Electrons are deflected on a curved path to the positive plate.
Protons are deflected on a curved path to the negative plate.
Electrons have greater angle of deflection than a proton due to smaller particle size of electron.

Question 4

Why do isotopes have the same chemical properties but different physical properties?

Answer 4

Isotopes have the same number of electrons hence they have the same chemical properties. They have different mass and density due to different nucleon number, so they have different physical properties.

Question 5

Define relative atomic mass.

Answer 5

The relative atomic mass of an element, Ar, is the average mass of its atoms in a naturally occuring sample compared to 1/12 the mass of a carbon 12 atom, where 12C is 12.0 exactly.

Question 6

Define relative molecular mass.

Answer 6

The relative molecular mass, Mr, of an element or covalent compound is the average mass of its molecules compared to 1/12 the mass of a carbon 12 atom, where 12C is 12.0 exactly.

Question 7

Describe the trend of Atomic radius across each period.

Answer 7

Generally decrease across each period. This is because the proton number increases across each period (increased positive nuclear charge) and at the same time extra electrons are added to the principal quantum shell. Shielding effect is relatively the same. Increased nuclear charge results in increased pull of nuclei on the electrons causing smaller atomic radius.

Question 8

Describe the trend of Atomic radius down each group.

Answer 8

Generally increase down the group. This is because electrons in inner shells repel valence electrons, shielding them from positive nuclear charge. There is a weaker pull of nuclei on the electrons resulting in larger atomic radius.

Question 9

Describe the trend of Ionic radius with increasing negative charge.

Answer 9

Ionic radius increases with increasing negative charge.
Ions with negative charges are caused by atoms accepting extra electrons while the nuclear charge remains the same. Outermost electrons are further away from positively charged nucleus so they are held weakly to the nucleus causing increased ionic radius.

Question 10

Describe the trend of Ionic radius with increasing positive charge.

Answer 10

Ionic radius decreases with increasing positive charge. Positively charged ions are formed by atoms losing electrons with the same nuclear charge. The remaining electrons undergo a greater electrostatic force of attraction to the nucleus which decreases the ionic radius.

Question 11

How to calculate the Relative Atomic mass of an atom using relative abundance of its isotopes?

Answer 11

( relative abundance 1 x mass 1 ) + ( relative abundance 2 x mass 2 ) / 100

Question 12

What is the full electronic configuration of Chromium? (24 electrons)

Answer 12

1s2 2s2 2p6 3s2 3p6 [3d5 4s1]

Question 13

What is the full electronic configuration of Copper? (29 electrons)

Answer 13

1s2 2s2 2p6 3s2 3p6 [3d10 4s1]

Question 14

Define the 1st Ionisation energy.

Answer 14

The 1st Ionisation Energy is the energy required to remove one mol of electrons from one mol of gaseous atoms to form one mol of gaseous ions.
eg.
Na (g) -> Na+ (g) + e-

Question 15

Describe the trend of Ionisation energy across the period.

Answer 15

Across the period Ionisation energy increases. This is because the nuclear charge increases across the period hence the outer shell of atoms are pulled closer to the nucleus causing atomic radius to decrease. The distance of the outer electrons from the nucleus decreases. Becomes harder to remove an electron as you move across a period, more energy is needed hence IE increases.

Question 16

Describe the trend of Ionisation energy down the group.

Answer 16

Across the period the Ionisation energy decreases. Nuclear charge increases down the group because proton number increases.
Atomic radius increases as the number of electron shells increases. Hence, the distance of the outer electrons from the nucleus increases. The shielding effect of inner shell electrons increases as there are more electron shells. These factors outweigh the increased nuclear charge. It becomes easier to remove an electron down the group.

Question 17

Describe the trend of successive ionisation energies of an element.

Answer 17

The successive ionisation energies of an element increases. This is because removing an electron from a positive ion is more difficult than from a neutral atom.
As more electrons are removed the attractive forces increase because of decreasing shielding effect and increase in proton to electron ratio.
The increase in IE, however, is not constant and may change depending on the atom’s electronic configuration.