Chapter 7: Trends in the Periodic Table (Definitions)

Question 1

Define atomic radius(covalent radius) of an atom.

4

Answer 1

• half the distance between the nuclei of two atoms
• of the same element
• that are joined together
• by a single covalent bond

Question 2

Why does the values of the atomic radius increase down the groups in the Periodic Table?

(a) =3
(b) =4

Answer 2

(a) the additional electrons are going into a new energy level which is further from the nucleus
- since the outer electrons are becoming further away from the nucleus
- the atomic radius increases
(b) the screening effect of inner electrons
- the nuclear charge increases but it is lessened because of the screening effect
- the electrons in the inner energy level(s) help to screen the outer electrons from the positive charge of the nucleus
- thus, some of the attractive force of the positive charges in the nucleus is lessened by the inner energy levels of electrons
- therefore, even though the nuclear charge increases, the increase is counteracted by the shielding effect of the inner energy levels of electrons.

Question 3

Why does the values of the atomic radius decrease across a period in the Periodic Table?

(a) =4
(b) =4

Answer 3

(a) Increase in effective nuclear energy:
- number of protons in the nucleus increases from left to right across a period
- there is greater attractive force on the outer electrons
- increased nuclear charge draws the energy levels closer to the nucleus
- this attractive force increases from left to right across the period.
(b) No increase in the screening effect
- the extra electron, moving left to right across a period, is being added into the same outer energy level
- therefore, there is no increase in the screening effect
- there is no additional energy level of electrons being added to counteract the increasing attractive force of the nucleus
- the size of the atom shrinks

Question 4

Define first ionisation energy of an atom

3

Answer 4

• the minimum energy required to completely remove the most loosely bound electron
• from a neutral gaseous atom
• in its ground state.

Question 5

What is the Equation for first ionisation energy(sodium)

Answer 5

Na(g)->Na+(g)+e-

Question 6

Why does the values of the first ionisation energy decrease down a group in the Periodic Table

(a) =3
(b) =5

Answer 6

(a) Increasing atomic radius
- the atomic radius increases down any one group
- the outermost electrons are becoming further away from the attractive force of the nucleus
- therefore, it becomes easier to remove an electron from the outer energy level
(b) Screening effect of inner electrons
- the increase in nuclear charge is cancelled out
- by the corresponding screening effect of the
- intervening energy levels of electrons.
- the outermost electrons are shielded from the attractive force of the positively charged nucleus
- and so are easier to remove

Question 7

Why does the values of the first ionisation energy increase across a period in the Periodic Table

(a) =5
(b) =4

Answer 7

(a) Increasing effective nuclear charge
- from left to right across a period the number of protons in the nucleus is increasing
- this means that the attraction between the nucleus and the outer electrons is steadily increasing
- without any corresponding increase in screening by electrons
- therefore, the electrons are being held more firmly
- and so it requires more energy to remove one of the electrons from the outermost energy level
(b) Decreasing atomic radius
- the atomic radius decreases from left to right across a period
- this means that an electron in the outermost energy level is becoming closer to the nucleus
- therefore, it is more difficult to remove the electron
- due to the increased attraction between it and the nucleus

Question 8

Define second ionisation energy of an atom.

4

Answer 8

• the energy required to remove an electron
• from an ion
• with one positive charge
• in the gaseous state

Question 9

What is the equation for second ionisation energy

Answer 9

X+(g)->X2+(g)+e-

Question 10

Why does the values of electronegativity decrease down the groups in the Periodic Table

(a) =3
(b) =4

Answer 10

(a) Increasing atomic radius
- the atomic radius increases as you move down any group
- this means the outermost electrons are becoming further away from the attractive force of the nucleus
- therefore, there is a smaller attraction between the nucleus and the shared pair of electrons.
(b) Screening effect of inner electrons
- the increase in nuclear charge down a group is cancelled by the screening effect of the intervening energy levels of electrons
- this means that the outermost electrons are shielded from the attractive force of the positively charged nucleus
- since it is these outermost electrons that are involved in bonding
- the attraction of the nucleus for these electrons decreases going down the group.

Question 11

Why does the values of electronegativity increase across the periods in the Periodic Table

(a) =3
(b) =3

Answer 11

(a) Increasing effective nuclear charge
- the number of protons in the nucleus increases from left to right across a period
- this means that the attraction between the nucleus and the outer electrons is steadily increasing
- therefore the electrons involved in bonding are being more strongly attracted to the nucleus
(b) decreasing atomic radius
- the atomic radius decreases from left to right across a period
- this means that the electrons in the outermost level are becoming closer to the nucleus
- therefore there is a greater attraction between the nucleus and these outer electrons