Atomic Structure

Question 1

What are the relative masses of protons, neutrons and electrons?

Answer 1

Proton = 1
Neutron = 1
Electron = 1/1836

Question 2

Why are electrons easily deflected?

Answer 2

• negative charge

- very small mass

Question 3

In the cathode-ray tube experiment, why does proton need huge voltages for it to work?

Answer 3

Protons are much heavier than electrons, thus requiring much higher voltages to show the deflection of the proton.

Question 4

What does an atom contain of?

Answer 4

• protons and neutrons in the nucleus

- electrons arranged in shells and subshells

Question 5

What are isotopes?

Answer 5

Atoms of the same element that have different masses.

OR

Atoms with the same number of protons but different number of neutrons.

Question 6

Define the first ionisation energy?

Answer 6

The energy needed to remove one mole of electrons from one mole of gaseous atoms to form one mole of ions with a single positive charge.

Question 7

How are electrons held?

Answer 7

Electrons are held by the positive charge on the nucleus, thus electrons closer to the nucleus are held more strongly than those in shells further away.

Question 8

Why does it take less energy to remove an electron from atoms further down the group?

Answer 8

• further away from the nucleus
• more shielding, more inner shells
• increase in positive charge on the nucleus has less effect than the increase of shielding.

Question 9

What happens when you remove more electrons?

Answer 9

More energy needed to remove more electron

Question 10

Why is each successive ionisation energy bigger than the last?

Answer 10

There are fewer electrons (less negative charge) but positive charge on the nucleus is the same, thus, the pull on each electron is greater.

Question 11

Why does the electrons closer to the nucleus need much more energy to be removed?

Answer 11

Less shielded by inner electrons from the charge on the nucleus.

Question 12

What are the factors influencing ionisation energies?

Answer 12

1. Size of the nuclear charge (i.e increases when proton no. Increases)
2. Distance of outer electrons from nucleus (further, lower i.e)
3. Shielding effect of inner electrons (inner shells shields outer shells from attraction force of nucleus)

Question 13

What does an orbital represent?

Answer 13

Space in which fast-moving electrons are most likely to be found.

Question 14

What are the electronic configurations of chromium and copper?

Answer 14

Chromium: [Ar] 3d5 4s1

(Half filled and fully filled orbitals are more stable)

Copper: [Ar] 3d9 4s2

Question 15

Why does the 1st I.E across the period generally increases?

Answer 15

• nuclear charge increases across the period
• distance between nucleus and outer electrons remains reasonably constant
• shielding effect by the inner shells will also be reasonably constant.

Question 16

Why is there rapid decrease between the last element one period and the first element in the next period?

Answer 16

• Distance between nucleus and outer electron increase.
• Shielding effect by inner shells increase.
• factors outweigh increase of nuclear charge

Question 17

Be & B or Mg & Al, why are they almost the same?

Answer 17

• distance between nucleus and outer electron SLIGHTLY increase
• shielding effect by inner shells also increases SLIGHTLY
• outweigh increase of nuclear charge

Question 18

Whats the pattern in I.E down the group?

Answer 18

Decreases down the group.
- nuclear energy increases, less attraction between outer electron and the nucleus because
~ distance between nucleus and outer electron increases
~ shielding by complete inner shells increases
~ outweigh increase of nuclear charge.