Atomic Structure

Question 1

Protons

Relative mass: ?

Relative charge: ?

Behaviour in electric field: ?

Answer 1

• Relative mass: 1
• Relative charge: +1
• Behaviour in electric field: deflected to negative plate (smaller deflection than electron due to greater mass)

Question 2

Neutrons

Relative mass: ?

Relative charge: ?

Behaviour in electric field: ?

Answer 2

• Relative mass: 1
• Relative charge: 0
• Behaviour in electric field: no deflection as no charge

Question 3

Electrons

Relative mass: ?

Relative charge: ?

Behaviour in electric field: ?

Answer 3

• Relative mass: 1/1836
• Relative charge: -1
• Behaviour in electric field: deflection to positive plate (large deflection due to small mass)

Question 4

Define

isotopes

Answer 4

Atoms of the same element/with the same atomic number/with the same number of protons but different masses/mass numbers/number of neutrons.

Question 5

What are principal energy levels/principal quantum shells?

Answer 5

These are regions at various distances from the nucleus that may contain up to a certain number of electrons (with similar energies and similar distances from the nucleus).

As the shell number increases, the energy of the electrons increases.

Question 6

What is a subshell?

Answer 6

A group of orbitals with the same energy level that differ in their orientation in space.

Each principal quantum shell contains a different number of subshells.

Question 7

What is an orbital?

Answer 7

An atomic orbital is a region of space around the nucleus where there is a maximum probabilty of electrons to be found.

Each orbital can contain up to two electrons. For orbitals with two electrons, the electrons have opposite spins two minimise repulsion.

When e^-s are placed in a subshell of orbitals of equal energy, they occupy them singly before pairing takes place.

Question 8

State the number of orbitals and the maximum number of electrons in these subshells:

s: ?
p: ?
d: ?

Answer 8

• s: 1 orbital, 2 electrons
• p: 3 orbitals, 6 electrons
• d: 5 orbitals, 10 electrons

Question 9

What are the shapes of the s and p orbitals?

Answer 9

Question 10

In what order are atomic orbitals filled?

Answer 10

Electrons will always occupy the lowest available energy level orbital first (Aufbau’s principle)

Question 11

State the exceptions to the expected pattern of electronic configuration

Answer 11

Chromium:

1s² 2s² 2p⁶ 3s² 3p⁶ 4s¹ 3d⁵

Copper:

1s² 2s² 2p⁶ 3s² 3p⁶ 4s¹ 3d¹⁰

Question 12

Define

ionisation energy

Answer 12

The minimum energy needed to remove 1 mole of e^-s from 1 mole of gaseous atoms to form 1 mole of unipositive ions.

Question 13

What 4 factors influence ionisation energy?

Answer 13

1. Distance from nucleus (atomic radius)
   
   • farther away means lower attraction
2. Nuclear charge
   
   • as nuclear charge increases, attraction increases (usually leading to a smaller radius, e.g. between isoelectronic ions of different elements)
3. Shielding
   
   • outer electrons are pushed away by full subshells of inner electrons
4. Spin-pair / Inter-electron repulsion
   
   • an electronic configuration having two electrons in an orbital is less stable than having one

Question 14

What trend in ionisation energy can be seen down a group?

Answer 14

Ionisation energy decreases down a group

• No. of shells increases, so atomic radius increases
• Larger number of shells means more inner electrons, so shielding increases
• Higher number of protons means higher attraction, but this factor is outweighed by shielding and distance from nucleus

Question 15

What trend in ionisation energy can be seen across a period?

Answer 15

Ionisation energy generally increases across a group

• Number of protons increases so positive nuclear charge and attraction increase
• No. of shells and shielding remain reasonably constant
• Higher attraction leads to smaller atomic radius

Question 16

Explain the first drop in ionisation energy when moving across periods 2 or 3.

(i.e. between Be and B or Mg and Al)

Answer 16

The electron to be removed is in a p subshell rather than an s subshell, and so is slightly further from the nucleus and has an additional subshell of inner electrons shielding it.

Question 17

Explain the second drop in ionisation energy when moving across periods 2 or 3.

(i.e. between N and O or P and S)

Answer 17

The previous element has half-subshell stability. The spin-pair repulsion in O and S leads to a less stable configuration, so the electron is more eaily removed.

Question 18

How can you interpret this information from a successive ionisation energy graph?

• Total no. of electrons
• No. of shells occupied
• No. of subshells occupied
• No. of electrons in each shell
• No. of electrons in each subshell

Answer 18

• Total no. of electrons: no. of plot points on graph
• No. of shells occupied: no. of sections between large jumps in energy
• No. of subshells occupied: no. of section between small jumps in energy
• No. of electrons in each shell: no. of electrons between two large jumps in energy
• No. of electrons in each subshell: no. of electrons between small jumps in energy

Question 19

Describe how ionic radius differs between positive and negative ions.

Answer 19

Ionic radius: describes the size of an ion

• Positive ion: smaller radius than original neutral atom because shell no. decreases, screening effect decreases but the attraction of nucleus increases.
• Negative ion: larger ionic radius than neutral atom because e^-s added while nuclear charge remains same
• Negative ions always larger than positive ions in the same period as they have one more shell