2.2.1 Electron structure

Question 1

Define orbital

Answer 1

region around a nucleus that can hold up to two electrons with opposite spins

Question 2

What is the shape of the s-orbital ?

Answer 2

spherical

Question 3

What is the shape of the p-orbital ?

Answer 3

dumbell shape

Question 4

How many orbitals in the s-sub shell ?
How many electrons ?

Answer 4

1 orbital
2 electrons

Question 5

How many orbitals in the p sub shell ?
How many electrons ?

Answer 5

3 orbitals
6 electrons

Question 6

How many orbitals in the d sub shell ?
How many electrons ?

Answer 6

5 orbitals
10 electrons

Question 7

How many orbitals in the f sub shell ?
How many electrons ?

Answer 7

7 orbitals
14 electrons

Question 8

What sub-shells make up the first shell ?
How many electrons are there ?

Answer 8

1s sub shell
Two electrons

Question 9

What sub-shells make up the 2nd shell ?
How many electrons are there ?

Answer 9

2s, 2p
8 electrons

Question 10

What sub-shells make up the 3rd shell ?
How many electrons are there ?

Answer 10

3p, 3p, 3d
18 electrons

Question 11

What sub-shells make up the 4th shell ?
How many electrons are there ?

Answer 11

4s, 4p, 4d, 4f
32 electrons

Question 12

What are the 2 rules for the filling of electron orbitals ?

Answer 12

• electrons fill orbitals in order of increasing energy
• orbitals (in the same sub shell) are** filled singularly **before pairing (Hunds rule) - Within a sub-shell all orbitals have the same energy

Question 13

How do you write electronic configurations ?

Answer 13

1S^2
* 1 = name of the main shell
* S = Type of sub-shell
* ^2 = number of electrons

Question 14

How is the electrons in boxes method used to represent orbitals ?

Answer 14

Box = the orbital
Arrows = two electrons with opposite spins
(Electrons in the same sub-shell will fill the boxes singularly before pairing)

Question 15

What is the rule for the 4s and 3d sub-shell ?

THIS IS A KEY EXCEPTION !!!

Answer 15

• electrons fill the 4s sub-shell before the 3d
• because 4s is lower in energy
  (Electrons will also be lost from the 4s sub-shell first)

Question 16

Write the electronic configuration for sodium ?

Answer 16

Na = 11 electrons
1S^2, 2S^2 2p^6, 3p^1
Sodium 1+ ion
1s^2, 2s^2 2p^6
- as one electron has been lost

Write out the electrons in boxes and make sure all the orbitals in each sub shell are filled **SINGULARLY **to prevent any electron repulsion

Question 17

Define transition metal

Answer 17

elements which form stable ions with an incomplete d-orbital
-Sc and Zn are not transition metals despite being in the d-orbital

Question 18

Why is Sc not a transition metal ?

Answer 18

• Forms a Sc 3+ ion
• so loses electrons from the 4s and 3d sub shell
• so forms an empty d-sub shell
• so cannot be a transition metal element

Question 19

Why is Zn not a transition metal ?

Answer 19

• Forms a Zn + ion
• Electrons are lost from the 4s orbital
• So Zn forms a full 3d orbital (10 electrons)
• so cannot be a transition metal element

Question 20

Define ionisation energy
Define second ionisation energy

Answer 20

• energy required to remove one mole of electrons from one mole of gaseous ions
• energy required to remove the second electron

Question 21

What 3 factors affect ionisation energy?
How do they increase it ?

Answer 21

• **Nuclear charge ** = more protons, increased charge, more nuclear attraction, greater ionisation energy
• **Atomic radius **= decreased radius, more nuclear attraction, greater ionisation energy
• Shielding = less electrons, outer electron is closer to nucleus, more nuclear attraction, greater ionisation energy

Question 22

How does ionisation energy change down the group ?

Answer 22

• Atomic radius increases
• Shielding increases
• weaker nuclear attraction
  * decreased ionisation energy (as its easier to lose the outer electron)

Question 23

How does ionisation energy change across the period ?

Answer 23

• Nuclear charge increases
• Atomic radius decreases
• Shielding is relatively constant (So when a new proton is added as you move along the period, nuclear attraction increases as the shielding effect of the electrons doesn’t change)
• So stronger nuclear attraction
  * increased ionisation energy