2.2.1 Electron structure Flashcards
Define orbital
region around a nucleus that can hold up to two electrons with opposite spins
What is the shape of the s-orbital ?
spherical
What is the shape of the p-orbital ?
dumbell shape
How many orbitals in the s-sub shell ?
How many electrons ?
1 orbital
2 electrons
How many orbitals in the p sub shell ?
How many electrons ?
3 orbitals
6 electrons
How many orbitals in the d sub shell ?
How many electrons ?
5 orbitals
10 electrons
How many orbitals in the f sub shell ?
How many electrons ?
7 orbitals
14 electrons
What sub-shells make up the first shell ?
How many electrons are there ?
1s sub shell
Two electrons
What sub-shells make up the 2nd shell ?
How many electrons are there ?
2s, 2p
8 electrons
What sub-shells make up the 3rd shell ?
How many electrons are there ?
3p, 3p, 3d
18 electrons
What sub-shells make up the 4th shell ?
How many electrons are there ?
4s, 4p, 4d, 4f
32 electrons
What are the 2 rules for the filling of electron orbitals ?
- 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
How do you write electronic configurations ?
1S^2
* 1 = name of the main shell
* S = Type of sub-shell
* ^2 = number of electrons
How is the electrons in boxes method used to represent orbitals ?
Box = the orbital
Arrows = two electrons with opposite spins
(Electrons in the same sub-shell will fill the boxes singularly before pairing)
What is the rule for the 4s and 3d sub-shell ?
THIS IS A KEY EXCEPTION !!!
- 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)
Write the electronic configuration for sodium ?
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
Define transition metal
elements which form stable ions with an incomplete d-orbital
-Sc and Zn are not transition metals despite being in the d-orbital
Why is Sc not a transition metal ?
- 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
Why is Zn not a transition metal ?
- 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
Define ionisation energy
Define second ionisation energy
- energy required to remove one mole of electrons from one mole of gaseous ions
- energy required to remove the second electron
What 3 factors affect ionisation energy?
How do they increase it ?
- **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
How does ionisation energy change down the group ?
- Atomic radius increases
- Shielding increases
- weaker nuclear attraction
* decreased ionisation energy (as its easier to lose the outer electron)
How does ionisation energy change across the period ?
- 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
