Electron configuration and Ionisation Energy

atomic number

number of protons

mass number

the total number of protons and neutrons in the nucleus of an atom

relative atomic mass

the mean mass of one atom of an element relative to 1/12th the mass of a carbon-12 atom

relative molecular mass

the mean mass of a molecule relative use relative to 1/12th the mass of a carbon-12 atom

isotope

atoms of the same element with a different number of neutrons

electron congifuration notation

1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 4d, 4f

How many electrons can a 'S' orbital hold?

2

How many electrons can a 'P' orbital hold?

6

How many electrons can a 'D' orbital hold?

10

How many electrons can a 'F' orbital hold?

14

Heisenburg's Principle of Uncertainty

You cannot determine the position and momentum of an electron at the same time

The Aufbau Principle

Electrons enter the lowest available energy level

Pauli's Exclusion Principle

No two electrons can have the same 4 quantum numbers

Hund's Rule of Maximum Multiplicity

When in orbitals of equal energy, electrons will try to remain unpaired

Where is the S block?

groups 1 and 2

Where is the D block?

transition metals

Where is the P block?

groups 3-8

Where is the F block?

lanthanides and actinides

Ionisation Energy

The energy required to remove one mole of electrons from one mole of gaseous atoms to form one mole of gaseous positive ions

Notation for first ionisation energy (eg. Aluminium)

Al (g) -> Al+ (g) + e-

How does distance from the nucleus affect ionisation energy?

• when the electron is further from the nucleus there is less attraction between the opposite charges

• thus the electron is less strongly held and lost easier

• so the ionisation energy is lower

How does nuclear charge affect ionisation energy?

• the higher the overall nuclear charge, the greater the attraction between the outer electrons and the nucleus

• thus the outer electrons are more strongly held so they are harder to lose

• so the ionisation energy is higher

How does shielding affect ionisation energy

• the more orbitals between the outer electrons and the nucleus, the more shielded the outer electrons are

• thus the attraction between the nucleus and the outer electrons is less so they are more easily lost

• so the ionisation energy is lower

Why does the first ionisation energy of atoms generally increase across a period?

• nuclear charge increases by one each time

• extra electron goes into the same main energy level

• the electrons are held increasingly stronger due to increasing nuclear charge so they are harder to lose and ionisation energy increases

Why does the first ionisation energy of atoms generally decrease down a group?

• as you go down, number of orbitals increase so outer electrons gradually become more shielded from the nucleus

• thus they are not held as strongly so they are lost easier so ionisation energy decreases

Why is the 2nd IE of an atom always greater than the 1st?

• the atom becomes a positive ion so the nuclear charge increases

• thus there is a stronger force of attraction between the nucleus and outer electrons

• so the electrons are more strongly held and easier to lose so ionisation energy increases

Why does atom size decrease across a period?

• elements have the same number of orbitals so shielding remains mostly constant

• nuclear charge increases across a period as atoms have more protons and electrons

• so outer electrons are increasingly attracted to the nucleus causing the atom's radius to decrease

Why does atom size increase down a group?

• number of orbitals increases causing outer electrons to be more shielded

• thus attraction between outer electrons and the nucleus is weaker so the radius increases

Why are cations always smaller than the corresponding atoms?

• cations lose electrons so the ion is positively charged

• thus the outer electrons are more attracted to the nucleus so the ion's radius decreases

Why are anions always larger than the corresponding atoms?

• anions gain electrons so the ion is negatively charged

• thus the outer electrons are less attracted to the nucleus so the ion's radius increases

Why is the first ionisation energy of boron less than that of beryllium?

• beryllium has 2 electrons in the 2s orbital

• boron has 1 electron in the 2p orbital

• beryllium loses its outer electron easier as it is lost from a pair so the electrons are already repelling each other

Why is the first ionisation energy of oxygen less than that of nitrogen?

• nitrogen’s outer electron is by itself in the 2p orbital

• oxygen’s outer electron is in a pair in the 2p orbital

• so oxygen’s is lost easier as the electrons are already repelling each other

Why does helium have the highest 1st ionisation energy if all the elements?

• it has the highest nuclear charge in the 1s orbital so strongest attraction between electrons and the nucleus

• there are no orbitals between the outer electron and the nucleus so it is the least shielded

• so its outer electron is the hardest to lose

Group 2 to 3 drop in 1st ionisation energy (Period 3)

• Aluminium’s outer electron is in the 3p orbital

• Magnesium’s outer electron is in the 3s orbital

• so Aluminium’s outer electron is more shielded so less attracted to the nucleus

• so the 1st ionisation energy decreases across the period

Group 5 to 6 drop in 1st ionisation energy (Period 3)

• Phosphorus has 3 electrons in the 3p orbital

• Sulfur has 4 electrons in the 3p orbital

• Sulfur loses its outer electron easier as it is lost from a pair which are already repelling each other

• so the 1st ionisation energy decreases across the period

Electron configuration of Chromium (atomic number: 24)

[Ar] 4s1 3d5

• 4s and 3d orbitals filled without paired electrons so the atom is more stable

Electron configuration of Copper (atomic number: 29)

[Ar] 4s1 3d10

• only 1 electron in 4s orbital and fills 3d orbital completely so atom is more stable