Structure 1.3 Electron configurations

Radio waves

Long wavelength

Low frequency

Low energy

Gamma rays

Short wavelength

High frequency

High energy

Visible light: red light

Long wavelength (700nm)

Low frequency

Low energy

Visible light: violet light

Short wavelength (400nm)

High frequency

High energy

Continuous spectrum

Shows all wavelengths of visible light

Absorption line spectrum

Black lines on a colored background (certain wavelengths missing).

Emission line spectrum

Coloured lines on a black background (only certain wavelengths visible)

Hydrogen emission spectrum

656 nm (red line): transition from n=3 to n=2

486 nm (turquoise): transition from n=4 to n=2

434 nm (blue line): transition from n=5 to n=2

410 nm (violet line): transition from n=6 to n=2

Absorbing photons of energy

Electrons are promoted to higher energy levels (excited state)

Emitting photons of energy

Electrons transition to lower energy levels

Electron transitions to n=1

Emit UV radiation (electron transitions emitting the greatest amount of energy)

Electron transitions to n=2

Emit visible light

Electron transitions to n=3

Emit infrared radiation (electron transitions emitting the greatest amount of energy)

Atomic orbitals

A region of space where there is a high probability of finding an electron.

S orbitals

Spherical

P orbitals

Dumbbell shaped. Px, Py and Pz are at 90 degrees to each other.

Sublevels organized in order of energy

Within a main energy level the order of energy is:

s < p < d < f

Electron configuration

A representation of the arrangement of electrons within an atom or ion.

E.g helium: 1s²

1 = main energy level

s = sub-level

2 = number of electrons in sub-level

Electron configuration for carbon, C

1s² 2s² 2p² or [He] 2s² 2p²

Electron configuration for sodium, Na

1s² 2s² 2p⁶ 3s¹ or [Ne] 3s¹

Condensed/abbreviated electron configurations

Use the symbol of a noble gas to represent the core electrons

[He] = 1s²

[Ne] = 1s² 2s² 2p⁶

[Ar] = 1s² 2s² 2p⁶ 3s² 3p⁶

The Aufbau principle

Used to determine the electron config. of an atom or ion.

Electrons fill lowest energy sub-levels first, 1s 2s 2p 3s 3p 4s 3d 4p etc.

Pauli exclusion principle

An atomic orbital can hold a maximum of two electrons and they must have opposite spins.

Hund’s rule

Degenerate orbitals (orbitals with the same energy) are filled singly with the same spin before being doubly occupied.

Electron configurations 4s 3d sublevels

K = [Ar] 4s¹

Ca = [Ar] 4s²

Sc= [Ar] 4s² 3d¹

Ti = [Ar] 4s²

V = [Ar] 4s² 3d³

Cr = [Ar] 4s¹ 3d⁵

Mn = [Ar] 4s² 3d⁵

Fe = [Ar] 4s² 3d⁶

Co = [Ar] 4s² 3d⁷

Ni = [Ar] 4s² 3d⁸

Cu = [Ar] 4s¹ 3d¹⁰

Zn = [Ar] 4s² 3d¹⁰

Positive ions (cations)

1. Are formed when atoms lose electrons.

2. The electrons are lost from the highest energy level first.

When transition elements form cations, they lose their 4s electrons first.

Negative ions (anions)

Are formed when atoms gain electrons.

Convergence limit

The frequency at which the spectral lines converge.

The energy levels converge when n=∞

The spectral lines for each series (UV, visible light and IR) also converge at higher energy.

Ionisation energy

The energy required for the electron transition from n=1 to n=∞

Can be calculated using the frequency (or wavelength) of the convergence limit.

E = hv

energy (J) = Planck’s constant (6.63 × 10⁻³⁴ J*s) x frequency (s⁻¹)

c = vλ

speed of light (3.00 × 10⁸ m s⁻¹) = frequency (s⁻¹) x wavelength (m)

Successive ionisation energies

The ionisation energy increases as electrons are removed from an increasingly positive ion → causing an increase in attraction between the nucleus and the remaining electrons.

Big ionisation energy caused by:

1. Different energy levels

2. Doubly/singly filled electron orbitals