· Learn the order R,M,I,V,U,X,G

1. Why do E-M waves become more dangerous as their frequencies increase?

2. The speed of light is 300,000, 000 m/s. What is the wavelength of a typical:

a. X-ray with a frequency of 6 x 10¹⁸ Hz

b. Microwave with a frequency of 3 x 10⁸ Hz

3. The speed of light is 300,000, 000 m/s. What is the frequency of a typical:

a. Infrared wave with a wavelength of 5 x 10^-2 Hz

b. Gamma wave with a wavelength of 3 x 10^-12 Hz

· A continuous spectrum contains all the frequencies of visible light without gaps

· Energised elements give out light in Emission Spectra. But they give out only a few frequencies so the spectra is mostly black.

· If white light hits the same element it absorbs only a few frequencies , giving an almost continuous Absorption Spectrum that is the opposite of the Emission Spectrum.

Using Emission spectra can allow a researcher to identify an unknown element from the frequencies of light it emits when subjected to high voltage.

So what?

Line Spectra are important evidence for shells (energy levels).

The theory relies on understanding the idea of Quantisation of Energy.

This states that light energy exists as units called photons (rather than waves).

The energy of the photon is given by Planck’s Equation

h – is the Planck Constant (you don’t need to learn it)

{Since, c=νλ we can write ν= c/λ}

When energy is given to an atom its electrons will be in the Ground State but will be raised to a higher energy level (the electron is “excited”).

When it loses energy the electron emits the energy as a photon with energy exactly equal to the difference in the Ground State and the excited state.

And because there are only a limited number of energy states to fall from, there are only a few different colours of light that can be given out.

Lyman, Balmer, Paschen.

· Energy levels converge (get closer as they get further from the nucleus.

Falling back to n=1 is always a bigger drop in energy than dropping back to any other level.

The Aufbau principle – says you must fill the lowest energy orbital before putting any electrons in a higher energy orbital.

So you can’t put any electrons in the 2s orbital until the 1 s orbital is full, and you can’t put any electrons in a 2p orbital until the 2s orbital is full

The Pauli Exclusion principle – says that every orbital holds a maximum of two electrons, and that these must have opposite spin.

Hund’s rule – says that when there are orbitals at the same energy they must be singly filled first.

Fill in the electron box models for ₁H to ₁₈Ar

Fill the electron boxes for elements 19 to 36

The electron configuration for Sodium is 1s² 2s² 2p⁶ 3s¹

The electron configuration for Aluminium is 1s² 2s² 2p⁶ 3s² 3p¹

Writing the whole configuration each time is a pain. So we sometimes write condensed configurations based on the nearest Nobel Gas and simply write down the incomplete shell.

So, the electron configuration for Neon is 1s² 2s² 2p⁶

We can write Sodium as [Ne] 3s¹ and Aluminium as [Ne] 3s² 3p¹

Write condensed electron configurations for:

1. Mg

2. Mg²⁺

3. Sc

4. Zn

5. S

6. F^-

7. O^2-

8. Na⁺

9. Al³⁺

10. P^3-

Everything about 4s and 3d orbitals is a bit of a pain.

LEARN – the first electron Transition Elements lose are always the 4s electrons

So, Iron is Fe = [Ar] 3d⁶ 4s²

But Fe²⁺ = [Ar] 3d⁶ 4s⁰ or just [Ar] 3d⁶

And Fe³⁺ = [Ar] 3d⁵ 4s⁰ or just [Ar] 3d⁵

Write Condensed electron configurations for:

1. Sc ³⁺

2. Zn²⁺

3. Cr³⁺

4. Mn²⁺

5. Cu⁺

6. Cu²⁺