1. Electronic Structure of the Atom

2. Energy levels
An energy level is a region of definite space that an electron can occupy
Energy levels are not the same things as shells
Each shell has between 1 – 4 energy sub-levels

3. Atoms like to be stable so electrons enter the lowest energy level possible, called the ground state.
When the atoms gain energy they jump to higher energy levels, called the excited state.
This makes the atom unstable so the electrons fall back to their ground state and emit the difference in energy as a photon of light.

4. Can see this light if it falls in the visible spectrum.
As all elements have different numbers of electrons, they will have different excited states, and therefore produce different photons of light.
So each element emits its own unique spectrum
These spectra can be used to identify what elements are present in unknown samples.

5. Spectra
White light breaks up into a continuous spectrum when passed through a prism
Red, Orange, Yellow, Green, Blue, Indigo, Violet
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6. 1) Emission Line Spectrum:
An electric current is passed through a gaseous sample of an element in a discharge tube.
The light produced in the discharge tube is passed through a prism and forms a line spectrum.

7. Called an emission line spectrum because the lines are formed from the wavelengths of light emitted by the element.
Instrument = Spectrometer.
Each element has its own unique spectrum so they can be used as an identification tool.

8. The visible line spectrum produced by hydrogen is called the Balmer series.

9. Examples:
- Sodium street lamps are yellow.
- Neon signs are red.
- Fireworks.

10. 2) Atomic Absorption Spectrum:
White light is passed through a gaseous sample of an element.
The light that passes through the sample is passed through a prism and forms a continuous spectrum with certain wavelengths missing from it.

11. Called an atomic absorption spectrum because the missing wavelengths in the spectrum are the wavelengths of light absorbed by the element.
Instrument = Atomic absorption spectrometer.
Used to detect both the presence and concentration of an element in an unknown substance, e.g., heavy metal detection.

12. The emission line spectrum of an element is like the photographic negative of the atomic absorption spectrum.
The wavelengths of light absorbed by an element in its ground state are the same as those emitted by the element in its excited state.

15. Bohr’s Theory States that –
1. Electrons revolve around the nucleus in fixed paths called shells/orbits.
2. Electrons in any one orbit have a fixed amount of energy.
3. Once an electron remains in a particular energy level it neither gains nor loses energy.

16. 4. When an atom absorbs energy the electrons jump to higher energy levels (but are unstable so don’t stay for long).
5. Energy is lost when an electron falls back to a lower energy level (electrons can only fall to a certain energy level so only a fixed amount of energy can be released).

17. So when an atom is in its ground state all its electrons are in the lowest energy levels possible.
When this atom receives some energy (electrically or by heating) the electrons gain energy so jump to higher energy levels and are now in their excited state.

18. They do not remain in this state for long as they are unstable.

19. The electrons then fall back to a lower energy level releasing a fixed amount of energy.
This amount of energy is equal to the difference in energy between the two energy levels.
E2 – E1 = hf where E2 – E1 = difference in energy levels
h = Plancks constant
f = frequency of light

20. Since definite amounts of energy are released electrons must occupy definite energy levels - so energy levels must exist in atoms.
Each element has its own unique number of electrons - so will have different amounts and types of electron transitions - so will have its own unique spectrum.

21. Atomic orbitals
An atomic orbital is a region in space where there is a high probability of finding an electron.
You cannot know for sure where an electron is – explained by Heisenberg’s Uncertainty principle & de Broglie’s wave nature

22. Heisenberg’s Uncertainty Principle states that you cannot measure the position and velocity of an electron at the same time.
de Broglie states that electrons have a wave nature (like light) so have the properties of both waves and particles

23. There are 3 types of orbitals – s, p and d
The order of orbitals is:
1s,2s,2px,2py,2pz,3s,3px,3py,3pz,4s,3d,4px,4py,4pz
The 4s orbital is lower in energy than the 3d orbital, so it comes first.

24. Electrons fill the orbitals singly and then in pairs.
The maximum number of electrons in the s-orbital is 2, p-orbital is 6 and d-orbital is 10.
Each p and d sub-orbital can hold a maximum of 2 electrons.
Electron configuration represents the arrangement of the electrons in the atom.

25. 2 exceptions – chromium & copper
Cr =
4s1, 3d5 is a more stable configuration because the outer shell is half full
Cu =
4s1, 3d10 is a more stable configuration because the outer shell is full

26. Electron configuration of ions changes.
If you lose an electron you gain a positive charge (positive ion)
If you gain an electron you gain a negative charge (negative ion)

27. Stability of electrons is important for configuration
Full orbital is most stable (full outer shell)
Half orbital is next most stable

28. Trends in the Periodic Table
Have to know 3 trends:
1) Atomic radii
2) Ionisation energy
3) Electronegativity

29. Atomic radii
The atomic radius of an atom is half the distance between the nuclei of two atoms of the same element joined by a single covalent bond.

30. Decreases across a period because:
(a) Increase in nuclear charge –
Protons being added to the nucleus but electrons are being added to the same shell so pull of the nucleus is stronger.
(b) No screening effect –
Electrons are all added to the same shell so there are no inner shells to screen them from the nuclear charge.

31. Increases down a group because:
Decrease in nuclear charge –
Electrons are added to a new shell so the effect of the nucleus is not as strong.
(b) Screening effect –
Electrons in inner shells help shield outer electrons from the nuclear charge.

32. Ionisation energies
First Ionisation Energy of an atom is the minimum amount of energy required to remove the first most loosely bound electron from a neutral gaseous atom.
Increases across a period:
(a) Increase in nuclear charge –
Protons being added to the nucleus but
electrons are being added to the same shell so pull of the nucleus is stronger.

33. (b) Decrease in atomic radius –
Electrons are all added to the same shell so they are closer to the nucleus so it’s harder to remove the outer electron.
(c) No screening effect –
Electrons are all added to the same shell so there are no inner shells to screen them from the nuclear charge.

34. Decreases down a group:
(a) Increase in atomic radius –
Electrons are added to a new shell so the effect of the nucleus is not as strong so it’s easier to remove the outer electron.
(b) Screening effect –
Electrons in the inner shells help shield the outer electrons from the nuclear charge so it’s easier to remove the outer electron.

35. (c) Decrease in nuclear charge –
Electrons are added to a new shell so the effect of the nucleus is not as strong so it’s easier to remove the outer electron.
Ionisation energies are proof of energy levels because if all electrons occupied the same energy level they would require the same amount of energy to remove them, not different ones.

36. As more electrons are removed (2nd,3rd, … ionisation energies) more energy is needed because:
(a) You are removing electrons from a positive ion so nuclear charge is very strong (more protons than electrons).
(b) Often removing from a full or half-full orbital which are very stable.
(c) Often jumping to a new shell or orbital which is closer to the nucleus so nuclear charge is stronger.