1. Arrangement of electrons in atom

2. Neils Bohr
Investigated the location of electrons in the atom

3. Neils Bohr Looked at the emission line spectrum of elements
When light passes through a prism it splits to give a continuous spectrum

4. When light passes through a prism it splits to give a spectrum

5. Bohr repeated this experiment using light from a hydrogen discharge bulb
hydrogen discharge bulb = glass tube filled with hydrogen gas and an electric current passed through
Instead of seeing a continuous spectrum he saw a series of narrow lines

6. Bohr repeated it with other elements
Saw that each element had a different emission line spectrum
Therefore, you can identify an element by looking at its line spectrum

7. Bohr concluded that the different arrangement of electrons in each atom must explain why each element has a different emission line spectrum

8. Spectrometer

9. Uses of Emission Spectra
Yellow street lights = sodium
Advertising signs = Neon

10. Uses of Emission Spectra
Fireworks!

11. Mandatory Experiment 1.1 Flame Tests

12. Bohr’s Theory
Electrons revolve around the nucleus in fixed paths called energy levels (or orbits/shells)
Energy levels are represented by the letter n
Lowest: n=1, n=2 n=3 etc
n=1
n=2
n=3

13. Bohr’s Theory
Each energy level has a fixed amount of energy
An electron in an energy level posses the same amount of energy as the energy level
An energy level is the fixed energy value that an electron in an atom may have

14. n=1
n=2
n=3

15. Atoms normally exist in the ground state (they occupy the lowest available energy level)
If an electron is given energy (from electricity/heat) they can jump from the lower energy levels to higher ones.
n=1
n=2

16. When this happens an electron is said to be in its excited state
The excited state of an atom is when the electrons occupy higher energy levels than those available in the ground state
n=1
n=2

17. The amount of energy absorbed by the electron as it ‘jumps’ from the lower to the higher energy level equals the difference in energy between the levels
Orbit of energy E1
Orbit of energy E2
n=1
n=2

18. Electrons in the excited state are unstable and they fall back down to the ground state
As they fall back down they emit fixed amounts of light energy (seen as colours)
The amount of light energy emitted is equal to the amount of energy absorbed
n=1
n=2

19. The energy corresponds to a certain wavelength of light
The amount of light energy emitted is equal to the amount of energy absorbed
The energy corresponds to a certain wavelength of light
E=hf
n=1
n=2

20. Excited State Spectrum Excited State Excited State UV IR Ground State
Excited State unstable and drops back down
Excited State UV
But only as far as n = 2 this time
n=2
Vi s ible
Energy released as a photon
Frequency proportional to energy drop IR
n=1
Ground State

21. The light emitted appears as a specific line on the spectrum

23. The frequency of the light emitted depends on the difference in energy between the two energy levels
E2 - E1 = hf
Frequency of light emitted
Energy of higher energy level
Planck’s constant
Energy of lower energy level

24. Bohr Theory- Key points
The electron in a hydrogen atom occupies fixed energy levels
in its ground state, electrons occupy the lowest available energy level
The elctron can move (jump/become excited) to a higher energy level if it receives a certain amount of energy
The energy (photon) absorbed must exactly equal the energy difference between the ground state (lower level) and excited state (higher level) according to E2-E1= hf
The excited state is unstable and the electron falls back to a lower level
It emits the excess energy in the form of a photon of light (hf) with definite frequency (wavelength) according to E2 – E1 = hf

25. Bohr Theory- Key points
Frequency depends on difference in energy levels
When an electron falls to:
n=1 level gives UV range (Paschen Series)
n=2 level gives Visible Range (Balmer Series)
n=3 level gives IR Range (Lyman Series)

27. Exam Q
How does the lines on emission spectrum provide evidence for the existence of energy levels?
The lines on the emission spectrum are produced when an electron falls from the excited state to the ground state
The line is the light energy produced when it falls
The light energy is the difference between the energy level, so shows the existence of energy levels

28. Atomic Absorption Spectrometry

29. White light is passed through a gaseous sample of the element and then through a prism
The spectrum of light that comes out has certain wavelengths of light missing
The absorption spectrum of an element is the exact opposite of its emission spectrum

31. Principle behind AAS based on 2 facts:
Electrons in the ground state absorb the same energies of light that they emit
The amount of light absorbed is directly proportional to the concentration of the element present in the sample
Emission and absorption spectrums can be used to identify specific elements
Uses: Atomic Absorption Spectrometer used to detect metals in water

33. Sub-levels
In many cases, it appeared that some lines were actually made of two or more lines close together on the spectrum
E.g: ELS for sodium consists of two yellow lines not one
Could not have resulted from electrons dropping to 2 different energy levels (lines would be much further apart)

34. Sub-levels
Proposed that each energy level (excluding the first) is made up of a number of sub-levels which are close in energy
Discovered that the number of sub-levels an energy level has is the same as the value for n for that sub-level
n = 1; 1 sub-level
n = 3; 3 sub-levels
n = 2; 2 sub-levels
n = 4; 4 sub-levels

35. Sub-level of lowest energy is s
The next lowest energy is p
Followed by d and f.

36. Energy sublevels
n=1 1s n=2 2s 2p n=3 3s 3p 3d n=4 4s 4p 4d 4f n=5 5s 5p 5d 5f n=6 6s 6p 6d 6f n=7 7s 7p 7d 7f
s=sharp
p=principal
d=diffuse
f=fundamental

37. Bohr theory and emission spectra https://www. youtube. com/watch
Absorption v emission spectra

38. Wave nature of the electron
Discovered that the electron travels in a wave motion rather than in fixed paths as proposed by Bohr.

39. Wave nature of the electron
Discovered that the electron travels in a wave motion rather than in fixed paths as proposed by Bohr.

40. Heisenberg’s Uncertainty Principle:
It is impossible to measure both the position and velocity of an electron at the same time

41. Bohr’s Theory did not take into account the fact that the electron had a wave motion
Electrons do not revolve around the nucleus in fixed paths as proposed by Bohr
-> Heisenberg’s Uncertainty Principle was in conflict with Bohr’s Theory.
Therefore, we can only talk about the probability of finding an electron in a region in space
They exist in orbitals

42. Schrodinger
Worked out the probability of finding an electron in any sublevel
Schrodinger’s equations allowed for the calculation of the SHAPES of these orbitals

43. An Orbital:
An orbital is a region in space where there is a high probability of finding an electron
A Sublevel:
Is a subdivision of a main energy level and consists of one or more orbitals of the same energy

44. S orbitals are spherical
An Orbital:
An orbital is a region in space where there is a high probability of finding an electron
S orbitals are spherical

45. An Orbital:
An orbital is a region in space where there is a high probability of finding an electron
P orbitals are ’dumbbell” shaped
Consist of three separate orbitals (each containing 2 electrons) on the three axis
All three have the same energy

46. Energy sublevels https://www.youtube.com/watch?v=9E3QaRxqXZc

47. Limitations to Bohr’s Theory
Bohr’s theory only worked to explain the emission spectrum of Hydrogen
Bohr’s theory did not take into account the fact that an electron had a wave motion
Heisenberg’s Uncertainty Principle was in conflict with Bohr’s theory
Bohr’s theory could not explain the splitting of certain lines in emission spectra and did not take into account the presence of sublevels

48. N.B: Do not confuse orbit and orbital
Key Words
Energy Level/Orbit
Emission spectrum
Ground state
Excited state
Atomic Emission Spectrometry
Atomic Absorption Spectrometry
Energy Sublevels
Wave nature
Orbital
N.B: Do not confuse orbit and orbital