1. 1.4 Energy Levels in Atoms

2. 1.4 a Learning Outcomes
use flame tests to provide evidence that energy is absorbed or released in discrete units when electrons move from one energy level to another
explain how flame tests provide evidence that energy is absorbed or released in discrete units when electrons move from one energy level to another
relate energy levels in atoms to everyday applications such as sodium street lights and fireworks
illustrate how line spectra provide evidence for energy levels

3. 1.4 a Learning Outcomes define and explain energy levels in atoms
define and explain energy levels in atoms
describe and explain the emission spectrum of the hydrogen atom using the Balmer series in the emission spectrum as an example
describe and explain the absorption spectrum
discuss the uses of atomic absorption spectrometry (AAS) as an analytical technique

4. 1.4 a Learning Outcomes
illustrate how line spectra provide evidence for energy levels
use a spectroscope or a spectrometer to view emission spectra of elements
define and explain energy sub-levels
state the Heisenberg uncertainty principle
state the dual wave-particle nature of the electron (mathematical treatment not required)
define and explain atomic orbitals
describe the shapes of s and p orbitals
build up the electronic structure of the first 36 elements
derive the electronic configurations of ions of s- and p block elements only

5. Introduction
When metal ions are heated strongly in a non-luminous Bunsen flame or excited by high voltage electric currents characteristic colours are produced.
These colours are so characteristic that they can be used to identify the ions present.
How are these colours produced?

6. Flame Tests Element Flame Colour Barium Yellow-green Copper Blue-green
Lithium
Deep red
Potassium
Lilac
Sodium
Yellow
Strontium
Red
Flame Tests
Street Lights -Sodium
Fireworks - Strontium

7. Emission spectra known for hundreds of years
Niels Bohr unlocked their secret
Electrons orbiting in shells around the nucleus
Won the 1922 Nobel Prize for Physics

8. Excited State unstable and drops back down
Spectrum
n=4
n=3
Excited State unstable and drops back down UV
Excited State
n=2
Vi s ible
Energy released as a photon
Frequency proportional to energy drop IR
Electron excited by just the right amount of heat or electricity
n=1
Ground State

9. Excited State unstable and drops back down
Spectrum
n=4
Excited State
n=3
Excited State unstable and drops back down UV
n=2
Vi s ible
Energy released as a photon
Frequency proportional to energy drop IR
Electron excited by just the right amount heat or electricity
n=1
Ground State

10. Excited State unstable and drops back down
Spectrum
n=4
n=3
Excited State unstable and drops back down 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
Electron excited by just the right amount of heat or electricity
n=1
Ground State

11. Excited State unstable and drops back down
Spectrum
n=4
n=3
Excited State unstable and drops back down UV
But only as far as n = 3 this time
n=2
Vi s ible
Energy released as a photon
Frequency proportional to energy drop IR
Electron excited by just the right amount of heat or electricity
n=1
Ground State

12. Energy Level
Shell which electrons of equal energy can occupy

13. Summary Electron normally in Ground State
Just the right amount of energy supplied [as heat or electricity]
Electron jumps to higher energy level
Now in Excited State
Unstable
Drops back to a lower energy level

14. Summary
Energy that was absorbed to make the jump up is now released as a photon
Frequency depends on difference in energy levels
[ E2 - E1 = hf ]
E2 = excited state
E1 = ground state
h is Plank’s Constant and
f is frequency of light
When electron falls to
n = 1 level gives UV Range
n = 2 level gives Visible Range - Balmer Series
n = 3,4 or 5 levels gives IR Range

15. Summary Why does each element have a unique line emission spectrum?
Each element has different electron configurations
This gives rise to different electron transitions

16. Bohr
Bohr’s model of the atom explained the existence of energy levels on the basis of atomic spectra. It was later modified to incorporate the idea of orbitals in recognition of the wave nature of the electron and Heisenberg’s uncertainty principle.
Outline Bohr’s Atomic Theory, based on the hydrogen emission spectrum
The electron in a hydrogen atom electron occupies ground state (which is the lowest energy level available)
The electron can move (become excited) to a higher energy level if it receives a photon of energy

17. Bohr
The photon (energy) must be exactly equal to the energy difference between the ground
state (a lower level) and a higher energy level (excited state)
The electron in an excited state (a higher level) is unstable
The excited electron falls back to a lower energy level
Emitting energy in the form of a photon of light according to E2 – E1 = hf (hν)

18. Bohr
Explain how the expression E2 – E1 = hf links the occurrence of the visible lines in the hydrogenspectrum to energy levels in a hydrogen atom. (12) Outline Bohr’s Atomic Theory, based on the hydrogen emission spectrum
E2 – E1 :energy difference between higher (e.g. E2) and lower (e.g. E1) level
f : frequency of line in spectrum
h is Planck’s constant and hf is a photon {quantumof energy}
the expression indicates that the energy difference (E2 – E1) is proportional to (varies
directly with) the frequency (f)

19. Bohr
Describe how Bohr used line emission spectra to explain the existence of energy levels in atoms. (13)
Electrons in ground state OR energy of electron quantised (4)
Photons of light is absorbed and electron jumps to a higher level(s) (excited state) (3)
Excited state unstable and electrons fall back to lower levels emitting energy as a photon of light (3)
Energy difference between levels gives energy of photon which is calculated by E2 – E1 = h) (3)

20. Spectral Evidence for Sublevels
Lines found to be much more complex than at first thought
Found to be divided into other lines
Divisions called sub-levels

21. Sub-Levels Increasing energy On Closer inspection Nucleus 4f
Main Energy Levels 4d
n=4 4p 3d 4s 3p
n=3 3s 2p 2s
n = 2 1s
n =1
Nucleus
Sub-Levels

22. Energy Sub-level
A sub-division of a main energy level consisting of one or more orbitals of the same energy

23. Heisenberg Uncertainty Principle
Louis de Broglie suggested Wave-particulate duality electrons [like light] have the properties of waves as well as particles
If this is true Bohr’s idea of a precise path a precise distance from nucleus cannot be true

24. Heisenberg Uncertainty Principle
Heisenberg treated path of electron around nucleus mathematically
Heisenberg Uncertainty Principle: The more accurately we determine the position of an electron (at any instant) the less accurately we can determine its velocity.

25. Heisenberg Uncertainty Principle – So What?
When we measure its position - the measuring procedure affects its velocity
Now have to talk about the probability of finding an electron at a particular position
Erwin Schrodinger [Austrian]: Worked out the probability of finding a particular electron in an atom – it was his mathematical equations that determined the shapes of the orbitals

26. Atomic Orbitals
Shapes

27. Principal Energy Levels or shells
Divided into sublevels of differing energy
Sublevels further divided into divisions of equal energy called orbitals
Region in space around the nucleus in which electrons have a 99%+ probability of being found.
An orbital can hold 2 electrons
There are four types of orbitals; classed according to their shapes
s, p, d, and f orbitals

28. s orbitals
Spherical
There is only one s orbital per main energy level [shell]
Each s orbital can hold 2 electrons
So each s sub-level can hold 2 electrons

29. p orbitals Dumb-bell shaped
Each p sublevel is made up of 3 orbitals px , py , pz [mutually at right angles]
Each p sublevel can hold a total of 6 e-
All the orbitals in a p sublevel have the same energy i.e. px = py = pz

30. d orbitals Complex in shape
Each d sublevel is made up of 5 orbitals (each orbital can hold 2 e-)
Therefore each d sublevel can hold 10 e-
All orbitals in a d sublevel have the same energy

31. Atomic
Models

32. Aufbau Principle Electrons occupy the lowest available energy level
(Aufbau is German for construction or building)
Fill sublevels from the nucleus outwards
This gives an s,p,d electron configuration

33. Order of filling of sublevels
1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f
Need to know s,p,d electron configuration of first 36 elements

34. 1s2 2s2 2p6 3s2 3p3 Example 11 5 3 Phosphorous 1531 P 15 electrons
Leaves 13 electrons 11 5 3
Considering Aufbau: Above example indicates n=1 energy level filled before n=2, which is filled before n=3. Next example shows exception of 4s sublevel

35. Scandium Sc
21 electrons
1s2 2s2 2p6 3s23p6 4s2 3d1
Iron Fe
26 electrons
1s2 2s2 2p6 3s23p6 4s2 3d6

36. Exceptions at 3d Chromium 2452 Cr 24 electrons
1s2 2s2 2p6 3s23p6 4s1 3d5
Copper Cu
29 electrons
1s2 2s2 2p6 3s23p6 4s1 3d10

37. The Electronic Configuration of Atoms

38. 1s

39. 2p 2s 1s

40. 3d 3p 3s 2p 2s 1s

41. 4f 4d 4p 3d 4s 3p 3s 2p 2s 1s

42. 4f 6s 5p 4d 5s 4p 3d 4s 3p 3s 2p 2s 1s

43. Aufbau Principle: Electrons fill the lowest available energy level
Click to add electrons 4p 3d 4s 3p
Cr an electron is promoted from 4s to 3d to give a half-filled 3d subshell
4s fills before 3d
Cu an electron is promoted from 4s to 3d to give a full 3d subshell 3s
Electrons remain unpaired as far as possible 2p 2s
When building up the electron configuration of an atom, in it’s ground state, the electrons occupy the lowest available energy level
Hund’s Rule of Maximum Multiplicity: When two or more orbitals of equal energy are available, the electrons occupy them singly before filling them in pairs. 1s

44. Electronic configuration in shorthand nomenclature
Click to add electrons 4p 3d 4s 3p 3s
Cr 1s2 2s2 2p6 3s2 3p6 4s1 3d5
Fe 1s2 2s2 2p6 3s2 3p6 4s2 3d6
V 1s2 2s2 2p6 3s2 3p6 4s2 3d3
Mn 1s2 2s2 2p6 3s2 3p6 4s2 3d5
K 1s2 2s2 2p6 3s2 3p6 4s1
Co 1s2 2s2 2p6 3s2 3p6 4s2 3d7
Ca 1s2 2s2 2p6 3s2 3p6 4s2
Sc 1s2 2s2 2p6 3s2 3p6 4s2 3d1
Ti 1s2 2s2 2p6 3s2 3p6 4s2 3d2
Ge 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p2
Se 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p4
Br 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p5
Kr 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6
As 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p3
Ar 1s2 2s2 2p6 3s2 3p6
Cu 1s2 2s2 2p6 3s2 3p6 4s1 3d10
Zn 1s2 2s2 2p6 3s2 3p6 4s2 3d10
Ga 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p1
Ni 1s2 2s2 2p6 3s2 3p6 4s2 3d8
Si 1s2 2s2 2p6 3s2 3p2
B 1s2 2s2 2p1
C 1s2 2s2 2p2
N 1s2 2s2 2p3
Be 1s2 2s2
Li 1s2 2s1
H 1s1
He 1s2
Cl 1s2 2s2 2p6 3s2 3p5
O 1s2 2s2 2p4
P 1s2 2s2 2p6 3s2 3p3
S 1s2 2s2 2p6 3s2 3p4
F 1s2 2s2 2p5
Al 1s2 2s2 2p6 3s2 3p1
Ne 1s2 2s2 2p6
Na 1s2 2s2 2p6 3s1
Mg 1s2 2s2 2p6 3s2 2p 2s
Pauli Exclusion Principle: No more than two electrons may occupy an orbital and they must have opposite spin 1s

45. Reason for Cu & Cr Anomaly
p, d and f sublevels that are
half filled or
completely filled
have extra stability
Therefore one of the electrons in the 4s sublevel flips over into the 3d sublevel to attain this stability
[You need to know the electron configuration for the first 36 elements.]

46. Electron configuration of ions
What is an ion?
It is an atom with a charge
2 types and –
Atoms that take in electrons form -ve ions
For each e- taken in they gain a –ve charge
Atoms that give away e- form +ve ions
For each e- lost they gain a +ve charge

47. Ionisation 3d 4s 3p Zn
Zn2+ 3s
4s electrons (outer shell) are removed before 3d (inner shell) 2p 2s 1s

48. Ionisation Fe Fe2+ Fe3+ 3d 4s 3p 3s
4s electrons (outer shell) are removed before 3d (inner shell)
Only then are 3d electrons removed 2p 2s 1s

49. Electron configuration of ions
Al3+
the 3+ tells us the Al has lost 3 electrons
1327Al has 13 e- so Al3+ has 10 e-
1s2 2s2 2p6
Normally written as [1s2 2s2 2p6 ] N.B. only ions of first 20 asked

50. Electron configuration of ions
the 3- tells us the N has gained 3 electrons
714N has 7 e- so N3- has 10 e-
1s2 2s2 2p6
Normally written as [1s2 2s2 2p6 ] 3-
This is the same configuration as Al3+
And is the same configuration as the nearest Nobel gas - Neon
The [ ] and charge are necessary to distinguish between them

51. Revision of Hund and Pauli
How are the electrons in the 2 p sublevel arranged or distributed among the three orbitals 2px 2py and 2pz ?
Hund’s Rule of Maximum Multiplicity
When two or more orbitals of equal energy are available, the electrons occupy them singly and then in pairs.

52. 2p3 px py pz 
How are the electrons arranged in these orbitals?
Which of these two setups is correct? px py pz 

53. Stable Boron 2p1 Carbon 2p2 Nitrogen 2p3 Oxygen 2p4 Fluorine 2p5
Neon 2p6
Very stable

54. Pauli Exclusion Principle
Arrows-in-boxes show how electrons are distributed in orbitals
Box represents orbitals and arrows represent electrons.
Arrows were first used when it was discovered that electrons ‘spin’ on their own axis.
(rather like the earth as it revolves around the sun)

55. Pauli Exclusion Principle
Electrons can spin
Clockwise or
Anticlockwise
Pauli deduced this from spectra in 1925
Pauli Exclusion Principle
No more than two electrons may occupy an orbital and to do so they must have opposite spin

56. Another way of showing distribution in orbitals
B 1s2 2s2 2px1 2py0 2pz0
C 1s2 2s2 2px1 2py1 2pz0
N 1s2 2s2 2px1 2py1 2pz1
O 1s2 2s2 2px2 2py1 2pz1
F 1s2 2s2 2px2 2py2 2pz1
Ne 1s2 2s2 2px2 2py2 2pz2
Give these if specifically asked for the distribution of electrons in the p sublevel or orbital – if in doubt give these