1. Atomic Theory
State the position of protons, electrons and neutrons in the atom
State the relative masses and relative charges of protons, neutrons and electrons
Define the terms mass number (A), atomic number (Z) and isotopes of an element
Deduce the symbol for an isotope given its mass number and atomic number
Calculate the number of protons, neutrons and electrons in atoms and ions from the mass number, atomic number and charge.
Compare the properties of the isotopes of an element
Discuss the uses of radioisotopes

2. History of the atom
Democritus (400 BC) suggested that the material world was made up of tiny, indivisible particles
atomos, Greek for “uncuttable”
Aristotle believed that all matter was made up of 4 elements, combined in different proportions
Fire - Hot
Earth - Cool, heavy
Water - Wet
Air - Light
The “atomic” view of matter faded for centuries, until early scientists attempted to explain the properties of gases

3. Re-emergence of Atomic Theory
John Dalton postulated that:
All matter is composed of extremely small, indivisible particles called atoms
All atoms of a given element are identical (same properties); the atoms of different elements are different

4. 3. Atoms are neither created nor destroyed in chemical reactions, only rearranged
4. Compounds are formed when atoms of more than one element combine
A given compound always has the same relative number and kind of atoms

5. Atoms are divisible!
By the 1850s, scientists began to realize that the atom was made up of subatomic particles
Thought to be positive and negative
How would we know this if we can’t see it or touch it?

6. Cathode Rays and Electrons
Mid-1800’s scientists began to study electrical discharge through cathode-ray tubes. Ex: neon signs
Partially evacuated tube in which a current passes through
Forms a beam of electrons which move from cathode to anode
Electrons themselves can’t be seen, but certain materials fluoresce (give off light) when energised

7. Oh there you are!
JJ Thompson observed that when a magnetic or electric field are placed near the electron beam, they influence the direction of flow
opposite charges attract each other, and like charges repel.
The beam is negatively charged so it was repelled by the negative end of the magnet

8. Magnetic field forces the beam to bend depending on orientation
Magnetic field forces the beam to bend depending on orientation
Thompson concluded that:
Cathode rays consist of beams of particles
The particles have a negative charge

9. Thompson understood that all matter was inherently neutral, so there must be a counter
A positively charged particle, but where to put it
It was suggested that the negative charges were balanced by a positive umbrella-charge
“Plum pudding model” “chocolate chip cookie model”

10. Rutherford and the Nucleus
This theory was replaced with another, more modern one
Ernest Rutherford (1910) studied angles at which a particles (nucleus of helium) were scattered as they passed through a thin gold foil

12. Rutherford expected …
Rutherford believed that the mass and positive charge was evenly distributed throughout the atom, allowing the a particles to pass through unhindered
a particles

13. Rutherford explained …
Atom is mostly empty space
Small, dense, and positive at the center
Alpha particles were deflected if they got close enough +
a particles

14. The modern atom is composed of two regions:
Nucleus: Containing protons and neutrons, it is the bulk of the atom and has a positive charge associated with it
Electron cloud: Responsible for the majority of the volume of the atom, it is here that the electrons can be found orbiting the nucleus (extranuclear)

15. Major Subatomic Particles
Name
Symbol
Charge
Relative Mass (amu)
Actual Mass (g)
Electron e- -1
1/1840
9.11x10-28
Proton p+ +1 1
1.67x10-24
Neutron no
Atoms are measured in picometers, meters
Hydrogen atom, 32 pm radius
Nucleus tiny compared to atom
If the atom were a stadium, the nucleus would be a marble
Radius of the nucleus is on the order of m
Density within the atom is near 1014 g/cm3

16. Elemental Classification
Atomic Number (Z) = number of protons (p+) in the nucleus
Determines the type of atom
Li atoms always have 3 protons in the nucleus, Hg always 80
Mass Number (A) = number of protons + neutrons [Sum of p+ and nº]
Electrons have a negligible contribution to overall mass
In a neutral atom there is the same number of electrons (e-) and protons (atomic number)

17. Nuclear Symbols
Every element is given a corresponding symbol which is composed of 1 or 2 letters (first letter upper case, second lower), as well as the mass number and atomic number E A Z
elemental symbol
mass number
atomic number

18. W F Br 184 74 19 9 80 35 Find the number of protons number of neutrons
number of electrons
atomic number
mass number W
184 74 F 19 9 Br 80 35

19. Ions
Cation is a positively charged particle. Electrons have been removed from the element to form the + charge.
ex: Na has 11 e-, Na+ has 10 e-
Anion is a negatively charged particle. Electrons have been added to the atom to form the – charge.
ex: F has 9 e-, F- has 10 e-

20. Isotopes
Atoms of the same element can have different numbers of neutrons and therefore have different mass numbers
The atoms of the same element that differ in the number of neutrons are called isotopes of that element
When naming, write the mass number after the name of the element H 1
Hydrogen-1 2
Hydrogen-2 3
Hydrogen-3

21. How heavy is an atom of oxygen?
There are different kinds of oxygen atoms (different isotopes)
16O, 17O, 18O
We are more concerned with average atomic masses, rather than exact ones
Based on abundance of each isotope found in nature
We can’t use grams as the unit of measure because the numbers would be too small
Instead we use Atomic Mass Units (u)
Standard u is 1/12 the mass of a carbon-12 atom
Each isotope has its own atomic mass

22. Calculating Averages
Average = (% as decimal) x (mass1) (% as decimal) x (mass2) (% as decimal) x (mass3) + …
Problem:
Silver has two naturally occurring isotopes, 107Ag with a mass of u and abundance of % ,and 109Ag with a mass of u and abundance of % What is the average atomic mass?
Average = (0.5184)( u) + (0.4816)( u)
= amu

23. Average Atomic Masses
If not told otherwise, the mass of the isotope is the mass number in amu
The average atomic masses are not whole numbers because they are an average mass value
Remember, the atomic masses are the decimal numbers on the periodic table

24. Properties of Isotopes
Chemical properties are primarily determined by the number of electrons
All isotopes has the same number of electrons, so they have nearly identical chemical properties even though they have different masses.
Physical properties often depend on the mass of the particle, so among isotopes they will have slightly different physical properties such as density, rate of diffusion, boiling point…
The isotopes of an element with fewer neutrons will have:
Lower masses • faster rate of diffusion
Lower densities • lower melting and boiling points

25. More Practice Calculating Averages
Calculate the atomic mass of copper if copper has two isotopes
69.1% has a mass of amu
The rest (30.9%) has a mass of amu
Magnesium has three isotopes
78.99% magnesium 24 with a mass of amu
10.00% magnesium 25 with a mass of amu
The rest magnesium 26 with a mass of amu
What is the atomic mass of magnesium?

26. Radioisotopes
Isotopes of atoms that have had an extra neutron attached to their nucleus.
Carbon-14 radioactive decay is used to measures the date of objects.
After 5700 years the amount of 14C will be half its original value.
Iodine-125 or 131 is used to monitor the activity of the thyroid gland (b/c the thyroid tends to absorb iodine)

27. Cobalt-60 produces gamma rays (intense radioactivity) and is used in radiation treatment of cancer.
Note: gamma rays are the shortest wavelength on the electromagnetic spectrum. They are the most dangerous and difficult to shield from.

28. 2.2 The Mass Spectrometer

29. Mass Spectrometer The mass spectrometer is an instrument used:
To measure the relative masses of isotopes
To find the relative abundance of the isotopes in a sample of an element
When charged particles pass through a magnetic field, the particles are deflected by the magnetic field, and the amount of deflection depends upon the mass/charge ratio of the charged particle.

30. Mass Spectrometer – 5 Stages
Once the sample of an element has been placed in the mass spectrometer, it undergoes five stages.
Vaporisation – the sample has to be in gaseous form. If the sample is a solid or liquid, a heater is used to vaporise some of the sample.
X (s)  X (g)
or X (l)  X (g)

31. Mass Spectrometer – 5 Stages
Ionization – sample is bombarded by a stream of high-energy electrons from an electron gun, which ‘knock’ an electron from an atom. This produces a positive ion:
X (g)  X + (g) + e-
Acceleration – an electric field is used to accelerate
the positive ions towards the magnetic field. The
accelerated ions are focused and passed through a
slit: this produces a narrow beam of ions.

32. Mass Spectrometer – 5 Stages
Deflection –
The accelerated ions are deflected into the magnetic field. The amount of deflection is greater when:
• the mass of the positive ion is less
• the charge on the positive ion is greater
• the velocity of the positive ion is less
• the strength of the magnetic field is
greater

33. Mass Spectrometer
If all the ions are travelling at the same velocity and carry the same charge, the amount of deflection in a given magnetic field depends upon the mass of the ion.
For a given magnetic field, only ions with a particular relative mass (m) to charge (z) ration – the m/z value – are deflected sufficiently to reach the detector.

34. Mass Spectrometer
Detection – ions that reach the detector cause electrons to be released in an ion-current detector
The number of electrons released, hence the current produced is proportional to the number of ions striking the detector.
The detector is linked to an amplifier and then to a recorder: this converts the current into a peak which is shown in the mass spectrum.

35. Atomic Structure – Mass Spectrometer
Isotopes of boron
m/z value 11 10
Relative abundance %
18.7
81.3
Ar of boron = (11 x 18.7) + (10 x 81.3)
( ) =
100
= = 10.2

36. Mass Spectrometer – Questions
A mass spec chart for a sample of neon shows that it contains:
90.9% 20Ne
0.17% 21Ne
8.93% 22Ne
Calculate the relative atomic mass of neon
You must show all your work!

37. Mass Spectrometer – Questions
90.9% 20Ne
0.17% 21Ne
8.93% 22Ne
(90.9 x 20u) + (0.17 x 21u) + (8.93 x 22u)
100
Ar= 20.18u

38. Mass Spectrometer – Questions
204
206
207
208
52.3
23.6
22.6
1.5
Calculate the relative atomic mass of lead
You must show all your work!

39. Mass Spectrometer – Questions
1.5% 204Pb
23.6% 206Pb
22.6% 207Pb
52.3% 208Pb
(1.5 x 204) + (23.6 x 206) + (22.6 x 207)+(52.3 x 208)
100
100
100 =
Ar=

40. 2.3 Electron Arrangement
Describe the electromagnetic spectrum
Distinguish between a continuous spectrum and a line spectrum
Explain how the lines in the emission spectrum of hydrogen are related to electron energy levels
Deduce the electron arrangement for atoms and ions up to Z=20

41. Electromagnetic radiation.

42. Electromagnetic Radiation
Most subatomic particles behave as PARTICLES and obey the physics of waves.

43. Electromagnetic Radiation
wavelength
Visible light
Ultaviolet radiation
Amplitude
Node

44. Wavelengths and energy
Understand that different wavelengths of electromagnetic radiation have different energies.
Waves have a frequency
c=νλ
c=velocity of wave (2.998 x 108 m/s)
ν=(nu) frequency of wave, units are “cycles per sec”
λ=(lambda) wavelength

45. Electromagnetic Spectrum
In increasing energy, ROY G BIV

46. Electromagnetic Spectrum
Long wavelength --> small frequency
Short wavelength --> high frequency
increasing frequency
increasing wavelength

47. Bohr’s Model Why don’t the electrons fall into the nucleus?
Move like planets around the sun.
In circular orbits at different levels.
Amounts of energy separate one level from another.

48. Bohr postulated that: Fixed energy related to the orbit
Electrons cannot exist between orbits
The higher the energy level, the further it is away from the nucleus
An atom with maximum number of electrons in the outermost orbital energy level is stable (unreactive)
Think of Noble gases

49. Those who are not shocked when they first come across quantum theory cannot possibly have understood it. (Niels Bohr on Quantum Physics)

50. Atomic Line Emission Spectra and Niels Bohr
Bohr’s greatest contribution to science was in building a simple model of the atom. It was based on an understanding of the LINE EMISSION SPECTRA of excited atoms.
Problem is that the model only works for Hydrogen
Niels Bohr
( )

51. How did he develop his theory?
He used mathematics to explain the visible spectrum of hydrogen gas
Lines are associated with the fall of an excited electron back down to its ground state energy level.

52. Spectrum of White Light

53. The line spectrum
electricity passed through a gaseous element emits light at a certain wavelength
Can be seen when passed through a prism
Every gas has a unique pattern (color)

54. Line Emission Spectra of Excited Atoms
Excited atoms emit light of only certain wavelengths
The wavelengths of emitted light depend on the element.

55. Spectrum of Excited Hydrogen Gas

56. Line Spectra of Other Elements

57. Continuous line spectrum
Carbon
Helium
Continuous line spectrum

58. Bohr also postulated that an atom would not emit radiation while it was in one of its stable states but rather only when it made a transition between states.
The frequency of the radiation emitted would be equal to the difference in energy between those states divided by Planck's constant.

59. ν = frequency of the photon of light
Ehigh-Elow= hν = hc/λ
h=6.63 × 10–34 J s = Planck’s constant
E= energy of the emitted light (photon)
ν = frequency of the photon of light
λ = is usually stated in nm, but for calculations use m.
This results in a unique emission spectra for each element, like a fingerprint.
electron could "jump" from one allowed energy state to another by absorbing/emitting photons of radiant energy of certain specific frequencies.

60. Energy must then be absorbed in order to "jump" to another energy state, and similarly, energy must be emitted to "jump" to a lower state.
The frequency, ν, of this radiant energy corresponds exactly to the energy difference between the two states.
In order for the emitted energy to be seen as light the wavelength of the energy must be in between 380 nm to 750 nm

61. Bohr’s Triumph
His theory helped to explain periodic law (the trends from the periodic table)
Halogens (gp.17 or group VII) are so reactive because it has one e- less than a full outer orbital
Alkali metals (gp. 1 or group I) are also reactive because they have only one e- in outer orbital

62. Drawback
Bohr’s theory did not explain or show the shape or the path traveled by the electrons.
His theory could only explain hydrogen and not the more complex atoms

63. Energy level populations
Electrons found per energy level of the atom.
The first energy level holds 2 electrons
The second energy level holds 8 electrons (2 in s and 6 in p)
The third energy level holds 18 electrons (2 in s, 6 in p and 10 in d) There is overlapping here, so when we do the populations there will be some changes.
That is as far as this course requires us to go!

64. Examples for group 1
Li 2.1 Na K

65. The Quantum Mechanical Model
Energy is quantized. It comes in chunks.
A quanta is the amount of energy needed to move from one energy level to another.
Since the energy of an atom is never “in between” there must be a quantum leap in energy.
Schrödinger derived an equation that described the energy and position of the electrons in an atom

66. Quantum or Wave Mechanics
Schrodinger applied idea of e- behaving as a wave to the problem of electrons in atoms.
He developed the WAVE EQUATION
Solution gives set of math expressions called WAVE FUNCTIONS, 
Each describes an allowed energy state of an e-
E. Schrodinger

67. Heisenberg Uncertainty Principle
The problem of defining nature of electrons in atoms was solved by W. Heisenberg.
He observed that one cannot simultaneously define the position and momentum (= m•v) of an electron.
If we define the energy exactly of an electron precisely we must accept limitation that we do not know exact position.
W. Heisenberg

68. A good site: http://www.chemguide.co.uk/basicorg/bonding/orbitals.html

69. Electron Configuration HL only
State the relative energies of s, p, d, and f orbitals in a single energy level
State the maximum number of orbitals in a given energy level.
Draw the shape of an s orbital and the shapes of px, py and pz orbitals
Apply the Aufbau principle, Hund’s rule and the Pauli exclusion principle to write electron configurations for atoms and ions up to Z=54.

70. S orbitals 1 s orbital for every energy level 1s 2s 3s
Spherical shaped
Each s orbital can hold 2 electrons
Called the 1s, 2s, 3s, etc.. orbitals

71. P orbitals Start at the second energy level 3 different directions
3 different shapes
Each orbital can hold 2 electrons

72. The D sublevel contains 5 D orbitals
The D sublevel starts in the 3rd energy level
5 different shapes (orbitals)
Each orbital can hold 2 electrons

73. The F sublevel has 7 F orbitals
The F sublevel starts in the fourth energy level
The F sublevel has seven different shapes (orbitals)
2 electrons per orbital

74. Summary
Starts at energy level

75. Electron Configurations
The way electrons are arranged in atoms.
Aufbau principle- electrons enter the lowest energy first.
This causes difficulties because of the overlap of orbitals of different energies.
Pauli Exclusion Principle- at most 2 electrons per orbital - different spins
Hund’s Rule- When electrons occupy orbitals of equal energy they don’t pair up until they have to .

76. Increasing energy 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p 3s 2p

77. Increasing energy 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p 3s 2p
Phosphorous, 15 e- to place
The first to electrons go into the 1s orbital
Notice the opposite spins
only 13 more

78. Increasing energy 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p 3s 2p
The next electrons go into the 2s orbital
only 11 more

79. Increasing energy 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p
The next electrons go into the 2p orbital
only 5 more

80. Increasing energy 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p
The next electrons go into the 3s orbital
only 3 more

81. Increasing energy 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s
The last three electrons go into the 3p orbitals.
They each go into separate shapes
3 unpaired electrons
1s22s22p63s23p3

82. Orbitals fill in order Lowest energy to higher energy.
Adding electrons can change the energy of the orbital.
Half filled orbitals have a lower energy.
Makes them more stable.
Changes the filling order

83. Write these electron configurations
Titanium - 22 electrons
1s22s22p63s23p64s23d2
Vanadium - 23 electrons
1s22s22p63s23p64s23d3
Chromium - 24 electrons
1s22s22p63s23p64s23d4 is expected
But this is wrong!!

84. Chromium is actually 1s22s22p63s23p64s13d5 Why?
This gives us two half filled orbitals.
Slightly lower in energy.
The same principal applies to copper.

85. Copper’s electron configuration
Copper has 29 electrons so we expect
1s22s22p63s23p64s23d9
But the actual configuration is
1s22s22p63s23p64s13d10
This gives one filled orbital and one half filled orbital.
Remember these exceptions

86. Electronic Structure – of transition metals
With the transition metals it is the 4s electrons that are lost first when they form ions:
Titanium (Ti) - loss of 2 e-
1s2 2s2 2p6 3s2 3p6 4s2 3d2 
1s2 2s2 2p6 3s2 3p6 3d2
Ti atom
Ti2+ ion
Chromium (Cr) - loss of 3 e-
1s2 2s2 2p6 3s2 3p6 4s1 3d5 
1s2 2s2 2p6 3s2 3p6 3d3
Cr atom
Cr3+ ion

87. Electronic Structure - Questions
Copy and complete the following table:
Atomic no.
Mass no.
No. of protons
No. of neutrons
No. of electrons
Electronic structure Mg 12
1s2 2s2 2p6 3s2
Al3+ 27 10
S2- 16
Sc3+ 21 45
Ni2+ 30 26

88. Ionization Energy
Explain how evidence from first ionization energies across periods accounts for the existence of main energy levels and sub-levels in atoms
Explain how successive ionization energy data is
related to the electron configuration of an atom

89. Ionization Energy
The amount of energy required to completely remove an electron from a gaseous atom.
An atom's 'desire' to grab another atom's electrons.
Removing one electron makes a +1 ion.
The energy required is called the first ionization energy.
X(g) + energy →X+ + e-

90. Ionization Energy
The second and third ionization energies can be represented as follows:
 X+ (g) + energy X2+ (g) + e-
X2+ (g) + energy X3+ (g) + e-
More energy required to remove 2nd electron, and still more energy required to remove 3rd electron

91. Group trends
Ionization energy decreases down the group.

92. Going from Be to Mg, IE decreases because:
Mg outer electron is in the 3s sub-shell rather than the 2s. This is higher in energy
The 3s electron is further from the nucleus and shielded by the inner electrons
So the 3s electron is more easily removed
A similar decrease occurs in every group in the periodic table.

93. Notice any trends? Any surprises?

94. General trend: Increasing I.E. as we go across a period
Look at the peak at Mg and the plateau between P and S. Can you explain why?

95. Why is there a fall from Mg to Al?
 Al has configuration 1s2 2s2 2p6 3s2 3p1, its outer electron is in a p sublevel
 Mg has electronic configuration s22s22p63s2.
 The p level is higher in energy and with Mg the s sub level is full – this gives it a slight stability advantage

96. Why is there a fall from P to S?
 This can be explained in terms of electron pairing.
 As the p sublevel fills up, electrons fill up the vacant sub levels and are unpaired.
This configuration is more energetically stable than S as all the electrons are unpaired. It requires more energy to pair up the electrons in S so it has a lower Ionization energy.
There is some repulsion between the paired electrons which lessens their attraction to the nucleus.
 It becomes easier to remove!

97. Driving Force Full Energy Levels are very low energy.
Noble Gases have full energy levels.
Atoms behave in ways to achieve noble gas configuration.

98. 2nd Ionization Energy
For elements that reach a filled or half filled sublevel by removing 2 electrons 2nd IE is lower than expected.
Makes it easier to achieve a full outer shell
True for s2
Alkaline earth metals form +2 ions.

99. 2nd IE and 3rd IE are always higher than 1st IE!!!
Using the same logic s2p1 atoms have an low 3rd IE.
Atoms in the aluminum family form +3 ions.
2nd IE and 3rd IE are always higher than 1st IE!!!