1. Atomic Structure

2. Subatomic particles Relative mass Actual mass (g) Name Symbol Charge
Electron e- -1
0.0005
9.11 x 10-28
Proton p+ +1 1
1.67 x 10-24
Neutron n0 1
1.67 x 10-24

3. Atomic Number = number of protons in the nucleus
Counting the Pieces
Atomic Number = number of protons in the nucleus
# of protons determines kind of atom (since all protons are alike!)
the same as the number of electrons in the neutral atom.
Mass Number = the number of protons + neutrons.
These account for virtually all mass

4. Symbols
Isotopic Symbols (Isotopic Notation) contain the symbol of the element, the mass number and the atomic number.

5. Symbols
Isotopic Symbols contain the symbol of the element, the mass number and the atomic number.
Mass
number X
Atomic
number

6. F 19 9 Symbols Find the number of protons number of neutrons
number of electrons
Atomic number
Mass Number 19 F 9

7. Isotopes
Atoms of the same element can have different numbers of neutrons and different mass numbers
These are called isotopes.

8. Naming Isotopes
We can also put the mass number after the name of the element.
carbon- 12 or C-12
carbon -14 or C-14
uranium-235 or U-235

9. Atomic Mass
How heavy is an atom of oxygen?
There are different kinds of oxygen atoms.
More concerned with average atomic mass.
Based on abundance of each element in nature.
Don’t use grams because the numbers would be too small.

10. Measuring Atomic Mass
Unit is the Atomic Mass Unit (amu)
One twelfth the mass of a carbon-12 atom.
Each isotope has its own atomic mass, thus we determine the average from percent abundance.

11. Calculating averages
Multiply the atomic mass of each isotope by it’s abundance (expressed as a decimal), then add the results.

12. Atomic Mass
Calculate the atomic mass of copper if copper has two isotopes. 69.1% has a mass of amu and the rest has a mass of amu.

13. Atomic Mass
Magnesium has three isotopes % magnesium -24 with a mass of amu, 10.00% magnesium -25 with a mass of amu, and the rest magnesium -25 with a mass of amu. What is the atomic mass of magnesium?
If not told otherwise, the mass of the isotope is the mass number in amu

14. Electromagnetic radiation.

16. Electromagnetic radiation.
Most subatomic particles behave as PARTICLES and obey the physics of waves.

17. Electromagnetic radiation.
wavelength
Visible light
Ultaviolet radiation
Amplitude
Node

18. Electromagnetic radiation.
Waves have a frequency
Use the Greek letter “nu”, , for frequency, and units are “cycles per sec”
All radiation:  •  = c where c = velocity of light = 3.00 x 108 m/sec

19. Electromagnetic radiation.
Long wavelength --> small frequency
Short wavelength --> high frequency
increasing frequency
increasing wavelength

20. Electromagnetic Spectrum
In increasing energy, ROY G BIV

22. Prism
White light is made up of all the colors of the visible spectrum.
Passing it through a prism separates it.

23. If the light is not white
By heating a gas with electricity we can get it to give off colors.
Passing this light through a prism does something different.

24. Atomic Spectrum Each element gives off its own characteristic colors.
Can be used to identify the atom.
How we know what stars are made of.

25. These are called discontinuous spectra, or line spectra unique to each element.
These are emission spectra

26. Emission: gained energy is given off as electron/molecule goes back to ground state. We can see which energies are given off.

27. We also consider absorption spectra
Absorption: gain in energy to higher energy state. We can determine wavelengths absorbed by looking at the wavelengths that are missing (comparing continuous spectra to transmitted wavelengths)

28. Only certain amounts of energy (“quantum”) can be absorbed by molecules, atoms, electrons.
Because of this we say energy is quantized. It is only transferred between species in certain sized packets.
We can use absorption and emission of energy to identify species.

29. The energy absorbed by a species is unique because of the unique energy levels for each atom/element/molecule.
Spectroscopy involves the absorption and emission of electromagnetic radiation.

30. Explanation of atomic spectra
When we write electron configurations, we are writing the lowest energy.
The energy level, and where the electron starts from, is called it’s ground state- the lowest energy level.

31. Changing the energy
Let’s look at a hydrogen atom

32. Changing the energy
Heat or electricity or light can move the electron up energy levels (“excited”)

33. Changing the energy
As the electron falls back to ground state, it gives the energy back as light

34. Changing the energy May fall down in steps
Each with a different energy

35. { { {

37. Further they fall, more energy, higher frequency. This is simplified
Ultraviolet
Visible
Infrared
Further they fall, more energy, higher frequency.
This is simplified
the orbitals also have different energies inside energy levels
All the electrons can move around.

38. Ultraviolet
Visible
Infrared

40. Series:
Lyman – occurs in ultraviolet region (high energy). Transition is from upper energy level to first energy level
Balmer – occurs in visible region. Transitions occur between upper energy level to second energy level.
Paschen – occurs in infrared region (low energy). Transitions occur between upper energy level to third energy level.

41. ELECTRONIC CONFIGURATIONS
Before you start configurations it would be helpful to…
Know that electrons can be found outside the nucleus in energy levels ( shells)
Know the electronic configurations of the first 20 elements in 2,8,1 notation OR Electron Arrangement

42. 2,8,1 Notation (Electron Arrangement)
Simply write the number of electrons in each energy level of the atom, separated by a comma
Carbon 2,4
Magnesium 2,8,2
Chlorine 2,8,7

43. Maximum electrons per shell
THE BOHR ATOM
Ideas about the structure of the atom have changed over the years. The Bohr theory thought of it as a small nucleus of protons and neutrons surrounded by circulating electrons.
Each shell or energy level could hold a maximum number of electrons.
The energy of levels became greater as they got further from the nucleus and electrons filled energy levels in order.
The theory couldn’t explain certain aspects of chemistry.
Maximum electrons per shell
1st shell
2nd shell
3rd shell
4th shell
5th shell

44. Maximum electrons per shell
Ideas about the structure of the atom have changed over the years. The Bohr theory thought of it as a small nucleus of protons and neutrons surrounded by circulating electrons.
Each shell or energy level could hold a maximum number of electrons.
The energy of levels became greater as they got further from the nucleus and electrons filled energy levels in order.
The theory couldn’t explain certain aspects of chemistry.
Maximum electrons per shell
1st shell
2nd shell
3rd shell
4th shell
5th shell

45. PRINCIPAL ENERGY LEVELS INCREASING ENERGY / DISTANCE FROM NUCLEUS
LEVELS AND SUB-LEVELS
PRINCIPAL ENERGY LEVELS
During studies of the spectrum of hydrogen it was shown that the energy levels were not equally spaced. The energy gap between successive levels got increasingly smaller as the levels got further from the nucleus. The importance of this is discussed later. 4 3
INCREASING ENERGY / DISTANCE FROM NUCLEUS 2 1

46. PRINCIPAL ENERGY LEVELS INCREASING ENERGY / DISTANCE FROM NUCLEUS
LEVELS AND SUB-LEVELS
PRINCIPAL ENERGY LEVELS
SUB LEVELS
During studies of the spectrum of hydrogen it was shown that the energy levels were not equally spaced. The energy gap between successive levels got increasingly smaller as the levels got further from the nucleus. The importance of this is discussed later.
A study of Ionization Energies and the periodic properties of elements suggested that the main energy levels were split into sub levels.
Level 1 was split into 1 sub level
Level 2 was split into 2 sub levels
Level 3 was split into 3 sub levels
Level 4 was split into 4 sub levels 4 3
INCREASING ENERGY / DISTANCE FROM NUCLEUS 2 1
CONTENTS

47. RULES AND PRINCIPLES HEISENBERG’S UNCERTAINTY PRINCIPLE
“You cannot determine the exact location of an electron at any given time.”
This means that you cannot say exactly where an electron is. It put paid to the idea of electrons orbiting the nucleus in rings and introduced the idea of orbitals.
THE AFBAU PRINCIPLE
“Electrons enter the lowest available energy level.”
PAULI’S EXCLUSION PRINCIPLE
“No more than two electrons can occupy any one orbital.”
Two electrons can go in each orbital, providing they are of opposite spin.
HUND’S RULE OF MAXIMUM MULTIPLICITY
“If more than one orbital in a sublevel is available, electrons will occupy different orbitals with parallel spins.”
Placing two electrons in one orbital means that, as they are both negatively charged, there will be some electrostatic repulsion between them. Placing each electron in a separate orbital reduces the repulsion and the system is more stable.

48. ORBITALS
An orbital is... a region in space where one is likely to find an electron.
Orbitals can hold up to two electrons as long as they have opposite spin; this is known as PAULI’S EXCLUSION PRINCIPAL.
Orbitals have different shapes...

49. ORBITALS
An orbital is... a region in space where one is likely to find an electron.
Orbitals can hold up to two electrons as long as they have opposite spin; this is known as PAULI’S EXCLUSION PRINCIPAL.
Orbitals have different shapes...
ORBITAL SHAPE OCCURRENCE
s spherical one in every principal level
p dumb-bell three in levels from 2 upwards
d various five in levels from 3 upwards
f various seven in levels from 4 upwards

50. DO NOT CONFUSE AN ORBITAL WITH AN ORBIT
ORBITALS
An orbital is... a region in space where one is likely to find an electron.
Orbitals can hold up to two electrons as long as they have opposite spin; this is known as PAULI’S EXCLUSION PRINCIPAL.
Orbitals have different shapes...
ORBITAL SHAPE OCCURRENCE
s spherical one in every principal level
p dumb-bell three in levels from 2 upwards
d various five in levels from 3 upwards
f various seven in levels from 4 upwards
An orbital is a 3-dimensional statistical shape showing where one is most likely to find an electron. Because, according to Heisenberg, you cannot say exactly where an electron is you are only able to say where it might be found.
DO NOT CONFUSE AN ORBITAL WITH AN ORBIT

51. s orbitals SHAPES OF ORBITALS spherical
one occurs in every principal energy level

52. p orbitals SHAPES OF ORBITALS dumb-bell shaped
three occur in energy levels except the first

53. d orbitals SHAPES OF ORBITALS various shapes
five occur in energy levels except the first and second

54. ORDER OF FILLING ORBITALS
INCREASING ENERGY / DISTANCE FROM NUCLEUS 1 1s 2 2s 2p 4s 3 3s 3p 3d 4 4p 4d 4f
PRINCIPAL ENERGY LEVELS
SUB LEVELS
Orbitals are not filled in numerical order because the principal energy levels get closer together as you get further from the nucleus. This results in overlap of sub levels. The first example occurs when the 4s orbital is filled before the 3d orbitals.

55. ORDER OF FILLING ORBITALS
INCREASING ENERGY / DISTANCE FROM NUCLEUS 1 1s 2 2s 2p 4s 3 3s 3p 3d 4 4p 4d 4f
PRINCIPAL ENERGY LEVELS
SUB LEVELS 1 1s 2 2s 2p 3d 3 3s 3p 4s 4 4p 4d 4f
PRINCIPAL ENERGY LEVELS
SUB LEVELS
Orbitals are not filled in numerical order because the principal energy levels get closer together as you get further from the nucleus. This results in overlap of sub levels. The first example occurs when the 4s orbital is filled before the 3d orbitals.

56. ORDER OF FILLING ORBITALS
INCREASING ENERGY / DISTANCE FROM NUCLEUS 1 1s 2 2s 2p 4s 3 3s 3p 3d 4 4p 4d 4f
PRINCIPAL ENERGY LEVELS
SUB LEVELS 1 1s 2 2s 2p 3d 3 3s 3p 4s 4 4p 4d 4f
PRINCIPAL ENERGY LEVELS
SUB LEVELS
THE FILLING ORDER 1s
2s 2p
3s 3p 3d
4s 4p 4d 4f
5s 5p 5d 5f
6s 6p 6d
7s 7p
HOW TO REMEMBER ...
Orbitals are not filled in numerical order because the principal energy levels get closer together as you get further from the nucleus. This results in overlap of sub levels. The first example occurs when the 4s orbital is filled before the 3d orbitals.

57. THE ‘AUFBAU’ PRINCIPAL4 4p 4d 4f
This states that…
“ELECTRONS ENTER THE LOWEST AVAILABLE ENERGY LEVEL” 3 3s 3p 3d
The following sequence will show the ‘building up’ of the electronic structures of the first 36 elements in the periodic table.
Electrons are shown as half headed arrows and can spin in one of two directions or
s orbitals
p orbitals
d orbitals 4s
INCREASING ENERGY / DISTANCE FROM NUCLEUS 2 2s 2p 1 1s

58. 1s1 THE ELECTRONIC CONFIGURATIONS OF THE FIRST 36 ELEMENTS 4f HYDROGEN4p 4d 4f
HYDROGEN
1s1
Hydrogen atoms have one electron. This goes into a vacant orbital in the lowest available energy level. 3 3s 3p 3d 4s
INCREASING ENERGY / DISTANCE FROM NUCLEUS
‘Aufbau’
Principle 2 2s 2p 1 1s

59. 1s2 THE ELECTRONIC CONFIGURATIONS OF THE FIRST 36 ELEMENTS 4 4p 4d 4f
HELIUM
1s2
Every orbital can contain 2 electrons, provided the electrons are spinning in opposite directions. This is based on...
PAULI’S EXCLUSION PRINCIPLE
The two electrons in a helium atom can both go in the 1s orbital. 3 3s 3p 3d 4s
INCREASING ENERGY / DISTANCE FROM NUCLEUS 2 2s 2p 1 1s
‘Aufbau’
Principle

60. 1s2 2s1 THE ELECTRONIC CONFIGURATIONS OF THE FIRST 36 ELEMENTS 4f4p 4d 4f
LITHIUM
1s2 2s1
1s orbitals can hold a maximum of two electrons so the third electron in a lithium atom must go into the next available orbital of higher energy. This will be further from the nucleus in the second principal energy level.
The second principal level has two types of orbital (s and p). An s orbital is lower in energy than a p. 3 3s 3p 3d 4s
INCREASING ENERGY / DISTANCE FROM NUCLEUS 2 2s 2p 1 1s
‘Aufbau’
Principle

61. 1s2 2s2 THE ELECTRONIC CONFIGURATIONS OF THE FIRST 36 ELEMENTS 4f4p 4d 4f
BERYLLIUM
1s2 2s2
Beryllium atoms have four electrons so the fourth electron pairs up in the 2s orbital. The 2s sub level is now full. 3 3s 3p 3d 4s
INCREASING ENERGY / DISTANCE FROM NUCLEUS 2 2s 2p
‘Aufbau’
Principle 1 1s

62. 1s2 2s2 2p1 THE ELECTRONIC CONFIGURATIONS OF THE FIRST 36 ELEMENTS 4f4d 4f
BORON
1s2 2s2 2p1
As the 2s sub level is now full, the fifth electron goes into one of the three p orbitals in the 2p sub level. The 2p orbitals are slightly higher in energy than the 2s orbital. 3 3s 3p 3d 4s
INCREASING ENERGY / DISTANCE FROM NUCLEUS 2 2s 2p
‘Aufbau’
Principle 1 1s

63. 1s2 2s2 2p2 THE ELECTRONIC CONFIGURATIONS OF THE FIRST 36 ELEMENTS 4f4d 4f
CARBON
1s2 2s2 2p2
The next electron in doesn’t pair up with the one already there. This would give rise to repulsion between the similarly charged species. Instead, it goes into another p orbital which means less repulsion, lower energy and more stability. 3 3s 3p 3d 4s
INCREASING ENERGY / DISTANCE FROM NUCLEUS 2 2s 2p
HUND’S RULE OF
MAXIMUM MULTIPLICITY 1 1s

64. 1s2 2s2 2p3 THE ELECTRONIC CONFIGURATIONS OF THE FIRST 36 ELEMENTS 4f4d 4f
NITROGEN
1s2 2s2 2p3
Following Hund’s Rule, the next electron will not pair up so goes into a vacant p orbital. All three electrons are now unpaired. This gives less repulsion, lower energy and therefore more stability. 3 3s 3p 3d 4s
INCREASING ENERGY / DISTANCE FROM NUCLEUS 2 2s 2p
HUND’S RULE OF
MAXIMUM MULTIPLICITY 1 1s

65. 1s2 2s2 2p4 THE ELECTRONIC CONFIGURATIONS OF THE FIRST 36 ELEMENTS 4f4d 4f
OXYGEN
1s2 2s2 2p4
With all three orbitals half-filled, the eighth electron in an oxygen atom must now pair up with one of the electrons already there. 3 3s 3p 3d 4s
INCREASING ENERGY / DISTANCE FROM NUCLEUS
‘Aufbau’
Principle 2 2s 2p 1 1s

66. 1s2 2s2 2p5 THE ELECTRONIC CONFIGURATIONS OF THE FIRST 36 ELEMENTS 4f4d 4f
FLUORINE
1s2 2s2 2p5
The electrons continue to pair up with those in the half-filled orbitals. 3 3s 3p 3d 4s
INCREASING ENERGY / DISTANCE FROM NUCLEUS 2 2s 2p 1 1s

67. 1s2 2s2 2p6 THE ELECTRONIC CONFIGURATIONS OF THE FIRST 36 ELEMENTS 4f4d 4f
NEON
1s2 2s2 2p6
The electrons continue to pair up with those in the half-filled orbitals. The 2p orbitals are now completely filled and so is the second principal energy level.
In the older system of describing electronic configurations, this would have been written as 2,8. 3 3s 3p 3d 4s
INCREASING ENERGY / DISTANCE FROM NUCLEUS 2 2s 2p 1 1s