1. Unit 1:Atomic Structure Part 2
CH1120

2. Subatomic Particles Atoms are composed of
Protons
Electrons
Neutrons
The amount of these subatomic particles alters the identity of the atom

3. Subatomic Particles

4. Protons
Positively charged
Relative mass of 1
Found in nucleus

5. Neutrons
Neutral
Relative mass of 1
Found in nucleus

6. Electrons Negatively charged Negligible mass
Found outside of the nucleus

7. Subatomic Particles
When an atom has an equal number of protons and electrons it is neutral
Has no charge

8. Ions When an atom has more protons than electrons it is a cation
Positively charged, name is same as element
When an atom has more electrons than protons it is an anion
Negatively charged, name ends in –ide (ie. Chloride)
An atom with a charge is called an ion

9. Ions
Atoms gain and lose electrons in interactions to have the same number of electrons as the closest noble gas
Na+ loses 1 electron to become like Ne
S2- gains 2 electrons to become like Ar
Carbon can gain or lose 4 electrons to become like Ne or He

10. Atomic Number Number of protons in the nucleus of an atom
Found on the periodic table

11. Mass Number
Total number of protons and neutrons in the nucleus of an atom
Mass number is always larger than atomic number

12. Isotopes
Isotopes are atoms of the same element that have varying masses
Remember we can change the number of neutrons without changing the element
This simply changes the mass of the atom

13. Isotopes Changing number of protons changes the element
Atomic number
Changing number of neutrons changes the mass
Creates different isotope
Changing the number of electrons changes the charge

14. Isotopes

15. Isotopes

16. Isotopes When given names like Carbon-13, 13 is the mass number
Neutrons and protons

17. Example Draw the isotope symbol of Hydrogen-3 with a charge of +1.
How many of each of the subatomic particles are present?

18. Isotopes

19. Isotopes

20. Isotopes Smoke detectors Archaeological dating Medical imaging
Americium-241
Archaeological dating
Carbon-14
Medical imaging
Medical treatment

21. Average Atomic Mass Atomic mass and mass number are NOT the same thing
Mass number = number of protons and neutrons
Whole number
Atomic mass = average atomic mass/weight

22. Average Atomic Mass
The average of all of the isotopes of a particular element found in nature
This can be found on the periodic table or can be calculated from data given

23. Average Atomic Mass
Average Atomic Mass = [M1 x %Abundance1] + [M2 x %Abundance2] + ...

24. Average Atomic Mass
Gallium has two naturally occurring isotopes: Ga-69 with mass amu and abundance of 60.11% and Ga-71 with mass amu and abundance of 39.89%. Calculate the atomic mass of gallium.

25. Average Atomic Mass
Bromine has two naturally occurring isotopes (Br-79 and Br-81) and an average atomic mass of amu. If the mass of Br-81 is amu, what is the mass of Br-79? The natural abundance of Br-79 is 50.69%.

26. Atomic Orbitals
Electrons of an atom are attracted to the positive nucleus of the atom
Electrons are not all the same distance from the nucleus
Therefore they have different energies
Further away = greater energy

27. Quantum Mechanical Model
Electrons can be grouped according to their ease of removal from the atom
Ease of removal of electrons depends on their distance from the nucleus

28. Quantum Mechanical Model

29. Quantum Mechanical Model
Electrons are found in energy levels called shells
Each shell is broken down into subshells
Each subshells has one or more atomic orbitals with a specific 3D shape and energy
Think of orbitals as clouds where electrons live

30. Quantum Mechanical Model

31. Quantum Mechanical Model
F Orbitals

32. Quantum Numbers Principle Quantum Number (n)
Identifies energy level (therefore number of subshells)
Positive integral numbers (1 and up)
As “n” increases, orbital size increases, electron energy increases

33. Quantum Numbers Angular momentum quantum number (l )
Defines the shape of the orbital
All value 0 to n-1
Letters are used to represent different values of l
0 = s
1 = p
2 = d
3 = f

34. Quantum Numbers Magnetic quantum number (ml )
Describes 3D orientation of orbital in space
Values from –l to +l

35. Quantum Numbers Spin quantum number (ms)
Identifies the rotation of the electron
No 2 electrons can have all 4 quantum numbers the same
Values are +1/2 and -1/2

36. Quantum Numbers

37. Quantum Mechanical Model
Electron shell: collection of orbitals with the same principle quantum number (n)
3s, 3p, 3d, and 3f, are all in the 3rd electron shell
Subshell: collection of orbitals within a shell
3d is a subshell of the 3rd electron shell

38. Quantum Mechanical Model
The Principle Quantum Number (n) determines how many subshells are present
n = 1 has one subshell, n = 2 has two subshells, etc.
Each subshell consists of a specific number of orbitals
s = 1, p = 3, d = 5; f = 7
The total number of orbitals in a given shell is n2
Ex. n = 1 has one orbital, n = 2 has four orbitals, etc.

39. Orbital Shapes s-orbitals Begin at the first energy level Spherical
Hold a maximum of 2 electrons

40. Orbital Shapes p-orbitals Begin at the second energy level
3 p-orbitals at each energy level
Each orbital holds 2 electrons (total of 6 at each energy level)

41. Orbitals Shapes d-orbitals Begin at third energy level
5 d-orbitals at each energy level
Each orbital holds 2 electrons (total of 10 at each energy level)

42. Orbital Shapes f-orbitals Begin at the fourth energy level
7 f-orbitals at each energy level
Each orbital hols 2 electrons (total of 14 at each energy level)

43. Filling Orbitals
When filling orbitals, electrons fill those with the lowest energy first
Orbitals with the same energy are said to be degenerate
Ex. All orbitals in the 3p subshell are degenerate
Repulsions between electrons cause the subshells in a given shell to be at different energies

44. Filling Orbitals
Generally, a higher n indicates a higher energy and s<p<d<f (s having the lowest and f the highest energy)
But when you move past p, things get more complicated

45. Aufbau Principle
Electrons will assume the most stable position based on the nucleus of an atom and the electrons already present

46. Pauli Exclusion Principle
No two electrons can have the same exact 4 quantum numbers
Orbitals can hold a maximum of 2 electrons therefore they must have opposite spins

47. Electron Configurations
Each component consists of:
A number denoting the energy level
A letter denoting the type of orbital
A superscript denoting the number of electrons in those orbitals

48. Electron Configurations
Remember each orbital holds a maximum of 2 electrons
s-orbitals have 1 subshell
p-orbitals have 3 subshells
d-orbitals have 5 subshells
f-orbitals have 7 subshells

49. Electron Configurations
1. Determine total number of electrons for atom
2. Start filling in orbitals according to the order of increasing energy (we will look at 2 methods)
3. Continue until all electrons are used up in the orbitals

50. Electron Configurations

51. Electron Configurations

52. Electron Configurations
Write the electron configurations for the following elements:
Oxygen
Iron

53. Electron Configurations
Write the electron configurations for the following elements: H B Cl

54. Orbital Diagrams Each box or line represents one orbital
Half-arrows represent the electrons
The direction of the arrow represents the relative spin of the electron
Example: lithium

55. Hund’s Rule
For degenerate orbitals, the lowest energy is attained when the number of electrons with the same spin is maximized
Example: oxygen

56. Orbital Diagrams 1. Determine total number of electrons for atom
2. Start filling in orbitals according to the order of increasing energy
3. Apply Hund’s rule for the subshells until all electrons are used up in the orbitals
Place 1 electron in each orbital until all contain 1 electron, then start pairing

57. Orbital Diagrams Draw orbital diagrams for the following elements: BeNa

58. Condensed Electron Configurations
Write the nearest noble gas with lower atomic number in square brackets
Begin the electron configuration from that point

59. Condensed Electron Configurations
Write the condensed electron configuration for Fluorine.

60. Condensed Electron Configurations
Write condensed electron configurations for the following elements:
Beryllium
Chlorine
Cobalt

61. Valence Electrons Electrons filling the outermost energy levels
Used for chemical bonding
Group A elements, the group number is the number of valence electrons

62. Valence Electrons
How many valence electrons are in the following elements:
Fluorine
Magnesium
Carbon

63. Valence Electrons