1. Atomic Structure and Bonding
Unit 2

2. 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

3. 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)

4. 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

5. ATOMIC NUMBER AND MASS NUMBER Number of electrons = Number of protonsHe
Mass Number 4
the number of protons and neutrons in an atom 2
Atomic Number
the number of protons in an atom
Number of electrons = Number of protons
in a neutral atom 5

6. 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

7. 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-

8. When naming, write the mass number after the name of the element
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

9. 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

10. Average Atomic Masses
If not told otherwise, the mass of the isotope is the mass number in ‘u’
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

11. 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?

12. Bohr
Proposed electrons (e-) orbit around the nucleus in circular paths
Said e- in a particular path have a fixed energy (energy levels)
e- can go from any energy level to another by gaining or losing a specific amount of energy = a “quantum of energy”
When e- absorbs a quantum of energy, it goes from it’s ground state (where it’s normally found) to an excited state
The excited state is at a higher energy level

13. 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

14. 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
( )

15. Spectrum of White Light

16. Spectrum of Excited Hydrogen Gas

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

18. Drawback to Bohr
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

19. Energy level populations (Science10)
Electrons found per energy level of the atom.
The first energy level holds 2 electrons
The second energy level holds 8 electrons
The third energy level holds 18 electrons

20. Examples for group 1
Li 2.1 Na K

21. 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 – an ORBITAL

22. Orbits (Bohr) vs Orbitals (Quantum Mechanics)
Bohr said electrons travel in an orbit – can predict exact location of electron at any point in time.
Schrodinger used mathematics (calculus) to find the region in space where an electron will be found 90% of the time - this region is called an orbital. There are 4 main types of orbitals – s, p, d, and f.

23. Modern View of the Atom
The modern view of the atom suggests that the atom is more like a cloud. Atomic orbitals around the nucleus define the places where electrons are most likely to be found. 23

24. 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

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

26. 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

27. 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

28. Electron Configuration
We use e- configuration as a shorthand to show how e- are arranged around a nucleus
Example: Carbon is …
1s2 2s2 2p2

29. 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 .

30. Summary
Sublevel
# of shapes
(Orbitals)
Max number of e-
Starts at energy level s 1 2 p 3 6 d 5 10 f 7 14 4

31. Electron Arrangement 1st Rule: The Aufbau Principle
e- fill orbitals of the lowest energy first
We can use the periodic table to help us!

32. The Diagonal Rule

33. Example #1
1 s
2 s
2 p
Oxygen
1s2
2s2
2p4

34. Example #2
1 s
2 s
2 p
3 s
Magnesium
1s2
2s2
2p6
3s2

35. Example #3
1 s
2 s 3d
2 p
3 s
3 p
4 s
Iron
1s2
2s2
2p6
3s2
3p6
4s2
3d6

36. Practice
Boron
Argon
Calcium
Iodine
Sodium
Zinc
Lead

37. Abbreviations
We can abbreviate electron configurations using the Noble Gases
Ex: Sulfur
1s2 2s2 2p6 3s2 3p4
[Ne] 3s2 3p4
Ex: Lead
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14 5d10 6p2
[Xe] 6s2 4f14 5d10 6p2

38. 2nd Rule: Pauli Exclusion Principle
Each orbital orientation can hold up to 2 e-
e- must have opposite spins (up/clockwise or down/counter clockwise)
Therefore:
s has up to 2 e- (1 orientation)
p has up to 6 e- (3 orientations)
d has up to 10 e- (5 orientations)
f has up to 14 e- (7 orientations)
We can use the 2nd rule to draw Orbital Diagrams

39. Example #1
Oxygen: 1s2 2s2 2p4
1s s p

40. Example #2
Magnesium: 1s2 2s2 2p6 3s2
1s s p s

41. Example #3
Iron
1s2
2s2
2p6
3s2
3p6
4s2
3d6

42. 1s 2s 2p 3rd Rule: Hund’s Rule
e- will not pair up until each orbital orientation has 1 e- in it
The first e- in a pair will spin up, the second will spin down
Example: Oxygen is 1s2 2s2 2p4
1s s p

43. Orbital Notation
Orbital Notation shows us visually the arrangement and spin of electrons
Example: Carbon is 1s2 2s2 2p2
1s s p

44. Energy Level Diagrams 2p 2s 1s
Energy Level Diagrams give us the same information as orbital diagrams, plus they show us the different energy levels of each orbital
Example: Carbon is 1s2 2s2 2p2 2p 2s 1s

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

46. 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

47. 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

48. 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

49. 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

50. 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

51. 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

52. Write these electron configurations
Titanium - 22 electrons
1s22s22p63s23p64s23d2
Vanadium - 23 electrons
1s22s22p63s23p64s23d3
Chromium - 24 electrons
1s22s22p63s23p64s23d4

53. 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