1. Chemistry Unit Three
Atomic Structure

2. A Brief History of the
Atomic Theory

3. Democritus 400 B.C. Greek philosopher
Coined the term “atomos” which means, Indivisible.
All matter is made of atoms.
Atoms are hard, solid particles, made of the same material but are of different shapes and sizes.

4. John Dalton 1803 English Chemist Atoms are solid, neutral spheres.
Atoms of same element are the same.
Atoms of different elements are different.
Compounds form from the joining of atoms of two or more elements.

5. J.J. Thomson Cathode Ray Experiments; discovered the electron.
1897 English Chemist.
Atoms are made of even smaller particles.
Called the Plum Pudding Model (Chocolate Chip Cookie Dough Model)
Positively charged material through which negative particles are scattered.
Since atoms are neutral, therefore, there must be (+) particles too, but Thomson never found them.
Cathode Ray Experiments; discovered the electron.

6. Ernest Rutherford . 1911 British physicist Gold Foil experiment.
Atom has a small, dense positively charged center called the Nucleus.
Negative electrons are scattered outside the nucleus.
Most of the atom is empty space
If an atom was the size of a baseball stadium, the nucleus would be the size of a marble. .

7. Rutherford’s
Gold Foil Experiment

8. A beam of + particles (alpha particles) shot through a thin sheet of gold foil.
Most particles passed straight through. (Most of atom is empty space.)
A few were deflected. (Positive core-similar charges repel each other.)
Very few bounced off. (Solid core is very small.)

9. Neils Bohr Labeled each energy level by a quantum number.
1913 Danish Scientist
Planetary model.
Electrons are held in place by the attraction between them and the + charged nucleus.
Each electron occupies a specific energy level and orbit the nucleus like planets circling the sun.
Labeled each energy level by a quantum number.

10. Wave Model
Electrons are not discreet particles moving in discreet orbits.
The probable location of an electron depends on how much energy it has.
Electrons seem to be everywhere at once, like the moving blades of a fan.

11. Electron Cloud Model Quantum Model
Positively charged protons and neutral neutrons are held together with a huge amount of energy forming the nucleus of the atom.
Negatively charged electrons move rapidly around the outside of the nucleus forming “clouds” of negative charge.
Most of the mass of the atom is in the nucleus.
Quantum Model

12. Summary of Atomic Models

13. Has no charge (neutral)
Atomic Structure
Protons
(p+)
Neutrons
(n)
Electrons
(e-)
Found in the Nucleus
Found in the Nucleus
Found outside the nucleus
Has 1 amu of mass
Has 1 amu of mass
Has 0 amu of mass
Has a positive charge
Has no charge (neutral)
Has a negative charge

14. Atomic Structure Atom Isotope Ion
The number of protons in an atom never changes.
Isotope
Atoms of the same element that have different numbers of neutrons.
Ion
An atom that has lost or gained electrons.

15. Atomic Mass
Average Atomic Mass – the average mass of all of the isotopes of an element. Decimal number.
Mass Number – the total number of protons and neutrons in the nucleus of an atom. Whole number.

16. Atomic Structure
Atomic Number = the number of protons. Equal to the number of electrons.
Atomic Mass = the number of protons and neutrons added together.
Atomic Mass – Atomic Number = the number of neutrons.

17. C Calculating Numbers of Protons, Neutrons and Electrons. 6 12.011
Atomic Number = 6
6 protons = 6 electrons
6 p+ = 6 e-
(atom is neutral) 6 C
12.011
Atomic Mass = 12
12 p+ and n
-6 p+
6 neutrons

18. Practice Calculating p+, n, e-
Element
Atomic number
Mass Number
# proton
# electron
# neutron
Silver 47 61
108 47 47
*Atomic # is number of protons so protons = 47
*number (+) charges (p+) must equal (–) charges to make the atom neutral so electrons = 47
*Mass Number is total particles with mass (p+ and n) so = 108

19. Practice Calculating p+, n, e-
Element
Atomic number
Mass Number
# proton
# electron
# neutron
Copper 64 29 35 29 29
*number of protons is the atomic # so atomic number is 29
*number (+) charges (p+) must equal (–) charges to make the atom neutral so electrons = 29
*Mass Number is total of all particles with mass (p+ and n) so subtract away the atomic number (#p+) and you will have just neutrons (64 – 29 = 35)

20. Practice Calculating p+, n, e-
Element
Atomic number
Mass Number
# proton
# electron
# neutron
Tin 50 69
Helium 4 2
Boron 11 6 50
119 50
p+ + n
+ = - (neutral)
Is # p+ 2 2 2
Is # p+
+ = - (neutral)
Mass # - atomic # 5 5 5
Is # p+
Mass #- n
+ = - (neutral)
Remember e- = p+ (to make atom neutral)
 #p+ is atomic number
 p+ + n = mass number
 Mass number – atomic number = n

21. Bohr Models e- e- e- e- e- e- e- e- e- e- e- e- e- e- e- e- e- e-
p+ & n in nucleus
e- in energy levels around nucleus
3 energy levels
-1st has up to 2e-
-2nd has up to 8e-
-3rd has up to 8e- e- e- e- e- e- e- e- e- e- e- e- e-

22. Bohr Model of Lithium e- e- e- 3 Li
6.941 3 4

23. e- e- e- e- e- e- e- e- e- e- e- e- e- e- e- e- e- e-
Bohr Model of Argon e- e- e- e- 18 Ar
39.948 e- e- e- e- 18 e- e- e- e- e- e- 22 e- e- e- e-

24. Electron Configuration
Shows the distribution of electrons among the orbitals of an atom.
Describes where the electrons are located and how much energy each one has.

25. Rules for Electron Configuration
Aufbau Principle - Electrons enter orbitals of lowest energy level first.
Pauli Exclusion Principle – An orbital can hold a maximum of 2 electrons. To occupy the same orbital, the 2 electrons must spin in opposite directions.
Hund’s Rule - one electron enters each orbital until each orbital contain one electron with parallel spins before a second electron is added.

26. Determining Electron Configurations
Quantum Numbers describe the amount of energy in that level. The lower the number, the less energy it has. (n = 1, 2, 3, 4, etc.)
Sublevels are divisions of the principle energy levels. The main sublevels are called s, p, d and f.
Each sublevel has a different shape caused by the different energy levels.

27. Number of Electrons per Sublevel
Sublevel Number of Maximum
Orbitals # of e- s p d f

28. s and p Orbital Electrons

29. d Orbital Electrons

30. f Orbital Electrons

31. Periodic Table to remember order
Sublevels (s,p,d,f) by columns - Energy levels by rows (1,2,3,4,5,6,7 except d(row-1) & f(row-2))
X – 1s22s22p63s23p4 (16 e-) s1
Y – 1s22s22p63s23p6 4s23d104p6 5s1 (36 e-) s2 s2 p1 p2 p3 p4 p5 1 p6 2 d1 d2 d3 d4 d5 d6 d7 d8 d9
d10 X 3 4 5 Y 6 7 f1 f2 f3 f4 f5 f6 f7 f8 f9
f10
f11
f12
f13
f14

32. Example 1 He (atomic # = 2) (Which means 2 p+ = 2e-)
1s2 (1=energy level; s=sublevel; 2=electrons) He

33. Example 2 Li (atomic #=3) (means 3p+ = 3e-)
1s22s1 (1,2=energy levels; s=sublevel; 2+1=electrons) Li

34. Example 3 Be (atomic #=4) (means 4p+ = 4e-)
1s22s2 (1,2=energy levels; s=sublevel; 2+2=electrons) Be

35. Example 4 Si (atomic #=14) (means 14p+ = 14e-) 1s22s22p63s23p2
(1,2,3=energy levels; s,p=sublevels; =electrons) Si

36. Lewis Dot Structures One more type of atomic model…
(In addition to Bohr models and electron configurations)
Consists of the element’s symbol and the atom’s valence electrons.
Symbol = kernel (represents the protons, neutrons and full electron shells).
Dots = valence electrons.

37. B B = Kernel (The protons, neutrons and full electron shells.)
Lewis Dot Structures Con’t
B = Kernel (The protons, neutrons and full electron shells.) B
Valence shell electrons

38. You can use the Electron Configuration to get the Lewis Dot Structure…
1s22s22p63s23p64s2
Locate the highest quantum number. (4)
Add the s and p orbital electrons, and place them around the element symbol. (2) Ca

39. One Final Example Sn Tin 1s22s22p63s23p64s23d104p65s24d105p2
Locate the highest quantum number (5)
Add the s and p orbital electrons (4) Sn

40. How to place electrons on a Lewis Dot
First two dots represent the s orbital electrons and are placed at the top of the element’s symbol.
Then the p orbital electrons are placed in this order: right, bottom, left, right, bottom, left.

41. So, it goes like this… 1 2 8 3 Ne 5 6 7 4

42. Percent Abundance
The percentage of how much one specific isotope of an element is found in nature.
FORMULA:
% abundance = amount of one isotope
total amount of all isotopes

43. Average Atomic Mass (How the number ends up on the periodic table!!)
1st  Mass of one isotope x % abundance
in decimal form (watch SIG FIGS!!)
2nd  Do this for each isotope of that element
3rd  Then add all individual isotopes
together to get the average atomic
mass.

44. 1. Calculate the average atomic mass of potassium using the following data:
Isotope
Mass
% abundance
Potassium-39
amu
93.12%
Potassium-41
amu
6.88 %
Potassium-39
amu x
0.9312 =
36.28 amu
Potassium-41
amu x
0.0688 =
2.82 amu +
Average atomic mass for K = amu

45. 2. Calculate the average atomic mass of magnesium using the following data:
Isotope
Mass
% abundance
Magnesium-24
amu
78.70%
Magnesium-25
amu
10.13 %
Magnesium-26
amu
11.17 %
Magnesium-24
amu x
0.7870 =
18.88 amu
Magnesium-25
amu x
0.1013 =
2.531 amu + x
Magnesium-26
amu
0.1117 =
2.902 amu +
Average atomic mass for K = amu