1. Atomic Structure

2. Introduction
Laws of nature that apply to atoms govern behavior and changes in the universe
We can not see atoms, so we get our facts about atoms from indirect evidence – behavior not appearance
Models of atoms were developed to make things easier
These models were revised every time we got new information

3. Democritus 400 BC Greek philosopher
Wondered if matter could be divided
First to come up with idea of atoms

4. Indivisible models
Dalton – an English scientist that lived
Studied meteorology and chemistry
Experimented with gases
Combined results and in 1803 developed the atomic theory of matter
All matter is made of atoms
All atoms of an element are alike.
Atoms cannot be created or destroyed
Atoms combine in whole – number ratios to form compounds

5. Divisible Models Thomson – English – 1897
Studied the uncharged gases and noticed that a negative charged particle was being emitted.
Decided this particle must have come from within the atom.
Named electrons.
Also concluded that positive particles must exist but never found them.
Measured charge to mass ratio of electrons

6. Thomson’s Atomic Model
Thomson believed that the electrons were like plums embedded in a positively charged “pudding,” thus it was called the “plum pudding” model.

8. Conclusions from the Study of the Electron
Cathode rays have identical properties regardless of the element used to produce them. All elements must contain identically charged electrons.
Atoms are neutral, so there must be positive particles in the atom to balance the negative charge of the electrons
Electrons have so little mass that atoms must contain other particles that account for most of the mass

9. Rutherford - 1910 Discovered nucleus and that it was positive
Gold foil experiment
1) most of atom is empty space (majority of particles went straight through)
2) nucleus is small, dense and positively charged (some positive charges were greatly deflected)

10. Rutherford’s Gold Foil Experiment
Alpha particles are helium nuclei
Particles were fired at a thin sheet of gold foil
Particle hits on the detecting screen (film) are recorded

11. Rutherford’s Findings
Most of the particles passed right through
A few particles were deflected
VERY FEW were greatly deflected
“Like howitzer shells bouncing off of tissue paper!”
Conclusions:
The nucleus is small
The nucleus is dense
The nucleus is positively charged

12. Niels Bohr - 1915 Proposed early model of atom
“Planetary model” electrons orbit nucleus like planets orbit sun
Lacks math of modern version
Has some errors/violates current theory
Radiation is emitted when electrons move from one orbit to another

13. Divisible Models Wave Model – 1926 Quantum Mechanics
Based on wave calculations and quantum mechanics
Cannot predict exact location of electrons, only region it is most likely to be in
Chadwick – 1932 – proves the existence of neutrons

14. Basic Atomic Structure – Subatomic Particles
Protons (p+)
Positive charge
Mass = 1 amu (atomic mass unit)
Identical to all other protons in all other elements
Located in the nucleus so it is a nucleon
Atomic number = # of protons in nucleus – identifies element

15. Basic Atomic Structure – Subatomic Particles
Neutrons (n)
Neutral charge
Mass = 1 amu
Identical to all other neutrons in all other elements
located in the nucleus so it is a nucleon
Mass number – sum of protons and neutrons in nucleus – always a whole number

16. Basic Atomic Structure – Subatomic Particles
Electrons e-
Outside nucleus in energy levels; aka electron cloud
Negative charge
Mass = 1/1836 amu (0 for calculation purposes)

17. Basic Atomic Structure – Subatomic Particles
Isotope
atoms of the same element that have the same number of protons but different number of neutrons
Protons + neutrons = mass number, so isotopes have different mass
gives different properties to element
Example: carbon -12 and carbon -14; carbon 14

18. Isotopes Isotope Protons Electrons Neutrons Nucleus Hydrogen–1
(protium) 1
Hydrogen-2
(deuterium)
Hydrogen-3
(tritium) 2

19. Isotope symbols
Hyphen notation Nuclear notation Carbon – 12 12C 6p+ and 6no Carbon – 13 13C 6p+ and 7no Carbon – 14 14C 6p+ and 8no 6 6 6

20. Atomic Number Element # of protons Atomic # (Z) Carbon 6 Phosphorus 15
Atomic number (Z) of an element is the number of protons in the nucleus of each atom of that element.
Element
# of protons
Atomic # (Z)
Carbon 6
Phosphorus 15
Gold 79

21. Mass Number Mass # = p+ + n0 Nuclide p+ n0 e- Mass #
Mass number is the number of protons and neutrons in the nucleus of an isotope.
Mass # = p+ + n0
Nuclide p+ n0 e-
Mass #
Oxygen - 10 - 33 42
- 31 15 18 8 8 18
Arsenic 75 33 75
Phosphorus 16 15 31

22. Composition of the nucleus
Atomic Masses
Atomic mass is the average of all the naturally isotopes of that element.
Carbon =
Isotope
Symbol
Composition of the nucleus
% in nature
Carbon-12
12C
6 protons
6 neutrons
98.89%
Carbon-13
13C
7 neutrons
1.11%
Carbon-14
14C
8 neutrons
<0.01%

23. How to Calculate the Average Mass
What is the average atomic mass of sample of Cesium with 3 isotopes:
75% 133Cs, 20% 132Cs, and 5% 134Cs.
0.75 x 133 = 99.75
0.20 x 132 =
0.05 x 134 =
Total =
avg. atomic mass

24. Basic Atomic Structure – Subatomic Particles
Ion
Charged atom – an atom with an unequal number of protons and electrons
Cation: + ion (has more protons)
Anion: - ion (has more electrons)

25. Charge number Charge # = p+ - e -
Charge number is the number of protons minus the number of electrons in an ion. This occurs because protons are positive and electrons are negative.
Charge # = p+ - e -
Ion p+ e-
Charge #
Oxide 10
Magnesium Ion 12 +2
Bromide 15 -1 8 -2 10 16

26. Energy Levels
Within the electron cloud, the location of the electrons depend on its energy level
Low energy electrons are close to the nucleus
High energy electrons are farther from the nucleus
Energy levels can only hold a certain number of electrons
Electrons have to gain or lose a specific amount of energy to move from one level to another
They have to gain a quanta of energy
The chemical properties of an element depend on how many electrons are in each energy level.
This is especially true about the outer level. The electrons in the outer level are called valence electrons.

27. 1s Up to 2 2 s p Up to 8 3 s p d Up to 18 4 s p d f Up to 32
Quantum Energy Level Name
Number of Electrons 1s
Up to 2
2 s p
Up to 8
3 s p d
Up to 18
4 s p d f
Up to 32