1. Introduction to the Atom and Atomic Models

2. Atomic - Molecular Theory of Matter
The Atomic - Molecular Theory of Matter states that all matter is composed of small, fast moving particles called atoms. These atoms can join together to form molecules.
This theory is really thousands of individual ideas and models that provide evidence for the whole theory.

3. Matter
Since the atom is too small to be seen even with the most powerful microscopes, scientists rely upon models to help us to understand the atom.
Believe it or not this is a microscope. Even with the world’s best microscopes we cannot see the structure or behavior of the atom.

4. Democritus 400 BC
Democritus believed all things consisted of tiny indivisible units.
He called these tiny units he called atomos. The Greek word for “can not be cut” or “indivisible”
Ancient philosopher: Father of the Atom

5. John Dalton (1799)
 Developed what is considered to be the 1st Atomic Theory
 Was born into a modest Quaker family in England
 Began lecturing in public at the age of 12

6. Dalton’s Model (1799)
Dalton's model was that the atoms were tiny, indivisible, indestructible particles and that each one had a certain mass, size, and chemical behavior that was determined by what kind of element they were.

7. Dalton’s Model
Dalton’s model of the atom was similar to a tiny billiard ball.
Dalton’s model of the atom was solid and had no internal structure.

8. Dalton’s Atomic Theory
elements consisted of tiny particles called atoms.
1. all atoms of an element are identical
2. atoms of each element are different from one another; they have different masses.
3. compounds consisted of atoms of different elements combined together.
4. chemical reactions involved the rearrangement of combinations of those atoms.

9. Flaws in Dalton’s Model
Dalton’s falsely believed that the atom was the most fundamental particle.
We now know the atom is made up of even smaller particles we call the proton, neutron and electron.
Dalton’s theory could also not account for the formation of ions (charged particles)

10. Daltons Atomic Model Summary
Called: Billiard Ball Model
Could account for
Atoms of different element have different atomic masses
Elements consists of tiny particles called atoms
Could NOT account for
Atom being the smallest particle
Atoms having an internal structure
Atoms having charged particles

11. John J. Thomson (1897)
Discovered the electron using the Cathod Ray Tube (CRT)
Thomson found that the beam of charge in the CRT was attracted to the positive end of a magnet and repelled by the negative end.

12. Thomson’s Hypothesis
Concluded that the cathode beam was a stream of negative particles (electrons).
He tested several cathode materials and found that all of them produced the same result.
He also found that the charge to mass ratio was the same for all electrons regardless of the material used in the cathode or the gas in the tube.
Thomson concluded that electrons must be part of all atoms.

13. Thomson’s atomic model
Called the “plum-pudding” it was the most popular and most wildly accepted model of the time.

14. Thompsons atomic model could account for…..
the atom having an internal structure
Light given off by atoms
Atom with different atomic masses

15. Thompsons atomic model could NOT account for…..
Empty space (had atom filled with positive pudding)
Formation of ions

16. Gold Foil Experiment Conducted by students of Rutherford.
Proved that all atoms had a tiny, positively charged center.
Confirmed that atom’s were mostly empty space.

17. Rutherford ~ early 1900s
α-particle interaction with matter studied in gold foil experiment

19. Rutherford's Nuclear Model
1. The atom contains a tiny dense center
the volume is about 1/10 trillionth the volume of the atom
2. The nucleus is essentially the entire mass of the atom
3. The nucleus is positively charged
the amount of positive charge of the nucleus balances the negative charge of the electrons
4. The electrons move around in the empty space of the atom surrounding the nucleus

20. Rutherford’s atomic model (1911)
Could account for:
Empty space
Ions
Internal structure
Light given off when heated to high temperature.
Could not account for:
Stability

21. Planetary model
Planetary model used to explain electrons moving around the tiny, but dense nucleus
Nucleus contains
Protons- existence proposed in 1900s
Neutrons- existence proposed in 1930s

22. Bohr Questioned ‘planetary model’ of atom
Electrons located in specific levels from nucleus (discontinuous model)
Proposed electron cloud model based on evidence collected with H emission spectra

23. Bohr’s Atomic Model (1913) Bohr was a student of Rutherford.
Improved Rutherford’s model by proposing electrons are found only in specific fixed orbits.
These orbits have fixed levels of energy
This explained how electrons could give off light (gain or lose energy)

24. BOHR MODEL
Electrons are placed in energy levels surrounding the nucleus
8e-
8e-
Nucleus
(p+ & n0)
2e-

25. Bohr’s Atomic Model Could account for Flaws Internal structure
Atoms of different masses
Atom being mostly empty space
Light given off
Formation of positive ions
Flaws
Only really worked for Hydrogen

26. Chadwick (1932)
Discovered the neutron by bombarding Be with beta radiation.
Nuclear fission released a neutron.
chart

27. Review Describe each of the 6 different atomic models. Give the
Scientist Name
Name of model
What they could account for
What they could not account for (flaws)

28. Subatomic particle summary
Discovery by
Year
experiment
Proton
Rutherford
1911
Gold Foil Experiment
Electron
Thompson
1897
The response of cathode ray tube to a magnetic and electric fields
Neutrons
Chadwich
1932
Bombarded Be with beta radiation and a neutron was released

29. Subatomic Particles Relative mass Actual mass (g) Name Symbol Charge
Electron e- -1
1/1840
9.11 x 10-28
Proton p+ +1 1
1.67 x 10-24
Neutron n0 1
1.67 x 10-24

30. Subatomic Particles (cont.)
All atoms of an element have the same # of protons
protons identify an atom  atomic #
Atoms are electrically neutral
#p = #e-
Only neutrons and protons contribute to an atoms mass
#n + #p = atomic mass

31. Subatomic Particles equal in a neutral atom Atomic Number
NUCLEUS
ELECTRONS
equal in a neutral atom
PROTONS
NEUTRONS
NEGATIVE CHARGE
QUARKS
Atomic Number
equals the # of...
1st period
POSITIVE CHARGE
NEUTRAL CHARGE
Most of the atom’s mass.

32. ISOTOPES
= atoms with the same number of protons but DIFFERENT numbers of neutrons
Atomic Number
Mass Number
Element Symbol
Ex. Na-23 or Sodium-23
C-14 or Carbon-14
F-19 or

33. B. Isotopes Mass # Atomic # Nuclear symbol: Hyphen notation: carbon-12
Atoms of the same element with different mass numbers.
Nuclear symbol:
Mass #
Atomic #
Hyphen notation: carbon-12
C. Johannesson

34. B. Isotopes
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C. Johannesson

35. B. Isotopes 17 37 20 Chlorine-37 atomic #: mass #: # of protons:
# of electrons:
# of neutrons: 17 37 20
C. Johannesson

36. C. Relative Atomic Mass atomic mass unit (amu)
12C atom = × g
atomic mass unit (amu)
1 amu = 1/12 the mass of a 12C atom
1 p = amu 1 n = amu 1 e- = amu
© Addison-Wesley Publishing Company, Inc.
C. Johannesson

37. D. Average Atomic Mass Avg. Atomic Mass
weighted average of all isotopes
on the Periodic Table
round to 2 decimal places
Avg.
Atomic
Mass
C. Johannesson

38. D. Average Atomic Mass
EX: Calculate the avg. atomic mass of oxygen if its abundance in nature is 99.76% 16O, 0.04% 17O, and 0.20% 18O.
Avg.
Atomic
Mass
16.00
amu

39. D. Average Atomic Mass
EX: Find chlorine’s average atomic mass if approximately 8 of every 10 atoms are chlorine-35 and 2 are chlorine-37.
Avg.
Atomic
Mass
35.40 amu
C. Johannesson