1. Introduction to the Atom and Atomic Models

2. 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

3. 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

4. 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.

5. 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.

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

7. 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)

8. Daltons Atomic Model Summary
Called: Billiard Ball Model
Could account for
Atoms of different atomic masses
Elements were tiny particles
Could NOT account for
Though atom was smallest particle
Did not have an internal structure
The formation of charged particles

9. 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.

10. 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.

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

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

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

14. 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.

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

17. 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

18. 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

19. Philipp Lenard (1903) Aluminum foil experiment
Lenard found that a beam of electrons was able to pass through a sheet of Al foil with almost no deflection.
Lenard correctly concluded that the majority of an atom’s volume is empty space.

20. Lenard’s atomic model Lenard’s model was composed of dynamids.
Lenard calculated the size of a dynamid based on his experimental results and found it to be 1 billionth (1/1,000,000,000) the size of the atom.

21. Lenard’s model could account for:
Different atomic masses (based on the number of “dynamids”).
The internal structure of an atom
The fact that most of the atom was empty space.

22. Flaws in Lenard’s model
Formation of Ions (gaining or losing charge)
The light given off by materials when heated to a high temperature.

23. Hantaro Nagaoka
First to present an atomic model close to the presently accepted model.
He came up with his model in 1903.

24. Nagoka’s Model Nagoka’s model of the atom was unstable.
According to the laws of planetary motion, the atom would collapse over time.

25. 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

26. Successes of Nagoka’s model
Atoms were able to give off electrons to form ions.
Accounted for the experimental fact that atom’s were mostly empty space
Explained different atomic weights.
Explained the light given off when heated to high temperatures.

27. 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

28. 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)

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

30. 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

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

32. 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)

33. Subatomic particle summary
Discovery by
Year
experiment
Proton
Rutherford
1911
Gold Foil Experiment
Electron
Thompson
1887
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

34. 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

35. 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

36. ISOTOPES
= atoms with the same number of protons but DIFFERENT numbers of neutrons
Mass Number
Atomic Number
Element Symbol
Ex. Na-23 or Sodium-23 or 23 Na or 23Na
C-14 or Carbon-14 or 14C or 14C
B-10 or Boron-10 or 10B or 10B 11 6 5

37. Isotope Practice Element has 6 p+, 8 n and 6 e- 6 p+, 6 n and 6 e-
Symbol .
Cu-65
H-2
C-14
C-12
K-40

38. Ions
Ion is an atom that has gained or lost one or more electrons and now has a charge
Metals always lose electrons to form POSITIVE ions
Non-metals always gain electrons to form NEGATIVE ions
The charge of the element is show on the right side of the symbol as a super script
Example: Na+1, Zn+1, S-2, N-3

39. Ions practice Element Ca+2 Br-1 Al+3 I-1
Number of electrons 20-2= = = =54

40. 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

41. 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

42. 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