1. Atomic Theory
Chapter 3
Sections 1 &2
9/18/14

2. WHAT’S A THEORY?

3. Atomic Theory

4. The Ancient Greeks
Democritus and other Ancient Greeks were the first to describe the atom around 400 B.C.
The atom was nature’s basic particle that makes up all matter.
The word atom comes from the Greek word “atomos” meaning indivisible.

5. Aristotle
The Greek philosopher Aristotle did not believe in atoms, and instead believed all matter was continuous.
Aristotle’s belief was accepted for ~2000 years, as there was no evidence to support either theory.

6. Antoine Lavoisier
First to state the law of conservation of mass, which states that mass is neither created nor destroyed during chemical or physical changes.

7. Joseph Louis Proust
Law of Definite Proportions = chemical compounds contain the same elements in the same proportions by mass regardless of the source or the amount.
i.e. Sodium Chloride (table salt) is always 39.34% Na and 60.66% Cl by mass.

8. “The BIG 3”
3 Major Theories on the ATOM
Dalton
Thompson
Rutherford

9. John Dalton’s Atomic Theory
All matter is composed of extremely small particles called atoms.
Atoms of an element are identical to each other and different from atoms of other elements.

10. 3) Atoms cannot be created or destroyed.
4) Atoms of different elements combine in simple whole-number ratios to form chemical compounds.
5) Atoms are combined, separated or rearranged when chemical reactions occur.

11. Impact of Dalton’s Atomic Theory
By relating atoms to mass, Dalton turns Democritus’ ideas into a scientific theory which could be tested.
Dalton’s model of the atom is a small, solid, and indivisible particle. (known as the “ball-bearing” model)

12. Flaws in Dalton’s Theory
Not all parts of Dalton’s Theory have proven to be correct. Today we know that…
Atoms are divisible into even smaller particles
Atoms of the same element can have different masses.

13. Cathode-Ray Tube Experiments
Electric current was passed through gases at low pressures. It was observed that the end of the tube near the cathode glowed and a ray traveled from there to the anode when current was passed through the tube.

14. In these experiments, it was also observed that the rays were deflected away from magnetic fields and negatively charged objects .

15. J.J. Thomson
Carries out experiments that support the hypothesis that cathode rays are made of negatively charged particles.
Measured the charge to mass ratio of cathode ray particles and determined that it was the same regardless of the metal or gas used.
Thomson concludes that all cathode rays are composed of identical negatively charged sub-atomic particles, which he calls electrons.

16. Thomson’s Model of the Atom
Thomson’s experiments show that (1) atoms are divisible and (2) they include negatively charged electrons.
Thomson’s Plum-Pudding
Model of the Atom has
negative particles spread
evenly throughout a solid,
positively charged sphere.

17. Robert Millikan
Conducted an Oil Drop Experiment that measured the charge of the electron.
Using this charge and Thomson’s charge to mass ratio, the mass of the electron is also discovered.
Mass of 1 electron = x10-31 kg

18. Ernest Rutherford Conducted the Gold-Foil Experiment
Positively charged alpha particles were fired at a thin piece of gold foil.
It was expected that the alpha particles would pass through with only a slight deflection.

19. However, it was discovered that about 1 out of every 8,000 particles were actually deflected back toward the source.
Rutherford concludes that there must be a very small and densely packed region of positive charge in the atom that would repel these alpha particles.
Rutherford is credited with discovering the nucleus.

20. The Nuclear Model of the Atom
Nucleus = very small and dense center of the atom, where protons and neutrons are found.
Protons have a positive charge.
Neutrons are electrically neutral.
Negatively charged electrons orbit the nucleus like the planets orbit the sun.
Most of the atom is empty space.

21. Mass of protons and neutrons = 1
*Mass of protons and neutrons = 1.67 x10-27 kg compared to the electron’s mass of x10-31 kg
Subatomic Particles
Proton
Neutron
Electron
Charge + —
Mass*
1 amu
~ 0
Location
Nucleus
Electron Cloud 21

22. Atomic Mass Units AMU = Atomic Mass Unit
Masses of atoms are very small
Carbon-12 atom is used as the standard atom for a mass reference.
1 amu = 1/12 of C-12 atom
Protons = amu ~= 1.0 amu
Neutrons = amu ~= 1.0 amu
Electrons = amu ~= 0.0 amu

23. Average Atomic Mass
Most elements occur in nature as a mixture of isotopes

24. Average Atomic Mass
Weighted average of the atomic masses of the naturally occurring isotopes
Think about your Warm Up Questions

25. Average Atomic Mass / Isotopes

26. Change in Charge (“Ions”)
Protons
+1 Charge for each Proton
Ex: Carbon12
6 protons (+6)
6 Neutrons (no charge)
Mass Number =
# protons + # Neutrons
Electrons
-1 Charge for each Electron
Ex: Carbon-12
6 electrons (-6)
C-12 is a neutral Atom
(+6) + (-6)
What if add an electron (-1)?