1. The Building Blocks of Matter
Chapter 3: Atoms
The Building Blocks of Matter
An atom is the smallest particle of an element that retains the chemical properties of that element.

2. The Atom: From Philosophical Idea to Scientific Theory
Section 1
The Atom: From Philosophical Idea to Scientific Theory
Page 64

3. The Early Atom
As early as 400 B.C., Democritus called nature’s basic particle the “atomon” based on the Greek word meaning “indivisible”.
Aristotle succeeded Democritus and did not believe in atoms. Instead, he thought that all matter was continuous. It was his theory that was accepted for the next years. (Read page 43 of your textbook.)

4. Fast Forward to the early1700s: What scientists knew
Definition of the word “element” was widely accepted
Elements combine to form compounds that have different properties from those elements
Controversy
Did elements always combine in the same ratio when forming a particular compound?

5. How to answer the controversy:
Late 1700’s, the study of matter was revolutionized by focusing on quantitative analysis of reactions
These studies were made possible by developing and using newer, more precise balances

6. Three Basic Laws of Matter:
Law of Conservation of Mass
Law of Definite Proportions
Law of Multiple Proportions

7. Basic Laws of Matter
Law of Conservation of Mass- mass is neither created nor destroyed during ordinary chemical reactions or physical changes.
CH4 + 2O2 → 2H2O + CO2
16g + 64g → 36g g
Antoine Lavoisier
stated this about 1785

8. Basic Laws of Matter
Law of Definite Proportions – no matter how much salt you have, it is always 39.34% Na and 60.66% Cl by mass.
Example: Sodium chloride always contains % Na and 60.66% Cl by mass.
2NaCl → 2Na Cl2
100g → 39.34g g
116.88g → ? ?
Joseph Louis Proust
stated this in 1794.

9. John Dalton proposed this in 1803.
Basic Laws of Matter
Law of Multiple Proportions- Two or more elements can combine to form different compounds in whole-number ratios.
Example
John Dalton proposed this in 1803.

10. Dalton’s Atomic Theory
In 1808, Dalton proposed a theory to summarize and explain the laws of conservation of mass, definite proportions, & multiple proportions.
I was a school teacher at the age of 12!

11. Dalton’s Atomic Theory
John Dalton
All matter is composed of extremely small particles called atoms.
2. Atoms of a given element are identical in size, mass, and other properties.**
3. Atoms cannot be subdivided, created, or destroyed.**
4. Atoms of different elements combine in simple whole-number ratios to form chemical compounds.
5. In chemical reactions, atoms are combined, separated, or rearranged.
**Today, we know these parts to have flaws.

12. Flaws of Dalton’s Theory…
2. Atoms of a given element are identical in size, mass, and other properties.
3. Atoms cannot be subdivided, created, or destroyed.
Isotopes – atoms with the same number of protons but a different number of neutrons
Subatomic particles – electrons, protons, neutrons, and more

13. The Structure of the Atom
Section 2
The Structure of the Atom
Page 70

14. The Atom
Atom - the smallest particle of an element that retains the chemical properties of that element.
CARBON

15. Discovery of the Subatomic Particles
The discovery of the subatomic particles came about from the study of electricity & matter.
Benjamin Franklin’s kite experiment in demonstrated that lightning was electrical.

16. Charged Particles
In 1832, Michael Faraday proposed that objects are made of positive and negative charges.

17. Discovery of the Electron
In the late 1870’s many experiments were performed in which electric current was passed through gases at low pressures due to the fact that gases at atmospheric pressure don’t conduct electricity well.
These experiments were carried out in glass tubes called cathode-ray tubes or Crookes tubes.
Sir William Crookes developed these tubes.

18. Discovery of the Electron
Cathode-Ray Tube or Crookes Tube

19. Discovery of the Electron
When current was passed through the cathode ray tube, the surface of the tube, directly opposite the cathode, glowed.
It was thought that this glow was caused by a stream of particles called cathode rays.
The rays traveled from cathode (negative) to anode (positive).

20. Discovery of the Electron
Negatively charged objects deflected the rays away.
Therefore, it was determined that the particles making up the cathode rays were negatively charged.

21. Joseph John Thomson
In 1897 the English physicist Joseph John Thomson was able to measure the ratio of charge of the cathode ray particles to their mass.
He found that the ratio was always the same regardless of the metal used to make the cathode or the nature of the gas inside the cathode ray tube.
Thomson concluded that cathode rays were composed of identical, negatively charged particles called electrons.

22. Joseph John Thomson
Thomson’s experiments revealed that the electron has a very large charge-to-mass ratio.
Thomson determined that electrons were present in all elements because he noted that cathode rays had identical properties regardless of the element used to produce them.

23. Cathode Ray Tube Experiment Accomplishments
Proved that the atom was divisible and that all atoms contain electrons.
This contradicted Dalton’s Atomic Theory.
This allowed a new model of the atom.

24. Plum-Pudding Model of the Atom

25. Discovery of X-Rays
In 1895 William Conrad Roentgen discovered X- rays, a form of radiation.

26. Radioactivity
In 1896, the French scientist Henri Becquerel was studying a Uranium mineral. He discovered it was spontaneously emitting high-energy radiation.
In 1898, Marie and Pierre Curie attempted to isolate radioactive components of the mineral.

27. Radioactivity
In 1899, Ernest Rutherford, a British scientist, began to classify radiation: alpha (a), beta (b), and gamma (g).

28. Radiation
Look closely at the paths of radiation. Do you notice something about the amount of deflection of each type of particles?

29. Radiation

30. Discovery of the Nucleus
In 1911, Ernest Rutherford performed a Gold Foil Experiment.
He and his colleagues bombarded a thin piece of gold foil with fast moving, positively charged alpha particles.

31. Alpha Particles Alpha (a) particles are Helium-4 nuclei.
This means they are two protons and two neutrons (with no electrons).
Thus, they are positive.

32. Gold Foil Experiment

33. Gold Foil Experiment
As expected, most of the alpha particles passed straight through with little or no deflection.
However, 1/8000 of the positively charged alpha particles were deflected, some back at the source.

34. (Po)
This slide is animated! Check out the website!

35. Gold Foil Experiment

36. This slide is animated! Check out the website!

37. Gold Foil Experiment
From this experiment, Rutherford discovered that there must be a very densely packed positively charged bundle of matter within the atom which caused the deflections.
He called this positive bundle the nucleus.
He tried this experiment with other metals and found the same results.

38. Gold Foil Experiment
The volume of the nucleus was very small compared to the volume of the atom.
Therefore, most of the atom was composed of empty space. Niels Bohr later found that this empty space was where the electrons were located.

39. Checking for Understanding Gold Foil Experiment
Why did some of the alpha particles come straight back to the source or deflect away from the nucleus?
Why did he conclude that the nucleus must be positive?
What three things did Rutherford conclude from the gold foil experiment?

40. Checking for Understanding Gold Foil Experiment
If gold atoms were solid spheres stacked together with no space between them, what would you expect would happen to particles shot at them?
What year did Ernest Rutherford perform this experiment?
Rutherford experimented with many kinds of metal foil as the target. The results were always similar. Why was it important to do this?

41. “It was about as believable as if you had fired a 15-inch shell at a piece of tissue paper, and it came back and hit you.”
-Ernest Rutherford

42. Bohr’s Model of The Atom - 1913
Neils Bohr discovered that electrons orbit around the nucleus at different distances.
These orbits have different amounts of energy based upon how far from the nucleus they are located.
More on this in Chapter 4!

43. Discovery of the Proton - 1918
Henry Mosley found that he could use alpha particles as “bullets” to knock positively charged particles from the nucleus of the atom.
His conclusion was that the nucleus must be a collection of protons.
Electrons in orbit
Nucleus made of protons

44. The Big Question:
Similar charges repel, so how can the nucleus be made of all positively charged particles (protons)?

45. The Answer: The discovery of the neutron - 1932
James Chadwick discovered that the nucleus also includes neutral particles that he called, neutrons.
Neutrons act like a “glue” to help hold the nucleus together.
The “electron cloud” – see next slide
Nucleus made of protons & neutrons

46. The electron cloud
In 1927, Werner Heisenberg proposed that electrons do not orbit the nucleus like planets orbit the sun.
He proposed that electrons form an “electron cloud” around the nucleus
The “electron cloud” – see next slide
More on this in Chapter 4!
Nucleus made of protons & neutrons

47. The electron cloud The Heisenberg Uncertainty Principle:
In 1927, Werner Heisenberg proposed that there is no way to know both the location and the velocity of an electron at any given time.
The electron cloud is a scatter plot of where you are most likely to find an electron at any given time.

48. The Simple Atomic Model
This simple atomic model combines the big ideas that all of the scientists contributed to the atom.
This is the model we will refer to because it is the simplest to understand 