1. Investigating Atoms and Atomic Theory
Students should be able to:
Describe the particle theory of matter. PS.2a
Use the Bohr model to differentiate among the three basic particles in the atom (proton, neutron, and electron) and their charges, relative masses, and locations. PS.3
Compare the Bohr atomic model to the electron cloud model with respect to their ability to represent accurately the structure of the atom.PS.3

2. Do Now:
On a clean sheet of paper, identify and list the following for your topic
Title or Topic idea:
Independent Variable: (list all three levels that you will be using)
Dependent Variable: (make sure it is quantitative)
Control group: (identify which of the 3 I.V. is your control group)
Who is going to benefit from your research?

3. Life of Pi
Solar Still
Read description pg 217

4. The History of Atomic Theory
Atomos: Not to Be Cut
The History of Atomic Theory

5. Atomic Models
A model uses familiar ideas to explain unfamiliar facts observed in nature.
A model can be changed as new information is collected.

6. The atomic model has changed throughout the centuries, starting in 400 BC, when it looked like a billiard ball →

7. Who are these men?
In this lesson, we’ll learn about the men whose quests for knowledge about the fundamental nature of the universe helped define our views.

8. Democritus 400 BC This is the Greek philosopher Democritus -
His theory: Matter could not be divided into smaller and smaller pieces forever, eventually the smallest possible piece would be obtained.
He asked: Could matter be divided into smaller and smaller pieces forever, or was there a limit to the number of times a piece of matter could be divided?

9. Atomos This piece would be indivisible.
He named the smallest piece of matter “atomos,” meaning “not to be cut.”

10. Atomos
To Democritus, atoms were small, hard particles that were all made of the same material but were different shapes and sizes.
Atoms were infinite in number, always moving and capable of joining together.

11. This theory was ignored and forgotten for more than 2000 years!

12. Why?
The eminent philosophers of the time, Aristotle and Plato, had a more respected, (and ultimately wrong) theory.
Aristotle and Plato favored the earth, fire, air and water approach to the nature of matter. Their ideas held sway because of their eminence as philosophers. The atomos idea was buried for approximately 2000 years.

14. Dalton’s Model
In the early 1800s, the English Chemist John Dalton performed a number of experiments that eventually led to the acceptance of the idea of atoms.

15. Dalton’s Theory
All elements are made of tiny, indivisible particles called atoms.
All atoms of the same element are identical. All atoms of different elements are different (unique).
Atoms combine with different atoms to form compounds. Atoms unite in whole number ratios (H2O, CO2).
Atoms are not created or destroyed, simply rearranged in chemical reactions.

16. .
This theory became one of the foundations of modern chemistry.

17. Thomson’s Plum Pudding Model
In 1897, the English scientist J.J. Thomson provided the first hint that an atom is made of even smaller particles.

18. Thomson Model
Thomson studied the passage of an electric current through a gas.
As the current passed through the gas, it gave off rays of negatively charged particles.

20. Thomson Model He proposed -“Plum Pudding” model.
Atoms were made from a positively charged substance with negatively charged electrons scattered about, like raisins in a pudding.

21. Thomson Model
Where did they come from?
This surprised Thomson, because the atoms of the gas were uncharged. Where had the negative charges come from?

22. Thomson concluded that the negative charges came from within the atom.
A particle smaller than an atom had to exist.
The atom was divisible!
Thomson called the negatively charged “corpuscles,” today known as electrons.
Since the gas was known to be neutral, having no charge, he reasoned that there must be positively charged particles in the atom.
But he could never find them.

23. Rutherford’s Gold Foil Experiment
In 1908, the English physicist Ernest Rutherford.
Rutherford’s experiment Involved firing a stream of tiny positively charged particles at a thin sheet of gold foil (2000 atoms thick)
was hard at work on an experiment that seemed to have little to do with unraveling the mysteries of the atomic structure

24. Gold Foil Experiment

25. Gold Foil Experiment
Most of the positively charged “bullets” passed right through the gold atoms in the sheet of gold foil without changing course at all.
Some of the positively charged “bullets,” however, did bounce away from the gold sheet as if they had hit something solid. He knew that positive charges repel positive charges.

27. This could only mean atoms were mostly open space.
Rutherford concluded that an atom had a small, dense, positively charged center that repelled his positively charged “bullets.”
He called the center of the atom the “nucleus”
The nucleus is tiny compared to the atom as a whole.

29. Rutherford
Rutherford reasoned that all of an atom’s positively charged particles were contained in the nucleus. The negatively charged particles were scattered outside the nucleus around the atom’s edge.

30. Bohr Model
In 1913, the Danish scientist Niels Bohr proposed that electrons were in a specific energy level.

31. Bohr Model
electrons move in definite orbits around the nucleus, much like planets circle the sun.
These orbits, or energy levels, are located at certain distances from the nucleus.

32. Wave Model

33. The Wave Model
Today’s atomic model is based on the principles of wave mechanics.
electrons do not move about an atom in a definite path, like the planets around the sun.
In fact, it is impossible to determine the exact location of an electron. The probable location of an electron is based on how much energy the electron has.

34. The Wave Model
According to the modern atomic model, at atom has a small positively charged nucleus surrounded by a large region (electron cloud) in which there are enough electrons to make an atom neutral.

35. Electron Cloud: A space in which electrons are likely to be found.
Electrons whirl about the nucleus billions of times in one second
They are not moving around in random patterns.
Location of electrons depends upon how much energy the electron has.

36. Electron Cloud:
Depending on their energy they are locked into a certain area in the cloud.
Electrons with the lowest energy are found in the energy level closest to the nucleus
Electrons with the highest energy are found in the outermost energy levels, farther from the nucleus.

37. Greek X Dalton Thomson Rutherford Bohr Wave Indivisible Electron
Nucleus
Orbit
Electron Cloud
Greek X
Dalton
Thomson
Rutherford
Bohr
Wave

38. The Atom
CHAPTER 4, PART I

40. Relative Sizes
What is the relative distance from the nucleus of an atom to the electrons? If the nucleus of an atom were the size of a marble how far away would the first electron be?

41. Answer

42. Subatomic Particles Proton Neutron Electron FYI: Quarks Up / Down
Charm / Strange
Top / Bottom

43. Proton Symbol = P+ Charge = Positive Mass = 1 atomic mass unit (amu)
Location = In the nucleus

44. Neutron Symbol = n0 Charge = Neutral Mass = 1 atomic mass unit (amu)
Location = In the nucleus

45. Electron Symbol = e- Charge = Negative Mass = 0 atomic mass unit (amu)
Location = Orbiting the nucleus

47. Atomic Calculations

48. X Atomic Number, (Z) = # of P+ in an atom’s nucleus
Mass Number, (A) = Sum of the # of N0 and # of P+ in a given nucleus.
  A Z
A = mass #, Z = atomic #, X = symbol X

49. Atomic Calculations 1) Atomic Number = Number of Protons
Z = Atomic Number = # p+
2) If neutral, then protons must equal electrons
3) Atomic Mass = Number of Protons + Neutrons
A = Atomic Mass
A = Z + n0

50. Atomic Calculations Game

51. Nucleons – particles that make up the atomic nucleus
Isotopes – atoms of the same element with the same # of P+, but different # of N0 and different mass numbers.
Ex.
Nuclide – a particular atom containing a definite number of protons and neutrons.
Ex. Carbon-12 Iron-56
Nucleons – particles that make up the atomic nucleus
Ex. protons and neutrons.

52. Atoms vs Ions
B-1
C+1 C

53. Ion – an atom that has lost or gained electrons and now has a charge
Number of protons = Z (atomic number)
Number of neutrons = A – Z (atomic mass – atomic number)
Number of electrons = Number of protons - charge

54. Average Atomic Mass
AMU – atomic mass unit, 1/12 the mass of a carbon-12 atom.
Mass on periodic table is a weighted average mass of all isotopes of an element.

55. Example Neon-12 relative abundance – 98%
Carbon-14, relative abundance – 2%
0.98 (12 amu) (14 amu) = amu

56. Example 2
What is the atomic mass of silicon if 92.21% of its atoms have mass of amu, 4.70% have mass of amu, and 3.09% have a mass of amu?
28.09 amu