1. Anything in black letters = write it in your notes (‘knowts’)
Chapter 4
Atomic Structure
Anything in black letters = write it in your notes (‘knowts’)

2. 4.1 – Defining the Atom Atom -
smallest particle of an element that still has the properties of the element
comes from the Greek word atomos which means uncuttable or indivisible

3. Democritus (460 B.C. – 370 B.C.)
one of the first to propose the idea of the atom; based on pure speculation

4. John Dalton (~1800)
proposed 1st atomic theory

5. Daltons Atomic Theory (~1800) p. 103)
All elements are composed of tiny indivisible particles called atoms.
Atoms of the same element are identical. The atoms of any one element are different from those of any other element.
Atoms of different elements can physically mix together or can chemically combine in simple whole-number ratios to form compounds.
Chemical reactions occur when atoms are separated from each other, joined, or rearranged in different combinations. Atoms of one element are never changed into atoms of another element as a result of a chemical reaction.
Atoms of element A
Atoms of element B
Compound made by atoms of elements A and B
Mixture of atoms of elements A and B

6. Caveats of Dalton’s Atomic Theory
All elements are composed of tiny indivisible particles called atoms.
Atoms are not indivisible – they are made of subatomic particles
Atoms of the same element are identical. The atoms of any one element are different from those of any other element.
Every atom has at least one isotope; one atom’s isotope is NOT identical to another isotope of the same atom.
Atoms of different elements can physically mix together or can chemically combine in simple whole-number ratios to form compounds.
This is known as the Law of Definite Proportions – very important.
Chemical reactions occur when atoms are separated from each other, joined, or rearranged in different combinations. Atoms of one element are never changed into atoms of another element as a result of a chemical reaction.
Atoms of one element can change into an atom of another element as a result of a nuclear reaction.

7. 4.2 – Structure of the Nuclear Atom
Subatomic Particles
Electron – discovered in 1897 by J.J. Thomson while experimenting with cathode ray tubes

8. Thomson performed experiments that involved passing electric current through gases at low pressure.
The result was a glowing beam, or cathode ray, that traveled from the cathode to the anode.
Thomson found that a cathode ray is deflected by electrically charged metal plates.
Thompson knew that opposite charges attract and like charges repel, so he hypothesized that a cathode ray is a stream of tiny negatively charged particles moving at high speed; now called electrons.
To test his hypothesis, Thompson set up an experiment to measure the ratio of an electron’s charge to its mass.
Also, the charge-to-mass ratio of electrons did not depend on the kind of gas in the cathode-ray tube or the type of metal used for the electrodes.

9. A cathode ray can also be deflected by a magnet.

10. Properties of Subatomic Particles Relative mass (mass of proton = 1)
Proton – observed in cathode ray tubes also in 1886 by Eugen Goldstein.
Neutron – confirmed in 1932 by James Chadwick.
Properties of Subatomic Particles
Particle
Symbol
Relative charge
Relative mass (mass of proton = 1)
Actual mass (g)
Electron e– 1–
1/1840
9.11  10–28
Proton p+ 1+ 1
1.67  10–24
Neutron n0

11. Thomson’s Plum Pudding Model
The Atomic Nucleus
How are atoms structured?
Dalton
Democritus
Thomson’s Plum Pudding Model

12. In 1911, Ernest Rutherford and others performed the Gold Foil Experiment to test the plum pudding model
Ernest Rutherford

13. The Gold Foil Experiment

14. The results… Expected Actual
It was expected that alpha particles would pass through the plum pudding model of the gold atom undisturbed.
It was observed that a small portion of the alpha particles were deflected, indicating a small, concentrated positive charge (the nucleus!)
Expected
Actual

15. It was quite the most incredible event that has ever happened to me in my life. It was almost as incredible as if you fired a 15-inch shell at a piece of tissue paper and it came back and hit you. On consideration, I realized that this scattering backward must be the result of a single collision, and when I made calculations I saw that it was impossible to get anything of that order of magnitude unless you took a system in which the greater part of the mass of the atom was concentrated in a minute nucleus. It was then that I had the idea of an atom with a minute massive center, carrying a charge.
—Ernest Rutherford

16. Nucleus – tiny positively charge core of an atom
Rutherford’s Nuclear Model of the Atom
Is this the current model of the atom?
NO…
If an atom were the size of a football stadium, the nucleus would be about the size of a marble

17. ASSIGN:
Read 4.2
Lesson Check 4.2; #9-15 (page 109)

18. 4.3 – Distinguishing Among Atoms
Atomic Number (Z) -
The number of protons in an atom; identifies the element.
In a neutrally charged atom, the number of protons (p+) equals the number of electrons (e-)

19. The number of protons (p+) and neutrons (n0) in an atom.
Mass Number (A) -
The number of protons (p+) and neutrons (n0) in an atom.
The mass number is NOT the atomic mass.
Element
Atomic Number (Z)
Protons (p+)
Electrons (e-)
Neutrons (n0) H O Ca 1 1 1
??? 8 8 8
??? 20 20 20
???
The number of n0 depends on the mass number of the isotope

20. Isotopes -
Atoms of an element that have a different number of neutrons.
Element
Atomic Number (Z)
Protons (p+)
Electrons (e-)
Neutrons (n0) H O Ca 1 1 1 8 8 8 20 20 20

21. Zoom for detail

22. Chemical Symbols for Isotopes
A is the superscript
Z is the subscript

23. Determining the Composition of an Atom
How many protons, electrons, and neutrons are in each atom?
a. Be b. Ne c. Na 9 4 20 10 23 11

24. Percent Abundance in Nature
Naturally Occurring Isotopes of Neon
Percent Abundance in Nature
90.48% % %

25. The masses of atoms are rarely expressed in grams.
The C-12 isotope has been given a mass of exactly 12 atomic mass units (amu)
The masses of all other elements are based on the mass of the C-12 isotope.
Why is the mass of a carbon atom amu?

26. Atomic Mass of Carbon = 12.011 amu
Weighted average of all the naturally occurring isotopes of the element.
amu
98.93 %
amu
1.07 %
amu %
Atomic Mass of Carbon = amu

27. Atomic masses are weighted averages.
amu
98.93 %
amu
1.07 %
amu %
Atomic Mass of Carbon = amu
No atom of carbon actually weighs amu. But a typical carbon atom averages amu.
Atomic masses are weighted averages.

28. There are 2 stable isotopes of silver
Silver-107; amu; 51.84%
Silver-109; amu; 48.16%
Calculate the atomic weight of silver.
Atomic Weight of Silver = amu

29. Despite differences in the number of neutrons, isotopes of an element are chemically similar.
Neutrons do not determine chemical reactivity; the electrons do.

30. ASSIGN:
Lesson Check 4.3 (#26-34); p. 119

31. ASSIGN:
#53-59, 61, 64-67