Chapter 4 Atomic Structure

1. 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. Compound made by atoms of elements A and B Atoms of element A Atoms of element B Mixture of atoms of elements A and B

6. What we know now 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 – negatively charged particle in all atoms,discovered in 1897 by J.J. Thomson

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

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

12. The Gold Foil Experiment The Gold Foil Experiment

13. The results… 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

14. “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

15. 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

16. Proton – positively charged particle in the nucleus of all atoms. Neutron – particle with no charge in the nucleus of all atoms except 1H

17. ASSIGN: Read 4.1 & 4.2 Lesson Check 4.1 & 4.2; #1-15 (page 104 & 109)

18. Chapter 4 Quick Quiz 1. Identify the idea or discovery these people are credited with. Thompson, Dalton, Democritus, Rutherford 2. Place the above ideas or discoveries in chronological order. 3. Describe the setup of the gold foil experiment. Also describe the expected and actual results of this experiment. What conclusion was made about the atom from the gold-foil experiment? 4. Name the three subatomic particles. Also give their symbol and charge. What particles are present in the nucleus of the atom? 5. What is an atom?

19. 4.3 – Distinguishing Among Atoms Atomic Number (Z) - The number of protons in an atom; identifies the element. Atoms have no net charge: # p+ = # e-

20. Mass Number (A) - The number of protons (p+) and neutrons (n0) in an atom. The mass number is NOT the atomic mass. 1 1 1 ??? 8 8 8 ??? 20 20 20 ??? The number of n0 depends on the mass number of the isotope

21. Isotopes - Atoms of the same element that have a different number of neutrons.

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

23. a. Be b. Ne c. Na 9 4 20 10 23 11 • Determining the Composition of an Atom • How many protons, electrons, and neutrons are in each atom? • p+ • e- • n0

24. Naturally Occurring Isotopes of Neon Percent Abundance in Nature 90.48% 0.27% 9.25%

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.

26. Atomic Mass – Weighted average of all the naturally occurring isotopes of the element. 12.000 amu 98.93 % 13.003 amu 1.07 % 14.003 amu 0.0000000001 % Atomic Mass of Carbon = 12.011 amu

27. 12.000 amu 98.93 % 13.003 amu 1.07 % 14.003 amu 0.0000000001 % Atomic Mass of Carbon = 12.011 amu No atom of carbon actually weighs 12.011 amu. But a typical carbon atom averages 12.011 amu. Atomic masses are weighted averages.

28. There are 2 stable isotopes of silver Silver-107; 106.905097 amu; 51.84% Silver-109; 108.904752 amu; 48.16% Calculate the atomic weight of silver. Atomic Weight of Silver = 107.868 amu

29. Weighted Averages are NOT just for Atomic Masses Your Chemistry grade is calculated as a weighted average. 50% Quizzes/Tests, 15% Homework, 15% Lab Reports, 10% Quarter Project, 10% Respect & Participation Calculate the un-weighted average grade Calculate the weighted average grade 60/90 50/50 70/75 45/50 60/60

30. Mass Spectrometer – separates isotopes by mass differences.

31. Mass Spectrum for Cadmium

32. Mass Spectrum for Zinc; Atomic Weight = 65.395

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

34. ASSIGN: Read 4.3 & answer #18-34 within 4.3 p. 114-119

35. Terms & Things to Know Democritus, Dalton, Thomson, Rutherford Atom Atomic Number (Z) Mass Number (A) Proton, Neutron, Electron Nucleus Isotope Atomic Mass (a.m.u.) Weighted Average Natural Abundance

36. How to determine the number of p+, n0, e- How to write isotope symbols How to calculate weighted averages Explain Gold Foil Experiment (setup, expected and actual results, conclusion) Different models of the atom (Democritus, Dalton, Thomson, Rutherford)

37. ASSIGN: Ch. 4 #35, 44, 47-57, 61, 64-71, 75 p. 122 -

39. Example… Calculate The average mass of zinc The weighted average mass of zinc