Chapter 4

1. Chapter 4 The Structure of the Atom

2. Section 4.1—Early Theories of Matter • Science as we know it did not exist several thousand years ago. Curiosity sparked the investigations of most scholarly thinkers and they based their thoughts on life experiences. • The prefix “phil” means loving or dear & is seen in many science & non-science words. “Philosophy” refers to a love of knowledge.

3. Democritus

4. Section 4.1—Early Theories of Matter • Democritus was a Greek philosopher and the 1st person to propose the idea that matter was not infinitely divisible. He also thought that: • Matter is composed of empty space through which atoms move • Atoms are solid, homogeneous, indestructible, & indivisible • Different kinds of atoms have different sizes & shapes • Apparent changes in matter result from changes in the atoms themselves • The differing properties of matter are due to the size, shape, & movement of the atoms

5. Section 4.1—Early Theories of Matter • Democritus idea of the atomic “theory” was rejected because he could not answer the question of what held the atoms together. Also, the most influential Greek philosopher Aristotle rejected Democritus’ theory just because it did not agree with his own ideas of nature.

6. Section 4.1—Early Theories of Matter • In 1808, another atomic theory was published by John Dalton, a schoolteacher in England. Dalton’s theory marks the beginning of the development of modern atomic theory.

7. Section 4.1—Early Theories of Matter • Dalton revised Democritus’s ideas based upon the results of scientific research he conducted. Dalton thought that: • All matter is composed of extremely small particles called atoms • All atoms of a given element are identical, having the same size, mass, & chemical properties. Atoms of a specific element are different from those of any other element • Atoms cannot be created, divided into smaller particles, or destroyed • Different atoms combine in simple whole-number ratios to form compounds • In a chemical reaction, atoms are separated, combined, or rearranged

8. Section 4.1—Early Theories of Matter • Dalton’s theory easily explained the law of conservation of mass • Dalton’s theory has received 2 revisions over time • atoms are divisible into several subatomic particles • atoms of an element may have slightly different masses

9. Section 4.1—Early Theories of Matter • Atom—smallest particle of an element that retains the properties of the element • Scanning tunneling microscope allows us to see atoms and to actually move atoms to form shapes, patterns, & simple machines. This has lead to the discovery of nanotechnology—the atom-by-atom building of machines the size of molecules.

10. Scanning tunneling microscope

11. Section 4.1—Early Theories of Matter • Molecule—is a group of atoms that are bonded together & act as a unit

12. Section 4.1—Early Theories of Matter **Draw a double bubble map of Dalton’s & Democritus’s atomic theory.

13. Section 4.2—Subatomic Particles & the Nuclear Atom • In the 1800s, scientists were looking for a relationship between matter & electric charge. • Cathode ray tube—glass tube of which most of the air and matter had been sealed out of with electrodes connected to the ends of the tube and then to a battery discovered by William Crookes • negative terminal end—cathode • positive terminal end—anode

14. Section 4.2—Subatomic Particles & the Nuclear Atom • As a result of continued research, Crookes was convinced that • cathode rays were actually a stream of charged particles • the particles carried a negative charge • These negatively charged particles were found to be in all matter and were called electrons.

16. Section 4.2—Subatomic Particles & the Nuclear Atom • As a result of William Crookes cathode ray tube, one of the most important technological & social developments was discovered—the television.

17. Section 4.2—Subatomic Particles & the Nuclear Atom • J.J. Thomson, an English physicist, began to run cathode ray experiments to measure the ratio of the electron’s charge to its mass. After obtaining the ratio, he compared it to other known ratios, for instance the hydrogen atom. Thomson found that the mass of the electron was lighter than that of hydrogen, which disproved part of Dalton’s atomic theory (atoms can’t be divided into smaller particles).

18. Section 4.2—Subatomic Particles & the Nuclear Atom • In 1909, American physicist Robert Millikan determined the actual charge of an electron, -1. Knowing the mass to charge ratio and the charge of an electron, Millikan was able to determine the mass of a single electron—9.1x10-28 g or 1/1840 mass of a hydrogen atom.

19. Section 4.2—Subatomic Particles & the Nuclear Atom • It was known that matter is neutral because you don’t go around getting shocked when you touch any object. So it was proposed that electrons were evenly spaced within a uniformly distributed positive charged. This was known as the “plum pudding” model. AKA chocolate-chip cookie dough model.

20. Section 4.2—Subatomic Particles & the Nuclear Atom • In 1911, Ernest Rutherford conducted an experiment which led to the rejection of the “plum pudding” model. Using gold and an alpha particle emitting source, Rutherford calculated that an atom consisted mostly of empty space through which the electrons move. He also coined the term nucleus--the tiny, dense region in the center of the atom that contained most of the atom’s positive charge & virtually all of its mass. Rutherford’s new theory was called the nuclear atomic theory.

21. Gold foil experiment

22. Section 4.2—Subatomic Particles & the Nuclear Atom • In 1919, Rutherford fine-tuned his concept of the nucleus, stating that it contained positively charged subatomic particles equal to but opposite that of an electron (+1 charge)—proton. • In 1932, Rutherford & his co-worker, James Chadwick, concluded that the nucleus contained yet another subatomic particle, a neutral particle called the neutron, having a mass nearly equal to that of a proton.

23. Section 4.2—Subatomic Particles & the Nuclear Atom • The mass of the nucleus (protons & neutrons) is about 99.7% of the atom’s total mass. **Draw a double bubble map about Thomson’s “plum pudding” atomic model and Rutherford’s nuclear atomic model.

24. Section 4.3—How Atoms Differ • Not long after Rutherford’s gold foil experiment, Henry Moseley discovered that atoms of each element contain a unique positive charge in their nuclei. The number of protons in an atom is referred to as the element’s atomic number. • The periodic table is arranged left-to-right, top-to-bottom by increasing atomic number.

25. Section 4.3—How Atoms Differ • Remember that all atoms are neutral—SO, if Atomic number = number of protons = number of electrons • How many protons are in each type of atom? Gold (Au) _______ Silver (Ag) _________ Potassium (K) ________

27. example

28. Practice problems 1.How many protons & electrons are in each of the following atoms? • a.Boron c.Platinum • b.Radon d.Magnesium • 2.An atom of an element contains 66 electrons. What element is it? • 3.An atom of an element contains 14 protons. What element is it?

29. **worksheet • Determine the element, symbol, # of p, n, e, atomic #, or atomic mass where needed.

30. Section 4.3—How Atoms Differ • Isotopes have the same number of protons & electrons but different numbers of neutrons. • Most elements in nature are a mixture of isotopes. Isotopes with more neutrons have higher mass numbers—sum of the number of protons & the number of neutrons in the nucleus. To make it easier to identify each of the various isotopes of the elements, chemists add the mass number after the name. Chemists also use abbreviations to represent isotopes also. Fig 4-15, p 100

32. Section 4.3—How Atoms Differ • Mass number – atomic number = number of neutrons 107 109 47 Ar 47 Ar

33. element Atomic # Mass number a. neon 10 22 b. calcium 20 46 c. oxygen 8 17 d. iron 26 57 e. zinc 30 64 f. mercury 80 204 Example 4-2 & practice b-f • Determine the number of protons, electrons, & neutrons in the isotope of neon. Also, name the isotope given.

34. Section 4.3—How Atoms Differ • You can calculate the atomic mass of any element if you know its number of isotopes, their masses, & their % abundances.

35. Isotope Mass (amu) Percent abundance 6X 6.015 7.5% 7X 7.016 92.5% Example 4-3 • Calculate the atomic mass of unknown element X. Then identify the unknown element, which is used medically to treat some mental disorders.

36. Practice • Boron has 2 naturally occurring isotopes: Boron-10 (abundance= 19.8%, mass= 10.013 amu), boron-11 (abundance= 80.2%, mass= 11.009 amu). Calculate the atomic mass of boron • Helium has 2 naturally occurring isotopes, helium-3 & helium-4. The atomic mass of helium is 4.003 amu. Which isotope is more abundant in nature? • Calculate the atomic mass of magnesium. The 3 magnesium isotopes have atomic masses & relative abundances of 23.985 amu (78.99%), 24.986 amu (10.00%), & 25.982 amu (11.01%).

37. Section 4.4—Unstable Nuclei & Radioactive Decay • In a chemical reaction, only the electrons are involved in the reaction—NOT the particles of the nucleus. This is the reason that an atoms identity does not change during a chemical reaction. • However, there are some reactions that DO involve the changing of an atoms nucleus—nuclear reactions. In the late 1890s, scientists noticed that some substances spontaneously emitted radiation in a process called radioactivity. The actual rays & particles emitted were called radiation.

38. Section 4.4—Unstable Nuclei & Radioactive Decay • Radioactive atoms emit radiation because their nuclei are unstable. To gain stability, they emit radiation so that they can lose energy, trying to gain stability—a process called radioactive decay. • There are 3 types of radiation:

39. Section 4.4—Unstable Nuclei & Radioactive Decay • A nuclear equation shows the atomic number & mass number of the particles involved. It is important to note that in a nuclear equation, both mass number & atomic number are conserved. • The primary reason that atoms are not stable is the neutron-to-proton ratio. Eventually, radioactive atoms undergo enough radioactive decay to form stable, non-reactive atoms—this explains their rare existence in nature.