Chapter 2

1. Chapter 2 Atoms and Elements

2. Visualizing Atoms • Binnig & Rohrer – development of the Scanning Tunneling Microscope (STM). • Produce images on the atomic level. • Iodine atoms on the surface of platinum metal

3. Modern Atomic Theory • Law of Conservation of Mass • Antoine Lavoisier – 1789 • In a chemical reaction, matter is neither created nor destroyed. 7.7 g Na + 11.9 g Cl2 19.6 g NaCl

4. Modern Atomic Theory • Law of Definite Proportions • Proust – 1797 • All samples of a given compound will have the same proportions of their constituent elements.

5. Modern Atomic Theory • The atomic theory of matter • Dalton, 1808 • Four postulates (main themes) • Each element is composed of tiny, indestructable particles called atoms. • All atoms of a given element are identical; The atoms of different elements are different and have different properties.

6. Modern Atomic Theory • Atoms combine in simple, whole number ratios to form compounds. • Atoms of one element cannot change into atoms of another element. In a chemical reaction, atoms only change the way they are bound together. • The evidence for the existence of atoms is overwhelming!

7. Discovery of Atomic Structure • Discovery of the electron. • J.J. Thompson ‑ cathode ray tube ‑ 1897 – determined that an electron was negatively charged.

8. Discovery of Atomic Structure • Robert Milliken ‑ oil drop ‑ 1909 – determined the charge value of an electron as well as its mass.

9. Oil Drop Experiment

10. Discovery of Atomic Structure • Radioactivity • Curie – 1900 • Rutherford – 1905 • Three main types of radioactive particles: Alpha, Beta, and Gamma. • Alpha particles • Essentially a helium nuclei. • Sources – many, but Curie’s used radium.

11. Nuclear Model of the Atom • Ernest Rutherford ‑ 1911. • Gold foil experiment. • Gold can be smashed into very thin sheets that are only a few atoms thick. • Alpha particles, from an alpha source, beamed at the gold foil. • Photographic paper placed as a detector in front and behind alpha source.

12. Gold Foil Experiment

13. Gold Foil Experiment

14. Nuclear Model of the Atom • Only 1 in 8000 alpha particles is scattered. • Scattering occurs when an alpha particle encounters a massive gold nuclei. • Rutherford proposed that: • Most of the atom’s mass and all of its positive charge were found in the small core of the atom called the nucleus. • Most of the volume of the atom is empty space.

15. Nuclear Model of the Atom • Rutherford’s model still had one problem. • H = 1 proton in nucleus. • He = 2 protons in nucleus. • He mass  4x mass of H mass. • Final piece of the puzzle is the neutron. • Neutrons have no charge and a mass of 1amu. • Discovered in 1932 by James Chadwick.

16. Modern View of Atomic Structure • Three subatomic particles exist in an atom. • Protons, neutrons, and electrons.

17. Modern View of Atomic Structure • Electrons have a negative charge of 1.602 x 10-19 C and a negligible mass. • Protons and neutrons reside in the nucleus which is extremely small. • Protons have a positive charge, equal in magnitude to an electron and a mass of about 1amu . • Neutrons have no charge and a mass of 1amu. • Over 99.9% of the mass of an atom resides in the nucleus.

18. Modern View of Atomic Structure • The atom is 100,000 times larger than the nucleus. If a golf ball represented the size of the nucleus, the atom would be about 3 miles in diameter. • Diameter of nucleus  10-15 m. • Diameter of atom  10-10 m. • The density of the nucleus is roughly 1013 to 1014 g/cm3.

19. Atomic and Mass Numbers • The Atomic Number (Z) is equal to the number of protons in the nucleus. • Each element has a unique atomic number and hence a unique number of protons. • The Mass Number (A) is the sum of the protons and neutrons found in the nucleus.

20. Isotopes & Symbols • Isotopes for an element occur when they have more than one mass number. • Isotopes of an element have the same number of protons, but a different number of neutrons. • An isotope can be designated by its mass number in the upper left corner – AX. • Or it can be designated after the symbol – X–A. • Example - 12C or C-12

21. Ions • Atoms quite often will gain or lose electrons when forming compounds. • When Lithium metal reacts, it loses one electron forming a +1 charge. Li+1. • When Fluorine gas reacts, it gains one electron forming a -1 charge. F-. • Positive ions are called cations. • Negative ions are called anions.

22. Periodic Law • Mendeleev (1869) – first to group the elements by similar properties. • First, he listed them in order of increasing atomic mass. • He then started a new row when elements had similar properties. • Thus, they are arranged according to horizontal rows which highlight the repeating properties of the elements in the vertical columns. • periods ‑ the elements in a horizontal row constitutes a period. • groups ‑ the elements in a vertical column constitutes a group.

23. The Periodic Table • Groups are numbered and labeled with A and B's. • Different conventions of numbering are used, however, we will use the traditional N.A. method with the A’s for the first two and last six groups.

24. The Periodic Table • The A groups are called main group elements and the B groups are called transition elements. • Some groups have special names: • Group 1A = Alkali metals • Group 2A = Alkaline Earth metals • Group 7A = Halogens • Group 8A = Noble gases or inert gases.

25. The Periodic Table

26. The Periodic Table • Metals ‑ a substance that has a characteristic of luster or shine, and is a good conductor of heat and electricity. • Metals are solids at room temperature and tend to lose electrons easily. • Metals are found to the left and below the diagonal line that runs through the right side of the main group elements. • The greatest majority of elements are metals.

27. The Periodic Table • Non‑metals ‑ either a gas at room temperature or a brittle solid, they are non‑conductors of heat and electricity. • Non‑metals tend to gain electrons easily. • Non-metals are found to the right and above the diagonal line. • Metalloids ‑ are elements that have the properties of both metals and non‑metals. • These fall along the diagonal line and include B, Si, Ge, As, Sb, and Te • Hydrogen – the one oddball on the periodic table.

28. Describing An Element • How would you describe the element: • C • Al • Sr • Fe

29. Ions • Metals form cations. Ex) Na  Na+ + 1e- • Non-metals form anions. Ex) Cl + 1e- Cl- • Predicting the charge of a species based on the periodic table.

30. Ions and Ionic Compounds • Main group elements usually form only one charge (valence) • Transition metals usually form many different charges. • Polyatomic ions are special groups of atoms that chemically combine to form a charged species.

31. Atomic Weights • The Atomic Mass Unit Scale is based on the 12C atom. • 1 amu = 1/12 the mass of the 12C atom. • Average atomic masses are based on the masses of each type of isotope a well as their abundance in nature. • Relationship of an amu to grams: • 1 amu = 1.66054 x 10-24 g • 1 g = 6.012214 x 1023 amu

32. Determining An Average Mass • A.W. = S(fract. abundance) x (isotopic mass) • Example • Chlorine has two isotopes, Cl-35 and Cl-37. Cl-35 has a mass of 34.969amu and an abundance of 75.78% Cl-37 has a mass of 36.966amu and an abundance of 24.22%. What is the atomic weight of Chlorine?

33. Determining An Average Mass A.W. = (34.969amu) x (0.7578) + (36.966amu) x (0.2422) A.W. = 35.45amu

34. Molar Mass • The Mole • Unit of quantity used in Chemistry. • Not convenient to count atoms. • 1 atom of C-12 = 12 amu (exact) • One mole of C-12 = 12 grams (exact) • Avogadro’s Number • Represents the number of C-12 atoms in 12 grams. • =6.02 x 1023 atoms.