1. Chapter 3 – Atoms: The Building Blocks of Matter 3.1: Atomic Theory History A. 1700s: quantitative studies of chemical reactions led to several laws: 1. law of conservation of mass – mass is neither created nor destroyed in a chemical reaction

2. 2. law of definite proportions – a chemical compound always contains the same elements in the same proportion by mass ex. NaCl is always 39.3% sodium and 60.7% chlorine by mass. 3. law of multiple proportions – if the same 2 elements are found in different compounds, then the ratio of the masses of the second element (with the first element’s mass being the same) is always a whole number

3. B. 1800s: Dalton’s Atomic Theory –He explained the above 3 laws in his theory: 1. All matter is composed of atoms 2. Atoms of a given element are identical. Atoms of different elements are different. 3. Atoms cannot be subdivided, created, or destroyed 4. Atoms of different elements combine in whole-number ratios 5. Atoms are combined, separated, or rearranged in a chemical reaction

4. C. The Modern Atomic Theory – A couple of Dalton’s points have been modified: Atoms are divisible Atoms of the same element can have different masses.

5. 3.2: The Structure of the Atom A. The Electron – discovered in 1897 by J.J. Thomson after experiments with a cathode-ray tube. Properties:  negatively charged  1/1837 the mass of a proton  symbol = e -

6. B. The Nucleus – discovered by Rutherford after doing his Gold Foil Experiment. Composition of the nucleus: 1. protons = positively charged subatomic particles 2.neutrons = subatomic particles with no charge Both protons and neutrons have a mass of 1amu (atomic mass unit) Rutherford's experiment

7. 3.3: Counting Atoms A. Atomic Number = the # of protons Each element has its own atomic # In a neutral atom, the # of protons = the # of e - B. Mass Number = the # of protons + the # of neutrons in an atom (e - do not contribute significantly to an atom’s mass because their mass is too small) C. Isotopes = atoms of the same element (same # of protons) that have different numbers of neutrons. 2 ways of writing: 1. element name – mass # (ex. Hydrogen - 3) 2.

8. D.Average Atomic Mass Most elements exist in nature as a mixture of isotopes. Average atomic mass is the weighted average mass of all the isotopes of an element. In calculating atomic mass, we must consider the abundance of each isotope. – Steps in calculating atomic mass: 1. Multiply the mass of each isotope by the relative abundance of that isotope 2. Total the answers from step #1

9. Chapter 21 – Nuclear Chemistry 21.2: Radioactive Decay A. Terms : radioactive decay = the process by which an unstable nucleus loses energy by emitting penetrating rays called radiation radioisotope (or radioactive nuclide) = an isotope of an element that has an unstable nucleus

10. B. Types of Radiation: 1. alpha (α): 2 protons + 2 neutrons (same thing as a helium nucleus) -not very penetrating (can be stopped by paper or skin) Results in the atomic number going down by 2 and the mass going down by 4

11. 2. beta (β): electrons formed from the decomposition of a neutron. The proton formed stays in the nucleus (atomic number inc. by 1) -more penetrating than alpha (can be stopped by wood or foil)

12. 3. gamma (γ): high-energy electromagnetic radiation. Often emitted along with alpha or beta. -has no mass or charge -very penetrating (need concrete or thick lead to stop

13. 21.1: Nuclear Stability The stability of a nucleus depends on its neutron to proton ratio. For atoms with an atomic # < 20, a 1:1 ratio is stable. For atoms with an atomic # > 20, more neutrons are needed for stability, and a ratio of 1.5:1 is most stable for heavy elements Note: all nuclei with atomic # >83 are radioactive (not stable)

14. Too many neutrons leads to beta decay, which results in a decrease in neutrons and an increase in protons (atomic # increases by 1) A nucleus with too many protons and neutrons will emit alpha particles (atomic # decreases by 2)

15. 21.2: Half-Life = the time required for half of a radioactive sample to decay – ex. half life of uranium-238 is 4.5 x 10 9 years – Half-life problems: see sample problem B on page 689

16. 3.3: Counting Atoms A. The Mole = 6.022 x 10 23 particles of something (also called Avogadro’s number) B. Molar mass = the mass of one mole of a substance. The units are g/mole. It’s numerically equal to the atomic mass in atomic mass units. Sometimes it’s called the formula mass. For ex., the molar mass of sodium is 23.0g/mole. Thus, 23g of sodium = 1 mole = 6.022 x 10 23 atoms.

17. C. Conversions : between mass and moles of an element: use the molar mass as a conversion factor between particles and moles of an element: use 6.022 x 10 23 particles/mole as a conversion factor See sample problems B-E on pages 84-86