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. law of definite proportions – a chemical 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

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

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.

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-

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

3.3: Counting Atoms – the basics 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.

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

3.3: Counting Atoms – involving the mole A. The Mole = 6.022 x 1023 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 1023 atoms.

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 1023 particles/mole as a conversion factor See sample problems B-E on pages 84-86