1. Chapter 2 Notes Atomic Structure

2. Atoms Democritus – Ancient Greek Science dude, 1 st proposed the idea of atoms, tiny indivisible particles Atomos – ancient Greek word for indivisible

3. John Dalton’s Atomic Theory 1. All elements are composed of submicroscopic particles called atoms 2. Atoms of the same element are identical. Atoms of one element are different from atoms of another element 3. Atoms of different elements can combine in simple whole number ratios to form compounds

4. Cont. 4. Chemical rxn’s occur when atoms are separated, joined or rearranged. However, atoms of one element are never changed into atoms of another element as the result of a chemical reaction. ATOM- the smallest particle of an element that retains the properties of that element

5. Electrons, Protons and Neutrons Electrons- are negatively charged particles –Discovered by J.J. Thomson in 1897 by bending a charged beam –Elektron (greek) shining beam –http://www.chem.uiuc.edu/clcwebsite/video/C ath.mov http://www.chem.uiuc.edu/clcwebsite/video/C ath.movhttp://www.chem.uiuc.edu/clcwebsite/video/C ath.mov

6. Protons Positively charged particle Discovered 1886 – E. Goldstein Discovered by Henry Mosley in 1911 Protos (greek) first

7. Neutron Subatomic particle with no charge Discovered by James Chadwick in 1932 SymbolCharge Mass (AMU) Electron e-e-e-e-1/1840 Proton p+p+p+p++11 Neutron n0n0n0n001

8. The Structure of the Nuclear Atom Nucleus – central core of the atom, composed of protons and neutrons Discovered by Earnest Rutherford in 1911 Gold foil Experiment Gold foil Experiment

10. Atomic Number The number of protons in the nucleus of an atom Neutral atoms – have equal # p + (+ charge) and e - (- charge)

11. Mass Number Number of p + and n 0 in the nucleus p + + n 0 = mass

12. Isotopes of Elements Isotopes: Atoms that have the same number of protons, but different numbers of neutrons Carbon – 12 Carbon – 13 Carbon - 14

13. Atomic Mass Atomic Mass Atomic Mass Unit (amu) – 1/12 the mass of a Carbon – 12 atom Atomic Mass – The weighted average mass of the isotopes in a naturally occurring sample

14. Calculating Atomic Mass Ex. Magnesium: Magnesium – 24 78.70% Magnesium – 25 10.13% Magnesium – 26 11.17% Multiply the mass number by the percent and add up the answers to get the average mass.