1. DOR Chemical/Physical Changes Classify each change as either chemical or physical. 1) Gasoline in your engine burns as you start the car. 2) Distilled water 3) Rust on a nail 4) Glow sticks 5) Medicine crushed into a powder

2. The Atom

3. Law of Conservation of Mass/Matter Matter cannot be created or destroyed Total mass is constant in chemical reactions. Originated with Antoine Lavoister (1700s) Quantitative mass data of reactants and products in mercury oxide decomposition.

4. Law of Definite Proportions Proposed by Joseph Proust (late 1700s) Decompositions and research with copper carbonate Compound composition and properties are fixed All compound samples have the same composition Same % of elements in the compound Ex. H 2 O

5. Law of Multiple Proportions 2+ compounds with same 2 elements, one element masses combined with second element’s mass in whole number ratio. Compositions of these compounds are related Proposed by John Dalton in addition to his atomic theory. Ex. CO 2 (2:1), CO (1:1)

6. Terminology Element– basic unit of a substance, contain only ONE type of atom, represented by symbol. Example: Ag, only contains Ag atoms. Atom—smallest particle of an element that still contains element properties. Example: One atom of Au, cannot have a smaller particle of gold and still be gold.

7. Compound vs. Molecule Compounds: more than one element elements combined in definite proportions Molecule: Smallest unit of a compound that still retains the properties of the compound.

8. How far back does the “atom” go? Democritus 400 B.C. Called the basic unit of matter an “atom”

9. The Atom and its Structure

10. Dalton Atomic Theory 1800s Atoms make up elements. Atoms form compounds as a whole and cannot be divided. Compounds formed from atoms joining in FIXED proportions

11. Dalton Atomic Theory (cont.) All matter made of atoms Atoms of an element have the same size, mass, etc. Different atoms have various sizes, mass, etc. Atoms cannot be divided, destroyed, or created. Atoms rearrange in chemical reactions.

12. John Thomson 1897 Cathode-Ray experiments. Discovered the electron particle and its possible charge. Stated electrons have a negative charge Determined ratio between mass and charge of an electron

13. Robert Millikan 1909, American Found the mass of an electron (VERY small) with Thompson’s data Currently, mass of electron = 9.109 x 10 -31 kg Discovered electron charge e = -1.602 x 10 -19 C Oil drop experiments.

15. Early Models of the Atom Thompson Must be a balance between negative and positive charges “Raisin-Pudding” model Uniform distribution of positive charge Positive cloud with stationary electrons

17. Early Models of the Atom Rutherford How are electrons distributed in an atom? Discovered alpha particles as 4 2 He Experiments with Au, Ag, and Pt foils bombarded with alpha particles

19. Early Models of the Atom Rutherford Mostly empty space Small, positive nucleus Contained protons Negative electrons scattered around the outside

20. James Chadwick SOOO we have protons and electrons…anything else? Experiments shooting alpha particles at Beryllium atoms Colleague of Rutherford Participated in Manhattan Project 1932 discovered neutrons contained in atom’s nucleus No charge Mass approximately same as proton mass

21. Early Models of the Atom Bohr 1913—hydrogen atom structure Physics + quantum theory Electrons move in definite orbits around the positively charged nucleus—planetary model Does not apply as atoms increase in electron number

22. Erwin Schrödinger Quantum mechanics 1926---wave equation Electrons behave more like waves than particles

23. Heisenberg’s Uncertainty Principle Electron’s location and direction cannot be known simultaneously Electron as cloud of negative charge

24. Modern Model of the Atom The electron cloud Sometimes called the wave model Electron as cloud of negative charge Spherical cloud of varying density Varying density shows where an electron is more or less likely to be

25. Homework Read pp. 36-39, 263-267, 276-280