1. Matter, and more!

2. Matter Anything that has MASS and takes up SPACE

3. How is Matter classified? 1) Pure Substances 2) Mixtures

4. 1) Pure Substances Composition remains the same, does not depend on a sample = Fixed composition Homogenous—same throughout. Example: Compound (NaCl) or Element (Fe)

5. 2) Mixtures 2+ substances combined together Substances do not change their properties or name. Able to be separated, not chemically combined. Possess a combination of properties based on the substances present.

6. Types of Mixtures  Homogenous Uniform composition Also known as “true solutions” Ex. Salt-water 2) Heterogeneous No uniform composition Can easily see the different components of the mixture Ex. Italian dressing

7. Mixture Types---More Detail True Solution What we normally think of as a “solution” Homogenous Solute/solvent completely dissolved

8. How are mixtures separated? Thin-Layer Chromatography Filtration Centrifuge

9. Identify whether a substance is pure/mixture and homogenous/heterogeneous. 1) Salad 2) Kool-Aid 3) Vegetable soup 4) Ca 5) Water

10. Properties of Matter  Chemical Ability to go through changes resulting in a different substance The substance is no longer the same, different identity Evidence of chemical reaction: color change, precipitate forms, gas formation, and/or temperature change Ex. Burning  Physical Observed or measured property Substance identity is not changed Ex. melting point, boiling point, density

11. Classify each change as either chemical or physical.  Gasoline in your engine burns as you start the car.  Distilled water  Rust on a nail  Glow sticks  Medicine crushed into a powder

12. The Atom

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

14. 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.

15. 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

16. 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.

17. 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.

18. 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

19. 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.

20. 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

21. 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

23. 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

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

26. Atomic Structure Nucleus Protons Neutrons Electrons

27. Atomic Structure Electrons Tiny, very light particles Have a negative electrical charge (-) Move around the outside of the nucleus

28. Atomic Structure Protons Much larger and heavier than electrons Protons have a positive charge (+) Located in the nucleus of the atom

29. Atomic Structure Neutrons Large and heavy like protons Neutrons have no electrical charge Located in the nucleus of the atom

30. Describing Atoms Atomic Number (Z) = number of protons In a neutral atom, the # of protons = the # of electrons Mass Number (A)= the number of protons + the number of neutrons

31. Isotopes The number of protons for a given atom never changes. The number of neutrons can change. Two atoms with different numbers of neutrons are called isotopes have the same atomic # have different atomic Mass # ’ s Behave the same chemically

32. Isotopes

33. Atomic Mass Weighted average of element’s natural isotopes Some isotopes are more abundant than others…. SO atomic mass leans towards more abundant mass

34. How do we calculate atomic mass? 1) Masses of Isotopes 2) Fraction of the abundance of each isotope usually a percentage Average atomic mass = mass contributed by all isotopes Fraction of abundance (isotope mass) = mass from a particular isotope

35. Example 1: Neon has 3 natural isotopes. Ne-20 (90.51%, 19.99244 u) Ne-21 (0.27%, 20.99395 u) Ne-22 (9.22%, 21.99138 u) What is the weighted average atomic mass for Ne?

36. Example 2: Two natural copper isotopes are Cu-63 (62.9298 u) and Cu-65 (64.9278 u) If copper’s atomic mass is given as 63.546 u, what are the percent abundances of these isotopes? Which isotope is the most abundant?

37. Homework pp. 66-67 #7, 31-33, 37, 39-42 Finish “Atomic Theory I” worksheet