1. Chapter 2.1-2.4 & definition of moles Atoms, Molecules, and Ions

2. Atomic Theory (Dalton) 1.Elements are made up of atoms (Na, K, H, etc.) 2.Atoms of same elements have same properties 3.Atoms are neither created nor destroyed in a chemical reaction (Law of Conservation of mass) 4.Compounds are formed by combinations of elements in fixed ratios. H 2 O  2 H, 1 O

3. What makes an atom? –How can one get that information? What makes an egg? –How can you get that information?

4. ATOM Electron (e) Proton (p) Neutron (n) Charge -1 +1 0 (au) Mass 1/1837 1 1 (au) 1 amu = 1.66 x 10 -24 g

5. Discovery & Characterization of the electron (e) Cathode Ray (-) charges electrons

6. Thompson’s Experiment With electrically charged plates and a magnet Charge/mass ratio for an electron –(e/m) = 1.76 x 10 8 Coulomb/g (or C/g)

7. Millikan oil-drop experiment Charge of e- = 1.60 x 10 -19 Coulomb (C) Mass of e- = 9.10 x 10 -28 g

8. Radioactivity Spontaneous emission of radiation α ( ) β ( )  electron γ ( γ ) Note: 238 is # protons + # neutrons, 92 is # protons α is attracted to (-) plate and hence must contain (+) charge

9. Rutherford’s Experiment Atom –Protons (p), Neutrons (n), and Electrons (e-) How are they together? Are they uniformly spread over?

10. Rutherford’s Experiment α-Particles are allowed to strike a gold plate »Most particles go through »A few go backward

11. Rutherford’s Experiment Can we predict what is going on in Rutherford’s Experiment? Why do α –particles (used as balls) go through undeflected?

12. Rutherford’s Experiment Conclusion –An atom is mostly made up of empty space (like an open door) with a massive nucleus (like a concrete wall), occupying a very small volume. –Nucleus has protons and neutrons, electrons revolve outside.

13. Isotopes, Atomic Number, and Mass Number Atomic number = Number of Protons Mass number = # Protons + # Neutrons Can you define “isotopes”? –Same elements? –How are they different?

14. Isotopes, Atomic Numbers, and Mass Numbers Here is an example of Isotopes. Remember that the top number is the mass # and the bottom number is the atomic #.

15. Average Atomic Mass Avg. mass = P 1 M 1 + P 2 M 2 + … –P 1 = % abundance for Isotope 1 –M 1 = Mass for Isotope 1 –P 2 = % abundance for Isotope 2 –M 2 = Mass for Isotope 2 Refer to Problem 29 (Pg. 72)

16. Mass Spectrometer Obtain atomic and molecular weights Need gaseous samples Ionization of sample by high energy electrons Pass gaseous ions through poles of a magnet For ions carrying the same charge, more massive particles will be less deflected than the less massive particles. Thus, mass separation takes place.

17. Mole & Number 1 thousand = 10 3 = 1000 1 million = 10 6 1 mole = 6.022 x 10 23 (Avogadro’s number) 1 mole H atoms = 6.022 x 10 23 H atoms 1 mole H 2 O molecules = 6.022 x 10 23 H 2 O molecules

18. Mole & Gram (Molar Mass) (Remember: M-g-P = mole-gram-Periodic Table) 1 mole C weighs 12.0107 g ≈ 12.0 g C atoms –Found on the Periodic Table 1 mole H weighs 1.00794 g ≈ 1.00 g H atoms 1 mole H 2 O = 2(1.00794) + 15.9994 ≈ 2 + 16 = 18 g H 2 O molecules