1. Laws to Know Conservation of Mass (matter) – Matter is neither created nor destroyed, only rearranged Definite Proportions – Compounds contain same elements with same proportions (H2O is always 2H:1O)

2. Dalton’s Atomic Theory Formed in early 1800’s All matter is composed of atoms Atoms cannot be made or destroyed All atoms of the same element are identical Different elements have different types of atoms Chemical reactions occur when atoms are rearranged Compounds are formed from atoms of the constituent elements

3. Atomic theory Smallest unit that maintains properties of a substance All matter is composed of atoms If atomic structure is changed, all chemical and physical properties will change Consist of 3 main subatomic particles – Proton, electron, neutron These consist of elementary particles (quarks, leptons, gauge bosons – list goes on an on: you do not need to know these)

4. Electron discovery Cathode Ray Tube (JJ Thomson 1897) – Found mass-charge ratio of electrons – Cathode ray consist of a stream of negative charges (electrons) – Flow from cathode (-) to the anode (+) – Degree of deflection led to the ratio

5. Example of Cathode Ray Tube

6. Mass of an electron Oil Drop Experiment (Millikan 1909) – Found the charge of an electron – Combined with Thomson’s work, one can find the mass of an electron – Atomizer used to spray fine mist of oil droplets – Used electric charge to suspend droplets

7. Nucleus Conclusions Gold Foil Experiment (Rutherford, Geiger, Marsden 1911) Fire alpha particles (He 2+ ) at thin gold foil Most pass through Some have small angle deflection Very few (1/8000) redirected back to the source

8. Gold Foil Cont. Conclude that most of the atom is empty space Dense positive center called the nucleus Most of the atoms mass found in the nucleus

9. Protons X-Ray Tube (Moseley 1914) Determined the charge of the nucleus (# of protons) Assigned atomic number to most atoms Beginning of periodic table being ordered off atomic number (was atomic mass before)

10. Isotopes and Neutrons Each atom has a unique number of protons – If number of protons are different, atom is different Atoms of the same element can have different masses – Difference caused by neutron Uncharged particle in the nucleus Roughly the same mass as a proton

11. Determining Mass Number All isotopes have different mass number Mass number is from Proton + Neutron – Number of Protons do not change – Neutrons Can change 12 C  6 N, 6P 14 C  8N, 6P

12. Mass of Protons, Neutrons, Electrons Proton = 1.67262158 x 10 -27 kilograms Neutron =1.67492729 x 10- 27 kilograms Electron = 9.10938188 × 10-31 kilograms Proton and neutron have about same mass – Near 1 atomic mass unit (amu) – Amu is based off of a Carbon-12 nucleus Carbon-12 = exactly 12 amu Electron mass is nearly negligible in comparison to protons and neutron

13. Average Atomic Mass Average taken of all natural isotopes Some isotopes are more common than other – Carbon-12, Oxygen-16, Hydrogen-1 etc – Average atomic mass will be near the most prevalent isotope

14. Ions Atoms that have gained or lost an electron Those that gain electrons – Negative charge – Nonmetals Those that lose electrons – Positive charge – Metals

15. Pure Substances All have fixed ratios Element: – Consist of only one type of atom Compound – Consist on 1 or more types of atoms Ionic: Metal and nonmetal Molecular: 2 or more nonmetals Metallic: 2 metals

16. Mixtures Do not have fixed ratios Homogeneous (solution) – Consist on only one phase – All parts have same physical/chemical properties Heterogeneous – May consist of multiple phases (not required) – Sections have different chemical/physical properties

17. Quantum Numbers Principal Quantum Number Angular Momentum Quantum Number Magnetic Quantum Number Spin Quantum Number

18. Principal Quantum Number (n) Main energy level Describes how far from the nucleus an electron is – Far away equals high energy Value range = 1-7 – Must be positive integer Total number of orbitals in a given energy level is = n 2

19. Angular Momentum ( l ) Describes orbital shape – s, p, d, f, (g) s = sphere P= dumbell D, f, g = complex shapes S=0, p=1, d=2, f=3

20. S and P Sublevel

21. D sublevel

22. Magnetic Quantum Number (m) Orientation of orbital around nucleus – ie. What orbital the electron is in Range from -3  +3 S, only one orbital  0 P, three orbitals  -1, 0 +1 D, five orbitals  -2,-1,0,+1,+2 F, seven orbitals  -3,-2,-1,0,+1,+2,+3

23. Spin Quantum Number (s) Indicates the 2 fundamental spin states of an electron Electrons in the same orbital must have opposite spin Values  +1/2 or -1/2

24. Principles and Rules for electrons Aufbau Principle – Electron will occupy lowest available energy level – They can move between energy levels Add energy, an electron goes up Remove energy, an electron falls to the ground state Falling releases energy in the form of light E=h (h=6.626 x10 -34 J*s = Plank’s constant C= (C=3.00x10 8 m/s)

25. Pauli Exclusion Principle No two electrons have the same exact set of quantum numbers

26. Hund’s Rule Orbitals are filled with one electron first Electrons in single filled orbitals must have the same spin Once all orbitals have 1 electron, then they can accept 2 Orbitals do not hold more than 2 electrons

27. Electron Stability 1. Full Octet (s and p) 2. Full sublevel 3. Half full sublevel 4. No order Atoms will shift electrons to achieve stability – Cu, Ag, Au, Cr, Mo, – Others that don’t follow rule exactly – Pt, Nb, Ru, Rh, Pd