1. Atomic Structure www.chemistryhelpers.com

2. Subatomic Particles An atom is the smallest unit of an element. It consists of three major particles: Note: amu = atomic mass unit

3. ParticleMassLocationCharge Proton 1 amuNucleus+1 Neutron 1 amuNucleus0 Electron 1/1836 amu Orbitals

4. Atomic Theory explains the structure of matter in terms of different combinations of very small particles called atoms.

5. Dalton’s Theory All elements are composed of indivisible atoms. All atoms of a given element are identical.

6. Atoms of different elements are different; that is, they have different masses. Compounds are formed by the combination of atoms of different elements.

7. JJ Thomson used a cathode ray tube to show smaller units that make up an atom. The ray was deflected a certain way by a magnetic field, so he concluded that the ray was formed by particles and that the particles were negatively charged. The only source available for the particles was the atoms present.

9. Therefore, Thomson theorized that an atom contains small, negatively charged particles called electrons. This theory is referred to as the Plum Pudding Model. In this model, the mass of the rest of the atom was evenly distributed and positively charged, taking up all of the space not occupied by the electrons.

11. Rutherford’s Gold Foil Experiment Alpha particles directed at a piece of gold foil. Some alpha particles went through, some were deflected, some were returned.

13. Conclusion: Majority of the volume of an atom is empty space. Atoms have a dense positively charged central core.

14. Do Now How many p +, n 0, e - are in: 35 17 Cl Draw the model (where are the p +, n 0, e - ) *look at periodic table

15. Quick Review Who are the 3 theorists we went over yesterday? What are the “nicknames” of their theories and why?

16. The Bohr Model 1 or K-shell = max 2 e- 2 or L-shell = max 8 e- 3 or M-shell = max 18 e- 4 or N-shell = max 32 e- Niels Bohr Also called the Planetary Model Use Periodic Table to determine how electrons are arranged.

17. Basic formula for determining the amount of electrons a shell can hold is: 2n2 where n is the principle energy level (level 4 = 2(42) = 32 electrons.

18. The Orbital Model Electron Cloud model Wave Mechanical Model

19. * Important Definitions: –Principle Energy Level: Region around the nucleus (the dense positively charged central core of an atom) in which electrons can be found. (The closer to the nucleus, the lower the energy).

20. –Quanta: Small amount of energy that a(n) electron can absorb or release as it moves through principle energy levels. –Ground State: all electrons fill lowest energy levels before higher energy levels begin to fill.

21. –Excited State: one or more electrons fills a higher energy level before the lower ones are filled.

22. –Spectral Lines: As electrons at higher energy levels (excited electrons) fall back to their normal energy levels (ground state) they release energy in the form of the spectrum. ROYGBIV

24. Electron Configurations Looking at the periodic table of elements, you will notice numbers at the bottom of each element. These numbers represent the electron configuration of the element (the address of the electrons). The Period represents the number of principle energy levels (orbitals) present. The Group represents the sublevel for each principle energy level.

25. Principle Energy Levels Period 1 = 1 shell Period 2 = 2 shells Period 3 = 3 shells Period 4 = 4 shells Period 5 = 5 shells Period 6 = 6 shells

27. SPDF Sublevels Group 1 & Group 2 = ‘s’ sublevel (max – 2 electrons) Group 13 – Group 18 = ‘p’ sublevel (max – 6 electrons) Group 3 – Group 12 = ‘d’ sublevel (max – 10 electrons) Lanthanum & Actinum Series = ‘f’ sublevel (max – 14 e) See examples

28. Valence Electrons Electrons that fill the outermost principle energy level of an atom. Ex: Mg 2-8-2 has 2 valence electrons Ex: Ne 2-8 has 8 valence electrons Valence electrons are largely responsible for an element’s chemical and physical properties.

29. Electron Dot Diagrams The term Kernel refers to all of the non-valence electrons and the nucleus of an atom. The Kernel is represented by the element’s symbol, valence electrons are represented by dots. Example: Oxygen

30. Atomic Number: the number of Protons in the nucleus of an atom. (Atom by definition is an electrically neutral particle, so this must also be equal to the number of Electrons). Mass Number: the number of protons plus neutrons.

31. Question: Why are there fractional mass numbers (decimals) on the periodic table? (ex: Na, O 2, …) Answer: Due to the existence of isotopes. Note: Atomic Symbols: One or two letters, 1st is always capital, the 2nd is always lower case.

32. Isotope: Atoms if the same element having the same number of protons, but different number of neutrons Example: Isotopes of Hydrogen

33. ParticleProtonsNeutronsMass #Symbol Protium 101 11H11H Deuterium 112 21H21H Tritium 123 31H31H

34. Calculating Isotopes (weighted atomic mass) 1) Take the percent of each isotope and convert it back to a decimal (  100) 2) Multiply the decimal by the mass number 3)Add the numbers together to get the Weighted Atomic Mass

35. C-12 99% C-14 1% (.99 x 12) + (.01 x 14) = 12.02 Mg-26 1.75 %Mg-24 98.25% Cl-35 75%Cl-37 25%

36. Ions Atoms of the same elements having the same number of protons, but different number of electrons

37. Ca+ions: positive charge formed by losing electrons Anions: negative charge formed by gaining electrons