1. ATOMIC THEORY

2. Defining the Atom  An atom is the smallest particle of an element that retains its identity in a reaction.  The basic building blocks of matter that make-up everyday objects.

3. DemoCritus  Early Greek Scholar-was the first to suggest the existence of atoms  He believed that atoms were indivisible and indestructible.  Never developed a theory  lacked experimental support

4. Dalton’s Atomic Theory  Put Democritus’s ideas into a scientific theory.  1) All matter is composed of tiny indivisible particles called atoms.  2) Atoms of the same element are identical.  3) Atoms of different elements can mix together to form compounds  4) Chemical reactions occur when atoms are separated, joined, or rearranged.

5. Dalton’s Theory Revised  Most of Dalton’s theory is still accepted today EXCEPT that atoms are known to be divisible.  Atoms can be broken down into 3 subatomic particles: electrons, protons and neutrons.

6. J.J. Thomson  Used a cathode ray tube to prove the smallest particles present must have a negative charge.  He discovered the ELECTRON! http://www.youtube.com/watch?v=RW_z fKOU9uM&feature=related

7. J.J. Thomson  Atoms were electrically neutral, so there must be a + particle to cancel out the – charge from the electron.  Developed the Plum Pudding Model (positive ball containing scattered electrons)

8. Ernest Rutherford  Former student of Thomson, disproved the Plum Pudding Model of the atom.  The Gold Foil Experiment: Sent a beam of + charges (alpha particles) through a piece of very thin gold foil.  Angles of deflection were measured.

9. Rutherford  Results: Most of the alpha particles passed straight through, most of the foil must be regions of “empty” space – not a + sphere like Thomson believed.  + charges and the atoms mass must be found in the center  discovered the nucleus

10. The Nuclear Atom  In Rutherford’s atomic model, the protons and neutrons are located in the nucleus.  The electrons are distributed around the nucleus and occupy almost all the volume of the atoms.

11. ATOMIC STRUCTURE Particle proton neutron electron Charge + ve charge -ve charge No charge 1 1 nil Mass

12. ATOMIC STRUCTURE the number of protons in an atom the number of protons and neutrons in an atom He 2 4 Atomic mass Atomic number number of electrons = number of protons

13. HELIUM ATOM + N N + - - proton electron neutron Shell Are atoms electrically neutral? Why?

14. Complete the following table in your notes Atomic # Mass # # of Protons # of Neutrons # of Electrons 910 1415 4722 5525

15. Isotopes  Atoms that have the same number of protons but different numbers of neutrons  Different mass numbers  Chemically alike because they have identical numbers for the characteristic chemical behavior of each element  Ex: Three known isotopes for H –H: Hydrogen (no neutrons, mass # of 1) –H-2: Deuterium (one neutron, mass # of 2) –H-3: Tritium (two neutrons, mass # of 3)

16. Calculating Average Atomic Mass of Isotopes  In nature, isotopes occur in various percentages. The higher the percent the more abundant.  In order to figure out the average mass of each element the percent abundance and mass of each isotope need to be considered  We can calculate average atomic mass in much the same way as we calculate your grade in this class…

17. Calculating Average Atomic Mass 1.Divide the percent abundances by 100. (natural occurrence) 2.Multiply each isotope mass by its natural occurrence. (maintain sigfigs)* 3.Add up all the masses (maintain place values) 4.Include a unit (amu)

18. What are the different categories that you are graded on in this class?  Classwork: 79pts  Practice: 12pts  Final: 14pts  What would your semester grade be if you received an 81% for classwork, 52% for practice, and 73% on your final? –0.81 x 79 = 64 –0.52 x 12 = 6.2 –0.73 x 14 = 10 –Add all answers together to get % semester grade –63.2 + 3.5 + 10.1 = 80.2  80 % (a B)

19. Now lets try with an element!  Copper has two isotopes: copper-63 and copper-65. The relative abundances of these isotopes are 69.2% and 30.8% respectively. Calculate the average atomic mass of copper. 0.692 x 63 = 43.6 0.308 x 65 = 20.0 43.6 + 20.0 = 63.6

20. One more example…  Uranium has three naturally occurring isotopes with the following percent abundances: U-234 (0.0058%), U-235 (0.71%), and U-238 (99.23%). –What do you expect the average atomic mass to be and why? –What is the average atomic mass?  237.9

21. SUMMARY 1. The Atomic Number of an atom = number of protons in the nucleus. 2. The Atomic Mass of an atom = number of Protons + Neutrons in the nucleus. 3. The number of Protons = Number of Electrons. 4. Electrons orbit the nucleus in shells. 5. Each shell can only carry a set number of electrons.

22. ATOMIC STRUCTURE There are two ways to represent the atomic structure of an element or compound; 1.Electronic Configuration 2.Dot & Cross Diagrams

23. ELECTRONIC CONFIGURATION With electronic configuration elements are represented numerically by the number of electrons in their shells and number of shells. For example; N Nitrogen 7 14 2 in 1 st shell 5 in 2 nd shell configuration = 2, 5 2 + 5 = 7

24. ELECTRONIC CONFIGURATION Write the electronic configuration for the following elements; Ca O ClSi Na 20 40 11 23 8 17 16 35 14 28 B 11 5 a)b)c) d)e)f) 2,8,8,22,8,1 2,8,72,8,42,3 2,6

25. DOT & CROSS DIAGRAMS With Dot & Cross diagrams elements and compounds are represented by Dots or Crosses to show electrons, and circles to show the shells. For example; Nitrogen N XX X X XX X N 7 14

26. DOT & CROSS DIAGRAMS Draw the Dot & Cross diagrams for the following elements; OCl 817 16 35 a)b) O X X X X X X X X Cl X X X XX X X X X X X X X X X X X