1. History of Atomic Theory

2. Scientists B.C. DemocritusAristotle 384-322 BC Believed matter is continuous 400 BC Coined the term “atom”

3. Dalton’s Atomic Theory -early 1800’s- All matter is composed of tiny, indivisible particles called atoms. All matter is composed of tiny, indivisible particles called atoms. Atoms of the same element have the same properties (mass, size, etc.). Atoms of the same element have the same properties (mass, size, etc.). In a chemical reaction, matter cannot be created or destroyed. (Law of Conservation of Mass) In a chemical reaction, matter cannot be created or destroyed. (Law of Conservation of Mass) Compounds always contain elements in the same ratio by mass (Law of Definite Proportions) Compounds always contain elements in the same ratio by mass (Law of Definite Proportions)

4. Atomic size A penny contains 2.4 x 10 22 atoms A penny contains 2.4 x 10 22 atoms Radius of one atom is around 2 x 10 -10 m Radius of one atom is around 2 x 10 -10 m or.2 nm Scanning tunneling microscope can generate images of individual atoms. Scanning tunneling microscope can generate images of individual atoms.

5. Thomson’s Cathode Ray Tube -late 1800’s- Showed that electrons are negatively charged particles. Image from Addison Wesley Chemistry

6. Thomson’s “plum pudding” model

7. Rutherford’s gold foil exp. -early 1900’s- Conclusion: Most of an atom’s volume is empty space.

8. Rutherford’s “planetary” model

9. 5 Models of the Atom (a) Dalton's model (1803) (b) Thomson's model (1897) (c) Rutherford's model (1909) (d) Bohr's model (1913) © Prentice-Hall, Inc. (e) Electron-cloud model (present)

10. Subatomic particle chargelocationmass Other feature proton+Nucleus 1 amu Defines the element -atomic no. neutron0Nucleus 1 amu Change no. to form isotopes electron-Electroncloud ~0~0~0~0 atom’s volume -dictates reactivity

11. Nuclear Forces Short-range forces that hold the nuclear particles together. Short-range forces that hold the nuclear particles together.

12. Isotopes Atoms of the same element that differ in mass Atoms of the same element that differ in mass Atomic no.=# protons Atomic no.=# protons #protons=#electrons #protons=#electrons Mass no.=#protons + # neutrons (nucleons) Mass no.=#protons + # neutrons (nucleons) Num f neutrons

13. Isotopes of Hydrogen NuclideProtonsNeutronsMassNumber Protium101 Deuterium112 tritium123

14. Isotopes can be written two ways or bromine-80

15. Electrons Found in an electron cloud outside of the nucleus (but not in paths like the planets) Found in an electron cloud outside of the nucleus (but not in paths like the planets) 1 st energy level holds 2 electrons 1 st energy level holds 2 electrons 2 nd energy level holds up to 8 2 nd energy level holds up to 8 3 rd energy level holds up to 18 3 rd energy level holds up to 18

16. Periodic Table Arranged by increasing atomic number Arranged by increasing atomic number Rows are called periods Rows are called periods Columns are called groups Columns are called groups

18. Average Atomic Mass An element’s atomic mass is the weighted average of its naturally occurring isotopes. An element’s atomic mass is the weighted average of its naturally occurring isotopes.

19. Average Atomic Mass Multiply the mass of each isotope by its abundance to get the weighted average. (% x mass)+ (% x mass) +... (% x mass)+ (% x mass) +... 100 100

20. Ex.: Boron is 80.20% boron-11 (atomic mass 11.01 amu) and 19.80% boron-10 (atomic mass 10.01 amu). What is the average atomic mass of boron? What is the average atomic mass of boron? (11.01amu)(80.20) + (10.01amu)(19.80) = 100 100 =10.81 amu

21. Sample Problem ex.: Neon has 2 isotopes. Neon-20 has a mass of 19.992amu and neon-22 has a mass of 21.991amu. In an average sample of neon atoms, 90% will be neon-20 and 10% will be neon-22. Calculate the average atomic mass. (90 x 19.992amu)+(10 x 21.991amu) (90 x 19.992amu)+(10 x 21.991amu) 100 100 = 20.192 amu