1. SOL Review Atomic Structure and Periodicity

2. Atomic structure 1802 – Dalton 1897 – Thompson 1911 – Rutherford 1913- Bohr 1920’s Quantum mechanical model Indestrucible Plum pudding Nuclear model Planetary Mathematical model electrons Dense + nucleus e- orbit e- treated as waves in a sea of surrounded by the e- can be anywhere positive charge e- in empty space nucleus

3. Why is Rutherford’s nuclear model of the atom more consistent with the results of the alpha particle scattering experiment that with Thomson’s “plum pudding” model ?

4. (c) 2006, Mark Rosengarten Rutherford Model The atom is made of a small, dense, positively charged nucleus with electrons at a distance, the vast majority of the volume of the atom is empty space. Alpha particles shot at a thin sheet of gold foil: most go through (empty space). Some deflect or bounce off (small + charged nucleus).

5. (c) 2006, Mark Rosengarten EXAMPLE SPECTRUM Light is formed when electrons drop from the excited state to the ground state. The lines on a bright-line spectrum come from specific energy level drops and are unique to each element. No other element has the same bright-line spectrum as hydrogen, so these spectra can be used to identify elements or mixtures of elements.

6. Niels Bohr Planetary model of the atom Electrons have certain energies which allow them to stay in certain orbits around the nucleus.

7. (c) 2006, Mark Rosengarten Quantum-Mechanical Model Electron energy levels are wave functions. Electrons are found in orbitals, regions of space where an electron is most likely to be found. You can’t know both where the electron is and where it is going at the same time. Electrons buzz around the nucleus like gnats buzzing around your head.

8. Subatomic particles _______ – Positive – In nucleus – Same as atomic # ________ – Negative – Outside nucleus – 1840x’s smaller than p _______ – Neutral – In nucleus – About same mass as p

9. Atomic Symbol Isotope: element name – mass # ex. Carbon - 14 #p #p + #n#p - #e Element symbol

10. Isotopes Mass number = ______ + _______ Number of neutrons may vary Ex. 14 C – 6p, 8n 12 C – 6p, 6n

11. Isotopes 52 Crp:n:e: 13 N -3 p:n:e: 23 Na + p:n:e: 34 p38n36e Copper 64 p:n:e:

12. Atomic Mass An element has two naturally occurring isotopes as given in the table. What is the average atomic mass of the element? What is the element? Mass (amu) % abunda nce 6.027.50% 7.0292.5% Remember you are basically averaging the mass of 100 atoms!

13. Calculate the atomic mass for chlorine, given the following information: Mass number Exact weight Percent abundanc e 3534.96885275.77% 3736.96590324.23%

14. Some groups are given special names: Group 1: Alkali metals Group 2: Alkaline earth metal Group 3-12: Transition elements Group 17: Halogens Group 18: Noble gases

15. Electron Configuration Shows 1 st quantum # - ______ Shows 2 nd quantum # - _______ Shows 3 rd quantum # - __________ s=1 orbital p=3 orbitals d=5 orbitals f=7 orbitals Electron Configuration for Nitrogen

16. Electron Configuration ┌ _______ 1s 2  ______ └ _______

17. The periodic table is arranged by electron configurations. What is the electron configuration for magnesium? What does the electron configuration for zirconium end in? What element ends in 5p 3 ?

18. Periodic trends Fr F Atomic radius Ionization energy Electronegativity – noble gases do not have an electronegativity Reactivity – noble gases are inert

19. Valence Electrons Electrons in the outer most energy level of the atom. These electrons participate in bonding. Always in “s” and “p” orbitals. Draw Lewis dot diagrams for: CaAsKrZr

20. Fill the blanks: 1.The symbol and atomic number of the lightest alkaline earth metal are ________________________and _______________________. 2.The symbol and atomic number if the heaviest metalloid in group 5A (15) are _______________________and _______________________. 3.Group 1B (11) consists of the coinage metals. The symbol and atomic mass of the coinage metal whose atoms have the fewest electrons are ________________________and_______________________. 4.The symbol and atomic mass of the halogen in period 2 are ________________________and _______________________. 5.The symbol and atomic number of the heaviest noble gas are ________________________and _______________________. 6.The symbol and group number of the transition element whose atoms have the fewest protons are ________________________and _______________________. 7.The elements in group 6A (16) are sometimes called the chalcogens. The symbol and atomic number of the only metallic chalcogen are ________________________and _______________________. 8.The symbol and number of protons of the period 5 alkali metal atom are ________________________and _______________________.

21. (c) 2006, Mark Rosengarten Nuclear Chemistry

22. (c) 2006, Mark Rosengarten Alpha Decay The nucleus ejects two protons and two neutrons. The atomic mass decreases by 4, the atomic number decreases by 2. 238 92 U 

23. (c) 2006, Mark Rosengarten Beta Decay A neutron decays into a proton and an electron. The electron is ejected from the nucleus as a beta particle. The atomic mass remains the same, but the atomic number increases by 1. 14 6 C 

24. (c) 2006, Mark Rosengarten Gamma Decay The nucleus has energy levels just like electrons, but the involve a lot more energy. When the nucleus becomes more stable, a gamma ray may be released. This is a photon of high-energy light, and has no mass or charge. The atomic mass and number do not change with gamma. Gamma may occur by itself, or in conjunction with any other decay type.

25. Half-Life Length of time required for ½ of an atom of a radioactive substance to decay # HL = total time /half life Remaining mass = Mass 2 #HL Ex. Manganese-56 is a beta emitter with a half life of 2.6 hours. What is the mass of the Manganese-56 in a 1mg sample after 10.4 hours?

26. (c) 2006, Mark Rosengarten a) Write the balanced nuclear reaction for iodine 131 undergoing beta decay. b) How many grams of a 10.0 gram sample of I-131 (half-life of 8 days) will remain in 24 days?