1. Atomic Number Number of Protons

2. Mass Number Number of Protons + Neutrons

3. 12 C Left Superscript = mass number 6

4. 12 C Left Subscript = atomic number 6

5. 80 Br 35 Atomic Number = ?

6. 20 Ne 20 10 Mass Number = ?

7. 27 Al 27 13 Mass Number = ?

8. 40 Ca 20 Atomic Number = ?

9. Isotope Atoms of the same element with a different # of neutrons

10. Characteristics of Proton Charge = +1, mass = 1 amu, location = inside nucleus

11. Characteristics of Neutron Charge = 0, mass = 1 amu, location = inside nucleus

12. Characteristics of Electron Charge = -1, mass = 1/1836 amu or 0.0005 amu, location = outside nucleus

13. Ion An atom that has gained or lost electrons & so carries charge

14. Nucleons Protons & Neutrons

15. atom Smallest bit of an element that retains the properties of the element.

16. atom Electrically neutral. # of protons = # of electrons.

17. Charge # protons - # electrons

18. # of neutrons Mass number – atomic number

19. 14 C 8 6 # of neutrons = ?

20. 9 Be 5 4 # of neutrons = ?

21. 40 Ar 22 18 # of neutrons = ?

22. 15 N 8 7 # of neutrons = ?

23. 24 Mg Right superscript = charge 12 2+

24. 24 Mg 10 electrons 12 2+ # of electrons?

25. 86 Rb 36 electrons 37 1+ # of electrons?

26. 127 Te 53 electrons 52 1- # of electrons?

27. 32 S 18 electrons 16 2- # of electrons?

28. Cation Positive ion: atom lost electrons

29. Anion Negative ion: atom gained electrons

30. Avg. Atomic Mass Weighted avg. of masses of naturally occurring isotopes of an element.

31. 2 isotopes of Cl: 75% Cl-35 & 25% Cl-37. Calculate avg. atomic mass. Avg. atomic mass =.75(35) +.25(37) = 35.5 amu

32. Dalton’s Model Billiard Ball Model

33. Thomson’s Model Plum Pudding Model - - - - - + + + + +

34. Rutherford’s Model Nuclear Model - - -

35. Rutherford’s Experiment Source: http://www.dlt.ncssm.edu/TIGER/chem1.htm#atomic

36. Rutherford’s Experiment: Results 1)Most of the alpha particles went straight through.  Most of the atom is empty space. 2)Some of the alpha particles were deflected back.  The nucleus was tiny, but contained most of the mass of the atom.

37. Bohr’s Model Planetary Model

38. Schrodinger’s Model Modern or Quantum Mechanical Model Source: http://www.dlt.ncssm.edu/TIGER/chem1.htm #atomic

39. Bohr Configuration Ground state configurations found in reference tables. Cannot be predicted.

40. Bohr Configuration of Na = 2-8-1 2 electrons in energy level 1 8 electrons in energy level 2 1 electron in energy level 3

41. Bohr Diagram of Na +11

42. Bohr Model Electrons are restricted to specific orbits or shells or principle energy levels. Each shell holds a specific # of electrons. Each shell has a specific energy & radius. Energy of electron must match energy of shell.

43. Maximum Capacity of Bohr Levels Shell #Max # of electrons 1 2 3 4 n 2 8 18 32 2n 2

44. Ground State Bohr model Every electron is in the lowest available orbit.

45. Excited State Bohr model An electron has absorbed heat, light, or electrical energy and moved to a higher energy level. Unstable. Returns to ground state quickly by emitting a photon.

46. Continuous Spectrum Spectrum produced by holding a prism in sunlight. Contains light at every wavelength.

47. Bright Line Spectrum Visible light produced by electrons in atom returning to ground state: light of only a few wavelengths is present. Each element has a unique bright line spectrum. Used to identify elements. Wavelengths of bright lines correspond to difference between energy levels. Source: http://www.dlt.ncssm.edu/TIGER/chem1.htm#atomic

48. Absorbtion of Energy h Ground state Excited state E1E1 E2E2 E3E3

49. Emission of Energy h Ground state Excited state E1E1 E2E2 E3E3

50. Orbital Modern Model Region of space that holds 2 electrons. Has a specific energy. Shapes vary.