1. Ch. 3 - Atomic Structure Subatomic Particles – show powers of ten ect…..

3. Development of the Atom 1 st Model: Dalton’s Theory 1. All matter is made of indivisible atoms. 2. All atoms of a given element are identical. 2 nd Model: Thomson’s model “Plum-pudding” model. Discovered electron (negative charges) through Cathode Ray Tube Experiment. 3 rd Model: Rutherford’s Model, Discovered the nucleus made of protons through Gold Foil Experiment.

4. Development of the Atom  4 th Model: Bohr’s Model “Planetary” model: Electrons orbit the nucleus on specific paths.  Chadwick: Discovered the neutrons in the nucleus. Neutrons are the “glue” of the atom, or strong nuclear force.

5. Development of the Atom  5 th Model: Electron Cloud Model, the current model of the atom. Developed by Schrodinger. Electrons are in “clouds” or regions of space outside the nucleus. Also called orbitals or energy levels.

6. Subatomic Particles Most of the atom’s mass. NUCLEUS ELECTRON CLOUD PROTONS NEUTRONS ELECTRONS POSITIVE CHARGE NEUTRAL CHARGE ATOM NEGATIVE CHARGE Atoms Size Smallest particle that retains the chemical identity of an element

7. Subatomic Particles  Quarks component of protons & neutrons 6 types 3 quarks = 1 proton or 1 neutron He

8. Subatomic particles  In a neutral atom, the # of protons equals the # of electrons.  A proton is more massive than an electron: 1 proton equals approx. 2000 electrons!

9. Subatomic particles  Atomic Number (Moseley): the # of protons in the nucleus of an atom Unique to each element For example: Oxygen: Atomic # 8 8 protons 8 electrons Positive charge = Negative charge in a neutral atom

10. Subatomic particles  Mn (manganese) # of protons: # of electrons:  Kr (Krypton) # of protons: # of electrons: 25 36

11. Mass Number  mass # (atomic mass or atomic weight)= protons + neutrons  round the atomic mass to the 2 nd decimal when calculating © Addison-Wesley Publishing Company, Inc.

12. Subatomic particles Charge (c)Mass (g)Mass (amu) Proton1.6 x 10 -19 1.67 x 10 -24 1 Neutron 01.67 x 10 -24 1 Electron -1.6 x 10 -19 9.11 x 10 -28 0

13. Subatomic particles  Diameters: 0.100 to 0.500 nm If you drew a line through the diameter of a penny (1.9 cm), it would cross 810 million Cu atoms!

14. Mass Number  Mass number – the atomic mass (weight) rounded to the nearest whole number when displayed like this. Mass # Atomic #  Nuclear symbol:  Hyphen notation: carbon-12

15. Ions  Ions – atoms with a net electrical charge (lose or gain one or more elctrons - # of protons is unchanged) Oxygen O –2 gains 2e- HydrogenH + loses 1e- MagnesiumMg +2 or Mg 2+

16. Ions  Write the symbol: 9 protons, 10 electrons (anion) 13 protons, 10 electrons (cation) S –2 has how many protons and electrons? F -1 or F 1- or F - Al +3 or Al 3+ 16 p + and 18 e -

17. Ions  Iron has how many protons, electrons and neutrons?  Fe +2 has how many protons, electrons, and neutrons? 56 26 Fe 26 protons, 26 electrons, 30 neutrons 26 protons, 24 electrons, 30 neutrons (h.o.)

18. Isotopes … …don’t have to be radioactive. Some isotopes are unstable and decay But, there are many stable isotopes that don’t decay. Atoms of the same element with different mass numbers ( diff. # of neutrons )

19. Isotopes © Addison-Wesley Publishing Company, Inc.

20. Isotopes  Hydrogen Isotopes: Hydrogen-1 : 1 p, 1 e, 0 n (protium) Hydrogen-2 : 1 p, 1 e, 1 n (deuterium) Hydrogen-3 : 1 p, 1 e, 2 n (tritium)

21. Some elements have several Isotopes Lead has four naturally occurring isotopes, Pb-204, Pb-206, Pb-207, and Pb-208; but there are 23 man- made isotopes of lead. Isotopes

22.  The periodic table gives the average atomic mass of all the naturally occurring isotopes according to how abundant the isotope is.  If the avg. atomic mass is rounded to the nearest whole number it will tell the most abundant isotope. Carbon: 12.0107 amu so…… Hydrogen: 1.00794 amu so….

23. Relative Atomic Mass  12 C atom = 1.992 × 10 -23 g © Addison-Wesley Publishing Company, Inc.  atomic mass unit (amu – Symbol:  )  1 amu= 1 / 12 the mass of a 12 C atom  1 amu = 1.66 × 10 -24 g  Carbon: 12.011 amu  NOT 12.011 g !!

24. Isotopes  Chlorine-37 atomic #: mass #: # of protons: # of electrons: # of neutrons: 17 37 17 20 (Is this the most abundant isotope?)

25. Average Atomic Mass  weighted average of all isotopes  on the Periodic Table (write this on p.t.)  round to 2 decimal places Avg. Atomic Mass

26. Avg. Atomic Mass Average Atomic Mass  EX: Calculate the avg. atomic mass of oxygen if its abundance in nature is 99.76% 16 O, 0.04% 17 O, and 0.20% 18 O. 16.00 amu

27. Average Atomic Mass  Find the average atomic mass of neon if : 90.48 % Neon-20 19.9924 amu, 0.27 % Neon-21 20.9938 amu, 9.25 % Neon-22 21.9914 amu. = 20.18 amu

28. Avg. Atomic Mass D. Average Atomic Mass  EX: Find chlorine’s average atomic mass if approximately 8 of every 10 atoms are chlorine-35 and 2 are chlorine-37. 35.40 amu