1. Atomic Theory
Chapter 4

2. History
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3. Democritus 400BC Greek philosopher
Thought all matter was made up of atoms (atomos), which are the basic, indivisible particles of matter

4. Aristotle Believed all matter was continuous-did not believe in atoms
His opinion was accepted for nearly years

5. Late 1700’s Scientists agreed that
Most natural materials are mixtures of pure substances.
Pure substances are elements or compounds.
Law of constant composition-a compound always has the same composition.

6. John Dalton
English schoolteacher who tied all three laws together in his atomic theory
Dalton turned Democritus’s idea into a scientific theory that could be tested by experimentation.

7. Dalton’s Atomic Theory
Elements are made of tiny particles called atoms.
All atoms of a given element are identical.
The atoms of an element are different from those of any other element.
Atoms of one element combine with others to form compounds.
Atoms are indivisible in chemical processes. A chemical reaction changes the way the atoms are grouped together.

8. Joseph John Thomson
Cathode-ray experiment: showed that the atoms of any element can be made to emit tiny negative particles
Determined the charge ratio of electrons

9. William Thomson
Plum pudding model-a bunch of positive stuff with the electrons scattered throughout.

10. Rutherford, Geiger, Marsden-nucleus
Gold foil experiment, which led to the discovery of the nucleus.
Like bullets through a tissue

11. Florescent
Screen
Lead block
Uranium
Gold Foil

12. What he expected

13. Because, he thought the mass was evenly distributed in the atom.

14. What he got

15. +

16. Atomic Structure

17. Protons
The number of protons in an atom determines the element’s identity
Nuclear forces hold the nuclear particles together
The atomic number equals the number of protons

18. Electrons Are very small.
If the nucleus is a grape, the electrons would be about one mile away.
Have a negative charge
The arrangements of electrons determines the element’s chemical properties.

19. Neutrons Mass number= protons+neutrons
Neutrons=mass number-atomic number
Isotope-atoms that have the same number of protons and electrons but different numbers of neutrons (disproves point 2 of Dalton’s theory)
Nuclide-any isotope of any element

20. Table 2.1 The Mass and Charge of the Electron, Proton, and Neutron

21. Practice- Give the protons, neutrons, and electrons for each
Mercury
Sodium
Carbon
13C 6

22. Answers Mercury 80p, 80e, 121n Sodium 11p, 11e, 12n Carbon 6p, 6e, 6n
13C p, 6e, 7n 6

23. Ions
An ion is formed when we remove or add an electron to a neutral atom.
Cation-a positive ion
Anion-a negative ion

24. Ion Example
Regular sodium has 11 electrons, 11 protons, and 12 neutrons. If we take away 1 electron, it would have 10 electrons (-), and 11 protons (+) so the charge would be +1. (Neutrons would stay the same).

25. Periodic Table

26. Organization
Groups/families-vertical columns
Periods-horizontal rows

27. Metals
Conductors
Lose electrons (cations +)
Malleable and ductile

28. Nonmetals
Brittle
Gain electrons (anions -)
Covalent bonds

29. Semi-metals or Metalloids

30. Alkali Metals

31. Alkaline Earth Metals

32. Halogens

33. Transition metals

34. Rare Earth Metals (Inner transition metals)

35. Noble Gases

36. Periodic Table Label the following on your periodic table
Group Charge 1 and and
Group Charge none

37. +1 -4 -3 -2 -1 +2 +3 +1 +2

38. Trends
Notice that metals tend to give up electrons while nonmetals tend to gain electrons.

39. Ionic Compounds
Contains a metal and a nonmetal (causes ions that is why it is called ionic)
The net charge of an ionic compound has to be zero.

40. Ionic Compound Examples
Sodium chloride Na +1 Cl NaCl
Magnesium chloride Mg +2 Cl -1 MgCl2