1. HISTORY OF THE ATOM 460 BC Democritus develops the idea of atoms He pounded up materials in his pestle and mortar until he had reduced them to smaller and smaller particles which he called ATOMA (Greek for indivisible)

2. HISTORY OF THE ATOM 1808 John Dalton suggested that all matter was made up of tiny spheres that were able to bounce around with perfect elasticity and called them ATOMS

3. HISTORY OF THE ATOM 1898 Joseph John Thompson found that atoms could sometimes eject a far smaller negative particle which he called an ELECTRON

4. HISTORY OF THE ATOM Thompson develops the idea that an atom was made up of electrons scattered unevenly within an elastic sphere surrounded by a soup of positive charge to balance the electron's charge 1904 like plums surrounded by pudding. PLUM PUDDING MODEL

5. HISTORY OF THE ATOM 1910 Ernest Rutherford Oversaw Geiger and Marsden carrying out his famous experiment. They fired Helium nuclei at a piece of gold foil which was only a few atoms thick. They found that although most of them passed through. About 1 in 10,000 hit

6. HISTORY OF THE ATOM gold foil helium nuclei They found that while most of the helium nuclei passed through the foil, a small number were deflected and, to their surprise, some helium nuclei bounced straight back. helium nuclei

7. HISTORY OF THE ATOM Rutherford’s new evidence allowed him to propose a more detailed model with a central nucleus. He suggested that the positive charge was all in a central nucleus. With this holding the electrons in place by electrical attraction However, this was not the end of the story.

8. HISTORY OF THE ATOM 1913 Niels Bohr studied under Rutherford at the Victoria University in Manchester. Bohr refined Rutherford's idea by adding that the electrons were in orbits. Rather like planets orbiting the sun. With each orbit only able to contain a set number of electrons.

9. Bohr’s Atom electrons in orbits nucleus

10. HELIUM ATOM + N N + - - proton electron neutron Shell

11. ATOMIC STRUCTURE Electrons are arranged in Energy Levels or Shells around the nucleus of an atom. first shella maximum of 2 electrons outer shella maximum of 8 electrons Valence electrons are the electrons in the last shell or energy level of an atom.

12. Periodic Table Group Valence Electrons Group 1 (I) (alkali metals) alkali metalsalkali metals1 Group 2 (II) (alkaline earth metals) alkaline earth metalsalkaline earth metals2 Groups 3-12 (transition metals) transition metalstransition metals#* Group 13 (III) (boron group) boron groupboron group3 Group 14 (IV) (carbon group) carbon groupcarbon group4 Group 15 (V) (nitrogen group) nitrogen groupnitrogen group5 Group 16 (VI) (chalcogens) chalcogens 6 Group 17 (VII) (halogens) halogens Group 18 (VIII) (noble gases) 78** #* The general method for counting valence electrons is generally not useful for transition metals. ** Except for helium, which has only two valence electrons.helium

13.  Atoms with a complete shell of valence electrons tend to be chemically inert.chemically inert  Atoms with one or two valence electrons more than a closed shell are highly reactive because the extra electrons are easily removed to form positive ions.ions  Atoms with one or two valence electrons less than a closed shell are also highly reactive because of a tendency either to gain the missing electrons and form negative ions, or to share electrons and form covalent bonds.covalent bonds

14. Octet Rule = atoms tend to gain, lose or share electrons so as to have 8 electrons H would like to N would like to O would like to Lose 1 electron Gain 3 electrons Gain 2 electrons

16. E lectron from Na is transferred to Cl, this causes a charge imbalance in each atom. The Na becomes (Na+) and the Cl becomes (Cl-), charged particles or ions. IONIC BONDING

18. bond formed by the sharing of electrons. Between nonmetallic elements of similar electronegativity. Covalent bonding-bond formed by the sharing of electrons. Between nonmetallic elements of similar electronegativity. Two Types of Covalent Bonds Non-polar covalent bond – equal sharing of electrons Polar covalent bond – unequal sharing of electrons What does the word polar mean?

19. Symbols of atoms with dots to represent the valence-shell electrons Bonds in all the polyatomic ions and diatomic molecules are all covalent bonds Ammonium NH 4 + Cl 2 H2H2

21. Oxygen Atom Oxygen Molecule (O 2 ) Oxygen Molecule (O 2 ) Electron Configuration Model

22. Polar Covalent Bonds: Unevenly matched, but willing to share.

23. - water is a polar molecule because oxygen is more electronegative than hydrogen, and therefore electrons are pulled closer to oxygen.

24. Molecular Formula Structural Formula Lewis Dots Ball-and-Stick Space-Filling

25.  (Greek isos = "equal", méros = "part") are compounds with the same molecular formula but different structural formulamolecular formulastructural formula A simple example of isomerism is given by propanol: it has the formula C 3 H 8 O (or C 3 H 7 OH) and occurs as two isomers: propan-1-ol (n-propyl alcohol; I) and propan-2-ol (isopropyl alcohol; II) (methyl- ethyl-ether; III) propanolC H O OHpropan-1-olpropan-2-ol(methyl- ethyl-ether; III)

26. The difference in electronegativity of the elements can be used to determine the type of bond that is present. 1) What is electronegativity? It is a measure of how strongly an atom attracts the bonding electrons in a chemical bond. The higher the electronegativity, the stronger an atom's attraction for bonding electrons. Intermolecular Forces-attraction between molecules

27. Pauling Scale of Electronegativities

28. Electronegativity difference (approx.) Type of Bond Example 0.0-0.4 Non-polar covalent H-H (0.0) 0.4-1.0 Moderately polar covalent H-Cl (0.9) 1.0-2.0 Very polar covalent H-F (1.9) ≥ 2.0 Ionic Na + Cl - (2.1)