2. Fourth century, B.C. – The Greek Philosopher Democritus suggested that the universe was made of indivisible units, which he called atoms. Democritus

3. Every element is made up of tiny particles called atoms that cannot be subdivided. Atoms of the same element are exactly alike. Atoms of different elements can join in simple whole number ratios to form molecules. Reactions occur when atoms are separated, joined, or rearranged. Dalton’s Atomic Theory (1800’s)

4. J.J. Thomson (1897)

5. Thomson’s model

6. Rutherford’s Gold-Foil Experiment (1911)

7. Rutherford Simulation

8. Electrons Protons Neutrons Subatomic Particles

9. Atomic number (Z) is the number of protons in the nucleus of an atom Mass number (A) is the total number of protons and neutrons in an atom. – # of neutrons = A – Z Atoms are identified using atomic number and mass number. – Example: 197 79 Au or gold-197 Distinguishing Among Atoms

10. Isotopes are different versions of the same atom. They have the same number of protons and a different number of neutrons. Isotopes

11. – The atomic mass is measured in atomic mass units (amu). One amu = 1/12 of the mass of the carbon-12 isotope. – The average atomic mass is a weighted average of the masses of all the naturally occurring isotopes of each element. Atomic Mass

12. Average Atomic Mass of Chlorine

13. 98.93% of all of the Carbon earth is Carbon-12 with a mass of 12.0000 amu, and the remaining 1.07% is Carbon-13 with a mass of 13.003355 amu. Calculate the average atomic mass. Examples

14. 69.15% of all of the copper on earth is copper-63 with a mass of 62.929601amu. The remaining 30.85% is copper -65 with a mass of 64.927794 amu. What is the average atomic mass of copper? Examples

15. The Bohr Model (1913) 1913 – A Danish scientist Niels Bohr suggested that the electrons revolve around the nucleus like the planets revolve around the sun. In Bohr’s model each electron has a certain amount of energy that is defined by its path around the nucleus. These defined orbits are called energy levels. Only worked for Hydrogen

16. Electrons can only be located within an energy level according to Bohr. They can never be found in between the energy levels. Electrons can either gain energy and move to outer energy levels, or lose energy by emitting photons of light and move to inner energy levels.

17. Atomic Spectra When electrons absorb energy they move to higher energy levels, and when electrons lose energy by emitting light they return to their ground state.

19. Quantum Mechanics De Broglie matter waves = h/mv h = 6.63 x 10 -34 Js Heisenberg uncertainty principle

20. Quantum Mechanical Model (1926) Erwin Schrödinger devised mathematical equations that could determine the allowed energy for an electron and the likelihood of finding an electron around the nucleus. Solving the equations gives the allowable energies an electron can have. An atomic orbital is a visual representation of Schrödinger's mathematical equation of the probability of finding an electron around the nucleus of an atom.

21. Quantum Numbers Each electron in an atom is defined by 4 quantum numbers: n, l, m, and s. nPrincipal quantum number corresponds to the energy level lThis quantum number corresponds to the shape of the atomic orbital: s,p,d, or f mThis number identifies the orientation of the orbital in space sThis number corresponds to the spin of the electron.

22. Atomic Orbitals

23. Electron Configurations Aufbau principle – electrons occupy the orbitals of lowest energy first.

24. Electron Configuration Pauli exclusion principle – each orbital can hold two electrons however they must have opposite spins. – Arrows are used to indicate spin Hund’s rule – electrons occupy orbitals of the same energy in a way that makes the number of electrons with the same spin as large as possible. – One electron goes in each suborbital before doubling up

25. Electron Filling Order As you move down the periodic table each atom contains more electrons. There are two ways of writing electron configuration. The first is spectroscopic notation. Example: Hydrogen atomic number 1 1s 1 Energy level (n) Orbital (l) Number of electrons

26. Electron Filling Order The second is orbital box diagrams. Example: Hydrogen atomic number 1 ___ 1s Energy Level (n) Orbital (l) Electron spin (+1/2 or –1/2)

27. Examples Use the orbital box diagrams and draw out the electron configuration for carbon in its ground state. Also write out carbons electron configuration using spectroscopic notation.

28. Examples Use the orbital box diagrams and draw out the electron configuration for vanadium in its ground state. Also write out vanadium electron configuration using spectroscopic notation.

29. Exceptions Full sublevels are the most stable, however half-filled sublevels are more stable than partially filled sublevels. Draw the orbital box diagram for copper in it’s ground state. What other elements would follow this exception?

30. Periodic Table Elements in the same group have similar properties, because they have the same number of valence electrons. The columns are also arranged on the periodic table according to the electron orbitals of the different atoms.