2. HISTORY OF THE ATOM 1910 Ernest Rutherford He oversaw his students, Geiger and Marsden, carrying out his famous Gold Foil experiment. They fired alpha particles (helium nuclei), from the radioactive decay of radium, at a piece of gold foil which was only a few atoms thick.

3. What did Rutherford expect?

4. What did Rutherford find?

6. According to Rutherford: “It was quite the most incredible event that has ever happened to me in my life. It was almost as incredible as if you fired a 15-inch shell at a piece of tissue paper and it came back and hit you. On consideration, I realized that this scattering backward must be the result of a single collision, and when I made calculations I saw that it was impossible to get anything of that order of magnitude unless you took a system in which the greater part of the mass of the atom was concentrated in a minute nucleus. It was then that I had the idea of an atom with a minute massive center, carrying a charge.” —Ernest Rutherford

7. Observations They found that: most of the alpha particles passed through the gold foil unaffected as they had predicted. some were deflected at small angles following a curved path – no real surprise. to the astonishment of all - 1 in 8000 - a few were deflected back at angles greater than 90 degrees.

8. Observations They found that: most of the alpha particles passed through the gold foil unaffected as they had predicted. some were deflected at small angles following a curved path – no real surprise. to the astonishment of all - 1 in 8000 - a few were deflected back at angles greater than 90 degrees.

9. Rutherford’s new evidence allowed him to propose a more detailed model of the atom with a central nucleus. The nucleus contains most of the mass of the atom. He suggested that the positive charge was all in a central nucleus. The moving negative electrons were held in place by electrostatic attraction. The atom is mainly empty space filled by the orbiting electrons – like planets around the Sun. However, this was not the end of the story.

10. Problems with the Rutherford Model An atomic nucleus composed of entirely positive charges should fly apart due to electrostatic repulsion. The model could not explain the total mass of the atom. Nineteenth century physics stated that an electron in motion around a central body must continuously give off radiation. This would mean that atoms should continuously glow.

11. Problems continued… If electrons are releasing radiation then they are losing energy. This loss of energy would cause their orbit radius to decrease and eventually the electron would spiral into the nucleus and matter would collapse. Along came Neils Bohr…

12. HISTORY OF THE ATOM 1913 Niels Bohr Bohr studied under Rutherford at Victoria University in Manchester. Bohr proposed a new model of the hydrogen atom. This model retained Rutherford’s nucleus but did not allow the electrons to move anywhere within the volume of space around the nucleus.

13. HISTORY OF THE ATOM The Bohr Atomic Model of the Hydrogen Atom Rutherford’s planetary model is correct. When an electron is in an "allowed" orbit it does not radiate. Bohr’s model implied that the classical electromagnetic theory did not apply at the atomic level. When an electron absorbs energy from incident electromagnetic radiation, it "quantum jumps" into a higher energy allowed state. This higher energy state corresponds to an allowed orbit with a higher value of the integer n. ( n = 1, 2, 3, 4, …) When an electron is in a higher energy state, it can quantum jump into a lower energy state, one with a smaller value of n, emitting all of its energy as a single photon of electromagnetic energy.

14. HISTORY OF THE ATOM

15. Emission spectra are produced by thin gases in which the atoms do not experience many collisions (because of the low density). The emission lines correspond to photons of discrete energies that are emitted when excited atomic states in the gas make transitions back to lower-lying levels. A continuum spectrum results when the gas pressures are higher. Generally, solids, liquids, or dense gases emit light at all wavelengths when heated. An absorption spectrum occurs when light passes through a cold, dilute gas and atoms in the gas absorb at characteristic frequencies; since the re-emitted light is unlikely to be emitted in the same direction as the absorbed photon, this gives rise to dark lines (absence of light) in the spectrum.

16. HISTORY OF THE ATOM The Bohr Model Successes The model successfully predicted the lines in the visible portion, uv portion, and infrared portion of the spectrum for hydrogen. The Bohr Model Limitations The model only explained the spectra of one-electron systems. It could not explain the emission spectra produced by atoms of two or more electrons.

17. ATOMIC STRUCTURE Particle Relative Charge (C) Mass Charge u nit of electric charge (  ) proton + 1 + 1.6 x 10 -19 1.00732 neutron no charge 0 1.00871 electron - 1 - 1.6 x 10 -19 0.00055

18. ATOMIC STRUCTURE Mass Number - the total number of protons and neutrons in an atom Atomic Number - the number of protons in an atom By definition, the word atom implies a neutral particle. Therefore, there are an equal number of protons and electrons in any atom.

19. SUMMARY 1. The Atomic Number of an atom = number of protons in the nucleus. 2. The Atomic Mass of an atom = number of Protons + Neutrons in the nucleus. 3. The number of Protons = Number of Electrons. 4. Electrons orbit the nucleus in shells. 5. Each shell can only carry a set number of electrons.

20. ATOMIC STRUCTURE Electrons are arranged in energy levels or shells around the nucleus of an atom. first energy level a maximum of 2 electrons second energy levela maximum of 8 electrons third energy level a maximum of 18 electrons (when it is the outer most energy level a max. of 8) max. # of electrons in an energy level = 2n 2

21. ATOMIC STRUCTURE There are two ways to represent the atomic structure of an element or compound; 1.Electronic Configurations: the arrangement of the electrons around the nucleus 2. Energy Level Diagrams

22. Carbon Electron Configuration – 1s 2 2s 2 2p 2 Lewis Dot Diagram for Carbon – Why can carbon form 4 bonds? Energy level diagram – Ground state/ excited state – http://www.youtube.com/watch?v=Vb6kAxwSWgU http://www.youtube.com/watch?v=Vb6kAxwSWgU – http://www.youtube.com/watch?v=K-jNgq16jEY http://www.youtube.com/watch?v=K-jNgq16jEY

23. Rules for making energy levels diagrams Pauli exclusion Principle: Electrons can not have the same 4 quantum numbers. Electrons with opposite spins can occupy the same atomic orbital. Aufbau principle: Electrons are placed in the orbitals, starting with the lowest energy orbitals first. (MAX 2 electrons per orbital). A sublevel must be filled before moving onto the next higher sublevel. Hund’s Rule: When electrons are placed in a set of orbitals of equal energy, they are spread out as much as possible to give as few paired electrons as possible.

24. Quantum Numbers (aka the electron’s address) The Principal quantum number, n is the main energy of an electron, where n= 1,2,3,4,…. (Energy LEVEL) The Secondary quantum number, l represents the shape of the electron orbit. This describes the additional electron energy sublevel. The number of values that l can have equals the principal quantum number. The values for l = 0,1,2,3 The Magnetic Quantum number, m l gives the direction of the electron orbit (or orientation in space) – orbital If l = 0, then m = 0 if l = 1, then m = -1, 0, 1 If l = 2, then m = -2,-1,0,1,2 etc…. The Spin Quantum number, m gives the electron’s spin The values for m = +1/2, -1/2 (clockwise or counter clockwise)

25. Quantum Numbers n Principal Quantum number (major shell or energy level) l Secondary Quantum Number- shape (subshell or sublevel) m Magnetic quantum number (orbitals: volume of space- 2 e /orbital) S Spin quantum number (spin) Subshell = quantum # max # of eSub- shell # of orbitals – quantum number 1s = 02 - total 2s1 0 Clockwise (+1/2) 2s = 0 p = 1 2 6 - total 8 spsp 1 0 3 -1,0,+1 or counterclockwise (- 1/2) 3s = 0 p = 1 d = 2 2 6 10 – total 18 spdspd 1 0 3 -1,0,+1 5 -2,-1,0,+1,+2 4s =0 p = 1 d = 2 f =3 2 6 10 14 – total 32 spdfspdf 1 0 3 -1,0,+1 5 -2,-1,0,+1,+2 7 -3-2-10+1+2+3

26. Some exceptions to energy level diagrams EXPECTED Cr [Ar]4s 2 3d 4 Cu [Ar]4s 2 3d 9 ACTUAL Cr [Ar]4s 1 3d 5 Cu [Ar]4s 1 3d 10 What happens? Electrons are ‘borrowed’ or ‘promoted’ from the 4s subshell to give a 3d subshell that is exactly half- filled for Cr or completely filled for Cu. A half filled or completely full subshell has a special stability as it is lower in energy.

27. Some exceptions to energy level diagrams Ag [Kr]5s 1 4d 10 Au [Xe]6s 1 4f 14 5d 10 What happens? Again the electrons are promoted from the s subshell to the d orbital as it is more stable, of lower energy… more energetically favourable to do this. U [Rn]7s 2 5f 3 6d 1 Why is Fe +2 and +3 ? Sn?? http://www.chemguide.co.uk/inorganic/group4/oxstates.html