1. Atomic Structure

3. Relative masses/charges of protons, neutrons and electrons Sub-atomic particle Relative Mass Relative Charge Proton (p)1+1 Neutron (n)10 Electron (e)5 x 10 -4

4. ProtonsNeutronsElectronsElectronic Configuration

5. Isotopes Isotopes are atoms of the same element with the same ________(same number of protons) but different __________(different no. of neutrons). IsotopeProtonsNeutronsElectrons Pg 57 Test yourself

6. Isotopes react the same way (same chemical properties) N 2 + 3H 2 2NH 3 N 2 + 3D 2 2ND 3 have different physical properties has a boiling pt of -253 0 C whereas has a boiling pt of -250 0 C Why? Chemical properties – the same number and arrangement of electrons Physical properties – different masses hence move at diferent speed

7. Radioactive Isotopes Many isotopes are radioactive as the nuclei of these atoms break down spontaneously, emitting radiation. 3 different forms of radiation -Gamma radiation : highly penetrating -Alpha radiaton : can be stopped by a few cm of air -Beta radiation : can be stopped by a thin sheet of aluminium.

8. The mass spectrometer Atoms/molecules are bombarded by high energy electron. Atoms are ionised by knocking one or more electrons off to give a positive ion. M(g) + e  M + (g) + 2e Ions are deflected by a magnetic field. The beam of ions is detected electrically Positive ions are accelerated in an electric field. http://www.chemguide.co.uk/analysis/masspec/howitworks.html

9. Different ions are deflected by the magnetic field by different amounts. The amount of deflection depends on: the mass positive charge on the ion. The lower the mass/charge (m/z) ratio, the more the ions are deflected Lighter ions are deflected more than heavier ones. Ions with 2 (or more) positive charges are deflected more than ones with only 1 positive charge. Lightest ions Heaviest ions

10. Which ion will deflect the most in a mass spectrometer?

11. The mass spectrum shows that iron has 4 isotopes as follows: Calculate the relative atomic mass of iron. The mass spectrometer measures the relative abundance of different isotopes (atoms) of an element. The output is a mass spectrum. Pg 61 Test yourself

12. An element has an atomic number of 24. The natural element consists of four isotopes. The mass spectrum of the element X produced the following peaks of three of its isotopes on the chart recorder. (a)What is element X classified in the Periodic Table? (b)Calculate the isotopic mass of the 4 th isotope if the relative atomic mass of element X is 52.06.

13. Rutherford’s planetary model Electron should emit energy and spirally fall on the nucleus. But the atome is stable.

14. Niels Henrik Bohr Tackled the question of why the electrons did not fall on the nucleus. Conducted experiments on the emission spectra of hydrogen atoms. The light emitted created a line spectrum.

15. Excitation of Electrons When an atom receives enough energy, the electron can jump to higher levels, farther from the nucleus. The electrons are in __________ states. Once excited (unstable), the electron rapidly falls to a lower level. The excited state had an excess energy. So, the atom must __________ energy. We observe the emission of a ______________________.

16. The Bohr Model The electron can circle the nucleus only in __________ orbits designated by a quantum number, n. Those orbits correspond to specific _______ distance from the nucleus and correspond to a very specific ________ state of the electron. The electron CANNOT exist at an other distances or have any other energy except the allowed ones.

17. The Bohr Model The quantum number can have integer values, n = 1 corresponding to the orbit __________to the nucleus. When an electron resides in the orbit designated by n = 1 it is said to be in the ___________ state.

18. Excited electrons give out energy when they return to their ground (stable) state. Loss of energy seen as visible light or other radiation. Captured on photograhic plates as emission spectral.

19. Page 56

20. The visible hydrogen spectrum The spectrum consists of discrete lines and that the lines converges towards the high energy (violet) end of the spectrum. The lines in the spectrum get closer together at higher frequency / energy Energy = h x frequency Energy = h x speed of light / wavelength Page 63

21. Continuous Spectrum White light is made up of all the colours of the spectrum. When it is passed through a prism, a continuous spectrum of all the colours can be obtained. A continuous spectrum contains all wavelengths from a band of the electromagnetic spectrum.

22. In summary The H atom has only ertain allowed energy levels called _________ states. The atom does not ______ energy while in one of these stationary states. The atom changes to another stationary state only by absorbing or emitting a ________ whose energy equals the difference between the 2 states. Bohr was thus able to explain the distinct lines (not continuous) in the atomic spectra of hydrogen

23. Precedents to Quantum Theory However, Bohr’s theory of distinct energy levels could explain electron behavaiour in atoms beyond __________. At this time, a lot of work about the nature of light was being investigated. Light was seen both as a wave and a particle – the idea of duality.

24. Quantum Mechanical Model Erwin Schrodinger Werner Heisenberg Wolfgang Pauli Friedrich Hund

25. In order to explain the nature of the electron, the idea of electrons being particles was changed to that of a dual nature (De Brogile)

26. Erwin Schrodinger The position of the electron was determined using a ________ equation (incoporate both the wave & particle properties of the electron)

27. Quantum Mechanical Model Erwin Schrodinger’s work showed that electrons do not circle orbits. Since the position of an electron varies, the term “orbital” was used, instead of shells. Electrons can be found with a high probability in specific regions of space called “orbitals” An orbital is defined as the region withn which there is a ______% probability of locating a particular electron in a free atom.

28. Quantum Mechanical Model In quantum theory, the electron shells are not fixed orbits but clouds of probability. Cannot measure the exact location of the electron.

29. Quantum Mechanical Model The electron is located within a sphere (or shell) around the nucleus – the probability of finding it near the nucleus is higher, but never 100%. The quantum shells are called orbitals.

30. Quantum Mechanical Model Modern quantum theories lead to stable locations of electrons, which are not exact planetary orbits but are characterized by specific quantum numbers.

31. Principal Quantum Number, n Values of n = 1,2,3,... (positive integers only) Determines orbital size i.e. the larger the value of n, the ___________ the orbital. Also, determines the __________ level of the electrons it contains and the average ___________ from the nucleus.

32. Principal Quantum Number, n Each shell is characterized by a different principal quantum number (n). Larger n => __________shell is from nucleus and ___________ in energy.

33. Principal Quantum Number, n n also determines the ____________ number of electrons in the shell : 2n 2 Shell12345 Max no. of electrons28183250

34. Modern Atomic Structure n = 1 n = 2 n = 3 n = 4 Hein, Arena, Foundations of College Chemistry, 2000, page 202 1s1s 2s2s 3s3s 4s4s 2p2p 3p3p 4p4p 3d3d 4d4d4f4f Sublevel designation An orbital for a hydrogen atom. The intensity of the dots shows that the electron spends more time closer to the nucleus. The first four principal energy levels in the hydrogen atom. Each level is assigned a principal quantum number n. The types of orbitals on each of the first four principal energy levels.

35. s,p, and d-orbitals

36. Energy Level Diagram of a Many-Electron Atom Arbitrary Energy Scale 18 32 8 8 2 1s 2s 2p 3s 3p 4s 4p 3d 5s 5p 4d 6s 6p 5d 4f NUCLEUS O’Connor, Davis, MacNab, McClellan, CHEMISTRY Experiments and Principles 1982, page 177

37. Atomic no. (electron config.) H1 He2 Li3 Be4 B5 C6 N7 O8 F9 Ne10 Na11 Mg12 Al13 Si14 P15