2. Atomic Structure HL and SL

3. 2.1 The Atom Atoms were thought to be uniform spheres like snooker balls. Experiments, however, have shown that atoms consist of a small nucleus, which contains particles called protons and neutrons, around which electrons orbit in energy levels.

4. 1+ 1 10 - 15 x 10 -4 In the nucleus Energy levels around nucleus Atoms are NEUTRAL, this means that they have an equal number of protons and electrons. The table below shows the relative masses and charges of the three sub-atomic particles: You can use Avogadro’s constant to determine the mass of a proton and an electron. The charge on an electron is 1.60 x 10 -19 C.

5. Important Definitions: The ATOMIC NUMBER is the number of protons in an atom The MASS NUMBER is the number of protons plus the number of neutrons in an atom. ISOTOPES are atoms of the same element with the same number of protons and electrons but differing numbers of neutrons.

6. Shorthand: X symbol of element A Z mass number atomic number

7. So: sodium atom Na 23 11 mass number atomic number

8. Na 23 11 This sodium atom has 11 protons 11 electrons 12 neutrons NUMBER OF NEUTRONS = MASS NUMBER – ATOMIC NUMBER

9. More on ISOTOPES: As the chemical properties of an atom depend upon their electron arrangements and isotopes have the same electron arrangement (as they are atoms of the same element), isotopes have the same chemical properties. However, their physical properties especially those that depend upon mass may vary. These include density and boiling point. Using isotopes: Find out about the use of 14 C in radiocarbon dating, 60 Co in radiotherapy, 131 I and 125 I as medical tracers. Do the benefits outweigh the dangers?

10. Ions are charged particles formed when atoms either lose or gain electrons. Metal atoms tend to lose electrons to form positive ions. e.g. a sodium ion (Na + ) is formed when a sodium atom loses one electron. A magnesium atom (Mg 2+ ) is formed when a magnesium atom loses 2 electrons. Non-metal atoms tend to gain electrons to form negative ions. e.g. a chlorine atom gains an electron to form a chloride ion (Cl - ) An oxygen atom gains two electrons to from an oxide ion (O 2- )

11. 2.2 Mass Spectrometry Measures mass of atoms and their relative abundance enabling relative atomic mass values to be calculated. 1. electron gun 2. electric field 3. strong electromagnet 4. detector vapourised sample

12. Mass spectrometry

13. Ionisation The vapourised sample (atoms or molecules) diffuses into the path of high energy electrons fired from the electron gun. These electrons knock out an electron from the sample producing positive ions. M (g) + e -  M + (g) + 2e - Some doubly charged ions (M 2+ (g) ) are also formed but in small amounts as it requires more energy to knock out 2 electrons. Molecules can be broken into ‘fragments’ by the high energy electrons as they break covalent bonds.

14. Acceleration The positive ions are accelerated by an electric field and then focussed into a beam by passing them through a series of slits.

15. Deflection The beam of fast moving positive ions is then deflected by a strong magnetic field. The magnitude of the deflection depends upon the mass to charge ratio (m/z) of the ion. When the m/z is small, the deflection is large. The magnetic field can be increased in order to deflect heavier ions into the detector. Explain why 40 Ca 2+ and 20 Ne + are deflected by the same amount.

16. Detection The ions are detected electrically. The ions hit a plate and this sends a current to an amplifier and then to a recorder. The chart produced by the recorder is called a ‘mass spectrum’. The spectrum on the next slide is for a sample of lead.

17. This shows that there are 4 isotopes of lead with mass numbers of 204, 206, 207 and 208. The height of the peak gives a measure of the relative abundance of that peak. This information can be used to calculate the A r of lead.

18. Using Mass Spectra to calculate A r values The A r is given by calculating the weighted mean of the individual relative atomic masses of the isotopes relative to 1/12 th of carbon-12. To do this you need to find the total mass of all of the isotopes and then divide this by the total abundance. The total mass of all the isotopes is found by adding together the mass present of each isotope in the sample which is determined by multiplying the mass of the isotope by its abundance.

19. A r = total mass / total abundance Total mass = (204 x 1.5) + (206 x 23.6) + (207 x 22.6) + (208 x 52.3) = 20724.2 Total abundance = 1.5 + 23.6 + 22.6 + 52.3 = 100 A r = total mass / total abundance = 20724.2 / 100 = 207.2

20. Time to use brainpower: Sketch the spectrum that would be obtained for a molecule of chlorine. Given that chlorine exists as 2 isotopes, one with mass number 35 and the other with mass number 37. These two isotopes occur in the ratio 3 : 1.

21. 2.3 Electron Arrangements Why are street lights orange? Why does copper give a green colour to a bunsen flame when it is being heated? Due to their electron arrangements! The different colours are electromagnetic radiation. Each colour has a particular wavelength and frequency and is associated with a particular amount of energy. To discover the cause of these colours we need to consider the electromagnetic spectrum.

22. The velocity of travel of electromagnetic waves is related to its wavelength and frequency by the equation: c = f c = velocity in ms -1, = wavelength in m, f = frequency in s -1 The energy of electromagnetic radiation is related to its frequency by the equation: E = hf E = energy in J, h is Planck’s constant (6.63 x 10 -34 Js) The smaller the wavelength, the higher the frequency and the more energy the wave possesses. Electromagnetic waves have a wide range of wavelengths varying from radio waves (10 3 m) to gamma radiation (10 -12 m).

23. The electromagnetic spectrum is a continuous spectrum. You can get another version of this diagram at www.chemsoc.org/Networks/Learnnet/data/ds_electromagnetic_spectrum.htm

24. If the orange light emitted by sodium is passed through a prism, it is seen as a series of lines, each at a fixed wavelength. Each element has its own characteristic set of lines. These lines are known as the line emission spectrum of the element. These lines become closer together (converge) at the high energy end of the spectrum. The simplest spectrum to consider is the hydrogen spectrum: This is the series of lines found in the visible region. There is another series in the uv region and several in the ir region.

25. How do these lines arise? Niels Bohr published a model in 1913 to explain. Electrons travel in orbits around the nucleus of the atom, each orbit is in a fixed energy level. If the electron is given energy it is promoted to a higher energy level. As it drops back down it emits a packet of light called a quantum with a particular amount of energy. This energy corresponds to light of a particular wavelength and shows up as one of the lines in the spectrum.

26. The spectrum is not continuous as the electrons only exist at certain fixed energy levels. Electrons dropping back down to the lowest energy (n = 1) emit most energy so this produces the series in the uv region of the hydrogen spectrum. Electrons dropping down to the third energy level (n = 3) cause the first series in the ir region. The visible spectrum arises due to electrons falling to the second energy level (n = 2). The lines get closer as the energy levels themselves get closer. The value n is known as the principal quantum number of the energy level.

27. Look at www.chemsoc.org/Networks/Learnnet/data/int_electron_energy_hydrogen.htm

28. Each of the energy levels described by the principal quantum numbers can only hold a certain number of electrons. The lowest fills first. When one energy level is full with electrons the next then begins to fill. Principal quantum number Maximum number of electrons 12 28 38 (18 for HL) The electron arrangements of the first 20 elements can be found on the department web pages.