1. Atomic Structure

2. Fundamental Particles Knowledge and understanding of atomic structure has evolved over time. Atoms are made from three types of particles: protons, neutrons and electrons. Small and extremely dense central nucleus contains protons and neutrons.

3. The absolute masses and charges of protons, neutron and electrons are so small that we usually compare them to each other. The numbers are called relative mass and relative charge ParticleRelative massRelative charge Proton1+1 Neutron10 Electron1/1836

4. Mass Number and Isotopes The atomic number of an atom is the number of protons in the nucleus. The atomic number is also called the proton number and it is given the symbol Z The mass number is the number of protons and neutrons in the nucleus of an atom. The mass number is given the symbol A X Z A Mass number Atomic number Element symbol Mass number = No. of protons + No. of neutrons

5. Atoms have no charge and are neutral. This means atoms must have an equal number of protons and electrons. Therefore the number of electrons is same as the number of protons (or atomic number) for atoms. You must be able to work out the number of protons, neutrons and electrons for atoms and ions in the exam. You should be able to complete the table below using the Periodic Table. In the table below, you can use the relative atomic mass in the periodic table as the mass number. This is correct for the most abundant isotope. AtomProtonsNeutronsElectrons Atomic number Mass number Boron565 Potassium1920 Chromium24 52 Mercury 101 Argon

6. FluorineUnknown Proton918 Neutron1020 Mass number1938

7. Isotopes Definition: Atoms of the same element with the same number of protons but different number of neutrons Example: Isotopes of carbon Carbon-12Carbon-13Carbon-14 Mass number121314 No. Protons666 No. Neutrons678 No. Electrons666 Properties of isotopes Chemical properties are dictated by the number of electrons. Chemical properties of isotopes are identical, because isotopes have the same number of electrons. Isotopes have different physical properties such as rates of diffusion and boiling point. This is because isotopes have different mass numbers.

8. Electron Configuration Main Energy Levels In the modern atomic model of the atom, electrons are arranged in energy levels (shells). Each main energy level is given a number (1, 2, 3 or 4). Energy level 1 is closest to the nucleus. The energy of each level increases the further away it is from the nucleus. An electron in isolation has zero energy. The attraction to the positive nucleus make an electron in an atom more stable, so the energy levels in an atom have negative energies.

9. Sub-Levels Each energy level contains sub-levels. Each sub-level is given a letter (s, p, d or f) The number of sub-levels depends on the number of the main energy level: the first has 1 sub-level, the second 2, the third 3 etc. Each sub-level (s, p, d or f) can hold a maximum number of electrons: s sub-level can hold 2 electrons p sub-level can hold 6 electrons d sub-level can hold 10 electrons The energy of the sub-levels increases from s to p to d

10. Orbitals Each sub-level consists of a different number of orbitals. s sub-levels have 1 orbital p sub-levels have 3 orbitals d sub-levels have 5 orbitals f sub-levels have 7 orbitals An orbital is the volume of space where the electron(s) are located. Each orbital can hold a maximum of two electrons. The different types of orbital have different shapes.

11. Spin Electrons have a property called spin, which is in one of two directions. These are usually called ‘spin up’ and ‘spin down’, often denoted with arrows ↑ or ↓. When two electrons occupy the same orbital, there spins will always be in opposite directions. Recap: The four properties (quantum numbers) that describe an electron in an atom are: Main energy level (shell) Sub-level (sub-shell) Orbital Spin

12. Energy level diagram The main energy level and sub-level of an electron affect its energy. The diagram shows the relative energies of the sublevels from 1s to 4d for a hydrogen atom. The exact energies will vary slightly, but this diagram is a good guide for all most atoms and ions.

13. Electrons will fill the orbitals from the lowest energy upwards. In sub-levels with more than one orbital (p, d, f), the electrons go one into each orbital with spins parallel before they start pairing up.

14. This diagram can help to remember the order of filling sub-levels: 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 5g 6s 1s 2, 2s 2, 2p 6, 3s 2, 3p 6, 4s 2, 3d 10, 4p 6, 5s 2, 4d 10, 5p 6, 6s 2 N.B. The 4s sub-level fills before the 3d. The blocks in the periodic table help to determine the number of electrons in each sub-level.

15. Electron configurations (arrangements) These are lists of the number of electrons in each of the occupied sub-levels. They give: The number of the main energy level The letter of the sub-level (always lower case) The number of electrons (as a superscript) Examples 1) Sodium (11 electrons)1s 2,2s 2,2p 6,3s 1 2) Phosphorus (15 electrons)1s 2,2s 2,2p 6,3s 2,3p 3

16. We can draw diagrams to show electron arrangements that give the detail of orbitals and spins of electrons. Example: Sodium 1s 2 2s 2 2p 6 3s 1

17. Elements with Z above 20 With these we have to remember to fill 4s before 3d Examles: Vanadium: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 3 Zinc 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 Krypton 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6

18. Shorthand: Often the inner electron arrangements are abbreviated by giving the symbol of the noble gas that has the same electron configuration as the inner energy levels. Examples The electron configuration for Zn: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 Becomes: [Ar] 4s 2 3d 10 The electron configuration for Sr: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 Becomes: [Kr] 5s 2

19. Ions An atom becomes an ion if it loses or gains electrons. An ion is a charged particle + means take 1 electron away (lost) - means add 1 electron (gained) Positive ions All electrons are lost from the highest energy level Na atom: 1s 2 2s 2 2p 6 3s 1 Na + ion (1s 2 2s 2 2p 6 ) + Al atom: 1s 2 2s 2 2p 6 3s 2 3p 1 Al 3+ ion (1s 2 2s 2 2p 6 ) 3+

20. Negative ions All electrons are gained into the highest energy level Cl atom: 1s 2 2s 2 2p 6 3s 2 3p 5 Cl - ion (1s 2 2s 2 2p 6 3s 2 3p 6 ) -

21. Unusual electron configurations There are three occasions when we come across unusual electron configurations, which all result from the 4s and 3d energy levels swapping in order. 1.Chromium atoms 2.Copper atoms 3.All transition metal ions

22. Chromium and copper atoms The swapping in order of energy for the 4s and 3d sub- levels gives: Cr[Ar] 4s 1 3d 5 Cu[Ar] 4s 1 3d 10 The swapping of s and d sub-levels also occurs in other periods of the d block, one electron before the sub-level is either filled or half-filled.

23. Transition metal ions The swapping in order of energy for the 4s and 3d sub- levels results in electrons being lost from the 4s before the 3d: Fe atom [Ar] 4s 2 3d 6 Fe 2+ ion ([Ar] 4s 0 3d 6 ) 2+ Fe 3+ ion ([Ar] 4s 0 3d 5 ) 3+ Ni atom [Ar] 4s 2 3d 8 Ni 2+ ion ([Ar] 4s 0 3d 8 ) 2+

24. Mass Spectrometer The mass spectrometer is a machine that gives you information about: Which isotopes are present in a sample and their relative amounts The A r (relative atomic mass) or M r (relative molecular mass) of the sample There are many different types of mass spectrometers and we will be covering the principles of a simple time of flight (TOF) mass spectrometer

25. What happens in a time of flight mass spectrometer? 1)Vacuum 2)Ionisation 3)Acceleration 4)Ion drift 5)Detection 6)Data analysis

26. What happens in a time of flight mass spectrometer? 1) Vacuum - The apparatus is kept in a vacuum to prevent the ions that are produced colliding with air molecules 2) Ionisation - The sample is dissolved in a volatile solvent and forced through a fine hollow needle that is connected to the positive terminal of a high voltage supply. This produces tiny positively charged droplets which have lost electrons to the positive charge of the voltage supply. The solvent evaporates from the droplets and they get smaller and smaller in size until they contain no more than a single positively charged ion.

27. 3) Acceleration - The positive ions are attracted and accelerated towards a negatively charged plate. Lighter ions and more highly charged ions achieve a higher speed 4) Ion drift - The ions pass through a hole in the negatively charged plate, forming a beam and travel along a tube, called the flight tube, to a detector

28. 5) Detection - When ions with the same charge arrive at the detector, the lighter ones are first as they travelled faster. The positive ions pick up an electron from the detector which causes a current to flow 6) Data analysis – The signal from the detector is passed to a computer which generates a mass spectrum

30. m/z242526 Relative abundance (%) 79.010.011.0

32. Mass spectrum of chlorine Chlorine has two isotopes, 35 Cl and 37 Cl which are present in the ratio 3:1. This means the peak height for 35 Cl is three times bigger than that of 37 Cl

33. Mass spectrum of chlorine Chlorine exists as a diatomic molecule, Cl 2 There are three peaks due to Cl 2 + ions with the ratio 9:6:1 35 Cl, 35 Cl m/z = 709 35 Cl, 37 Cl m/z = 726 37 Cl, 37 Cl m/z = 741

34. Mass spectra of molecules Definition: Relative molecular mass, M r is the average mass of a molecule compared to one-twelfth the mass of a 12 C atom. The peak with the highest m/z ratio is of the molecular ion. The m/z ratio of the molecular ion = the M r of the molecule.

35. Mass spectra of molecules Example: Mass spectrum of ethanol The mass spectra shows a peak for the molecular ion at m/z 46 and therefore the M r of the molecule is 46

36. Ionisation Energies Definition: The first ionisation energy is the enthalpy change for the removal of one mole of electrons from one mole of atoms of the element in the gas phase M (g) → M + (g) + e - The magnitude of the 1 st ionisation energy of an element is a measure of how strongly attracted the outer electron is to the nucleus.

37. Ionisation Energies Example: First ionisation of magnesium Mg (g) → Mg + (g) + e - is +738 kJ mol -1 Ionisation energies are endothermic and they have positive values. Electrons are negatively charged and attracted to the positively charged nucleus. Therefore energy is required to remove an electron.

38. First Ionisation energy down a group The first ionisation energy decreases as you go down a group. As you go down a group the atoms get bigger as more main energy levels of electrons are filled. The distance between the nucleus and the outer electron increases. The outer electron is also more shielded from the nucleus and is less strongly attracted. Therefore less energy is needed to remove the outer electron.

39. Group 2 (Be-Ba) The first ionisation energy decreases as you move from beryllium to barium. Magnesium has a lower first ionisation energy than beryllium. The outer electron in magnesium is in a 3s sub-level and in beryllium it is in a 2s sub-level. The outer electron in the 3s sub-level is further from the nucleus and it is more shielded from the nucleus. Therefore less energy is needed to remove it Mg (12 electrons): 1s 2 2s 2 2p 6 3s 2 Be (4 electrons): 1s 2 2s 2

40. Group 2 (Be-Ba)

41. Ionisation Energies across Period 3 (Na-Ar) Evidence for the existence of sub-levels and orbitals comes from looking at the first ionisation energies of the period 3 elements

42. There is a general increase in the first ionisation energy across a period. The nuclear charge increases (as the No. protons increases) and the shielding remains constant. The outer electron is more strongly attracted to the nucleus. Therefore more energy is required to remove the outer electron.

43. There is a fall in the first ionisation energy from magnesium to aluminium. The outer electron in aluminium is in a 3p sub-level and in magnesium it is in a 3s sub-level. The 3p sub- level is higher in energy than the 3s sub-level. Therefore less energy is needed to remove the outer electron in aluminium. This is evidence for the existence of sub-levels.

44. There is a fall in the first ionisation energy from phosphorus to sulfur. The 3p electrons in phosphorus are unpaired. In sulfur two of the 3p electrons are paired. There is some repulsion between the paired electrons in the 3p sub-level, therefore less energy is needed to remove one of these. 3s3p x 3p y 3p z phosphorus [Ne] 3s 2 3p 3     sulfur[Ne] 3s 2 3p 4    

45. There is a fall in the first ionisation energy from phosphorus to sulfur. The 3p electrons in phosphorus are unpaired. In sulfur two of the 3p electrons are paired. There is some repulsion between the paired electrons in the 3p sub-level, therefore less energy is needed to remove one of these. 3s3p x 3p y 3p z phosphorus [Ne] 3s 2 3p 3     sulfur[Ne] 3s 2 3p 4     This is evidence for the existence of orbitals.

46. Successive ionisation energies The equation for the first ionisation energy is: M (g) → M + (g) + e - The equation for the second ionisation energy is: M + (g) → M 2+ (g) + e - The equation for the third ionisation energy is: M 2+ (g) → M 3+ (g) + e - These are the enthalpy changes for the removal of the 1 st, 2 nd and 3 rd outermost electrons respectively.

47. Successive ionisation energies Example: Sodium The second ionisation energy is larger than the first. This is because the second electron is removed from a positive charged ion. IonisationEquationEnergy (kJ mol -1 ) FirstNa(g) → Na + (g) + e - 496 SecondNa + (g) → Na 2+ (g) + e - 4563 ThirdNa 2+ (g) → Na 3+ (g) + e - 6913

48. Successive ionisation energies Successive ionisation energies tell you the number of electrons in the outer shell (energy level) of an element. There is a big jump in the ionisation energy when the electron is being removed from a new shell (energy level). This is a result of the large increase in the attraction of the outer electron in the next energy level. It has one less inner energy level shielding the charge of the nucleus.

49. Example: Aluminium There is a big jump after the third electron is removed. This means there are 3 outer electrons and it belongs to group 3. The fourth electron is removed from a new shell (energy level).

50. Example: There is a big jump after the second electron is removed. This means there are 2 outer electrons and the element belongs to group 2. IonisationEnergy (kJ mol -1 ) 1 st 738 2 nd 1451 3 rd 7733 4 th 10,543 5 th 13,630 6 th 18,020