# 1. Atomic Structure and Periodic Table Details of the three Sub-atomic (fundamental) Particles

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1 1. Atomic Structure and Periodic Table Details of the three Sub-atomic (fundamental) Particles Particle Position Relative Mass Relative Charge Proton Nucleus 1 +1 Neutron Nucleus 1 Electron Orbitals 1/18-1 An atom of Lithium (Li) can be represented as follows: Isotopes Mass Number Atomic Number 7 3 Li Atomic Symbol The atomic number, Z, is the number of protons in the nucleus. The mass number,a, is the total number of protons and neutrons in the atom. Number of neutrons = A - Z There are various models for atomic structure Isotopes are atoms with the same number of protons, but different numbers of neutrons. DEFINITION: Relative isotopic mass is the mass of one atom of an isotope compared to one twelfth of the mass of one atom of carbon-12 Isotopes have similar chemical properties because they have the same electronic structure. They may have slightly varying physical properties because they have different masses. DEFINITION: Relative atomic mass is the average mass of one atom compared to one twelfth of the mass of one atom of carbon-12 DEFINITION: Relative molecular mass is the average mass of a molecule compared to one twelfth of the mass of one atom of carbon-12 THE MASS SPECTROMETER The mass spectrometer can be used to determine all the isotopes present in a sample of an element and to therefore identify elements. Calculating relative atomic mass The relative atomic mass quoted on the periodic table is a weighted average of all the isotopes % abundance % 24 Mg + 25 Mg + 26 Mg % 11.17% m/z Fig: spectra for Magnesium from mass spectrometer For each isotope the mass spectrometer can measure a m/z (mass/charge ratio)and an abundance If asked to give the species for a peak in a mass spectrum then give charge and mass number e.g. 24 Mg + 1

2 R.A.M = Σ (isotopic mass x % abundance) 1 For above example of Mg R.A.M = [(78.7 x 24) + (1.13 x 25) + (11.17 x 26)] /1 = 24.3 R.A.M = Σ (isotopic mass x relative abundance) total relative abundance Sometimes two electrons may be removed from a particle forming a 2+ ion. 24 Mg 2+ with a 2+ charge would have a m/z of 12 Use these equations to work out the R.A.M If relative abundance is used instead of percentage abundance use this equation Mass spectra for Cl 2 and Br 2 Cl has two isotopes Cl 35 (75%) and Cl 37 (25%) Br has two isotopes Br 79 (5%) and Br 81 (5%) These lead to the following spectra caused by the diatomic molecules relative abundance Cl 35 Cl 35 + Cl 35 Cl 37 + Cl 37 Cl 37 + relative abundance Br 79 Br 79 + Br 79 Br 81 + Br 81 Br 79 + Br 81 Br m/z m/z Measuring the M r of a molecule If a molecule is put through a mass spectrometer it will often break up and give a series of peaks caused by the fragments. The peak with the largest m/z, however, will be due to the complete molecule and will be equal to the M r of the molecule. This peak is called the parent ion or molecular ion Spectra for C 4 H 1 Mass spectrum for butane Molecular ion C 4 H Uses of Mass spectrometers Mass spectrometers have been included in planetary space probes so that elements on other planets can be identified. Elements on other planets can have a different composition of isotopes. Drug testing in sport to identify chemicals in the blood and to identify breakdown products from drugs in body quality control in pharmaceutical industry and to identify molecules from sample with potential biological activity radioactive dating to determine age of fossils or human remains 2

3 Ionisation Energies Definition :First ionisation energy The first ionisation energy is the energy required when one mole of gaseous atoms forms one mole of gaseous ions with a single positive charge This is represented by the equation: H (g) H + (g) + e - Definition :Second ionisation energy Always gaseous Remember these definitions very carefully The equation for 1st ionisation energy always follows the same pattern. It does not matter if the atom does not normally form a +1 ion or is not gaseous The second ionisation energy is the energy required when one mole of gaseous ions with a single positive charge forms one mole of gaseous ions with a double positive charge This is represented by the equation: Ti + (g) Ti 2+ (g)+ e - Factors that affect Ionisation energy There are three main factors 1.The attraction of the nucleus (The more protons in the nucleus the greater the attraction) 2. The distance of the electrons from the nucleus (The bigger the atom the further the outer electrons are from the nucleus and the weaker the attraction to the nucleus) 3. Shielding of the attraction of the nucleus (An electron in an outer shell is repelled by electrons in complete inner shells, weakening the attraction of the nucleus) Many questions can be answered by application of these factors Successive ionisation energies The patterns in successive ionisation energies for an element give us important information about the electronic structure for that element. Why are successive ionisation energies always larger? The second ionisation energy of an element is always bigger than the first ionisation energy. When the first electron is removed a positive ion is formed. The ion increases the attraction on the remaining electrons and so the energy required to remove the next electron is larger. How are ionisation energies linked to electronic structure? Ionisation energy No of electrons removed Notice the big jump between 4 and 5. Explanation The fifth electron is in a inner shell closer to the nucleus and therefore attracted much more strongly by the nucleus than the fourth electron. It also does not have any shielding by inner complete shells of electron Example: What group must this element be in? Ionisation energy kj mol -1 Here there is a big jump between the 2nd and 3 rd ionisations energies which means that this element must be in group 2 of the periodic table as the 3rd electron is removed from an electron shell closer to the nucleus with less shielding and so has a larger ionisation energy 3

4 The first Ionisation energy of the elements Ionisation energy kj mol Atomic number The shape of the graph for periods two and three is similar. A repeating pattern across a period is called periodicity. The pattern in the first ionisation energy gives us useful information about electronic structure You need to carefully learn the patterns Q. Why has Helium the largest first ionisation energy? A. Its first electron is in the first shell closest to the nucleus and has no shielding effects from inner shells. He has a bigger first ionisation energy than H as it has one more proton Q. Why do first ionisation energies decrease down a group? Many questions can be answered by application of the 3 factors that control ionisation energy A. As one goes down a group, the outer electrons are found in shells further from the nucleus and are more shielded so the attraction of the nucleus becomes smaller Q. Why is there a general increase in first ionisation energy across a period? A. As one goes across a period, the number of protons increases making the effective attraction of the nucleus greater. The electrons are being added to the same shell which has the same shielding effect and the electrons are pulled in closer to the nucleus. Q. Why has Na a much lower first ionisation energy than Neon? This is because Na will have its outer electron in a 3s shell further from the nucleus and is more shielded. So Na s outer electron is easier to remove and has a lower ionisation energy. Q. Why is there a small drop from Mg to Al? Al is starting to fill a 3p sub shell, whereas Mg has its outer electrons in the 3s sub shell. The electrons in the 3p subshell are slightly easier to remove because the 3p electrons are higher in energy and are also slightly shielded by the 3s electrons Q. Why is there a small drop from P to S? With sulphur there are 4 electrons in the 3p sub shell and the 4th is starting to doubly fill the first 3p orbital. When the second electron is added to a 3p orbital there is a slight repulsion between the two negatively charged electrons which makes the second electron easier to remove. 3s 3p phosphorus 1s 2 2s 2 2p 6 3s 2 3p 3 sulphur 1s 2 2s 2 2p 6 3s 2 3p 4 3s 3p Two electrons of opposite spin in the same orbital Learn carefully the explanations for these two small drops as they are different to the usual factors 4

5 Electronic Structure Models of the atom An early model of the atom was the Bohr model (GCSE model) (2 electrons in first shell, 8 in second etc.) with electrons in spherical orbits. Early models of atomic structure predicted that atoms and ions with noble gas electron arrangements should be stable. The A-level model Electrons are arranged on: Principle energy levels numbered 1,2,3,4.. 1 is closest to nucleus Split into Sub energy levels labelled s, p, d and f sholds up to 2 electrons pholds up to 6 electrons d holds up to 1 electrons fholds up to 14 electrons Principle level Sub-level 1s 2s, 2p 3s, 3p, 3d 4s, 4p, 4d, 4f An atom fills up the sub shells in order of increasing energy (note 3d is higher in energy than 4s and so gets filled after the 4s 1s2s2p3s3p 4s3d4p5s4d5p Writing electronic structure using letters and numbers For oxygen 1s 2 2s 2 2p 4 Number of electrons in sub-level Split into Orbitals which hold up to 2 electrons of opposite spin Shapes of orbitals Orbitals represent the mathematical probabilities of finding an electron at any point within certain spatial distributions around the nucleus. Each orbital has its own approximate, three dimensional shape. It is not possible to draw the shape of orbitals precisely. s sublevels are spherical Number of main energy level Using spin diagrams For fluorine 1s 2s 2p Name of type of sub-level The arrows going in the opposite direction represents the different spins of the electrons in the orbital An arrow is one electron Box represents one orbital The periodic table is split into blocks. A s block element is one whose outer electron is filling a s-sub shell p sublevels are shaped like dumbbells When filling up sub levels with several orbitals, fill each orbital singly before starting to pair up the electrons 2p Electronic structure for ions When a positive ion is formed electrons are lost Mg is 1s 2 2s 2 2p 6 3s 2 but Mg 2+ is 1s 2 2s 2 2p 6 When a negative ion is formed electrons are gained O is 1s 2 2s 2 2p 4 but O 2- is 1s 2 2s 2 2p 6 5

6 PERIODICITY Classification of elements in s, p, d blocks Elements are classified as s, p or d block, according to which orbitals the highest energy electrons are in. Period 2 = Li, Be, B, C, N, O, F, Ne Period 3 = Na, Mg, Al, Si, P, S, Cl, Ar Atomic radius Atomic radii decrease as you move from left to right across a period, because the increased number of protons create more positive charge attraction for electrons which are in the same shell with similar shielding. Exactly the same trend in period 2 atomic radius (nm) Na Mg Al Si P S Cl Ar 1st ionisation energy There is a general trend across is to increase. This is due to increasing number of protons as the electrons are being added to the same shell There is a small drop between Mg + Al. Mg has its outer electrons in the 3s sub shell, whereas Al is starting to fill the 3p subshell. Al s electron is slightly easier to remove because the 3p electrons are higher in energy. There is a small drop between phosphorous and sulphur. Sulphur s outer electron is being paired up with an another electron in the same 3p orbital. When the second electron is added to an orbital there is a slight repulsion between the two negatively charged electrons which makes the second electron easier to remove. 1st ionisation energy (kj/mol) Na Mg Al Si P S Cl Ar Exactly the same trend in period 2 with drops between Be & B and N to O for same reasons- make sure change 3s and 3p to 2s and 2p in explanation! Melting and boiling points For Na, Mg, Al- Metallic bonding : strong bonding gets stronger the more electrons there are in the outer shell that are released to the sea of electrons. A smaller positive centre also makes the bonding stronger. High energy is needed to break bonds. Si is Macromolecular: many strong covalent bonds between atoms high energy needed to break covalent bonds very high mp +bp Cl 2 (g), S 8 (s), P 4 (S) - simple Molecular : weak London forces between molecules, so little energy is needed to break them low mp+ bp S 8 has a higher mp than P 4 because it has more electrons (S 8 =128)(P 4 =6) so has stronger London forces between molecules Ar is monoatomic weak London Forces between atoms Melting and boiling points (K) Na Mg Al Si P S Cl Ar Similar trend in period 2 Li,Be metallic bonding (high mp) B,C macromolecular (very high mp) N 2,O 2 molecular (gases! Low mp as small London Forces) Ne monoatomic gas (very low mp)

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