Atomic Structure and Periodic Table PDF - IGCSE

2. Atomic Structure and Periodic Table Details of the three Sub-atomic (fundamental) Particles Particle Position Relative Mass Relative Charge Proton Nucleus 1 +1 Neutron...

2. 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 0 There are various Electron Orbitals 1/1840 -1 models for atomic structure An atom of Lithium (Li) can be represented as follows: Mass Number 7 Li Atomic Symbol Atomic Number 3 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 Isotopes 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 weighted mean 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 Fig: spectra for 100 Magnesium from mass spectrometer 80 78.70% For each isotope the mass % abundance 60 24Mg+ spectrometer can measure a m/z 40 (mass/charge ratio) and an abundance 25Mg+ 26Mg+ 20 10.13% If asked to give the species for a peak 11.17% in a mass spectrum then give charge and mass number e.g. 24Mg+ m/z 24 25 26 N Goalby chemrevise.org 1  Sometimes two electrons may be R.A.M =  (isotopic mass x % abundance) removed from a particle forming a 2+ ion. 24Mg2+ with a 2+ charge would 100 have a m/z of 12 For above example of Mg R.A.M = [(78.7 x 24) + (10.13 x 25) + (11.17 x 26)] /100 = 24.3 Use these equations to work out the R.A.M R.A.M =  (isotopic mass x relative abundance) If relative abundance is used instead of total relative abundance percentage abundance use this equation Mass spectra for Cl2 and Br2 Cl has two isotopes Cl35 (75%) and Cl37(25%) Br has two isotopes Br79 (50%) and Br81(50%) These lead to the following spectra caused by the diatomic molecules Br79Br81 + Br81Br79 + Cl35Cl35 + relative relative abundance abundance Cl35Cl37 + Br79Br79 + Br81Br81 + Cl37Cl37 + 70 72 74 m/z m/z 158 160 162 Measuring the Mr of a molecule Spectra for C4H10 If a molecule is put through a mass spectrometer it will often break up and give a series of peaks caused Mass spectrum for butane by the fragments. The peak with the largest m/z, however, will be due to the complete molecule and 43 will be equal to the Mr of the molecule. This peak is Molecular ion called the parent ion or molecular ion C4H10+ 29 58 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 N Goalby chemrevise.org 2  Ionisation Energies Definition : First ionisation energy Remember these definitions very carefully 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 The equation for 1st ionisation This is represented by the equation: H(g)  H+ (g) + e- energy always follows the same pattern. It does not matter if the atom does Always gaseous not normally form a +1 ion or is not gaseous Definition : Second ionisation energy 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)  Ti2+(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) Many questions can be 2. The distance of the electrons from the nucleus answered by application (The bigger the atom the further the outer electrons are from the nucleus and the of these factors 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) 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? Explanation Ionisation The fifth electron is in a inner energy shell closer to the nucleus and therefore attracted much more Notice the big strongly by the nucleus than the jump between 4 fourth electron. and 5. It also does not have any shielding by inner complete shells 1 2 3 4 5 6 of electron No of electrons removed Example: What group must this element be in? Here there is a big jump between the 2nd and 3rd ionisations energies which means that this 1 2 3 4 5 element must be in group 2 of the periodic table Ionisation 590 1150 4940 6480 8120 as the 3rd electron is removed from an electron energy kJ mol-1 shell closer to the nucleus with less shielding and so has a larger ionisation energy N Goalby chemrevise.org 3 The first Ionisation energy of the elements The shape of the graph for periods two and three is similar. A repeating pattern across a Ionisation energy kJ mol-1 period is called periodicity. 2000 1500 The pattern in the first ionisation energy gives us useful information about 1000 electronic structure 500 You need to carefully learn the 0 patterns 5 10 15 20 Atomic number 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 Many questions can be energy than H as it has one more proton answered by application of the 3 factors that control Q. Why do first ionisation energies decrease down a group? 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 Learn carefully the explanations for Q. Why is there a small drop from P to S? these two small With sulfur there are 4 electrons in the 3p sub shell and the 4th is starting to doubly fill drops as they are the first 3p orbital. different to the When the second electron is added to a 3p orbital there is a slight repulsion between usual factors the two negatively charged electrons which makes the second electron easier to remove. 3p 3p 3s 3s Two electrons of opposite spin in the same orbital phosphorus 1s2 2s2 2p63s23p3 sulfur 1s2 2s2 2p63s23p4 N Goalby chemrevise.org 4 Patterns in the second ionisation energy. If the graph of second ionisation or each successive element is plotted then a similar pattern to the first ionisation energy is observed but all the elements will have shifted one to the left. 5000 Na 2nd Ionisation energy 4500 4000 (kJ/mol) 3500 3000 Ar 2500 S P 2000 Al Cl 1500 Mg Si 1000 10 12 14 16 18 20 Atomic Number The group 1 elements are now at the peaks of the graph Lithium would now have the second largest ionisation of all elements as its second electron would be removed from the first 1s shell closest to the nucleus and has no shielding effects from inner shells. Li has a bigger second ionisation energy than He as it has more protons. N Goalby chemrevise.org 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: Sub energy levels labelled s , p, d and f Principle energy levels Split s holds up to 2 electrons Split Orbitals which hold up numbered 1,2,3,4.. into to 2 electrons of p holds up to 6 electrons into 1 is closest to nucleus opposite spin d holds up to 10 electrons f holds up to 14 electrons Shapes of orbitals Orbitals represent the Principle level 1 2 3 4 mathematical probabilities of finding an electron at any point within certain spatial Sub-level distributions around the 1s 2s, 2p 3s, 3p, 3d 4s, 4p, 4d, 4f nucleus. Each orbital has its own approximate, three An atom fills up the sub shells in order of increasing energy (note 3d is dimensional shape. higher in energy than 4s and so gets filled after the 4s It is not possible to draw the 1s2s2p3s3p 4s3d4p5s4d5p shape of orbitals precisely. Writing electronic structure using letters and numbers s sublevels are Number of electrons spherical in sub-level For oxygen 1s2 2s2 2p4 Number of main Name of energy level type of p sublevels are shaped sub-level like dumbbells Using spin diagrams For fluorine An arrow is one electron 2p Box represents one 2s orbital The arrows going in the The periodic table is split into opposite direction represents 1s the different spins of the blocks. A s block element is electrons in the orbital one whose outer electron is filling a s-sub shell 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 When a negative ion is formed electrons are gained Mg is 1s2 2s2 2p6 3s2 but Mg2+ is 1s2 2s2 2p6 O is 1s2 2s2 2p4 but O2- is 1s2 2s2 2p6 N Goalby chemrevise.org 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 0.18 0.16 atomic radius (nm) Atomic radii decrease as you move from left to right 0.14 across a period, because the increased number of 0.12 protons create more positive charge attraction for 0.1 electrons which are in the same shell with similar 0.08 shielding. 0.06 0.04 0.02 Exactly the same trend in period 2 0 Na Mg Al Si P S Cl Ar 1st ionisation energy The general trend across is to increase. This is due to 1st ionisation energy 1600 increasing number of protons as the electrons are being 1400 added to the same shell 1200 (kJ/mol) 1000 There is a small drop between Mg + Al. Mg has its outer 800 600 electrons in the 3s sub shell, whereas Al is starting to fill the 400 3p subshell. Al’s electron is slightly easier to remove 200 because the 3p electrons are higher in energy. 0 Na Mg Al Si P S Cl Ar There is a small drop between phosphorous and sulfur. Sulfur’s outer electron is paired up with an another electron in Exactly the same trend in period 2 with the same 3p orbital. drops between Be & B and N to O for When the second electron is added to an orbital there is a same reasons- make sure change 3s slight repulsion between the two negatively charged and 3p to 2s and 2p in explanation! electrons which makes the second electron easier to remove. Melting and boiling points For Na, Mg, Al- Metallic bonding : strong bonding – gets 3000 stronger the more electrons there are in the outer shell that are Melting and boiling 2500 released to the sea of electrons. A smaller sized ion with a 2000 points (K) greater positive charge also makes the bonding stronger. High 1500 energy is needed to break bonds. 1000 Si is Macromolecular: many strong covalent bonds between 500 atoms high energy needed to break covalent bonds– very high 0 mp +bp Na Mg Al Si P S Cl Ar Cl2 (g), S8 (s), P4 (S)- simple Molecular : weak London forces between molecules, so little energy is needed to break them – Similar trend in period 2 low mp+ bp Li,Be metallic bonding (high mp) S8 has a higher mp than P4 because it has more electrons (S8 B,C macromolecular (very high mp) =128)(P4=60) so has stronger London forces between N2,O2 molecular (gases! Low mp as molecules small London Forces) Ne monoatomic gas (very low mp) Ar is monoatomic weak London Forces between atoms N Goalby chemrevise.org