ÐÏà¡±á>þÿ =?þÿÿÿ:;<ÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿÿì¥Áq` ø¿_LbjbjqPqP 4š::\Dÿÿÿÿÿÿ¤nnnàÄ 4444¤°4Œ¤4YŽH5Ø 7 7 7 7ÿ7>=8Q8³XµXµXµXµXµXµX$Â[h*^zÙXn]8û7ÿ7]8]8ÙX 7 7ÛîXK?K?K?]8j8 7n 7³XK?]8³XK?K?â%T¤VnõT 7<5Oº,#É4Ç;öÉT·VüY04YßT¤^½=ô¤^,õTõTÚ¤^nÏUè]8]8K?]8]8]8]8]8ÙXÙX±>š]8]8]84Y]8]8]8]8¤¤¤Äh'¤¤¤¤h'‚ÔV†Üÿÿÿÿ1.5 Atomic structure and periodic table 1.5a Relative massesThe unit of mass is 1/12 of the mass of an atom of the isotope Carbon-12 (12C=12 exactly) The Relative Atomic Mass (Ar)of an element is the ratio of the mass of an average atom of that element to 1/12 of the mass of an atom of the nuclide Carbon-12. The Relative Isotopic Mass of a nuclide is the ratio of the mass of one atom of that nuclide to 1/12 of the mass of a Carbon-12 atom. The Relative Molecular Mass (Mr)of a substance is the ratio of the mass of an average molecule of that substance to 1/12 of the mass of an atom of the nuclide Carbon-12. Task1.5a Explain what is meant by The relative atomic mass of helium is 4.002602. The relative isopotic mass of deuterium is 2.0000. The relative molecular mass of water is 18. 1.5b mass spectrometer INCLUDEPICTURE "http://www.chemguide.co.uk/analysis/masspec/masspec.GIF" \* MERGEFORMATINET (source: HYPERLINK "http://www.chemguide.co.uk/analysis/masspec/howitworks.html" http://www.chemguide.co.uk/analysis/masspec/howitworks.html) Task 1.5b.1 Write a sentence for each stage in the formation of a mass spectrum. Isotopes are atoms of the same element with different numbers of neutrons in the nucleusThe Relative Atomic Mass of an element is the weighted (to take account of relative abundance) average of the the Relative Isotopic Masses of all of the isotopes of that element.E.g. Chlorine has two isotopes with mass numbers (and relative isotopic masses) 35 and 37 35 37 75% is Cl and 25% is Cl 17 17Let there be 100 atomsTotal mass of 100 atoms = (75 * 35) + (25 * 37) = 3550Average mass of an atom (relative atomic mass of chlorine) = Total mass /Number of atoms=3550/100so relative atomic mass of chlorine = 35.5 Task 1.5b.2 1. Bromine exists as two isotopes with mass numbers 81 and 79. If there is 50% of each isotope in a sample what is the relative atomic mass of bromine? 2. Calculate the relative atomic mass of the following:(a) gallium 60% Ga 69 and 40% Ga 71(b) neon 90% Ne 20 and 10% Ne 22(c) silver 50% Ag 107 and 50% Ag 109(d) boron 20% B 10 and 80% B 11(e) neon 90.92% Ne 20, 8.82% Ne 22 and 0.26% Ne 21.(f) lead 52.3% Pb 208, 22.6% Pb 207, 23.6% Pb 206 and 1.5% Pb 204. Answers(1) 80, (2a) 69.8, (b) 20.2, (c) 108, (d) 10.8, (e) 20.18 The mass spectrum of an element shows the relative isotopic mass (m/e) and relative abundance of each isotope of the element being tested. In a mass spectrum the height of each peak = the relative abundance. The Relative Atomic Mass of an element can be found by finding the sum of the products of the relative abundance of each isotope and its relative isotopic mass and the dividing by the total relative abundance. INCLUDEPICTURE "http://www.drbateman.net/asa2sums/sum1.1/sum1.11.gif" \* MERGEFORMATINET For 100 neon atoms the total mass is (90.5*20)+(0.3*21)+(9.2*22) = 2018.7relative atomic mass of neon 2018.7/100 = 20.2Very accurate masses can be read from the spectrum if needed e.g. 20.994 for neon-21. The peak in the spectrum on the far right has the highest value of m/e and is called the molecular ion. This peak gives the Relative Molecular Mass of a compound. Below ethanol can be seen to have a relative molecular mass of 46. INCLUDEPICTURE "http://www.drbateman.net/asa2sums/sum1.1/sum1.12.gif" \* MERGEFORMATINET 1.5c Uses of mass spectrometers The mass spectrometer can detect the abundance of isotopes such as U-238 and Pb-206. In any one place the Earth concentrates some minerals like uranium but not others like lead. So some rocks form containing uranium but not lead. U-238 is radioactive and decays to form Pb-206 over time. This process is controlled by the half life of U-238 which is 4.5*109 years. The relative abundance of U-238 to Pb-206 can be used to calculate the age of the rock. Rock is found to have ages up to about 4 thousand million years, but not older. This is the age of the Earth when the first solid rock formed from the magma on the surface of the cooling planet. Radiocarbon dating using C-14 can be used to similarly date organic matter. Space research can use mass spectrometers to analyse samples on the surface of a planet HYPERLINK "http://www.sciencedirect.com/science?_ob=ArticleURL&_udi=B6V3S-4MBCBPV-1&_user=10&_coverDate=12%2F31%2F2007&_rdoc=1&_fmt=high&_orig=browse&_sort=d&view=c&_acct=C000050221&_version=1&_urlVersion=0&_userid=10&md5=6816b27a8554265411e3759c51c12c92" example1. HYPERLINK "http://cat.inist.fr/?aModele=afficheN&cpsidt=1048652" example2. Mass spectrometers can be used to detect anabolic steroids in urine samples provided by athletes. HYPERLINK "http://www.informaworld.com/smpp/content~content=a794022113~db=all~jumptype=rss" Example study. The synthesis of new drugs is far removed from what happens in a school laboratory. Hundreds of likely compounds at once can be made, identified using mass spectrometers and tested for effectiveness. A flavour of this process can be found at HYPERLINK "http://www.sciencemag.org/products/life_031706.dtl" Life science technologies 1.5d ionization energies The first ionisation energy is the enthalpy change when 1 mole of gaseous atoms of an element each lose an electron to form gaseous ions each with a single positive charge. It is the enthalpy change for the reaction: M(g) ---> M+(g) + e-The second ionisation energy is the enthalpy change when 1 mole of gaseous ions of an element each with a single positive charge each lose an electron to form gaseous ions each with a double positive charge. It is the enthalpy change for the reaction: M+(g)---> M2+(g) + e- All of these ionizations involve the removal of a negative electron from a positive nucleus. The attraction between these must be over come by supplying energy. The process is always endothermic. Example 1st IE of Na = +500kJmol-1 Task1.5d.1 Define the third ionisation enthalpy for the element M and give an equation.Task1.5d.2 Define the first ionisation enthalpy for sodium, magnesium, chlorine and neon.Task1.5d.3 Define the second ionisation enthalpy for lithium, aluminium and oxygen. 1.5e Electronic structure The graph of successive ionisation energies against ionisation number shows electrons grouped into three energy levels or quantum shells with similar energies. Electron 1 (ionisation number = 1, removed first) from the third shell, electrons 2-9 in second shell, electrons 10 and 11(removed last) from first shell closest to the nucleus. Electron 1 is on its own in the outer shell indicating that this element is in group 1. INCLUDEPICTURE "../../asa2sums/sum1.1/sum1.13.gif" \* MERGEFORMAT Task 1.5e.1Sketch similar graphs for carbon, magnesium and potassium (atomic numbers 6, 12 and 19) The graph of first ionisation energy against atomic number shows the grouping of electrons into s, p (and d) subshells (orbitals) within the (quantum shells) energy levels. The general trend is for ionisation energy to increase with increasing atomic number across a period but 1st IE B < 1st IE Be and 1st IE O < 1st IE N. INCLUDEPICTURE "../../asa2sums/sum1.1/sum1.14.gif" \* MERGEFORMAT Task1.5e.2 Label the diagram with symbols for the elements.Task1.5e.3 Label electrons removed as s or p.Task1.5e.4 Write two sentences using the frame: The ionisation energy (increases/decreases) (from left to right across a period/down a group) because (distance from outer electrons to nucleus is increasing/the nuclear charge in increasing) 1.5f Electron density maps X-ray diffraction can be used to form images which show areas of high and low electron density in compounds. It is possible to describe electrons as waves and the calculate, in simple cases, like a hydrogen atom, where high and low electron density should be found in an atom. These places match up to s and p sub=shells or orbitals. Cross-sections of an s-subshell orbital A spherical s subshell orbital with high (very blue) and decreasing (less blue) electron density INCLUDEPICTURE "http://faculty.colostate-pueblo.edu/linda.wilkes/111/3c.10.jpg" \* MERGEFORMATINET INCLUDEPICTURE "http://faculty.colostate-pueblo.edu/linda.wilkes/111/3c.11.jpg" \* MERGEFORMATINET All s subshell orbitals are spherical, the main difference between different major shells is in the size of the sphere INCLUDEPICTURE "http://faculty.colostate-pueblo.edu/linda.wilkes/111/3c.12.jpg" \* MERGEFORMATINET electron density model p subshell orbitals (source: HYPERLINK "http://faculty.colostate-pueblo.edu/linda.wilkes/111/3c.htm" http://faculty.colostate-pueblo.edu/linda.wilkes/111/3c.htm ) 1.5g Electronic configurations from hydrogen to kryptonElectronic Configuration can be predicted. If the following rules are followed:Electrons go into lower energy levels before higher ones. Electrons go into lower subshells before higher ones. Electrons occupy orbitals with 1 electron rather than 2 if possible.Electrons can only occupy the same orbital if they have opposite spins.When filling d-orbitals electrons create 3d5 and 3d10 by losing a 4s electron as half filled or filled d subshells are more stable than other arrangements. An orbital is represented by a box The orbital order 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p Max electrons in shell s=2, p=6( 3 orbitals/squares), d=10( 5 orbitals/squares) e.g Kr (the biggest you need) 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p6Sc1s2 2s2 2p6 3s2 3p6 3d1 4s2Cr 1s2 2s2 2p6 3s2 3p6 3d5 4s1 1s2s2p3s3p3d4s|||||| || |||||| || ||| | | | || HYPERLINK "buildatom.ppt" Build atom Task 1.5g.1 Draw s,p,d electronic configurations for Be, F, Mg, S and K.Task 1.5g.2 Draw electrons in boxes diagrams for N, Ne, Na, Ar and Ca.Task 1.5g.3 Draw s,p,d electronic configurations for Sc, V, Mn, Ni, Zn and Cu.Task 1.5g.4 Draw electrons in boxes diagrams for Ti, Fe, Ni, Se, Kr and CrTask 1.5g.5 Draw sketches of a 2s and a 3p subshell orbital. 1.1h Electronic structure and chemical propertiesAtoms of group 1 elements easily lose 1 electron to form stable ions. Atoms of group 2 elements easily lose 2 electrons to form stable positive ions. All of these elements are therefore strongly metallic. Group 6 and 7 elements are strongly non-metallic. They may gain 1 or 2 electrons to form negative ions with a stable inert gas configuration. Electron sharing is also common. Group 0 have atoms with complete outer shells therefore they are chemically inert. d-block elements have atoms with 1 or two outer electrons. These are easily lost and therefore these elements have metallic properties.Task 1.1h Explain why sodium easily reacts with HCl to form a salt,iron also reacts with HCl to give a salt,chlorine however does not react with acids. 1.1i s,p and d-block elementsINCLUDEPICTURE "../../asa2sums/sum1.1/sum1.15.gif" \* MERGEFORMAT All s-block elements have their outer electrons in s-orbitals.All p-block elements have their outer electrons in p-orbitals.All d-block elements have electrons in their d-orbitals which are in the process of being filled. Task 1.1i Draw electrons in boxes for any atom in each of the 3 periodic table blocks above 1.1j Periodic propertyThe value of a property, for example melting point, changes from element to element. If the property is plotted on a graph against atomic number then a pattern is revealed. Typically there are peaks and troughs at regular places. For example the melting points of group 4 elements are high but for group 0 the values are low. This repeating pattern is said to be periodic. Melting point is a periodic property. INCLUDEPICTURE "http://www.webelements.com/_media/periodicity/tables/line/melting_point.gif" \* MERGEFORMATINET HYPERLINK "http://www.webelements.com/periodicity/" http://www.webelements.com/periodicity/ Task 1.1j Use data tables to draw graphs of a property against atomic number and describe the periodic nature of the property. 1.5k Explaining periodic trends LiBeBCNOFNeNaMgAlSiPSClArGiant metallic structure, metallic bondingGiant atomic structure, covalent bondingSimple molecular structure, covalent bonding and Van der Waals forces Simple molecular structures. substancedensity/gcm-3melting point/Kelectrical conductivityO2 1.15 at 90K55very lowH2 0.07 at 20K14very lowS2 2.0390very lowI2 4.93387very lowStrong covalent bonds between atoms in molecules. Weak Van der Waals’ forces between molecules. Low melting points and boiling points because of weak forces between molecules. They require very little energy to be broken; therefore we have melting points and boiling points at low temperatures. Brittle (elements) because the weak forces between the molecules are unable to withstand a large external force. Electrical insulators because there are no charged particles to carry the charge. Often insoluble in water, especially the non-polar molecules. There is little interaction between polar water molecules and non-polar molecules. The exceptions are polar molecules. Low density because molecules are not pulled strongly together in solids or liquids and many are gases in which particles are far apart. HYPERLINK "http://www.drbateman.net/asa2sums/sum1.3/task1.3e.1.htm" Task 1.3e.1 Ionic structures. INCLUDEPICTURE "http://www.drbateman.net/asa2sums/sum1.3/image17.gif" \* MERGEFORMATINET Caesium chloride INCLUDEPICTURE "http://www.drbateman.net/asa2sums/sum1.3/image18.gif" \* MERGEFORMATINET Sodium Chloride 6:6 co-ordination: each chloride ion has 6 sodium ions as it’s nearest neighbours. Each sodium ion is also surrounded by six chloride ions as it’s nearest neighbours. The electrostatic forces holding the ions in place are not directional. mp = 1081K, density = 2.17gcm-3 All ionic compounds: Have high melting points and boiling points and are solids at room temperature because each ion is held firmly in place by strong ionic electrostatic forces. Have densities higher than water but much lower than typical metals. Although efficiently packed there is some empty space between ions. Are often soluble in water because polar H2O molecules are attracted to ions and the attractive forces of many H2O molecules can pull ions away from their crystal structures. HYPERLINK "http://www.drbateman.net/asa2sums/sum1.3/wpe1.jpg" INCLUDEPICTURE "http://www.drbateman.net/asa2sums/sum1.3/sum1.31.jpg" \* MERGEFORMATINET Do not conduct electricity while solid as no particles are free to carry the charge. ‚ Conduct electricity when molten or in a solution as the ions are free to carry the charge. Giant atomic structures. INCLUDEPICTURE "http://www.drbateman.net/asa2sums/sum1.3/image21.gif" \* MERGEFORMATINET The structure of Diamond. It is hard and has a very high melting point and boiling point because each atom is held firmly in place by 4 strong, short, covalent bonds and a lot of energy is required to break these strong bonds. Doesn’t conduct electricity even when molten as no charged particles to carry charge. Insoluble in water as forces between solvent and carbon atoms are too weak. Thermal conductor as rigid structure allows heat to be passed through vibrations. High density (3.51gcm-3) because atoms are packed tightly together. HYPERLINK "http://www.drbateman.net/asa2sums/sum1.3/task1.3e.htm" Task 1.3e.2 INCLUDEPICTURE "http://www.drbateman.net/asa2sums/sum1.3/image22.gif" \* MERGEFORMATINET The structure of graphite. Graphite is soft as there are weak bonds between layers, thus allowing layers to slide over each other. (Large distances between layers imply weak bonds.) ‚ Graphite has a high melting point and boiling point as 3 strong covalent bonds hold each atom in place. Graphite conducts heat and electricity in one direction due to delocalised electrons between the layers.Low density (2.25gcm-3) because the layers are far apart. Task 1.5k.1 Explain the difference in melting points for any two consecutive elements in periods 2 or 3: (a) Mg and Al, (b) C and N, (c) Cl and Ar, (d) Si and P, (e) Al and Si 1st ionisation energy increases across a period from left to right as there is an increasing nuclear charge and so an increasing pull of the nucleus on an outer electron. 1st ionisation energy decreases down a group as outer electrons are in quantum shells further from the nucleus and the greater distance reduces the pull of the nucleus for an outer electron. Shielding increases down a group and this has a further small effect on decreasing 1st ionisation energy. 1st IE B < 1st IE Be because Be completes the 2s sub-shell. B has 1 electron in the 2p subshell. The 2p electron is further from the nucleus than a 2s electron so although the nuclear charge increases from Be to B the 2p electron is easier to remove than the 2 s electron in Be. 1st IE of O < 1st IE of N because the p subshell for N has singly occupied orbitals. For O an electron is being added to an occupied orbital. The repulsion involved is enough to outweigh the increase in nuclear charge. Task 1.5k.2 Explain the differences in 1st IE between: (a) H and He, (b) He and Li, (c) F and Ne, (d) Na and Mg, (e) Mg and Al, and (f) P and S. 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