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http://www.youtube.com/watch?v=nsbXp64YPRQ Topic 3: Periodicity ● ● Chemistry is not study of a random collection of elements. Periodicity = repeating patterns of chemical and physical properties. 3.1 The periodic table ● ● ● The development of the periodic table took many years and involved scientists from several countries. Mendeleev grouped the known elements into families on the basis of their relative atomic masses and their chemical properties. Dimitri Mendeleev (1834-1907) He left gaps where no known elements fitted in and predicted the physical and chemical properties of these missing elements. http://www.youtube.com/watch?v=157KVUQXRJA ● In the modern periodic table the elements are arranged in order of increasing atomic number , Z. http://education.jlab.org/itselemental/ http://www.rsc.org/periodic-table/ Groups ● ● ● Elements with similar chemical and physical properties are placed underneath each other in vertical columns called groups. The groups are numbered 1-18. The number of valence electrons of the elements in group 1-2 and 13-18 can be found from the group number . ● Certain groups have their own names: - alkali metals - alkaline earth metals - pnictogens - chalcogens - halogens - noble gases Periods ● The period number is equal to the principal quantum number, n, of the highest occupied energy level in the elements of the period. Metals, non-metals and metalloids http://www.ptable.com/ Metals ● Shiny ● Malleable and ductile ● Good conductors of heat and electricity Non-metals Sulphur Carbon Selenium Phosphorus Metalloids B Si Ge Both metallic and non-metallic properties. ● They resemble metals physically and non-metals chemically. ● Silicon and germanium are semiconductors. ● Transition metals ● They have very similar chemical and physical properties: – Relatively high mp, high bp and high densities – They form more than one stable cation – They often have coloured compounds and coloured solutions. Lanthanoides ● ● Elements from Z=57 to Z=71 Ex. Europium (Eu) is a hard silvery-white metallic element that is used in the security marking of euronotes. Actinoids ● Elements from Z=89 to Z=103 A collection of uranium glassware ● ● Ex. Naturally occuring uranium is a silvery-white metal. It is a mixture of U-238 (99.27%) U-235 (0.72%) and U234 (0.006%). Uranium decays slowly by emitting an alpha particle. The half-life of U-238 is about 4.47 billion years and that of U-235 is 704 million years. 3.2 Periodic trends Effective nuclear charge ● ● In every atom there is a balance between the attraction of the positively charged nucleus for the negatively charged electrons and repulsion between the electrons. The outer electrons (valence electrons) are shielded from the nucleus and repelled by the inner electrons. ● ● The outer electrons do not experience the full attraction of the positive nucleus because of the presence of inner electrons. The effective charge experienced by the outer electrons is less than the full nuclear charge. Atomic radius ● Atomic radii increase down a group and decrease across a period. ● ● The effective charge increases as a period is crossed from the left to the right: – One proton is added to the nucleus and one electron to the outermost electron shell. – There is no change in the number of inner electrons. The effective charge remains almost the same down a group, because: – The increase in the nuclear charge is offset by the increase in the number of inner electrons. Ionic radius Ionic radii ● ● ● ● Positive ions are smaller than their parent atoms (because of loss of the outer shell). Negative ions are larger than their parent atoms (because of increased electron repulsion by addition of electrons). The ionic radii decrease as a period is crossed from the left to the right (because of increased attraction between the nucleus and the electrons). The ionic radii increase down a group as the number of electron shells increase. Ionization energies ● The energy required to remove one mole of electrons from one mole of gaseous atoms. Electronegativity ● The ability of an atom to attract the shared pairs of electrons in a covalent bond. Electron affinity ● The energy change when an electron is added to an isolated atom in the gaseous state. X(g) + e- → X- (g) Ex. F (g) + e- → F- (g) Ea = -328 kJ mol-1 Melting points ● ● The nature of the bonding between the particles of an element determines its melting point. Strong bonds require higher energy to break: - The stronger the bonds, the higher the melting point. Chemical properties http://www.webelements.com/ Metallic and non-metallic character ● ● Metals have a tendency to lose electrons during a chemical reaction = they tend to be oxidized. Non-metals have a tendency to gain electrons during a chemical reaction = they tend to be reduced. http://www.periodicvideos.com/ Group 18: the noble gases ● Colourless gases ● They exist as single atoms ● They are very unreactive Group 1: the alkali metals ● Shiny, silvery, soft ● Good conductors of electricity and heat ● Low densities ● Too reactive to be found in nature Reaction with water ● ● The alkali metals react with water to produce hydrogen and a metal hydroxide = the resulting solution is alkaline. Reactivity increases down the group. Li: Na: K: https://www.youtube.com/watch?v=uixxJtJPVXk Reaction with halogens 2 Na(s) + Cl2 (g) → 2 NaCl (s) 2 K(s) + Br2 (l) →2 KBr(s) ● The most vigorous reaction occurs between the most reactive alkali metal and the most reactive halogen: francium with fluorine. Group 17: the halogens ● Diatomic molecules Reaction between halogens and alkali metals ● The outer electron moves from the alkali metal to the halogen atom and an ionic halide is formed. 2 Na (s) + Cl2 (g) → 2NaCl (s) ● The most vigorous reaction occurs between elements which are furthest apart in the Periodic Table (Fr and F). https://www.youtube.com/watch?v=Mx5JJWI2aaw Reactions between halogens and halides ● ● ● Reactivity decreases down a group as the atomic radius increases and the attraction for outer electrons decreases. A solution of a more reactive halogen, X2(g), will react with a solution of halide ions of a less reactive halogen, X-(g). Ex. Which of the following chemical reactions are possible? Silver halide precipitate ● The halogens form insoluble salts (= precipitate) with silver: AgNO3 (aq) + NaI (aq) → AgI (s) Metal oxides Metal oxides (Na2O, MgO) are ionic compounds: - solids in room temperature - high mp & bp - conduct electricity when molten (or in aqueous solutions) - basic in aqueous solutions: Na2O (s) + H2O (l) → 2 NaOH (aq) MgO (s) + H2O (l) → Mg(OH)2 (s) Aluminium oxide ● ● Al2O3 is an ionic oxide with some covalent character. It is amphoteric as it acts as a base when it reacts with acids and acts as an acid when it reacts with bases: Silicon oxide ● ● ● Silicon oxide, SiO2, has a giant covalent structure with very high melting and boiling points. It idoes not dissolve in water. It is classified as an an acidic oxide, because it reacts with NaOH at temperatures above 350° C. Non-metallic oxides ● Nonmetal oxides are covalently bonded because of the small difference in the elements' electronegativity values. ● sulfur: SO2, SO3 ● chlorine: Cl2O, Cl2O7 ● phosphorus: P4O6, P4O10 ● They have low mp and bp and do not conduct electricity. ● Non-metallic oxides are acidic in aqueous solutions: P4O10 + 6 H2O (l) → 4 H3PO4 (aq) phosphoric(V)acid SO3 (g) + H2O (l) → H2SO4 (aq) sulfuric(VI)acid Chemical properties of elements in the Period 3 ● ● ● Na, Mg and Al are metals. They are shiny and good conductors of heat and electricity. Si is a semi-conductor and is called a metalloid since it has some of the properties of a metal and some of a nonmetal. P, S, Cl and Ar are all non-metals and do not conduct electricity. Formula of oxide Na2O(s) MgO(s) Al2O3(s) SiO2(s) P4O10(s) SO3(l) and SO2(g) Nature of oxide basic basic amphoteric acidic acidic acidic Acidic rain • Rain is naturally acidic, because the water molecules react with the CO2 in the air and form the weak acid H2CO3. • Acid rain: precipitation (rain, snow) with pH lower than 5.6. ● The main acids present in acid rain are sulfuric acid nitric acid. ● The sulfuric acid in the rain reacts with calcium carbonate (in limestone or marble) to create calcium sulfate, which then flakes off. CaCO3(s) + H2SO4(aq) → CaSO4(aq) + CO2(g) + H2O(l) 13.1 First row d-block elements d-block elements: Transition elements ● ● An element whose atoms have an incomplete d subshell, or which can give rise to cations with an incomplete d sub-shell. Zinc is a d-block element but not a transition metal. Physical properties ● ● ● ● Small atomic radii compared to the neighbouring sblock elements Only a small increase in atomic radii across the period Metallic bonding:high electrical and thermal conductivity, high mp, malleable, ductile, high density Magnetic properties Trends in the first IE ● The rate of increase in the first IE across the period is much lower for transition elements compared to that for the main-group elements. Chemical properties Oxidation number (oxidation state) ● ● The oxidation number of an element keeps track of the number of electrons it has lost or gained. Some elements always have the same oxidation number: Variable oxidation states ● ● When transition metals lose electrons they lose the 4s electrons first. All transition elements (except for Cr and Cu) contain two 4s electrons, which means that they all have an oxidation state of +2 ● In an positive ion all energy levels are closer to the nucleus (more p than e) and as the energy levels move down in energy the order of the 4s and 3d sub-levels are reversed. ● ● ● 3+ The M ion is the most stable for scandium to chromium. The M2+ ion is the most stable for Mn to Zn (the increased nuclear charge makes it more difficult to remove a third electron). In the higher oxidation states the elements usually not exist as a free metal ions, but covalently bonded or as a oxyanions (MnO4-). Complexes ● ● ● Because of their small size and relatively high charge, the transition metal ions have a high charge density. They attract species that are rich in electrons: ligands. A ligand has at least one atom with a lone pair of electrons, e.g. H2O, NH3, Cl-, CN- ● ● These electron pairs form co-ordinate covalent bonds with the metal ion and a complex ion is formed. The number of co-ordinate covalent bonds from ligands to the central metal ion is called the coordination number. ● ● In a co-ordinate bond the shared pair of electrons orginates from the same atom. The charge on a complex ion is the sum of the charge of the d-block metal ion and the charge of the ligand (if they are ions). Ligand exchange (replacement) ● In aqueous solutions the water molecule usually act as a ligand, but it can be replaced with another ligand. Catalytic behaviour ● ● ● A catalyst increases the rate of a reaction without participating in the reaction itself. A catalyst does not become chemically changed at the end of the reaction and can be reused. Many transition elements and their compounds are very efficient catalysts. Heterogeneous catalysts ● The catalyst is in a different state from the reactants. ● Ex. Iron (Fe) in the Haber-Bosch process: N2(g) + 3H2 (g) ● 2NH3 (g) Nickel (Ni) in the conversion of alkenes to alkanes ● Vanadium (V) oxide in the contact process ● Manganese (IV) oxide with hydrogen peroxide Homogenous catalysts ● ● Homogenous catalysts are in the same state of matter as the reactants and products. Transition metals are effective homogenous catalysts in redox reactions since they can relatively easily be oxidized and reduced due to their variable oxidation states. Magnetic properties ● ● ● Every spinning electron can behave as a tiny magnet (= a magnetic dipole). Electrons with opposite spins cancel each other out. Most substances have all electrons paired up and so are non-magnetic. Materials ca be classified as diamagnetic, paramagnetic or ferromagnetic based on their behaviour in an external magnetic field. a) Diamagnetism A property of all materials that do not contain unpaired electrons. ● ● In an external magnetic field the paired electrons orientate themselves such that the field created by their spin opposes the applied field → They are weakly repelled by an external magnetic field. Ex. Ne b) Paramagnetism ● ● ● ● Stronger than diamagnetism. Increases with the number of unpaired electrons (from left to right in the periodic table). The spins of unpaired electrons in an atom or ion can temporarily be aligned in an external magnetic field. These unpaired electrons behave as tiny magnets and are attracted by an external magnetic field. https://www.youtube.com/watch?v=hK_Yi5nxuKQ c) Ferromagnetism ● ● ● ● The largest effect. Only occurs in materials where unpaired d electrons in a large number of atoms can line up with parallell spins to form regions or domains of parallell spins. In an external magnetic field the unpaired electrons in the domains align themselves such that the magnetic field created by their spin is aligned with the applied field. After the external magnetic field is removed, the domains remain aligned making a permanent magnet. Colour of transition metal complexes ● White light is the combination of equal brightness of red, green and blue light (they are primary colours since all other colours can be formed from them). ● When white light is shone on a substance, some light is absorbed and some is reflected. www.youtube.com/watch? v=EHMH0uQDEOU ● ● ● If all light is absorbed, the substance appears black. If only certain wavelengths are absorbed, the compound appears coloured. If all light is reflected, the compound appears white. Transition metals absorb visible light ● ● When white light falls on an aqueous solution of a transition metal complex, the ions absorb some colours. The Fe3+ ion appears yellow because it absorbs light in the blue region of the spectrum. Yellow is the complementary colour to blue (= they are opposite each other in the colour wheel.) The colour of transition metal ions Splitting of the d-orbitals ● ● In an isolated gaseous transition metal atom, the five 3 d orbitals are degenerate since they all have the same energy. However, since the 3 d sub-shells all have different orientations in space they will be orientated differently relative to the ligands in a complex ion. ● ● The 3d electrons close to a ligand will experience repulsion and be raised in energy. The 3d electrons further from the ligand will be reduced in energy. → The 3d sub/shell splits into two energy levels. http://www.chemguide.co.uk/inorganic/complexions/whatis.html ● ● The amount the orbitals are split (∆E) and hence the colour of the complex depends on: – The nature of the transition metal – The oxidation state – The shape of the complex – The nature of the ligand If the orbital is completely empty (Sc3+) or completely full (Cu+, Zn2+), the complexes are colourless. Nuclear charge ● ● Ions with a higher nuclear charge form stronger coordinate bonds with the ligands lone pair of electrons (= because of stronger electrostatic attraction). This stronger attraction causes a larger split in the 3d orbitals (larger ΔE = higher energy light is absorbed). ● Mn2+ is pale pink (green light absorbed) ● Fe3+ is yellow (blue light is absorbed) Charge density of the ligand ● ● ● Ligands can be arranged in order of their ability to split the d orbitals in octahedral complexes. Spectrochemical series: Ammonia has greater charge density than water and produces a larger split in the d-orbitals. UV-vis absorbtion spectrum of some copper complexes