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Certificate in Plasma and Vacuum Studies – PVS103 - Unit 6 Lecture 6 Chemical Properties and the Periodic Table ________________________________________________________________________________ Part 1 Part 2 Part 3 Part 4 Part 5 Periodic Properties of the Elements Groups 1a & 2a Metals and Non-metals Groups 3a to 8a; the Non-metals Groups 3b to 12b; the Transition Metals ________________________________________________________________________________ Part 1 Periodic Properties of the Elements We have already seen how electronegativity increases on moving across the periodic table. Can you remember why? As a result electronegativity is a periodic property, so for example, the halogens (Group 7a elements) are all highly electronegative. Many other atomic properties are periodic because of the periodic nature of electronic configurations of the atoms. Note: As you go down a group, the 1s orbital is moves closer to the nucleus, because of greater attraction due to the increase in the nuclear charge, Z. For many purposes we can consider the nucleus and inner (non-valence) electrons as being an effective field in which the valence electrons move. DCU©2002 1 Certificate in Plasma and Vacuum Studies – PVS103 - Unit 6 In other words the effective nuclear charge, Z*, experienced by any given electron is given by: Z* = Z - σ Where Z is the actual nuclear charge σ is the screening constant Slater’s Rules are used to estimate σ 1. Write down electron configuration : eg (1s) (2s, 2p) (3s, 3p, 3d) (4s, 4p, 4d) 2. Electrons to the right of the electron of interest contribute zero to σ 3. Electrons in the same (s,p) group contribute 0.35 to σ, for 1s the value used is 0.30 4. All electrons in the (n-1) group contribute 0.85 to σ 5. All electrons in the (n-2) and lower contribute 1.00 to σ Consider He, Ne, Ar: He 1s2 1s electrons Z* Ne = 2 - 0.30 = 1.7 = 9.7 1s2 2s2 2p6 Z* = 10 - σ For a 2s or 2p electron in Ne Z* Ar = 1s2 2s2 DCU©2002 10 -( 7 x 0.35) - (2 x 0.85) = 5.85 2p6 3s2 3p6 2 Certificate in Plasma and Vacuum Studies – PVS103 - Unit 6 1s electron Z* = 18 - 0.30 = 17.70 For a 2s or 2p electron in Ar Z* = 18 - 4.15 = 13.85 Z* = 18 - ( 7 x 0.35) - (2 x 0.85) = 13.85 For a 3s or 3p electron in Ar Z* = 18 - ( 7 x 0.35) - (8 x 0.85) - 2 = 6.75 Consequences: 1. Quantum mechanics suggests that electrons in an atom can be described in terms of their most probable distance from the nucleus 2. For a given shell, the electron distribution function moves closer to the nucleus as Z increases 3. As Z increases, the inner shells are more tightly held and only the outer more diffuse electrons can be disturbed These findings affect many atomic properties, so there are several periodic properties: DCU©2002 3 Certificate in Plasma and Vacuum Studies – PVS103 - Unit 6 (a) Atomic Radii: The atomic radius can be calculated by measuring the internuclear distance in compounds of like elements e.g. H2, N2, Cl2 Where the atomic radius is taken to be half the internuclear distance between like atoms. Trends in the atomic radius are observed: 1. Across a period e.g. Li → F, size of the atom decreases. This is because Z increases, but each electron added does not fully shield the outer electrons from the nuclear charge. ⇒ Z* increases and the size reduces. 2. Going down a group the size of the atom increases This is because we are putting electrons into orbitals of higher principal quantum number, which are further from the nucleus. DCU©2002 4 Certificate in Plasma and Vacuum Studies – PVS103 - Unit 6 (b) Ionisation Energy: Is defined as the energy required to remove an electron from an atom or ion M(g) → M+ (g) + e M+ (g) → M2+ (g) + e - 1st IE - 2nd IE etc. For any element 1st IE < 2nd IE < 3rd IE …etc... The repulsive forces between electrons help in the removal of electrons On removing an electron, the screening of one electron by the others is reduced, therefore the effective nuclear charge (Z*) increases and further removal of electrons becomes more difficult. For example silicon, Si [Ne] 3s2 3p2 Ionisation energies (kJ/mol) 1st 780 2nd 1,575 3rd 3,220 4th 4,350 5th 16,100 IE (1) → IE (4) gradual increase as 3s and 3p electrons removed Then, a large increase in IE (5) as a 2p electron is removed. This electron is closer to the nucleus (lower n) and much more tightly held. DCU©2002 5 Certificate in Plasma and Vacuum Studies – PVS103 - Unit 6 Across a Period e.g. Li → Ne Alkali elements have lowest IE, while inert gas elements have highest IE. Note There are a number of discrepancies in going across a period; i.e. the trends are not perfect. 1. For B, C, N the IE increases, but the increase is lower than might be expected. Because we are now removing p electrons, these don’t penetrate as close to the nucleus as do s electrons; therefore they are more easily removed. 2. For O and F the IE are again lower than predicted, because we are now removing electrons that are spin paired. It is easier to remove these because the electrons are repelling each other, which aids the removal of electrons. The trend going down any Group: • The 1st IE decreases significantly. • This is because of the increasing number of inner filled shells screening the valence electrons, which become easier to remove. • So the following process becomes more favourable energetically: M → M+ + 1e- Later we will see that this is associated with the generally observed increase in metallic character going down a group. DCU©2002 6 Certificate in Plasma and Vacuum Studies – PVS103 - Unit 6 (c) Electron Affinity: Is the energy change that takes place when an electron is added to a gaseous atom or ion: M(g) + e- → M-(g) M(g)+ + e- → M(g) reverse of I.E. A negative value for EA indicates a release of energy when the anion is formed, i.e. that the anion is more stable than the neutral atom. Consider the halogens: Hal (ns2 np5) + e → hal- (ns2 np6) This should be favoured since X- has a full octet. E.A (kJ/mol) → F- -330 Cl Cl- -350 Br Br- -325 I I- -295 F DCU©2002 7 Certificate in Plasma and Vacuum Studies – PVS103 - Unit 6 The electron goes into an orbital further from the nucleus as we go down the group (F→I), you would predict that the EA should decrease. So you would expect F to have highest EA It turns out that electron-electron repulsion terms are more important in this case and they counterbalance the above stabilisation. ⇒ the EA of all the halogens are similar. DCU©2002 8 Certificate in Plasma and Vacuum Studies – PVS103 - Unit 6 Part 2 Groups 1a & 2a The chemistry of the elements in Groups 1a and 2a are determined by their low electronegativities & ionisation energies and also by their relatively large atomic radii. Group 1a; hydrogen Having said this, the chemistry of hydrogen is completely different from the other elements in Group 1a: Unlike the other elements it has a relatively high IE and electronegativity, it commonly forms covalent bonds, particularly to carbon. These properties arise because of its very small atomic radius. Exercise: Explain (i) why hydrogen is so small, (ii) why its small size results in high IE and electronegativity. Note: When we refer to the chemistry of hydrogen, we are actually talking about; H 1s1 the hydrogen atom; 1 proton and 1 electron 1s2 the hydride anion; 1 proton and 2 electrons or H- The latter can and does occur, as the hydrogen atom has moderate electron affinity. DCU©2002 9 Certificate in Plasma and Vacuum Studies – PVS103 - Unit 6 On the other hand, the hydrogen cation: H+ 1s0 consists of just a proton This can occur (though always associated with other ions). It is the entity, which is transferred in acid-base chemistry. So an acid is a proton (or H+) donor and a base is a proton acceptor. This type of behaviour will be dealt with later in terms of the molecules, which have this ability. The chemistry of hydrogen is dominated by covalent bond formation in which hydrogen shares electrons with other elements H• + •X → H:X If X not equal to hydrogen, the covalent bond will be polar If X has high electronegativity e.g. fluorine, F HF δ+ δ- H F -the δ refers to the molecular centre of charge hydrofluoric acid If X is less electronegative than hydrogen (i.e. is electropositive), e.g. Na δ- δ+ H Na HNa sodium hydride. Metal hydrides generally are salt-like materials, i.e. like the ionic NaCl. Covalent hydrides: DCU©2002 10 Certificate in Plasma and Vacuum Studies – PVS103 - Unit 6 Hydrogen forms strong stable single covalent bonds with elements including: B, C, N; DCU©2002 eg. BH3, CH4, NH3 11 Certificate in Plasma and Vacuum Studies – PVS103 - Unit 6 Group 1a continued, the alkali metals Li [He] 2s1 Na [Ne] 3s1 K [Ar] 4s1 Rb [Kr] 5s1 Cs [Xe] 6s1 Fr [Rn] 7s1 • Na and K are the most abundant • Due to low electronegativity and ionisation energy the compounds formed are mainly ionic although some covalent bonding does occur. i.e. the dominant process in their chemistry is: M → M+ + 1e- The loss of an electron is termed oxidation; so alkali metals are easily oxidised. Reactions of Alkali Metals • Reaction with Halogens (Group 7a elements) Due to the large electronegativity difference, all known halides are ionic (CsF). So these reactions invariably result in the formation of salts. 2M DCU©2002 + X2 → 2MX 12 Certificate in Plasma and Vacuum Studies – PVS103 - Unit 6 • Oxygen compounds They react rapidly with oxygen; different types of oxides can exist Li + O2 → Li2O Na + O2 → Na2O2 O2 → KO2 K + Exercise: Draw Lewis structures for the three oxides. Hint: The lithium compound contains the oxide ion [O]2-, the sodium compound contains the peroxide ion [O2]2- and the potassium compound contains the super-oxide ion [O2]-. Hydroxides Strongly basic hydroxides are formed when the metals react with water: Li + H2O → LiOH + H2 Na + H2O → NaOH + H2 K + H2O → KOH + H2 Exercise: On the basis of the known trends in electronegativity and ionisation energy going down the group; which of these reactions leading to the formation of Group 1a hydroxides is the most vigorous? -In fact it is explosive! Reaction with ammonia DCU©2002 13 Certificate in Plasma and Vacuum Studies – PVS103 - Unit 6 The alkali metals react with ammonia, NH3, to yield H2 gas and a metal amide M DCU©2002 + NH3 → MNH2 + H2 14 Certificate in Plasma and Vacuum Studies – PVS103 - Unit 6 Group 2a, the alkaline earth metals Be [He] 2s2 Mg [Ne] 3s2 Ca [Ar] 4s2 Sr [Kr] 5s2 Ba [Xe] 6s2 Ra [Rn] 7s2 • Like the alkali metals they tend to oxidise, in this case they usually lose two electrons to yield dipositive ions: → M M2+ + 2e- • They tend to be less reactive than the group 1a elements -this is because the first ionization energies are larger • They react with halogens to form ionic halides, which are usually salts. Ca + Cl2 → CaCl2 Beryllium is exceptional in the Group; BeF2 and BeCl2 are molecular substances. This is because of Be’s higher electronegativity. • They react with oxygen to form oxides 2 Mg + O2 → 2 MgO • Again, with the exception of Be the alkaline earth metals react with water to yield metal hydroxides M(OH)2, which are weakly basic: Mg DCU©2002 + 2H2O → MgOH + H2 15 Certificate in Plasma and Vacuum Studies – PVS103 - Unit 6 Part 3 1a 1 3 11 Na 19 37 2a H Li 4 4b 5b 6b 7b 8b 9b 12 Rb 38 Ca 56 Sr Ba B 13 Mg Sc 39 57 22 Y 40 La 72 Ti 23 V 24 Nb 42 Zr 41 Hf 73 Ta Cr 25 Mo 43 74 W Mn 75 26 Fe 27 Co 28 45 Rh 46 Tc 44 Ru Re 76 Os 77 5a 6a 7a Ir Ni Pd 47 Ag 48 Cd 49 79 Au 80 Hg 81 78 Pt Zn 31 Al 29 Cu 30 8a 2 5 21 4a 10b 11b 12b Be K Cs 3a 3b 20 55 Metals and Non-metals Ga 6 C 14 32 Si Ge 7 N 15 33 P 8 16 17 Se 35 As Sb 52 50 Sn 51 Th 82 Pb 83 Bi 84 F S 34 In O 9 Te Po He 10 Ne Cl 18 Ar Br 36 Kr 53 85 I 54 At 86 Xer Rn metals non-metals There is an approximate division between metals and non-metals. This is shown in the PTE above. However, around the division the metallic elements have some non-metallic properties, and visa versa. Metals • In a metal there are no molecules. • Instead the atoms are arranged in ordered arrays (or grids), similar to the arrays of atoms in the crystals of NaCl. • However, unlike in silicon, in metals the electrons in the bonds are not fixed in positions. DCU©2002 16 Certificate in Plasma and Vacuum Studies – PVS103 - Unit 6 each sphere represents an atom in the metal each atom is bonded to the atoms around it the electrons in each bond (three are drawn in here) are continuously swapping positions, randomly The electrons in metals are moving around continuously. In fact although the atoms in a metal have fixed positions, the electrons behave as if they were in a “gas”, i.e. they move about (diffuse) at random. the electrons are better thought of as being a “gas”, shown here as a blue haze The electrical conductivity of metals arises from the fact that these electrons are mobile and so can carry electric charge through the metal; i.e. there are “mobile charge carriers” available. DCU©2002 17 Certificate in Plasma and Vacuum Studies – PVS103 - Unit 6 Orbital view of metals: • In a diatomic molecule, eg H2, the two AOs give rise to two MOs; a bonding and antibonding pair. • So initially there are two isolated AOs of the same energy, but on bonding two states of different energy are formed. • It is always the case that the more orbitals overlap the more states are formed. • In a metal there are so many orbitals overlapping that effectively an infinite number of states with different energies arise; a band. 2nd row diatomic molecule H2 metal empty anti-bonding band σ∗2pz σ∗1s * π*2px, π2p y σ1s more orbitals π2px, π2py σ2pz band gap more orbitals σ∗2s partially filled conducting band σ2s • This band is a state (orbital) extending over the metallic solid. • An occupied band is called the valence band. If it is full, or the gap to the empty upper band (antibonding states) is significant, the electrons are not free to move. • In metals one of these criteria is met, so metals conduct electricity. DCU©2002 18 Certificate in Plasma and Vacuum Studies – PVS103 - Unit 6 This simple model can be used to rationalise the electrical conductivities of a range of materials: metal metal semi-conductor insulator ∆E ∆E valence band is not full, so it serves as the conducting band conducting band full valence band as in semiconductors overlaps the doesn’t overlap with but the energy valence band the empty conducting gap is large band but the energy gap is small • Metals conduct electricity due to the presence of a partially filled conducting band. • Semiconductors conduct electricity if the energy gap to the conducting band can be jumped, so if the temperature is raised conductivity increases. More commonly a low concentration of an impurity is added which introduces some extra electrons into the system (an n-type implant). • Insulators do not conduct electricity, as very few electrons can jump the energy gap. DCU©2002 19 Certificate in Plasma and Vacuum Studies – PVS103 - Unit 6 The Physical Properties of Metals: • Under normal conditions they are usually solids, with high melting and boiling points. For instance copper, Cu, melts at 3000°C. There is one exception; mercury, Hg, is a liquid at room temperature. • Because of the close packing of the atoms, they are dense materials. So a given volume of a metal will be heavier than other materials. • They usually are hard materials, but they are also usually malleable, i.e. they can be shaped. There are some exceptions, in particular the Group 1a elements (alkali metals), which are relatively soft solids. • They are good conductors of heat and electricity. The Chemical Properties of Metals: These arise because they tend oxidise (lose electrons), for example: Fe → • Metals react with oxygen to form metal oxides, which are bases. 4Fe(s) + 3O2(g) → • Fe3+ + 3e- 2Fe2O3(s) RUST Metals react with water to form Metal hydroxides, which are bases, and hydrogen. Na(s) + 2H2O (l) → 2NaOH (l) • + H2(g) Metals react with acids to form metal salts and hydrogen. Zn(s) + H2SO4 (l) → ZnSO4 (l) W(s) + H2SO4 (l) → + H2(g) no reaction so tungsten is a less active metal • Metals form chlorides, which are ionic materials. 2Na(s) + 2HCl(g) → 2NaCl (l) DCU©2002 + H2(g) 20 Certificate in Plasma and Vacuum Studies – PVS103 - Unit 6 Non-metals As we have seen previously, non-metallic elements tend to form molecular substances. These molecules do not pack as closely as the atoms in a metal. H Si H weak force of attraction H Si H H H H H Nearby molecules do not approach each other very closely. The only attractive forces are quite weak, these weak intermolecular forces are called “Van der Waals interactions”, they are only important for very large molecules. This is in contrast to the strong forces which hold metals and ionic solids together, so metals and ionic solids are dense materials. The Physical Properties of Non-Metals: • • • • They have low melting and boiling points. They have low density compared for instance to metals. They form soft solids liquids or gases (exceptions include diamond). They are poor conductors of heat and electricity. e a poor conductor, so in Intel it is used as an insulator between metallic layers (exceptions include graphite, a form of carbon, which conducts heat and electricity). DCU©2002 21 Certificate in Plasma and Vacuum Studies – PVS103 - Unit 6 The Chemical Properties of Non-Metals: • They usually gain electrons to form negative ions (anions). F + 1e- →F- • When reacted with oxygen they form oxides, which are acidic C(s) + O2(g) CO2 (g) + H2O(g) • • → CO2 (g) H2CO3 (g) They do not react with dilute acids. They form chlorides, which are liquids or gases. So CCl4, SiF4 and SiH2Cl2 are volatile liquids or gases at standard temperature and pressure (STP). DCU©2002 22 Certificate in Plasma and Vacuum Studies – PVS103 - Unit 6 Part 4 Groups 3a to 8a; the Non-metals The chemistry of the elements from Group 3a to 8a are determined by their moderate to high electronegativities & ionisation energies and to a lesser extent by their larger atomic radii & moderate electron affinities. So they tend to form covalent bonds which result in the formation of molecular substances. Group 3a, boron and the electron deficient elements 5 [He] 2s22p1 B 13 Al [Ne] 3s23p1 31 Ga [Ar] 4s24p1 49 In [Kr] 5s25p1 • On going down the group, the metallic character of the elements increases. • Boron is unique in the group in that it is clearly a non-metal, we will concentrate on its properties, as it is very interesting. • The molecules boron forms are unique in that they do not conform fully to Lewis theory, for instance BH3 is a stable molecule, but there is no octet of electrons on boron. In fact there are only 6 electrons formally. • For this reason we say that boron forms electron-deficient molecules. DCU©2002 23 Certificate in Plasma and Vacuum Studies – PVS103 - Unit 6 Exercise: Firstly draw the Lewis structure of ethane, C2H6, and then of diborane, B2H6. Do you agree that the only possible structure for B2H6 is the following? H H H B B H H H 2- However, both ethane and diborane exist; they are neutral (uncharged) and are covalent. If you draw B2H6 as a neutral molecule with a boron-boron bond there are not enough electrons to form all the bonds But B2H6 exists and like C2H6 it is neutral (uncharged) and covalent. The actual structure is now accepted as: H H B H H B H H • Here there are only really six bonds (12 electrons); the two bridging B-HB groups are held together by a pair of three-centre two-electron bonds and there is no direct boron-boron bond. • This is a classical case of the failure of the VB/Lewis approach. • However MO theory can easily accommodate this type of bonding; an MO containing two electrons arising from the overlap of AOs on three atoms is perfectly acceptable in MO theory. DCU©2002 24 Certificate in Plasma and Vacuum Studies – PVS103 - Unit 6 Group 4a, the carbon group 6 [He] 2s22p2 C 14 Si [Ne] 3s23p2 32 Ge [Ar] 4s24p2 50 Sn [Kr] 5s25p2 • Carbon is unique in the group in that it is clearly a non-metal. • It is also the element, from which by far the greatest number of molecules are formed; the chemistry of carbon is called organic chemistry. • This great diversity exists because of the carbon-carbon and carbonhydrogen bonds are stable covalent bonds. • Organic chemistry does not form part of this course. • The chemistry of silicon is less rich than that of carbon, partially because Si-Si and Si-H bonds are weaker. • On going down the group, the metallic character of the elements increases, so tin, Sn, is clearly a metal. Some organic molecules: NH2 N H N OH O HO N H O H H2N C OH OH H CH3 glucose - a sugar is a six-membered ring involving five carbons and an oxygen H H OH H H HO DCU©2002 N HO O H HO CH alanine - an amino acid is one of the building blocks of proteins it has a carbon-nitrogen backbone H the adenine nucleoside - is one of the building blocks of DNA 25 Certificate in Plasma and Vacuum Studies – PVS103 - Unit 6 The chemistry of silicon, Si: • Unlike carbon, silicon does not tend to form large molecules as S-S and S-H bonds are weaker than the carbon counterparts. • The major applications of silicon are in the semiconductor industry. Silicon oxidation: Si(s) + O2 (g) → SiO2 (s) Silicon oxide is an electrical insulator, so it is used to electrically isolates individual transistors and layers on an IC. Furthermore it acts as a contaminant sink, i.e. contaminants from various parts of the chip tend to gather in the oxide, where they are subsequently removed (sacrificial oxide). Formation of silicon nitride: 3SiH2Cl2 (g) + 4NH3(g) → Si3N4 (s) + 6HCl(g) + 6H2(g) Silicon nitride (Si3N4) is a white, chemically inert (i.e. unreactive), amorphous powder (i.e. unlike silicon, it has no defined crystal structure). It has the ability to form thin films. • These reactions take place using chemical vapour deposition (CVD), the reactants are gases which coat the surface of the solid silicon wafer with a silicon derivative. DCU©2002 26 Certificate in Plasma and Vacuum Studies – PVS103 - Unit 6 Group 5a, the nitrogen group 7 [He] 2s22p3 N 15 P [Ne] 3s23p3 33 As [Ar] 4s24p3 51 Sb [Kr] 5s25p3 • Again, on going down the group, the metallic character of the elements increases. • Nitrogen and phosphorus exhibit valencies of 3 or 5; thus PCl3 and PCl5 both exist, and in both cases the phosphous atom is neutral. Chemistry of Nitrogen • Dinitrogen, N2, is an extremely stable molecule. This is because of the presence of a covalent triple bond (draw the Lewis structure to convince yourself). Air in 80% N2. • Ammonia, NH3, is a very important molecule, it is produced commercially on a huge scale in the Haber process: N2 (g) + 3H2(g) 2NH3(g) - historically this was one of the first industrial scale chemical reactions. The ammonia produced is used in the fertiliser, explosive and polymer industries. • There are at least six stable oxides of nitrogen; NO2, NO, N2O, N2O3, N2O4 and N2O5, in which the nitrogen atoms exhibit a range of oxidation states (formal charges) from +1 to +5. As an exercise you could try drawing some of these. DCU©2002 27 Certificate in Plasma and Vacuum Studies – PVS103 - Unit 6 Chemistry of Phosphorus • The most important phosphorus molecules are phosphates or phosphorous oxides, which (like nitrates) are important fertilisers. • The most common form is the (PO4)3- anion, which is relatively stable due to delocalisation of the negative charge. Can you draw its Lewis structure? • The true structure (the average of the four possible equivalent Lewis structures) is a tetrahedron with the “3-” charge is distributed evenly over the four oxygen atoms, which are positioned at the apices of the tetrahedron (the phosphorus atom is at the centre). • On heating, phosphate groups have a tendency to lose water and form polyphosphates. 2H3PO4 → H4P2O7 + H2O Pyrophosphoric acid, H4P2O7, is used for etching in the semiconductor industry. • Long-chain polyphosphates can also form; their structure is a series of tetrahedrons, linked at the apices. DCU©2002 28 Certificate in Plasma and Vacuum Studies – PVS103 - Unit 6 Group 6a, the oxygen group 8 [He] 2s22p4 O 16 S [Ne] 3s23p4 34 Se [Ar] 4s24p4 52 Te [Kr] 5s25p4 • Again, on going down the group, the metallic character of the elements increases. • Oxygen gas, O2 is a highly oxidising substance: i.e. atoms to which oxygen adds are usually oxidised (they lose electrons) due to the high electronegativity of oxygen. • The most important compound on the planet is probably water, H2O. As we saw earlier, this is a liquid at STP (standard temperature and pressure), even though it is a very light molecule because of the extensive hydrogen-bonding. Life on earth evolved in water and is dependent on it; all animal and planet cells are composed mostly of water. • The difference between oxygen and sulphur is well illustrated by the different properties of H2O and H2S. As we saw earlier H2S has a different shape (because of the larger 3p orbitals) and it is a gas at STP because of the reduced hydrogen bonding (S is less electronegative). • The major compounds of sulphur are its two oxides; SO2 and SO3. In water sulphur dioxide is in equilibrium with sulphurous acid SO2 + H2O H2SO3 In water sulphur trioxide is in equilibrium with sulphuric acid, SO3 + DCU©2002 H2O H2SO4 29 Certificate in Plasma and Vacuum Studies – PVS103 - Unit 6 This is a much stronger acid than sulphurous acid, because the [HSO4]- anion is more stable than the [HSO3]- anion. DCU©2002 30 Certificate in Plasma and Vacuum Studies – PVS103 - Unit 6 Group 7a, the halides 9 [He] 2s22p5 F 17 Cl [Ne] 3s23p5 35 Br [Ar] 4s24p5 53 I [Kr] 5s25p5 • The chemistry of this group is dominated by their very high electronegativity, which decreases going down the group. • They have a strong tendency to form very stable salts with Group 1a and 2a elements, e.g. NaF, NaCl etc... • They form some of the strongest acids known, acids are proton donors. Consider what happens to these molecules in solution: HF + H2O F- + H3O+ HCl + H2O Cl- + H3O+ HF is a stronger acid than HCl, i.e. the proton concentration,[H+], is higher (the equilibrium is even further to the right than for HCl). This is due to the greater electronegativity of fluorine. Chemical equilibria and acid-base chemistry will be dealt with in the 2nd chemistry module. DCU©2002 31 Certificate in Plasma and Vacuum Studies – PVS103 - Unit 6 Group 8a, the inert (or noble) gases 2 He 1s2 10 Ne [He] 2s22p6 18 Ar [Ne] 3s23p6 36 Kr [Ar] 4s24p6 54 Xe [Kr] 5s25p6 • The chemistry of this group is dominated by their full outer shell, which means they tend not to react as they already have an octet. • There are some exceptions to this, particularly going down the group; xenon which has a full 5th shell which is very large. • Xe does form some compounds with strong oxidising agents (materials which remove an electron) or electronegative elements such as fluorine. • Thus PtF6 + Xe → [Xe]+PtF6- … and compounds such as XeF2 and XeF4 exist. DCU©2002 32 Certificate in Plasma and Vacuum Studies – PVS103 - Unit 6 Part 5 Groups 3b to 12b; the Transition Metals The chemistry of the elements in Groups 3b to 12b is influenced by their moderately low electronegativity & ionisation energies. In this sense they are metallic like the elements of Group 1a and 2a. However their chemistry is dominated by the presence of partially filled d-orbitals. Remember the order of orbital energy in atoms: 1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p This means that the 4s is usually full. So the electronic configuration of (26Fe)2+, which has 24 electrons, is: (26Fe)2+; 1s22s22p63s23p64s23d4 • Transition metals can exist with variable oxidation state and have strong catalytic activity. • So materials such as FeCl2 and FeCl3 can co-exist, in which iron has a formal charge (oxidation state) of +2 and +3 respectively. • The chemistry and bonding of transition metal compounds (more generally called transition metal complexes) is a huge subject. The theories of bonding we have met in this course are severely tested by these types of materials. • A detailed examination of transition metal chemistry is beyond the scope of this course. DCU©2002 33 Certificate in Plasma and Vacuum Studies – PVS103 - Unit 6 Summary Lecture 6 Chemical Properties and the Periodic Table • A number of properties including electronegativity, atomic radii, ionisation energies and electron affinities are periodic in nature. • For this reason the chemistry of most of the elements is periodic, i.e. elements on the same Group react in similar ways to form similar compounds. • The major division of the periodic table is into metals and the nonmetals. The properties of each of these two classes of materials are readily understood using simple arguments based on the periodic properties and MO theory. • The chemistry of the transition metals is dominated by the presence of partially filled d-orbitals. DCU©2002 34