Why transition elements have a high energy of hydration

('- exercises in the experimental lecture


1 Note This file is a protocol that contains a lecture as part of the chemistry teacher training course at the University of Marburg. For better searchability, text recognition was also carried out and placed behind the scanned image, so that copy & paste is possible, but be careful, the text recognition has not been corrected and is particularly flawed in files that are difficult to read. All of the more than 700 protocols at the beginning of 2007 can be viewed and downloaded from the site. In addition, further experiments, learning circles and state examination papers are available on the site. Dr. Ph. Reiss, July 2007

2 _.- - ~ --- Exercises in the experimental lecture 2. Lecture: K ~ er as a transition element SS 80 Michael Thie A broad definition of the transition elements means that they are characterized by the possession of partially occupied d or f shells. Theoretically, as a transition element of the first transition series, the element copper should have the electron configuration 4s 2 3d 9 in the ground state according to the laws governing the structure of the electron shell. One finds, however, a configuration in which the outermost s-shell is fully occupied with only one and the d-shell underneath with ten electrons. This is explained by the fact that, on the one hand, the energy difference between the 3d and 4s states is only small and, on the other hand, a d-shell fully occupied with ten electrons is more stable than one occupied with nine electrons, so that the more favorable arrangement 3d 9 4s with little energy gain 1 results. see slide) Because of this electron configuration, a comparison with the elements of the alkali group is appropriate, since they also only have one electron in the outermost s-shell; Here it should be remembered that chemical reactions are to be interpreted as interactions between the electron shells of atoms ~ u. see slide 11) In my lecture I will present some reactions of copper and either show parallels to the chemical behavior of the alkali elements or work out the independence of copper as a transition element. First I would like to deal with the metallic state of copper - the metallic state is distinguished by a number of characteristic properties, thermal and electrical conductivity, metallic luster, mechanical deformability and hardness, the structural principle of the closest packing of spheres and the possibility of forming positive metal ions. I will present the dissolving behavior of copper in more detail in the following: Experiment 1: Reaction of Cu with acids 3 test tubes 30 X 200 mm) are mixed with the same amount of Cu chips and with conc. HC, conc. H 2S04 and conc. HN0 heated. After the reaction mixture has cooled, the solutions are then added to 2N NaOH in each case 3 '- - 2. Cu chips react with conc. HG does not develop a gradual solution of Cu in conc. HCl is due to the presence of atmospheric oxygen as an oxidizing agent); also with kanz. H 2S0 4 does not react at room temperature, whereas strong S02 development and dissolution of the Cu chips can be recorded after heating; with conc. HN0, on the other hand, starts a reaction 3 at room temperature and the development of N0 2 is intensified when heated. In order to also detect the dissolved Cu 2+ ones, I pour the reaction mixture into alkali, whereupon a light blue precipitate occurs in the presence of Cu 2+ ones, which should be used as evidence and will be discussed in detail later in the lecture. see slide) If you want to take the reaction of a metal with acids as a criterion for classifying noble and base metals, it should be noted that copper is only concentrated from oxidizing acids. HN0 and hot conc. H 3 2S04) is attacked, but is not dissolved by non-oxidizing acids. A comparison of the solution-seeking copper with other metals shows the following picture: Versu ~ 2: Position of the Cu in the voltage series 3 beakers with 1m CuS0 4, 1m FeS0 4 and 1m AgN0 3 and in the CuS0 4 solution an iron sheet, and in the Immerse one copper sheet each in FeS0 4 or AgN0 solution. 3 At the same time, a CuS0 4 solution is mixed with iron filings - to be prepared beforehand). The iron sheet is immediately covered with a thin layer of metallic copper; So Cu 2+ ones have been reduced to Cu atoms - the electrons required for this are supplied by Fe atoms, which have gone into solution as Fe 2+ ones. In the opposite case, however, the metallic copper is not able to supply the electrons necessary for the reduction of Fe 2+ ones. If, on the other hand, copper is brought into a solution of AgN0 3, it is immediately silver-plated. In those cases in which reactions occur on the immersed metal sheets, the corresponding ones in the solution can also be detected analytically. Upon addition of conc. NH to the AgN0 solution turns blue 3 3

3 - 3 - of the solution, which should be taken as evidence of Cu 2+ ones. Ag + ons do not interfere with this copper detection, since they form a colorless [AgNB 3) 2] + complex; I will discuss the copper proof in more detail later. The iron ions are detected with [Fe CN) e;] 3- as Prussian blue in the almost Cu 2+ one ~ free solution of the parallel experiment, since Cu 2+ ones form a yellow-black complex with red blood liquor salt, which would mask the color of Prussian blue. see slide 111) In order to be able to quantify the attempt to solve the problem, one only needs to separate the two partial reactions of the redox process, which is achieved by a diaphragm and which means that the electron transfer can now be measured directly: Experiment 3: Galvanic Element 1m FeS0 4 solution with iron electrode, 1m CuS0 4 solution with copper electrode; Voltmeter see slide V) The specified galvanic element shows a voltage of approx. 0.7 volts. Assuming the normal potential of the redox couple Fe / Fe 2 + with -0.44 volts as known, then a normal potential of +0.35 volt is calculated for the redox couple Cu / Cu 2 +. The creation of the voltage of the galvanic element can be explained as follows: Between a metal rod and its corresponding metal salt solution there is an equilibrium - metal atoms go into solution as ones, so that the free electrons remaining in the metal rod are negatively charged against the solution. However, metal ions are again attracted to this negatively charged electrode and discharge Me ~ Me n + + n- e-; in the case of the redox couple Cu / Cu 2 + this equilibrium is strongly on the side of the copper, while in the redox couple Fe / Fe 2 + the equilibrium is more on the side of the iron ions. As a result, electrons are shifted from the iron electrode, the location of higher electron pressure, to the copper electrode, the location of lower electron pressure, whereupon Cu 2+ ions from the solution are reduced to metallic copper. See slide V) A comparison of the normal potential with that of the alkali elements and other metals now shows that copper is more noble than the alkali elements and iron and less noble than e.g. Silver is. see slide V) The more noble character of copper compared to the alkali elements is primarily due to the higher onization energy of copper - with copper the fully occupied d-shell can the s-electron much less effective against the attractive effect of the positive - to the shield nuclear charge far greater than the noble gas shell of the alkali elements can, so that this electron is much more firmly bonded to L; A comparison of the first onization energies of copper and the elements of the first main group confirms this theory (see slide V) I would now like to briefly look at the other properties of the metallic state already mentioned (see slide V): be a metal crystal consists of a solid lattice of positive metal ions, with the valence electrons moving freely in this lattice, like gas particles. Since there are only one type of lattice building block in such a crystal, highly symmetrical crystal structures are possible - the alkali metals consistently have body-centered cubic structures, while copper has face-centered cubic structures. According to this electron gas model, the actual metallic bond is a consequence of the strong delocalization of the valence electrons, which are freely movable within the entire crystal. In the case of a metallic bond, the individual atoms of the alkali metal12 in the metal lattice have only one binding electron available - the bond is relatively weak, the metals have low melting and boiling points and are so soft that they can be cut with the metal In addition to the electron in the outer shell, copper can also be involved in the bond between the electrons in the inner orbitals, so that higher melting and boiling points, greater hardness and better conductivity result. - The main thing is that the Alkalimet all occur in the +1 oxidation state, since further onization would affect a very stable noble gas shell. In the oxidation state they also occur in aqueous solution. In the +1 oxidation state, copper has the electron configuration 3d 10, if the atoms form ons, electrons are first given off from the s orbital, since the d orbitals in the on are apparently lower in energy than the s orbital) . See slide V The compounds of the monovalent copper are generally difficult to represent in aqueous solution because most of them are easily oxidized to the compounds of the divalent state. The equilibrium is found in aqueous solution:

4 - 5-2 Cu + aq) ~ Cu + Cu + aq) on the side of the Cu and Cu + aq), a phenomenon that can be explained by considering the potential of the redox couples Cu / Cu + and Cu + / Cu 2+: because the potential of Cu / Cu + is more positive than that of Cu + / Cu 2+, Cu + from the first redox couple can act as an oxidizing agent and there is disproportionation of Cu + into Cu and Cu2 +, basically the higher hydration energy of Cu 2+ Cu + decisive). However, by forming Cu compounds that are sparingly soluble in water, the +1 oxidation state can be stabilized in aqueous solution: Experiment 4: Formation of CuI 0.5 m CuSO solution, add the same volume of equal molar KI 4 solution and heat; then add sulphurous acid. Cu-1 reacts with I- to form the corresponding Cu - compound, the iodine which precipitates out initially causes the reaction mixture to turn brown when heated, violet "iodine vapors" appear); sulphurous acid reduces iodine to iodide. From a chemical point of view, the Cu 2+ one has distorted the large, easily deformable iodide ion so strongly that the electron of the iodide ion is transferred entirely to the copper ion and this is reduced, the ion being oxidized to iodine. Cu reacts immediately with further J to form CuI, which precipitates out as an insoluble precipitate. See slide V) The colored oxide represents a further stable form of Cu: Experiment 5: Precipitation of red Cu 20 Fehling 7g CuS04 '5 H 20 to 100 ml distilled water) and Fehling 11 35g potassium sodium Mix tart ~ at and 11g NaOH to 100 ml A. the) in equal parts of the volume, add 5% glucose solution and place in a water bath. First of all, finely dispersed yellow CU 20 is produced, which turns into coarser crystalline red Cu 20 when heated. In an alkaline medium, aldehydes are selectively oxidized to acids, the Cu-l-ons being reduced to Cu-ons and finally precipitating as red Ou 20. In analysis, this characteristic precipitation of red Cu, the color of which comes from a charge-transfer 20 band, is used in the Fehling's sample to detect reducing agents such as aldehydes, hydroxylamine, hydrazine or sugar in human urine, see slide V. Rule that the highest positive oxidation state of an element is equal to its group number; This rule of thumb applies to the elements of the alkali group - they occur at most in the positive monovalent state. However, the oxidation state +11 is found in copper and the oxidation state +111 in strongly electronegative partners, both of which are only conceivable due to the participation of the 3d orbitals in the formation of the compound. and to be attributed to the already mentioned low energy difference between the 4s and 3d orbitals ~ i ~ d. If one compares the second onization energies of copper and alkali metals with each other (see slide V), it makes sense why the alkali metals normally do not exceed the monovalent oxidation level q. Copper in oxidation state +11 has the electron configuration 3d 9 (see slide V) and is the most stable form in aqueous solution (see above). The best-known Cu-1 compound is the blue, water-containing sulfate, CuS0 4 '5 H 20, the structure of which in the crystallized state is imagined as follows: four water molecules are arranged in a planar square around the central copper ion, while two oxygen atoms of the SO- tetrahedron are in an axial position to the central ion; the fifth water molecule is tebrahedral bound by hydrogen bridges in the crystal lattice (see slide V). The water-containing sulfate can be practically concentrated to white, anhydrous CuS0 4 either by heating or the middle. H 2S04 drainage: Test 6: Representation of anhydrous CuS0 4 in a test tube 30 X 200 mm) CuS0 4 '5 H 20 with conc. Move H. 2S04 A transition from blue, water-containing to white, anhydrous CuS0 4 is recorded; the water of crystallization is released in three stages when heated (see slide V). The conc. As a result of its great affinity for water, H changes into the 2S04 hydrate H 2S04.x H 20. The anhydrous, colorless CuS0 4 easily takes water again under blue coloration on Exot her me reaction J, so that anhydrous CuS0 4 is used to detect smaller amounts of water: Experiment 7: CuS04 spatula tip anhydrous CuS0 4 to anhydrous and anhydrous

Give 5 - 7 - half a liter in methanol. The blue colouration in the latter case is due to water uptake. That the color of aqueous Cu-salt solutions does not depend on the anion, but is solely due to the structure of the cation, shows the fact that Various salt solutions without Cu-lones in A. de sto show the same color: Trial 8: 0.5 m CuC 2, 0.5 m CUN0 3) 2 '0.5 m CuS0 4 In order to be able to clarify this behavior, I would like to elaborate on: The atomic and ionic radii of copper are smaller than those of the alkali metals, so that the electrons are built into an inner d-shell where the The effect of the nuclear charge is greater than when it is built into an outer shell. This relatively small size of the ions and a high positive charge of the core result in a high charge density which, together with the availability of the d orbitals, leads to a great tendency towards the formation of stable complexes caused. By comparison, only a relatively small number of compl exes of the alkali metals is known - the solvaths of the cations, the structure of which is still not known exactly; Cryptate complexes of the cations with synthetic polyethers (see slide V). m Cu 2 + o n are initially the d orbits of the same energy, they are degenerate; when the metal ion starts to form a complex, the orbitals are no longer equi val ent and the degeneracy is completely undone. In the octahedral complex, the five ene rgally equal d states are split into two groups (see slide X). I do not want to go into the cause and the exact mechanism of this split here. Now, however, the electronic configuration 3d 9 of the Cu 2 + leads to an effect that / according to the J ahn-Telle r theorem, to which I just don't come any closer n ~ e he n would like the regular octahedron to merge into a distorted octahedron and thus into a stable state; This results in a planar arrangement of four short Cu ligand bonds and two long transients. In the extreme case, this distortion of the structure leads to a square-planar splitting of the d orbitals; The Cu 2 + on is thus assigned, depending on the ligand field, a splitting between octahedral and square-planar field (see Folio X). Now for the color: Imagine that the irradiation is not with visible light to raise it, e.g. in the octahedral field), a d electron from c of the lower-energy group into a gap in the higher-energy group takes place - as a result of this, an absorption band results in the visible area of ​​the Spe ctrums, so that the complex appears in color. From the absorption spectrum one can generally infer the strength and type of energy splitting of the electrons and the spatial structure of the complexes, so that l the erection of a so-called.A spectroch emic series becomes possible, in which the league is classified according to the elements of the league and nf eleven are split up (see slide X). ~ The colourfulness of the aqueous Cu 2 + solution is thus due to several factors - a hexaquocomplex is formed with a distorted octahedral structure, which leads to a ligand field splitting, which is ultimately one specific absorption band at 800 nm and so causes the light blue color. i The Bexaqu okompl ex has the property of a caustic acid. Experiment 9: Complex formation with alkali, add CuS0 4 solution with alkali and heat. It separates out curly, voluminous, blue Cu-hydroxide. The Cu 2 + detection in experiment 1 is based on this reaction. The process can be explained by the fact that H 20 molecules of the hydrate shell give off protons through the positive charge of the central ion could n, since these are repelled by the center alion. In addition, "complex particles" are agglomerated under the splitting off of water, so that crosslinking takes place and CuOH) 2 finally precipitates. When heated, the precipitate turns black, with the separation of water and formation of CU 20 (see slide X). When adding conc. NE to CuOH) 2, 3 the precipitate dissolves: Experiment 10: Tet rammincompl ex precipitate from Experiment 9 with co nz. NE 3 in excess. t In addition to the dissolution of the bottom line, there is a shift in color! too deep blue. In principle, the ligand H 2 0 has been gradually replaced by the ligand NB, whereby a tetrammine complex 3 has formed, which in turn has a distorted octahedral structure f - i has four NH Li gan de n in a square-planar arrangement and 3 two H 0 molecules in an axial 2 position; see slide X). The deepening of color is due to the fact that the NE 3 ligands produce a stronger ligand than the H 2 0 molecules, which causes the! Absorbance bands are shifted from 8 00 nm to 600 nm. n attempt.

6 - 9 2 I demonstrated Cu 2 + in this way. The solution of cu-tetrammine hydroxide [? Unh 3) 4] OH) 2 is called Schweizer's reagent and is a very good solvent for cellulose; The viscous solution of the copper cellulose complex is injected into a precipitation bath, in which one can then produce fine cellulose threads for artificial silk: Experiment 1 ~ Schweizer's reagent Dissolve cellulose in Schweizer's reagent a few days beforehand, filter the solution and inject 2N H 2S04 with a pipette . In the copper-cellulose complex, the cellulose is bound to the central copper atom via the hydroxyl groups; the precipitated cellulose has a different fine structure than the native cellulose (see slide X). In order to clear up an apparent contradiction, I have to introduce another complex: In experiment 5 I combined Cu2 + solution and alkali lye without CuOH) 2 precipitating, as happened in experiment 9: Experiment 12: Tartrate complex a) 0, 5 ml 0.5 m CuS0 4 5 H ml A.dest. + 5ml 11% NaOn b) 0.5 ml 0.5 M Cus0 4 5 H ml 0.5 M tartaric acid + 10 ml 11% NaOH in case a) precipitate of CUOH) 2 'in case b) no precipitate. A copper tartrate complex was formed which is soluble in aqueous solution and due to which the Cu 2+ ion concentration is no longer sufficient for the formation of a CuOH) 2 precipitate. This soluble copper tartrate complex is used in the Fehling's test because the aldehydes must be oxidized in an alkaline medium. See slide X) Finally, I would like to demonstrate that Cu compounds can also occur in the form of complexes: Experiment 13: Add a little 1) KCN solution to tetracyano complex CuS0 solution, heat 4 and then add KCN solution in excess Deduction! 1 !!) At first a brownish-yellow precipitate develops - 'of CueN) 2' which on heating with elimination of cyan 11) turns into white CuCN, which dissolves in excess cyanide to form the colorless tetracyano complex. Due to the fading of an electron gap in the d orbital, this complex is inevitably colorless (see slide X) The complex is assigned a tetrahedral structure, which results in a tetrahedral ligand field splitting. In summary, the following can be stated: Despite the agreement of the electron shell in the ground state established in the introduction, copper shows no parallels in chemical behavior to the alkali metals - firstly, copper proves to be a more noble metal, secondly, copper occurs in several positive oxidation states Although the divalent is the most stable in aqueous solution, while the alkali elements only occur in the monovalent stage, and finally the tendency of copper to form complexes is very pronounced in both oxidation stages, whereas the complex chemistry of the alkali metals is negligibly small. ~ r ~ t f ~ J f i i t i r i ..

7 _._ Literature: Christians, H. R., Fundamentals of the general i nen and a nor g. Chemistry, Cotton, Wi l kins on, Inorgani cal Chemistry, Glöckner, Die K o ~ ple xv e rb ind un g, Praxis Chemie Bd 7, Hol leman, Wiberg, Textbook of Inorganic Chemi e, Jander - Blasius, textbook of a na l ytic and nräparat i ve n ano r ganic chemistry, Kemper, Flad t, Chemi e - Köller, Meineke, Pfr iem, Würthwein, Chemie in Expe r imen ten, Master t on, Slowinski, Chemical Principles with Qualiti ve Ana lysis, Mahr, Inorganis che s basic internship, Mortimer, Chemie, Schmi dt, Anorga nische Chemie Vol. 1 and 2, Seel, Ato mbau und Chemical Bond St apf, Chemisc he Schulv ersuch e part 2 cf ~~ "o ~ v \ 9ts ~~~ \ AltQ.,", O \ l \ ilm Gtt4 ~ a ~ U. ~ {~ M:. !! # / 0 S $ W ~ 3cl 9 4-l B / - [lli-ffij-! M-rrr-it -m - lt-s ~~ 3r -! L 3! » --1U-ßH-ill] i ------ Z, -ill] ls -LD: --- A ~ ~ 3 cl.110 ft .. 5 J \ lil-illl-ll] -ill1- ~ - ~ ': t. - [1] ... 5 ~ D-illl ~ 3p 3S o -ffil- [d-ffil a o2 p -lill fd A ~

8 "[l & j..ho ~~~ \ .o" "bwq. \ o ~ '' '' '' \ Jw ~: Cu.: Li N ~ ~ Rb C!» 3. 't; i [lk] 2 ~ "[N ~ 3 ~ A [A \'] '1-5' [~~ 5 ~" [X ~; & ~, Mt.to.t! Ä ~ l . ~ e ", ZAAS \ o." '- e: t. r ~ l ~~ i \ ~ l -... "' - l1i ~ '" .. 6fa \ l \ l: V.tretcn ~ b ~ lu4l JA. l-töklt .. -, cl \ ~ lt ~ k \

9 1, al ~ u. ~~~~ 1t ~~ VOll \. ",: e-il. ~ -Type Fe / Fi ḷ + A. ~ ocle: Fe. ~ Fe u + le- ~ c" "1. +. ko." "ojc: le- .. C ': f . -40 u ... '! O + Co 4 ~ 4 ~ O.11V "- O, r. +" J \ V., l' Me. ~ ~ ,, - + + "..e-- L , ..- Li + - -3,0 '"[V] ~ ~ e K ki c-- t e- -2.n N, ~ Ne: - -11 ~ + e ~ Felt- Fe r- .. le - - O.ltlf. Hl. ~ Lh t le- to + .le- i- o. ~ 5 '".--. t ~ _ 14 A \ r == & A ..,. ~ e ~ lls Sc. ~~ t \ lf'J: .. {o, ~ C. ~ i «at., ... \. 0: ls ~ 5-c. ~ "Jl .. h.e \ .l!» t "k tia; ll & ~ i ~ l, f_1" i! ji .. ~ -t '.. ~ ,, "al. ~ "J ~; Jl ASO- l'-c, 113 \ 1 -" & o.e. ~~ M..1..Jt ... -iif. "R- ~~~", ~ K ~ [A \ r] 4-1 ~ "? A

10 O}, ~ lO \ l \ ~ s ~ "'te ...: OXicl.Q \ iO \ 4. \ s" "' l + ;;:, t ~ ho \ l \ ~", b ", t '\ "' ,, ~~ o ~ '.- A ~ 1 lt rro 3tp' <1. 10 llilillljhldoo 2 .. Jo. '~ Uu .. ~ fO ~ \. L .: tl1ho ~~ lo' - '. Rt' "~ ko'ot.: I ~~ 0. \) ~ c" ' . f ~ l. a \) D, ~ t "'Ofn" "C> \' \, RA '-' \ ~ ... 0.52. ~ .. 0.11 V -4-T 1 .. ~ u ~" 1- ~ 1 ~} ~ 11 ll. i H ~ O ~ "" 10 ~ l "} -... ~ O ~ - + 3H + 1+ ~ O ​​- 1. ~ + R-. + 't 0" ~' l- ~ "'~ _ ~! = k. \ j .. \ ",, ~! +1 ~ o .. luo + Rc, L. ~ 'oh CMo A ", ~~ c. ~ ,,' U, \ Gc, ~ \ alt ~ i [~ HP ~ 50,.] - Hp ~ 50 5 HO P ~~ o ~ "O ... L" .0 2 &. Bt. ",. ~~ O .. 3lt.L0 ~ ~~ O ..- A ~ o .. lh.l0 ~~ '~ "lo ~ ~ 50 ..." ~ O ~~ ", ~ v" uako \ mr ~ i ~~ \ 10 '"M1t -'-'- ~~" ,, Rb o N G.' rb obl ~ ~ 17U \ c. \ C.

11 ~ J !. -_._ - _._-.- .. / '-. J ~~ l. + - F ... be. : ~~ ~ b ~ "-'1 ~ [G .. KLO) J ~ \, ~ please \ ~~: lriolu b ~ cflr tlt ~ \ \ J..1ft", lci \ l \ ": - t- lo lt- ~ <; 00. ~ "") 5P.9t \ w, \ t_iS ~ c. Rtik: + CWbi ~ \ <& .;

12 c, ~~~ S ~!: A ~ [~ ~ "3)",] OH) l i! ~ / k. ~, yom.ol ~ fc..t: G.l ~ f ~ +1 ~ N) l ~:. ~ r -4) l \. \ CN -l ~ N) l.1 ~! t t ~ blo ~ o H 'ÖM' 0 l .. ,, \. ~ 0 '/, x loh H 0 U.i ~ - \) ~ l. e: c. -CO-c. -ex e - + [~ N) ~] 3- t- - "<