Grassroot Posted July 29, 2011 Report Share Posted July 29, 2011 (edited) I am a rising senior and I am reviewing the topics for IB.I had a question about chapter3 Periodicity (according to my book it is chapter 3 IBID 2007) regarding complex ions.It is stated that it is formed by the unfilled p orbital that the transition metals poccess. But how does the isomers form???According to my understanding that the ligands are the same... then how come it is possible to have isomers in complex ions?BTW, how is my textbook version? I had another confusion in P88, why the iron presented in the complex ion has a oxidation number of 1 or 2, in the contrary of 2 or 3 as general concensus??One last question, can lab reports be self-made i.e self-written and self-evaluated? because currently I do not have a teacher. Even if I will in next semester he will be a new one.I have another question after viewing the upcoming text. The textbook stated thay Chromium has an oxidation state of +3, however according to electron configuration it is not possible. Some problem exists to manganese whereas oxidation state of +4&+6 do not make sense....Hopefully I can get help, thanks! Edited July 31, 2011 by Grassroot Reply Link to post Share on other sites More sharing options...
loltim Posted July 29, 2011 Report Share Posted July 29, 2011 Iron, as a transition metal characteristic, has a variable oxidation state Iron (ii) or iron (iii). Transition metals also have incomplete d-orbitals, and splitting of these in a complex ion causes colouration. Splitting of the d-orbital is caused by ligands(a species that uses a lone pair to form a dative covalent bond with a metal ion). I don't know about isomers though sorry. Hope that helps Reply Link to post Share on other sites More sharing options...
Drake Glau Posted July 29, 2011 Report Share Posted July 29, 2011 I am a rising senior and I am reviewing the topics for IB. I had a question about chapter3 Periodicity (according to my book it is chapter 3 IBID 2007) regarding complex ions. It is stated that it is formed by the unfilled p orbital that the transition metals poccess. But how does the isomers form??? According to my understanding that the ligands are the same... then how come it is possible to have isomers in complex ions? BTW, how is my textbook version? I had another confusion in P88, why the iron presented in the complex ion has a oxidation number of 1 or 2, in the contrary of 2 or 3 as general concensus?? One last question, can lab reports be self-made i.e self-written and self-evaluated? because currently I do not have a teacher. Even if I will in next semester he will be a new one. Easy question first: No you can't grade your own labs. You can write your own though Isomer just means that the molecule has some sort of different structure to it. This could be caused by a lot of things. The environment that the complex ion is in could contribute to this, the charge of the central atom (iron in your case I'm guessing) could contribute to this as well. Structures of molecules is based off of VESPR theory which is all about electrons and how they repel each other thus forming the structures. If the iron is +3 instead of +2 then it has one less electron contributing to the repulsion so it would change the shape of the overall molecule. Same in reverse, +2 would have one more electron than +3. Complex ions to being with are formed using the p-orbital because it's the outer most orbital so you would need to figure out the electron configuration of the central atom to see if the p-orbital is effected. Sometimes if you have let's say a +2 transition metal ion with S2 D5 and P3 (I just made that up by the way) and then you have a +3, due to the tendancy of the D orbital to stay empty, full, or half full, it is likely to simply lose an S electron giving you S1 D5 P3 which shows that the p-orbital was unchanged. To be honest, I don't remember needing to know anything about isomers when it came to complex ions and all of what I just said came through some reasoning in my head so please don't take this is perfectly accurate Reply Link to post Share on other sites More sharing options...
Grassroot Posted July 30, 2011 Author Report Share Posted July 30, 2011 I am a rising senior and I am reviewing the topics for IB. I had a question about chapter3 Periodicity (according to my book it is chapter 3 IBID 2007) regarding complex ions. It is stated that it is formed by the unfilled p orbital that the transition metals poccess. But how does the isomers form??? According to my understanding that the ligands are the same... then how come it is possible to have isomers in complex ions? BTW, how is my textbook version? I had another confusion in P88, why the iron presented in the complex ion has a oxidation number of 1 or 2, in the contrary of 2 or 3 as general concensus?? One last question, can lab reports be self-made i.e self-written and self-evaluated? because currently I do not have a teacher. Even if I will in next semester he will be a new one. Easy question first: No you can't grade your own labs. You can write your own though Isomer just means that the molecule has some sort of different structure to it. This could be caused by a lot of things. The environment that the complex ion is in could contribute to this, the charge of the central atom (iron in your case I'm guessing) could contribute to this as well. Structures of molecules is based off of VESPR theory which is all about electrons and how they repel each other thus forming the structures. If the iron is +3 instead of +2 then it has one less electron contributing to the repulsion so it would change the shape of the overall molecule. Same in reverse, +2 would have one more electron than +3. Complex ions to being with are formed using the p-orbital because it's the outer most orbital so you would need to figure out the electron configuration of the central atom to see if the p-orbital is effected. Sometimes if you have let's say a +2 transition metal ion with S2 D5 and P3 (I just made that up by the way) and then you have a +3, due to the tendancy of the D orbital to stay empty, full, or half full, it is likely to simply lose an S electron giving you S1 D5 P3 which shows that the p-orbital was unchanged. To be honest, I don't remember needing to know anything about isomers when it came to complex ions and all of what I just said came through some reasoning in my head so please don't take this is perfectly accurate Thanks for both questions.... I did not mean to grade my report. I will just select some of the activities in the activity list and write reports for them. Reply Link to post Share on other sites More sharing options...
Grassroot Posted July 31, 2011 Author Report Share Posted July 31, 2011 (edited) Updated. Please see the red partI have another question after viewing the upcoming text. The textbook stated thay Chromium has an oxidation state of +3, however according to electron configuration it is not possible. Some problem exists to manganese whereas oxidation state of +4&+6 do not make sense....Hopefully I can get help, thanks! Edited August 1, 2011 by Keel Added quote Reply Link to post Share on other sites More sharing options...
Sammie Backman Posted July 31, 2011 Report Share Posted July 31, 2011 (edited) I have another question after viewing the upcoming text. The textbook stated thay Chromium has an oxidation state of +3, however according to electron configuration it is not possible. Some problem exists to manganese whereas oxidation state of +4&+6 do not make sense....Hopefully I can get help, thanks!Chromium can have any oxidation state between -2 and +6, the most common ones being +3 and +6. While you can sometimes deduce the oxidation state from the electron configuration, that is generally not the case with transition metals. Edited July 31, 2011 by Sammie Backman Reply Link to post Share on other sites More sharing options...
Drake Glau Posted August 1, 2011 Report Share Posted August 1, 2011 Updated. Please see the red part I have another question after viewing the upcoming text. The textbook stated thay Chromium has an oxidation state of +3, however according to electron configuration it is not possible. Some problem exists to manganese whereas oxidation state of +4&+6 do not make sense.... Hopefully I can get help, thanks! 1s2 2s2 2p6 3s2 3p6 4s2 3d4 This is just looking at it, but the d orbital likes to be half full so it'll take one from the s orbital "above" it making it: 1s2 2s2 2p6 3s2 3p6 4s1 3d5 Now this is charge of 0 and you have 1 electron in the 4th shell and 6 more in the outer p shell (each energy level goes s>d>p making p in the outside if you want to look at it like that). This means that the one electron from the 4s can go away giving you +1, then it can lose pretty much all of it's 3p electrons which can make it go up to +7, D orbital doesn't really ever change, it likes to be empty, half full, or full, and as Sammie stated it only goes from -2 to +6 and that's probably because the 7th ionization energy is just WAY too high to make it +7. For -1 and -2, well an electron can fill the 4s orbital, I'm not really sure on where the 2nd one would fit honestly Reply Link to post Share on other sites More sharing options...
daxx Posted September 17, 2011 Report Share Posted September 17, 2011 There's several types of isomerism for inorganic chemistry. The most important is stereoisomerism (which you study for the Organic Topic). There's two types of stereoisomerism: geometric, and optical. Reply Link to post Share on other sites More sharing options...
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