I would like to ask your thoughts about a question that I am trying to explain to a group of students whom I am helping to prepare for the Chemistry Olympiad.

Right answer: B

My explanation (I am trying to link it to periodic trends):

Reduction is the inverse of ionization energy.

In ionization energy, we have M(g) → M⁺(g) + e; the reduction runs backward: M⁺(g) + e⁻ → M(g).

The half-reaction in this question is M²⁺(aq) + 2e⁻ → M(s). Strip away the “aq” and “s” and look at the electrons: the cation M²⁺ gains two electrons to go from M²⁺ to neutral M (reduction process). It is the reverse of the element M losing two electrons to go from neutral M to M²⁺ (ionization energy).

Therefore, an element with a high ionization energy (IE) is hard to pull electrons off; this same element has a high reduction potential (E°), meaning electrons want to come back on. So, the element with the highest ionization energy has the higher reduction potential.

Now, let's analyze how ionization energy varies among Ni, Pd, and Pt.

In general, ionization energy increases from bottom to top; however, because we are dealing with transition metals, we must account for the anomalous behavior of d and f orbitals.

The reason why ionization energy increases going up a group is because each added shell brings more shielding while Zeff stays roughly the same. By that logic, we’d expect Ni to have the highest IE of the three (fewest shells, least shielding) and Pt to have the lowest (most shells, most shielding), with Pd in between.

That’s not what happens. Ni is the lowest; Pd and Pt are both higher than Ni. The reason is that between each pair, a subshell fills in that shields unusually poorly (d and f orbitals).

Between Ni and Pd, Ni’s 3d subshell is filling right before Pd. d electrons already shield worse than s or p electrons, so Pd’s valence electrons feel more of the nuclear charge.

Between Pd and Pt, the effect repeats, but worse. Before you reach Pt, all 14 lanthanides have to fill their 4f subshell, and f electrons shield even more poorly than d electrons. This lanthanide contraction raises Pt’s Zeff further still, which is why Pt’s IE stays elevated instead of dropping the way you’d expect from adding an entire extra shell.

Since IE and E° are linked, Ni’s low IE becomes Ni’s low E° (least willing to be reduced), and Pd and Pt’s anomalously high IE values become their anomalously high E° values.

Group 4 elements on the @compoundchem@mstdn.social infographic

Carbon, Silicon etc,

https://www.compoundchem.com/2013/12/29/group-4/

Whereas Wikipedia

https://en.wikipedia.org/wiki/Group_4_element

Suggests group 4 is Titanium, Zirconium, Most periodic tables are group 1,2 on the left, then there is a jump over the transition to group 3, Boron, Aluminum, Transition element groups are numbered IIIB, IVB (which wikipedia lists as group 4) 2nd group of the transition elements.

So which is correct here, or does it depend on the periodic table you are looking at?

Paul

this can mean 2 things

1. heating a colder object causes more randomness or

2. supplying same amount of heat slowly causes more randomness

I think the 2nd one because ∆S = qrev/T

"4. Carbon molecule (C 2 ): The electronic configuration of carbon is 1s2 2s2 2p2. There are 12 electrons in C2. The electronic configuration of C2 molecule, therefore, is

The bond order of C2 is ½ (8 – 4) = 2 and C 2 should be diamagnetic. Diamagnetic C2 molecules have indeed been detected in vapour phase. It is important to note that double bond in C 2 consists of both pi bonds because of the presence of four electrons in two pi molecular orbitals. In most of the other molecules a double bond is made up of a sigma bond and a pi bond. In a similar fashion the bonding in N2 molecule can be discussed.

O2

5. Oxygen molecule (O 2 ): The electronic configuration of oxygen atom is 1s2 2s2 2p4. Each oxygen atom has 8 electrons, hence, in O 2 molecule there are 16 electrons. The electronic configuration of O 2 molecule, therefore, is

From the electronic configuration of O2 molecule it is clear that ten electrons are present in bonding molecular orbitals and six electrons are present in antibonding molecular orbitals. Its bond order, therefore, is Bond order = [Nb – Na] = [10 – 6] =2. So in oxygen molecule, atoms are held by a double bond. Moreover, it may be noted that it contains two unpaired electrons in π ∗2px and π ∗2py molecular orbitals, therefore, O 2 molecule should be Paramagnetic, a p r e d i c t i o n t h a t c o r r e s p o n d s t o experimental observation. In this way, the theory successfully explains the paramagnetic nature of oxygen. Similarly, the electronic configurations of other homonuclear diatomic molecules of the second row of the periodic table can be written. In Fig. 4.21 are given the molecular orbital occupancy and molecular properties for B2 through Ne2. The sequence of MOs and their electron population are shown. The bond energy, bond length, bond order, magnetic properties and valence electron configuration appear below the orbital diagrams. "

Question

1. Seeing it in atomic orbital perspective the molecules electronic config should be KK 2s2 2p4 ie 1 filled and 2 single filled orbitals. But according to the book there is no single filled molecular bond or anti-bond so the molecule is diamagnetic. Is the atomic bonding and molecular orbitals totally unrelated ?

2. When Pie bond is formed doesn't both atoms' orbitals needs to be acting i.e. Px of Atmos A acts with Px of Atom B ? There Px of 1 Atom and Py of other Atmos forms 1 molecular bond each ?

O2 Question

1. Why does O2 form 1 bond and 1 antibond ?

2. O2 form a double bond and complete the valance but the book shows 3 bonds but says double..

3. Explain "Moreover, it may be noted that it contains two unpaired electrons in π ∗2px and π ∗2py molecular orbitals"

This O3 Lewis Diagram looks weird -

why does atom with 8 electrons positive,

why is atom with 4 electrons neutral,

why is atom with 6 electrons negative

In CO3 -2 Why there is a -2 charge ?

What is happening in HNO3 ?

N is positive that means it is 5 electrons short of Nobel config.

Why there is a negative charge on bottom oxygen ?

Are these Lewis Diagram possible or covalent and electrovalent bonds ?

[NO2]-

is there 1 electron extra on oxygen ?

[NO2]-

I have tried an expedient of first air tightening the tin and then adding an amount of acetone that I happened to have.

Can I hope that ,in time the resin (which has not been in contact with the hardener but hardened on its own) will combine the the acetone and become useable again?

I don't mind waiting up to a year as I can purchase a new tin of resin to use for now...

Periodicity in valence - among representative elements, the valence is either equal to the number of electrons in the outermost orbitals or 8 - minus this number

What is 8 ? There can be a max of 6 electroins untill p shell and 3 in s shell

I had assumed the colour would be due to copper and was a bit shocked to find the formula is: 

Na₇Ca(Al₆Si₆O₂₄)(SO₄)(S₃).H₂O , i.e no Cu in sight!  

Turns out the clue is in the S₃. This is present in the form of the trisulphide radical anion, S₃⁻•,  a curious species that breaks the school-level rules for stability and bonding - and so is automatically interesting to me. This radical anion apparently has an absorption band in the orange region of the visible spectrum, and thus reflects mainly blue light. I haven't managed to find a molecular orbital diagram for it on the internet but presume the odd electron may be in a relatively high energy orbital, from which it can be promoted to another one that is only slightly higher, i.e. with relatively little energy and this will be why it absorbs in the visible rather than the UV. S₃ itself is regarded as having a similar bonding scheme to ozone, i.e. the centre atom sp2 hybridised with one lone pair, but I presume the extra odd electron must go into either an sp2 hybrid antibonding orbital or else something involving participation by d orbitals, which obviously is possible in sulphur, unlike oxygen. 

I suppose one should expect this ion to be paramagnetic. From what I have found on the internet this radical anion has some applications in synthesis of organosulphur compounds. There seems to be a guy  called Tristram Chivers at Calgary (now retired and emeritus) who has done a lot of work on it. 

If anyone knows more about this anion I'd be interested to learn more about it, especially the bonding and electronic structure. 

I have some tartaric acid crystal. Can anyone help me determine if my calculations are correct for making the rochelle salt? I figured for every mole in weight of the tartaric, I add each a mole of sodium hydroxide and potassium hydroxide, and stir in DI water. Maybe check for pH. Does that make sense? 

Thanks for your advice and expertise.

I'm trying to make a tris HCL buffer of 1.2M with 10.5 pH. My preferred volume is 7.5 mL, and I know the amount of tris base (1.09g ) and HCl (3uLwith HCL being 37.2%(12.1M) )I need, but I don't know the amount of deionized water I need to dissolve the tris base in. Is it possible to use a recipe for 1L and add 800mL of deionized water or does that change the outcome for the tris HCL buffer?

total alkalinity is 190ppm and hardness is 120ppm , my thinking is if I increase hardness by using gypsum I'll possibly increase overall Ph as it falls in level with alkalinity , so I should reduce Alkalinity by around 70ppm .

My thoughts were to use muriatic acid but unsure of a. the volume of acid required b. how much free chlorine would be released from the breakdown as there is aquatic life present.

I'm thankful in advance to any advice !

What forms of iron might typically be found in a city’s water supply? In particular, what form of iron could cause my whole-house water filter, time after time over 2+ years, to turn a dark “rust” brown color within 2 to 3 weeks?  Please see attached photo.

Our Pasco County, FL water, of course, meets the EPA limit of 0.3 mg/L iron, so what “other iron”, if iron is what I’m seeing, is causing this rapid? discoloration of my filter elements?  

My store-bought test strips indicate iron at less than 0.3 mg/L, but are these strips only picking up “dissolved” iron, ie iron ions.  If there are minute “flakes” of rust from iron pipes in the county’s water supply system, I’m guessing my test strips will not detect that iron, unless ions of iron migrate from the iron particles into the surrounding water.

Also, I have been told that if phosphates are used in the water treatment process, which is apparently not uncommon, that this can cause the rust colored “stuff” that my filter is capturing.

Obviously, I’m shooting in the dark here, but any explanations for what I’m seeing here will be much appreciated.

ABR2955 Dr AT Odularu Oral (1).pdf

Have some chromic acid and it is very damp as it is very hygroscopic. Does anyone know of a way to dry it out? Maybe over conc sulfuric acid or anhydrous calcium chloride? Could not find anything on the internet. 

Thanks.

So the water has a pH of 6.35. Which is the pKa value of H₂CO₃/HCO₃^-. So in solution at this pH there should be 50% H₂CO₃, 50% HCO₃^- and 0% CO₃^2-.

So in general they now start calculating stuff with the electroneutrality equation which is:

H⁺= HCO₃^- + 2*CO₃^2- + OH^- .

Lets assume that the total of carbonates species is 0.03 mol/L at a pH of 6.35. Than 50%*0,03 = 0.015 mol/L of H₂CO₃ and 0.015 mol/L of HCO₃^- in solution. There is also kw/10^-6.35=10^-7.65 mol/L OH^- present in solution.

So filling this in the electroneutrality equation: H⁺= 0.015 + 10^-7.65 = 0.015 mol/L.

So this is a bit confusing. Since pH is 6,35 it should be H⁺= 10^-6.35 mol/L. Not 0.015 mol/L. 

My question is. What am I missing here. Can somebody help clear the confusion.

- butyric acid

- isovaleric acid

- propionic acid

I am not interested in any ppm values or thelike, but only in the values for the maximal possible concentrations of the acids in solutions so that they generally can't cause any harm any more if you inhale them (e. g. if the solution is then sprayed into the air or on a piece of cloth to be smelled...).

How do I get to know these values anyways? I couldn't find anything on this online.

Possible applications for this are manifold: creating prank sprays, scent generators, tests with sniffing sticks...

Nacl + H20 + current = naoh+HCl 

Sodium chloride+ water + current = sodium hydroxide + hydrochloric acid 

To my knowledge these are the ways hypochlorous can be made.

• salt + water + electrolysis= some hocl % some chorite and? 
• salt + water + vinegar + electrolysis= a higher hocl %, chorite, and? 
• salt + water + hcl + electrolysis= would this yield an even higher percent of hocl? 
• chlorine + h20 = hocl and hcl  << can make stable hocl by taking chlorine gas and dissolving it into a sodium hydroxide solution?? >> 
• dichlorine monoxide + water = hocl 

1. Which formula would yield the highest amount of HOCL? 

2. Which formula would produce the most stable HOCL?

Also, if your ph and ppm read fine, but your solution reeks of bleach-what happened? (Assuming you used salt, water, vinegar + electrolysis) (correct ratio & run time to a liter of water)

I would pay someone for their time. I'm not looking for a free hand out, but I really need help understanding all of this better. 

Sorry for the complete butchering of the formulas. 😐  Is there anywhere that I can go that someone would be able to help me in person or over the phone? 

Thank you in advance!

Sincerely, a lost and confused individual.