Sciencemadness Discussion Board

When to use excess of reagents?

Turner - 25-12-2013 at 15:59

Using 4Ag+6HNO3 --> 4AgNO3 + 3H2O + NO2 + NO

54g of Silver needs 48g of HNO3, or 38ml approx. I dilute it to 55% and use hot for the reaction. Would 80ml of 55% Nitric Acid really fully react with the silver or would the products dilute the remaining acid towards the end of the reaction to require and excess of nitric to fully react the silver?

Thanks

Random - 26-12-2013 at 00:18

It needs 6 moles of nitric acid, 6/4 times molar amount the silver.

[Edited on 26-12-2013 by Random]

HeYBrO - 26-12-2013 at 03:48


54 g Ag = 0.504 mol

n(HNO3)/n(Ag) = 6/4 x 0.504 = 0.756 mol

M(HNO3)= 63.01 g mol−1

m(HNO3)= 0.756 x 63.01 = 47.63 g

55/100 x 80ml = 44 g HNO3, so, by my calculations ( oh god i hope they're right ) you're not even in excess. And no the products won't dilute it so much that it won't work.

edit fixed typo

[Edited on 26-12-2013 by HeYBrO]

blogfast25 - 26-12-2013 at 07:19

To dissolve silver (not to mention some of the crap you often find in it) use at least 20 - 30 % excess acid, nitric with respect to stoichiometric requirements)

Why not dilute it to about 70 %, instead of 55 %? Why the 55? Magic number? ;)

Turner - 26-12-2013 at 08:00

90% won't react, I dilute it to around 45-65% and that works well.

47g of HNO3 is the stoic amount I need. 47/1.5 (density of HNO3) = 31.5ml of 100%, or 60 ml of 55% Nitric Acid.

I'll use 30% more Nitric acid so 80ml of 55% is what I need?

blogfast25 - 26-12-2013 at 08:12

Yes, 80 ml of 55 w% HNO3 is just above 30 % in excess, with respect to stoichiometry. 41 ml pure HNO3 plus 33 ml water, is exactly 30 % excess.

I've used 70 w% HNO3 straight from the jerry and it worked well.

[Edited on 26-12-2013 by blogfast25]

AJKOER - 26-12-2013 at 11:44

Yes, an excess of a reagent is advisable in some reaction to move the reaction to the right. In general, however, I would not recommend employing an excess without further research.

One reason relates to undesirable reactions between products and the excess reagent. For example, the action of Oxalic acid on an aqueous chlorate can form Chloric acid and then ClO2. Here is extract from a prior thread (see http://www.sciencemadness.org/talk/viewthread.php?tid=20109 )
Quote: Originally posted by chemretd  
A "safe" way of generating chlorine dioxide, mixed with an equal volume of carbon dioxide was published (I can't remember where) as warming a mixture of potassium chlorate and oxalic acid on a water- bath. In my spotty youth (MANY years ago), I tried this with sodium chlorate. It generated a green gas alright, but with a very alarming LOUD crackling noise. I suspect that this safe method is a serious accident waiting to happen.


Here is the reaction equations and source for Chemretd comment:

2 NaClO3 + H2C2O4 --> 2 HClO3 + Na2C2O4 (s)

and with more excess Oxalic acid:

2 HClO3 + H2C2O4 --> 2 ClO2 + 2 H2O + 2 CO2

so the net reaction is:

2 NaClO3 + 2 H2C2O4 --> 2 ClO2 (g) + 2 H2O + 2 CO2 (g) + Na2C2O4 (s)

See http://books.google.com/books?id=6wUmteTIc18C&pg=PA334&a...

as was reported by Chemretd. This reaction is surprising to some and may precipitate an accident.

Another example of this possible dual role of Oxalic acid was reported by Watt's (see page 649 http://books.google.com/books?pg=PA649&lpg=PA649&sig... ) as the reduction of HOCl by Oxalic acid. So the addition of very dilute Oxalic acid (and not concentrated H2C2O4 which can produce a very violent reaction with a hypochlorite and an explosion with a dry hypochlorite) to even pure NaOCl would first form HOCl, which could subsequently reacts with the excess H2C2O4 to form HCl and CO2 (and then some Cl2). Speculated reaction sequence:

H2C2O4 + 2 NaOCl + xH2O --> 2 HOCl + xH2O + Na2C2O4 (s)

2 HOCl + 2 H2C2O4 --> 2 HCl + 4 CO2 (g) + 2 H2O

2 HCl + 2 HOCl <----> 2 Cl2 + 2 H2O

Another example, the action of a dilute solution of Citric acid on NaOCl forms HOCl. However, in the event an excess of Citric acid (and even to some extent without an excess), the Hypochlorous acid formed apparently actively attacks the Citric acid liberating the scent of newly created Chloroform. Yet another example occurs when trying to create highly concentrated H2SO4 by the action of Oxalic acid on a sulfate. Using an excess of H2C2O4 to drive the reaction together with heating indeed results in a very high strength of sulfuric acid. However, at a certain point, the excess H2C2O4 is violently (and usually unexpectedly) decomposed (into CO2, CO and water vapor, see http://escholarship.org/uc/item/1s96t1nf#page-5 ) by the now the concentrated H2SO4. In other words, a large violent ejection of highly corrosive acid occurs (see comments by Formatik at http://www.sciencemadness.org/talk/viewthread.php?tid=18963&... ).

Predicting these, at times, dangerous secondary reaction with an excess of a reagent is not easy, but doing some research can spare one potentially a lot of pain. When unclear, try to avoid employing an excess of the reagent.

[Edited on 26-12-2013 by AJKOER]

Random - 29-12-2013 at 17:24

Speaking of safety and silver, mixing silver salts with ammonia could end up in unexpected explosion after it's left standing for some time.

[Edited on 30-12-2013 by Random]

AJKOER - 30-12-2013 at 14:12

In my opinion, some points on Ag3N formation:

First, Silver oxide dissolves in an excess of aqueous NH3:

Ag2O(s) + 4NH3 + 4H2O <--> 2 [Ag(NH3)2]OH + 3 H2O [1]

Note, I have written this reaction as reversible (see as a source, "Second year college chemistry" by William Henry Chapin, page 255):

"As might be expected, the silver-ammonium complex dissociates slightly into its constituents as indicated by the equation

Ag(NH3)2+ <=> Ag+ + 2 NH3

This is a reversible reaction, very much like the ionization of a very weak acid or base."

So:

2 Ag(NH3)2OH <---> 2 AgOH + 4 NH3

But, 2 AgOH --> Ag2O (s) + H2O

This reaction would move to the right with the loss of ammonia and/or water with, for instance, the evaporation over time from an open container of aqueous [Ag(NH3)2]OH.

Next, the equation for the creation of Silver nitride can be written for one mole of Ag2O as:

2/3 NH3 + Ag2O (s) --Energy--> 2/3 Ag3N (s) + H2O

where interestingly, some but less than the original quantity of ammonia is required together with an energy source to produce the problematic Ag3N. Again, loss of water (but not of ammonia!) helps to move this reaction to the right.

A reference on the formation of Silver nitride per "Comprehensive Inorganic Chemistry Series" Ag3N can be prepared by adding Ag2O to concentrated ammonia and letting stand in air, heating in a hot water bath (or, one can precipitate it with the addition of alcohol) or, per other sources, with the addition of KOH or NaOH.

So my take on the involuntary mechanism of Ag3N formation is that it is not so much a case of a loss of ammonia as perhaps loss of water combined with a heat source (or possibly light given the sensitivity of Silver salts to it, or perhaps the general propensity of Silver salts to undergo hydrolysis on standing).

As such, relative to the current topic, this is not so much a possible illustration of what could happen with an excess of a reagent, but more so, what happens with the loss of the excess (possibly water and/or ammonia, in the current example) and the addition of an energy source (for example, heat created with the action of NaOH added to water).

[Edited on 31-12-2013 by AJKOER]