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Submitted by chem5th on Thu, 12/27/2007 - 17:17

Here are some homework questions that I have attempted, but could not figure out / not sure if it is correct.

1. Use an equation to illustrate the fact that O2- is soluble in water.

Is it      O2-    + H2O    >>>  OH    + OH-      ???

2. Write balanced (1) formula unit, (2) total ionic, and (3) net ionic equation for the reaction between the acid-base pairs.
H2SO4    +  NH3    >>>>>

Is it    2H+  + SO42-    + NH3  >>>>>    NH4+    +  HSO4-      for both the total ionic and net ionic equations?

3. Which is the stronger acid? NH4+ or NH3

4. The following salts are components of fertilizers. They are made by reacting gaseous NH3 with concentrated solutions of acids. The heat produced by the reactions evaporates most of the water. Write balanced formula unit equations that show the formation of each.
a) (NH4)2HPO4
b) (NH4)3PO4
c) (NH4)2SO4

Much help would be appreciated! Thanks

1. Use an equation to illustrate the fact that O2- is soluble in water.
      O2-   +  H2O     -->     2 OH-(aq)
2. Write balanced (1) formula unit, (2) total ionic, and (3) net ionic equation for the reaction between the acid-base pairs.
H2SO4    +   NH3    >>>>>

Formula unit     H2SO4  +  2NH3   -->   SO4 2-    +   2 NH4+
Ionic:      2H+ SO42-  +  2 NH3    --> 2 NH4+   +  SO42-
Net Ionic:    H+   +   NH3   -->  NH4

3. Which is the stronger acid? NH4+ or NH3
NH4+  reacts with water to form H3O+ while NH3 reacts with water to form OH-.  That makes NH4+ the stronger acid.

4.  All salts can be formed from the neutralization reaction between an acid and a base. 

In a and be the base is NH3  (or NH4OH) and the acid is phosphoric acid (H3PO4).  The different salts in a and b are generated by using different amounts of NH3.  In c the base would be NH3 and the acid would be sulfuric (H2SO4).

Submitted by spock on Thu, 12/27/2007 - 21:01 Permalink

Thanks for the help on the first three problems! But, I still don't understand number four.
Would these equations work?

a)    NH3  +  H2PO4-    --->    (NH4)2HPO4

b)    NH3  +  HPO42-    --->    (NH4)3PO4

c)    NH3    +  HSO4-      --->    (NH4)2SO4

Submitted by chem5th on Fri, 12/28/2007 - 20:31 Permalink

Also, can someone explain to me why HI, HBr, and HCl are stronger than HNO3?

Order of increasing acid strengths
HNO3 < HCl < HBr < HI < HClO4

I understand that for binary acids, the weaker the electronegativity, the stronger the acid for HCl, HBr, and HI.
I also know that for ternary acids such as HNO3 and HClO4-- HClO4 is the stronger acid because it has the higher oxidation state.
But, I just don't know how HNO3 is randomly weaker than HCl, HBr, and HI.

Also, can someone explain why when SO42- reacts as a base with H+, it forms the conjugate acid HSO4-? Isn't it supposed to be H2SO4?

Thanks!

Submitted by chem5th on Sat, 12/29/2007 - 00:52 Permalink

just remember that an strong acid is one that completely dissociates in an aqueous sol. except for sulphuric acid.  or in a term pKa < -1.74 .
also the strength of an acid can be determined by observing its electronegativity like higher the EN of a conjugate base , the less acidic.
Also more positively charged a species is , more acidic.
And also with increasing atomic radius the acidity increases.
like for example if we take the case of HCl and Hl .....both are strong acids, and ionise 100 % in water, to their respective ionic constituents.  but Hl is stronger than HCl. just because the atomic radius of iodine is larger than chlorine atom. so as a result of this the negative charge over the l- anion is dispersed over a large e-1 cloud and its attraction for H+ is not as strong as the same attraction in HCl. .....hope this helps.

Submitted by Chemistry_Guru (not verified) on Sat, 12/29/2007 - 08:14 Permalink

Thanks for answering, but I already understand why HCl, HBr, and HI are in that order. I just don't understand how HNO3 is the weakest of those strong acids listed. Doesn't N have an oxidation state of +5, which is higher than the oxidation states of Cl, Br, and I?

Submitted by chem5th on Sat, 12/29/2007 - 09:45 Permalink

In the first two cases, the different salts can be generated by using the same acid H3PO4 but varying the amount of NH3.
a)    2NH3   +   H3PO4     --->     (NH4)2HPO4 

b)    3NH3   +   H3PO4     --->     (NH4)3PO4

c)    2NH3    +  H2SO4       --->     (NH4)2SO4

Submitted by spock on Sat, 12/29/2007 - 15:37 Permalink

I just don't understand how HNO3 is the weakest of those strong acids listed. Doesn't N have an oxidation state of +5, which is higher than the oxidation states of Cl, Br, and I?

This one sent me back to the books!!!!

Perhaps if we think of the strongest acid is being the one that has the weakest conjugate base it might be easier to see.

Sharma does a nice job of explaining the trend for the halogens by stating that the electon cloud becomes more diffuse as it is spread over a larger volume.  Let's call that concept charge density, as the charge density on a negative ion becomes less dense it becomes less likely to accept a proton.

I think that we can use the same idea when we look at the polyatomic ions.  Cl04- has its negative charge distributed across whole ion (mainly the 4 Oxygen atoms), each oxygen atom has an effective charge of -1/4.  This effective charge is further reduced by the fact that Cl has an oxidation state of +7, which means that it will tend to pull the bonding electrons toward the center of the ion. As a result, the perchlorate ion is not going to be very good at snagging protons which makes it a very poor base, and makes perchloric acid VERY strong.

In the case of  sulfuric acid, the negative charge is again distributed over four O atoms, but this time the S has an oxidation state of +6 (and sulfur atoms have a lower EN than Cl).  This would tend to make the charge density on the oxygen atoms larger, making HSO4- ion a little better as a base (and H2SO4 a little bit weaker as an acid).

Finally we get to NO3-.  Here the negative charge is distributed over only the three oxygens (instead of four as in the others).  N is also present in the +5 oxidation state in the nitrate ion.  Both of these should contribute to a higher effective negative charge on the O atoms in the ions, and thus an increased ability to act as a base (and reducing HNO3's ability to act as an acid).

The above is consistent with the following: 
H2SO4 is stronger than H2SO3  and HNO3 is stronger than HNO2

Submitted by spock on Sat, 12/29/2007 - 16:51 Permalink

Yes, but what about the oxidation states of HCl, HBr, and HI in comparison to HNO3 and H2SO4? Would it affect acid strength?

How would you determine which one is the stronger acid if you are given a binary acid and a ternary acid (ex.  HCl versus H2SO4)?

Submitted by chem5th on Sun, 12/30/2007 - 01:52 Permalink

I found a table that indicates that HClO4 is the strongest followed by HI, HBr, HCl, H2SO4, and HNO3.

I found the following information on Wikopedia (http://en.wikipedia.org/wiki/Strong_acid), but it really doesn't answer your question.  The oxidation states in the ternary acids were important because they affected the negative charge density on the oxygen atoms.  The higher the oxidation state of the central atom, the more the electrons were pulled away from the oxygens thus reducing the charge density.  The oxidation states of the atoms in the binary acids are all -1, but would not act in the same way, so I don't see how they can be compared.

Bottom line is that they are all strong acids, and short of experimentally determining them, I'm not sure how you could accurately predict their relative strengths.

Submitted by spock on Sun, 12/30/2007 - 02:31 Permalink

Thanks for answering the previous questions, I understood them now... but I have some new questions as well.

Identify each reactant and product in the following chemcial reactions as a Bronsted-Lowry acid, a Bronsted-lowry base, or neither. Arrange the species in each reaction as conjugate acid-base pairs.

HCl + NH3 ---> NH4Cl

Would the answer to that particular problem be neither?

Write the chemical equations for the stepwise ionization of citric acid, C3H5O(COOH)3, a triprotic acid.

Label the acids and bases using Lewis theory terminology.
a) NH3 + HBr ---> NH4Br

How can you tell which one is the Lewis acid or Lewis base?

Identify the Lewis acida nd base and the donor and acceptor atoms in each of the following reactions.

6[Cl]- + [Pt]4+ ---> [PtCl6]2-

Is 6[Cl]- the lewis base while [Pt]4+ is the lewis acid or is it the other way around?

Indicate which of the following substances can act as an acid, base, or both according to the Arrhenius theory or the Bronsted-Lowry theory.
a) H2S
b) PO(OH)3
c) H2CaO2
d) ClO3(OH)
e) Sb(OH)3

This is what I got: a) acid  b) base  c) acid  d)base  e) both
I'm not quite sure, can someone explain?

Submitted by chem5th on Sat, 01/05/2008 - 23:13 Permalink

HCl + NH3 ---> NH4Cl

Would the answer to that particular problem be neither?  HCl is the acid (a H+ or proton donor) while NH3 is the Bronsted-Lowry base (a H+ or proton acceptor).

6[Cl]- + [Pt]4+ ---> [PtCl6]2-

Is 6[Cl]- the lewis base while [Pt]4+ is the lewis acid or is it the other way around?
A Lewis acid is an electron pair acceptor Pt 4+  while Cl- is the Lewis base (an electron pair donor)

Somebody is being tricky with this one.  The formulas have been modified from what you would normally see:
a) H2S    -  acid
b) PO(OH)3  -  H3PO4  -  phosphoric acid
c) H2CaO2    -  Ca(OH)2 - calcium hydroxide (base)
d) ClO3(OH)  -  HCl04 - perchloric acid
e) Sb(OH)3  - -  Not positive, but I would guess amphoteric (or both as you said).

You might want to remember that acid anhydrides (oxides of non-metals) react with water to form acids as in b and d, while base anhydrides (oxides of metals) react with water to form bases.

Submitted by spock on Sat, 01/05/2008 - 23:52 Permalink