Making Normal Solutions
from Concentrated Acids
Tim Loftus
The last article covered the concept of Normal solutions in the
laboratory and how to calculate the equivalent mass of a compound.
Then I described how to use the equivalent mass to make a solution of
a predetermined Normality. However, the article did not address making
Normal solutions from concentrated mineral acids like sulfuric acid,
nitric acid, and hydrochloric acid. Unlike using powdered chemicals
where the chemical is simply weighed out then diluted to volume, the
use of liquid chemicals to make Normal solutions requires the addition
of a few more calculations. This article will address these extra
calculations.
First, it is important to describe a few aspects of concentrated
mineral acids (as well as that of many other solutions). Most of us
buy concentrated acids to use as stock solutions in the laboratory.
None of these acids are one hundred percent pure. Sulfuric acid is
only about 97% pure, nitric is about 69.5%, and hydrochloric acid is
about 37.5% pure. Manufacturers of these acids simply cannot
economically make these acids more concentrated than these respective
percentages.
Another important aspect of these solutions is their specific
gravities. The specific gravity of a liquid is, in most cases,
synonymous with the more familiar term of density. Water has a
specific gravity of 1. If the specific gravity of a liquid is greater
than 1, then the liquid is heavier than water. Less than 1, and the
liquid is lighter than water. The specific gravity for concentrated
sulfuric acid is about 1.84, or 1.84 times heavier than an equal
volume of water. The specific gravity of concentrated nitric acid is
about 1.42 and that of concentrated hydrochloric acid is about 1.19.
Both the percent concentration and specific gravity values of the acid
are required to determine the amount of concentrated acid needed when
making a Normal solution. This information is usually printed on a
label attached to the bottle of acid. Specific values vary depending
on the manufacturer and lot of acid.
To make a solution of a predetermined Normality, you must first
determine the equivalent mass of the chemical and then determine the
grams needed of that chemical. These calculations were described in
the last article, “Normality.” Then you must convert the number of
grams into its volume equivalent. Once this volume is determined, it
is a simple dilution after that.
Here is an example:
You want to make only 250 mL of a 1 N H2SO4 solution that will be used
to adjust the pH of BOD samples prior to analysis. How many
milliliters of concentrated sulfuric acid do you need to make 250 mL
of a 1 N solution?
To determine how many grams of sulfuric acid you will need, you will
first need to calculate the equivalent mass of H2SO4. This is the
gramformula weight divided by the number of acid hydrogens in the
compound. It is 98/2 = 49.
Then you can calculate the amount of grams of H2SO4 that are needed.
The formula to calculate this is:
Grams of compound needed = (N desired)(equivalent mass)(volume in
liters desired).
Substituting the above numbers into the equation, we get:
grams of compound needed = (1 N)(49)(0.250 liters) = 12.25 grams.
A 1 N solution requires 12.25 g of a pure sulfuric acid powder (if one
existed) diluted to 250 mL. But the acid is a liquid and it is not one
hundred percent pure active sulfuric acid. You will need to calculate
what volume of the concentrated acid that contains 12.25 grams of
sulfuric acid. The formula for this is:
Volume of concentrated acid needed = (grams of acid needed)/(percent
concentration x specific gravity)
Continuing with the sulfuric acid example, plug into the formula the
percent concentration and specific gravity from the label on the acid
container. For this example, I am using those values previously
mentioned in this article: volume of concentrated acid needed = (12.25
grams)/(0.97 x 1.84) = 6.9 mL
If you took 6.9 mL of concentrated sulfuric acid and diluted it to 250
mL, you would have a 1 N H2SO4 solution.
(Important note: Always add the acid (or base) to water, in that
order. Pour slowly with constant mixing. This will help prevent rapid
heat generation and spattering of the mixture. Fill a container about
half way or more with distilled water, add the acid, and then bring up
to volume with more water. In the example above, fill a flask with
about 150 mL or more with distilled water, add 6.9 mL of concentrated
sulfuric acid, then continue to dilute with water to the 250 mL mark.)
As with any acid or base made from a concentrated stock solution, the
resulting Normality will be an approximate value, which won’t be
accurate enough for analytical work. However, it will, in conjunction
with a pH meter, be good for adjusting the pH of samples. For
analytical procedures where the Normality needs to be accurately
known, as in alkalinity titrations, acidity titrations, and volatile
acid titrations, you will need to standardize the acid or base. An
overview of standardization and the shelf life of acids and bases will
be covered in a future article.
The information in this article is very general. As usual, check your
federal, state, and local regulations. You may have additional
regulations or requirements that you must meet.
If you have any questions, suggestions, or comments, contact NEWEA Lab
Practices Committee Chair Tim Loftus at (508) 9493865 timloftus@msn.com.
For more information on the NEWEA Laboratory Practices Committee,
please contact Tim Loftus or Elizabeth Cutone, NEWEA Executive
Director, 100 Tower Office Park, Woburn, MA 01801, (781) 9390908,
ecutone@newea.org.
All past articles are posted on our website. Go to www.NEWEA.org and
follow the link to the Committee Pages then to the Laboratory
Practices page.
