Lesson 7:
Acids and Bases

 

Objective

In this lesson we will answer the following questions:

  • What are acids and bases?
  • How do acids and bases participate in chemical reactions?
  • How are acids and bases measured?


 

Reading Assignment

In addition to the online lecture, read chapter 6 in Basic Chemistry for Water and Wastewater Operators and Chapter 2 and Chapter 21 in Simplified Procedures for Water Examination.




Lecture

Acids and Bases

Throughout history, chemists have created different definitions of acids and bases. Today, many people use the Brønsted-Lowry version. It describes an acid as a molecule that will give away a proton from one of its hydrogen atoms. At a minimum, that tells us that all Brønsted-Lowry acids must contain hydrogen as one of their building blocks. Hydrogen, the simplest atom, is made up of one proton and one electron. When an acid gives away its proton, it hangs on to the hydrogen atom's electron. This is why scientists sometimes call acids "proton donors". Brønsted-Lowry bases, in contrast, are good at stealing protons, and they'll gladly take them from acids.

Not to confuse you, but scientists sometimes use another scheme, the Lewis system, to define acids and bases. Instead of protons, this Lewis definition describes what molecules do with their electrons. In fact, a Lewis acid doesn't need to contain any hydrogen atoms at all. Lewis acids only need to be able to accept electron pairs.

Water is chemically neutral. That means it is neither an acid nor a base. But mix an acid with water and the water molecules will act as bases. They'll snag hydrogen protons from the acid. The altered water molecules are now called hydronium. Mix water with a base and that water will play the part of the acid. Now the water molecules give up their own protons to the base and become what are known as hydroxide molecules.

Acid-base Chemistry

Acid-base chemistry is based on one simple reaction - the ionization of water:


H2O double-sided arrow H+ + OH-


As we saw in the last lesson, this equation means that water can break apart into a hydrogen ion and a hydroxide ion.  A hydrogen and a hydroxide ion can also join together to form a water molecule. 

In pure water, hydrogen and hydroxide ions are present in a 1:1 ratio since the only source of these ions is the ionization of water.  Acids and bases are substances which change the balance of hydrogen and hydroxide ions in water.  In this lesson, we will first explore the chemistry of acids and bases, then we will show how to calculate their concentration mathematically.  In the next lesson, we will find out how to calculate the concentration of acids and bases in the lab.



Acids

vinegar and onion

Vinegar contains acetic acid, which makes it taste sour. 
Onions release a gas which turns into sulfuric acid when
it reaches your eyes, making them burn. 


What do you think of when you think of acids?  You might think of sour-tasting acidic foods such as lemons.  Or you might think of strong acids, such as battery acid, which can burn your skin or corrode metal.  In this lesson, we will be concerned with the less visible properties of acids - their chemical properties. 

Chemically, acids are substances which increase the concentration of hydrogen ions when they are placed in water.  In this section we will be concerned only with strong acids.  Weak acids act slightly differently and will be discussed in a later section. 

You can recognize strong acids because they are ionic compounds which contain hydrogen ions.  For example, consider the strong acids listed in the table below.  Notice that each compound consists of a hydrogen ion bonded to some sort of anion. 

Name
Formula
Cation
Anion
Nitric acid
HNO3
H+
NO3-
Perchloric acid
HClO4
H+ ClO4-
Sulfuric acid
H2SO4
H+ SO42-
Hydrochloric acid
HCl
H+ Cl-
Hydrobromic acid
HBr
H+ Br-
Hydriodic acid
HI
H+ I-


When a strong acid is placed in water, it will ionize completely, breaking down into its constituent ions.  For example, hydrochloric acid reacts as shown below:

HCl yields H+ + Cl-


The ionization of a molecule of hydrochloric acid introduces a hydrogen ion and a chloride ion to the solution.  The chloride ion has no effect on the acidity of the water, but the hydrogen ion makes the solution more acidic. 

Diagrams of solutions with and without hydrochloric acid.

The diagrams above show what happens when hydrochloric acid is added to water.  The top diagram merely contains water.  Notice that most of the water is present in its un-ionized state, but that two water molecules have ionized into hydroxide and hydrogen ions.  Despite the presence of hydrogen ions, this solution is neutral (meaning that it is neither acidic nor basic) because the number of hydroxide ions equals the number of hydrogen ions. 

The bottom diagram shows what happens to the water after the addition of hydrochloric acid.  Notice that all of the hydrochloric acid has ionized, so now there are five hydrogen ions and only two hydroxide ions in the solution.  Since there are more hydrogen ions than hydroxide ions present, the solution has become acidic. 




Bases

Common bases.


You probably have less familiarity with bases than with acids, so you may be surprised to learn how many bases you deal with in your everyday life.  Soap, baking soda, milk of magnesia, and ammonia all contain bases.  These substances exhibit some of the physical properties of bases, such as feeling slippery, tasting bitter, and dissolving greases.  Bases also irritate the skin and eyes just as acids do. 

Chemically, a base is a substance which decreases the concentration of hydrogen ions when it is placed in water.  Strong bases decrease the hydrogen ion concentration by increasing the hydroxide ion concentration. 

Let's consider what happens when we put sodium hydroxide (a strong base) in water.  Like strong acids, strong bases ionize completely, so the sodium hydroxide breaks down into a sodium ion and a hydroxide ion:

NaOH yields Na+ + OH-


The hydroxide ion released by the sodium hydroxide then reacts with a hydrogen ion from the water, forming a water molecule:

OH- + H+ yields H2O


The reaction between the hydroxide ion and the hydrogen ion removes the hydrogen ion from the solution, making the solution less acidic and more basic.  The two diagrams below show the net reaction of sodium hydroxide with water schematically:

Reaction of sodium hydroxide with water.


Notice that in the top diagram, three water molecules have ionized to produce hydrogen and hydroxide ions.  When three sodium hydroxide molecules are added (in the bottom diagram) the hydroxide ions from the base's dissociation combine with the three hydrogen ions in the solution, forming water.  So the net result of the addition of sodium hydroxide to the water is that the concentration of hydrogen ions becomes lower and the acidity of the solution is decreased. 

In the table below, I have listed some of the strong bases which you may run into in the lab:

Name
Formula
Sodium hydroxide NaOH
Lithium hydroxide LiOH
Potassium hydroxide KOH
Rubidium hydroxide RbOH
Cesium hydroxide CsOH
Calcium hydroxide Ca(OH)2
Strontium hydroxide Sr(OH)2
Barium hydroxide Ba(OH)2





Weak Acids and Bases

So far we have talked about strong acids which contain a hydrogen ion and strong bases which contain a hydroxide ion.  However, some acids and bases work slightly differently.  Technically, an acid is any substance which donates a hydrogen ion to a solution and a base is any substance which accepts a hydrogen ion. 

The table below lists some weak acids and bases:

Weak Acids Weak Bases
Name
Formula
Acetic acid
HC2H3O2
Phosphoric acid
H3PO4
Carbonic acid
H2CO3
Hydrofluoric acid
HF
Fluosilicic acid
H2SiF6
Name
Formula
Ammonia
NH4OH
Magnesium hydroxide
Mg(OH)2
Aluminum hydroxide
Al(OH)3
Lime
CaO
Sodium silicate Na2SiO2
HTH
NaOCl and Ca(OCl)2
Soda ash
Na2CO3


Some of these weak acids and bases ionize just like their strong counterparts, donating hydrogen or hydroxide ions to the solution.  For example, ammonia will ionize as follows:


NH4OH
yields NH4+ + OH-


The difference between strong and weak bases (and between strong and weak acids) is that weak bases do not ionize completely when placed in solution.  So, while every molecule of sodium hydroxide will break down into a sodium ion and a hydroxide ion when placed in solution, only some of the ammonia molecules will ionize under the same conditions.   The illustration below  gives a schematic representation of what happens under these two circumstances. 

comparison of a strong and a weak acid


The top rectangle is an example of a solution containing sodium hydroxide. 
All of the molecules have ionized into their constituent ions.  The bottom rectangle is an example of a solution containing ammonia.  Most of the ammonia stayed in its current state, with only a small percentage breaking apart into ammonium and hydroxide ions. 


Some weak bases not only do not ionize fully, they also do not release a hydroxide ion.  For example, ammonia in its gaseous form reacts as follows:

H2O + NH3 yields NH4+ + OH-


Notice that the ammonia accepted a hydrogen ion from the water, decreasing the acidity of the solution.  So it acted as a base even though it didn't produce hydroxide by ionization. 

 

 

pH

Introduction

You should already be familiar with pH, which is the scale we use to measure the acidity or alkalinity of water.  You will remember that the pH scale runs from 0 to 14, with numbers less than 7 being acidic and with numbers more than 7 being basic (also known as alkaline.)  A pH of 7 means that the substance is neutral. 

pH scale



Although this elementary understanding of the pH scale is enough for many water and wastewater treatment processes, you will need to know more in order to fully understand the chemistry of acids and bases.  pH is actually a measure of the concentration of hydrogen ions in the solution.  Since acids donate hydrogen ions to a solution, they tend to make the pH lower.  Bases, by accepting hydrogen ions, make the pH higher.  A neutral pH of 7, which is the pH of distilled water, contains the same number of hydrogen ions as hydroxide ions. 

The math used to calculate a solution's pH is relatively complicated, and we will not present it here.  Instead, you just need to remember that the lower the pH value, the higher the concentration of hydrogen ions.  For example, stomach acid with a pH of 1 has a much higher concentration of hydrogen ions than tomatoes do with a pH of 4.  Baking soda with a pH of 8 has more hydrogen ions than household ammonia with a pH of 11. 

You shouldn't find it surprising that pH is a measure of the hydrogen ion concentration of a solution, since we have already explained that acids and bases change that concentration.  What you may not realize is that the pH scale only covers a small range of acidity and alkalinity.  In fact, the pH scale is meant to mimic nature by covering the acidity and alkalinity values which might be found in natural waters.  Strong acids and bases, like those discussed on the last page, have pH values at the far ends of the scale or even off the scale.  Concentrated hydrochloric acid has a pH of 0, and one drop of 33% hydrochloric acid in a liter of distilled water can lower the pH from neutral to about 3.  We will discuss ways to measure the concentration of strong acids and bases in a later section.



Neutralization

A neutral pH of 7 may mean that you are dealing with distilled water containing no acids and bases.  In this case, the amount of hydrogen ions and hydroxide ions will be equivalent because they will both be due to the ionization of water.  However, a neutral pH can also be achieved in a solution containing acids and bases as long as the acids and bases have neutralized each other, meaning that the acids have donated as many hydrogen ions as have been accepted by the bases. 

Neutralization reactions occur whenever acids and bases are placed in proximity.  An acid combines with a base to create water and a salt, as shown below:

HCl + NaOH yields H2O + NaCl

Titration, which we will introduce in a later lesson, is based on this neutralization reaction.  You will also need to understand neutralization if you spill an acid or base in lab and want to clean it up safely.  Neutralizing an acid with a base (or vice versa) can aid in cleanup, but you should also be aware that strong acids and bases can react explosively.  Always use weak or low concentration acids or bases for neutralization reactions. 

Neutralization occurs in nature as well.  For example, organisms living in very acidic environments tend to excrete basic wastes which bring the environment back into equilibrium.  Organisms living in basic environments excrete acids instead.  This is one of the reasons that most natural waters, including septic tanks and wastewater treatment ponds which have been allowed to work for some time, tend to have a pH near 7.  In addition, buffers (which we will explain in the next section) neutralize acids and bases both in natural waters and in the laboratory. 



Buffers

A buffer is a solution containing a weak acid and one of its salts or a weak base and one of its salts.  This solution is able to neutralize acids and bases without allowing the pH of the solution to change greatly.  In lab, buffers are used when the pH of a solution must remain stable. 

Some examples of the pairs which make up buffer solutions are shown in the table below.

Acid or Base
Salt
Acetic acid
Sodium acetate
Phosphoric acid
Potassium phosphate
Oxalic acid
Lithium oxalate
Carbonic acid
Sodium carbonate
Ammonium hydroxide
Ammonium nitrate

 

In order for a buffer to "resist" the effect of adding strong acids or bases, it must have both an acidic and a basic component. However, you cannot mix any two acid/base combination together and get a buffer. If you mix HCl and NaOH, for example, you will simply neutralize the acid with the base and obtain a neutral salt, not a buffer. For a buffer to work, both the acid and the base component must be part of the same equilibrium system - that way, neutralizing one or the other component (by adding strong acid or base) will transform it into the other component, and maintain the buffer mixture. Therefore, a buffer must consist of a mixture of a weak conjugate acid-base pair.

Of course, a buffer will not continue to neutralize the solution indefinitely.  Eventually, the acid or salt will be used up, and the pH of the solution will begin to change.  The amount of acid or base which a buffer solution is able to neutralize is known as the buffer capacity

Buffer solutions are not limited to the lab.  In natural water systems, carbon dioxide from the air often enters the water, forming carbonic acid.  A salt of carbonic acid, such as calcium carbonate (limestone), may become dissolved in the water from the surrounding rocks and soil.  Thus, a natural buffer solution is formed. 

 

 

Normality

While pH is used to record the acidity or alkalinity of natural waters, we use a measurement known as normality to show the concentration of the much stronger acid and base solutions we use in the lab.  Normality is based on molarity, but also takes into account a characteristic of acids and bases which we will call "equivalents" and will describe in the next section. 

Normality, equivalents and equivalent weight are all related terms typically used in titrations when the titration reaction is unknown or just not used. Consequently, definitions for these terms vary depending on the type of chemical reaction that is being used for the titration. The two most common types of reactions for which normality is used are acid-base reactions and redox (reduction-oxidation) reactions.

The basic unit for normality related conventions is the equivalent. Equivalents are comparable to moles and used to relate one substance to another. Normality is a measure of concentration equal to the gram equivalent weight per liter of solution. Gram equivalent weight is the measure of the reactive capacity of a molecule. The solute's role in the reaction determines the solution's normality. Normality is also known as the equivalent concentration of a solution.

We have already determined in a previous lesson that molarity of a solution refers to its concentration (the solute dissolved in the solution). The normality of a solutionr refers to the number of equivalents of solute per Liter of solution. The definition of chemical equivalent depends on the substance or type of chemical reaction under consideration. Because the concept of equivalents is based on the reacting power of an element or compound, it follows that a specific number of equivalents of one substance will react with the same number of equivalents of another substance. When the concept of equivalents is taken into consideration, it is less likely that chemicals will be wasted as excess amounts. Keeping in mind that normality is a measure of the reacting power of a solution, we use the following equation to determine normality:

 

 

Example 1:

If 2.0 equivalents of a chemical are dissolved in 1.5 L of solution, what is the normality of the solution?

 

 

Example 2:

A 800-mL solution contains 1.6 equivalents of a chemical. What is the normality of the solution?

First convert 800 mL to Liters:

800 mL / 1000 mL = 0.8 L

 

Now calculate the normality of the solution:

 



Equivalents

The reactive capacity of a chemical species, the ions or electrons, depends on what is being transferred in a chemical reaction. In acid-base reactions, an equivalent is the amount of a substance that will react with one mole of hydrogen ions. In oxidation-reduction (redox) reactions, where electrons are either gained or lost in a chemical reaction, it is one mole of electrons. Finding equivalents depends on the chemical species under consideration.

The oxidation state of an element describes the number of electrons transferred in reactions. For example, the oxidation, or valence states, of the following elements are equal to the number of equivalents:

Calcium: (Ca+2) ion: valence of 2; number of equivalents = 2

Aluminum: (Al+3) ion: valence of 3; number of equivalents = 3


For acids, an equivalent is the number of hydrogen ions a molecule transfers. In acids it is straightforward to find equivalent units. Look at the number directly after the hydrogen (H) in the chemical formulas below. The number provides the number of equivalents per mole of that acid:

Hydrochloric acid (HCl): equivalents = 1

Sulfuric acid (H2SO4): equivalents = 2

Phosphoric acid (H3PO4): equivalents = 3

Nitric acid (HNO3): equivalents = 1

 

For bases, it is the number of hydroxide ions (OH-) provided for a reaction, such as:

Sodium hydroxide (NaOH): equivalents = 1

Barium hydroxide (Ba(OH)2): equivalent = 2

 

One equivalent of an acid reacts with one equivalent of a base. For the acid HCl and base NaOH, both with one equivalent, they have the same reactivity. For H2SO4, with two equivalents, and NaOH, it will take twice the amount of the NaOH to react with the sulfuric acid. Mixing equal equivalents of acidic and basic solutions will result in a neutral solution.

Altough molarity can be used to measure the concentration of acids, it is a relatively unuseful measurement for understanding neutralization reactions.  Why?  Because not every acid or base can add (or remove) the same number of hydrogen ions from solution. 

 

 

Equivalent Weight

The equivalent weight can be thought of as the weight (or mass, to be precise) of a substance that will contain a single reactive proton (or hydrogen ion) or a single reactive hydroxide ion. The former case applies to acids, which are proton donors, while the second applies to bases, which are proton acceptors. The reason the concept of equivalent weight is needed is that some compounds can donate or accept more than one proton, meaning that for every mole present, the substance is in effect doubly reactive. The equivalent weight can be determined by:


 

Equivalent weight is defined as the ratio of molar mass of a substance to the valence of the substance. Valence is also denoted as equivalence factor. The valence is the number of hydrogen atoms in an acid, or hydroxide atoms in a base, and for salt, charge present in ionic forms. For reduction/oxidation (redox) reactions, it is the number of electrons than an oxidizing or reducing agent can accept or dontate that are counted as valence or equivalence factor. When the equivalent weight is expressed in grams using molar mass in grams, it is called gram equivalent weight.

Let's get some practice determining the equivalent weight for the following formulas:

H2SO4:

Molar mass of H2SO4 = 98 g/mol

Looking at the formula, there are 2 hydrogen atoms, so "n" will be 2 when determine the equivalent weight:

 

NaCl:

Molar mass of NaCl = 58.5 g/mol

Looking at the formula, because there are no hydrogen or hydroxide atoms, the number of equivalents is 1, because there will always be at least one equivalent before a rection can occur.

 

NaOH:

Molar mass of NaOH = 40 g/mol

In looking at the formula, there is 1 hydroxide (OH) atom, so the equivalent is 1.

 

Now let's look at a salt (a salt determines its equivalents differently because there are no hydrogen or hydroxide atoms involved, so we look at the charge):

Na2CO3:

Molar mass of Na2CO3 = 106 g/mol

The salt Na2CO3 ionizes to form 2Na+ and CO3-2, so the charge present on both is 2.

 

 


Making Normal Solutions (N)


Normality is the most common measurement used for showing the concentration of acids and bases.  Normality takes into account both the molarity of the solution and the equivalent content of the acid or base. It is defined as the number of gram equivalent present in per liter solution and can be determined through the following formula:

 

To calculate normality of the solution, follow these steps:

  • Find the equivalent weight of the solute based on the chemical reaction it is going to be using
  • Calculate the number of gram equivalents of the solute
  • Calculate the volume in liters
  • Calculate the normality using the formula given above

 

Making normal solutions can be a bit confusing. Aqueous solutions of acids and bases are often described in terms of their normality rather than their molarity. In order to properly make a Normal solution, the student must understand the difference between a pure reagent and a diluted reagent.

A "1 Normal" solution (1 N) contains 1 "gram equivalent weight" of solute, topped-off to one liter of solution. The gram equivalent weight is equal to the solute's molecular weight (molar mass), expressed as grams, divided by the valence (n) of the solute:

 

After the equivalent weight (or milliequivalent weight) has been calculated, then the following equation is used:

or

 

The equivalent weight of a substance depends upon the type of reaction in which the substance is taking part. Some different types of chemical reactions, along with how to determine a solute's equivalent weight for each reaction, are given below.

 

 

 

Making a Normal Solution With Salts

Example:

Calculate the normality of a sodium chloride (NaCl) solution prepared by dissolved 2.9216 grams of NaCl in water and then topping it off with more water to a total volume of 500.0 mL.

First, check the periodic table to determine the molar mass, or molecular weight, of NaCl, which is 58.44.

In looking at the formula (NaCl), the equivalent is 1 because there is room in the molecule for only one replaceable (H+) ion. In other words, one hydrogen atom can replace the sodium atom in NaCl.

Now determine the equivalent weight of NaCl:

 

You will need to know the milliequivalent weight of NaCl (since the solute volume is in mL):

 

Next calculate the normality of the sodium chloride solution:

 


Making Normal Solutions With Pure (Non-Aqueous) Acids

The equivalent weight of an acid is its molecular weight, divided by the number of replaceable hydrogen atoms in the reaction. To clarify this concept, consider the following acids:

Hydrochloric acid (HCl) has one replaceable hydrogen ion (H+). Sulfuric acid (H2SO4) has two replaceable hydrogen ions (2H+). The valences of these acids are determined by their respective replaceable hydrogen ions, displayed below.

Acid Valence
(Replaceable hydrogen ions)
HCl n = 1
HNO3 n = 1
H2SO4 n = 2
HF n = 1

 

So, for pure HCl, its molecular weight is 36.46, its equivalent weight is 36.46 and therefore a 1N solution would be 36.46 grams of the pure chemical per liter. Note that, in the case of HCl, a 1N solution has the same concentration as a 1M solution.

To make a 1N H2SO4 solution from pure H2SO4, its molecular weight is 98.08, and its equivalent weight is (98.08/2 = 49.04 grams/Liter or 49.04 grams per 1000 mL). So a 1N solution would be 49.04 grams of the pure chemical per liter.

 

 

 

Making Normal Solutions From Pure Alkalis (Bases)

The equivalent weight of a base is defined as "its molecular weight divided by the number of hydrogen ions that are required to neutralize the base". To understand the valences of alkalis, consider the following examples:

The (OH-) ion in Sodium hydroxide (NaOH) can be neutralized by one hydrogen ion. the (OH)2-- ions in Calcium hydroxide (Ca(OH)2) can be neutralized by two hydrogen ions. As was the case with acids, the valences (n) of these bases are determined by their respective replaceable hydrogen ions, displayed below.

Base Valence
(Replaceable hydrogen ions)
NaOH n = 1
Ca(OH)2 n = 2

 

So, for NaOH, its molecular weight is 40, its equivalent weight is 40, and therefore a 1N solution would be 40 grams of the pure chemical per Liter of water. You will also note, in the case of NaOH, a 1N solution is the same concentration as a 1M solution.

For Ca(OH)2, its molecular weight is 74, its equivalent weight is (74/2 = 37, because n = 2). Therefore, a 1N aqueous solution of Ca(OH)2 is 37 grams of the pure chemical per Liter of water. Of course, many acidic reagents and basic reagents come from the factory in a diluted aqueous form, which is a form of preparation we will not cover in this lesson.

 

Example:

Calculate the normality of a NaOH solution formed by dissolving 0.2 g NaOH to make a 250 mL solution.

First we need to determine the molar mass, which is 40 g/mol.

Then determine the milliequivalent weight (since the solution is in mL):


 

Now, calculate the normality of the NaOH solution:

 

Relationship Between Normality and Molarity

Here is Normality in terms of molarity:

Normality = n x Molarity

Where n = number of Hydrogen in acids, or Hydroxides in bases and for salt, charge present in ionic forms.

 



Calculating Dilutions (Acid-Base Titration)

Once you have calculated the normality of an acid or base solution, you can easily calculate the concentration of any dilutions of that solution.  The formula used is essentially the same as that used for any other dilution calculation:

N1V1 = N2V2

Where:

N1 = normality of the first solution
V1 = volume of the first solution
N2 = normality of the second solution
V2 = volume of the second solution

 

Example 1:

After producing the 0.5 L of a 0.37N solution of calcium hydroxide in the last section, how would you dilute it to form a 0.25N solution? 

(0.37N) (0.50 L) = (0.25N) V2

(0.37N)(0.50 L) / 0.25N = V2

0.74 L = V2


Based on the calculations above, we know that we have to add enough water to the 0.37N solution so that the total volume reaches 0.74 L.  Then we would have a 0.25N solution. 

 

 

Example 2:

How many milliliters of 2N NaOH are needed to prepare 300 mL of 1.2N NaOH?

N1V1 = N2V2

(2N)(V1) = (1.2N)(300 mL)

V1 = (1.2N)(300 mL) / 2N

V1 = 180 mL

So, to prepare the 1.2N NaOH solution, you pour 180 mL of 2N NaOH into your container and add water to get 300 mL total volume.


 

 

Safety Tips

  • Always use a fume hood when handling highly concentrated acids and bases. Wear a platic apron, plastic gloves, and eye goggles. This is particular important when working with hydrofluoric acid.
  • If you need a reagent that can produce highly precise results in quantitative analysis, then you should store all newly prepared strong basic solutions in labeled plastic containers.
  • Highly acidic reagents can be stored in either glass containers or in plastic containers. The exception is hydrofluoric acid, which etches glass and weakens it. Store hydrofluoric acid in a labeled plastic container.
  • It is advisable to store concentrated reagents that come straight from the supplier in their original containers, with their labels intact.
  • When preparing solutions, students sometimes make the mistake of adding the solute to a set volume of solvent. Doing this will produce a solution of the wrong concentration. Example: when making one liter of a 1N solution of NaCl, do not add 58.5 grams of NaCl to 1 Liter of water. Instead, the correct way to make the solution is to add 58.5 grams of NaCl into a container and then top it off with water to a total volume of 1 Liter.
  • Never add water into a large volume of concentrated acid! You risk creating an explosion! The rule is:

    "Acid into water = you're doing what ya oughta."

    "Water into acid = you might get blasted!"

     

     

Alkalinity

The final topic we will introduce in this lesson is alkalinity, which is the capacity of a solution to neutralize a strong acid.  Despite the name, you should not think of alkalinity as the amount of base found in the solution; instead, alkalinity is a measurement of the buffering ability of the solution.  Water high in alkalinity will be able to maintain its pH, despite the addition of acids. 

Alkalinity in natural waters is caused primarily by carbonate (CO3-2) and bicarbonate (HCO3-) ions.  A few other causes of alkalinity may also be present in water, but usually at much lower concentrations.  These other types of alkalinity can include hydroxide, borate, silicate, phosphate, ammonium, and sulfide.  Bicarbonates are the major components because of carbon dioxide action on basic materials of the soil. The alkalinity of raw water may also contain salts formed from organic acids such as humic acids.

Alkalinity in water acts as a buffer that tends to stabilize and prevent fluctuations in pH. In fact, alkalinity is closely related to pH, but the two must not be confused. Total alkalinity is a measure of the amount of alkaline materials in the water. The alkaline materials act as buffers to changes in the pH. If the alkalinity is too low (below 80 ppm or mg/L), the pH can fluctuate rapidly because of insufficient buffer. High alkalinity (above 200 ppm) results in the water being too buffered. Thus having significant alkalinity in water is usually beneficial because it tends to prevent quick changes in pH that interfere with the effectiveness of common water treatment processes. Low alkalinity also contributes to the corrosive tendencies of water.

Acidity in water is usually due to carbon dioxide, mineral acids, or hydrolysis of some heavy metal salts, such as aluminum sulfate. Hydrolysis is decomposition of a substance by reacting with water.

How are hardness and alkalinity related? They are both expressed as mg/L CaCO3. In tap water, when the concentration of hardness and alkalinity are the same, both are probably due to dissolved calcium carbonate. Calcium carbonate does not dissolve well above pH 7.0, so other chemicals like calcium chloride and sodium bicarbonate can be used to adjust hardness and alkalinity independently.

Alkalinity usually enters the water as salts, such as calcium carbonate (also known as limestone).  Carbonate and bicarbonate can also be formed when carbon dioxide from the air is dissolved in the water.  Hydrolized metal salts, like aluminum sulfate and various iron salt coagulants, produce acidity and, thus, consume alkalinity. Although the exact amount may vary, 1 mg/L of alum requires about 0.5 mg/L of alkalinity. It is essential that operators understand this reaction because coagulation may be affected if enough alkalinity is not present.

In any case, all forms of alkalinity operate similarly to neutralize acids.  Let's see what happens when we add sulfuric acid to water containing bicarbonate alkalinity:

2HCO3- + H2SO4 → 2H2CO3 + SO42-



First, the sulfuric acid loses its hydrogen ions, which would usually increase the acidity of the water.  However, the bicarbonate accepts the hydrogen ions, turning into carbonic acid.  Since carbonic acid is a weak acid, it tends to stay in its un-ionized state, so the pH of the water remains neutral.

The general formula for alkalinity derived from a titration is:

 

Example:

A 100 mL sample of water is tested for alkalinity. The normality of the sulfuric acid used for titrating is 0.02N. If 7.6 mL of titrant is used, what is the alkalinity of the sample?

 

 

 

Alkalinity in Natural Waters

Alkalinity is a measure of a river's buffering capacity, or its ability to neutralize acids. Alkaline compounds in the water, such as bicarbonates (baking soda is one type), carbonates, and hydroxides remove hydrogen ions and lower the acidity of the water, which means increased pH. They do this usually be combining with the hydrogen ions to make new compounds. Without this acid neutralizing capacity, any acid added to a river would cause an immediate change in the pH. Measuring alkalinity is important to determining a river's ability to neutralize acidic pollution from rainfall or snowmelt. It's one of the best measures of the sensitivity of the river to acid inputs. Alkalinity comes from rocks and soils, salts, certain plant activities, and certain industrial wastewater discharges.

 




The Carbonate System

The relationship between alkalinity and pH is relatively complex.  As mentioned in the last section, higher alkalinity tends to prevent water from becoming acidic.  In addition, pH influences the type of alkalinity found in water.  In this section, we will consider the carbonate system, which is the relationship between pH and the different forms of carbonate - carbonic acid, bicarbonate, and carbonate. 



The graph above summarizes the carbonate system.  Notice that the form of carbonate in the water is very dependent upon pH.  At a low pH, the carbonate is present as carbonic acid.  Bicarbonate can be found in water with a pH between 4.3 and 12.3.  Above a pH of 8.3, carbonate is also present. 

What causes this relationship between carbonate species and pH?  Carbon dioxide, carbonic acid, bicarbonate, and carbonate can transform back and forth by gaining or losing hydrogen ions in the reactions shown below:

CO2 + H2O double-sided arrow H2CO3

H2CO3 double-sided arrow HCO3- + H+

HCO3- double-sided arrow CO32- + H+


Based on the concentration of hydrogen ions in the solution (the acidity of the solution), various of these reactions will be more or less likely to take place.  For example, when the water is very acidic (containing a high concentration of hydrogen ions), the hydrogen ions tend to attach to carbonate or bicarbonate, forming carbonic acid.  However, if the water is very basic (containing a lower concentration of hydrogen ions), then carbonic acid tends to break apart, adding hydrogen ions to the solution. Do not concern yourself with the reactions above, it is just to show you how gaining and losing hydrogen ions can transform the carbonate species.

 




In the Treatment Plant

Alkalinity is of little sanitary significance, although extremely high levels of alkalinity in water can be problematic.  It has been reported that a carbonate concentration of 350 ppm is unhealthy.  In addition, free caustic (hydroxide) alkalinity may impart a bitter taste to the water and can even cause a burning sensation at high concentrations.  However, in most water treatment plants, operators are more concerned with ensuring that there is enough alkalinity in the water to promote coagulation and prevent corrosion than they are with removing alkalinity from water.

Alkalinity is essential for proper coagulation.  When alum is added to raw water, the alum reacts with the alkalinity present to form floc.
  As a rule of thumb, coagulation usually requires an alkalinity equal to half of the amount of alum used.  So, if your plant treats raw water with 25 ppm of alum, then a minimum of 12.5 ppm of alkalinity must be present in the water for floc to form.  If insufficient alkalinity is present in the water during coagulation, then dense floc with not form and soluble alum will be left in the water. 

Since coagulation uses up alkalinity in the water, treated water often tends to be corrosive (acidic).  The operator must make sure that an adequate amount of alkalinity is present in the treated water as well as in the raw water to prevent corrosion of pipes in the distribution system.  We will discuss the math used to determine the appropriate alkalinity of treated water in the alkalinity lab. 

Water treatment may require the addition of artificial alkalinity to enhance coagulation or to prevent corrosion.  If the source water does not contain an adequate amount of natural alkalinity, treatment will usually include the addition of lime or soda ash to artificially increase the amount of alkalinity in the water. 
Water with high pH is generally depositing and water with low pH is corroding.





Review

Acids and bases are substances which change the balance of hydrogen and  hydroxide ions in water.  Both are divided into two categories - strong acids or bases which ionize completely in water, and weak acids or bases which ionize only partially in water. 

Acids and bases undergo a variety of chemical reactions.  Acids donate hydrogen ions to water while bases remove hydrogen ions from water.  When acids and bases are brought together, they neutralize each other.  Buffers can neutralize both acids and bases. 

pH can be used to measure the balance of acidity and alkalinity in water.  Normality, in contrast, is used to measure the concentration of a specific acid or base in a relatively strong solution. 

Alkalinity is the buffer capacity of water, usually caused by carbonate and bicarbonate ions.  The type of alkalinity found in water will depend on the pH of the water.  Regardless of the type of alkalinity, a certain concentration of alkalinity is required in the water treatment plant to promote good coagulation and to prevent corrosive water. 



 

New Formulas Used

   Normality1Volume1 = Normality2Volume2 or (N1V1 = N2V2)









Assignments

Complete the math worksheet for this lesson and return to instructor via email, fax or mail. (Normal layout)

If you would like to download the worksheet as a word document to fill out it can be found here: Worksheet 7

 


 

Lab

Read the pH and alkalinity labs and do the assignment listed above, there are questions concerning the virtual labs included.


 

Quiz

Answer the questions in the Lesson 7 quiz The grade will be submitted directly to the gradebook. If you are unhappy with your grade, close the quiz out and re log back in for the new grade to submit.