Acids and Bases
Acids and bases have been used for millennia. The term alkali for base dates back to ancient
times (particularly the Romans). They have many uses based on their universal properties.
Acids have a sour taste. So, lemons, sour patch kids, citrus, etc. have a large (relatively)
concentration of acid, although it may be partially masked by a large amount of sugar.
They react with metals to produce hydrogen gas. In the lab, you have reacted magnesium ribbon
or zinc with an acid to produce hydrogen gas. You confirmed the presence of hydrogen with a
burning splint that produced a loud POP as the hydrogen reacted with oxygen on the air to form
water. This is why we do not pour acids down the sink, they could react with the metal pipes
causing them to form soluble salts and weaken the pipes while making explosive hydrogen.
Acids react with carbonates (or bicarbonates) to from carbon dioxide. This is the test for
limestone or dolomite for geologists. A geologist will test a rock by placing a few drops of dilute
hydrochloric acid on a rock. If it is limestone or dolomite, it will fizz.
Acid turn litmus red.
Aci
ase
D. Conversely, bases turn litmus blue. B
Re
lue
Bases emulsify fats. This means that they dissolve oils that are not normally soluble in water. As
a result, bases are used as cleaning agents and degreasers. The most commonly used household
base for cleaning is ammonia (NH3). Bases taste bitter. So alkali compounds like sodium
carbonate and sodium bicarbonate (baking powder and baking soda, respectively) have a bitter
taste. So, bittersweet chocolate masks its bitterness with sugar, just as citrus fruits mask the
sour acid taste with sugar. Finally, bases feel slippery to the touch.
Fundamentally, in water, acids have a pH less than 7. Bases have a pH greater than 7. Any
solution that has a pH of 7 is considered neutral. Distilled water has a pH of 7. As the pH gets
farther from seven, the solution is a more concentrated (and therefore more dangerous) acid
(lower than 7) or base (higher than 7).
Acid/Base Theories
The first major theory of acids and bases was developed by Arrhenius. He stated that acids
contain (and therefore produce) H+1. He also stated that bases contain OH-1 ions. So, chemicals
like HCl, H2SO4, HCH3COO are acids and NaOH, Ca(OH)2, Ti(OH)3 are bases. This theory is neat
and simple. Unfortunately, we know it is wrong. ARRHENIUS WAS ERRONEOUS. Ammonia,
NH3, is a base. It has a pH larger than seven; it emulsifies fats; it feels slippery to the touch. But
notice the formula, there is no hydroxide ion!
That brings us to Bronsted and Lowry who independently created a theory of acids and bases
based on the movement of protons (H+1 ions). Acids are chemicals that donate protons in a
solution. Bases are chemicals that receive a proton from an acid. Therefore, you cannot have an
acid in a reaction without a base for it to react to. The second part of the theory is that acids
and bases always occur in conjugate pairs. It is important to note that the conjugates always
occur as products. To find the conjugate base of an acid, you remove an H and decrease the
charge by 1. For example,
. HCl is the acid and Cl-1 is its conjugate base.
Conversely, to find a conjugate acid of a base you add an H and increase the charge by 1. For
example, NH3, ammonia, is a base that has a conjugate acid that is NH 4+1. If you look closely at
most pharmaceuticals (prescription drugs), which are bases, they use the acidic salt form (HCl).
Another piece of Bronsted-Lowry theory is that a weak acid has a strong conjugate base and vice
versa. The smaller the Ka value is the weaker the acid. So, acids that have very small K a have
conjugate bases that have a large value for Kb.
Bronsted and Lowry also defined some substances as amphoteric. This refers to chemicals that
can act as either an acid or a base depending on what you react the chemical with. For example,
if you put an acid such as HCl or HCH3COO in water, water will behave as a base and accept a
proton from the acid to form H3O+1, hydronium ion, as a conjugate acid.
On the other hand, if you put a base in water, then it will behave as an acid by donating a proton
(H+1) to the base.
There is a fundamental problem with the Bronsted-Lowry theory. Protons are not transferred in
a chemical reaction! If a proton were transferred from one atom to another, the identities of
the atoms would change. Simply put, we would be performing ALCHEMY not CHEMISTRY.
This flaw was accounted for by Lewis (yes, the electron dot guy). Lewis stated that acids are
electron receivers (or electron poor, which means positive). Bases would therefore be electron
donors (i.e. electron rich or negatively charged).
Fundamentally, this means that nearly every chemical reaction can be defined as an acid-base
reaction. It also means that every time you have a central atom with a lone pair it can act as a
base. It also means that any atom that is the positive end of a polar bond (or a metal cation) can
act as an acid. This explains such things as protein synthesis, digestion, etc. Unfortunately,
Lewis theory cannot be used to calculate the molarity of H+1, [H+1], or pH.
The Math of Acids and Bases
Chemists would like to describe the acidity of a solution with the molarity of the H +1 or H3O+1
ion. However, these numbers are usually extremely small from 1.0 E-3 (.001) M to 1.0 E-11
(0.00000000001). Therefore, to make the numbers more manageable, chemists use a logarithmic
scale called pH which stands for the ‘power of the Hydronium ion.’
pH = -log[H +1]
[
For a base, it is more convenient to measure the [OH-1] and use pOH. pOH = -log OH
The two are easily interchanged using the equation pH + pOH =14
The key is knowing how to find the [H+1] or [OH-1] to start.
-1
]
A. STRONG ACID
For any of the strong acids (HCl, HBr, HI, HNO3,HClO4,HClO3, and H2SO4), the [H+1] is equal to
the molarity of the acid. This flows from the definition of a strong acid. Strong acids ionize
completely in water so the concentration of the hydronium ion, as well as the concentration of
the conjugate base, are equal to the molarity of the acid.
B. WEAK ACID
For any weak acid, it only partially ionizes in water. Consequently, the molarity of the hydronium
ion is always less than the original molarity of the acid.
1. Given percent ionization:
[H+1]=M*%ionization
You must remember to change the percentage ionization to a decimal by
dividing by 100. For example, if you had a 0.25 M weak acid that ionizes at 3.2%,
[H ] = .25 * .032 = .008M
+1
2. Given a Ka value
The acid equilibrium constant, Ka, is more widely used than percent ionization.
Ka =
x2
M
The ‘x’ is the [H+1]. The ‘M’ is the original molarity of the acid.
For example, a 0.125 M solution of acetic acid has a Ka value of 1.8 E-5.
x2
.125
.125(1.8E - 5) = x 2 = 2.25E - 6
1.8E - 5 =
2.25E - 6 = x 2
0.0015 M = x
C.STRONG BASE
The strong bases are the only hydroxides that were soluble according to the rules we learned for
double replacement. Column 1 hydroxides and Calcium, Barium, and Strontium (bottom of
column 2) are completely soluble and are strong bases.
For the column 1 hydroxides, the [OH-1] equals the molarity of the base. For barium, strontium,
and calcium hydroxides the [OH-1] is equal to twice the molarity of the base.
D. WEAK BASE
For any weak base, it only partially ionizes in water. Consequently, the molarity of the hydroxide
ion is always less than the original molarity of the base.
1. Given percent ionization:
[OH-1]=M*%ionization You must remember to change the percentage ionization to a decimal by
dividing by 100.
2. Given a KB value
The base equilibrium constant, KB, is more widely used than percent ionization.
Ka =
x2
M
The ‘x’ is the [OH-1]. The ‘M’ is the original molarity of the base.
Notice, that weak acids and bases use the same formulae; however, the unknown ‘x’ represents
something different (x is the [H+1] for acids and [OH-1] for bases). The math is identical and
lends itself to easy calculator manipulations.
HOW TO FIND pH, pOH, [H+1], and [OH-1]
These four things are all related, so if you know one of them you can find the other three using
the set of equations.
pH = - log[H+]
pOH = - log[OH-]
pH + pOH = 14
Where you start depends on what you know. For example, you have 0.0035 M HCl. You KNOW
the [H+]. So, the only place you can start is with the top equation.
pH = -log[H+] = -log [.0035] = 2.46
Now you know pH as well, so that allows you to drop to the bottom equation.
pH +pOH = 14
2.46 + pOH =14
So, pOH =11.54.
The only equation left is the middle one.
pOH = - log[OH-]
11.54 = - log [OH-]
To solve for something inside the log, you MUST use the
inverse log function, (commonly referred to as the ANTILOG). But, first, you must get all the
numbers, including the negative one on the right, to one side and the log function by itself on the
other. So, simply stated, just reverse the sign of both sides.
-11.54 = +log[OH-]
Now take the antilog of both sides.
ANTILOG(-11.54) = ANTILOG(log[OH ])
Log and antilog annihilate each other. Antilog on
your calculator is just [2nd] [log]. So hit [2nd][log](-11.54)
[enter]
2.86E-12 M = [OH-]