Introduction
Acids and bases are everywhere. Formic acid is an acid produced by ants, and acetic acid is
the acid in vinegar. Lactic acid is in sore muscles, citric acid is found in lemons, and sulfurous
acid is produced by volcanic gases and water. Chalk is a base and so is limestone. Natural
substances called alkaloids found in chocolate and tea are bases too. Bones are composed of
bases, and DNA is composed of four bases that are paired across the double helix. Sulfuric
acid, hydrochloric acid, and nitric acid are used extensively in the production of fertilizers,
making of plastics, purification of metals, synthesis of drugs, composition of batteries and
much more. Lye, or sodium hydroxide, and carbonates are bases commonly used in industry.
In this unit we will define acids and bases, identify common properties of acids and bases,
classify acids and bases as weak and strong, discuss pH and solutions, determine the products
for and balance acid-base reactions, and learn about experimental procedures and calculations
for a titration of an acid and base.
Acids and Bases
Acids and bases have specific properties that can be used to identify each type of substance.
Below is a table of the major properties of each.
Property
Chemical Formula
is Recognizable
Acid
Base
H–X
for example, HCl, HNO3, HC2H3O2,
H2SO4, & H3PO4,
Y–OH, Y-CO3 , & NR3
for example, NaOH, Ca(OH)2,
Al(OH)3, NaHCO3, CaCO3,
NH3, & NH(CH3)2
Most acids have only nonmetal
elements in the chemical formula
Most bases have a metallic
element in the chemical formula
Taste
sour
bitter
Feel
may sting
slippery
Common Reaction
Acid + Metal give hydrogen gas,
H2
Acid + Carbonates gives
carbon dioxide gas, CO2
Bases + fats/oils gives soaps
React with acids to from water
and a salt
pH
<7
>7
Indicator Colors
Litmus - Red
Bromothymol Blue - Yellow
Phenophthalein - Clear
Litmus - Blue
Bromothymol Blue - Blue
Phenophthalein - Magenta
Names of Acids and Bases
Most bases are named just like ionic compounds, by their cation and anion. For example,
NaOH is sodium hydroxide; Fe(OH)3 is iron(III) hydroxide; NaHCO3 is sodium hydrogen
carbonate (also sodium bicarbonate); CaCO3 is calcium carbonate. Amines, which are
compounds with carbon and hydrogen attached to N (NR3 where R can be a C or H group like
NH3 or NH(CH3)2 ) often have their own systematic names, NH3 is ammonia, NH(CH3)2 is
dimethyl amine.
Acids are named by a series of rules but always end with the term acid.
Rules for Naming Acids:
a. An acid with an element and hydrogen is called: hydro-element name-ic acid;
HCl = hydrochloric acid, H3P = hydrophosphoric acid, H2Se = hydroselenicic acid
(Many times these acids are named like diatomic covalent compounds, which they are;
for example, H3P = trihydrogen phosphorous (phosgene is the common name); or
H2S is dihydrogen sulfide)
b. An acid with a polyatomic ion ending in -ate is called: element name-ic acid;
HNO3 = nitric acid (NO3– is nitrate), H2SO4 = sulfuric acid (SO42– is sulfate),
H2CO3 = carbonic acid (CO32– is carbonate), H3PO4 = phophoric acid (PO43– is phosphate),
HCH3COO = acetic acid (CH3COO– is acetate, which can be written as,
C2H3O2– or CH3CO2–).
c. An acid with a polyatomic ion ending in -ite is called: element name-ous acid (these ions
have fewer oxygens than the ions ending in -ate, so the acid has the lowest number of
oxygens; I remember the the name because lowest rhymes with -ous).
HNO2 = nitrous acid; H2SO3 = sulfurous acid.
d. Some other acid names: HCN = hydrocyanic acid (CN– = cyanide, it is named like an
element); HOCl = hypochlorous acid (OCl– is hypochlorite) & HOCl4 = perchloric acid
(OCl4– is perchlorate).
Reactions with Water that Form Acids and Bases
The oxides of elements, like SO2 and Na2O, will react with water to form acids and bases.
When a nonmetal oxide such as SO2, NO2, or CO2 are combined with water the reaction
produces an acid. For example SO2 + H2O ✿H2SO3 or CO2 + H2O ✿H2CO3 . These reactions are
responsible for the acid precipitation, or acid rain, that leads to harmful affects to forests and
vegetation and the increased dissolving of limestone and marble statues and other structures.
When the oxides of metals react with water bases are formed; for example,
Na2O + H2O
2NaOH or CaO + H2O ✿Ca(OH)2.
This is why the chemical formula of acids usually only has nonmetals in the formula and the
chemical formula of bases contains a metal (H2CrO4 is chromic acid and it is an exception to
this pattern).
Acid and Base Solutions.
All acids react in a similar manner, despite the wide range in chemical formulas like HF,
H3PO3, or C7H6O3. This is because when an acid is dissolved in water, the reactive, a hydrogen
ion, H+, combines with water to make a hydronium ion, H3O+ : HA + H2O ✿A– + H3O+. The
hydronium ion is the reactive substance in an acid solution (there are substances called Lewis
acids that do not in this manner).
Many bases react with water to make a different reactive substance, hydroxide ion, OH–.
When the base has a chemical formula like BOH, the cation, B+, and the hydroxide ion, OH–,
separate to make the reactive OH–. When the base has a chemical formula like BCO3 or NR3 ,
these types of bases may react with water to form the hydroxide ion; for example, BCO3 + H2O
✿2 OH– + B+ + CO2 or NR3 + H2O ✿OH– + HNR3+, but they are also Brønsted-Lowry base that
don’t have to form OH–.
When an acid solution and a base solution combine, the two reactive ions, the hydroxide ion
and hydronium ion, combine to make water: H3O+ + OH– ✿ 2 H2O. The remaining ions, B+
and A–, form a salt which is dissolved but can be isolated by evaporating the water. When
this reaction occurs with acids and bases that form an ionic compound, called a salt, and water
the reaction is called a neutralization reaction.
The chemical equation for a reaction of an acid and a base shows the reactants and the
products but does not show the hydronium ion: HCl + NaOH ✿H2O + NaCl. When written
this way the neutralization reaction seems to be between the hydrogen ion, H+, and the
hydroxide ion, OH–: H+ + OH– ✿H2O. For this reason it is common to see the H+ replace the
H3O+ in acid reactions: HA ✿H+ + A–. Moreover, the hydrogen ion has lost its only electron,
so only one proton remains in the ion; therefore, H+ is commonly referred to as a proton. That
is why an acid can be defined as a proton donor. By the way, in the reaction above NaCl is the
salt.
Definitions of Acids and Bases
The earliest definition of acids and bases was developed by Svante Arrhenius. An Arrhenius
acid is a substance that forms H3O+ (or H+ in a shortened form) when added to water. An
Arrhenius base is a substance that forms OH– when added to water. Even though this covered
nearly all acid-base reactions, it was insufficient for reactions between acids and bases as gases
without water and for reactions of certain substances like NH3 that form very little OH– in
water and react directly with the acid: NH3 + HCl ✿NH4Cl.
Brønsted-Lowry acids and bases are defined more broadly, and this definition has become the
most common definition for an acid and base. A Brønsted-Lowry acid is a proton, H+, donor
and a Brønsted-Lowry proton acceptor. You must be able to identify the acid and base in a
chemical equation to use the Brønsted-Lowry definition.
Examples of acid and base reactions.
Acid-Base Reaction
Brønsted-Lowry
Acid
Brønsted-Lowry
Base
2 HF + Ca(OH)2 ✿2H2O + CaF2
HF
Ca(OH)2
Fe(OH)3 + H3PO4 ✿3H2O + FePO4
H3PO4
Fe(OH)3
Acid-Base Reaction
Brønsted-Lowry
Acid
Brønsted-Lowry
Base
H2S + NH3 ✿ NH4S
H2S
NH3
2 NaOH + H2CO3 + ✿Na2CO3 + H2O
H2CO3
NH(CH3)2
H2SO4 + Na2CO3 ✿Na2SO4 + H2O + CO2
H2SO4
Na2CO3
For these four reactions notice that the acid is not always at the beginning of the equation.
Most acids start with “H” although the number of H’s varies (the formula for an acid is based
on balancing charge between the H+ and the anion F– or PO43–). The bases have more varied
chemical formulae that include –OH, –CO3, and NR3 (where R is three connecting groups, in
particular R = H or CH3). However, what most of the bases do have in common is a metal
cation in the chemical formula (ammonia, NH3, is the exception). Finally, notice that while
most carbonates are bases, carbonic acid, H2CO3, ends in –CO3,
The products of an acid base reaction are usually water and an ionic compound called a salt.
For the five reactions above the salt is CaF2, FePO4, NH4S, Na2 CO3, and Na2SO4. For
carbonates, the reaction usually generates carbon dioxide, CO2. This is what occurs when
baking soda, NaHCO3 and vinegar, a solution containing acetic acid, HCH3COO:
NaHCO3 + HCH3COO ┪ NaCH3COO + H2O + CO2.
In the following equations the acids and bases can not be identified by common characteristics
in the chemical formulae. Instead, the identification requires using the Brønsted-Lowry
definition of an acid and a base and finding the proton, H+, donor and proton acceptor,
Acid-Base Reaction
Brønsted-Lowry
Acid
proton donor
Brønsted-Lowry
Base
proton acceptor
H2O + NH3 ┪ HO– + NH4+
H2O
NH3
H2SO3 + PO43– ┪ HSO3– + HPO42–
H2SO4
PO43–
H2S + HNO3 ✿ H3S+ + NO3–
HNO3
H2S
HSO3– + HCO3– ┪ SO32– + H2CO3
HSO3–
HCO3–
H2O + C6H5OH ┪ H3O+ + C6H5O–
C6H5OH
H2O
In the above reactions, it difficult to tell from inspecting the chemical formula of the reactants
which substance is the acid. An acid is not always the substance with the most H’s (H2O vs.
NH3 or H2S vs. HNO3), nor is it always the substance with an H in front (C6H5OH). To
determine which substance is an acid and which is a base, you have to decide which substance
is gaining an H+ and which is losing an H+ (H+ can be called a proton for these reactions). For
example, in the first reaction H2O becomes OH– in the reaction, this means H2O is an acid,
because it lost, or donated, an H+. NH3 in the same reaction becomes NH4+; thus it gains an H+
and is a base. In the last chemical equation, H2O becomes H3O+ so H2O is a proton acceptor
and a Brønsted-Lowry base; while C6H5OH is an acid, because it donates the H+. These two
examples, the first and last equations, show water can be either an acid or a base. Substances
that can act as either an acid or as a base are called amphoteric or amphiprotic. Other substances
like HSO3– , HCO3– , and H2PO4– are also amphoteric, since they have an H to donate and a “–”
charge that makes accepting an H+ energetically favorable.
Each acid-base reaction contains both an acid and a base. Check your identification of one by
identifying the other reactant; if you have identified one acid and one base then you have
probably made the correct identification. Moreover, all reactions conserve charge. By
transferring a hydrogen ion, H+, a “+” charge is transferred. But charge is conserved, because
the product of the acid has less “+” charge and the product of the base has more “+” charge in
exactly the same amount. For example, for the reaction PO43– + H2SO3 ┪ H2PO– + SO32– the
charge on the reactant side is 3– + 0 = 3– and the charge on the product side is 1– + 2– = 3–;
since both sides have the same charge, then charge is conserved.
Conjugate Acids and Conjugate Bases
The ┪ represent a solution in equilibrium (this is explored in more detail in another unit).
Essentially, it means that the reaction can go forward as written, but once products begin to
form, the reaction can go back to the reactants. The two reactions compete for as long as the
solution exists, but the amount of each substance reaches an equilibrium and no further change
in amount is observed (although the reaction is changing the individual pieces the amount
remains constant). The reverse reaction of an acid and base reaction also transfers an H+
between substances so the reverse reaction represents a second acid-base reaction. However,
since the reactants are labeled the acid and the base, different labels for the reverse reaction
starting materials are used. The conjugate acid is the acid of the reverse reaction and the
conjugate base is the base of the reverse reaction.
Examples of conjugate acids and conjugate bases.
Acid-Base Reaction
Acid
proton
donor
Base
proton
acceptor
Conjugate
Acid
Conjugate
Base
H2O + NH3 ┪ HO– + NH4+
H2O
NH3
NH4+
OH–
HSO3– + HCO3– ┪ SO32– + H2CO3
HSO3–
HCO3–
SO32–
H2CO3
H2O + C6H5OH ┪ H3O+ + C6H5O–
C6H5OH
H2O
H3O+
C6H5O–
Conjugate acids and bases are only present when a reaction is at equilibrium, which are
equations having the double arrows, ┪. An equilibrium solution forms when weak acids and
bases are the starting materials. So the reaction in the previous table between dihydrogen
sulfide and nitric acid, H2S + HNO3 ✿H3S+ + NO3– does not have a conjugate acid or conjugate
base because HNO3 is a strong acid.
It is also important to see the pattern between the chemical formulae of the original acid and
base and the conjugate acid and base. A conjugate acid forms from the original base when it
gains an H+, since the product with an extra H+ can be a proton, H+, donor, or an acid.
Likewise the acid forms the conjugate base, since it loses an H+ and the new substance will
accept that H+ back.
Strong and Weak Acids and Strong and Weak Bases.
There are only six common strong acids and all other acids are classified as weak. The six
strong acids to memorize are: HCl, HBr, HI, HNO3, H2SO4, and HClO4 .
Weak acids are not labeled weak because they are not reactive or dangerous. Concentrated
weak acids like HF and HCH3CO2, are extremely dangerous, but they are still weak. A weak
acid is an acid that remains, in large part, a molecule rather separate into the cation, H+, and
anion (remember, H+ always forms the hydronium ion, H3O+, in water). For example, 2.0 g of
HF, the weak acid hydrofluoric acid, in 1 liter of water forms only 8% H+, while 92% of the
original 2.0 g of HF remains a molecule in the solution. To show that the HF is the major
substance in the solution, the equation showing the dissolving of the acid (called the
dissociation of the acid) uses unequal arrows in the equilibrium chemical equation:
_✿
HF H+ + F–. Consequently, a strong acid dissociates 100% in water, for example
HCl ✿H+ + Cl–. A single arrow is used to denote the complete dissociation of a strong acid.
Bases are also classified as strong and weak. The strong bases are hydroxide compounds with
the alkali metals and some of the alkaline earth metals: LiOH, NaOH, KOH, RbOH, Mg(OH)2,
Ca(OH)2, & Ba(OH)2. All other bases are weak and remain mostly connected as molecules or
as complex ions (complex ions are metal ions combined with ligands like CN–, NH3, H2O, OH–
_✿
_✿
and more). For example, NH3 + H2O NH4+ + OH– or Al(OH)4– Al(OH)3 + OH–(Al(OH)4– is
a complex ion, a metal with ligands, it dissolves as individual unit; it is formed when the
crystal lattice of Al(OH)3 is dissolved in water).
Strong and weak acid solutions react in the same way, but the strong acid reacts more quickly.
This is because the reactive substance in an acid solution is the H3O+ or H+, and strong acids
dissociated 100% to give as much H3O+ as acid that was added. If a strong acid spills on your
hand, it will react on every surface of the skin and needs to be washed off quickly. Weak acids
react more slowly than strong acids so you may not even feel them, but can be equally or more
dangerous to work with because they still can form large amounts of H+, but not all at once.
The danger for hydrofluoric acid, HF, is that it does not react very much with your protective
layer of skin, but moves through the skin layer down into the bones where it cannot be
washed off. In the bones it dissociates slowly to reacts with the bones until the acid is used up
or the bones are dissolved.
Solubility of Hydroxides
Only the hydroxides of alkali metals are completely soluble, which means that the cation and
anion have separated to float freely in the solvent, usually water. Most hydroxides, even the
strong bases, Mg(OH)2, Ca(OH)2, and Ba(OH)2 are mostly insoluble. For example, less than 1%
of Mg(OH)2 will only dissolve in water no matter how much you put into the solution.
Nevertheless, the hydroxides of the alkaline earth metals are strong bases, because when they
do dissolve they dissociate completely (no complex ion is formed). Whereas weak bases that
are metal hydroxides, like Al(OH)3, will also barely dissolve and what does dissolve is mostly
found as the complex ion, like Al(OH)3(H2O)3 , and not free OH–.
pH Scale
The pH scale is a measure of the acidity of a
solution. pH stands for potential hydrogen and
is based on the moles of H+ in a 1 liter solution.
The pH scale, like the Richter scale, which
measures the magnitude of an earthquake, is a
logarithmic scale. pH equals –log[H+], where
the bracketed H+, [H+], is the moles of H+, in a
1 liter of solution (the unit for moles per liter is
M called molar). This calculation changes the
concentration of H+ from 1X10–1 M to
pH = 1 ( –log(1X10–1) = 1). The table to the
right shows more conversions, and it shows a
pattern that can be used to change from
concentration of H+ (more correctly H3O+) to
pH and visa versa. When the calculation is in
the form 1X10–x, then the pH = x, and if pH is x
then then [H+] = 1X10–x. Use the pattern to
find many answers, but when the numbers that
are not of the form 1 X 10–x, you’ll have to use a
calculator to calculate the answer.
Furthermore the table shows that when the pH
changes by 1 the concentration is multiplied or
divided by 10.
Solutions are acidic when pH < 7 and
solutions are basic when pH > 7. Any
solution that is exactly 7 is neutral. Solutions
with a pH of 1 or 2 have high concentrations
of H+ and are either concentrated or strong
acids. Solutions with a pH of about 4 are
dilute strong acids, with 1 mole of acid in
10,000 liters of solution, or they are weak
acid solutions with very little H+ dissociated
from the acid molecule. Solutions of bases
that are strong and have high concentrations
have pH’s of 13 or 14. Weak bases or dilute
strong base solutions have pH’s near 9.
[H3O+]
1 X 101
1 X 100
1 X 10–1
1 X 10–2
1 X 10–3
1 X 10–4
1 X 10–5
1 X 10–6
1 X 10–7
1 X 10–8
1 X 10–9
1 X 10–10
1 X 10–11
1 X 10–12
1 X 10–13
1 X 10–14
pH
-1
0
1
2
3
4
5
6
7
8
9
10
11
12
13
14
pOH
15
14
13
12
11
10
9
8
7
6
5
4
3
2
1
0
[OH–]
1 X 10–15
1 X 10–14
1 X 10–13
1 X 10–12
1 X 10–11
1 X 10–10
1 X 10–9
1 X 10–8
1 X 10–7
1 X 10–6
1 X 10–5
1 X 10–4
1 X 10–3
1 X 10–2
1 X 10–1
1 X 10–0
pOH is a similar calculation, –log[OH–].
This is a measurement of the basic
nature of the solution.
The pH of Common Solutions
Strongl
y Acidic
Weakly Acidic
Weakly
Basic
http://commons.wikimedia.org/wiki/File:PH_scale.
png
Strongl
y Basic
Calculation pH, pOH, [H3O+], and [OH–]
A neutral solution would be pure water with neither acid nor base added or a solution of
water with the salt of a strong acid and strong base. Careful measurements of water show
water dissociates to form H3O+ and OH–: 2 H2O ┪ H3O+ + OH– (or H2O ┪ H+ + OH–). The
dissociation is very limited and only 1 water molecule for every 10 million breaks apart, but
this equals 1X10–7 M H3O+, so water has a pH = 7. And since for every H3O+ there is an OH–
produced then the concentration of OH– is also 1X10–7 M. These measurements lead to the
water dissociation constant, KW: KW = [H3O+]•[OH–] = 1X10-7•1X10–7 = 1X10–14.
KW is a constant for all solutions—neutral, acid and base. So the equation KW = [H3O+]•[OH–]
will always give the product 1X10-14. Moreover, by taking the function –log to all parts of the
KW equation we get: –log { KW = [H3O+]•[OH–] = 1X10-7•1X10–7 = 1X10–14 } which becomes
–logKW = –log[H3O+] + –log[OH–] = –log(1X10-7) + –log(1X10–7) = –log(1X10–14) which becomes
pKW = pH + pOH = 7 + 7 = 14. This is another equation that is constant for all solutions:
pH + pOH = 14. Below are three example problems that show how to determine pH, pOH,
[H3O+], and [OH–] with the two equations: KW = [H3O+]•[OH–] & pH + pOH = 14.
Example Problems “Calculate the pH of a solution that has 1X10–4 M H3O+. Then calculate the
pOH and the concentration of OH–.”
pH = –log[H3O+] so pH = –log(1X10–4 ) = 4
pH + pOH = 14 so 4 + pOH = 14 and pOH = 10
pOH = –log[OH–] so 10 = –log[OH–] which gives [OH–] = 1X10–10 M.
“Find the concentration of [H3O+] for a solution that has a [OH–] = 1X 10–6 M. Determine the
pH and pOH for the solution too.”
KW = [H3O+]•[OH–] so 1X10–14 = x • 1X10–6 which is solved to give [H3O+] = 1X10–8 M.
pOH = –log[OH–] so pOH = –log(1X10–6) thus pOH = 6
pH + pOH = 14 so pH = 8.
If you need to carry out a calculation using your calculator because the number is not in the
format 1X10–x, such as 2.2X10–4 M or 0.000033 M, just type in log the number and ± to get the
pH or pOH If you have a pH or pOH like 4.5 or 9.33 then type in the number ± then 10x
(usually the inverse operation of log on the calculator) to get the [H3O+] or [OH–].
Titration
An important experimental procedure used to determine the concentration of an unknown
acid or base is a titration. A titration uses a buret, or another instrument that measures exact
volumes, to deliver a measured volume of a reactive solution to a known volume of another
solution. If one of the concentrations of the solutions is known then stoichiometry can be used
to determine the concentration of the other solution. For example, Given the reaction: 1A + 2B
✿ 1 C. If the acid, A, has an unknown concentration and the base, B has a concentration of
0.10M and if 22.2 mL of base are added with a buret to 10.0 mL of acid in a Erlenmeyer flask.
What is the concentration of A?
0.0222 L B • 0.10 moles • 1 mol A • 1 = 0.111 mol A/L A sol’n= 0.111 M A
1 L B sol’n 2 mol B
0.010 L B
For more worked problems and a simulation of a titration see:
Lecture and worksheet: http://www.saskschools.ca/curr_content/chem30_05/5_acids_bases/acids3_2.htm
Lecture & practice problems: Strong Acid and Strong Base Titration
(http://www.wwnorton.com/college/chemistry/gilbert2/tutorials/interface.asp?chapter=chapter_16&folder=st
rong_titrations ) & Titrations of Weak Acids
(http://www.wwnorton.com/college/chemistry/gilbert2/tutorials/interface.asp?chapter=chapter_16&folder=
weak_titrations )
Simulation:
http://www.chem.iastate.edu/group/Greenbowe/sections/projectfolder/flashfiles/stoichiometry/acid_base.ht
ml & http://lrs.ed.uiuc.edu/students/mihyewon/chemlab_experiment.html &
http://www.wfu.edu/~ylwong/chem/titrationsimulator/index.html
Solutions of Salts.
Most salt solutions are not neutral. They are either weakly acidic or basic. Only salts formed in
a neutralization reaction of a strong acid with a strong base give solutions with a pH of 7. For
example, NaOH + HCl ✿H2O + NaCl, the salt NaCl dissolves to make a neutral solution,
neither Na+ or Cl– will interact with water to change the pH.
However, the salt of a strong base and a weak acid forms a solution that is acidic. For
example, HF + NaOH ✿ H2O + NaF, the salt NaF dissolves to give both the sodium cation,
Na+, and the fluoride anion, F–. The anion, F–, is the conjugate base of the weak acid HF
(HF ┪ H+ + F–) so F– accepts a proton from water to reform the molecule:
F– + H2O ┪ HF + OH–. When this happens the solution has an increased concentration of OH–
and the solution has a pH > 7 and is basic.
When the salt of a strong acid combines with a weak base the solution is basic. So a salt like,
NH4Cl, dissolves then the NH4+ acts as a conjugate acid and donates an H+ to water:
NH4+ + H2O ┪ NH3 + H3O+. This gives an increase in the concentration of H3O+ and the
solution has a pH < 7 and is acidic. When a metal cation from a weak acid is dissolved in
water it forms a complex ion, like the solution of a salt containing Al3+ :
Al3+ + 7H2O ┪ [Al(H2O)5(OH)]2+ + H3O+. Complex ions differ between metals so determining
the identity of the complex ion is beyond the scope of this text, but be able to recognize that
metal cations of salts of strong acids and weak bases form acidic solutions. For information on
complex ions see http://www.docbrown.info/page07/appendixtrans01.htm .
Buffers
A buffer is a solution that resists changes in pH when acid, base, or neutral solutions are
added. Buffers can be acidic or basic and are very important in natural systems. The human
digestive system uses enzymes that require specific pH’s: a neutral pH in the mouth for
amylase, an acidic pH in the stomach for pepsin, and a basic pH in the intestines for lipase. In
the blood, carbonic acid, formed from absorbed carbon dioxide, is part of the buffer system
(recall nonmetal oxides dissolved in water lead to acids). The blood buffer is one of the most
important buffers since the pH of blood must remain between 7.0 and 7.8 or the circulatory
system is compromised, which can lead to death.
In order to maintain a specific pH, a buffer must contain a proton acceptor and a proton donor.
If the proton acceptor is a strong base, then it will accept protons even from water and the pH
will be high and stay there. Similarly, solutions with a strong acid will donate protons
indiscriminately and solutions will have low pH’s. A good buffer contains a weak acid and a
weak base in equilibrium with both substances acting to maintain the concentration of H3O+
and the pH. The most common way of creating a buffer is to use a weak base and a salt with
the base’s cation, which will act as a weak acid, or to use a weak acid and a salt with the acid’s
anion, which will act as a weak base.
In the three examples of buffer solutions a weak acid is on the left side of the equilibrium
arrows, ┪ , and a weak base on the right. Both substances are present in solution, so the buffer
solution has a base to counteract added H+ and an acid to counteract OH– and the pH will
remain constant.
1) In blood the buffer equilibrium is HCO3– + H+ ┪ H2CO3 , which has a weak acid, H2CO3,
and the anion that will act as a base, HCO3–.
2) For the cytoplasm in and around cells, the buffer equilibrium is H2PO4– ┪ H+ + HPO42- ,
which has a weak acid, H2PO4– and the anion HPO42– that is the conjugate base.
3) A weak base buffer system with a pH near 10 uses ammonia, NH3, as a base and an
ammonium, NH4+, salt as an acid: NH4+ + OH– ┪ NH3 + H2O. This is similar to how amines
like in amino acids can help maintain the pH.
All buffer systems have a buffering capacity, which indicates how much acid or base can be
added before the buffering system is no longer effective, since much of the weak base or weak
acid has reacted with the added substance.
Lewis Acids and Bases
While classifying reactions, chemists noticed that some reactions don’t transfer protons, H+,
but also don’t transfer electrons like redox reactions. Rather, they react by donating and
accepting electrons, but the electrons are not transferred completely to another atom. The
most common example of this type of reaction is NH3 + BF3 ✿H3N–BF3. Here’s the Lewis
structures so that you can see the electrons:
http://commons.wikimedia.org/wiki/File:NH3-BF3-adduct-bond-lengthening-2D.png
Examining the Brønsted-Lowry acid and base interaction shows a similar movement of
electrons. For example
H
H
|
••
|
••
+
H—N : __ H∇O∇H
_
H∇N∇H
+
:
O—H
••
••
|
|
H
H
As the H+ is transferred, the
–
electrons are donated by the N
and accepted by the O of H2O.
Unlike a redox reaction, like Zn + Cu2+ ✿Zn2+ + Cu, where electrons are transferred from one
atom to another, the electrons involved are bonding electrons. In the ammonia, NH3, example
the nonbonding electrons on N become bonded electrons between N and H, and the bonding
electrons between H and O on water become nonbonding electrons. In this way, the definition
of acid and bases expanded to include the NH3 + BF3 example and others that are called Lewis
acids and Lewis bases.
Lewis acids are electron acceptors (the opposite of proton donors which is the Brønsted-Lowry
acid definition). Lewis bases are electron donors. All Arrhenius acids are Brønsted-Lowry
acids and all Brønsted-Lowry acids are Lewis acids. But, the opposite is not true. So not all
Lewis acids are Brønsted-Lowry acids and not all Arrhenius acids are Brønsted-Lowry acids.
For another discussion, problems, and links see:
http://chemwiki.ucdavis.edu/Physical_Chemistry/Acids_and_Bases/Acid/Lewis_Concept_
of_Acids_and_Bases .
Summary
Acid-base reactions are found throughout nature and are important reactions in industry and
for use in household products like cleansers and vitamins. Acids taste sour, react with metals
to produce hydrogen gas, H2, react with carbonates to produce carbon dioxide, react with
litmus indicator to give a red color and with phenophthalein to give a colorless solution.
Acids react with basic solutions to make the solution more neutral or acidic. Bases taste bitter,
react with fats or oils to produce soap, react with litmus indicator to give a blue color and with
phenophthalein to give a magenta color. They also react with acidic solutions to make the
solutions more neutral or basic.
Six strong acids—HCl, HBr, HI, HNO3, H2SO4, and HClO4—show the common chemical
formula of acids but the most common definition of acids, the Brønsted-Lowry definition,
states that acids are substances that can donate a proton: HCl + OH– ✿Cl– + HOH. The seven
strong bases—LiOH, NaOH, KOH, RbOH, Mg(OH)2, Ca(OH)2, and Ba(OH)2—also show a
common chemical formula (OH), but a Brønsted-Lowry base is a substance that accepts a
proton: NaOH + H3O+ ✿Na+ + 2 H2O.
Any acid that is not a strong acid will be a weak acid such as HF, hydrofluoric acid, and
HCH3COO, acetic acid. Weak acids react more slowly than strong acids because they remain
mostly in their molecular form in solution and the reactive H+ is released over time instead of
_✿
all at once like the strong acid: HF H+ + F–. Moreover, weak bases are any base that is not a
strong base like NH3, ammonia, and NaHCO3, sodium hydrogen carbonate. The weak base
also partially dissociates in solution.
Reactions of weak acid and weak base are in equilibrium. This means that the products react in
a reverse reaction to reform the reactants. The reverse reaction is also an acid-base reaction, so
the products are classified as the conjugate acid, which is formed from the base, and the
conjugate base, which is formed from the acid. Solutions of weak acids and their salts or
weak bases and their salts are in equilibrium so that the solution contains both acids and bases.
These solutions are called buffers and maintain a constant acidity or pH.
pH is a measure of the acidity of a solution. The formula for pH is: -log[H+] = pH, where
[H+] is the concentration of hydrogen ion or hydronium ion, H3O+. Solutions are acidic when
the pH < 7 and basic when pH > 7. Solutions of salts of strong acids and strong bases like
NaCl are neutral, as is pure water, but salts from weak acids and weak bases will be basic or
acidic respectively.
Equations used to determine pH, pOH (-log[OH] = pOH), [H3O+]. and [OH–] when one of
these amounts is given are: KW = 1X10–14 = [H3O+]•[OH–] and 14 = pH + pOH.
Titrations are an important experimental procedure used to determine unknown concentration
of an acid or a base.
Simulations of a titration can be found at:
http://www.chem.iastate.edu/group/Greenbowe/sections/projectfolder/flashfiles/stoichio
metry/acid_base.html &
http://www.wfu.edu/~ylwong/chem/titrationsimulator/index.html
Acids and Bases
5.
Acids, bases, and salts are three classes of compounds that form ions in water solutions.
As a basis for understanding this concept:
a.Students know the observable properties of acids, bases, and salt solutions.
b.Students know acids are hydrogen-ion-donating and bases are hydrogen-ion-accepting
substances.
c.Students know strong acids and bases fully dissociate and weak acids and bases partially
dissociate.
d.Students know how to use the pH scale to characterize acid and base solutions.
e.* Students know the Arrhenius, Bronsted-Lowry, and Lewis acid-base definitions.
f. * Students know how to calculate pH from the hydrogen-ion concentration.
g.* Students know buffers stabilize pH in acid-base reactions.
Starred standards are non-tested standards on the California Standards Test
Contributed by Kenneth Pringle
Edited by Kathleen Duhl
Formatted and Wiki Contribution by Christine Mytko