I. Acid and Base Theory
Acids and bases are among the most important solutes in solution. They are compounds which have certain characteristic properties in common. In general, acids are compounds which, when dissolved in water,
* form solutions that conduct electricity
* react with metals like zinc and magnesium to produce salts and a gas - usually hydrogen
* taste sour (but not all are safe to taste)
* turn blue litmus paper to red
* react with bases to form salt and water.
On the other hand, bases are compounds which, when dissolve in water,
* yield solutions that also conduct electric current
* feel slippery or soapy on the skin
* taste bitter (but not all are safe to taste)
* turn red litmus paper to blue
* reacts with acids to form salt and water.
* form solutions that conduct electricity
* react with metals like zinc and magnesium to produce salts and a gas - usually hydrogen
* taste sour (but not all are safe to taste)
* turn blue litmus paper to red
* react with bases to form salt and water.
On the other hand, bases are compounds which, when dissolve in water,
* yield solutions that also conduct electric current
* feel slippery or soapy on the skin
* taste bitter (but not all are safe to taste)
* turn red litmus paper to blue
* reacts with acids to form salt and water.
II. Conceptual Definition of Acids and bases
a. Arrhenius Theory
Proposed by Swedish physical chemist Svante Augustus Arrhenius (1859-1927). His study of electrolytic dissociation that in an aqueous solution, a strong electrolyte exists partly as ions and partly as molecules. Three year later, he extended this theory by suggesting that acids are neutral compounds that ionize when dissolved in water to give hydrogen ions and a corresponding negative ion. Bases are neutral compounds that either dissociate or ionize in water to give hydroxide ions and a positive ion.
Arrhenius acid is any substance that ionizes when it dissolves in water to give H+.
HCl ---> H+(aq) + Cl-(aq)
Arrhenius base is any substance that gives OH- when it dissolves in water.
NaOH ---> Na+(aq) + OH-
Note: this theory has several disadvantages. It is only applicable for aqueous solutions.
b. Bronsted-Lowry Theory
Johannes Bronsted (1879-1947) and Thomas Lowry (1874-1936) suggested one chabge was the existence of hydronium ion. They proposed that hydronium ions cannot exist in water but rather covalently bond with water. They based their definition of acids and bases on the behavior of hydrogen ions.
Bronsted-Lowry acid proton donor [H+]
Bronsted-Lowry base proton acceptor [H-]
Note: protons are not actually being given but rather the bond between the proton and another element is being broken to form a new bond to the base.
This new definition broadens the range of acids and bases but still requires a protic acid.
Conjugate base (base 1) is the species remaining when a proton is removed from an acid.
Conjugate acid (acid 2) is the species formed when a proton is transferred to a base.
Conjugate acid-base is consists of two substances related to acids after by donating and accepting of a single proton.
Example;
1. CH3 COOH(aq) + H2O(l) yields to CH3Coo-(aq) + HO-(aq)
acid 1 base 2 base 1 acid 2
2. HClO3(aq) + NH3(aq) yields to Clo-3(aq) + NH+4(aq)
acid 1 base 2 base 1 acid 2
c. Lewis Theory
This theory proposed by Gilbert Lewis further widens the classification of acids and bases to non-aqueous solutions. He proposed that the electrons and not the protons are the ones transferring. The acids and bases can be charged or unionized.
Lewis Acid is electron pair acceptor
Lewis Base is electron pair donor
Example;
1. Ag- + NH3 ---> [Ag:NH3]+
acid base
2. BCl3 + Cl- ---> [BCl4]-
acid base
III. Classification of Acids according to the number of protons being dissociated
Certain compounds have more than one proton to give. These are called polyprotic acids. Some acids like sulfurous acid (H2SO3) possess two protons making them diprotic while some, like phosphoric acid (H3PO4), have three making them tripotic acids. It is important to discuss these acids because like water and other special compounds, their conjugates are amphoteric. Ampotherism is the capacity of a compound to acts as an acid or as a base. A good example of an amphoteric compound is water. Water can self-ionize by acting as an acid and a base. This process is called auto-ionization.
2H2O(l) yields to H3O+(aq) + OH-(aq)
CH3CH2COOH is an example of propanoic acid, the [H+] can be found in the last part of the equation.
Ionization of a Polyprotic Acid
H3PO4 ---> H+ + H2PO4-
H2PO- ---> H+ + HPO42-
HPO42- ---> H+ + PO43-
H3PO4 ---> 3H+ + PO43- -Net Ionization Reaction
IV. Strong and Weak Acids and Bases
The degree of ionization, not the concentration of an acid or a base, classifies it into weak or strong. Strong acids and bases completely ionize to produce proton and conjugate equal to its initial concentration. For example,
HCl(aq) ---> H3O+(aq) + Cl-(aq)
Where [HCl] = 1.0M [H3O+] = [Cl-] = 1.0 M
Unlike their counterparts, weak acids and bases have a more complex equation in determining the concentration of hydronium ions.
Common Weak Acids
Formic HCOOH
Acetic CH3COOH
Trichloroacetic CCl3COOH
Hydrofluoric HF
Hydrocyanic HCN
Hydrogen sulfide H2S
Water H2O
Conjugate acids of weak bases NH4+
Common Weak Bases
Ammonia NH3
Trimethyl ammonia N(CH3)3
Pyridine C5H5N
Ammonium hydroxide NH4OH
Water H2O
HS- ion HS-
Conjugate bases of weak acids HCOO-
Strong acids
Hydrogen halides HCl HBr HI
Oxyacids of halogens, HClO3, HClO4, HBrO3, HBrO4, HIO3, HIO4
Sulfuric acid H2SO4
Nitric acid HNO3
Strong bases
Sodium hydroxide NaOH
Potassium hydroxide KOH
Cesium hydroxide CsOH
Calcium hydroxide Ca(OH)
V. Common Acid, Base, and Salt and their uses
Common Acids and their Uses
1. Sulfuric Acid (H2SO4) –acid found in automobile battery, used as dehydrating agent.
2. Hydrochloric Acid (HCl) –gastric juice in the stomach, used to removed stain and paint from metals and concrete.
3. Nitric Acid (HNO3) –used to make fertilizers and explosives.
4. Phosphoric Acid (H3PO4) –used in dilute form in soft drinks and detergents and fertilizers.
5. Acetic Acid (CH3COOH) – acid in vinegar is dilute acetic acid.
6. Citric Acid (C6H8O7) –the acid in citrus fruits.
7. Carbonic Acid (H2CO3) –formulation of carbonic drinks
Common Bases and their Uses
1. Sodium hydroxide (NaOH) –known as lye; used in soap manufacture and paper production; drain cleaner.
2. Potassium hydroxide (KOH) –like NaOH, it is strong base and used to dissolved grease and hair in clogged drainers.
3. Magnesium hydroxide (Mg (OH) 2) –antacid with no dosage restriction.
4. Ammonium hydroxide (Al (OH) 3) –revive patients who’ve fainted fertilizer production.
Common Salts and their Uses
1. Sodium Chloride (NaCl) –common table salt.
2. Zinc Oxide (ZnO) –pigment, cosmetic, molds inhibitor.
3. Potassium Iodide (KI) _found in Iodized salt, source of Iodine
4. Tin (II) Fluoride (SnF2) –found in toothpaste, source of Iodine.
5. Cadmium Sulfide (CdS) –solar cells, phosphors, pigment.
6. Magnesium Carbonate (MgCO3) –antacid.
7. Iron (II) Sulfate (FeSO4) –medicine, source of iron.
8. Sodium Nitrate (NaNO2) –for curing meat (salitre).
9. Calcium Carbonate (CaCO3) –marble, limestone.
VI. Acid-Base Equilibria
a. pH and pOH
pH Concept
-the symbol pH, stands for some German words which literally means “the power of Hydrogen ion”, and defined as the negative logarithm of the hydronium ion concentration.
Soren Sorensen
-Danish Biochemist who proposed the pH scale as a more convenient way of expressing hydronium ion concentration.
Acid-Base Equilibria
H2O yields to H+ + OH- or 2H2O yields to H3o+ + OH-
Ionization Constant of H2O (kw)
-also known as equilibrium constant of ion-product constant
Kw = [H+] [OH-] or kw = 1.1 x 10-14M
pH –a measure of acidity
pOH –a measure of basicity or alkalinity
Ways of expressing acidity and basicity in terms of pH and pOH
pH + pOH = 14
10(-pH) = [H+]
10(-pOH) = [OH-]
-log([H+]) = pH
-log([OH-]) = pOH
Since pH is simply a way to express hydrogen ion concentration, acidic and basic solutions at 25 degrees Celsius can be determined by their pH values, as follows
Acidic Solutions = [H+] > 1.0 x 10-7 M, pH < 7.00
Basic Solutions = [H+] < 1.0 x 10-7 M, pH > 7.00
Neutral Solutions = [H+] = 1.0 x 10-7 M, pH = 7.00
Notice that pH increases as [H+] decreases
b. Neutralization and Titration
The process of neutralization is a method of producing a neutral solution by mixing acid and base. The equation simplifies the process of neutralization.
MoleAcid = Molebase or MaVa = MbVb
Volumetric Analysis
-the technique in determining the concentration of a second reactant if its volume is known and the 1st reactant has given concentration and volume.
Titration
-is a process of measuring the concentration of an acid or base in one solution by adding a base or acid solution of known concentration.
Titrant – known solution
-added slowly to a measured volume of an acid or a base solutions of unknown concentration.
Neutralization point/Endpoint/Equivalence point
-the point at which the titrant is enough to fully neutralize the acid or base.
Buffer
-is a solution consisting of a weak acid and its conjugate base, or of a weak base and its conjugate acid. It resists a change in pH when a moderate amount of acid or a base is added to it.
Example problem;
A 25 ml H2SO4 solution requires 32.58 ml of 0.5 NaOH to neutralize it. What is the molarity of H2SO4?
H2SO4 + 2NaOH ---> 2H2O + Na2SO4 1: 2
Given;
VA= 25 ml, MA = ?,
VB = 32.58 ml, MB = 0.5 M
Nb = (0.03258 L) (0.5 M) = 0.01629 mol NaOH
MA = 0.01629 mol (NaOH) x (1H2SO4 / 2NaOH) = 0.3258 M
