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 / Acid / Base /Chemical Formula is Recognizable / H–X
for example, HCl, HNO3, HC2H3O2, H2SO4, & H3PO4,
Most acids have only nonmetal elements in the chemical formula / Y–OH, Y-CO3 , & NR3
for example, NaOH, Ca(OH)2, Al(OH)3, NaHCO3, CaCO3,
NH3, & NH(CH3)2
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 carbondioxide 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 H2Sisdihydrogen sulfide)
b. An acid with a polyatomic ion ending in -ate is called: element name-ic acid; HNO3=nitricacid (NO3– is nitrate), H2SO4 = sulfuric acid (SO42– is sulfate), H2CO3=carbonicacid (CO32– is carbonate), H3PO4=phophoric acid (PO43– is phosphate), HCH3COO =acetic acid (CH3COO– is acetate, which can be written as, C2H3O2–orCH3CO2–).
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=nitrousacid;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 ➔ 2OH– + 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
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:
HFH++ 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.