5.1
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5.1 Identifying Acids and Bases Warm Up A student tested a number of unknown solutions and recorded the following observations. Based on each observation, place a check mark in the corresponding column that identifies the unknown solution as containing an acid, a base, or either one. Observation
Acid
Base
Turns phenolphthalein pink Feels slippery Has pH = 5.0 Tastes sour Conducts electricity Reacts with metal to produce a gas
The Arrhenius Theory of Acids and Bases
In previous years, you learned how to identify an acid or base using a concept developed by a chemist named Svante Arrhenius. According to the Arrhenius theory, acids release H+ ions in solution, and bases release OH– ions. Typically, when an Arrhenius acid and base react together, a salt and water form. A salt is an ionic compound that does not contain H+ or OH− ions. The cation from the base and the anion from the acid make up the salt. Example: HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l) acid base salt water
Sample Problem 5.1.1 — Identifying Arrhenius Acids and Bases Classify each of the following substances as an Arrhenius acid, an Arrhenius base, a salt or a molecular compound. (a) HNO3 (b) Al(OH)3 (c) Al(NO3)3 (d) NO2
What to Think About
How to Do It
(a) An acid releases H+ ions.
HNO3 is an acid.
(b) A base releases OH− ions.
Al(OH)3 is a base.
(c) A salt is an ionic compound not containing H+ or OH− ions.
Al(NO3)3 is a salt.
(d) A molecule is not made up of ions.
NO2 is a molecular compound.
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Practice Problems 5.1.1 — Identifying Arrhenius Acids and Bases 1. Classify each of the following as an Arrhenius acid, an Arrhenius base, a salt, or a molecular compound. (a) H2SO4
___________________________________________
(b) XeF6 ___________________________________________ (c) CH3COOH __________________________________________ (d) NaCH3COO _________________________________________ (e) KOH
___________________________________________
(f) NH3 ___________________________________________ 2. Complete the following neutralization equations. Make sure each equation is balanced, and circle the salt produced. (a) CH3COOH + LiOH → (b) HI + Ca(OH)2 → (c) Mg(OH)2 + H3PO4 → 3. Write the formula of the parent acid and the parent base that react to form each salt listed. Parent Acid Parent Base (a) KNO2 (b) NH4Cl (c) CuC2O4 (d) NaCH3COO
Brønsted-Lowry Acids and Bases
Arrhenius’ definition worked well to classify a number of substances that displayed acidic or basic characteristics. However, some substances that acted like acids or bases could not be classified using this definition. A broader definition of acids and bases was required. In the practice problem above, you may have classified NH3 as molecular. While it is molecular, a solution of NH3 also feels slippery, has a pH greater than 7, and turns phenolphthalein pink. It clearly has basic characteristics, but it is not an Arrhenius base. Chemists Johannes Brønsted and Thomas Lowry suggested a broader definition of acids and bases: Brønsted-Lowry acid — a substance or species that donates a hydrogen ion, H+ (a proton) Brønsted-Lowry base — a substance or species that accepts a hydrogen ion, H+
Consider the reaction that occurs in aqueous HCl: HCl(aq) + H2O(l) → H3O+(aq) + Cl−(aq) acid base hydronium In this example, the HCl donates a hydrogen ion (H+) to the water molecule. The HCl is therefore acting as a Brønsted-Lowry acid. The water accepts the H+ ion so it is acting as a Brønsted-Lowry base.
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Chapter 5 Acid-Base Equilibrium 265
The H3O+ ion is called a hydronium ion. It is simply a water molecule with an extra H+ ion (Figure 5.1.1). It is also called a protonated water molecule.
hydrogen ion (H+)
water (H2O)
hydronium (H3O+)
before
after
Figure 5.1.1 A hydronium ion is a water molecule with an
extra H+ ion.
Let’s look at NH3 now. In a solution of ammonia (NH3) the following reaction occurs: NH3(aq) + H2O(l) base acid
NH4+(aq) + OH−(aq)
The ammonia molecule accepted a H+ ion from the water, so NH3 is acting as a Brønsted-Lowry base, and water is acting as a Brønsted-Lowry acid. You may notice that in the equation for HCl, a one-way arrow was used, but in the equation for NH3, an equilibrium arrow was used. The reasons for this will be explained in the next section. Just remember that if you see an equilibrium arrow, then equilibrium is established. More importantly, in an equilibrium, you have both reactants and products present in the system. Both forward and reverse reactions occur. If we look at the reverse reaction for the ammonia system, we can identify BrønstedLowry acids and bases: NH3(aq) + H2O(l) NH4+(aq) + OH−(aq) base acid acid base In the reverse reaction, the NH4+ ion donates a proton to the OH− ion. According to BrønstedLowry definitions, the NH4+ ion is acting as an acid, and the OH− ion is acting as a base. In a Brønsted-Lowry equilibrium, there are two acids and two bases: one for the forward reaction, and one for the reverse reaction. An acid and a base react to form a different acid and base. Two substances that differ by one H+ ion are called a conjugate acid-base pair.
In the example above, the NH3 and NH4+ together form one conjugate acid-base pair, and the H2O and OH− form the other conjugate acid-base pair.
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Sample Problems 5.1.2 — Identifying Conjugate Acid-Base Pairs 1. In the following equilibrium, identify the acids and bases, and the two conjugate acid-base pairs: HF(aq) + CN−(aq)
HCN(aq) + F−(aq)
2. Complete the following table: Conjugate Acid
Conjugate Base
H2C2O4 SO32– HCO3− H2O 3. Complete the following equilibrium, which represents the reaction of a Brønsted-Lowry acid and base. Circle the substances that make up one of the conjugate acid-base pairs. NO2− (aq) + H2CO3(aq)
What to Think About 1. An acid donates a proton and a base accepts a proton.
In the forward reaction, the HF is the acid and the CN− is the base. In the reverse reaction, the HCN is the acid and the F− is the base.
The substances in a conjugate acid-base pair differ by one H+ ion.
How to Do It
The two conjugate acid-base pairs are: HF/F− and CN−/HCN. HF(aq) + CN−(aq) HCN(aq) + F−(aq) acid base acid base
2. An acid donates a proton. Find the conjugate base of an acid by removing one H+. Be careful of your charges on the ions! A base accepts a proton. Find the conjugate acid of a given base by adding one H+.
Conjugate acid base H2C2O4 → − HSO3 ← HCO3− → + H3O ←
3. An acid must be able to donate a H+ ion. Only the H2CO3 has a H+ to donate. When it donates a H+ ion, HCO3− forms. The NO2− ion is forced to accept the H+, making the NO2− a base. When the NO2− accepts a proton, it forms HNO2. The total charge on the reactant side should equal the total charge on the product side. Balance reactions for both. Note: H2CO3 only donates ONE H+ ion. Substances in a conjugate acid-base pair differ by only ONE H+ ion. Balance both sides of the equation for number of atoms and charge.
NO2−(aq) + H2CO3(aq) HNO2(aq) + HCO3−(aq)
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Conjugate remove H+ add H+ remove H+ add H+
HC2O4− SO32– CO32– H2O
Circle around either of the following: NO2−/HNO2 or H2CO3/HCO3−
Chapter 5 Acid-Base Equilibrium 267
Practice Problems 5.1.2 — Identifying Conjugate Acid-Base Pairs 1. For the following equilibria, label the acids and bases for the forward and reverse reactions. (a) HIO3 + NO2−
(c) Al(H2O)63+ + SO32–
HNO2 + IO3−
(b) HF + HC2O4−
H2C2O4 + F−
HSO3− + Al(H2O)5OH2+
2. Complete the following table: Conjugate Acid H2O2
Conjugate Base H2BO3−
HCOOH C6H5O73– 3. Complete the following equilibria. Label the acids and bases for the forward and reverse reactions. Circle one conjugate acidbase pair in each equilibria. (a) HNO2 + NH3
(b) H3C6H5O7 + CN−
(c) PO43– + H2S
Consider the two equilibria below:
Amphiprotic Species
NH3(aq) + H2O(l) NH4+(aq) + OH− (aq) HF(aq) + H2O(l) acid base
H3O+(aq) + F− (aq)
In the first reaction, water acts as a Brønsted-Lowry acid. In the second reaction, water acts as a Brønsted-Lowry base. An amphiprotic substance has the ability to act as an acid or a base, depending on what it is reacting with. Water is a common amphiprotic substance. Many anions also display amphiprotic tendencies. For a substance to be amphiprotic, it must have a proton to donate and be able to accept a proton. Examples of amphiprotic anions include HCO3−, HC2O4−, and H2PO4−. Uncharged species, with the exception of water, are generally not amphiprotic. For example, HCl will donate one H+ ion (proton) to form Cl−, but will not accept a proton to form H2Cl+. You should recognize many of the species that form.
Quick Check 1. Write an equation for a reaction between HCO3− and CN− where HCO3− acts as an acid.
2. Write an equation for a reaction between HCO3− and H2O where HCO3− acts as a base.
3. Circle amphiprotic substances in the following list: (a) CH3COOH (b) H2PO4−
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(c) PO43–
(d) H2C2O4
(e) HC2O4−
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