Unit 7 Study Guide and Review: Acid-Base and Oxidation-Reduction Reactions

7a: Use the Arrhenius and Brønsted–Lowry definitions to identify acids and bases.

There are more than one definition of acids and bases which can be useful in different situations.

    • Define Arrhenius acids and bases
    • Use the Arrhenius definition to identify acids and bases
    • Define Bronsted–Lowry acids and bases
    • Use the Bronsted–Lowry definition to identify acids and bases

An Arrhenius acid is a substance with at least one hydrogen that can dissociate or ionize when dissolved in water. This produces a hydrated hydrogen ion and a counter ion. An Arrhenius base is a substance that has at least one hydroxide (OH) group that can dissociate or ionize when dissolved in water.

      • Examples of Arrhenius acids include hydrochloric acid, HCl, sulfuric acid, H2SO4, and acetic acid (vinegar), CH3COOH.

Review the table in Section 2: Acids and the Hydrogen Ion in Acids and Bases: An Introduction for the dissociation reactions of these Arrhenius acids in water. You will clearly see the dissociated hydrogen ion in the products.

      • Examples of Arrhenius bases include sodium hydroxide (also known as lye), NaOH, and ammonia, NH3.

To review the dissociation reactions of these bases see Equations (3–1) and (3–3) in Section 3: What is a Base?, in Acids and Bases: An Introduction. You will clearly see the dissociated hydroxide ion in the products.

The Bronsted–Lowry definition of acids and bases is more general than the Arrhenius definition.

      • A Bronsted–Lowry acid is a proton (hydrogen ion) donator.
      • A Bronsted–Lowry base is a proton (hydrogen ion) acceptor. This definition is important because it shows that acids and bases must exist together. An acid cannot donate a proton without a base to accept it. Consequently, acids and bases always exist in pairs.

In cases of an acid, such as HCl in water, the HCl acts as the acid and the water acts as the base.

HCl(aq) + H2O → Cl(aq) + H3O+(aq)

Here, HCl is the acid because it gives a proton to the water, and becomes a Cl ion in the products. The water is the base because it accepts the proton and becomes H3O+, which is called the hydronium ion.

In cases such as NH3 in water, the NH3 acts as the base and the water acts as the acid.

NH3(aq) + H2O → NH4+(aq) + OH(aq)

Here, NH3 is the base because it accepts a proton from the water to become NH4+ in the products. The water is the acid because it donates a proton to the ammonia and becomes the hydroxide ion, OH.


7b: Write and balance equations for neutralization reactions.

Neutralization reactions are the simplest type of acid base reaction. In this type of reaction, an acid and a base react with each other to form water and a salt.

    • Write and balance a neutralization reaction for a given acid and base.
    • Define salt.

An acid and base will react to form water and a salt. A salt is an ionic compound, containing the cation of the base and the anion of the acid.

To write a neutralization reaction, begin by writing the molecular reaction. For an example, let’s write the neutralization reaction of HCl and NaOH.

We put the acid and base on the reactant side, with their states of matter.

HCl(aq) + NaOH(aq) →

We know that HCl is the acid because it has a proton that can dissociate. We know NaOH is the base because it has the OH that can dissociate. Therefore, we know the H + from the HCl will dissociate and the OH from the NaOH will dissociate. When H + and OH come together, they form water. The leftover parts are the Cl from the acid and the Na+ from the base. These come together to form NaCl.

HCl(aq) + NaOH(aq) → NaCl(aq) + H2O

Now, we can convert our reaction to what we call a net ionic equation. In a net ionic equation we eliminate ions that appear on both sides of the equation, so we only see the net change in the reaction. Here, all aqueous (aq) substances are known to dissociate into their ions in solution. Note that “aqueous substances” are substances that are dissolved in water.

First, we write a total ionic equation which includes each individual ion in the equation:

H+(aq) + Cl(aq)+ Na+(aq)+ OH(aq)→ Na+(aq)+ Cl(aq)+ H2O

We can see that Cl and Na+ appear on both sides of the equation. Just like in math, that means we can cancel them. This gets us to our net ionic equation for this reaction:

H+(aq) + OH (aq) → H2O

Review a step-by-step example of how to balance an acid–base neutralization reaction in Section 4: Neutralization: The Simplest Acid–Base Reaction, in Acids and Bases: An Introduction.


7c: Explain the process of self-ionization of water molecules.

Water is a unique molecule in many ways. One of its important properties that we saw above is that it can behave as either an acid or a base.

    • Write an equation for the self-ionization or autoprotolysis of water.
    • How is water both an acid and a base?

In the reaction of HCl with water, we saw that water acted as a base, accepting the proton from HCl to form the hydronium ion.

HCl(aq) + H2O → Cl (aq) + H3O+ (aq)

In the reaction of ammonia with water, we saw that water acted as an acid, donating a proton to the ammonia to form the hydroxide ion.

NH3(aq) + H2O → NH4+(aq) + OH (aq)

Notice that, since water can be both an acid and a base, water can react with itself in an acid–base reaction to form hydronium ion and hydroxide ion.

This is called the self-ionization of water or the self-protolysis of water:

H2O + H2O → H3O+(aq) + OH(aq)

In the self-ionization of water, one water molecule in the reactants serves as an acid, or proton donor. The other water molecule acts as a base, or proton acceptor.

Note that the proton from the water always has become the hydronium ion rather than the H+ ion in solution.

Review a detailed description of this important ion in Section 2: The Hydronium Ion, in Proton Donors and Acceptors: Acid–base Reactions à la Brønsted.


7d: Conduct pH calculations and use pH scale to classify solutions as acidic, basic, or neutral.

The pH scale gives us information about how acidic or basic a solution is.

    • Write the ion product of water.
    • Describe acidic, basic, and neutral solutions based on the [H3O+] and [OH].
    • Determine the concentration of [H3O+] or [OH] for a given solution.
    • Calculate pH for a given solution.
    • Use pH to classify a solution as acidic, neutral, or basic.

As we learned above, water self-ionizes or dissociates into the hydronium ion and hydroxide ion. This only occurs to a small extent. The majority of water stays as water, but a small percent does dissociate.

We can write this as the ion product of water, Kw: [H+][OH] = 1.00 x 10–14, where the amounts in the square brackets are concentrations of the ions.

We can use the ion product of water to determine if a solution is acidic, neutral, or basic:

Acidic

[H+] > [OH]

Neutral

[H+] = [OH]

Basic

[H+] < [OH]


This chart is reproduced from Section 1: The Dissociation of Water, in pH and Titration.

If we know the concentration of acid or base in a solution, we can determine the concentration of the H+ and OH ions by using the Kw expression.

Review the example involving HCl at the end of, Section 1: The Dissociation of Water, in pH and Titration.

Using Kw, we can derive the equation for pH: the common scale used to determine if a solution is acidic, neutral, or basic.

There are two important equations from pH that you should know (see Equations 2–1a and 2–1b in pH and Titration.):

pH = –log [H+] and [H+] = 10–pH

If we know the [H+], we can easily determine pH. If we know the pH, we can easily determine [H+].

Another important scale is the pOH scale. This is analogous to the pH scale, except with hydroxide ions:

pOH = –log [OH] and [OH] = 10–pOH

Importantly, based on the relationship between [H+] and [OH], we know:

pH + pOH = 14

Therefore, just by knowing one piece of information (pH, pOH, [H+] or [OH]), we can determine all other values.

Review an example of using pH and pOH to determine ion concentrations in Problem Example 2 in pH and Titration.

Finally, you must be able to determine if a solution is acidic, neutral, or basic from its pH value:

Acidic

pH < 7

Neutral

pH = 7

Basic

pH > 7

 

7e: Explain the process of titration.

Titrations are practical experiments to determine the concentration of an unknown sample of an acid or base.

    • Define titration.
    • Explain the steps of a titration.

Titrations are experiments we conduct to determine the concentration of an acid or base by reacting it with a known concentration of acid or base until it is neutralized. By determining the number of moles of acid or base required to neutralize the unknown sample, you can determine the number of moles and concentration in the unknown sample.

We will use an example to explain how a titration experiment works. Assume we want to determine the concentration of an unknown solution of HCl. Begin with a known volume of your HCl solution. Then, you will titrate the HCl solution with a solution of NaOH of known concentration. You slowly add NaOH solution to the HCl until the neutralization reaction has completed.

By knowing the volume of NaOH you added, you can determine the number of moles of NaOH added. Then, based on the mole ratio of the neutralization reaction, you can determine the number of moles and therefore the molarity, or concentration, of the HCl solution. The point at which the neutralization reaction has completed is called the equivalence point. Generally, the equivalence point is found using color-changing indicator solutions, which change color when the equivalence point is reached.

Review a worked example of a titration problem in Problem Example 3 in pH and Titration.

We can interpret a titration using a titration curve. The titration curve shows the pH of the solution as moles of the titrant, or solution of known concentration, are added. When the curve rises vertically, you have reached the equivalence point of the titration.

Image of a titration curve.

 

7f: Explain the relationship between conjugate acids and bases.

    • Define conjugate acid and bases.
    • For a given reaction, identify the conjugate acid–base pairs.

From the Bronsted–Lowry definition of acids and bases, we saw that acids and bases must exist together.

Review Section 1: Proton Donors and Acceptors, in Proton Donors and Acceptors: Acid–base Reactions à la Brønsted.

In any acid–base reaction, the acid donates a proton and the base accepts a proton. The products of the reaction are the “leftover” parts of the acid and base.

For example, in a generic acid dissociation reaction:

HA + B → A+ HB+

HA is the acid and B is the base. The product A is what the acid becomes, and the product HB+ is what the base becomes. We say that A is the conjugate base of the acid HA. HA and A are a conjugate acid base pair. Likewise, we say that HB+ is the conjugate acid of B. B and HB+ are a conjugate acid base pair.

We can summarize this as:

Image of conjugate base and conjugate acid.


7g: Compare and contrast processes of oxidation and reduction.

    • Define oxidation.
    • Define reduction.
    • For a given reaction, determine if it is oxidation or reduction.

Oxidation–reduction reactions, or redox reactions, are an important class of chemical reactions. Redox reactions involve two half–reactions: an oxidation reaction and a reduction reaction.

Let’s use an example to define oxidation and reduction from the text, Redox Reactions.

Consider this (unbalanced) reaction:

Cu(s) + 2Ag+(aq) → Cu2+(aq) + Ag(s)

We can think of this reaction in terms of electron transfer. Electrons from the solid copper must be removed to form the positive copper ion. We can think of this half reaction as:

Cu(s) → Cu2+(aq) + 2e where e is an electron.

Likewise, we can think of the half reaction involving the silver. Here, the silver ion gains two electrons from the copper to form silver solid:

2Ag+(aq) + 2e → Ag(s)

An oxidation reaction is a reaction in which electrons are lost. In this case, the copper is oxidized.

A reduction reaction is a reaction in which electrons are gained. In this case, silver is reduced.

We say copper is a reducing agent in this reaction because it causes the reduction reaction of silver. We say silver is an oxidizing agent in this reaction because it causes the oxidation reaction of copper.

Review how to write half reactions from a redox reaction in Example 11.14.1 in Redox Reactions.


7h: Write and balance equations for oxidation–reduction reactions.

In writing and balancing redox reactions, we need to use a different set of rules from general balancing rules. This is because we need to account for the electrons being transferred in the reaction.

    • Write and balance redox reactions in acidic solution.
    • Write and balance redox reactions in basic solution.

First, we will describe the rules for writing a redox reaction in acidic solution.

The first step is to write the oxidation numbers for each element in each compound in the reaction.

Review the rules for determining oxidation numbers in the section Determining Oxidation States, in Oxidation Numbers and Redox Reactions.

Secondly, write the unbalanced half reactions for oxidation and reduction. Note which elements are being oxidized or reduced.

Then, balance each half reaction. First, balance the elements being oxidized or reduced in each half reaction. Balance oxygen atoms by adding water as needed to either side of the half reactions. This is okay to do because water is the solvent in these reactions, so there is water present. Lastly, balance the hydrogen atoms by adding H+ ions as needed. This is okay to do because in an acidic solution, there are extra H+ ions in solution. Lastly, balance electronic charge by adding electrons as needed to the half reactions.

When you add the balanced half reactions up, cross out any terms that appear on both sides, and ensure that the number of atoms and the charges balance.

Now, we will describe the rules for writing a redox reaction in basic solution.

As above, the first step is to assign oxidation states to each atom in the compounds in the reaction.

The second step is to write the unbalanced half reactions for oxidation and reduction. Note which elements are being oxidized or reduced.

Then, balance the half reactions. First, balance the elements being oxidized or reduced in each half reaction. Balance oxygen by adding hydroxide, OH as needed. This is acceptable because a basic solution has excess OH in solution. Then, balance the hydrogen atoms by adding water as needed. Lastly, balance electronic charge by adding electrons as needed to the half reactions.

When you add the balanced half reactions up, cross out any terms that appear on both sides, and ensure that the number of atoms and the charges balance.

Review detailed, step-by-step examples of balancing redox reactions in acidic and basic solutions in the section Balancing Redox Reactions, in Balancing Redox Equations.


7i: Identify common oxidizing agents, common reducing agents, and substances that can act both as oxidizing and reducing agents.

    • Identify common oxidizing agents.
    • Identify common reducing agents.
    • Identify common substances that can act as either oxidizing or reducing agents.

Oxygen is the most common oxidizing agent. Oxygen readily oxidizes metals, creating metal oxides. Most metals exist as oxides in nature because of the abundance of oxygen in the air.

This picture shows a common example we are familiar with, rust, which is a hydrated iron oxide.

Image of a rusty pipe.

Metals, particularly those on the left of the periodic table, are strong reducing agents.

Water and hydrogen peroxide, H2O2, can act as both oxidizing agents and reducing agents.


Unit 7 Vocabulary
 

      • Arrhenius acid
      • Arrhenius base
      • Aqueous
      • Bronsted–Lowry acid
      • Bronsted–Lowry base
      • Conjugate acid
      • Conjugate base
      • Conjugate base pair
      • Equivalence point
      • Half reaction
      • Hydronium ion
      • Hydroxide ion
      • Indicator solution
      • Ion product of water
      • Net ionic equation
      • Neutralization
      • Oxidation reaction
      • Oxidizing agent
      • pH
      • pOH
      • Redox
      • Reducing agent
      • Reduction reaction
      • Salt
      • Self-ionization/self protolysis of water
      • Titrant
      • Titration
      • Titration curve
Last modified: Tuesday, June 11, 2019, 8:30 AM