CHEM101 Study Guide

Unit 7: Acid-Base and Oxidation-Reduction Reactions

 

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

  • What are Arrhenius acids and bases?
  • Can you use the Arrhenius definition to identify acids and bases?
  • What are Brønsted-Lowry acids and bases?
  • Can you use the Brønsted-Lowry definition to identify acids and bases?
  • Can you write an equation for the self-ionization or autoprotolysis of water?
  • How is water both an acid and a base?
  • What are conjugate acids and bases?
  • For a given reaction, can you identify the conjugate acid-base pairs?

There is more than one definition of acids and bases which can be useful in different situations. 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 are hydrochloric acid, HCl, sulfuric acid, H2SO4, and acetic acid (vinegar), CH3COOH. Examples of Arrhenius bases are sodium hydroxide (also known as lye), NaOH, and ammonia, NH3.

The Brønsted-Lowry definition of acids and bases is more general than the Arrhenius definition. A Brønsted-Lowry acid is a proton (hydrogen ion) donator. A Brønsted-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-.

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. 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)

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

From the Brønsted-Lowry definition of acids and bases, we saw that acids and bases must exist together. 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:

Review this material in Acids and Bases: An Introduction, What are Acids and Bases?, Arrhenius Definition of Acids and Bases, Brønsted-Lowry Definition of Acids and Bases, Proton Donors and Acceptors: Acid-base Reactions à la Brønsted, and Conjugate Acids and Bases.

 

7b. Write and balance equations for neutralization reactions

  • Can you write and balance a neutralization reaction for a given acid and base?
  • What is a salt?

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. 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 those that are dissolved in water.

So, we first write a total ionic equation in which we write 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 Acids and Bases: An Introduction. Also, review Acid-Base Neutralization Reaction.

 

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

  • Can you write the ion product of water?
  • What are acidic, basic, and neutral solutions based on the [H3O+] and [OH-]?
  • Can you determine the concentration of [H3O+] or [OH-] for a given solution?
  • Can you calculate pH for a given solution?
  • Can you use pH to classify a solution as acidic, neutral, or basic?

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

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 K_{w} expression. Using K_{w}, 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:

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.

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

Review this material in Aqueous Solution, pH, and Titration and Definition of pH.

 

7d. Contrast strong acids and bases with weak acids and bases

  • Can you identify the strong acid or base in a reaction?
  • Can you write the neutralization reaction between a strong acid and a strong base?
  • How would you organize acids based on their strengths?

Acids and bases are either strong or weak. Strong acids are good at donating protons because of the stability of the anions they form. In each strong acid, the proton is bonded to a very strong electronegative atom (O, Cl, Br, I). When the proton is donated, the leftover bonding electrons remain on the electronegative atom.

There are 7 strong acids: HCl, HBr, HI, HNO3, HClO3, HClO4, and H2SO4.

There are 8 strong bases: LiOH, NaOH, KOH, RbOH, CsOH, Ca(OH)2, Sr(OH)2, and Ba(OH)2.

Any other acid or base is considered weak.

Strong acids completely ionize in solution, whereas weak acids only partially ionize. Notice the difference between the reaction arrows used here:

Strong bases will also completely ionize in solution. Weak acids and bases have much smaller ionization constants (Ka) than their stronger counterparts.

Review this material in Strong Acids and Bases and Ka and Acid Strength.

 

7e. Perform calculations for a titration of a strong acid with a strong base

  • What is titration?
  • What are the steps of a titration?

Titrations are practical 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 the 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.

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.

Review this material in Titration Introduction.

 

7f. Compare and contrast processes of oxidation and reduction

  • What is oxidation?
  • What is reduction?
  • For a given reaction, can you 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. 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 Redox Reactions.

 

7g. Calculate oxidation numbers for each element in a given compound

  • How do you assign oxidation numbers for a given element?
  • How do you determine the oxidation states of all atoms in a compound?

Oxidation numbers are used to assign the most common electronic charge for a given element. Here are the general rules:

  1. The sum of oxidation states in a molecule must equal the molecule's charge. The sum is zero for a neutral molecule.

  2. The atom that is more electronegative in a bond gets a negative oxidation state. The atom that is more electropositive gets a positive oxidation state.

  3. Certain elements always have the same oxidation state.

Determining the oxidation number of elements in a formula gives us important information about the bonding and reactivity of the compound. We use these rules to assign oxidation numbers for each atom in the compound. It is important to apply these rules systematically and be sure to note any exceptions to the rules that may apply.

For example, let's explore the nitrate ion NO3.

  • We know this compound is an ion with a -1 charge, so the overall oxidation numbers must add up to -1.

  • We know oxygen is more electronegative than nitrogen, so we know oxygen will have a negative oxidation number and nitrogen will have a positive one.

  • We know oxygen almost always has a -2 oxidation state (with a few exceptions).

  • Since this is not one of the exceptions, the oxidation state for each oxygen is -2.

  • Now we can determine the oxidation state of nitrogen.

  • There are three oxygen atoms, each with an oxidation state of -2.

  • So, the total oxidation state from the three oxygen atoms is -6.

  • To get the overall compound oxidation state to be -1, the oxidation state of nitrogen must be +5.

Review this material in Oxidation Numbers and Redox Reactions.

 

7h. Write and balance equations for oxidation-reduction reactions

  • Can you write and balance redox reactions in an acidic solution?
  • Can you write and balance redox reactions in a basic solution?

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.

First, we will describe the rules for writing a redox reaction in an acidic solution. The first step is to write the oxidation numbers for each element in each compound in the reaction.

Second, 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 the 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 solutions. 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 the rules for determining oxidation numbers in Oxidation Numbers and Redox Reactions. Review detailed, step-by-step examples of balancing redox reactions in acidic and basic solutions in Balancing Redox Equations and Write a Balanced Redox Reaction.

 

Unit 7 Vocabulary

  • Arrhenius acid
  • Arrhenius base
  • Aqueous
  • Aqueous substance
  • Brønsted-Lowry acid
  • Brønsted-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 of water
  • Self-protolysis of water
  • Titrant
  • Titration
  • Titration curve