M3Q3-4: Acids, Bases, Neutralization, and Gas-Forming Reactions

Introduction

This section continues to explores aqueous chemical reactions. First, we examine acids and bases, both strong and weak. Second, we examine neutralization reactions. Third, we examine gas-forming reactions. The section below provides a more detailed description of these topics, worked examples, practice problems and a glossary of important terms.

Learning Objectives for Acid, Base, and Neutralization Reactions

| Key Concepts and Summary | Glossary | End of Section Exercises |

Acid-Base Reactions

An acid-base reaction is one in which a hydrogen ion, H+, is transferred from one chemical species to another. Such reactions are of central importance to numerous natural and technological processes, ranging from the chemical transformations that take place within cells and the lakes and oceans, to the industrial-scale production of fertilizers, pharmaceuticals, and other substances essential to society. The subject of acid-base chemistry, therefore, is worthy of discussion.

For purposes of this brief introduction, we will consider only the more common types of acid-base reactions that take place in aqueous solutions. In this context, an acid is a substance that will dissolve in water to yield hydronium ions, H3O+. As an example, consider the equation shown here:

HCl(aq) + H2O(ℓ)  ⟶  Cl–(aq) + H3O+(aq)

The process represented by this equation confirms that hydrogen chloride (also commonly known as hydrochloric acid) is an acid. When dissolved in water, H3O+ ions are produced by a chemical reaction in which H+ ions are transferred from HCl molecules to H2O molecules (Figure 1b).

This figure shows two flasks, labeled a and b. The flasks are both sealed with stoppers and are nearly three-quarters full of a liquid. Flask a is labeled H C l followed by g in parentheses. In the liquid there are approximately twenty space-filling molecular models composed of one red sphere and two smaller attached white spheres. The label H subscript 2 O followed by a q in parentheses is connected with a line to one of these models. In the space above the liquid in the flask, four space filling molecular models composed of one larger green sphere to which a smaller white sphere is bonded are shown. To one of these models, the label H C l followed by g in parentheses is attached with a line segment. An arrow is drawn from the space above the liquid pointing down into the liquid below. Flask b is labeled H subscript 3 O superscript positive sign followed by a q in parentheses. This is followed by a plus sign and C l superscript negative sign which is also followed by a q in parentheses. In this flask, no molecules are shown in the open space above the liquid. A label, C l superscript negative sign followed by a q in parentheses, is connected with a line segment to a green sphere. This sphere is surrounded by four molecules composed each of one red sphere and two white smaller spheres. A few of these same molecules appear separate from the green spheres in the liquid. A line segment connects one of them to the label H subscript 2 O which is followed by l in parentheses. There are a few molecules formed from one central larger red sphere to which three smaller white spheres are bonded. A line segment is drawn from one of these to the label H subscript 3 O superscript positive sign, followed by a q in parentheses.
Figure 1. When hydrogen chloride gas dissolves in water, (a) it reacts as an acid, transferring protons to water molecules to yield (b) hydronium ions (and solvated chloride ions).

The nature of HCl is such that its reaction with water as just described is essentially 100% efficient: Virtually every HCl molecule that dissolves in water will undergo this reaction. Acids that completely react in this fashion are called strong acids, and HCl is one among just a handful of common acid compounds that are classified as strong (Table 1). Memorizing these strong acids is highly recommended! Even though a bottle of hydrochloric acid is labeled as HCl(aq), there are essentially no molecules of HCl present in solution due to the complete dissociation of the molecule to produce ions. HCl is therefore classified as a strong electrolyte (see the Dissolving and Electrolytes section of a previous section) and HCl(aq) is a conductive solution.

Table 1. Common Strong Acids To Know 
Compound Formula Name in Aqueous Solution Compound Formula Name in Aqueous Solution
HBr hydrobromic acid HNO3 nitric acid
HCl hydrochloric acid HClO4 perchloric acid
HI hydroiodic acid H2SO4 sulfuric acid

A far greater number of compounds behave as weak acids and only partially react with water, leaving a large majority of dissolved molecules in their original form and generating a relatively small amount of hydronium ions. Weak acids are commonly encountered in nature, being the substances partly responsible for the tangy taste of citrus fruits, the stinging sensation of insect bites, and the unpleasant smells associated with body odor. A familiar example of a weak acid is acetic acid, the main ingredient in food vinegars:

CH3COOH(aq) + H2O(ℓ)  ⇌  CH3COO–(aq) + H3O+(aq)

When dissolved in water under typical conditions, only about 1% of acetic acid molecules are present in the ionized form, CH3COO–(aq) (Figure 2). (The use of a double-arrow in the equation above denotes the partial reaction aspect of this process.) Weak acids, as their name suggests, are classified as weak electrolytes due to the low concentration of ions in solution. To see how these conduct electricity, refer back to the conductivity movies presented in a previous section.

This figure contains two images, each with an associated structural formula provided in the lower left corner of the image. The first image is a photograph of a variety of thinly sliced, circular cross sections of citrus fruits ranging in color for green to yellow, to orange and reddish-orange. The slices are closely packed on a white background. The structural formula with this picture shows a central chain of five C atoms. The leftmost C atom has an O atom double bonded above and to the left and a singly bonded O atom below and to the left. This single bonded O atom has an H atom indicated in red on its left side which is highlighted in pink. The second C atom moving to the right has H atoms bonded above and below. The third C atom has a single bonded O atom above which has an H atom on its right. This third C atom has a C atom bonded below it which has an O atom double bonded below and to the left and a singly bonded O atom below and to the right. An H atom appears in red and is highlighted in pink to the right of the singly bonded O atom. The fourth C atom has H atoms bonded above and below. The fifth C atom is at the right end of the structure. It has an O atom double bonded above and to the right and a singly bonded O atom below and to the right. This single bonded O atom has a red H atom on its right side which is highlighted in pink. The second image is a photograph of bottles of vinegar. The bottles are labeled, “Balsamic Vinegar,” and appear to be clear and colorless. The liquid in this bottle appears to be brown. The structural formula that appears with this image shows a chain of two C atoms. The leftmost C atom has H atoms bonded above, below, and to the left. The C atom on the right has a doubly bonded O atom above and to the right and a singly bonded O atom below and to the right. This O atom has an H atom bonded to its right which is highlighted in pink.
Figure 2. (a) Fruits such as oranges, lemons, and grapefruit contain the weak acid citric acid. (b) Vinegars contain the weak acid acetic acid. (credit a: modification of work by Scott Bauer; credit b: modification of work by Brücke-Osteuropa/Wikimedia Commons)

Not only is sulfuric acid, H2SO4, a strong acid, it is also a diprotic acid as it contains two protons, which is a common way to refer to H+ ions since H+ contains 1 proton and 0 electrons.. The dissociation of diprotic acids in water is best described using two separate equations with the first equation describing the transfer of one proton to water, and the second equation describing the transfer of the second proton from the HSO4–(aq) produced in the first equation:

H2SO4(aq) + H2O(ℓ)  ⟶  HSO4–(aq) + H3O+(aq)  strong acid, fully dissociated

HSO4–(aq) + H2O(ℓ)  ⇌  SO42-(aq) + H3O+(aq)  weak acid, partially dissociated

Example 1

Writing Dissociation equations for Acids in Water

Write the chemical equation that describes the dissociation of nitric acid, HNO3, in water.

Solution

HNO3 is a strong acid (Table 1) and therefore completely dissociates in water:

HNO3(aq) + H2O(ℓ)  ⟶  H3O+(aq) + NO3–(aq)

Check Your Learning

Write the chemical equation that describes the dissociation of perchloric acid in water.

Answer:

HClO4(aq) + H2O(ℓ)  ⟶  H3O+(aq) + ClO4–(aq)

Example 2

Writing Dissociation equations for Acids in Water

Write the chemical equation that describes the dissociation of nitrous acid, HNO2, in water.

Solution

HNO2 is not in Table 1 and is therefore a weak acid and will only partially dissociate in water, hence the double arrow in the equation below.

HNO2(aq) + H2O(ℓ)  ⇌  H3O+(aq) + NO2–(aq)

Check Your Learning

Write the chemical equation that describes the dissociation of hypochlorous acid, HClO(aq), acid in water.

Answer:

HClO(aq) + H2O(ℓ)  ⇌  H3O+(aq) + ClO–(aq)

A base is a substance that will dissolve in water to yield hydroxide ions, OH−. The most common bases are ionic compounds composed of alkali or alkaline earth metal cations (groups 1 and 2) combined with the hydroxide ion—for example, NaOH and Ca(OH)2. When these compounds dissolve in water, hydroxide ions are released directly into the solution. For example, KOH and Ba(OH)2 dissolve in water and dissociate completely to produce cations (K+ and Ba2+, respectively) and hydroxide ions, OH−. These bases, along with other hydroxides that completely dissociate in water, are considered strong bases.

Table 2. Common Strong Bases 
Compound Formula Name in Aqueous Solution Compound Formula Name in Aqueous Solution
LiOH lithium hydroxide Ca(OH)2 calcium hydroxide
NaOH sodium hydroxide Sr(OH)2 strontium hydroxide
KOH potassium hydroxide Ba(OH)2 barium hydroxide

Consider as an example the dissolution of lye (sodium hydroxide) in water:

NaOH(s)  ⟶  Na+(aq) + OH–(aq)

This equation confirms that sodium hydroxide is a base. When dissolved in water, NaOH dissociates to yield Na+ and OH− ions. This is also true for any other ionic compound containing hydroxide ions. Since the dissociation process is essentially complete when ionic compounds dissolve in water under typical conditions, NaOH and other ionic hydroxides are all classified as strong bases.

Unlike ionic hydroxides, some compounds produce hydroxide ions when dissolved by chemically reacting with water molecules. In all cases, these compounds react only partially and are therefore classified as weak bases. These types of compounds are also abundant in nature and important commodities in various technologies. For example, global production of the weak base ammonia is typically well over 100 metric tons annually, being widely used as an agricultural fertilizer, a raw material for chemical synthesis of other compounds, and an active ingredient in household cleaners (Figure 3). When dissolved in water, ammonia reacts partially to yield hydroxide ions, as shown here:

NH3(aq) + H2O(ℓ)  ⇌  NH4+(aq) + OH–(aq)

This is, by definition, an acid-base reaction, in this case involving the transfer of H+ ions from water molecules to ammonia molecules. Under typical conditions, only about 1% of the dissolved ammonia is present as NH4+ ions, and NH3(aq) is a weak electrolyte.

This photograph shows a large agricultural tractor in a field pulling a field sprayer and a large, white cylindrical tank which is labeled “Caution Ammonia.”
Figure 3. Ammonia is a weak base used in a variety of applications. (a) Pure ammonia is commonly applied as an agricultural fertilizer. (b) Dilute solutions of ammonia are effective household cleansers. (credit a: modification of work by National Resources Conservation Service; credit b: modification of work by pat00139)

Neutralization Reactions

Reactions between strong acids and bases

The chemical reactions described in which acids and bases dissolved in water produce hydronium and hydroxide ions, respectively, are, by definition, acid-base reactions. In these reactions, water serves as both a solvent and a reactant. A neutralization reaction is a specific type of acid-base reaction in which the reactants are an acid and a base, the products are often a salt and water, and neither reactant is the water itself:

acid + base  ⟶  salt + water

To illustrate a neutralization reaction, which are another category of double displacement reactions, consider what happens when a typical antacid such as milk of magnesia (an aqueous suspension of solid Mg(OH)2) is ingested to ease symptoms associated with excess stomach acid (HCl):

Mg(OH)2(s) + 2 HCl(aq)  ⟶  MgCl2(aq) + 2 H2O(ℓ)

Note that in addition to water, this reaction produces a salt, magnesium chloride. When writing these equations it is essential that you recognize the ions involved in the chemical reaction and their appropriate charges to ensure the correct formula for the products is written before any attempt is made to balance the equation.

Here, the magnesium hydroxide is solid, so it must be written as Mg(OH)2(s). The complete ionic equation for this reaction is:

Mg(OH)2(s) + 2 H+(aq) + 2 Cl–(aq)  ⟶  Mg2+(aq) + 2 Cl–(aq) + 2 H2O(ℓ)

Canceling out the spectator Cl–(aq) ions gives the net ionic equation:

Mg(OH)2(s) + 2 H+(aq)  ⟶  Mg2+(aq) + 2 H2O(ℓ)

Demonstration: Conductivity titration of a strong base with a strong acid

Set up. This demonstration shows the titration of a strong base, Ba(OH)2, with a strong acid (H2SO4). The buret contains 0.1 M H2SO4 and the dish contains 0.1 M Ba(OH)2. Phenolphthalein indicator is also added to the dish. This indicator, as we saw in a previous demo, is colorless when the solution is acidic and pink when the solution is basic. Before watching the video, write the reaction equation and predict whether the solution will transition from colorless to pink, or from pink to colorless.

In this demonstration, we are also observing the conductivity using a similar set-up to what we used in the electrolyte demonstration earlier in this module. Remember that electricity will flow (and light the bulb) when there are electrolytes in solution.

Prediction. Before watching the video, write the reaction equation for this experiment and form a hypothesis about if and when the bulb will be lit.

Explanation. The reaction occurring in this demonstration is:

H2SO4(aq) + Ba(OH)2(aq)  ⟶  BaSO4(s) + 2 H2O(ℓ)

This video shows that the strong base, Ba(OH)2, is a strong electrolyte and conducts electricity in order to light the bulb. This is in agreement with what we saw in our earlier electrolyte demonstration. As H2SO4 is added, the solution fades from pink to colorless and forms solid barium sulfate. The light bulb begins to dim as the equivalence point is approached. The equivalence point is the point at which all of the base in the dish is completely neutralized by the addition of the strong acid. At this point, all ions are neutralized and are present as a solid salt and therefore there are no more mobile ions in solution to carry the charge. The bulb is off. When more H2SO4 is added after the equivalence point, there are excess H+ and SO42- ions in solution to carry the charge, and the bulb lights up again.

Reactions involving weak acids and bases

Neutralization reactions between weak acids and strong bases also produce a salt and water as products, but care must be taken when writing the complete and net ionic equations as weak acids are only slightly dissociated in solution.

To illustrate, consider the reaction between weak acid HNO2(aq) and strong base NaOH(aq) described below:

HNO2(aq) + NaOH(aq)  ⟶  NaNO2(aq) + H2O(ℓ)  overall equation

Because HNO2(aq) is only partially dissociated and it exists mainly in its molecular form in solution, it is not logical to write it as separate ions in the complete ionic equation. It is therefore unchanged in the ionic equations. However, NaOH(aq) is a strong base and should be represented as separate ions, as should NaNO2, which is a soluble salt.

HNO2(aq) + Na+(aq) + OH–(aq)  ⟶  Na+(aq) + NO2–(aq) + H2O(ℓ)  complete ionic equation

Canceling out the Na+(aq) spectator ions produces the net ionic equation:

HNO2(aq) + OH–(aq)  ⟶  NO2–(aq) + H2O(ℓ)  net ionic equation

Example 3

Writing Equations for Acid-Base Reactions
Write balanced chemical equations (overall, complete ionic and net ionic) for the reaction when a solution of barium hydroxide is neutralized with a solution of nitric acid.

Solution

The two reactants are Ba(OH)2 and HNO3. Since this is a neutralization reaction between a strong acid and strong base, the two products will be water and a salt composed of the cation of the ionic hydroxide (Ba2+) and the anion generated when the acid transfers its hydrogen ion (NO3−).

Ba(OH)2(aq) + 2 HNO3(aq)  ⟶  Ba(NO3)2(aq) + 2 H2O(ℓ)

The complete ionic equation is:

Ba2+(aq) + 2 OH–(aq) + 2 H+(aq) + 2 NO3–(aq)  ⟶  Ba2+(aq) + 2 NO3–(aq) + 2 H2O(ℓ)

The spectator ions are Ba2+ and NO3– leading to the net ionic equation:

2 OH–(aq) + 2 H+(aq)  ⟶  2 H2O(ℓ)

Which in its smallest whole number ratio reduces to the following:

OH–(aq) + H+(aq)  ⟶  H2O(ℓ)

This is the net ionic equation for any strong acid strong base reaction when all reactants and products are soluble.

Check Your Learning

Write the overall, complete ionic and net ionic equation representing the neutralization of solid calcium hydroxide and perchloric acid.

Answer:

Ca(OH)2(s) + 2 HClO4(aq)  ⟶  Ca(ClO4)2(aq) + 2 H2O(ℓ)  overall equation

Ca(OH)2(s) + 2 H+(aq) + 2 ClO4–(aq)  ⟶  Ca2+(aq) + 2 ClO4–(aq) + 2 H2O(ℓ)  complete ionic equation

Ca(OH)2(s) + 2 H+(aq)  ⟶  Ca2+(aq) + 2 H2O(ℓ)  net ionic equation

Check Your Learning
Write the overall, complete ionic, and net ionic equation that describes the reaction between the weak acid CH3COOH(aq) and KOH(aq).

Answer:

CH3COOH(aq) + KOH(aq)  ⟶  CH3COOK(aq) + H2O(ℓ)  overall equation

CH3COOH(aq) + K+(aq) + OH–(aq)  ⟶  CH3COO–(aq) + K+(aq) + H2O(ℓ)  complete ionic equation

CH3COOH(aq) + OH–(aq)  ⟶  CH3COO–(aq) + H2O(ℓ)  net ionic equation

 

Explore the microscopic view of strong and weak acids and bases.

Acid Base Gas-Forming Reactions

Ionic compounds containing carbonate, sulfite, and sulfide anions are bases that react with acids to produce a salt, water, and a gas. Carbonates produce gaseous carbon dioxide, metal sulfites produce gaseous sulfur dioxide, and metal sulfides produce gaseous hydrogen sulfide as outlined in the ionic reaction schemes below:

CO32-(aq) + 2 H+(aq)  →  H2CO3(aq, Unstable)  →  CO2(g) + H2O(ℓ)

SO32-(aq) + 2 H+(aq)  →  H2SO3(aq, Unstable)  →  SO2(g) + H2O(ℓ)

S2-(aq) + 2 H+(aq)  →  H2S(g)

The H2CO3 (carbonic acid) and H2SO3 (sulfurous acid) formed when carbonates and sulfites react with acids are generally unstable when formed in solution in an open container and immediately decompose to form the respective gas and water. For example, when hydrochloric acid is added to solid calcium carbonate, bubbles of carbon dioxide are immediately observed. This chemical reaction is described by the following overall, complete ionic, and net ionic equations:

CaCO3(s) + 2 HCl(aq)  →  CaCl2(aq) + CO2(g) + H2O(ℓ) overall equation

CaCO3(s) + 2 H+(aq) + 2 Cl–(aq)  →  Ca2+(aq) + 2 Cl–(aq) + CO2(g) + H2O(ℓ) complete ionic equation

CaCO3(s) + 2 H+(aq)  →  Ca2+(aq) + CO2(g) + H2O(ℓ)  net ionic equation

The reaction of calcium sulfite with hydrochloric acid is very similar to the above reaction, the overall equation given below:

CaSO3(s) + 2 HCl(aq)  →  CaCl2(aq) + SO2(g) + H2O(ℓ) overall equation

The overall equation describing the reaction between sodium sulfide and hydrochloric acid is written below:

Na2S(aq) + 2 HCl(aq)  →  2 NaCl(aq) + H2S(g)

Key Concepts and Summary

Chemical reactions are classified according to similar patterns of behavior. This section discussed basic properties of acids, bases, neutralization reactions, and gas-forming reactions. Acids produce H+ ions in solution, while bases produce OH– ions in solution. Neutralization reactions occur between acids and bases and produce a salt and water. The term “salt” is used to refer to any ionic compound. When writing complete and net ionic equations for neutralization reactions between a strong base and weak acid or between a weak base and strong acid, it is important to remember that weak acids and bases do not dissociate to a large extent in solution. Weak acids and bases remain intact in a complete and net ionic equation.

Ionic compounds containing carbonate (CO32-), sulfite (SO32-), and sulfide (S2-) are bases that form salt, water and a gas when reacted with acids.

Glossary

acid
substance that produces H3O+ when dissolved in water
acid-base reaction
reaction involving the transfer of a hydrogen ion between reactant species
base
substance that produces OH− when dissolved in water
neutralization reaction
reaction between an acid and a base to produce salt and water
salt
ionic compound that can be formed by the reaction of an acid with a base that contains a cation and an anion other than hydroxide or oxide
strong acid
acid that reacts completely when dissolved in water to yield hydronium ions
strong base
base that reacts completely when dissolved in water to yield hydroxide ions
weak acid
acid that reacts only to a slight extent when dissolved in water to yield hydronium ions
weak base
base that reacts only to a slight extent when dissolved in water to yield hydroxide ions

Chemistry End of Section Exercises

  1. The reaction 2 HCl(aq) + Pb(NO3)2(aq)  →  PbCl2(s) + 2 HNO3(aq) is an example of a(n) _______ reaction. Select all that apply.
    1. Precipitation
    2. Acid-base
    3. Neutralization Reaction
    4. Gas-forming
  2. Fill in the table by writing what you would observe when the two solutions are mixed together. Options are: “No observed change”, “precipitate is formed”, “gas is formed”.
    Na2CO3(aq) NH3(aq) H2SO4(aq) NiCl2(aq)
    Na2CO3(aq)
    NH3(aq)
    H2SO4(aq)
    NiCl2(aq)
    1. Write the chemical formula(s) for all of the precipitate(s) that formed.
    2. Write the balanced overall equation for one reaction that produced a gas. Include states (g, ℓ, s, aq) in your equation.
  3. Which of the following reactions will produce at least one gaseous product? Select all that apply.
    1. The reaction between sodium sulfate and nitric acid.
    2. The reaction between nitric acid and potassium hydroxide.
    3. The reaction between calcium sulfide and hydrochloric acid.
  4. You have a salt solution that contains either potassium sulfate or lithium carbonate, and you need to determine which of the two salts is present.
    1. You choose ______ to determine which of the two salts is present.
      1. barium acetate
      2. ammonium acetate
      3. hydrochloric acid
      4. none of the above
    2. The reason for your choice above is
      1. It forms a precipitate with one unknown and has no visible reaction with the other.
      2. It forms a gas with one unknown and has no visible reaction with the other.
      3. It forms a gas with one unknown and a precipitate with the other.
      4. Reagents (a)-(c) will react the same with both potassium sulfate and lithium carbonate.
  5. Complete and balance the following acid-base equations:
    1. HCl gas reacts with solid Ca(OH)2.
    2. A solution of Sr(OH)2 is added to a solution of HNO3.
  6. Complete and balance the following acid-base equations:
    1. A solution of HClO4 is added to a solution of LiOH.
    2. Aqueous H2SO4 reacts with NaOH.
    3. Ba(OH)2 reacts with HF gas.
  7. Complete and balance the equations for the following acid-base neutralization reaction. If water is used as a solvent, write the reactants and products as aqueous ions. In some cases, there may be more than one correct answer, depending on the amounts of reactants used.
    Mg(OH)2(s) + HClO4(aq)  →
  8. Sr(ClO4)2 is a salt that can be formed in a neutralization reaction between an acid and a base. Write the balanced chemical equation for this reaction. Include states (g, ℓ, s, aq) in your equation.
  9. 1.0 mol carbonic acid (H2CO3) and 1.0 mol nitric acid (HNO3) are dissolved in equal volumes of water, and each solution is subjected to a conductivity test.  Explain at the particle-level why one solution generates a dim lighting-up of the bulb (and is a weak conductor) while the other generates a bright lighting-up of the bulb (and is a strong conductor).

Answers to Chemistry End of Section Exercises

  1. A
  2. Na2CO3(aq) NH3(aq) H2SO4(aq) NiCl2(aq)
    Na2CO3(aq)
    NH3(aq) No observed change
    H2SO4(aq) gas is formed No observed change
    NiCl2(aq) precipitate is formed precipitate is formed No observed change

    (a) NiCO3, Ni(OH)2
    (b) Na2CO3(aq) + H2SO4(aq) → Na2SO4(aq) + CO2(g) + H2O(ℓ)

  3. C
  4. (a) C;  (b) F
    Left-click here to watch Exercise 4 problem solving video.

  5. (a) 2 HCl(g) + Ca(OH)2(s)  →  CaCl2(s) + 2 H2O(ℓ)
    (b) Sr(OH)2(aq) + 2 HNO3(aq)  →  Sr(NO3)2(aq) + 2 H2O(ℓ)
  6. (a) HClO4(aq) + LiOH(aq)  →  LiClO4(aq) + H2O(ℓ)
    (b) H2SO4(aq) + 2 NaOH(aq)  →  Na2SO4(aq) + 2 H2O(ℓ)
    (c) Ba(OH)2(aq) + 2 HF(g)  →  BaF2(s) + 2 H2O(ℓ)
  7. Mg(OH)2(s) + 2 HClO4(aq)  →  Mg2+(aq) + 2 ClO4–(aq) + 2 H2O(ℓ)
  8. Sr(OH)2(aq) + 2 HClO4(aq)  →  Sr(ClO4)2(aq) + 2 H2O(ℓ)
  9. Carbonic acid is a weak acid that will only partially dissociate to H3O+ and HCO3–, while nitric acid is a strong acid that will fully dissociate to H3O+ and NO3–. More dissociation means that more charged particles are in solution.  The circuit will be complete in both cases, but the light bulb will be brighter if there are more charged particles in solution.
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