Acids And Bases Chapter Assessment 17 Answers

Acids and bases are fundamental concepts in chemistry, underpinning a vast array of chemical reactions and biological processes. Understanding their properties, behavior, and interactions is crucial for anyone venturing into the realms of chemistry, biology, or related fields. Chapter 17, commonly found in introductory chemistry textbooks, typically delves into the intricacies of acids and bases, exploring their definitions, strengths, pH calculations, and applications. Mastering this chapter requires not only memorizing definitions but also applying these concepts to solve problems and analyze real-world scenarios.

Delving into Acid-Base Theories

The journey into understanding acids and bases begins with exploring different theoretical frameworks that define them. While several theories exist, three stand out as particularly significant: the Arrhenius theory, the Brønsted-Lowry theory, and the Lewis theory.

Arrhenius Theory: A Foundation

The Arrhenius theory, proposed by Svante Arrhenius, provides a foundational understanding of acids and bases. According to this theory:

• An Arrhenius acid is a substance that increases the concentration of hydrogen ions (H+) in aqueous solution.
• An Arrhenius base is a substance that increases the concentration of hydroxide ions (OH-) in aqueous solution.

While straightforward, the Arrhenius theory has limitations. It only applies to aqueous solutions and does not account for substances that exhibit acidic or basic properties without donating or accepting protons (H+).

Brønsted-Lowry Theory: Expanding the Scope

The Brønsted-Lowry theory, developed independently by Johannes Brønsted and Thomas Lowry, broadens the definition of acids and bases beyond aqueous solutions.

• A Brønsted-Lowry acid is a proton (H+) donor.
• A Brønsted-Lowry base is a proton (H+) acceptor.

This theory introduces the concept of conjugate acid-base pairs. When an acid donates a proton, the remaining species becomes its conjugate base. Conversely, when a base accepts a proton, it becomes its conjugate acid. For example, in the reaction:

HCl (acid) + H2O (base) ⇌ H3O+ (conjugate acid) + Cl- (conjugate base)

HCl donates a proton to H2O, forming H3O+ (hydronium ion) and Cl-. HCl and Cl- are a conjugate acid-base pair, as are H2O and H3O+.

Lewis Theory: The Most Inclusive Definition

The Lewis theory, proposed by Gilbert N. Lewis, provides the most comprehensive definition of acids and bases.

• A Lewis acid is an electron pair acceptor.
• A Lewis base is an electron pair donor.

This definition encompasses all Brønsted-Lowry acids and bases, but also includes substances that can accept or donate electron pairs without involving proton transfer. For example, BF3 (boron trifluoride) is a Lewis acid because it can accept an electron pair from a Lewis base like NH3 (ammonia).

Acid Strength, Base Strength, and pH

The strength of an acid or base refers to its ability to donate or accept protons, respectively. Strong acids and bases completely dissociate in solution, while weak acids and bases only partially dissociate.

Strong Acids and Bases

• Strong acids completely ionize in water, meaning that every molecule of the acid donates its proton to water, forming hydronium ions (H3O+). Common examples include hydrochloric acid (HCl), sulfuric acid (H2SO4), and nitric acid (HNO3).
• Strong bases completely dissociate in water, releasing hydroxide ions (OH-). Examples include sodium hydroxide (NaOH) and potassium hydroxide (KOH).

Because strong acids and bases dissociate completely, their concentrations directly determine the concentration of H3O+ or OH- ions in solution.

Weak Acids and Bases

• Weak acids only partially ionize in water, establishing an equilibrium between the undissociated acid, its conjugate base, and hydronium ions. Acetic acid (CH3COOH) and hydrofluoric acid (HF) are common examples. The extent of ionization is quantified by the acid dissociation constant, Ka. A smaller Ka value indicates a weaker acid.
• Weak bases only partially react with water to produce hydroxide ions (OH-), also establishing an equilibrium. Ammonia (NH3) and pyridine (C5H5N) are examples of weak bases. The extent of this reaction is quantified by the base dissociation constant, Kb. A smaller Kb value indicates a weaker base.

The pH Scale: Quantifying Acidity and Basicity

The pH scale is a logarithmic scale used to express the acidity or basicity of a solution. It ranges from 0 to 14, with:

• pH < 7 indicating an acidic solution.
• pH = 7 indicating a neutral solution.
• pH > 7 indicating a basic solution.

The pH is defined as the negative logarithm (base 10) of the hydronium ion concentration:

pH = -log[H3O+]

Similarly, the pOH is defined as the negative logarithm of the hydroxide ion concentration:

pOH = -log[OH-]

In aqueous solutions at 25°C, the following relationship holds:

pH + pOH = 14

This relationship allows for the calculation of pH from pOH, and vice versa.

Calculating pH for Strong Acids and Bases

For strong acids, the hydronium ion concentration ([H3O+]) is equal to the acid's concentration. Therefore, the pH can be calculated directly using the pH formula. For example, if a solution is 0.01 M HCl, then [H3O+] = 0.01 M, and pH = -log(0.01) = 2.

Similarly, for strong bases, the hydroxide ion concentration ([OH-]) is equal to the base's concentration. The pOH can be calculated, and then the pH can be found using the relationship pH + pOH = 14. For example, if a solution is 0.005 M NaOH, then [OH-] = 0.005 M, pOH = -log(0.005) ≈ 2.3, and pH = 14 - 2.3 = 11.7.

Calculating pH for Weak Acids and Bases

Calculating the pH of weak acid or base solutions is more complex because they only partially dissociate. It requires using the Ka or Kb value and setting up an ICE (Initial, Change, Equilibrium) table to determine the equilibrium concentrations of the ions.

Example: Calculating the pH of a Weak Acid Solution

Let's calculate the pH of a 0.1 M solution of acetic acid (CH3COOH), given that its Ka is 1.8 x 10-5.

1. Write the equilibrium reaction:

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

2. Set up the ICE table:

   	CH3COOH	H3O+	CH3COO-
   Initial (I)	0.1	0	0
   Change (C)	-x	+x	+x
   Equilibrium (E)	0.1-x	x	x

3. Write the Ka expression:

   Ka = [H3O+][CH3COO-] / [CH3COOH] = x2 / (0.1 - x)

4. Solve for x:

   Since Ka is small, we can assume that x is much smaller than 0.1, so 0.1 - x ≈ 0.1.

   1.8 x 10-5 = x2 / 0.1

   x2 = 1.8 x 10-6

   x = √(1.8 x 10-6) ≈ 1.34 x 10-3 M

   Therefore, [H3O+] ≈ 1.34 x 10-3 M

5. Calculate the pH:

   pH = -log[H3O+] = -log(1.34 x 10-3) ≈ 2.87

Therefore, the pH of a 0.1 M acetic acid solution is approximately 2.87.

Similar steps are involved in calculating the pH of a weak base solution, using the Kb value and determining the equilibrium concentration of OH- ions.

Acid-Base Reactions: Neutralization and Titration

Acid-base reactions involve the transfer of protons from an acid to a base. The most common type of acid-base reaction is neutralization, where an acid and a base react to form salt and water.

Neutralization Reactions

A neutralization reaction occurs when an acid and a base react in stoichiometrically equivalent amounts, resulting in a solution that is neither acidic nor basic (pH ≈ 7). For example:

HCl (aq) + NaOH (aq) → NaCl (aq) + H2O (l)

In this reaction, hydrochloric acid (HCl) reacts with sodium hydroxide (NaOH) to form sodium chloride (NaCl) and water (H2O). The H+ from HCl combines with the OH- from NaOH to form water, neutralizing the solution.

Titration: A Quantitative Analysis Technique

Titration is a laboratory technique used to determine the concentration of an acid or base by reacting it with a solution of known concentration (the titrant). The titrant is added to the analyte (the solution being analyzed) until the reaction is complete, which is indicated by a color change of an indicator or a sudden change in pH.

The point at which the reaction is complete is called the equivalence point. At the equivalence point, the number of moles of acid is equal to the number of moles of base (or vice versa). By knowing the volume and concentration of the titrant, the concentration of the analyte can be calculated.

Titration Curve:

A titration curve is a graph that plots the pH of the solution as a function of the volume of titrant added. The shape of the titration curve depends on the strength of the acid and base involved.

• Strong Acid-Strong Base Titration: The pH changes gradually at the beginning and end of the titration, with a sharp change in pH near the equivalence point (pH = 7).
• Weak Acid-Strong Base Titration: The initial pH is higher than that of a strong acid, and there is a buffering region before the equivalence point. The pH at the equivalence point is greater than 7.
• Strong Acid-Weak Base Titration: The initial pH is lower than that of a strong base, and the pH at the equivalence point is less than 7.
• Weak Acid-Weak Base Titration: The pH change near the equivalence point is less pronounced, making it difficult to determine the equivalence point accurately.

Indicators:

Indicators are weak acids or bases that change color depending on the pH of the solution. They are used to visually signal the endpoint of a titration, which is the point at which the indicator changes color. The ideal indicator should change color close to the equivalence point of the titration. Common indicators include phenolphthalein (colorless in acidic solutions, pink in basic solutions) and methyl orange (red in acidic solutions, yellow in basic solutions).

Buffers: Resisting pH Changes

Buffers are solutions that resist changes in pH when small amounts of acid or base are added. They are crucial in biological systems, where maintaining a stable pH is essential for proper function.

A buffer typically consists of a weak acid and its conjugate base, or a weak base and its conjugate acid. The weak acid neutralizes added base, and the conjugate base neutralizes added acid. This keeps the pH relatively constant.

How Buffers Work:

Consider a buffer solution containing acetic acid (CH3COOH) and its conjugate base, acetate (CH3COO-).

• If a strong acid is added: The acetate ions (CH3COO-) react with the added H+ ions to form acetic acid (CH3COOH), minimizing the change in pH.

  CH3COO- (aq) + H+ (aq) ⇌ CH3COOH (aq)

• If a strong base is added: The acetic acid (CH3COOH) reacts with the added OH- ions to form acetate ions (CH3COO-) and water, minimizing the change in pH.

  CH3COOH (aq) + OH- (aq) ⇌ CH3COO- (aq) + H2O (l)

The Henderson-Hasselbalch Equation:

The Henderson-Hasselbalch equation is used to calculate the pH of a buffer solution:

pH = pKa + log([A-] / [HA])

where:

• pKa is the negative logarithm of the acid dissociation constant (Ka) of the weak acid.
• [A-] is the concentration of the conjugate base.
• [HA] is the concentration of the weak acid.

This equation shows that the pH of a buffer solution depends on the pKa of the weak acid and the ratio of the concentrations of the conjugate base and the weak acid.

Buffer Capacity:

Buffer capacity refers to the amount of acid or base a buffer can neutralize before the pH begins to change significantly. A buffer has the highest capacity when the concentrations of the weak acid and its conjugate base are equal, meaning that pH = pKa. The buffer capacity is also affected by the overall concentration of the buffer components; higher concentrations lead to a greater buffer capacity.

Applications of Acids and Bases

Acids and bases play a vital role in a wide range of applications, spanning from industrial processes to biological functions.

• Industrial Applications: Sulfuric acid (H2SO4) is a crucial industrial chemical used in the production of fertilizers, detergents, and various other chemicals. Hydrochloric acid (HCl) is used in metal cleaning and processing. Sodium hydroxide (NaOH) is used in the manufacture of paper, textiles, and soaps.
• Biological Systems: Acids and bases are essential for maintaining the pH balance in biological systems. Enzymes, which catalyze biochemical reactions, are highly sensitive to pH changes. Buffers in blood and other bodily fluids help maintain a stable pH, ensuring proper physiological function.
• Environmental Chemistry: Acid rain, caused by the release of sulfur dioxide (SO2) and nitrogen oxides (NOx) from industrial processes and combustion, can acidify lakes and streams, harming aquatic life. Understanding acid-base chemistry is crucial for addressing environmental problems related to acidity.
• Pharmaceuticals: Many drugs are either acids or bases, and their effectiveness depends on their ionization state, which is influenced by pH. The pH of the gastrointestinal tract can affect the absorption and distribution of drugs.
• Food Chemistry: Acids and bases are used in food processing and preservation. Acetic acid (vinegar) is used as a preservative and flavoring agent. Baking soda (sodium bicarbonate, NaHCO3) is a base used as a leavening agent in baking.

Practice Problems and Chapter Assessment

Mastering the concepts presented in Chapter 17 requires practice. Working through a variety of problems will solidify your understanding and improve your problem-solving skills. Here are some types of problems you might encounter in a chapter assessment:

• Identifying Acids and Bases: Given a chemical reaction, identify the Brønsted-Lowry acids and bases, and their conjugate pairs.
• pH Calculations: Calculate the pH of strong and weak acid or base solutions, using Ka, Kb, and ICE tables.
• Titration Calculations: Determine the concentration of an unknown acid or base using titration data.
• Buffer Calculations: Calculate the pH of a buffer solution using the Henderson-Hasselbalch equation.
• Conceptual Questions: Explain the differences between Arrhenius, Brønsted-Lowry, and Lewis definitions of acids and bases. Describe how a buffer works. Discuss the applications of acids and bases in various fields.

By understanding the fundamental principles of acids and bases, and by practicing problem-solving, you can successfully navigate Chapter 17 and build a solid foundation for further studies in chemistry and related disciplines.