Chapter 8: Acids and Bases Study Notes

Chapter 8: Acids and Bases Notes

8.1 Introduction to Acids and Bases

• The earliest definition of acids and bases was provided by Arrhenius.

  • Acid Definition: An acid contains a hydrogen atom and dissolves in water to form a hydrogen ion.

  • Example: $ext{HCl(g)} <br>ightarrow ext{H}^+(aq) + ext{Cl}^-(aq)$

  • Base Definition: A base contains hydroxide and dissolves in water to form hydroxide ions.

  • Example: $ext{NaOH(s)} <br>ightarrow ext{Na}^+(aq) + ext{OH}^-(aq)$

8.1A Brønsted–Lowry Definition

• The Arrhenius definition accurately predicts the behavior of many acids and bases but has limitations.

• Limitations:

  • For instance, hydrogen ions do not exist independently in water. Instead, they react with water to form hydronium ions ($ext{H}_3 ext{O}^+$).

• Brønsted–Lowry Acid-Base Definitions:

  • A Brønsted–Lowry acid is defined as a proton donor.

  • A Brønsted–Lowry base is defined as a proton acceptor.

8.1B Brønsted–Lowry Acids

• A Brønsted–Lowry acid must contain a hydrogen atom.

• Example of a Brønsted–Lowry Acid Reaction:

  • $ext{HCl(g)} + ext{H}<em>2 ext{O}(l) ightarrow ext{H}</em>3 ext{O}^+(aq) + ext{Cl}^-(aq)$

  • In this reaction, HCl donates a proton to water, thereby forming the hydronium ion (H₃O⁺).

• General Representation:

  • Brønsted–Lowry acids are often symbolized as $ext{HA}$ where $A$ can be elements like Cl or Br in the acids $ext{HCl}$ and $ext{HBr}$.

  • Other examples include sulfuric acid ($ext{H}<em>2 ext{SO}</em>4$) and nitric acid ($ext{HNO}_3$).

  • Carboxylic acids have the functional group $ext{C(=O)OH}$, where the hydrogen from the hydroxyl group is the acidic proton that gets donated.

8.1C Brønsted–Lowry Bases

• A Brønsted–Lowry base is defined as a proton acceptor, meaning it forms a bond to a proton.

• Characteristics of Brønsted–Lowry Bases:

  • A base must possess a lone pair of electrons that may be used to form a new bond with the proton.

• General Representation:

  • Brønsted–Lowry bases can be denoted as $ext{B:}$ to indicate the presence of lone pairs.

8.2 The Reaction of a Brønsted–Lowry Acid with a Brønsted–Lowry Base

8.2.1 Reaction Overview

• In an acid-base reaction:

  • One bond is broken while another is formed.

  • The product formed from the acid losing a proton is called its conjugate base.

  • The product formed from the base gaining a proton is referred to as its conjugate acid.

8.2.2 Conjugate Acid-Base Pairs

• Two entities that differ by the presence of a proton are termed a conjugate acid-base pair.

  • Example: $ext{HBr}$ (strong acid) and $ext{Br}^-$ (conjugate base).

  • Example: $ext{H}_2 ext{O}$ (acting as an acid) and $ext{OH}^-$ (conjugate base).

• Important Concept:

  • The total net charge must remain constant on both sides of the reaction equation.

8.2.3 Charge Changes During Reactions

• When a species gains a proton, it results in a gain of +1 charge.

• When a species loses a proton, it results in an effective gain of -1 charge.

  • Example: ( \text{Base} + ext{H}^+ \rightarrow \text{Acid} )

  • Example: ( \text{Acid} \rightarrow \text{Conjugate Base} + ext{H}^+ )

8.2.4 Amphoteric Compounds

• An amphoteric compound is one that can act as both an acid and a base.

• Water ($ext{H}_2 ext{O}$) is a prime example, as it can accept a proton to become a hydronium ion and can donate a proton to form hydroxide.

8.3 Acid and Base Strength

8.3.1 Dissociation Explained

• The process of a covalent molecule or ionic compound separating into its ions in water is called dissociation.

• Strong Acids:

  • Strong acids fully dissociate in water (100%).

  • A single arrow representation is used, indicating product predominance in equilibrium.

  • Common strong acids include: $ext{HI}$, $ext{HBr}$, $ext{HCl}$, $ext{H}<em>2 ext{SO}</em>4$, and $ext{HNO}_3$.

  • Example: $ext{HCl} + ext{H}<em>2 ext{O} ightarrow ext{H}</em>3 ext{O}^+ + ext{Cl}^-$

8.3.2 Weak Acids

• Weak acids partially dissociate in water.

  • Unequal reaction arrows denote that the reactants are favored at equilibrium.

  • Common weak acids include: $ext{H}<em>3 ext{PO}</em>4$, $ext{HF}$, $ext{H}<em>2 ext{CO}</em>3$, and $ext{HCN}$.

8.3.3 Strong vs Weak Acids

• A strong acid, such as $ext{HCl}$, is entirely dissociated into ions.

• A weak acid, such as $ext{CH}_3 ext{COOH}$, predominantly exists in its undissociated form.

8.3.4 Strong Bases

• Strong bases dissolve completely in water (100% dissociation).

  • Common examples include sodium hydroxide ($ext{NaOH}$) and potassium hydroxide ($ext{KOH}$).

8.3.5 Weak Bases

• Weak bases dissolve with partial dissociation in water.

8.3.6 Conjugate Bases

• A strong acid readily donates a proton, producing a weak conjugate base.

• Conversely, a strong base readily accepts a proton, yielding a weak conjugate acid.

8.4 Dissociation of Water

8.4.1 Behavior of Water

• Water can act as both a Brønsted–Lowry acid and base, permitting two water molecules to engage in an acid-base reaction:

  • Example: $ext{H}<em>2 ext{O} + ext{H}</em>2 ext{O} <br>ightarrow ext{H}_3 ext{O}^+ + ext{OH}^-$

8.4.2 Ion-Product Constant for Water

• In pure water at 25°C, the concentrations of hydronium and hydroxide are $[H_3O^+] = [OH^-] = 1.0 imes 10^{-7} ext{ M}$.

• The ion-product constant for water ($K_w$) is dependent on these concentrations:

  • $K<em>w = [H</em>3O^+][OH^-] = 1.0 imes 10^{-14}$

8.4.3 Calculating Ion Concentration

• To calculate $[ ext{H}_3 ext{O}^+]$ or $[ ext{OH}^-]$ based on known quantities, utilize the relation derived from the constant:

  • $K_w = 1.0 imes 10^{-14}$

8.4.4 Applications to Blood pH

• Example: Blood with $[ ext{H}_3 ext{O}^+] = 3.98 imes 10^{-8} ext{ M}$ is considered basic.

• With the constant, calculate associated ion concentrations.

8.5 The pH Scale

8.5.1 pH Overview

• The pH scale indicates the acidity or basicity of solutions:

  • Acidic solution: pH < 7

  • Neutral solution: pH = 7

  • Basic solution: pH > 7

8.5.2 pH Examples

• Different substances demonstrate various pH values:

  • Common fruits are generally acidic (pH < 7).

  • Household ammonia and bleach are basic (pH > 7).

8.5.3 Calculating pH

• To calculate pH:

  • Formula: $ext{pH} = - ext{log}[ ext{H}_3 ext{O}^+]$

8.5.4 Using a Calculator

• Note that logarithmic calculations maintain the same decimal places as the digit count of the original number.

8.6 A. Reaction of Acids with Hydroxide Bases

8.6.1 Neutralization Overview

• A neutralization reaction involves an acid and a base producing salt and water.

  • Here, an acid $ext{HA}$ donates a proton to the base $ext{OH}^-$, yielding water and a salt, represented as:

  • $ext{HA} + ext{OH}^- <br>ightarrow ext{H}_2 ext{O} + ext{MA}$

8.6.2 Writing Balanced Equations

• Example Reaction:

  • $ext{HCl} + ext{Mg(OH)}<em>2 ightarrow ext{MgCl}</em>2 + ext{H}_2 ext{O}$

8.6.3 Balanced Equation Elements

• The salt formed consists of the remaining parts of the acid and base not used to generate water.

8.6.4 Identifying Net Ionic Equations

• A net ionic equation encapsulates only the species participating in the reaction, omitting spectator ions.

8.6B. Reaction of Acids with Bicarbonate and Carbonate

8.6B.1 Reactions with Bicarbonate

• The reaction of a bicarbonate base, such as $ext{HCO}_3^-$, with an acid yields carbonic acid, which subsequently decomposes.

  • $ext{HCO}<em>3^- + ext{H}^+ ightarrow ext{H}</em>2 ext{CO}<em>3 ightarrow ext{H}</em>2 ext{O} + ext{CO}_2(g)$

8.6B.2 Reactions with Carbonate

• A carbonate base can react similarly:

  • $ext{CO}<em>3^{2-} + 2 ext{H}^+ ightarrow ext{H}</em>2 ext{CO}<em>3 ightarrow ext{H}</em>2 ext{O} + ext{CO}_2(g)$

8.7 Titration

8.7.1 Titration Process

• Titration is a method to determine the concentration of an acid or base within a solution.

  • A base of known concentration is added gradually to an acid until neutralization is achieved when moles of acid equal moles of base.

8.7.2 Determining Unknown Concentrations

• Steps involved to calculate unknown molarity:

  1. Measure the volume of the NaOH solution used (known concentration).

  2. Calculate moles of NaOH.

  3. Find moles of HCl using stoichiometric conversion. 4. Calculate the molarity of the HCl solution.

8.7.3 Sample Problem

• Example:

  • Volume of base (NaOH): 22.5 mL

  • Concentration of base (NaOH): 0.100 M

  • Volume of acid (HCl): 25.0 mL

  • Calculate the molarity of the HCl solution based on titration data.

8.8 Buffers

8.8.1 Buffer Overview

• A buffer is a solution that minimizes changes in pH upon addition of acids or bases.

  • Typically composed of roughly equal amounts of a weak acid and its conjugate base salt.

  • Buffer Functionality:

  • Excess base reacts with weak acid.

  • Excess acid reacts with the conjugate base.

8.8.2 Buffer Behavior

• Illustrative examples showcase how buffer systems maintain a stable pH regardless of the addition of strong acids or bases.

8.8A General Characteristics of a Buffer

• When acid is added, it reacts with conjugate base, thus minimal pH change occurs.

• When base is added, excess reacts with conjugate acid, sustaining equilibrium in pH.

8.8B Focus on the Environment

• Acid Rain: As acid precipitation adds $ext{H}_3 ext{O}^+$ to water, it can lower pH levels.

• Buffered lakes: Lakes surrounded by limestone act as natural buffers, preventing drastic pH changes despite acid rain.

8.9 Buffers in the Human Body

8.9.1 Blood pH Regulation

• Normal blood pH fluctuates between 7.35 and 7.45, with carbonic acid and bicarbonate serving as the principal buffers.

• The produciton of carbon dioxide ($ext{CO}_2$) during metabolic processes directly influences blood pH.

8.9.2 Respiratory Disorders

• Respiratory Acidosis: Results from inadequate elimination of $ext{CO}_2$, often due to lung disease or failure.

• Respiratory Alkalosis: Caused by hyperventilation, leading to reduced production of $ext{CO}_2$ in the body.