Acid-Base Equilibria Lecture Notes PDF

Previously we learned Chapter 13. Fundamental Equilibrium Concepts 13.1 Chemical Equilibrium mA + nB xC + yD 13.2 Equilibrium Constants 𝑄𝐶 at equilibrium = 𝐾𝑐 =¿¿ 13.3 Shifting Equilibria: Le Châtelier’s Principle 13.4 Equilibrium Calculations  Previous...

Previously we learned Chapter 13. Fundamental Equilibrium Concepts 13.1 Chemical Equilibrium mA + nB xC + yD 13.2 Equilibrium Constants 𝑄𝐶 at equilibrium = 𝐾𝑐 =¿¿ 13.3 Shifting Equilibria: Le Châtelier’s Principle 13.4 Equilibrium Calculations  Previously we learned Chapter 13. Fundamental Equilibrium Concepts 13.1 Chemical Equilibrium mA + nB xC + yD 13.2 Equilibrium Constants 𝑄𝐶 at equilibrium = 𝐾𝑐 =¿¿ 13.3 Shifting Equilibria: Le Châtelier’s Principle Qc vs. Kc 13.4 Equilibrium Calculations Chapter 10: Molecular solid Chapter 11: Dissolution, electrolyte Chapter 12: Equilibrium Chapter 13: Acid  Today we will learn Chapter 14 Acid-Base Equilibria 14.1 Brønsted-Lowry Acids and Bases 14.2 pH and pOH 14.3 Relative Strengths of Acids and Bases 14.4 Hydrolysis of Salts 14.5 Polyprotic Acids 14.6 Buffers 14.7 Acid-Base Titrations Sinkholes are the result of reactions between acidic groundwaters and basic rock formations (limestone). H2CO3(aq) + CaCO3(s) ⇌ Ca2+(aq) + 2HCO3-(aq)  Objectives: 14.1 Brønsted-Lowry Acids and Bases Definition of Brønsted-Lowry Identify acids, bases, and conjugate acid-base pairs Describe the acid-base behavior of amphiprotic substances Write equations for acid/base ionization reactions calculate hydronium (H3O+) and hydroxide (OH-) concentrations Acid and Base Definitions Johannes Brønsted and Thomas Lowry (1923) definition: Focused on the reactions between acids and bases. Brønsted-Lowry Acid: A compound that donates a proton to another compound. Brønsted-Lowry Base: A compound that accepts a proton from another compound. What Types of Compounds are Acids? Acids: A compound that donates a proton to another compound.  Molecules H2O, HF, H2CO3, H2SO3, HNO3  Cations that contain H H3O+, NH4+  Anions that contain H HCO3-, HS- What Types of Compounds are Bases? Bases: A compound that accepts a proton from another compound. Ionic compounds that contain -OH HCO3-, HSO3-, H2PO4-, Molecules H2O, NH3 Anions – With or without H OH-, ClO4-, HS- Cations – Few examples Brønsted-Lowry Acid/Base Theory The species formed when a proton is removed from an acid is the conjugate base of that acid. The species formed when a proton is added to a base is the conjugate acid of that base. Brønsted-Lowry Acid/Base Theory The species formed when a proton is removed from an acid is the conjugate base of that acid. The species formed when a proton is added to a base is the conjugate acid of that base.  Brønsted-Lowry Acid/Base Theory The species formed when a proton is removed from an acid is the conjugate base of that acid. conjugate acid conjugate base base acid The species formed when a proton is added to a base is the conjugate acid of that base. Amphiprotic Species A species that can either accept or donate a proton is called amphiprotic (or amphoteric). An amphiprotic species can serve as an acid or a base. Example: Water In the presence of a base, water acts as an acid. Amphiprotic Species A species that can either accept or donate a proton is called amphiprotic (or amphoteric). An amphiprotic species can serve as an acid or a base. Example: Water In the presence of a base, water acts as an acid. In the presence of an acid, water acts as a base.  Amphiprotic Species Examples: H2O HCO3- HS- HSO3- … Amphiprotic Species  Amphiprotic Species: Water as an Acid The reaction between a Brønsted-Lowry base and water is called base ionization.  Amphiprotic Species: Water as a Base The reaction between a Brønsted-Lowry acid and water is called acid ionization.  The Autoionization of Water In pure water, water also acts both as an acid and a base. Autoionization of water: The equilibrium constant for this reaction is called the ion product constant for water, Kw Kw = [H3O+] [OH-] = 1.0 × 10–14 (25 °C) Notice that H2O is not a part of the expression for Kw. Why? Hint:  Today we will learn Chapter 14 Acid-Base Equilibria 14.1 Brønsted-Lowry Acids and Bases proton donor vs. proton acceptor Kw = [H3O+] [OH-] = 1.0 × 10–14 (25 °C) 14.2 pH and pOH 14.3 Relative Strengths of Acids and Bases 14.4 Hydrolysis of Salts 14.5 Polyprotic Acids 14.6 Buffers 14.7 Acid-Base Titrations  H3O+ and OH– Concentrations Kw = [H3O+][OH–] = 1.0 × 10–14 (25 °C) [H3O+] and [OH–] are inversely related to each other. In pure water, [H3O+] and [OH–] are equal to each other. This solution is neutral. [H3O+] = [OH–] = 1.0 × 10–7 M When an acid or base is dissolved, [H3O+] and [OH–] are not equal to each other. If [H3O+] > [OH–], the solution is called acidic. If [OH–] > [H3O+], the solution is called basic. H3O+ and OH– Concentrations Kw = [H3O+][OH–] = 1.0 × 10–14 (25 °C) [H3O+] and [OH–] are inversely related to each other. The acidity and basicity of a solution can be expressed in terms of its [H3O+] or [OH–]. These concentrations can span many orders of magnitude, therefore logarithmic values are commonly used. H3O+ and OH– Concentrations Kw = [H3O+][OH–] = 1.0 × 10–14 (25 °C) [H3O+] and [OH–] are inversely related to each other. pH & pOH pH = -log[H3O+] [H3O+] = 10-pH (this is the antilog function) pOH = -log[OH-] [OH-] = 10-pOH (this is the antilog function) pH = - pH + pOH = pOH = - log[H3O+] log[OH-] 14 Table 14.1. Summary of Relations for Acidic, Basic, and Neutral Solutions Classification Relative Ion Concentrations pH at 25 °C acidic [H3O+] > 10-7 M > [OH–] pH < 7 neutral [H3O+] = 10-7 M = [OH–] pH = 7 basic [H3O+] < 10-7 M < [OH–] pH > 7 acidic neutral basic https://www.differencebetween.com/difference-between-ph- and-vs-poh/ pH = -log[H3O+] pH = -log[H3O+] [H3O+] = 10-pH pH = -log[H3O+] pH + pOH = 14  pH = 2.5 3.1 3.1 3.4 https:// springsoralhealth.wordpress.co m/2012/02/01/ph-of-soft-drinks/ Figure 14.2 The pH and pOH values of some common substances at 25 °C. Figure 14.3. Rain water is acidic (a) Acid rain makes trees more susceptible to drought and insect infestation, and depletes nutrients in the soil. (b) It also corrodes statues that are carved from marble or limestone. H2CO3(aq) + CaCO3(s) ⇌ Ca2+(aq) + 2HCO3-(aq) Figure 14.4 A research-grade pH meter used in a laboratory can have a resolution of 0.001 pH units, an accuracy of ± 0.002 pH units, and may cost in excess of $1000. Convenient test strips, called pH paper, contain embedded indicator dyes that yield pH-dependent color changes on contact with aqueous solutions. Images from Amazon.com Today we learned Chapter 14 Acid-Base Equilibria 14.1 Brønsted-Lowry Acids and Bases proton donor vs. proton acceptor Kw = [H3O+] [OH-] = 1.0 × 10–14 (25 °C) 14.2 pH and pOH pH = -log[H3O+] pOH = -log[OH-] pH + pOH = 14 14.3 Relative Strengths of Acids and Bases 14.4 Hydrolysis of Salts 14.5 Polyprotic Acids 14.6 Buffers 14.7 Acid-Base Titrations Next time we will learn Chapter 14 Acid-Base Equilibria 14.1 Brønsted-Lowry Acids and Bases proton donor vs. proton acceptor Kw = [H3O+] [OH-] = 1.0 × 10–14 (25 °C) 14.2 pH and pOH pH = -log[H3O+] pOH = -log[OH-] pH + pOH = 14 14.3 Relative Strengths of Acids and Bases 14.4 Hydrolysis of Salts 14.5 Polyprotic Acids 14.6 Buffers 14.7 Acid-Base Titrations