Unit 13 cem acid equlib

Classification and pH of Common Substances

• The pH Scale: Substances are categorized by their pH value ranging from 0 to 14.     * Acidic: pH from 0 to less than 7.     * Neutral: pH of exactly 7.0.     * Alkaline or Basic: pH from greater than 7 to 14.

• pH Values of Common Acidic Substances:     * 0: Battery Acid.     * 1: Stomach Acid (Hydrochloric acid).     * 2: Lemon Juice, Vinegar.     * 2.5 to 3.0: Coke and Pepsi.     * 3: Grapefruit and Orange Juice.     * 3.5 to 4.0: Apples, Dr. Pepper Soda.     * 4: Tomato Juice, Beer.     * 4.5 to 5.0: Acid Rain, 7-UP Soda.     * 5: Black Coffee, Pepto Bismol.     * ~5.5: Healthy Skin, Hair, and Nails.     * 6: Urine, Saliva, Milk.

• pH Values of Neutral Substances:     * 7.0: Pure Water, Blood.     * 7.0 to 10.0: Shampoos.

• pH Values of Basic (Alkaline) Substances:     * 8: Baking Soda, Seawater, Eggs.     * 8.5 to 9.5: Perm Solutions.     * 9: Toothpaste, Hand Soap.     * 10: Milk of Magnesia, Mild Detergent.     * 11: Household Ammonia and Cleaners.     * 12: Soapy Water.     * 11.5 to 14.0: Hair Straighteners.     * 13: Bleach, Oven Cleaner.     * 14: Liquid Drain Cleaner, Caustic Soda.

Naming Conventions for Acids

• General Rule: Acid formulas begin with the element Hydrogen ($H$ ).

• Binary Acids: Formulated as $HX$ where $X$ is another element.     * Naming Format: hydro____ic acid.     * Example: $HF$ is hydrofluoric acid.

• Oxyacids: Formulated as $HX_O Y$ where $X_O Y$ is a polyatomic ion.     * Naming by Polyatomic Suffix:         * If the polyatomic ion ends with -ite, the acid name ends with -ous acid.         * If the polyatomic ion ends with -ate, the acid name ends with -ic acid.     * Mnemonics:         * "-ic" you "-ate" it!         * Spr-ite is delici-ous!     * Examples:         * $HNO_3$ (Nitrate ion) = nitric acid.         * $H_2SO_3$ (Sulfite ion) = sulfurous acid.

• Writing Formulas from Names: Charges must be balanced similarly to ionic compounds ($H^+$ and the anion).     * Sulfuric Acid: $H^+$ and $SO_4^{-2} \rightarrow H_2SO_4$.     * Hydrobromic Acid: $H^+$ and $Br^- \rightarrow HBr$.     * Nitrous Acid: $H^+$ and $NO_2^- \rightarrow HNO_2$.

Properties of Acids and Bases

• Properties of Acids:     * Taste: Sour.     * Skin Contact: Irritating and corrosive.     * Metal Reactivity: Reacts with many metals via single replacement reactions.     * pH Level: \text{pH} < 7.     * Conductivity: They are electrolytes.

• Properties of Bases:     * Taste: Bitter.     * Skin Contact: Irritating and caustic.     * Physical Feel: Slippery to the touch.     * pH Level: \text{pH} > 7.     * Conductivity: They are electrolytes.

Theoretical Definitions of Acids and Bases

• Arrhenius Theory:     * Acids: Formulas start with $H$ and produce Hydrogen ions ($H^+$) when dissolved in water.         * Example: $HNO_3 \rightarrow H^+ + NO_3^-$.     * Bases: Formulas end with $OH$ and produce hydroxide ions ($OH^-$) when dissolved in water.         * Example: $LiOH \rightarrow Li^+ + OH^-$.     * Neutralization Reaction: An Arrhenius acid and base react to form liquid water and a salt.         * Reaction: $HNO_3(aq) + LiOH(aq) \rightarrow H_2O(l) + LiNO_3(aq)$ .

• Brønsted-Lowry Theory:     * Acids: Proton ($H^+$) donors that form hydronium ions ($H_3O^+$) in water.         * Example: $HCl + H_2O \rightarrow H_3O^+ + Cl^-$.     * Bases: Proton ($H^+$) acceptors that form hydroxide ions ($OH^-$) in water.         * Example: $NH_3 + H_2O \rightarrow NH_4^+ + OH^-$.

Conjugate Acid-Base Pairs

• Definition: Products of a Brønsted-Lowry reaction are "conjugate" acids or bases. A conjugate pair differs by exactly one proton ($H^+$).

• Examples of Pairs:     * In the reaction $HCl + H_2O \rightarrow Cl^- + H_3O^+$:         * $HCl$ (Acid) / $Cl^-$ (Conjugate Base).         * $H_2O$ (Base) / $H_3O^+$ (Conjugate Acid).     * In the reaction $NH_3 + H_2O \rightarrow NH_4^+ + OH^-$:         * $NH_3$ (Base) / $NH_4^+$ (Conjugate Acid).         * $H_2O$ (Acid) / $OH^-$ (Conjugate Base).

• Practice Identifications:     * Conjugate Bases of Acids:         * $HBr \rightarrow Br^-$         * $HNO_3 \rightarrow NO_3^-$         * $H_2SO_4 \rightarrow HSO_4^-$     * Conjugate Acids of Bases:         * $F^- \rightarrow HF$         * $HSO_3^- \rightarrow H_2SO_3$         * $SO_3^{2-} \rightarrow HSO_3^-$

Buffers and pH Regulation

• Definition: A solution that resists changes in pH when small amounts of hydronium ($H_3O^+$) or hydroxide ($OH^-$) are added.

• Composition:     * A buffer is made from a weak acid and its conjugate base (e.g., $H_2CO_3$ and $HCO_3^-$).     * Alternatively, it can be made from a weak base and its conjugate acid.

• Applications: Blood, Kidneys, Swimming Pools, Beverages, Lotion, and Shampoo.

• Mechanism of Action:     * Neutralizing Added Acid: If $H_3O^+$ is added, the conjugate base reacts with it.         * Reaction: $H_3O^+ + HCO_3^- \rightarrow H_2O + H_2CO_3$.     * Neutralizing Added Base: If $OH^-$ is added, the weak acid reacts with it.         * Reaction: $OH^- + H_2CO_3 \rightarrow H_2O + HCO_3^-$.

• Buffer Preparation: Combine a weak acid with an ionic compound (salt) containing its conjugate base.     * Example: Combining $H_2CO_3$ and $NaHCO_3$ (which provides $HCO_3^-$).     * HF Buffer Example: To make a buffer with $HF$, combine it with $NaF$, $KF$, or $LiF$.     * Reactions for HF/NaF Buffer:         * Added Acid: $H_3O^+ + F^- \rightarrow H_2O + HF$.         * Added Base: $OH^- + HF \rightarrow H_2O + F^-$.

Quantifying Acid-Base Solutions

• Self-Ionization of Water: Water can behave as both an acid and a base.     * Equilibrium Equation: $H_2O(l) + H_2O(l) \rightleftharpoons H_3O^+(aq) + OH^-(aq)$.     * Ion Product Constant ($K_w$): At $25^oC$, $K_w = [H_3O^+][OH^-] = 1 \times 10^{-14}$.

• Relative Concentrations and Solution Type:     * Neutral: $[H_3O^+] = [OH^-] = 1 \times 10^{-7}\,M$.     * Acidic: [H_3O^+] > [OH^-].     * Basic: [OH^-] > [H_3O^+].

• Logarithmic Scales:     * pH Calculation: $\text{pH} = -\log_{10}[H_3O^+]$.     * pOH Calculation: $\text{pOH} = -\log_{10}[OH^-]$.     * Conversions:         * $[H_3O^+] = 10^{-\text{pH}}$.         * $[OH^-] = 10^{-\text{pOH}}$.         * $\text{pH} + \text{pOH} = 14$.     * Relationship of Scale: Every change of 1 in pH represents a 10x change in concentration ($[H_3O^+]$).

• Calculated Examples:     * Vinegar: Given $[H_3O^+] = 3.3 \times 10^{-3}\,M$.         * Calculation: $[OH^-] = \frac{1 \times 10^{-14}}{3.3 \times 10^{-3}} = 3.0 \times 10^{-12}\,M$.     * Multi-Factor Relationships Table:         * $[H_3O^+] = 1\,M \rightarrow [OH^-] = 1 \times 10^{-14}\,M, \text{pH} = 0.0, \text{pOH} = 14.0$ (Acidic).         * $[H_3O^+] = 1 \times 10^{-7}\,M \rightarrow [OH^-] = 1 \times 10^{-7}\,M, \text{pH} = 7.0, \text{pOH} = 7.0$ (Neutral).         * $[H_3O^+] = 1 \times 10^{-14}\,M \rightarrow [OH^-] = 1\,M, \text{pH} = 14.0, \text{pOH} = 0.0$ (Basic).

Strength of Acids and Bases

• Strong Acids: Completely dissociate/donate $H^+$ ($100\%$).     * The Big 6: $HCl$, $HBr$, $HI$, $HNO_3$, $HClO_4$, and $H_2SO_4$.     * Concentration Relationship: $1\,M\,HCl$ yields $1\,M\,H^+$.

• Weak Acids: Do not completely dissociate (< 100\%).     * Example: $H_3PO_4 + H_2O \rightleftharpoons H_2PO_4^- + H_3O^+$.     * Concentration Relationship: $1\,M\,H_3PO_4$ yields approximately $0.1\,M\,H^+$.

• Acid Dissociation Constant ($K_a$):     * Expression: $K_a = \frac{[H_3O^+][ ext{Conjugate Base}]}{[ ext{Acid}]}$.     * Strength correlation: Larger $K_a$ = stronger acid; Smaller $K_a$ = weaker acid.     * Ranking Example ($K_a$ values): HF ($7.2 \times 10^{-4}$) > HClO ($3.5 \times 10^{-8}$) > HCN ($6.2 \times 10^{-10}$).

• Strong Bases: Completely dissociate into $OH^-$ ($100\%$).     * Examples: Arrhenius bases such as $NaOH$, $LiOH$, $KOH$, $Ba(OH)_2$.     * Concentration Relationship: $1.0\,M\,NaOH$ yields $1.0\,M\,OH^-$.

• Weak Bases: Do not completely form $OH^-$ (< 100\%).     * Examples: Brønsted-Lowry bases such as $NH_3$, $F^-$, $CO_3^{2-}$.     * Base Dissociation Constant ($K_b$):         * Expression: $K_b = \frac{[NH_4^+][OH^-]}{[NH_3]}$.         * Strength correlation: Larger $K_b$ = stronger base; Smaller $K_b$ = weaker base.         * Ranking Example ($K_b$ values): $NH_3$ ($1.8 \times 10^{-5}$) > $CN^-$ ($1.6 \times 10^{-5}$) > $ClO^-$ ($2.9 \times 10^{-7}$).

• Strength of Conjugates: A stronger acid ($K_a$) corresponds to a weaker conjugate base.

Acid-Base Titrations

• Purpose: To determine the unknown concentration of an acid or base using a neutralization reaction.     * Reaction: Acid + Base $\rightarrow$ Water + Salt (Ionic Compound).     * Net Ionic focus: $H^+ + OH^- \rightarrow H_2O$.

• Key Terms:     * Equivalence Point: Point where moles of acid ($H^+$) are stoichiometrically equal to moles of base ($OH^-$). For strong acid/base, pH = 7.0.     * End Point: Point where the chemical indicator changes color (e.g., phenolphthalein turns light pink). In a successful titration, the end point approximates the equivalence point (differing by only one drop).

• Indicators and Color Change Ranges:     * Methyl violet: pH ~0 to 2.     * Thymol blue: pH ~1 to 3 and ~8 to 9.     * Methyl orange: pH ~3 to 4.5.     * Litmus: pH ~5 to 8.     * Bromothymol blue: pH ~6 to 7.5.     * Phenolphthalein: Turns pink in basic solutions (used for acid-base titrations).     * Alizarin yellow: pH ~10 to 12.

• Experimental Procedure:     1. Fill a buret with a known concentration of base (titrant).     2. Fill a flask with an unknown concentration of acid (analyte) and indicator.     3. Add base dropwise until every $H^+$ is neutralized by an $OH^-$.     4. Once neutralized, one extra drop of base makes the solution basic, triggering the indicator color change.

Laboratory Skills and Buret Management

• Reading a Buret:     * Burets are marked to the $0.1\,mL$ place; measurements must be estimated to the $0.01\,mL$ place.     * Volume dispensed = Final Volume - Initial Volume.     * Example: Initial 0.00 mL to Final 15.75 mL indicates $15.75\,mL$ dispensed.     * Example: An end point between 28 and 29 might lead to a reading like $28.85\,mL$.

• Preparing the Buret:     1. Fill with $NaOH$.     2. Drain a small amount to remove air bubbles from the tip.     3. Align the meniscus with the $0.00\,mL$ mark.

• Cleaning the Buret:     1. Drain remaining titrant.     2. Fill with distilled water and drain; repeat 3 times.     3. Rinse the outside and the tip with distilled water.     4. Dry by flipping the buret upside down with the stopcock open.