Chapter 8 Notes - Rachel Lusby PDF

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ACIDS AND BASES AND OXIDATION-REDUCTION Acids and Bases ● Acid: taste sour, dissolve some metals, cause plant dye to change color ● Bases: taste bitter, slippery, corrosive ● Two theories that helped us understand the chemistry of acids and bases ○ Arrhenius Theory ○ Bronsted-Lowry Theory ● Arrhenius Theory ○ Acid: a substance that dissociates to produce hydrogen ions when dissolved in water ■ Hydrogen ion: H+ also called proton ○ Base: a substance that dissociates to produce hydroxide ions when dissolved in water ■ NaOH is a base ○ Where does NH3 fit? ■ When it dissolves in water it has basic properties but it does not have OHions in it ■ The next acid-base theory gives us a broader view of acids and bases ● Bronsted-Lowry Theory ○ Acid – proton (H+) donor ○ Base - proton (H+) acceptor ■ Notice that acid and base are not defined using water ○ When writing the reactions, both accepting and donation are evident ○ HCl(aq) + H2O(l) → Cl−(aq) + H3O+(aq) ■ HCL donated the proton? ● Is it an acid or base? Acid ■ H2O accepted the proton ● Is it an acid or base? Base ■ Did NH3 donate or accept a proton ● Accept ■ Is it an acid or base ● Base ■ What is water in this reaction ● Acid ● Acid-Base Properties of Water ○ Water possesses both acid and base properties ■ Amphiprotic – a substance possessing both acid and base properties ■ Water is the most commonly used solvent for both acids and bases ■ Solute-solvent interactions between water and both acids and bases promote solubility and dissociation ● Acid and Base Strength ○ Acid and base strength – degree of dissociation ■ Not a measure of concentration

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Strong acids and bases – reaction with water is virtually 100% (Strong electrolytes) Strong Acids and Bases ○ Acids ■ HCl ■ HBr ■ HI ■ HNO3 ■ H2SO4 ■ HClO4 ○ Bases ■ NaOH ■ KOH ■ Ba(OH)2 ■ All metal hydroxides Weak Acids and Bases ○ Only a small amount disassociates ■ Weak electrolytes ○ Acid Examples ■ Acetic Acid ■ Carbonic Acid ○ Base Examples ■ Ammonia: NH3(aq) + H2O(l) ⇄ NH4+(aq) + OH-(aq) ■ Pyridine: C5H5NH2(aq) + H2O(l) ⇄ C5H5NH3+(aq) + OH-(aq) ■ Aniline: C6H5NH2(aq) + H2O(l) ⇄ C6H5NH3+(aq) + OH-(aq) Conjugate Acids and Bases ○ The acid base reaction can be written in the general form: ■ HA (Acid) + B (Base) ⇄ A- + BH+ ○ Notice the reversible arrows ○ The products are also an acid and base called the conjugate acid and base ○ Conjugate Acid – what the base becomes after it accepts a proton. ■ BH+ is the conjugate acid of the base B ○ Conjugate Base – what the acid becomes after it donates its proton ■ A− is the conjugate base of the acid HA ○ Conjugate Acid-Base Pair – the acid and base on the opposite sides of the equation ■ B and BH+ constitute a conjugate acid-base pair ■ HA and A − constitute a conjugate acid-base pair ■ HF is a stronger acid than HCN ■ CN- is a stronger base than H20 Acid-Base Dissociation ○ HA + B ⇆ A- + HB+ ○ The reversible arrow isn’t always written ■ Some acids or bases essentially dissociate 100%

■ One way arrow is used HCl + H2O → Cl- + H3O+ ■ All of the HCl is converted to Cl■ HCl is called a strong acid – an acid that dissociates 100% ○ Weak acid - one which does not dissociate 100% ● The Dissociation of Water ○ Pure water is virtually 100% molecular ○ Very small number of molecules dissociate ■ Dissociation of acids and bases is often called ionization ○ H2O(l) + H2O(l) ⇆ H3O+(aq) + OH-(aq) ■ Called autoionization ■ Very weak electrolyte ● Hydronium Ion ○ H3O+ is called the hydronium ion ○ In pure water at room temperature: ■ [H3O+] = 1 ◊ 10−7 M ■ [OH−] = 1 ◊ 10−7 M ○ What is the equilibrium expression for H2O(l) + H2O(l) ⇆ H3O+(aq) + OH−(aq) ■ Keq = [H3O+][OH−] ○ Remember, liquids are not included in equilibrium expressions ● Ion Product of Water ○ This constant is called the ion product of water and has the symbol Kw ○ Kw = [H3O+][OH−] ○ Since [H3O+] = [OH−] = 1.0 ◊ 10−7 M, what is the value for Kw? ■ 1.0 ◊ 10-14 ■ It is unitless pH: A Measurement Scale ● pH scale – a scale that indicates the acidity or basicity of a solution ○ Ranges from 0 (very acidic) to 14 (very basic) ● The pH scale is rather similar to the temperature scale assigning relative values of hot and cold ● The pH of a solution is defined as ○ pH = -log [H3O+]: ● Definition of pH ○ Use these observations to develop a concept of pH ■ if add an acid, [H3O+] ↑ and [OH-] ↓ ■ if add a base, [OH-] ↑ and [H3O+] ↓ ■ [H3O+] = [OH-] when equal amounts of acid and base are present ○ In each of these cases: 1.0 ◊ 10−14 = [H3O+][OH−] ● Measuring pH ○ pH of a solution can be: ○

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Calculated if the concentration of either ion is known Approximated using indicator / pH paper that develops a color related to the solution pH Measured using a pH meter whose sensor measures an electrical property of the solution that is proportional to pH

Calculate pH ○ pH = -log [H3O+]: ○ 1.10 x 10-14 = [H3O+][OH-]

Reactions Between Acids and Bases ● Neutralization reaction – the reaction of an acid with a base to produce a salt and water ○ HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l) ○ Acid Base Salt Water ● Break apart into ions: ● H+(aq) + Cl−(aq) + Na+(aq) + OH−(aq) → Na+(aq) + Cl−(aq) + H2O(l) ○ Net ionic equation ○ Show only the changed components ○ Omit any ions appearing the same on both sides of equation = Spectator Ions ■ H+(aq) + OH−(aq) → H2O(l) ● Net Ionic Neutralization Reaction ○ The net ionic neutralization reaction is more accurately written: ■ H3O+(aq) + OH−(aq) → 2H2O(l) ○ This equation applies to any strong acid / strong base neutralization reaction ○ An analytical technique to determine the concentration of an acid or base is titration ○ Titration involves the addition of measured amount of a standard solution to neutralize the second, unknown solution ○ Standard solution – solution of known concentration ● Acid Base Titration ○ Standard solution is slowly added until the color changes

○ The equivalence point is when the moles of H3O+ and OH- are equal ○ Buret: long glass tube calibrated in mL which contains the standard solution ○ Indicator: a substance which changes color as pH changes ○ Flask contains a solution of unknown concentration plus indicator ● Determine the Concentration of a Solution of Hydrochloric Acid ○ Place a known volume of acid whose concentration is not known into a flask ○ Add an indicator, here phenol red is good ○ Known concentration of NaOH is placed in a buret ○ Drip NaOH into the flask until the indicator changes color ○ Indicator changes color ■ equivalence point is reached ■ mol OH− = mol H3O+ present in the unknown acid ○ Volume dispensed from buret is determined ○ Calculate acid concentration from the following data: ■ Volume of HCl: 25.00 mL ■ Volume of NaOH added: 35.00 mL ■ Concentration of NaOH: 0.1000 M ■ Balanced reaction shows that 1 mol HCl reacts with 1 mol NaOH (a 1:1 ratio) ■ 3.500 ◊ 10−3 mol HCl is contained in 25.00 mL = 0.1400 M HCl ● Polyprotic Substances ○ The previous example have the acid and base at a 1:1 combining ratio ■ Not all acid-base pairs do this ○ Polyprotic substance – donates or accepts more than one proton per formula unit ■ HCl is monoprotic, producing one H+ ion for each unit of HCl ■ Sulfuric acid is diprotic, each unit of H2SO4 produces 2 H+ ions ○ H2SO4(aq) + 2NaOH(aq) → Na2SO4(aq) + 2 H2O(l) ■ Step 1: H2SO4(aq) + H2O(l) →HSO4−(aq) + H3O+(aq) ■ Step 2: HSO4−(aq) + H2O(l) ⇌ SO42−(aq) + H3O+(aq) ● In Step 1 H2SO4 behaves as a strong acid – dissociating completely ● In Step 2 HSO4− (behaves as a weak acid--reversibly dissociating; note the double arrow) Acid-Base Buffers ● Buffer solution - solution which resists large changes in pH when either acids or bases are added ● These solutions are frequently prepared in laboratories to maintain optimum conditions for chemical reactions ● Blood is a complex natural buffer solution maintaining a pH of ~7.4 using mainly carbonic acid (H2CO3) and bicarbonate (HCO3−) ions. ● Buffer Capacity ○ A measure of the ability of a solution to resist large changes in pH when a strong acid or strong base is added

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Also described as the amount of strong acid or strong base that a buffer can neutralize without significantly changing pH Addition of Base (OH-) to a Buffer Solution ○ Adding a basic substance to a buffer causes changes: ■ The OH- will react with the H3O+ producing water ■ Acid in the buffer system dissociates to replace the H3O+ consumed by the added base ■ Net result is to maintain the pH close to the initial level ○ The loss of H3O+ (the stress) is compensated by the dissociation of the acid to produce more H3O+ Addition of Acid (H 3O+) to a Buffer Solution ○ Adding an acidic substance to a buffer causes changes ■ The H3O+ from the acid will increase the overall [H3O+] ■ Conjugate base in the buffer system reacts with the H3O+ to form more acid ■ Net result is to maintain the H3O+ concentration and the pH close to the initial level ○ The gain of H3O+ (the stress) is compensated by the reaction of the conjugate base to produce more acid

○ Preparation of a Buffer Solution ○ Buffering process is an equilibrium reaction described by an equilibrium-constant expression ○ If you want to know the pH of the buffer, solve for H3O+ then calculate the pH ○ CH3COOH (aq) + H2O (l) CH3COO- (aq) + H3O+ (aq) ○ In acids, this constant is Ka ■

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𝐾 = 𝑎

[𝐻3𝑂+][𝐶𝐻3𝐶𝑂𝑂−] [𝐶𝐻3𝐶𝑂𝑂𝐻]

Henderson-Hasselbalch Equation ○ Solution of equilibrium-constant expression and pH can be combined into one operation ○ Henderson-Hasselbalch Equation is this combined expression ○ For the acetic acid/sodium acetate buffer system: ■

pH = pKa + log

[𝑐𝑜𝑛𝑗𝑢𝑔𝑎𝑡𝑒 𝑏𝑎𝑠𝑒] [𝑤𝑒𝑎𝑘 𝑎𝑐𝑖𝑑]...