Chapter 8: Acids, Bases, Oxidation Reduction
Chapter 8: Acids, Bases, Oxidation Reduction Chem 110
Popular in Chem
Popular in Chemistry
Mariah J. Bibbs
verified elite notetaker
verified elite notetaker
verified elite notetaker
verified elite notetaker
verified elite notetaker
Chem 471 (Biological Chemistry Lab)
verified elite notetaker
This 10 page Class Notes was uploaded by Justyna Jaworski on Wednesday December 2, 2015. The Class Notes belongs to Chem 110 at Northern Illinois University taught by M. Leifker in Summer 2015. Since its upload, it has received 13 views. For similar materials see Chem in Chemistry at Northern Illinois University.
Reviews for Chapter 8: Acids, Bases, Oxidation Reduction
Report this Material
What is Karma?
Karma is the currency of StudySoup.
You can buy or earn more Karma at anytime and redeem it for class notes, study guides, flashcards, and more!
Date Created: 12/02/15
8.1 Acids and Bases ● Acids taste sour, dissolve some metals, cause plant dye to change color ● Bases taste bitter, are slippery, corrosive ● Two theories that help us to understand the chemistry of acids and bases ○ Arrhenius Theory ○ Bronsted Lowry Theory Arrhenius Theory of Acids and Bases ● Acid a substance, when dissolved in water, dissociates to produce hydrogen ions: ○ Hydrogen ion H+ also called “proton” ○ HCl is an acid:HCl(aq) → + (aq) Cl (aq) ● Base a substance, when dissolved in water, dissociates to produce hydroxide ions ○ NaOH is a base: NaH (aq) → Na+ (aq) + OH (aq) Where does NH3 fit? When it dissolves in water it has basic properties but it does not have OH ions in it Bronsted Lowry Theory of Acids and Bases ● 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) What donated the proton? HCl because Cl has no H and it is Cl Is it an acid or baseAcid What accepted the proton? H2O Is it an acid or baseBase NH3 (aq) + H2O (l) double arrow NH4+ (aq) + OH (aq) Did NH3 donate or accept a proton? Accept, making it a base What is water in this reactioACID water can be both because its weak at both AcidBase Properties of Water ● Water possesses both acid and base properties ○ amphiprotic a substance processing 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 ○ strong acids and bases reaction with water is virtually 100% (strong electrolytes) ○ Example: ■ HCl (aq) + H2O (l) → H3O + (aq) Cl (aq) ● HCl is a strong acid ■ KOH (aq) → K+ (aq) + OH (aq) ● KOH is a strong base Strong Acids and Bases Strong Acids: ● HCl, HBr, HI Hydrochloric acid etc. ● HNO3 Nitric acid ● H2SO4 Sulfuric Acid ● HClO4 Perchloric Acid Strong Bases: ● NaOH, KOH, Ba(OH)2 ● All metal hydroxides Weak Acids ● Weak acids and bases only a small percent dissociates (weak electrolytes) ● Weak acid examples: ○ Acetic Acid ■ CH3COOH (aq) + H2O (l) double arrow CH3COO (aq)+H3O+ (aq) ○ Carbonic Acid: ■ H2CO3 (aq) + H2O (l) double arrow HCO3 (aq) + H3O + (aq) Conjugate Acids and Bases ● The acid base reaction can be written in the general form: ○ HA +B double arrow 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 form of the base B ● Conjugate base what the acid becomes after it donates its proton ○ A is the conjugate base of the acid HA AcidBase Dissociation HA + B double arrow A + HB+ ● The reversible arrow isn’t always written ○ some acids or bases essentially dissociate 100% ○ one way the 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% Acid Base Practice Write chemical reaction for following acids or bases in water Identify conjugate acid base pairs 1. HF (a weak acid) HF + H2O double arrow because weak F + H3O+ 2. H2S (weak acid) H2S + H2O double arrow HS + H3O + 3. HNO3 (strong acid) HNO3 + H2O → because strong acid NO3 + H3O+ 4. ch3 always weak CH3NH2 (weak base) CH3NH2 + H2O double CH3NH3+ OH Dissociation of Water ● Pure water is virtually 100% molecular ● Very small number of molecules dissociate ○ dissociation of acids and bases is often callionization ○ H2O (l) + H2O (l) double H3O + (aq) + OH (aq) ■ Called autoionization ■ very weak electrolyte Hydronium Ion ● H3O+ is called the hydronium ion ● In pure water at room temperature: ○ [H3O+] = 1x10^7 M ○ [OH] = 1x10^7 M ● What is the equilibrium expression for H2O(l) + H2O(l) double H3O+ (aq) + OH (aq) ○ Keq= [H3O+][OH] ○ Remember, liquids are not included in the equilibrium expressions Ion Product of Water ● This constant is called the ion product for water and has the symbol of Kw ○ Kw = [H3O+][OH] ● Since [H3O+] = [OH] = 1.0 x 10^7 M, what is the value for Kw? 8.2 pH: A Measurement Scale for Acids and Bases ● 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+] A Definition of pH ● Use these observations to develop a concept of pH ○ if add an acid, [H3O+] increases and [OH] decreases ○ if add base, [OH] increases and other decreases ○ [H3O+] = [OH] when equal amounts of acid and base are present ● in each of these cases 1.0x10^14 = [H3O+][OH] Measuring pH ● pH of a solution can be: ○ calculated if the concentration of either ion is known ○ approximated using indicator/ pH paper that develops a color related to the solution ○ Measured using a pH meter whose sensor measures an electrical property of the solution that is proportional to pH Calculating pH ● How do we calculate the pH of a solution when either the H3O+ or OH ion concentrations is known? ● how do we calculate the H3O+ or OH ion concentration when the pH is known? ● use these facts: ○ 1.0x10^14 = [H3O+][OH] (WEAK ACID ) ○ pH= log[H3O+] FOR A STRONG ACID USE THIS) Calculating pH from Acid Molarity What is the pH of a 1.0 x 10^4 M HCl solution? ● HCl is a strong acid and dissociates in water ● 1 mol HCl produces 1 mol [H3O+] in solution ● Therefore, 1.0 x 10^4 M HCl solution has [H3O+] = 1.0 x 10^4 M ● pH= log[H3O+] ● pH= log (1.0x10^4) ● PH = 4.0 What is the [H3O+] of a solution with pH= 6.0? ● pH=log[H3O+] ● An alternative math form of eq is [H3O+] = 10^pH ○ [H3O+]= 10^6 ■ 1.0 x 10^6 M Calculating the pH of a base What is the pH of a 1.0 x 10^3 M KOH solution? ● KOH is a strong base, so it dissociates completely ● 1 mol KOH produces 1 mol OH ion in solution ● Therefore, 1.0 x 10^3 M KOH has [OH] = 1.0 x 10p^3 M ● To use the pH formula, we need to know the [H3O+] instead: ○ pH= log[H3O+] ○ Kw= [H3O+][OH] ■ Kw= 1.0 x 10^14 Acid 1 step base 2 steps The importance of pH and pH Control ● any change that takes place in aqueous solution generally has at least some pH dependence ○ Agriculturecrops grow best in soil with proper pH ○ Physiology blood pH shift of 1 pH is fatal ○ Acid rain lowers pH of water in aquatic system causing problems for native fishes ○ Municipal services sewage treatment and water purification require optimal pH ○ Industry many processes require strict pH control for costeffective production 8.3 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) ○ when you mix these, you always get a salt and 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 Determine concentration of solution of HCl acid ● indicator changes color and equivalence point is reached ○ mol OH = mol H3O+ present in unknown acid ● volume dispensed from buret is determined ● calculate acid concentration from the following data: ○ volume of HCl: 25.00 mL ○ V of NaOH added: 35.00 mL ○ Concentration of NaOH: 0.1 M ○ Balanced reaction shows that 1 mol HCl reacts with 1 mol NaOH (a 1:1 ratio) Polyprotic Substances ● The previous example have the acid base at 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 ● HCl(aq) + NaOH(aq) → H2O(l) + Na+ (aq) + Cl(aq) reactions of polyprotic substances ● sulfiric acid is diprotic, each unit of H2SO4 produces 2 H+ ions ○ H2SO4+2NaOH(aq) → Na2SO4+2H2O ● Phosphoric acid is triprotic, each unit of H3PO4 produces 3H+ ions ○ H3PO4+3NaOH→ H2O+3Na……. egrwdgvc Dissociation of polyprotic substances step1: H2SO4+H2O→ HSO4+H3O+ step2: HSO4 +H2O double arrow SO4^2 +H3O+ in step 1 HSO4 behaves as a strong acid dissociating completely In step 2 HSO4 (behaves as weak acid, reversibly dissociating, note the double arrow EX Phosphoric acid dissociates in 3 steps all behaving as weak acids step 1: H3PO4 + H2O double H2PO4 + H3O+ step 2: H2PO4 + H2O double HPO4 2 + H3O+ Step 3: HPO4 2 + H2O double PO4 3 + H3O+ 8.4 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 larbs to maintain optimum conditions for chemical reactions ● blood is a complex natural buffer solution maintaining pH of 7.4 using mainly carbonic acid H2CO3 and bicarbonate HCO3 ^ ions The buffer process ● buffers act to establish an equilibrium between a conjugate acid/base pair ● buffers consist of either a weak acid and its salt( conjugate base) or a weak base and its salt( conjugate acid) ● an equilibrium is established in solution between the acid and base Buffer Capacity ● 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 ● also described as the amount of strong acid or strong base that a buffer can neutralize without significantly changing pH Preparation of a Buffer Solution ● buffering process is an equilibrium reaction described by an equilibriumcontrast expression ○ CH3COOH(aq) + H2Ol double arrow CH3COO(aq) + H3O+(aq) ■ In acids, this constant is Ka ■ ka= [H3O+][CH3COO]/[CH3COOH] ○ If you want to know the ph of the buffer, solve for [H3O+], then calculate pH Calculating the pH of a buffer solution assume the [CH3COOH] represents the concentration of the acid component of the buffer assume the [CH3COO] represents the concentration of the conjugate base(salt) component of the buffer [H3O+] = Ka[CH3COOH]/[CH3COO] [H3O+]= ka[acid]/[conjugate base] Calculating the pH of a buffer solution in which: both acetic acid and sodium acetate concentrations are 2.0 x 10^2 M the Ka for acetic acid is 1.75 x 10^5 use equation and get [H3O+]= 1.75 x 10^5 M pH= log[H3O+] final pH of 4.76 HendersonHasselbalch Equation ● solution of equilibriumconstant expression and pH can be combined into one operation ● henderson hasselbalch eq is this combined expression ● for acetic acid/sodium acetate buffer system: ○ Ka= [H3O+][CH3COO]/[CH3COOH] ● The expression is generalized as ○ pH= pKa + log [ (conjugate base)/(weak acid)] Using this eq calc pH of buffer sol where acetic acid = 2.0 x 10^1 M, sodium acetate = 1.0x 10^1 M, Ka= 1.8x10^5 Control of Blood pH ● ph of 7.4 maintained in blood partly by a complex carbon dioxide carbonic acid bicarbonate buffer system ● CO2(aq)+2H2O(l) double arrow H2CO3(aq) + H2O(l) double arrow HCO3(aq)+H3O+(aq) ● Higher than normal CO2 shift the equilibrium to the right, increasing H3O+ which lowers the pH ○ termed acidosis ○ leads to emphysema, pneumonia ● Lower than normal CO2 shift the equilibrium to left, decreasing H3O+ which raises the pH ○ termed alkalosis ○ leads to hyperventilation 8.5 Oxidation Reduction Processes ● Oxidation reduction processes are responsible for many types of chemical change ● Oxidation defined by one of following ○ loss of electrons ○ loss of hydrogen atoms ○ gain of oxygen atoms ● ex) Mg → Mg2+ + 2e ○ oxidation half reaction ● Cannot have oxidation without reduction Oxidation and Reduction as Complementary process Oxidizing agent is reduced gains electrons causes oxidation Reducing agent is oxidized loses electrons causes reduction Application of Oxidation and Reduction ● corrosion the deterioration of metals caused by an oxidation reduction process ○ ex) rust (oxidation of iron) ○ 4Fe(s) +3O2(g) → 2Fe2O3(s) ● Combustion of fossil fuels ○ ex) natural gas furnaces ● Bleaching ○ most bleaching agents are oxidizing agents ○ the oxidation of the stains produce compounds that do not have color Voltaic Cells electrochemical cell that converts stored chemical energy into electrical energy consider: Zn(s) + Cu2+(aq) → Zn2+ (aq) + Cu(s) Zn is being oxidized Cu is reduced ● If the two reactants are placed in the same flask they cannot produce electrical current ● a voltaic cell separates the two half reactions ● this makes the electrons flow through a wire to allow the oxidation and reduction to occur Voltaic cell generating electrical current ● Zn → Zn2+ +2e ○ oxidation ■ anode electrode where oxidation occurs ● Cu2+ + 2e → Cu ○ Reduction ■ cathodeelectrode where reduction occurs Silver battery ● batteries use the concept of voltaic cell ● modern batteries are smaller, safer, more dependable Electrolysis ● electrolysis reactions uses electrical energy to cause non spontaneous oxidation reduction reactions to occur ● these reactions are the reverse of a voltaic cell ○ rechargeable battery ■ when powering device behaves as voltaic cell ■ with time the chem reaction nears completion ■ battery appears run down ■ cell reaction is reversible when battery attached to charger
Are you sure you want to buy this material for
You're already Subscribed!
Looks like you've already subscribed to StudySoup, you won't need to purchase another subscription to get this material. To access this material simply click 'View Full Document'