71% of the earth s surface is covered by water, and water makes up ~70% of the mass of a human body.

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1 The Major Classes of Chemical Reactions Chapter 4 Learning Goals: Chapter The Role of Water as a Solvent: Terms 4.2 Writing Equations for Aqueous Ionic Reactions 4.3 Precipitation Reactions 4.4 AcidBase Reactions 4.5 OxidationReduction (Redox) Reactions 4.6 Elements in Redox Reactions 4.7 Reversible Reactions: An Introduction to Chemical Equilibrium 71% of the earth s surface is covered by water, and water makes up ~70% of the mass of a human body. Water is a major constituent in all life forms. Understanding the chemistry of water provides useful information used in many real world applications. Learn some general aqueous solution classes of chemical reactions: precipitation, acidbase (neutralization), oxidationreduction. Composition of A Bacterium Substance % Mass of Number Types Cell of Molecule Water X Ions 1 20? Water is a polar molecule due to oxygen s ability to attract electrons (called electronegativity) more so than hydrogen. e O e It s just symbolic language Sugars X 10 8 H H Amino Acids X 10 7 Lipids X 10 7 Nucleotides X 10 7 Small molecules Large Molecules 0.2 ~200? 23 ~ X 10 6 H O H Oxygen atom sucks electron density from hydrogen atoms leading to a polar molecule.

2 Ionic compounds dissolve when ionwater forces are larger than ionion forces holding the ionic crystal together. Likes dissolve likes Ionic substances conduct electricity, covalent substances do not. Strong Electrolyte 100% dissociation into ions NaCl (s) H 2O Na+ (aq) + Cl (aq) Weak Electrolyte not completely dissociated CH3COOH CH3COO (aq) + H+ (aq) NonElectrolyte no dissociation CH3OH The dissolution of an ionic compound in water An electrolyte is a substance that when dissolved in water results in a solution that can conduct electricity. A nonelectrolyte is a substance that when dissolved, results in a solution that does not conduct electricity. Solubility (S) is the maximum amount of a solute that can dissolve in a fixed quantity of a solvent at a specified temperature. (Units of g solute/100 g water) Examples: Sucrose (sugar) 203 g per 100 g H2O NaCl g per 100 g H2O (very soluble) AgCl g per 100 g H2O (insoluble) Nonelectrolyte No Dissociation Strong Electrolyte 100% Dissociation Weak Electrolyte Little dissociation Concentration of Solute Chemists also describe solutions as being 1 of 3 kinds. A supersaturated solution contains more solute than is present in a saturated solution at a specific temperature. A saturated solution contains the maximum amount of a solute that will dissolve in a given solvent at a specific temperature. An unsaturated solution contains less solute than the solvent has the capacity to dissolve at a specific temperature. Ionic compounds dissociate into ions thus its chemical formula tells us the number of moles of different ions in solution. How many moles of each ion are in the following solutions? (a) 5.0 mol of ammonium sulfate dissolved in water (b) 78.5 g of cesium bromide dissolved in water (c) 7.42 x 1022 formula units of copper(ii) nitrate dissolved in water Covalent compounds do not dissociate! (e) 35 ml of 0.84 M glucose (C6H12O6)

3 H 2 O (a) (NH 4 ) 2 SO 4 (s) (b) CsBr(s) Cs + (aq) + Br (aq) mol CsBr 78.5 g CsBr = mol CsBr g CsBr (c) Cu(NO 3 ) 2 (s) H 2 O 2NH 4 + (aq) + SO 4 2 (aq) 5.0 mol (NH 4 ) 2 SO 4 2 mol NH mol (NH 4 ) H 2 O 2 SO 4 = 10. mol NH 4 + Cu 2+ (aq) + 2NO 3 (aq) 5.0 mol SO 4 2 = mol Cs + = mol Br = mol Cu x10 22 formula mol Cu(NO 3 ) 2 = mol Cu(NO 3 ) 2 units Cu(NO 3 ) x10 23 formula units = mol NO 3 H 2 O (d) ZnCl 2 (aq) Zn 2+ (aq) + 2Cl (aq) 1L 0.84 mol ZnCl 2 35 ml ZnCl = 2.9x110 ml L 2 mol ZnCl 2 = 2.9x110 2 mol Zn 2+ = 5.8x10 2 mol Cl What does 3.5 M FeCl3 mean? 3.5 M FeCl3 = 3.5 moles FeCl3 1 Liter solution (1) a homogeneous solution of 3.5 moles of dry 100% pure FeCl3 dissolved in 1.00 Liter total solution volume (not 1 L of liquid!). (3) Note: It does not mean 3.5 moles of FeCl3 is dissolved in 1.00 liter of water! (4) [Fe 3+ ] = 3.5M and [Cl ] = 3 x 3.5 M (5) It can be used as a conversion factor There are 3classes of chemical reactions that occur in aqueous solution. Precipitation AcidBase Neutralization Oxidation Reduction 1. Precipitation Reaction an insoluble solid is formed from specific cationanion combinations. Involves Substances Involves Substances Cations Anions H + ions OH ions Involves Substances Oxidation Reduction 2. AcidBase Reaction a protons donor substance reacts with a hydroxide donor substance forming a salt and water. Combine to Form Insoluble Precipitate Combine to Form Salt and H2O Loss of e Which is the Gain of e 3. OxidationReduction Reactionelectron donor substances react with react with substances that accept electrons. Predicted by Solubility Rules Reducing Agent Which is called Oxidizing Agent In a precipitation reactions a metal cation and nonmetal anion combine & form an insoluble solid that precipitates. AgNO3(aq) + Na2CrO4(aq) => Ag2CrO4 (s) + NaNO3(aq) How Do We Predict If A Precipitate Rx Will Occur? We memorize solubility rules and apply the rule if it can happen, it will happen. Precipitation occurs when ionic attractive forces overcome water s tendency to hydrate and dissolve. Spectator ions: any species that remains soluble.

4 Solubility Rules For Ionic Compounds Silberberg Soluble Ionic Compounds 1. Salts of Group 1A and ammonium ion (NH 4+ ) 2. NO 3, CH 3 COO or C 2 H 3 O 2, ClO 4 3. Cl, Br I except those of Ag +, Pb 2+, Cu +, and Hg The driving force of a precipitation reaction is the act of PRECIPITATION. If it can happen, it will. Insoluble Ionic Compounds 1. All OH, except those of Group 1A and the larger members of Group 2A (beginning with Ca 2+ ). 2. CO 3 2 and PO 4 3 ) are insoluble, except those of Group 1A(1) and NH All S 2 except those of Group 1A, Group 2A and NH 4+. How To Predict Whether A Precipitation Reaction Example: Pb(NO 3 ) 2 (aq) + NaI(aq) ==>? Look at ions that are possible from the reactants, then use Solubility Rules! 1. Note the ions present in the reactants. 2. Consider the possible cationanion combinations. 3. Refer to the table of solubility rules and decide whether any of the ion combinations is insoluble. 4. If a combination is insoluble, that RXN will occur. 5. Write the molecular, ionic and net ionic equation for the reaction. Example: The reaction of the salts: Pb(NO 3 ) 2 (aq) + NaI(aq) ==> PbI2(s) 1. Recognize => IONIC => Solubility Rules Apply 2. Consult the Solubility table to see if PbI 2 or NaNO3 are insoluble. 3. PbI 2 is insoluble so that reaction will occur. 4. Now write 3equations 1)molecular, 2) ionic and 3) net equation as: 2NaI(aq) + Pb(NO 3 ) 2 (aq) 2I (aq) + Pb 2+ (aq) PbI 2 (s) PbI 2 (s) + 2NaNO 3 (aq) 2Na + + 2I + Pb NO 3 PbI2 + 2Na + + 2NO 3 There are 3types of equations that are written for precipitation, acidbase and redox reactions 1. The molecular equation shows all reactants and products as undissociated compounds. precipitate 2NaI(aq) + Pb(NO 3 ) 2 (aq) 2. The total ionic equation shows all of the soluble ionic substances dissociated into ions. Pb NO 3 + 2Na + + 2I PbI 2 (s) + 2Na + + 2NO 3 Na + and NO 3 are spectator ionsthey don t do much but sit there! 3. The net ionic equation PbI 2 (s) + 2NaNO 3 (aq) eliminates the spectator ions and shows the actual chemical change taking place. Pb I PbI 2 (s) Putting It All Together: Predicting Precipitation 0. Know how to read a Table of Solubility Rules 1. Know ionic nomenclature so you can write the correct ionic formula of reactants and products. 2. Write the molecular equation by writing the chemical formula for reactants and products. 3. Break the compounds into their ions and write the ionic equation for the reaction. 3. Refer to the table of solubility rules and decide whether any of the ion combinations is insoluble. 4. If a candidate is insoluble, that reaction will occur. 5. Remove the spectator ions and write the net ionic equation that summarized the reaction.

5 Learning Check: Precipitation Reactions Predict whether a reaction occurs when each of the following pairs of solutions are mixed. If a reaction does occur, write balanced molecular, total ionic, and net ionic equations, and identify the spectator ions. Solubility Rules For Ionic Compounds in Water Soluble Ionic Compounds 1. Salts of Group 1A and ammonium ion (NH 4+ ) 2. NO 3, CH 3 COO or C 2 H 3 O 2, ClO 4 3. Cl, Br I except those of Ag +, Pb 2+, Cu +, and Hg (a) sodium sulfate(aq) + strontium nitrate(aq) (b) ammonium perchlorate(aq) + sodium bromide(aq) (c) silver nitrate(aq) + sodium chromate(aq) Insoluble Ionic Compounds 1. All OH, except those of Group 1A and the larger members of Group 2A (beginning with Ca 2+ ). 2. CO 3 2 and PO 4 3 ) are insoluble, except those of Group 1A(1) and NH All S 2 except those of Group 1A, Group 2A and NH Convert names formulas, write a balanced equation showing intact compounds (not dissociated). Molecular Equation Na 2 SO 4 (aq) + Sr(NO 3 ) 2 (aq) 2Na + (aq) +SO 4 2 (aq)+ Sr 2+ (aq)+2no 3 (aq) SO 4 2 (aq)+ Sr 2+ (aq) 2NaNO 3 (aq) + SrSO 4 (s) 2. Dissociate the formula equation to an ionic equation showing cations and anions with charge and phase. Ionic Equation Net Ionic Equation SrSO 4 (s) 2Na + (aq) +2NO 3 (aq)+ SrSO 4 (s) 3. Cancel the spectator ions to obtain the net ionic equation Molecular Equation NH 4 ClO 4 (aq) + NaBr (aq) NH 4 Br (aq) + NaClO 4 (aq) Examing the solubility table shows that none of the possible combinations of ions would result in an insoluble precipitate (NH4ClO4 is soluble as is NH4Br, NaBr and NaClO4) In this case all ions are spectator ions. Ionic Equation NH 4+ + ClO 4 +(aq) + Na + + Br (aq) NH 4+ + Br (aq) + Na + + ClO 4 (aq) Does a precipitate form when silver nitrate is mixed with a solution of sodium chromate? Write the molecular, ionic and net ionic equations for the reaction. Write the net ionic equation for the reaction of silver nitrate with sodium chromate. 1. Write the balanced molecular equation. 2. Write the ionic equation showing the strong electrolytes completely dissociated into cations and anions. 3. Cancel the spectator ions on both sides of the ionic equation to obtain the net ionic equation

6 2. AcidBase Reactions There are 3classes of chemical reactions that occur in aqueous solution. 1. Precipitation Reaction an insoluble solid is formed from specific cationanion combinations. 2. AcidBase Reaction a protons donor substance reacts with a hydroxide donor substance forming a salt and water. The effects of acid rain on a statue of George Washington taken in 1935 (left) and 2001 (right) marble. 3. OxidationReduction Reactionelectron donor substances react with react with substances that accept electrons. Precipitation Involves Substances AcidBase Neutralization Involves Substances Oxidation Reduction Involves Substances Cations Anions H + ions OH ions Oxidation Reduction Combine to Form Insoluble Precipitate Predicted by Combine to Form Salt and H2O Which is the Loss of e Gain of e Which is called Solubility Rules Reducing Agent Oxidizing Agent Chemistry uses 3definitions of acids and bases: 1) Arrehenius 2) BronstedLowry 3) Lewis Chem 7 Chem 11 Name Acid Definition Base Definition Acids are substances that produce H + when dissolved in water. H2O Acid donates H + HA(g) ==> H + + A A is symbol for halide anion Arrhenius Substance that increases H + Substances that increase OH Bases are substances that produce OH when dissolved in water. Brønsted Lowry Lewis Substances that donate H + Electronpair acceptor Substances that accept H + Electronpair donar H2O MOH(s) ==> OH + M + Bases donate OH M + is symbol for metal cation Both are strong electrolytes because 100% dissociation!

7 Acid: substance that has a covalent H atom in its formula, and releases a proton H + when dissolved in water. Chemists symbolize the reactions of acids with water two confusing ways. Don t let it bug you. H2O HA(g) ==> H + (aq) + A (aq) H + representation H 3O + Representation HCl(g), HBr(g), HI (g), HNO3(l), H2SO4(l), HClO4(l) HCl(aq), HBr(aq), HI (aq), HNO3(aq), H2SO4(aq), HClO4(aq) H2O HCl(g) ===> H + + Cl HCl(g) + H2O => H3O + + Cl Base: a substance that contains OH in its formula, and releases hydroxide ions (OH ) when dissolved in water. H2O MOH(s) ==> OH + M + NaOH(s), KOH (s), LiOH (s), Mg(OH)2(s), Ca(OH)2(s) NaOH(aq), KOH (aq), LiOH (aq), Mg(OH)2(aq), Ca(OH)2(aq) H2O HNO3(l) ===> H + + NO3 H2O HA(g) ===> H + + A Generalized Acid HA HNO3(g) + H2O => H3O + + NO3 HA(g) + H2O => H3O + + A Generalized Acid HA Acids and bases are classified as either strong or weak. We use arrows to symbolize the difference. Strong acids and strong bases completely dissociate or completely ionize in water: 100% ionized = strong electrolyte = ==> arrow Strong Acid H2SO4(aq) ==> 2H + (aq) + SO4 2 (aq) HA H + + A HA(g) + H 2 O(l) H 3 O + (aq) + A (aq) Strong Base KOH(aq) ==> OH (aq) + K + (aq) Before dissociation After dissociation A H3O + <20% ionized = weak electrolyte = <==> arrow Weak Acid Weak Base HNO2(aq) <==> H + (aq) + NO2 (aq) NH4OH(aq) <==> OH (aq) + NH4 + (aq) HA HA H3O + A Strong acids and strong bases dissociate completely, conduct electricity and are strong electrolytes. We must memorize common strong acids and strong bases. All are strong electrolytes that dissociate completely in solution. Unidirectional arrow used to heavily productfavored Strong Acids Strong Bases HCl(aq) H + (aq) + Cl (aq) Strong Strong hydrochloric acid HCl hydrobromic acid HBr lithium hydroxide: sodium hydroxide: LiOH NaOH HNO 3 (aq) H + (aq) + NO 3 (aq) hydroiodic acid HI potassium hydroxide: KOH H2SO 4 (aq) H + (aq) + HSO 4 (aq) nitric acid HNO 3 sulfuric acid H 2 SO 4 perchloric acid HClO 4 calcium hydroxide Ca(OH) 2 strontium hydroxide Sr(OH) 2 NaOH(aq) Ca(OH)2(aq) Na + (aq) + OH (aq) Ca 2+ (aq) + 2OH (aq) HX acids Oxide containing acids barium hydroxide Ba(OH) 2 Group I and II Hydroxides

8 Remember anything not strong is weak! Common Weak Acids and Their Anions Acid HF CH3COOH HCN HNO2 H2CO3 H2SO3 H3PO4 (COOH)2 Anion FCH3COOCNNO2CO32SO32PO43(COO)22 Anion Name fluoride ion acetate ion cyanide ion nitrite ion carbonate ion sulfite ion phosphate ion oxalate ion Weak acids and weak bases dissociate only to a slight extent in water Ka << 1. HA HA(aq) + H2O(l) reactantfavored Before dissociation After dissociation HA HA All not strong is weak! HNO2 (aq) <=> NO2 (aq) + H+(aq) CH3CO2H(aq) <=> CH3CO2 (aq) + H+ H2SO3 (aq) <=> HSO3 (aq) + H+(aq) HA HA HA HA HA H3O+ A HA Weak acids and weak bases do not dissociate or ionize to a large extent in solution, and so are weak electrolytes (small value of Ka). H+ + AH3O+(aq) + A(aq) Acids can have one, two or three acidic protons depending on their structure. Monoprotic acidsonly one H+ available HCl H+ + ClHNO3 H+ + NO3CH3COOH H+ + CH3COO Strong electrolyte, strong acid Strong electrolyte, strong acid Weak electrolyte, weak acid Diprotic acidstwo acidic H+ available for reaction H2SO4 H+ + HSO4 Strong electrolyte, strong acid HSO4 H+ Weak electrolyte, weak acid + SO4 2 Triprotic acidsthree acidic H+ H3PO4 H2PO4HPO42 Acids and bases have distinct properties. Acids: Acid Acrid sour taste React with metals (Group I,II) to yield H2 gas Changes plant dye litmus from blue to red Common Acids, soft drinks Most antiperspirants, water treatment plants, paper Common Bases Bases: 50 Million MT NaOH/yr Bitter taste 3 million containers Slippery feel Changes plant dye litmus from red to blue React and neutralizes the effects of acids Weak electrolyte, weak acid Acids and bases are everywhere. React with carbonates and bicarbonates to produce CO2 gas 200 Million MT H2SO4 Base H+ + H2PO4H+ + HPO42H+ + PO43

9 Acids and bases are named systematically. 1. Naming Binary Acids (no oxygen atoms) Use the prefix hydro + name of nonmetal + ic acid. 2. Naming Ternary Acids or Oxyacids Ternary acids contain hydrogen and oxygen and one other (usually) nonmetallic element. sulfuric acid nitric acid HCl(aq) = hydrochloric acid HF(aq) = hydrofluoric acid HBr(aq) = hydrobromic acid HI(aq) = hydroiodic acid H2S(aq) = hydrosulfuric acid H2SO4 H2SO3 HNO3HNO2 sulfurous acid nitrous acid Determining the Molarity of H + Ions in Aqueous Solutions of Acids Nitric acid is a major chemical in the fertilizer and explosives industries. In aqueous solution, each molecule dissociates and the H becomes a solvated H + ion. What is the molarity of H + (aq) in 1.4M nitric acid? What is the H + molarity of 0.70 M H2SO4? Of H M H3PO4? Of 2.5 M NaOH? Determining the Molarity of H + Ions in Aqueous Solutions of Strong Acids or Bases Nitric acid is a major chemical in the fertilizer and explosives industries. In aqueous solution, each molecule dissociates and the H becomes a solvated H + ion. What is the molarity of H + (aq) in 1.4M nitric acid? One mole of H + (aq) is released per mole of nitric acid (HNO 3 ) HNO 3 (l) H 2 O H + (aq) + NO 3 (aq) Acids react with bases in a chemical reaction called a neutralization reaction forming a salt and water. NaOH (aq) + HCl (aq) ==> H2O + Na + + Cl base + acid ==> Water + salt 1.4M HNO 3 (aq) is 1.4M H + (aq). What is the molarity of H + in a 0.70 M H2SO4? What is the molarity of H + in a M H3PO4? A salt s cation comes from a parent base and anion comes from a parent acid! Chemists use general symbols to represent the neutralization reaction to form a salt and water. NaOH (aq) + HCl (aq) ==> H2O + Na + + Cl base + acid ==> Water + salt Learning check: write the molecular, ionic and net equation for the neutralization between calcium hydroxide and sulfuric acid. 1. The molecular equation shows all reactants and products as undissociated compounds. Ca(OH)2 (aq) + H2SO4 (aq) ===> 2H2O + CaSO4 MOH (aq) + HX (aq) ==> H2O + M + + X Alkali Metal Cation Halide Anion 2. The total ionic equation Shows all soluble ionic substances dissociated into ions. Ca OH + 2H + + SO4 2 ===> 2H2O + Ca 2+ + SO4 2 Ca 2+ and SO 4 2 are spectator ionsthey just watch! 3. The net ionic equation Eliminate spectator ions and show actual chemical change H + (aq) + OH (aq) ===> H 2 O(l)

10 Learning check: Write molecular, ionic and net equations for the following acid base neutralization reactions. strontium hydroxide(aq) + perchloric acid(aq) barium hydroxide(aq) + sulfuric acid(aq) Nitric acid(aq) + barium hydroxide(aq) Acetic acid(aq) + potassium hydroxide Writing Ionic Equations for AcidBase Reactions strontium hydroxide(aq) + perchloric acid(aq) (a) Sr(OH) 2 (aq)+2hclo 4 (aq) (b) Sr 2+ (aq)+2oh (aq)+ 2H + (aq)+2clo 4 (aq) (c) 2OH (aq)+ 2H + (aq) barium hydroxide(aq) + sulfuric acid(aq) (b) Ba(OH) 2 (aq) + H 2 SO 4 (aq) 2H 2 O(l)+Sr(ClO 4 ) 2 (aq) 2H 2 O(l)+Sr 2+ (aq)+2clo 4 (aq) 2H 2 O(l) 2H 2 O(l) + BaSO 4 (aq) Ba 2+ (aq) + 2OH (aq)+ 2H + (aq)+ SO 2 4 (aq) 2H 2 O(l)+Ba 2+ (aq)+so 2 4 (aq) (c) 2OH (aq)+ 2H + (aq) 2H 2 O(l) Weak acids dissociate to a very small extent and this fact is reflected in their equations as well using a doublearrow (<==>). Molecular equation NaOH(aq) + CH 3 COOH(aq) Total ionic equation Na + (aq)+ OH (aq) + CH 3 COOH(aq) CH 3 COONa(aq) + H 2 O Weak acid is not dissociated! CH 3 COO (aq) + Na + (aq) + H 2 O(l) Learning Check: Acids and Bases 1. Name and characterize each as a base, acid or salt? HF(g), HI(aq), LiOH(aq), Mg(OH)2, Na2SO4 CH3COONH4 2. Name and classify the following as strong, weak acid or base? HClO4, Sr(OH)2, HClO2, NH3(g), H3PO4(aq), H2SO4(aq), HNO3 (aq) 3. Write the M/I/NI equation for the reactions of a) hydrochloric acid with calcium hydroxide and b) phosphoric acid with sodium hydroxide Net ionic equation OH (aq) + CH 3 COOH(aq) CH 3 COO (aq) + H 2 O(l) 4. What acids and bases were reacted to form the following salts? Show using balanced equations. 1) NaNO2 2) CaSO4 3) Mg(PO4)2 Learning check: 1. Identify the following as a strong or weak acid or base a salt. If a salt what is the parent acid and base? HF, HI, LiOH, Ca(OH)2, Na2SO4 CH3COO, NH Classify the following as strong, weak acid or base? HClO4, Sr(OH)2, HClO2, NH3, H3PO4, H2SO4 3. What is the correct formula of the salt formed in the neutralization reaction of hydrochloric acid with calcium hydroxide? Name the type of reaction and write M/I/NI ionic equations for the following: 1. hydrochloric acid and potassium hydroxide 2. sodium carbonate and hydrochloric acid 3. aluminum nitrate and sodium phosphate 4. potassium chloride + iron(ii) nitrate 5. Iron sulfide and hydrochloric acid 6. acetic acid and magnesium hydroxide 7. sodium sulfite and hydrochloric acid 4. What is the chemical formula of the salt produced by the neutralization of sodium hydroxide with sulfurous acid?

11 Reactions between acids and carbonate containing compounds yield gaseous CO2 and H2O + salt. Molecular equation NaHCO 3 (aq) + CH 3 COOH(aq) CH 3 COONa(aq) + CO 2 (g) + H 2 O(l) A titration is the lab technique in which a solution of known molarity is added to another solution of unknown concentration until the chemical reaction between the two solutions is stoichiometric. Total ionic equation Na + (aq)+ HCO 3 (aq) + CH 3 COOH(aq) CH 3 COO (aq) + Na + (aq) + CO 2 (g) + H 2 O(l) Net ionic equation HCO 3 (aq) + CH 3 COOH(aq) CH 3 COO (aq) + CO 2 (g) + H 2 O(l) Start of titration Equivalence Point Slight excess of base There are many different types of indicator dyes used in titrations. 1. Pipet 20.0 ml of an acid solution of unknown molarity 2. Titrate with a solution of known [C] 3. When just the right volume of titrant is added, the solution changes color = end pt Buzzwords To Understand & Know Equivalence point: The point in the titration at which both acid and base have been consumed (it s stoichiometric). Neither acid nor base is present in excess and nothing left over. End point: The point at which the ph indicator changes color. It is not identical to the equivalence pointbut close enough! Titrant: The solution of known concentration (often standardized) and added to the solution of unknown concentration. Titration Curve: A graph or plot of Solution ph vs. volume of titrant added to the unknown solution. You perform an acidbase titration to standardize an HCl solution by placing ml of HCl in a flask with a few drops of indicator solution. You load M NaOH into the buret, and you record the initial reading as 0.55 ml. At the end point, the buret reading is ml. What is the concentration of the initial HCl solution? 1. Write and balance equation 2. Use volumes, molarities, and stoich to answer problem. 3. Watch sig figs.

12 You perform an acidbase titration to standardize an HCl solution by placing ml of HCl in a flask with a few drops of indicator solution. You load M NaOH into the buret, and you record the initial reading as 0.55 ml. At the end point, the buret reading is ml. What is the concentration of the HCl solution? NaOH(aq) + HCl(aq) mol HCl = ( ) ml X mol HCl = 5.078x10 3 [HCl] = 5.078x10 3 mol HCl L 1L 10 3 ml NaCl(aq) + H 2 O(l) X mol NaOH 1 L NaOH = M HCl What volume of a M NaOH solution is required to titrate ml of a 4.50 M H 2 SO 4 solution? H 2 SO 4 + 2NaOH 2H 2 O + Na 2 SO 4 M rx M volume acid moles acid moles base volume base acid coef. base 4.50 mol H 2 SO 4 2 mol NaOH 1000 ml soln ml x x x = 158 ml 1000 ml soln 1 mol H 2 SO mol NaOH How many ml of 0.300M HCl are required to titrate (neutralize) 6.25 x 10 3 mol NaOH in this acidbase neutralization reaction? How many ml of 0.300M HCl are required to neutralize 6.25 x 10 3 mol NaOH in this acidbase neutralization reaction? 1. Write a balanced equation NaOH + HCl!" NaCl + H 2 O 2. Use stoichiometry and definition of molarity to solve You perform an acidbase titration to standardize an HCl solution by placing ml of HCl in a flask with a few drops of indicator solution. You load M NaOH into the buret, and you record the initial reading as 0.55 ml. At the end point, the buret reading is ml. What is the concentration of the HCl solution? Involves simultaneous Oxidation Loss of e Reducing Agent Increases oxidation number Oxidation Reduction Which is the Reduction Which is called Which Gain of e Oxidizing Agent Decreases oxidation number

13 Redox reactions trends can be understood by applying the properties of elements in the periodic table. An oxidationreduction reaction occurs when there is a net movement of electrons from one reactant to another. One reactant gives the other receives! Metals tend to loose electrons (oxidized) forming cations Metals are reducing agents Nonmetals tend to gain electrons (be reduced) forming anions which are oxidizing agents. 2Mg(s) + O 2 (g) Oxidizing Perspective Mg loses electron(s) Mg(s) is oxidized Mg is the reducing agent Mg increases its oxidation number Oxidized 2MgO (s) Reduced Reduction Perspective O gains electron(s) O is reduced O is the oxidizing agent O decreases its oxidation number We learn rules to figure out what is oxidized and what is reduced! Oxidation and reduction occur simultaneously. 2Mg(s) + O 2 (g) 2MgO (s) Oxidized Reduced 2Mg 2Mg e Oxidation halfreaction (lose e ) O 2 + 4e 2O 2 Reduction halfreaction (gain e ) 2Mg + O 2 + 4e 2Mg O 2 + 4e Ionic Equation 2Mg + O 2 2MgO Net Equation A set of rules called assigning oxidation numbers is used to identify what species is being oxidized and reduced in a chemical reaction. 1. The oxidation state of a free elements in atomic or molecular for are assigned an oxidation number of 0. Na, Be, K, Pb, H 2, O 2, P 4 = 0 2. The oxidation number of a an monoatomic ion is equal to the charge on the ion (Group Number Rule) Li +, Li = +1; Fe 3+, Fe = +3; O 2, O = 2 3. The sum of the oxidation numbers of all the atoms in a molecule or ion is equal to the charge on the molecule or ion. H2O (H = +1 and O = 2) Sum = 0 4. Fluorine has an oxidation state of 1 in its compounds 5. The oxidation number of hydrogen is +1 except when it is bonded to metals in binary compounds. In these cases, its oxidation number is The oxidation number of oxygen is usually 2. In H 2 O 2 and O 2 2 it is 1 (exceptions). 7. When there is a conflict the lower numbered rules take priority over the higher numbered rules. Oxidation numbers of all the elements in HCO 3? HCO 3 O = 2 H = +1 3x(2) + 1 +? = 1 C = +4 Quick Table Assigning Oxidation Numbers 1. The oxidation state of any free element = 0 2. In monatomic ions, the oxidation number is equal to the charge on the ion (Use Group Number) 3. The sum of the oxidation numbers of all the atoms in a molecule or ion is equal to the charge on the molecule or ion. 4. Fluorine has an oxidation state of 1 in its compounds 5. The oxidation number of hydrogen is +1 except when it is bonded to metals in binary compounds ( 1). 6. The oxidation number of oxygen is usually 2. In H 2 O 2 and O 2 2 it is 1 (exceptions). 7. When there is a conflict the lower numbered rules take priority over the higher numbered rules.

14 Assign oxidation numbers to the following ions and the elements within a compound. HCO 3 IF 7 O = 2 H = +1 F = 1 3x(2) + 1 +? = 1 C = +4 7x(1) +? = 0 I = +7 K 2 Cr 2 O 7 O = 2 K = +1 7x(2) + 2x(+1) + 2x(?) = 0 Cr = +6 NaIO 3 Na = +1 O = 2 3x(2) + 1 +? = 0 I = +5 Assign oxidation numbers to the following ions and the elements within a compound. 1.! Na2SO3! 1.!Na +1! 2.!S +4! 3.!O 2 2.! PF3! 1.!P +3! 2.!F 1 3.! CrO 3 2! 1.!Cr +4! 2.!O 2 4.! Cr 2O 7 2! 1.!Cr +6! 2.!O 2 5.! (NH 4) 3PO 4! 1.!N 3! 2.!H +1! 3.!P +5! 4.!O 2 Elements that make up compounds can take on a range of oxidation states. Example: N Chemists use oxidation rules to identify what is being oxidized and what is being reduced in a chemical reaction. The reactions below occur as written. Use the rules for assigning oxidation numbers to the oxidizing and reducing agents in the reactions below? (increase in ON = reduced, decrease = oxidized) CH4 + 2O2 CO2 + 2H 2 O HgO 2Hg + O Zn(s) + Cu 2+ (aq)! Zn 2+ (aq) + Cu(s) Identify the oxidizing agent and reducing agent in each of the following: Zn (s) + CuSO 4 (aq) ZnSO 4 (aq) + Cu (s) 1. Assign oxidation numbers 2. Identify the species being oxidized (reactant that increases in oxidation number) and the specie being reduced (reactant that decrease in oxidation number). 3. The species that is oxidized is the reducing agent, the species that is reduced is called the oxidizing agent. Identify the oxidizing agent and reducing agent in each of the following: Zn (s) + CuSO 4 (aq) ZnSO 4 (aq) + Cu (s) Oxidized Reduced Zn Zn e Zn is oxidized Zn is the reducing agent Cu e Cu Cu 2+ is reduced Cu 2+ is the oxidizing agent

15 Copper wire reacts with silver nitrate to form silver metal and soluble cupric nitrate. What is the oxidizing agent and the reducing agent in this reaction? Copper wire reacts with silver nitrate to form silver metal and soluble cupric nitrate. What is the oxidizing agent and the reducing agent in this reaction? (Translate words to equations) Cu (s) + 2AgNO 3 (aq) Cu(NO 3 ) 2 (aq) + 2Ag (s) Oxidized Reduced Cu is oxidized Cu is the reducing agent Ag + is reduced Ag + is the oxidizing agent Identify the oxidizing agent and reducing agent in each of the following: Identify the oxidizing agent and reducing agent in each of the following: (a) 2Al(s) + 3H 2 SO 4 (aq) (b) PbO(s) + CO(g) (c) 2H 2 (g) + O 2 (g) Al 2 (SO 4 ) 3 (aq) + 3H 2 (g) Pb(s) + CO 2 (g) 2H 2 O(g) (b) PbO(s) + CO(g) Pb(s) + CO 2 (g) The O.N. of C increases; it is oxidized; CO is the reducing agent. The O.N. of Pb decreases; it is reduced; PbO is the oxidizing agent. 1. Know the rules for assigning oxidation numbers 2. Assign oxidation numbers to all reactants and products 3. Increasing ON = Reactant Oxidized (reducing agent) Decreasing ON = Reactant Reduced (oxidizing agent) (c) 2H 2 (g) + O 2 (g) 2H 2 O(g) The O.N. of H increases; it is oxidized; it is the reducing agent. The O.N. of O decreases; it is reduced; it is the oxidizing agent. Using oxidation numbers 1. When an element combines with O2 the element has been oxidized and O2 reduced. 2. When an element combines with a halogen the element is oxidized and the halogen reduced. 3. When a metal combines with something, the metal is been oxidized. Whatever the metal combined with, is reduced. There are 5Main Types of Redox Reactions 1. Combination A + B => AB produce fewer products than reactants. i. Element + Element => Compound ii. Element + Compound => New Compound iii. Compound + Compound => New Compound 2. Decomposition AB => A + B produce more products than reactants. 3. Displacement AB + C => AC + B methathesis (switching places) 4. Disproportionationcommon element or element in compound is both oxidized and reduced 5. Combustion CH4 + O2 => CO2 + H2O

16 1. Combination Reactions: A + B => C Combination reactions produce fewer products than reactants. 1. Metal + Nonmetal Binary Ionic Compound 2Na(s)+ Cl 2 2NaCl 2. NonMetal + Nonmetal Binary Molec Cmpd S + O 2 SO 2 3. Compound + Element New Compound PCl3+ Cl 2 4. Compound + Compound New Compound CaO(s)+ H 2 O PCl5 Ca(OH)2 2. Decomposition Reactions produces more products than reactants through the action of heat (thermal decomposition) or electricity (electrolytic decomposition). 1. Compound Element A + Element B electricity 2H2O 2H2 + O 2 heat 2HgO 2Hg + O 2 2. Compound Compound A + Element B heat KClO 3 2KCl + 3O 2 3. Displacement Reactions A more reactive metal will displace a less reactive metal or H2 within their compounds to form the oxidized metal and the reduced form of the less active metal or H2. A + BC AC + B Ca(s) + 2H 2 O(aq)! Ca(OH) 2 + H 2 H2 Displacement Mg(s) + 2HCl(aq)! MgCl 2 (aq) + H 2 (g) TiCl 4(aq) + 2Mg(s) Ti + 2MgCl Zn(s) + CuSO4 (aq) => ZnSO4 (aq) + Cu (s) double Methathesis = displacement 0 Metal Displacement Activity Series For Metals Pure metals higher in the series will be oxidized by a dissolved elemental ion below it in the series which is reduced. H2 DISPLACEMENT Metals above H2 react with acids (H + ), or H2O to displace H2(g). H + reduced to H2. Al(s) + H + => Al 3+ + H2(g) Ag(s) + H + => No reaction METAL DISPLACEMENT More active metal + Salt of less active metal => Reduced Less active metal + Salt of more active metal. Al(s) + Fe 3+ (aq) => Al 3+ + Fe(s) Zn(s) + CuSO4 => ZnSO4 + Cu(s) Metal Displacement (Activity Series) Zn Rod CuSO4 Solid Zn rod dunked A solid Zn rod into dunked a solution into a of CuSO4. solution of copper sulfate (CuSO4) CuSO4 Use activity series: Zn metal is higher than Cu in the series. It will displace Cu 2+ from solution. Zn(s) + Cu 2+ (aq)! Zn 2+ (aq) + Cu(s) The Zn is oxidized and the Cu 2+ is reduced to solid Cu(s) 1. Fe(s) + 3Cu 2+ => 2Fe 3+ + Cu(s) 2. Sn(s) + Ca 2+ => Sn 2+ + Ca(s) 3. Au(s) + 3AgNO 3 (aq) => Au(NO 3 ) Ag(s) Predict whether the following reactions will occur as written Al(s) + Cr 2 O 3 (s)! Al 2 O 3 (s) + 2 Cr(s) 5. Pt(s) + 4 HCl(aq)! PtCl 4 (aq) + 2 H 2 (g) More active metal + Salt of less active metal => Reduced Less active metal + Salt of more active metal.

17 A more active halogen will displace a less active (heavier) halide from their binary salts. Strength as oxidizing agent F 2 Cl 2 Br Cl2 + 2NaI ===> I2 + 2NaCl Br2 + 2NaI ===> I2 + 2NaBr I2(s) + 2F ===> No reaction Cl2 and Br2 both are higher and will displace I from its compounds, forming I2 and the chloride and bromide salt respectively. A less active halogen will not displace a more active halide from their salts. 4. Disproportionation Reaction A single element in a compound is both oxidized and reduced (both oxidizing and reducing agent in one) H2O2 + 2OH Cl 2 + 2OH 2H 2 O + O ClO + Cl + H 2 O I 2 Halogen + Salt of Less Active Halide gives Less Active Halide + salt of more active halide. 5. Combustion Reaction In a combustion reaction a fuel or typically a hydrocarbon reacts with oxygen forming carbon dioxide and water. The oxygen is increases in O.N. and is therefore oxidized (reducing agent) CH4 + 2O2 CO2 + 2H 2 O Identifying the Type of Redox Reaction Classify each of the following redox reactions as a combination, decomposition, or displacement reaction, write a balanced molecular equation for each, as well as total and net ionic equations for part (c), and identify the oxidizing and reducing agents: (a) magnesium(s) + nitrogen(g) (b) hydrogen peroxide(l) (c) aluminum(s) + lead(ii) nitrate(aq) magnesium nitride (aq) water(l) + oxygen gas aluminum nitrate(aq) + lead(s) 2C4H O2 8CO2 + 10H 2 O Identifying the Type of Redox Reaction (a) Combination 3Mg(s) + N 2 (g) Mg 3 N 2 (aq) Mg is the reducing agent; N 2 is the oxidizing agent H 2 O 2 (l) H 2 O(l) + 1/2 O 2 (g) or 2 H 2 O 2 (l) 2 H 2 O(l) + O 2 (g) H 2 O 2 is the oxidizing and reducing agent (b)decomposition (c)displacement Al(s) + Pb(NO 3 ) 2( aq) Al(NO 3 ) 3 (aq) + Pb(s) 2Al(s) + 3Pb(NO 3 ) 2( aq) 2Al(NO 3 ) 3 (aq) + 3Pb(s) Pb(NO 3 ) 2 is the oxidizing and Al is the reducing agent.