Chapter 4 Types of Chemical Reactions and Solution Stoichiometry

Chapter 4: Types of Chemical Reactions and Solution Stoichiometry 4.1 Water, the Common Solvent 4.2 The Nature of Aqueous Solutions: Strong and Weak Electrolytes 4.3 The Composition of Solutions 4.4 Types of Chemical Reactions 4.5 Precipitation Reactions 4.6 Describing Reactions in Solution 4.7 Selective Precipitation 4.8 Stoichiometry of Precipitation Reactions 4.9 Acid-Base Reactions 4.10 Oxidation-Reduction Reactions 4.11 Balancing Oxidation-Reduction Equations 4.12 Simple Oxidation-Reduction Titrations

Figure 4.1: A space-filling model of the water molecule.

The Water molecule has two sets of electrons and can hydrogen bond with a total of four other water molecules. H H

When Water freezes, it sets up a structure with Hexagons formed by water molecules hydrogen bonding with other water molecules. These hexagons have space inside, and make the solid form of water less dense than the liquid form. Because of this, ice floats on liquid water. Water is unique in this regard, as in all other the solid form is more dense than the liquid form. This hydrogen bonding is also very important in Biomolecules such as DNA and Proteins.

If one looks at the boiling points of the hydrides of many elements, water, ammonia, and hydrogen fluoride have uniquely High boiling points. For example water is projected to have a Boiling point of only about -80 o C in stead of 100 o C!

Figure 4.2: Polar water molecules interact with the positive and negative ions of a salt, assisting with the dissolving process.

Figure 4.3(a) The ethanol molecule contains a polar O-H bond similar to those in the water molecule. (b) The polar water molecule interacts strongly with the polar O-H bond in ethanol.

The Role of Water as a Solvent: The solubility of Ionic Compounds Electrical conductivity - The flow of electrical current in a solution is a measure of the solubility of ionic compounds or a measurement of the presence of ions in solution. Electrolyte - A substance that conducts a current when dissolved in water. Soluble ionic compound dissociate completely and may conduct a large current, and are called Strong Electrolytes. NaCl (s) + H 2 O (l) Na + (aq) + Cl - (aq) When Sodium Chloride dissolves into water the ions become solvated, and are surrounded by water molecules. These ions are called “aqueous” and are free to move through out the solution, and are conducting electricity, or helping electrons to move through out the solution

Electrical Conductivity of Ionic Solutions

Figure 4.4: Electrical Conductivity

Figure 4.5: HCL (aq) is completely ionized.

Carbohydrates Molecules that contain carbon and water! C x H 2y O y CH 2 OH C C C C CC CC C O O O H H H H H H H H OH HO Sucrose C 12 H 22 O 11, C 12 (H 2 O) 11 a disaccharide

Hydrogen-Bonding of Sucrose in Water O-H CH 2 H-O H-C O H CH 2 O H H-O-C C C H H C C C O C C CH 2 -O-H H-O O-H H-O O-H H-O H O-H H O-H H O-H H O-H H O-H H O-H H O-H H O-H H O-H H H-O H H-O H H-O H O-H Some of the water molecules that can hydrogen bond to the sucrose molecule (shown in red), only 16 are shown here! Theoretically there can be 30 water molecules associated with one sucrose molecule, three on each –OH group, and two on each oxygen atom. This is why an organic molecule like sucrose are so soluble in water, all of the sugars do this extensive H-bonding.

Figure 4.6: An aqueous solution of sodium hydroxide.

Figure 4.7: Acetic acid (HC 2 H 3 O 2 ) exists in water mostly as undissociated molecules.

Figure 4.8: The reaction of NH 3 in water.

Molarity (Concentration of Solutions)= M M = = Moles of Solute Moles Liters of Solution L solute = material dissolved into the solvent In air, Nitrogen is the solvent and oxygen, carbon dioxide, etc. are the solutes. In sea water, Water is the solvent, and salt, magnesium chloride, etc. are the solutes. In brass, Copper is the solvent (90%), and Zinc is the solute(10%)

LIKE EXAMPLE 4.1 (P 97) Calculate the Molarity of a solution prepared by bubbling 3.68g of Gaseous ammonia into 75.7 ml of solution. Solution: Calculate the number of moles of ammonia: 3.68g NH 3 X = mol NH 3 1 mol NH g Change the volume of the solution into liters: 75.7 ml X = L 1 L 1000 mL Finally, we divide the number of moles of solute by the volume of the solution: Molarity = = ____________ M NH mol NH L

Preparing a Solution - I Prepare a solution of Sodium Phosphate by dissolving 3.95g of Sodium Phosphate into water and diluting it to ml or l ! What is the Molarity of the salt and each of the ions? Na 3 PO 4 (s) + H 2 O (solvent) = 3 Na + (aq) + PO 4 -3 (aq)

Preparing a Solution - II Mol wt of Na 3 PO 4 = g / mol 3.95 g / g/mol = mol Na 3 PO 4 dissolve and dilute to ml M = mol Na 3 PO 4 / l = M Na 3 PO 4 for PO 4 -3 ions = ______________ M for Na + ions = 3 x M = ___________ M

Like Example 4.3 (P 98) An isotonic solution, one with the same ionic content as blood is about 0.14 M NaCl. Calculate the volume of blood that would contain 2.5 mg Of NaCl? Find the moles in 1.0 mg NaCl: 2.5 mg NaCl x x = 4.28 x mol NaCl 1 g NaCl 1000 mg NaCl 1 mol NaCl 58.45g NaCl What volume of 0.14 M NaCl that would contain the amount of NaCl (4.28 x mol NaCl): V x = 4.28 x mol NaCl 0.14 M NaCl L solution Solving for Volume gives: V = = ______________________ L 4.28 x mol NaCl 0.14 mol NaCl L solution Or _________ ml of Blood!

Figure 4.9: Steps involved in the preparation of a standard solution.

Like Example 4.4 (P 98) A Chemist must prepare a 1.00 L of a M solution of Ammonium Carbonate, what mass of (NH 4 ) 2 CO 3 must be weighed out to prepare this solution? First, determine the moles of Ammonium Carbonate required: 1.00 L x = M (NH 4 ) 2 CO M (NH 4 ) 2 CO 3 L solution This amount can be converted to grams by using the molar mass: M (NH 4 ) 2 CO 3 x = g (NH 4 ) 2 CO g (NH 4 ) 2 CO 3 mol (NH 4 ) 2 CO 3 Or, to make 1.00L of solution, one must weigh out 35.3 g of (NH 4 ) 2 CO 3, put this into a 1.00 L volumetric flask, and add water to the mark on the flask.

Make a Solution of Potassium Permanganate Potassium Permanganate is KMnO 4 and has a molecular mass of g / mole Problem: Prepare a solution by dissolving 1.58 grams of KMnO 4 into sufficient water to make ml of solution g KMnO 4 x = moles KMnO 4 1 mole KMnO g KMnO 4 Molarity = = ______________ M moles KMnO liters Molarity of K + ion = [K + ] ion = [MnO 4 - ] ion = _____________ M

Figure 4.10: (a) A measuring pipette (b) A volumetric pipette.

Figure 4.11: (a) A measuring pipette (b) Water is added to the flask. (c) The resulting solution is 1 M acetic acid.

Dilution of Solutions Take ml of the M KMnO 4 Dilute the ml to l - What is the resulting Molarity of the diluted solution? # moles = Vol x M l x M = Moles Mol / 1.00 l = _______________ M

Figure 4.13: Reactant solutions: (a) Ba(NO 3 ) 3 (aq)

Figure 4.13: Reactant solutions: (b) K 2 CrO 4 (aq).

Figure 4.12: When yellow aqueous potassium chromate is added to a colorless barium nitrate solution, yellow barium chromate precipitates.

Figure 4.14: Reaction of K 2 CrO 4 (aq) and Ba(NO 3 ) 2 (aq).

Figure 4.15: Precipitation of silver chloride by mixing solutions of silver nitrate and potassium chloride.

Figure 4.16: Photos and molecular-level representations illustrating the reaction of KCL(aq) with AgNO3(aq) to form AgCl(s).

Precipitation of silver chromate by adding potassium chromate to a solution of silver nitrate. K 2 CrO 4 (aq) + 2 AgNO 3 (aq) Ag 2 CrO 4 (s) + 2 KNO 3 (aq)

Table 4.1 (P105) Simple Rules for Solubility of Salts in Water 1.Most nitrate (NO 3 - ) salts are soluble. 2.Most salts of Na +, K +, and NH 4 + are soluble. 3.Most chloride salts are soluble. Notable exceptions are AgCl, PbCl 2, and Hg 2 Cl 2. 4.Most sulfate salts are soluble. Notable exceptions are BaSO 4, PbSO 4, and CaSO 4. 5.Most hydroxide salts are only slightly soluble. The important soluble hydroxides are NaOH, KOH, and Ca(OH) 2 (marginally soluble). 6.Most sulfide (S 2- ), carbonate (CO 3 2- ), and phosphate (PO 4 3- ) salts are only slightly soluble.

The Solubility of Ionic Compounds in Water The solubility of Ionic Compounds in water depends upon the relative strengths of the electrostatic forces between ions in the ionic compound and the attractive forces between the ions and water molecules in the solvent. There is a tremendous range in the solubility of ionic compounds in water! The solubility of so called “insoluble” compounds may be several orders of magnitude less than ones that are called “soluble” in water, for example: Solubility of NaCl in water at 20 o C = 365 g/L Solubility of MgCl 2 in water at 20 o C = g/L Solubility of AlCl 3 in water at 20 o C = 699 g/L Solubility of PbCl 2 in water at 20 o C = 9.9 g/L Solubility of AgCl in water at 20 o C = g/L Solubility of CuCl in water at 20 o C = g/L

The Solubility of Covalent Compounds in Water The covalent compounds that are very soluble in water are the ones with -OH group in them and are called “Polar” and can have strong polar (electrostatic)interactions with water. Examples are compound such as table sugar, sucrose (C 12 H 22 O 11 ); beverage alcohol, ethanol (C 2 H 5 -OH); and ethylene glycol (C 2 H 6 O 2 ) in antifreeze. CH H H O H Methanol = Methyl Alcohol Other covalent compounds that do not contain a polar center, or the -OH group are considered “Non-Polar”, and have little or no interactions with water molecules. Examples are the hydrocarbons in Gasoline and Oil. This leads to the obvious problems in Oil spills, where the oil will not mix with the water and forms a layer on the surface! Octane = C 8 H 18 and / or Benzene = C 6 H 6

When a solution of Na 2 SO 4 (aq) is added to a solution of Pb(NO 3 ) 2, the white solid PbSO 4 (s) forms.

Determining Moles of Ions in Aqueous Solutions of Ionic Compounds - I Problem: How many moles of each ion are in each of the following: a) 4.0 moles of sodium carbonate dissolved in water b) 46.5 g of rubidium fluoride dissolved in water c) 5.14 x formula units of iron (III) chloride dissolved in water d) 75.0 ml of 0.56M scandium bromide dissolved in water e) 7.8 moles of ammonium sulfate dissolved in water a) Na 2 CO 3 (s) 2 Na + (aq) + CO 3 -2 (aq) moles of Na + = 4.0 moles Na 2 CO 3 x = 8.0 moles Na + and 4.0 moles of CO 3 -2 are present H2OH2O 2 mol Na + 1 mol Na 2 CO 3

Determining Moles of Ions in Aqueous Solutions of Ionic Compounds - II b) RbF (s) Rb + (aq) + F - (aq) H2OH2O moles of RbF = 46.5 g RbF x = moles RbF 1 mol RbF g RbF thus, mol Rb + and mol F - are present c) FeCl 3 (s) Fe +3 (aq) + 3 Cl - (aq) H2OH2O moles of FeCl 3 = 9.32 x formula units 1 mol FeCl x formula units FeCl 3 x = mol FeCl 3 moles of Cl - = mol FeCl 3 x = _________ mol Cl - 3 mol Cl - 1 mol FeCl 3 and ____________ mol Fe +3 are also present.

Determining Moles of Ions in Aqueous Solutions of Ionic Compounds - III d) ScBr 3 (s) Sc +3 (aq) + 3 Br - (aq) H2OH2O Converting from volume to moles: Moles of ScBr 3 = 75.0 ml x x = mol ScBr 3 1 L 10 3 ml 0.56 mol ScBr 3 1 L Moles of Br - = mol ScBr 3 x = mol Br - 3 mol Br - 1 mol ScBr mol Sc +3 are also present e) (NH 4 ) 2 SO 4 (s) 2 NH 4 + (aq) + SO (aq) H2OH2O Moles of NH 4 + = 7.8 moles (NH 4 ) 2 SO 4 x = ____ mol NH mol NH mol(NH 4 ) 2 SO 4 and ______ mol SO are also present.

Solid Fe(OH) 3 forms when aqueous KOH and Fe(NO 3 ) 3 are mixed.

Precipitation Reactions: Will a Precipitate form? If we add a solution containing Potassium Chloride to a solution containing Ammonium Nitrate, will we get a precipitate? KCl (aq) + NH 4 NO 3 (aq) = K + (aq) + Cl - (aq) + NH 4 + (aq) + NO 3 - (aq) By exchanging cations and anions we see that we could have Potassium Chloride and Ammonium Nitrate, or Potassium Nitrate and Ammonium Chloride. In looking at the solubility table it shows all possible products as soluble, so there is no net reaction! KCl (aq) + NH 4 NO 3 (aq) = No. Reaction! If we mix a solution of Sodium sulfate with a solution of Barium Nitrate, will we get a precipitate? From the solubility table it shows that Barium Sulfate is insoluble, therefore we will get a precipitate! Na 2 SO 4 (aq) + Ba(NO 3 ) 2 (aq) BaSO 4 (s) + 2 NaNO 3 (aq)

Precipitation Reactions: A solid product is formed When ever two aqueous solutions are mixed, there is the possibility of forming an insoluble compound. Let us look at some examples to see how we can predict the result of adding two different solutions together. Pb(NO 3 ) 2 (aq) + NaI (aq) Pb +2 (aq) + 2 NO 3 - (aq) + Na + (aq) + I - (aq) When we add These two solutions together, the ions can combine in the way they came into the solution, or they can exchange partners. In this case we could have Lead Nitrate and Sodium Iodide, or Lead Iodide and Sodium Nitrate formed, to determine which will happen we must look at the solubility table(P 141) to determine what could form. The table indicates that Lead Iodide will be insoluble, so a precipitate will form! Pb(NO 3 ) 2 (aq) + 2 NaI (aq) PbI 2 (s) + 2 NaNO 3 (aq)

Predicting Whether a Precipitation Reaction Occurs; Writing Equations: a) Calcium Nitrate and Sodium Sulfate solutions are added together. Ca(NO 3 ) 2 (aq) + Na 2 SO 4 (aq) CaSO 4 (s) +2 NaNO 3 (aq) Ca 2+ (aq) +2 NO 3 - (aq) + 2 Na + (aq) + SO 4 -2 (aq) CaSO 4 (s) + 2 Na + (aq + ) 2 NO 3 - (aq) Molecular Equation Total Ionic Equation Net Ionic Equation Ca 2+ (aq) + SO 4 -2 (aq) CaSO 4 (s) Spectator Ions are Na + and NO 3 - b) Ammonium Sulfate and Magnesium Chloride are added together. In exchanging ions, no precipitates will be formed, so there will be no Chemical reactions occurring! All ions are spectator ions!

Figure 4.17: Selective precipitation of Ag +, Ba 2+, and Fe 3+ ions.

Species present, Balanced net ionic equation.

Like Example 4.7 (P 110) Calculate the mass of solid sodium iodide that must be added to 2.50 L of a M lead nitrate solution to precipitate all of the lead as PbI 2 (s) ! The chemical equation for the reaction is: Pb(NO 3 ) 2 (aq) + 2 NaI (aq) PbI 2 (s) + 2 NaNO 3 (aq) Two times as much sodium iodide is needed to precipitate the Lead ions. The number of moles of sodium iodide needed is: 2.50 L x x = mol I Mol Pb L soln. 2 mol I - 1 mol Pb 2+ The mass of sodium iodide is: mol I - x x = __________ g NaI 1 mol NaI 1 mol I g NaI 1 mol NaI

Like Example 4.8 (P 111) When aqueous solutions of silver nitrate and sodium chloride are mixed, silver chloride is precipitated. What mass of silver chloride would be formed by the addition of ml to 3.17 M NaCl and 128 ml of 2.44 M silver nitrate? The stoichiometric relationship comes from the chemical equation: AgNO 3 (aq) + NaCl (aq) AgCl (s) + NaNO 3 (aq) There is a one to one relationship, therefore the number of moles are the same, but which is in the lowest quantity? V AgNO3 x M AgNO3 = L x 2.44 M = mol Ag + V NaCl x M NaCl = L x 3.17 M = mol Cl - Since the Chloride ion is smaller, it is limiting, and we use it to calculate the mass of AgCl, since we can only obtain mol of AgCl: Mass AgCl = mol x g AgCl/ mol = _________ g

Like Example 4.9 (P 111) What mass of Pb 2+ could by precipitated from a solution by the addition of L of M Sodium Iodide solution? Find the stoichiometric relationship from the chemical equation: Pb 2+ (aq) + 2 I - (aq) PbI 2 (s) It will take twice the iodide ion to precipitate the Lead ions: Moles I - = V NaI x M NaI = L x Moles = mol I - Liter Moles Lead ion = = mol Pb 2+ Mass of Lead = 207.2g Pb x moles = ____________ g Pb mol Pb mol I - 2 mol I - / mol Pb 2+

The gravimetric procedure.(P112) g CaC 2 O 4 H 2 O X = x mol CaC 2 O 4 H 2 O x mol Ca 2+ X = x g Ca 2+ Mass % Ca is: x 100% = ______________%. 1 mol CaC 2 O 4 H 2 O g CaC 2 O 4 H 2 O g Ca 2+ 1 mol Ca x g

Acids - A group of Covalent molecules which lose Hydrogen ions to water molecules in solution HI (g) + H 2 O (L) H + (aq) + I - (aq) When gaseous hydrogen Iodide dissolves in water, the attraction of the oxygen atom of the water molecule for the hydrogen atom in HI is greater that the attraction of the of the Iodide ion for the hydrogen atom, and it is lost to the water molecule to form an Hydronium ion and an Iodide ion in solution. We can write the Hydrogen atom in solution as either H + (aq) or as H 3 O + (aq) they mean the same thing in solution. The presence of a Hydrogen atom that is easily lost in solution is an “Acid” and is called an “acidic” solution. The water (H 2 O) could also be written above the arrow indicating that the solvent was water in which the HI was dissolved. HI (g) + H 2 O (L) H 3 O + (aq) + I - (aq) HI (g) H + (aq) + I - (aq) H2OH2O

Strong Acids and the Molarity of H + Ions in Aqueous Solutions of Acids Problem: In aqueous solutions, each molecule of sulfuric acid will loose two protons to yield two Hydronium ions, and one sulfate ion. What is the molarity of the sulfate and Hydronium ions in a solution prepared by dissolving 155g of concentrate sulfuric acid into sufficient water to produce 2.30 Liters of acid solution? Plan: Determine the number of moles of sulfuric acid, divide the moles by the volume to get the molarity of the acid and the sulfate ion. The hydronium ions concentration will be twice the acid molarity. Solution: Two moles of H + are released for every mole of acid: H 2 SO 4 (l) + 2 H 2 O (l) 2 H 3 O + (aq) + SO (aq) Molarity of H + = 2 x mol H + = 1.37 Molar in H + Moles H 2 SO 4 = = 1.58 moles H 2 SO g H 2 SO 4 1 mole H 2 SO g H 2 SO 4 x 1.58 mol SO l solution Molarity of SO = = Molar in SO 4 - 2

Species present, Balanced net ionic equation.

Acid - Base Reactions : Neutralization Rxns. An Acid is a substance that produces H + (H 3 O + ) ions when dissolved in water, and is a proton donor A Base is a substance that produces OH - ions when dissolved in water. the OH - ions react with the H + ions to produce water, H 2 O, and are therefore proton acceptors. Acids and Bases are electrolytes, and their strength is categorized in terms of their degree of dissociation in water to make hydronium or hydroxide ions. Strong acids and bases dissociate completely, and are strong electrolytes. Weak acids and bases dissociate weakly and are weak electrolytes. The generalized reaction between an Acid and a Base is: HX (aq) + MOH (aq) MX (aq) + H 2 O (L) Acid + Base = Salt + Water

Selected Acids and Bases Acids Bases Strong Hydrochloric, HCl Sodium hydroxide, NaOH Hydrobromic, HBr Potassium hydroxide, KOH Hydroiodoic, HI Calcium hydroxide, Ca(OH) 2 Nitric acid, HNO 3 Strontium hydroxide, Sr(OH) 2 Sulfuric acid, H 2 SO 4 Barium hydroxide, Ba(OH) 2 Perchloric acid, HClO 4 Weak Hydrofluoric, HF Ammonia, NH 3 Phosphoric acid, H 3 PO 4 Acetic acid, CH 3 COOH (or HC 2 H 3 O 2 )

Writing Balanced Equations for Neutralization Reactions - I Problem: Write balanced chemical reactions (molecular, total ionic, and net ionic) for the following Chemical reactions: a) Calcium Hydroxide(aq) and Hydroiodoic acid(aq) b) Lithium Hydroxide(aq) and Nitric acid(aq) c) Barium Hydroxide(aq) and Sulfuric acid(aq) Plan: These are all strong acids and bases, therefore they will make water and the corresponding salts. Solution: a) Ca(OH) 2 (aq) + 2HI (aq) CaI 2 (aq) + 2H 2 O (l) Ca 2+ (aq) + 2 OH - (aq) + 2 H + (aq) + 2 I - (aq) Ca 2+ (aq) + 2 I - (aq) + 2 H 2 O (l) 2 OH - (aq) + 2 H + (aq) 2 H 2 O (l)

Writing Balanced Equations for Neutralization Reactions - II b) LiOH (aq) + HNO 3 (aq) LiNO 3 (aq) + H 2 O (l) Li + (aq) + OH - (aq) + H + (aq) + NO 3 - (aq) Li + (aq) + NO 3 - (aq) + H 2 O (l) OH - (aq) + H + (aq) H 2 O (l) c) Ba(OH) 2 (aq) + H 2 SO 4 (aq) BaSO 4 (s) + 2 H 2 O (l) Ba 2+ (aq) + 2 OH - (aq) + 2 H + (aq) + SO 4 2- (aq) BaSO 4 (s) + 2 H 2 O (l)

Figure 4.18: The titration of an acid with a base.

Like Example 4.10 (P 115) What volume of M H 2 SO 4 is needed to neutralize ml of a M LiOH solution? Calculate the number of moles of base: V base x M base = L x M = mol LiOH From the balance equation find the moles of acid needed: 2 LiOH (aq) + H 2 SO 4 (aq) 2 H 2 O (l) + Li 2 SO 4 (aq) Since there are two protons per molecule, we will need half as much sulfuric acid as we have lithium hydroxide: or mol H 2 SO 4 Volume of acid: V acid = = = L H 2 SO 4 Moles acid M acid moles Mol L

A firefighter in a protective suit with an oxygen tank neutralizes an acid spill.

Finding the Concentration of Base from an Acid - Base Titration - I Problem: A titration is performed between Sodium Hydroxide and Potassium Hydrogenphthalate (KHP) to standardize the base solution, by placing mg of solid Potassium Hydrogenphthalate in a flask with a few drops of an indicator. A buret is filled with the base, and the initial buret reading is 0.55 ml; at the end of the titration the buret reading is ml. What is the concentration of the base? Plan: Use the molar mass of KHP (204.2 g/mol) to calculate the number of moles of the acid, from the balanced chemical equation, the reaction is equal molar, so we know the moles of base, and from the difference in the buret readings, we can calculate the molarity of the base. Solution: HKC 8 H 4 O 4 (aq) + OH - (aq) KC 8 H 4 O 4 - (aq) + H 2 O (aq)

Potassium Hydrogenphthalate KHC 8 H 4 O 4 C O OK+K+ C O H O K+K+ O C O O C O H+H+

Finding the Concentration of Base from an Acid - Base Titration - II moles KHP = x = mol KHP mg KHP g KHP 1 mol KHP 1.00 g 1000 mg Volume of base = Final buret reading - Initial buret reading = ml ml = ml of base one mole of acid = one mole of base; therefore moles of acid will yield moles of base in a volume of ml. molarity of base = = __________ moles per liter moles L molarity of base = ___________________ M

Like Example 4.12 (P117) A powered residue contains some ascorbic acid(Vitamin C, mol wt = 176g/mol) and the rest is a non acidic compound. If 10.0g of the powder is neutralized by ml of 1.5 M sodium hydroxide, a strong base, and the remaining base titrated with hydrochloric acid using ml of 1.80 M. What is the percentage of ascorbic acid? Mol acid = L x 1.80 Mol/L = mol HCl Mol base = L x 1.50 Mol/L = mol NaOH The difference between the base and acid will be the moles of ascorbic acid! Reacted base = – = mol Ascorbic acid Mass ascorbic acid = mol x 176g/mol = 1.95 g Ascorbic acid % ascorbic acid = x 100% = ____________% 1.95g 10.00g

Figure 4.19: Reaction of solid sodium and gaseous chlorine to form solid sodium chloride.

Highest and Lowest oxidation numbers of Chemically reactive main-group Elements 1 +1 H non-metals metalloids metals F Cl Br I At ONCB 1A +1 2A LiBe 3A A5A6A7A S Se Te Po P As Sb Bi Si Ge Sn Pb Al Ga In Tl NaMg KCa RbSr CsBa RaFr Period

Rn Xe Kr Ar NeFONCBBeLi HeH Period IA IIAIIIAIVAVAVIAVIIA VIIIA Main Group Elements Na K Rb Cs Mg Ca Sr Ba Al Ga In Tl Si Ge Sn Pb P As Sb Bi S Se Te Po Cl Br I At , ,+1 +4,+2 -1,-4 +4,-4 +4, ,+2, -4 +4,+2 +7,+5 +3,+1 +7,+5 +3,+1 +7,+5 +3,+1 +7,+5 +3, , ,-2 +6,+4 +2,-2 +6, , ,+4 +2, , , ,+3 -3 all from +5 -3

Transition Metals Possible Oxidation States IIIBIVBVBVIBVIIBIBIIB VIIIB ScTiVCrMnFeCoNiCuZn YZrNbMoTcRuRhPdAgCd HgAuPtIrOsReWTaHfLa +3 +4, , , , ,+2 +3,+2 +7,+6 +4, , , , , ,+5 +4,+3 +7, ,+5 +4,+3 +5, , , ,+6 +4,+3 +2

Determining the Oxidation Number of an Element in a Compound Problem: Determine the oxidation number (Ox. No.) of each element in the following compounds. a) Iron III Chloride b) Nitrogen Dioxide c) Sulfuric acid Plan: We apply the rules in Table 4.3, always making sure that the Ox. No. values in a compound add up to zero, and in a polyatomic ion, to the ion’s charge. Solution: a) FeCl 3 This compound is composed of monoatomic ions. The Ox. No. of Cl - is -1, for a total of -3. Therefore the Fe is +3. b) NO 2 The Ox. No. of oxygen is -2 for a total of -4. Since the Ox. No. in a compound must add up to zero, the Ox. No. of N is +4. c) H 2 SO 4 The Ox. No. of H is +1, so the SO 4 2- group must sum to -2. The Ox. No. of each O is -2 for a total of -8. So the Sulfur atom is +6.

Figure 4.20: A summary of an oxidation- reduction process, in which M is oxidized and X is reduced.

Recognizing Oxidizing and Reducing Agents - I Problem: Identify the oxidizing and reducing agent in each of the Rx: a) Zn (s) + 2 HCl (aq) ZnCl 2 (aq) + H 2 (g) b) S 8 (s) + 12 O 2 (g) 8 SO 3 (g) c) NiO (s) + CO (g) Ni (s) + CO 2 (g) Plan: First we assign an oxidation number (O.N.) to each atom (or ion) based on the rules in Table 4.3. The reactant is the reducing agent if it contains an atom that is oxidized (O.N. increased in the reaction). The reactant is the oxidizing agent if it contains an atom that is reduced ( O.N. decreased). Solution: a) Assigning oxidation numbers: Zn (s) + 2 HCl (aq) ZnCl 2 (aq) + H 2 (g) HCl is the oxidizing agent, and Zn is the reducing agent!

Recognizing Oxidizing and Reducing Agents - II b) Assigning oxidation numbers: S 8 (s) + 12 O 2 (g) 8 SO 3 (g) S 8 is the reducing agent and O 2 is the oxidizing agent c) Assigning oxidation numbers: S [0] S[+6] S is Oxidized O[0] O[-2] O is Reduced NiO (s) + CO (g) Ni (s) + CO 2 (g) Ni[+2] Ni[0] Ni is Reduced C[+2] C[+4] C is oxidized CO is the reducing agent and NiO is the oxidizing agent

Activity Series of the Metals Li Li + + e - K K + + e - These elements react rapidly with aqueous H + ions Ba Ba e - (acid) or with liquid H 2 O to release H 2 gas. Ca Ca e - Na Na + + e - Mg Mg e - Al Al e - These elements react with aqueous H + ions or with Mn Mn e - steam to release H 2 gas. Zn Zn e - Cr Cr e - Fe Fe e - Co Co e - Ni Ni e - These elements react with aqueous H + ions to Sn Sn e - release H 2 gas. H 2 2 H + + 2e - Cu Cu e - Ag Ag + + e - These elements do not react with aqueous H + ions Hg Hg e - to release H 2 gas. Pt Pt e - Au Au e - Strongly reducing Weakly reducing

Oxidation of copper metal by nitric acid.

Examples of Activity Series Problems Cu (s) + 2 Ag + (aq) Cu 2+ (aq) + 2 Ag (s) 2 Fe (s) + 3 Cu 2+ (aq) 2 Fe 3+ (aq) + 3 Cu (s) Mg (s) + Zn 2+ (aq) Mg 2+ (aq) + Zn (s) 2 K (s) + Sn 2+ (aq) 2 K + (aq) + Sn (s) Pt (s) + Ni 2+ (aq) N. R. 2 Al (s) + 6 H + (aq) 2 Al 3+ (aq) + 3 H 2 (g) Au (s) + H + (aq) N. R.

Problem: Calculate the mass of metallic Iron that must be added to liters of a solution containing M of Pt 2+ (aq) ions in solution to reclaim all of the Platinum. Solution: V x M = # moles 500.0L x Mol/L = 0.20 Mol Pt 2+ assume that the Iron goes to Fe 3+ therefore we will need only 2 moles of Iron for every 3 moles of Platinum 0.20 mol Pt 2+ x = mol Fe mol Fe x = ____________ g Fe 2 mol Fe 3 mol Pt g Fe mol Fe

Balancing REDOX Equations: The oxidation number method Step 1) Assign oxidation numbers to all elements in the equation. Step 2) From the changes in oxidation numbers, identify the oxidized and reduced species. Step 3) Compute the number of electrons lost in the oxidation and gained in the reduction from the oxidation number changes. Draw tie-lines between these atoms to show electron changes. Step 4) Multiply one or both of these numbers by appropriate factors to make the electrons lost equal the electrons gained, and use the factors as balancing coefficients. Step 5) Complete the balancing by inspection, adding states of matter.

REDOX Balancing using Ox. No. Method - I ___ H 2 (g) +___ O 2 (g) ___ H 2 O (g) e e - electrons lost must = electrons gained therefore multiply Hydrogen reaction by 2! and we are balanced! 22

REDOX Balancing Using Ox. No. Method - II Fe +2 (aq) + MnO 4 - (aq) + H 3 O + (aq) Fe +3 (aq) + Mn +2 (aq) + H 2 O (aq) -1e e - 5 Fe +2 (aq) + MnO 4 - (aq) +8 H 3 O + (aq) 5 Fe +3 (aq) + Mn +2 (aq) +12 H 2 O (aq) Multiply Fe +2 & Fe +3 by five to correct for the electrons gained by the Manganese. 5 Fe +2 (aq) + MnO 4 - (aq) + H 3 O + (aq) 5 Fe +3 (aq) + Mn +2 (aq) + H 2 O (aq) Make four water molecules from protons from the acid, and the oxygen from the MnO 4 -, this will require 8 protons, or Hydronium ions. This will give a total of 12 water molecules formed.

Balancing Oxidation-Reduction Equations Occurring in Acidic Solution by the Half- Reaction Method.

Balancing Oxidation- Reduction Equations Occurring in Basic Solution by the Half- Reaction Method.

Balancing Redox Equations in Aqueous Acid and Base Solutions : ACID : You may add either H + ( H 3 O + ), or water ( H 2 O ) to either side of the chemical equation. BASE : You may add either OH -, or water to either side of the chemical equation. H + + OH - H 2 O H + + H 2 O H 3 O + H + + OH - H 2 O

REDOX Balancing by Half-Reaction Method-I Fe +2 (aq) + MnO 4 - (aq) Fe +3 (aq) + Mn +2 (aq) [acid solution] Identify Oxidation and Reduction Half Reactions Fe +2 (aq) Fe +3 (aq) + e - [oxidation half-reaction] MnO 4 - (aq) Mn +2 (aq) add H + to the reactants and that will give water as a product! MnO 4 - (aq) + 8H 3 O + (aq) +5e - Mn +2 (aq) + 12H 2 O (l) [reduction half-reaction] Sum the two half-reactions { Fe +2 (aq) Fe +3 (aq) +e - } x5 MnO 4 - (aq) + 8H 3 O + (aq) +5e - Mn +2 (aq) + 12H 2 O (l) MnO 4 - (aq) + 8H 3 O + (aq) +5e - +5Fe +2 (aq) 5Fe +3 (aq) +5e - + Mn +2 (aq) + 12H 2 O (l)

REDOX Balancing by Half-Reaction Method -II MnO 4 - (aq) + SO 3 2- (aq) MnO 2 (s) + SO 4 2- (aq) [basic solution] Oxidation: SO 3 2- SO 4 2- (aq) + 2e - Add OH - to the reactant side, and water to the product side to get oxygen to balance since we have one more oxygen on sulfate than on sulfite. SO 3 2- (aq) + 2 OH - (aq) SO 4 2- (aq) + H 2 O (l) + 2e - Reduction: MnO 4 - (aq) + 3e - MnO 2 (s) Add water to the reactant side and OH - to the product side to take up the oxygen lost when MnO 4 - goes to MnO 2 and looses two oxygen atoms. MnO 4 - (aq) + 2 H 2 O (l) + 3e - MnO 2 (s) + 4 OH - (aq) Multiply the oxidation equation by 3 to make the electrons 6. Multiply the reduction equation by 2 to make the electrons 6, and add the two. 3 SO 3 -2 (aq) + 2 MnO 4 - (aq) + H 2 O (l) 3 SO 4 -2 (aq) + 2 MnO 2 (s) + 2 OH - (aq)

REDOX Balancing by Half-Reaction Method-III MnO 4 - (aq) + SO 3 2- (aq) MnO 2 (s) + SO 4 2- (aq) [acidic solution] Oxidation: SO 3 2- (aq) SO 4 2- (aq) + 2 e - Add water to the reactant side to supply an oxygen and add two protons to the product side that will remain plus the two electrons. SO 3 2- (aq) + H 2 O (l) SO 4 2- (aq) + 2 H + (aq) + 2 e - Reduction: MnO 4 - (aq) + 3 e - MnO 2 (s) Add water to the product side to take up the extra oxygen from Mn cpds, and add Hydrogen to the reactant side. MnO 4 - (aq) + 3 e - + 4H + MnO 2 (s) + 2 H 2 O (l) Multiply the oxidation equation by 3, and the reduction equation by 2, and add them canceling out the electrons, protons and water molecules. 3SO 3 2- (aq) + 2MnO 4 - (aq) + 2H + (aq) 3 SO 4 2- (aq) + 2MnO 2 (s) + H 2 O (l)

REDOX Balancing using Ox. No. Method - III MnO 4 - (aq) + SO 3 2- (aq) MnO 2 (s) + SO 4 2- (aq) e e - To balance the electrons, we must multiply the sulfite by 3, and the permanganate by 2. We then have to account for the oxygen imbalance by adding acid to the reactant side, and water to the product side. 2 MnO 4 - (aq) + 3 SO 3 2- (aq) + H 3 O + (aq) 2 MnO 2 (s) + 3 SO 4 2- (aq) + H 2 O (aq) ( Acidic Solution ) 2 MnO 4 - (aq) + 3 SO 3 2- (aq) +2 H 3 O + (aq) 2 MnO 2 (s) + 3 SO 4 2- (aq) +3 H 2 O (aq) For the final balance it is necessary to realize that protons needed to bind up the oxygen atoms must be balanced, and since we have called H + ion - hydronium ions,therefore water will be formed!

REDOX Balancing by Half-Reaction Method-IV MnO 4 - (aq) +SO 3 2- (aq) MnO 2(s) + SO 4 2- (aq) [basic solution] balance the equation as if it were in acid, and then convert it to base: 2MnO 4 - (aq) + 3SO 3 2- (aq) + 2H + (aq) 2MnO 2(s) + 3SO 4 2- (aq) + H 2 O (l) To convert to base, add two OH - to each side of the equation: 2MnO 4 - (aq) + 3SO 3 2- (aq) + 2H + (aq) 2MnO 2(s) + 3SO 4 2- (aq) + H 2 O (l) 2MnO 4 - (aq) + 3SO 3 2- (aq) +2 H 2 O (l) 2MnO 2(s) + 3SO 4 2- (aq) + H 2 O (l) +2OH - (aq) On the reactant side, the H + and the OH - cancel to give water. Cancel out the water on each side of the equation, and you are done! 2MnO 4 - (aq) + 3SO 3 2- (aq) + H 2 O (l) 2MnO 2(s) + 3SO 4 2- (aq) +2OH - (aq)

REDOX Balancing Using Ox. No. Method-IV Zinc metal is dissolved in Nitric Acid to give Zn 2+ and the ammonium ion from the reduced Nitric acid, write the balanced chemical equation! Zn (s) + H + (aq) + NO 3 - (aq) Zn 2+ (aq) + NH 4 + (aq) Oxidation # method - 2 e e - Multiply Zinc and Zn 2+ by 4, and ammonia by unity. Since we have no oxygen on the product side, add 3 water molecules to the product side, requiring 10 H + on the reactant side. 4 Zn (s) +10 H + (aq) + NO 3 - (aq) 4 Zn 2+ (aq) + NH 4 + (aq) + 3 H 2 O (l)

REDOX Balancing by Half-Reaction Method-V Zn (s) + H 3 O + (aq) + NO 3 - (aq) Zn 2+ (aq) + NH 4 + (aq) Given: Oxidation: Zn (s) Zn e - Reduction: H 3 O + (aq) + NO 3 - (aq) + 8 e - NH 4 + (aq) + H 2 O (l) We will need three waters to pick up the oxygen from the nitrate ion, and for the hydrogen, we will need to have 10 hydrogen ions. Because the Hydrogen ions came as hydronium ions, we will need 10 more water molecules. 10 H 3 O + (aq) + NO 3 - (aq) + 8 e - NH 4 + (aq) + 13 H 2 O (l) Finally, if we are to add the two equations, we must multiply the Ox. one by 4 to be able to cancel out the electrons, so the final balanced equation is: 10 H 3 O + (aq) + NO 3 - (aq) + 4 Zn (s) 4 Zn +2 (aq) + NH 4 + (aq) + 13 H 2 O (l)

REDOX Balancing by Half-Reaction Method -VI - A In acid Potassium dichromate reacts with ethanol(C 2 H 5 OH) to yield the blue-green solution of Cr +3, the reaction used in “breathalyzers”. H 3 O + (aq) + Cr 2 O 7 2- (aq) + C 2 H 5 OH (l) Cr 3+ (aq) + CO 2 (g) + H 2 O (l) Oxidation: C 2 H 5 OH (l) CO 2 (g) We need to balance oxygen by adding water to the reactant side, and balance Hydrogen by adding protons to the product side. C 2 H 5 OH (l) + 3 H 2 O (l) 2 CO 2 (g) + 12 H + (aq) C 2 H 5 OH (l) + 15 H 2 O (l) 2 CO 2 (g) + 12 H 3 O + (aq) + 12 e - Since we wish to consider H + as the Hydronium ion - H 3 O +, we must add 12 water molecules to the reactant side, and make the H + into H 3 O +.

REDOX Balancing by Half-Reaction Method - VI - B Reduction: Cr 2 O 7 2- (aq) Cr +3 (aq) Dichromate has two chromium atoms, therefore the products need to have two Cr +3, and 3 electrons per atom. The oxygen atoms from the dichromate need to be taken up as water on the product side by adding protons to the reactant side. 14H + (aq) + Cr 2 O 7 2- (aq) Cr +3 (aq) + 7 H 2 O (l) Each Chromium atom changes oxidation from a +6 to a +3 there by accepting 6 electrons, so we add 6 electrons to the reactant side. 6e H 3 O + (aq) + Cr 2 O 7 2- (aq) 2 Cr +3 (aq) + 21 H 2 O (l) Adding the two equations will give the final equation: Ox: C 2 H 5 OH (l) + 15 H 2 O (l) 2 CO 2 (g) + 12 H 3 O + (aq) + 12 e - [6e H 3 O + (aq) + Cr 2 O 7 2- (aq) 2 Cr +3 (aq) + 21 H 2 O (l) ] x 2 C 2 H 5 OH (l) + 16 H 3 O + (aq) + 2 Cr 2 O 7 2- (aq) 2 CO 2 (g) + 4 Cr +3 (aq) + 27 H 2 O (l) Rd:

REDOX Balancing by Half-Reaction Method -VII - A Silver is reclaimed from ores by extraction using basic Cyanide ion. Ag (s) + CN - (aq) + O 2 (g) Ag(CN) 2 - (aq) OH - Oxidation: CN - (aq) + Ag (s) Ag(CN) 2 - (aq) Since we need two cyanide ions to form the complex, add two to the reactant side of the equation. Silver is also oxidized, so it looses an electron, so we add one electron to the product side. 2 CN - (aq) + Ag (s) Ag(CN) 2 - (aq) + e - Reduction: O 2 (g) + H 2 O (aq) OH - (l) Since oxygen is to form oxide ions, 4 electrons need to be added to the reactant side, and 2 water molecules are needed to supply the hydrogen to make hydroxide ions, yielding 4 OH - ions. 4 e - + O 2 (g) + 2 H 2 O (aq) 4 OH - (l)

REDOX Balancing by Half-Reaction Method - VII - B Adding the Reduction equation to the Oxidation equation will require the Oxidation one to be multiplied by 4 to eliminate the electrons. Ox (x4) 8CN - (aq) + 4 Ag (s) 4 Ag(CN) 2 - (aq) + 4 e - Rd 4 e - + O 2 (g) + 2 H 2 O (l) 4 OH - (aq) 8 CN - (aq) + 4 Ag (s) + O 2 (g) + 2 H 2 O (l) 4 Ag(CN) 2 - (aq) + 4 OH - (aq)

REDOX Balancing Using Ox. No. Method -V Ag (s) + CN - (aq) + O 2 (g) Ag(CN) 2 - (aq) + OH - (aq) e e - To balance electrons we must put a 4 in front of the Ag, since each oxygen looses two electrons, and they come two at a time! That requires us to put a 4 in front of the silver complex, yielding 8 cyanide ions. 4 Ag (s) + 8 CN - (aq) + O 2 (g) 4 Ag(CN) 2 - (aq) + OH - (aq) We have no hydrogen on the reactant side therefore we must add water as a reactant, and since we also add oxygen, we must add two water molecules, that well give us 4 hydroxide anions, giving us a balanced chemical equation. 4 Ag (s) + 8 CN - (aq) + O 2 (g) + 2 H 2 O (l) 4 Ag(CN) 2 - (aq) + 4 OH - (aq)

Silver Precipitates 2 Ag + (aq) + Na 2 CO 3(aq) Ag 2 CO 3(s) + 2Na + (aq) [White] Ag 2 CO 3(s) + 2NaOH (aq) 2AgOH (s) + Na 2 CO 3 (aq) [Brown] AgOH (s) + NaCl (aq) AgCl (s) + NaOH (aq) [White] AgCl (s) + 2NH 3(aq) Ag(NH 3 ) 2 + (aq) + Cl - (aq) [Clear] Ag(NH 3 ) 2 + (aq) + NaBr (aq) AgBr (s) + 2NH 3(aq) + Na + (aq) [Yellow] AgBr (s) + 2Na 2 S 2 O 3(aq) Ag(S 2 O 3 ) 2 -3 (aq) + Br - (aq) + 4Na + (aq) [Clear] Ag(S 2 O 3 ) 2 -3 (aq) + NaI (aq) AgI (s) + 2S 2 O 3 -2 (aq) Na + (aq) [Yellow] 2AgI (s) + Na 2 S (aq) Ag 2 S (s) + 2I - (aq) [Black]

Redox Titration- Calculation outline - I Volume (L) of KMnO 4 Solution Moles of KMnO 4 M (mol/L) Molar ratio Chemical Formulas Moles of CaC 2 O 4 Problem: Calcium Oxalate was precipitated from blood by the addition of Sodium Oxalate so that calcium ion could be determined. In the blood sample. The sulfuric acid solution that the precipitate was dissolved in required 2.05 ml of 4.88 x M KMnO 4 to reach the endpoint. a) calculate the amount (mol) of Ca +2. b) calculate the Ca +2 ion conc. Plan: a) Calculate the molarity of Ca +2 in the H 2 SO 4 solution. b) Convert the Ca +2 concentration into units of mg Ca +2 / 100 ml blood. Moles of Ca +2 a) b) c)

Figure 4.21: Permanganate being introduced into a flask of reducing agent.

Redox Titration - Calculation - I Equation: 2 KMnO 4 (aq) + 5 CaC 2 O 4 (aq) + 8 H 2 SO 4 (aq) 2 MnSO 4 (aq) + K 2 SO 4 (aq) + 5 CaSO 4 (aq) + 10 CO 2 (g) + 8 H 2 O (L) a) Moles of KMnO 4 Mol = Vol x Molarity Mol = L x 4.88 x mol/L Mol = 1.00 x mol KMnO 4 b) Moles of CaC 2 O 4 Mol CaC 2 O 4 = 1.00 x mol KMnO 4 x = 5 mol CaC 2 O 4 2 mol KMnO 4 Mol CaC 2 O 4 = 2.50 x mol CaC 2 O 4 c) Moles of Ca +2 Mol Ca +2 = 2.50 x mol CaC 2 O 4 x = 1 mol Ca +2 1 mol CaC 2 O 4 Mol Ca +2 = 2.50 x mol Ca +2

Redox Titration - Calculation Outline - II Moles of Ca 2+ / 1 ml of blood Moles of Ca 2+ / 100 ml blood Mass (g) of Ca 2+ / 100 ml blood Mass (mg) of Ca 2+ / 100 ml blood multiply by 100 a) Calc of mol Ca +2 per 100 ml M (g/mol) b) Calc of mass of Ca +2 per 100 ml 1g = 1000mg c) convert g to mg!

Redox Titration - Calculation - II a) Mol Ca +2 per 100 ml Blood b) mass (g) of Ca +2 Mol Ca ml Blood = x 100 ml Blood = Mol Ca ml Blood Mol Ca ml Blood = x 100 ml Blood = 2.50 x mol Ca ml Blood Mol Ca ml Blood = 2.50 x mol Ca +2 Mass Ca +2 = Mol Ca +2 x Mol Mass Ca/ mol = Mass Ca +2 = 2.50 x mol Ca +2 x 40.08g Ca/mol = g Ca +2 c) mass (mg) of Ca +2 Mass Ca +2 = g Ca +2 x 1000mg Ca +2 /g Ca +2 = 10.0 mg Ca ml Blood

Types of Chemical Reactions - I I) Combination Reactions that are Redox Reactions a) A Metal and a Non-Metal form an Ionic compound b) Two Non-Metals form a Covalent compound c) Combination of an Element and a Compound II) Combination Reactions that are not Redox Reactions a) A Metal oxide and a Non-Metal form an ionic compound with a polyatomic anion b) Metal Oxides and water form Bases c) Non-Metal Oxides and water form Acids III) Decomposition Reactions a) Thermal Decomposition i) Many ionic compounds with oxoanions form a metal oxide and a gaseous non-metal ii) Many Metal oxides, Chlorates, and Perchlorates release Oxygen b) Electrolytic Decomposition

Types of Chemical Reactions - II IV) Displacement Reactions a) Single -Displacement Reactions - Activity Series of the Metals i) A metal displaces hydrogen from water or an acid ii) A metal displaces another metal ion from solution iii) A halogen displaces a halide ion from solution b) Double -Displacement Reactions - i) In Precipitation Reactions- A precipitate forms ii) In acid-Base Reactions - Acid + Base form a salt & water V) Combustion Reactions - All are Redox Processes a) Combustion of an element with oxygen to form oxides b) Combustion of Hydrocarbons to yield Water & Carbon Dioxide ReactantsProducts