Chapter 4 Aqueous Reactions and Solution Stoichiometry Prentice Hall © 2003 Chapter 4

Solutions Solutions are defined as homogeneous mixtures of two or more pure substances. The solvent is present in greatest abundance. All other substances are solutes. When water is the solvent, the solution is called an aqueous solution.

General Properties of Aqueous Solutions Electrolytic Properties Aqueous solutions, solutions in water, have the potential to conduct electricity. The ability of the solution to conduct depends on the number of ions in solution. There are three types of solution: Strong electrolytes, Weak electrolytes, and Nonelectrolytes. Prentice Hall © 2003 Chapter 4

General Properties of Aqueous Solutions Ionic Compounds in Water Ions dissociate in water. In solution, each ion is surrounded by water molecules. Transport of ions through solution causes flow of current. Prentice Hall © 2003 Chapter 4

General Properties of Aqueous Solutions Molecular Compounds in Water Molecular compounds in water (e.g., CH3OH): no ions are formed. If there are no ions in solution, there is nothing to transport electric charge. Prentice Hall © 2003 Chapter 4

Electrolytic Properties electrolyte nonelectrolyte Prentice Hall © 2003 Chapter 4

Strong and Weak Electrolytes Strong electrolytes: completely dissociate in solution Acids ionize Soluble Salts Strong Acids (HCl, HBr, HI, HNO3, HClO4, HClO3, H2SO4) HCl(aq)  H+(aq) + Cl-(aq) KNOW THE 7 STRONG ACIDS! Prentice Hall © 2003 Chapter 4

These ions exist in equilibrium with the un-ionized substance Weak electrolytes: produce a small concentration of ions when they dissolve (weak acids) These ions exist in equilibrium with the un-ionized substance For example: Prentice Hall © 2003 Chapter 4

Sample Problems p. 158: # 17, 19 Prentice Hall © 2003 Chapter 4

4.2 Precipitation Reactions When two solutions containing soluble salts are mixed, sometimes an insoluble salt will be produced. A salt “falls” out of solution, like snow out of the sky. This solid is called a precipitate.

Solubility of Ionic Compounds Not all ionic compounds dissolve in water. A list of solubility rules is used to decide what combination of ions will dissolve.

Text, P. 118

Sample Problem p. 158: # 21, 23 Prentice Hall © 2003 Chapter 4

Metathesis (Exchange) Reactions Metathesis comes from a Greek word that means “to transpose” AgNO3 (aq) + KCl (aq)  AgCl (s) + KNO3 (aq)

Metathesis (Exchange) Reactions Metathesis comes from a Greek word that means “to transpose” It appears the ions in the reactant compounds exchange, or transpose, ions AgNO3 (aq) + KCl (aq)  AgCl (s) + KNO3 (aq)

Metathesis (Exchange) Reactions Metathesis comes from a Greek word that means “to transpose” It appears the ions in the reactant compounds exchange, or transpose, ions AgNO3 (aq) + KCl (aq)  AgCl (s) + KNO3 (aq)

Signs of a chemical reaction! Exchange (Metathesis) Reactions Metathesis reactions involve swapping ions in solution: (anions/cations switch partners) AX + BY  AY + BX Metathesis reactions will lead to a change in solution if one of three things occurs: an insoluble solid is formed (precipitate) weak or nonelectrolytes are formed an insoluble gas is formed Signs of a chemical reaction! Prentice Hall © 2003 Chapter 4

Sample Problems p. 158: # 25, 27, 29 Prentice Hall © 2003 Chapter 4

Solution Chemistry It is helpful to pay attention to exactly what species are present in a reaction mixture (i.e., solid, liquid, gas, aqueous solution). If we are to understand reactivity, we must be aware of just what is changing during the course of a reaction.

AgNO3 (aq) + KCl (aq)  AgCl (s) + KNO3 (aq) Molecular Equation The molecular equation lists the reactants and products in their molecular form. AgNO3 (aq) + KCl (aq)  AgCl (s) + KNO3 (aq)

Ionic Equation In the ionic equation all strong electrolytes (strong acids, strong bases, and soluble ionic salts) are dissociated into their ions. This more accurately reflects the species that are found in the reaction mixture. Ag+ (aq) + NO3- (aq) + K+ (aq) + Cl- (aq)  AgCl (s) + K+ (aq) + NO3- (aq)

Net Ionic Equation To form the net ionic equation, cross out anything that does not change from the left side of the equation to the right. Ag+(aq) + NO3-(aq) + K+(aq) + Cl-(aq)  AgCl (s) + K+(aq) + NO3-(aq)

Ag+(aq) + Cl-(aq)  AgCl (s) Net Ionic Equation To form the net ionic equation, cross out anything that does not change from the left side of the equation to the right. The only things left in the equation are those things that change (i.e., react) during the course of the reaction. Ag+(aq) + Cl-(aq)  AgCl (s)

Net Ionic Equation To form the net ionic equation, cross out anything that does not change from the left side of the equation to the right. The only things left in the equation are those things that change (i.e., react) during the course of the reaction. Those things that didn’t change (and were deleted from the net ionic equation) are called spectator ions. Ag+(aq) + NO3-(aq) + K+(aq) + Cl-(aq)  AgCl (s) + K+(aq) + NO3-(aq)

(ions that appear on both sides of the equation) Ionic Equations Ionic equation: used to highlight a reaction between ions (formation of an insoluble product, molecular compound or a gas) Molecular equation: all species listed as molecules (“the reaction”): HCl(aq) + NaOH(aq)  H2O(l) + NaCl(aq) Complete ionic equation: lists all ions (soluble compounds are broken into ions): H+(aq) + Cl-(aq) + Na+(aq) + (OH)-(aq)  H2O(l) + Na+(aq) + Cl-(aq) Net ionic equation: lists only unique ions (the insoluble product, molecular compound or a gas) H+(aq) + (OH)-(aq)  H2O(l) Cancel Spectator Ions (ions that appear on both sides of the equation) Prentice Hall © 2003 Chapter 4

Sample Problem # 31, 33 Prentice Hall © 2003 Chapter 4

4.3 Acid-Base Reactions Acids Dissociation = ions in a solid move apart in solution. Ionization = a neutral substance forms ions in solution. Acid = substances that ionize to form H+ in solution monoprotic(HCl) diprotic(H2SO4) many acidic protons = polyprotic Prentice Hall © 2003 Chapter 4

Bases Bases = substances that react with the H+ ions formed by acids (e.g. NH3, Drano™, Milk of Magnesia™) Prentice Hall © 2003 Chapter 4

Strong and Weak Acids and Bases Strong acids and bases are strong electrolytes They are completely ionized in solution Weak acids and bases are weak electrolytes They are partially ionized in solution Prentice Hall © 2003 Chapter 4

Text, P. 134 Prentice Hall © 2003 Chapter 4

Review: Identifying Strong and Weak Electrolytes STRONG ELECTROLYTES Water soluble and ionic Strong acid or strong base WEAK ELECTROLYTES Water soluble and not ionic Weak acid or weak base Otherwise, the compound is a NONELECTROLYTE Prentice Hall © 2003 Chapter 4

Sample Problems: # 35, 37 Prentice Hall © 2003 Chapter 4

Neutralization Reactions and Salts Acid + Base  Salt + Water (Neutralization Reaction) HCl(aq) + NaOH(aq)  H2O(l) + NaCl(aq) SA/SB net ionic: H+(aq) + (OH)-(aq)  H2O(l) HC2H3O2(aq) + NaOH(aq)  NaC2H3O2 (aq) + H2O(l) WA/SB net ionic: HA(aq) + (OH)-(aq)  A- (aq) + H2O(l) Salt Ionic compound Composed of the cation from the base and the anion from the acid Prentice Hall © 2003 Chapter 4

Is this assault?

What are salts? In a acid-base reaction, a salt is formed when a metallic ion or an ammonium ion replaces one or more hydrogen ions in an acid. E.g. Zinc hydroxide + sulfuric acid  zinc sulfate + water Zn(OH)2 (s) + H2SO4 (aq)  ZnSO4 (aq) + H20 (l) Zn2+ comes from the base, Zn(OH)2 ZnSO4 SO42- comes from sulfuric acid, H2SO4

Salts A salt is an ionic compound that: comes from the anion of an acid comes from the cation of a base is formed from a neutralization reaction some neutral; others acidic or basic

Acid-Base Reactions with Gas Formation Sulfide and carbonate ions can react with H+ in a similar way to OH- 2HCl(aq) + Na2S(aq)  H2S(g) + 2NaCl(aq) net ionic: 2H+(aq) + S2-(aq)  H2S(g) HCl(aq) + NaHCO3(aq)  NaCl(aq) + H2O(l) + CO2(g) net ionic: H+(aq) + (HCO3)-(aq) H2O(l) + CO2(g) Prentice Hall © 2003 Chapter 4

Sample Problems # 39,41,43 Prentice Hall © 2003 Chapter 4

4.4 Oxidation-Reduction Reactions Oxidation and Reduction Oxidized: an atom, molecule, or ion becomes more positively charged Oxidation is the loss of electrons Reduced: an atom, molecule, or ion becomes less positively charged Reduction is the gain of electrons Oxidation and reduction always occur simultaneously Prentice Hall © 2003 Chapter 4

Oxidizing Agent: makes it possible for something else to be oxidized The agent is reduced Reducing Agent: makes it possible for something else to be reduced The agent is oxidized Prentice Hall © 2003 Chapter 4

Oxidation Is a Loss of electrons Reduction Is a Gain of electrons And Remember: OIL RIG Oxidation Is a Loss of electrons Reduction Is a Gain of electrons And LEO the lion says GER Loss of Electrons is Oxidation Gain of Electrons is Reduction LEO the GERbil Prentice Hall © 2003 Chapter 4

Oxidation number for an ion: the charge on the ion Oxidation Numbers Oxidation number for an ion: the charge on the ion Oxidation number for an atom: the hypothetical charge that atom would have if it was an ion It keeps track of electron transfers Prentice Hall © 2003 Chapter 4

Oxidation numbers are assigned by a series of rules: Text, P. 140 - 141 Oxidation numbers are assigned by a series of rules: If the atom is in its elemental form, the oxidation number is zero. ( Cl2 , H2 , P4 ) For a monoatomic ion, the charge on the ion is the oxidation state. (F- , O-2 , Ca+2 ) Prentice Hall © 2003 Chapter 4

Oxidation Numbers: Rules Nonmetals usually have negative oxidation numbers: Oxidation number of O is usually –2 In the peroxide ion, O22-, the oxygen atom has an oxidation number of –1 b) Oxidation number of H is +1 when bonded to nonmetals and –1 when bonded to metals c) The oxidation number of F is –1. The other halogens have an oxidation number of -1 in most binary compounds. When combined with oxygen, as in oxyanions, however, they have positive oxidation states. Prentice Hall © 2003 Chapter 4

Oxidation Numbers: Rules The sum of the oxidation numbers for a polyatomic ion is equal to the charge on the molecule The sum of the oxidation numbers is zero for a neutral compound Prentice Hall © 2003 Chapter 4

Sample Problems # 47, 49, 51 Prentice Hall © 2003 Chapter 4

Oxidation of Metals by Acids and Salts Metals are oxidized by acids to form salts: Zn(s) +2HCl(aq)  ZnCl2(aq) + H2(g) During the reaction, 2H+(aq) is reduced to H2(g) Metals can also be oxidized by other salts: Net ionic: Fe(s) +Ni2+(aq)  Fe2+(aq) + Ni(s) Prentice Hall © 2003 Chapter 4

Some metals are easily oxidized whereas others are not Activity Series Some metals are easily oxidized whereas others are not Activity series: a list of metals arranged in decreasing ease of oxidation Notice the TYPES of elements & their location on the series Prentice Hall © 2003 Chapter 4

Text, P. 144

From Your Handout Prentice Hall © 2003 Chapter 4

“Activity Series” for halogens: F2 > Cl2 > Br2 > I2 Activity Series, Continued The higher the metal on the activity series, the more active that metal (“No Reaction” is possible) An element can only “replace” an element that is below it on the activity series. Form: AX + B  BX + A a “spontaneous reaction” “Activity Series” for halogens: F2 > Cl2 > Br2 > I2 Form: AX + Z2  AZ + X2 “Like replaces like” Prentice Hall © 2003 Chapter 4

Identifying Redox Equations In general, all chemical reactions can be assigned to one of two classes: oxidation-reduction, in which electrons are transferred: Single-replacement, combination, decomposition, and combustion this second class has no electron transfer, and includes all others: Double-replacement and acid-base reactions

Sample Problems p. 159 #53, 55, 57 Prentice Hall © 2003 Chapter 4

4.5 Concentrations of Solutions Molarity Solution = solute dissolved in solvent Solute: present in smallest amount Water as solvent (aqueous solutions) Concentration: the amount of solute dissolved in a given quantity of solvent or solution. Change concentration by using different amounts of solute and solvent Molarity: Moles of solute per liter of solution If we know molarity and liters of solution, we can calculate moles (and mass) of solute Prentice Hall © 2003 Chapter 4

Molarity Remember: V * M = moles Prentice Hall © 2003 Chapter 4

Sample Problems P. 160: #59, 61, 63, 69, 71 Can you EXPLAIN how to prepare these solutions? Prentice Hall © 2003 Chapter 4

4.5

Always Add Acid to Water Dilution We recognize that the number of moles are the same in dilute and concentrated solutions. So: MdiluteVdilute = moles = MconcentratedVconcentrated Always Add Acid to Water Prentice Hall © 2003 Chapter 4

Dilution is the procedure for preparing a less concentrated solution from a more concentrated solution. Dilution Add Solvent Moles of solute before dilution (i) after dilution (f) = MiVi MfVf = 4.5

Sample Problems P. 160 # 73, 75, 77 (also do an HCl dilution and explain how to prepare the solution), Prentice Hall © 2003 Chapter 4

4.6: Solution Stoichiometry and Chemical Analysis There are two different types of units: laboratory units (macroscopic units) chemical units (microscopic units) Convert! Grams  moles (using molar mass) Volume or molarity  moles (using M = mol/L) Use the stoichiometric coefficients to move between reactants and products Prentice Hall © 2003 Chapter 4

4.6: Solution Stoichiometry and Chemical Analysis Figure 4.18, P. 139 of text Prentice Hall © 2003 Chapter 4

Titration Terms Titration: Controlled combination of a solution of unknown concentration with a solution of known concentration Primary Standard: A relatively pure compound with a high molar mass known pattern of reactivity used as a basis of comparison for the determination of the concentration of an unknown Secondary Standard: A solution with a molarity that is very accurately known Standard Solution: A solution of known concentration. Prentice Hall © 2003 Chapter 4

Indicator: Dyes used to indicate the equivalence point (color change) Titration Terms, continued Equivalence Point: Stoichiometrically equivalent amounts of substances are combined Indicator: Dyes used to indicate the equivalence point (color change) Phenolphthalein: colorless in an acid, pink in a base End Point: the point of color change that coincides with the equivalence point. Prentice Hall © 2003 Chapter 4

Titrations

Sample problems P. 160 #79, 81, 83 Prentice Hall © 2003 Chapter 4