Chapter 4: Chemical Reactions in Aqueous Solutions

Solute Concentrations, Molarity  Solution: homogeneous mixture of two or more substances.  Solute: the substance being dissolved.  Solvent: the substance doing the dissolving.

 Concentration of a solution: the quantity of a solute in a given quantity of solution (or solvent). A concentrated solution contains a relatively large amount of solute vs. the solvent (or solution). A dilute solution contains a relatively small concentration of solute vs. the solvent (or solution). “Concentrated” and “dilute” aren’t very quantitative.

Molar Concentration Molarity (M), or molar concentration, is the amount of solute, in moles, per liter of solution: Formula:  Molarity = moles of solute liters of solution  A solution that is 0.35 M sucrose contains 0.35 moles of sucrose in each liter of solution.  Keep in mind that molarity signifies moles of solute per liter of solution, not liters of solvent.

Volumetric Flasks Weigh mol (1.580 g) KMnO 4. Dissolve in water. How much water? Doesn’t matter, as long as we don’t go over a liter. Add more water to reach the liter mark.

Calculating Molarity a. Given moles of solute and volume of solution Example 4.1 What is the molarity of a solution prepared by dissolving 5.0 mol of NaCl in enough water to make 2 L of solution?

Sample Problem b. Given mass of solute and volume of solution Example 4.2 What is the molarity of a solution prepared by dissolving g of AgNO 3 in enough water to make 1500 mL of solution?

Sample Problem c. Calculating Mass from Molarity and Volume Example 4.3 How many grams of sodium carbonate are needed to prepare L of a M solution?

 Why do some solutions conduct electricity?  Arrhenius’s theory: Certain substances dissociate into cations and anions when dissolved in water. These ions allow electricity to flow. Arrhenius’s Theory of Electrolytic Dissociation

Electrolytes dissociate to produce ions. Review: Ch 2 Ionic Compounds The more the electrolyte dissociates, the more ions it produces.

Types of Electrolytes  A strong electrolyte dissociates completely. A strong electrolyte is present in solution almost exclusively as ions. Strong electrolyte solutions are good conductors. Ex: most ionic compounds

Types of Electrolytes  A weak electrolyte dissociates partially. Weak electrolyte solutions are poor conductors. Different weak electrolytes dissociate to different extents. Ex: weak acids and bases

Nonelectrolytes  A nonelectrolyte does not dissociate. A nonelectrolyte is present in solution almost exclusively as molecules. Nonelectrolyte solutions do not conduct electricity. Ex: sugar, ethanol

 Strong electrolytes include: Strong acids & Strong bases Most water-soluble ionic compounds  Weak electrolytes include: Weak acids and weak bases A few ionic compounds  Nonelectrolytes include: Most molecular compounds Most organic compounds (most are molecular) Is it a strong electrolyte, a weak electrolyte, or a nonelectrolyte? How do we tell whether an acid (or base) is weak? Memorize the strong acids/bases!

Ion Concentrations in Solution A. Dissolution equations -For a strong electrolyte  write cations and anions as individual particles NaCl(aq)  MgSO 4  Na 2 SO 4 

Ion Concentrations in Solution A. Dissolution equations -For a strong electrolyte  write cations and anions as individual particles NaCl(aq)  Na + + Cl - MgSO 4  Mg +2 + SO 4 2- Na 2 SO 4  2Na + + SO 4 2-

Multiply molarity (M) x COEFFICIENT of each ion to find the ion concentration!  EXAMPLE: Na 2 SO 4  2Na + + SO M M M Calculating Ion Concentrations in Solution

Example Calculate the concentrations of all the ions present in the following solutions: a. 0.75M solution of NaCl b. 1.5M solution of AlF 3 c. 0.5M solution of Al 2 (CO 3 ) 3

Example a. 0.75M solution of NaCl NaCl(s)  Na + + Cl - b. 1.5M solution of AlF 3 AlF 3  Al F - c. 0.5M solution of Al 2 (CO 3 ) 3 Al 2 (CO 3 ) 3  2Al CO 3 -2

Chemical Reactions in Water We will look at DOUBLE REPLACEMENT REACTIONS Pb(NO 3 ) 2 (aq) + 2 KI(aq) ----> PbI 2 (s) + 2 KNO 3 (aq) The cations change places.

Definition  Precipitate: insoluble product (solid formed in a DR reaction)  Example: PbI 2 is a precipitate: Pb(NO 3 ) 2 (aq) + 2 KI(aq) ---> PbI 2 (s) + 2 KNO 3 (aq)

The “driving force” is the formation of an insoluble compound — a precipitate. Pb(NO 3 ) 2 (aq) + 2 KI (aq) -----> 2 KNO 3 (aq) + PbI 2 (s) Precipitation Reactions

Solubility Rules See handout!!

Solubility Rules Are the following compounds soluble or insoluble? NaNO 3 Ba(OH) 2 NaSO 4 CaSO 4

Use of Solubility Tables—See Reference Booklet s = soluble -- dissolves in water, is (aq) i = ss = insoluble or slightly soluble--forms a precipitate, is (s)

Water Solubility of Ionic Compounds Common minerals are often formed with anions that lead to insolubility: sulfidefluoridecarbonate oxide Azurite, a copper carbonate Iron pyrite, a sulfide Orpiment, arsenic sulfide

Net Ionic Equations Pb(NO 3 ) 2 (aq) + 2 KI (aq) -----> 2 KNO 3 (aq) + PbI 2 (s) Net ionic equation (See handout!) Pb 2+ (aq) + 2 I - (aq) ---> PbI 2 (s)

Sample Problem Predict whether or not the following pairs of reactants will form precipitates. Then write the double replacement reaction, the full ionic equation, and the net ionic equation.  a. CuCl 2 and (NH 4 ) 2 SO 4  b. Ba(NO 3 ) 2 and Na 2 CO 3  c. MgCl 2 and AgNO 3

Acids and Bases

Properties of Acids Taste sour React with indicators: turn blue litmus red turn phenolphthalein colorless React with certain metals to form H 2 Corrosive to metals and skin (strong acids) Neutralize bases to form water and salts

Properties of Bases Taste bitter Feel slippery React with indicators: turn red litmus blue turn phenolphthalein pink Corrosive to skin (strong bases) Neutralize acids to form water and salts

Arrhenius Definitions of Acids and Bases  Acids produce H + in water solution  Bases produce OH - in water solution Examples: HCl  H + + Cl - NaOH  Na + + OH -

Arrhenius Definitions: Acids produce H + in water solution Bases produce OH - in water solution

Arrhenius definitions are limited!!!  Not all bases contain OH -  H + does not exist by itself in aqueous systems

Definition: Hydronium Ion In aqueous solution, H + does NOT exist! Note: In problems, [H + ] = [H 3 O + ] H + + H 2 O  H 3 O + (hydronium ion)

An acid ----> H + in water SIX strong acids are HClhydrochloric HBrhydrobromic HIhydroiodic H 2 SO 4 sulfuric HClO 4 perchloric HNO 3 nitric HNO 3 STRONG ACIDS = STRONG ELECTROLYTES

Strong Acids/Bases Strong Acids: THERE ARE ONLY SIX! HCl, HBr, HI, HNO 3, HClO 4, H 2 SO 4 Strong acids are 100% ionized in solution! (Use single arrow; no equilibrium established) HNO 3  H + + NO 3 -

 Strong acids are strong electrolytes; completely ionized in water. In water:HCl(aq) → H + (aq) + Cl – (aq) No HCl in solution, only H + and Cl – ions. Reactions of Acids and Bases: Strong and Weak Acids

Strong Bases Strong Bases: Group I metals + OH -; some Group II metals + OH – Some examples: NaOH, LiOH, Sr(OH) 2 Strong bases are 100% ionized in solution! (Use single arrow; no equilibrium established) Sr(OH) 2  Sr OH -

Common Strong Acids and Strong Bases- memorize!!! MEMORIZE the strong acids and know how to recognize the strong bases!

 Weak acids are not fully ionized in water.  Weak organic acids contain the – COOH group  One of the best known is acetic acid = CH 3 COOH Weak Acids/Bases

 Weak bases are not fully ionized in water.(use double arrow; equilibrium established)  One of the best known weak bases is ammonia  NH 3 (aq) + H 2 O(liq)  NH 4 + (aq) + OH - (aq) Weak Acids/Bases

ACID-BASE THEORIES  The most general theory for common aqueous acids and bases is the BRØNSTED - LOWRY theory DEFINITIONS:  ACIDS DONATE H + IONS  BASES ACCEPT H + IONS

The Brønsted definition means NH 3 is a BASE in water — and water is itself an ACID ACID-BASE THEORIES

BRØNSTED - LOWRY theory : ACIDS DONATE H + IONS BASES ACCEPT H + IONS

NH 3 / NH 4 + is a conjugate pair — related by the gain or loss of H +NH 3 / NH 4 + is a conjugate pair — related by the gain or loss of H + Every acid has a conjugate base, formed when H + is removed from the acid.Every acid has a conjugate base, formed when H + is removed from the acid. Every base has a conjugate acid, formed when H + is added to the base.Every base has a conjugate acid, formed when H + is added to the base. Conjugate Pairs

Generalized equation: HB (aq) + A (aq)  HA (aq) + B - (aq)

Conjugate Pairs

Sample Problem Example Write the Bronsted Lowry equations for the weak acid HNO 2 and the weak base NH 3, identifying the conjugate acid-base pairs in each equilibrium.

Sample Problem Example Write the Bronsted Lowry equations for the weak acid HCO 3 - and the weak base NH 3, identifying the conjugate acid-base pairs in each equilibrium.

Some substances can function as both an ACID OR a BASE, depending on what they are reacted with. (can donate OR accept H + ) They are called amphiprotic or amphoteric E.g., H 2 O + H 2 O  H 3 O + + OH - Amphiprotic or Amphoteric Substances

For any sample of water molecules: H 2 O (liq) + H 2 O (liq)  H 3 O + (aq) + OH - (aq) K w = [H 3 O + ] [OH - ] = 1.00 x (at 25 o C) Water dissociation constant, or the ion product constant of water

Neutral, Acid and Basic Solution Neutral solution: [H + ] = [OH - ] = 1.0 x M Acid solution: [H + ] > [OH - ]; [H + ] > 1.0 x M Basic solution: [H + ] 1.0 x M

Sample Problem Example 15.4 A sample of tap water has a [H+] = 2.8 x M. What is the [OH-]?

The pH Scale A common way to express acidity and basicity is with pH (the “power of hydrogen”) pH = - log [H 3 O + ] In a neutral solution, [HO + ] = [OH - ] = 1.00 x at 25 o C [H 3 O + ] = [OH - ] = 1.00 x at 25 o C pH = -log (1.00 x ) = - (-7) = 7

The pH Scale: pH = - log [H 3 O + ]

Size of pH Basic solution pH > 7 Basic solution pH > 7 Neutral pH = 7 Neutral pH = 7 Acidic solutionpH < 7 Acidic solutionpH < 7 Big number = Basic

pH of Common Substances

pOH Definition - log [OH - ] pOH = - log [OH - ] pH + pOH = 14.0

The pH Scale

Sample Problem Example Calculate the pH, pOH and [OH-] of an acid solution whose [H+] is 1.8 x M. Strategy: 1. Calculate pH using formula and [H + ]. 2. Calculate pOH from pH. 3. Can calculate [OH-] from pOH or by using K w and [H + ].

If the pH of Coke is 3.12, it is acidic. Because pH = - log [HO + ] then Because pH = - log [H 3 O + ] then log [HO + ] = - pH log [H 3 O + ] = - pH Take antilog and get [HO + ] = 10 -pH [H 3 O + ] = 10 -pH [HO + ] = = 7.6 x M [H 3 O + ] = = 7.6 x M pH to [H + ] Calculations

Strong Acids Strong Acids dissociate completely in aqueous solution: E.g. HCl  H + + Cl - In a strong acid, [H+] can be calculated from the molarity of the acid. 2.0 M 2.0 M 2.0 M

Sample Problem Example Calculate the [H+], pH, and [OH-] of a 0.15M solution of the strong acid, HNO 3. Strategy: 1. Write the disassociation equation and calculate [H+]. 2. Calculate pH from [H+]. 3. Use Kw to find [OH-].

Strong Bases Strong Bases dissociate completely in aqueous solution: E.g. Sr(OH) 2  Sr OH - In a strong base, [OH-] can be calculated from the molarity of the base. 0.5 M 0.5 M 1.0 M

Sample Problem Example State the pH, pOH, [H+], and [OH-] of a solution made by dissolving mol of Ba(OH)2 - a strong base - in 5.00 L of water. Strategy: 1. Calculate molarity of Ba(OH) Write disassociation equation; calculate molarity of OH-. 3. Calculate pOH from [OH-]. 4. Calculate pH from pOH. 5. Calculate [OH-] from K w.

Neutralization Reactions and Titration for Strong Acids and Bases 1. Definitions titration titrant indicator endpoint equivalence point

Acid-Base Indicators and Acid-Base Titrations for Strong Acids and Bases 1. Neutralization reactions Neutralization is the (usually complete) reaction of an ACID + BASE The products of this neutralization are a “salt” (ionic compound) + water It is a double replacement reaction. Example: HCl + NaOH  H 2 O + NaCl

Definitions A titration is a carefully controlled neutralization reaction. A buret is used in a titration. Why? To determine the concentration of an unknown acid or base.

Definitions The titrant is the substance of known concentration used to determine the unknown concentration of the other substance. An indicator--substance that changes color at a certain pH—is added to tell us when the neutralization is complete. Example: Phenolphthalein undergoes a color change between pH 8 and 10 clear in acid Light pink in neutral Dark pink in base

Definitions The equivalence point, is the point in the titration where the neutralization is complete: [H 3 O + ] = [OH - ] The endpoint is the point where the indicator changes color. If indicator chosen correctly, two points are identical!

Acid–Base Titration

Strong acid and base titration Chemical reaction: ACID + BASE  SALT + WATER HCl + NaOH  NaCl + H 2 O A double replacement reaction!

 In the reaction above, the HCl, NaOH, and NaCl all are strong electrolytes and dissociate completely.  The actual reaction occurs between ions. Acid–Base Reactions: Example HCl + NaOH  H 2 O + NaCl H + + Cl – + Na + + OH –  H 2 O + Na + + Cl – H + + OH –  H 2 O A net ionic equation shows the species actually involved in the reaction. Na + and Cl – are spectator ions.

Strong acid and base titration Short-cut equation: ( # H + )M A V A = ( # OH - ) M B V B Where #H + and #OH - are obtained from the formula Volumes can be in mL or L, but must be same units on both sides of equation!

Procedure for Titration Total mols of H + from the acid Total mols of OH - from the base = = *Remember: (conc)(vol in L) = moles At the equivalence point (end point):

Sample Problem Example How many milliliters of M HCl are required to titrate mL of M KOH?

Sample Problem Example How many milliliters of M Ba(OH) 2 are required to titrate mL of M HCl?

 Oxidation: Loss of electrons  Reduction: Gain of electrons  Both oxidation and reduction must occur simultaneously. A species that loses electrons must lose them to something else (something that gains them). A species that gains electrons must gain them from something else (something that loses them).  Historical: “oxidation” used to mean “combines with oxygen”; the modern definition is much more general. Reactions Involving Oxidation and Reduction

 An oxidation number is the charge on an ion, or a hypothetical charge assigned to an atom in a molecule or polyatomic ion.  Examples: in NaCl, the oxidation number of Na is +1, that of Cl is – 1 (the actual charge).  In CO 2 (a molecular compound, no ions) the oxidation number of oxygen is – 2, because oxygen as an ion would be expected to have a 2 – charge.  The carbon in CO 2 has an oxidation number of + 4 (Why?) Oxidation Numbers

In a redox reaction, the oxidation number of a species changes during the reaction. Oxidation occurs when the oxidation number increases (species loses electrons). Reduction occurs when the oxidation number decreases (species gains electrons). If any species is oxidized or reduced in a reaction, that reaction is a redox reaction. The two processes occur at the same time Identifying Oxidation–Reduction Reactions

A Redox Reaction: Mg + Cu 2+  Mg 2+ + Cu … the products are Cu metal and Mg 2+ ions. Electrons are transferred from Mg metal to Cu 2+ ions and …

 An oxidizing agent causes another substance to be oxidized.  The oxidizing agent is reduced.  A reducing agent causes another substance to be reduced.  The reducing agent is oxidized. Mg + Cu 2+  Mg 2+ + Cu What is the oxidizing agent? What is the reducing agent? Oxidizing and Reducing Agents

Examples of redox reactions:  Displacement of an element by another element  Combustion  Extraction of metal from ores  Metabolic reactions in living organisms  Manufacturing chemicals  Mg(s) + Ca 2+( aq)  Mg 2+ (aq) + Ca(s) Ox: Mg(s)  Mg 2+ (aq) +2e - Red: Cu 2+ (aq) + 2e -  Cu(s)

 Redox equations must be balanced according to both mass and electric charge.  For now, our main goals will be to: Identify oxidation–reduction reactions. Balance certain simple redox equations by inspection. Recognize, in all cases, whether a redox equation is properly balanced. Oxidation–Reduction Equations

Neutral Species For an isolated atom, a molecule, or a formula unit—the oxidation numbers is 0. ex: Cl 2, Fe Monatomic Ions (Groups 1, 2, 17) Group 1 elements all have an oxidation number of +1 Group 2 elements all have an oxidation number of +2. Fluorine always has an oxidation number or -1. Oxygen In most compounds, oxygen has an oxidation number of –2. Rules for Assigning Oxidation Numbers

Sum of oxidation numbers for neutral and charged species: Hydrogen  +1 when bonded to a nonmetal, -1 when bonded to a metal  The sum of the oxidation numbers of all atoms (or ions) in a neutral compound = 0.  The sum of the oxidation numbers of all atoms in a polyatomic ion = charge on the polyatomic ion. ex: CaCl 2

Example: Assign oxidation numbers to the elements in the following species: CaC 2 O 4 Cr 2 O 7 2- N 2 ON 2 O 4 ClO 1- ClO 4 1-

6. Identifying redox reactions  Look for changes in oxidation numbers Al (s) + Fe 2 O 3 (s)  2 Fe (l) + Al 2 O

Example 4.7 Identify the element reduced, the element oxidized, the reducing agent and the oxidizing agent : A) Fe 2+ (aq) + Cr 2 O 7 2- (aq) + H + (aq)  Fe 3+ (aq) + Cr 3+ (aq) + H 2 O(l) B) 3 Cl 2 (g) + 2 Cr(OH) 3 (aq) + 10 OH -  2CrO 4 - (aq) + 6Cl - (aq) + 8H 2 O(l)

Applications of Oxidation and Reduction  Analytical Chemistry  KMnO 4 is the most commonly used oxidizing agent in chemical laboratories.  5Fe 2+ (aq) + MnO 4- (aq) + 8H + (aq)  5Fe 3+ (aq) + Mn 2+ (aq) + 4H 2 O (l)

Oxidation and Reduction in Organic Chem. Potassium dichromate Ethanol Initially the solution turns the orange of Cr 2 O 7 2– After a while the alcohol is oxidized to a ketone, and the Cr 2 O 7 2– is reduced to Cr 3+

 In industry: to produce iron, steel, other metals, and consumer goods.  In foods and nutrition: redox reactions “burn” the foods we eat; antioxidants react with undesirable free radicals. Applications of Oxidation and Reduction

Everyday life: to clean (bleach) our clothes, sanitize our swimming pools (“chlorine”), and to whiten teeth (peroxide).