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Chapter 4Aqueous Reactions and Solution Stoichiometry

John D. Bookstaver

St. Charles Community College

St. Peters, MO

2006, Prentice Hall, Inc.

Chemistry, The Central Science, 10th editionTheodore L. Brown; H. Eugene LeMay, Jr.;

and Bruce E. Bursten

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Chapter 4 HW

• Visualizing: 1,3,5,7,9,10

• Electrolytes: 11,13,15,17

• Pp rx and net ionic eq: 19,23,25

• Acid-base rx: 29,35,37,39

• Redox: 49,51,53

• Molarity: 61,65,67,71

• Titrations : 77,81,85

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Aqueous Solutions

Water is the dissolving medium, or solvent.

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Figure 4.1 The Water Molecule is Polar

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Some Properties of Water

Water is “bent” or V-shaped. The O-H bonds are covalent. Water is a polar molecule. Hydration occurs when salts

dissolve in water.

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Solutions:

• Homogeneous mixtures of two or more pure substances.

• The solvent is present in greatest abundance.

• All other substances are solutes.

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A Solute

dissolves in water (or other “solvent”)

changes phase (if different from the solvent)

is present in lesser amount (if the same phase as the solvent)

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A Solvent

retains its phase (if different from the solute)

is present in greater amount (if the same phase as the solute)

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Polar Water Molecules Interact with the Positive and Negative Ions of a

Salt

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BaCI2 Dissolving

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Dissociation

• When an ionic substance dissolves in water, the solvent pulls the individual ions from the crystal and solvates them.

• This process is called dissociation.

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Electrolytes

• Substances that dissociate into ions when dissolved in water.

• A nonelectrolyte may dissolve in water, but it does not dissociate into ions when it does so.

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Electrolytes and Nonelectrolytes

Soluble ionic compounds tend to be electrolytes.

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Electrolytes and Nonelectrolytes

Molecular compounds tend to be nonelectrolytes, except for acids and bases.

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Electrolytes

• A strong electrolyte dissociates completely when dissolved in water.

• A weak electrolyte only dissociates partially when dissolved in water.

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Strong Electrolytes Are…

• Strong acids

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Strong Electrolytes Are…

• Strong acids• Strong bases

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Strong Electrolytes Are…

• Strong acids• Strong bases• Soluble ionic salts

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Electrolytes

Strong - conduct current efficiently

NaCl, HNO3

Weak - conduct only a small current

vinegar, tap water

Non - no current flows

pure water, sugar solution

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Strong Electrolytes

1. Soluble salts- eg. NaCl, Pb(ClO3)2

2. Strong acids- eg. HCl, H2SO4

3. Strong bases- eg. NaOH, KOH

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Weak or Nonelectrolytes

1. Insoluble or only slightly soluble salts. Eg. AgCl, CaSO4

2. Weak Acids- eg. CH3COOH, HF

3. Weak Bases- eg. NH3, amines

4. Water

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September 23-Section 4.2Precipitation Reactions

*SOLUBILITY

*PRECIPITATION REACTIONS

*DOUBLE REPLACEMENT (METATHESIS –EXCHANGE ) REACTION

• NET IONIC EQUATIONS

• SPECTATOR IONS

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Precipitation Reactions

When one mixes ions that form compounds that are insoluble (as could be predicted by the solubility guidelines), a precipitate is formed.

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SOLUBILITY

• The amount of a substance that can be dissolved in a given quantity of solvent at a given temperature.

Example

Solubility of KNO3

65 g/100 ml H2O of KNO3 at 40 C

*It is determined experimentally.

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Solubility Rules – Used to determine what reaction occurs, if any:

1. Separate all ions.

2. Determine all possible compounds formed.

3. Determine which, if any, of the compounds will precipitate.

4. Write the appropriate chemical equation.

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Rules for Solubility of Ionic Compounds in Water1. All common Group 1 and ammonium salts are

soluble.2. All nitrates, chlorates, and acetates are soluble except silver

acetate.3. All halide salts (except fluorides) are soluble except those

of silver, mercury (I), and lead.4. All sulfates are soluble except those of silver, mercury (I or

II), lead, calcium, strontium, and barium.5. Calcium, strontium, and barium hydroxides (as well as

group 1 hydroxides) are soluble; other hydroxides generally are not.

6. Most other ionic compounds are generally insoluble.

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Double Replacement (Metathesis) Reactions• Metathesis or double displacement/replacement

reactions involve swapping ions in solution: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, or– an insoluble gas is formed.

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Molecular Equation

The molecular equation lists the reactants and products in their molecular form.

AgNO3 (aq) + KCl (aq) AgCl (s) + KNO3 (aq)

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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)

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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)

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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)

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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)

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Writing Net Ionic Equations

1. Write a balanced molecular equation.

2. Dissociate all strong electrolytes.

3. Cross out anything that remains unchanged from the left side to the right side of the equation.

4. Write the net ionic equation with the species that remain.

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Examples: Write the molecular, complete ionic and net ionic equations for each of the following reactions.

1. Aqueous solutions of sodium sulfide and calcium nitrate are mixed.

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2. Aqueous solutions of barium chloride and potassium sulfate are mixed.

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3. Aqueous solutions of silver nitrate and ammonium chloride are mixed.

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Homework

• Page 157 answer: 4.7 to 4.9

• Page 158 4.19 to 4.28 odd

• Read Section 4.3

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September 24- Section 4.3 Acid – Bases Reactions

• Strong and weak Acid and Bases

• Neutralization Reactions

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Classify as strong, weak or nonelectrolyte

CaCl2

HNO3

C2H5OH

HC2H3O2

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Rank the following solutions in order of increasing electrical

conductivity

• Ca(NO3)2

• C6 H12 O6

• NaC2 H3O2

• HC2 H3O2

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Acids:

• Substances that increase the concentration of H+ when dissolved in water (Arrhenius).

• Proton donors (Brønsted–Lowry).

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Acids

There are only seven strong acids:• Hydrochloric (HCl)• Hydrobromic (HBr)• Hydroiodic (HI)

• Nitric (HNO3)

• Sulfuric (H2SO4)

• Chloric (HClO3)

• Perchloric (HClO4)

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Acids• Dissociation = pre-formed ions in solid move

apart in solution.• Ionization = neutral substance forms ions in

solution.• Acids = substances that ionize to form H+ in

solution (e.g. HCl, HNO3, CH3CO2H, lemon, lime, vitamin C). Proton donors.

• Acids with one acidic proton are called monoprotic (e.g., HCl).

• Acids with two acidic protons are called diprotic (e.g., H2SO4).

• Acids with many acidic protons are called polyprotic.

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Bases:

• Substances that increase the concentration of OH− when dissolved in water (Arrhenius).

• Proton acceptors (Brønsted–Lowry).

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Bases• Bases = substances that react with the H+ ions

formed by acids (e.g. NH3, Drano™, Milk of Magnesia™). Proton acceptors (Bronsted-Lowry).

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Strong and Weak Acids and Bases• Strong acids and bases are strong electrolytes.

– They are completely ionized in solution.• Strong acids are HCl, HBr, HI, HNO3, H2SO4,

HClO3 and HClO4.• Strong bases are group 1 hydroxides and

soluble group 2 hydroxides.• Weak acids and bases are weak electrolytes.

– They are partially ionized in solution.

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Acid-Base Reactions

In an acid-base reaction, the acid donates a proton (H+) to the base.

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Neutralization Reactions and Salts• Neutralization occurs when a solution of an acid

and a base are mixed:HCl(aq) + NaOH(aq) H2O(l) + NaCl(aq)

• We form a salt (NaCl) and water.• Salt = ionic compound whose cation comes from a

base and anion from an acid.• Neutralization between acid and metal hydroxide

produces water and a salt.• In net ionic equations, strong acids and bases are

written dissociated, while weak acids and bases are written associated (non-dissociated)

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Neutralization Reactions

Observe the reaction between Milk of Magnesia, Mg(OH)2, and HCl.

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September 27 – Section 4.4Oxidation-Reduction Reactions

• Acid-Base reactions with gas formation• Redox concept LEO GER• Oxidation number • Oxidizing-Reducing Agents• Single replacement reactions (oxidation of

metals by acids and salts)- Displacement reactions

• Activity Series• Balancing redox in acid and alkaline media.

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Metathasis Reactions that produce gas

GAS REACTANTS

H2S (g) Any sulfide plus any acid

CO2 (g) Any carbonate plus acid

SO2 (g) Any sulfite plus acid

NH3 (g) Any ammonium salt plus strong hydroxide and heat

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Acid-Base Reactions with Gas Formation

• Sulfide ions can react with H+ producing H2S gas.

2HCl(aq) + Na2S(aq) H2S(g) + 2NaCl(aq)

2H+(aq) + S2-(aq) H2S(g)

Na2S (aq) + H2SO4 (aq) Na2SO4 (aq) + H2S (g)

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Carbonate ions produce CO2(g)and H2O

• HCl(aq) + NaHCO3(aq) NaCl(aq) + H2O(l) + CO2(g)

CaCO3 (s) + HCl (aq) CaCl2 (aq) + CO2

(g) + H2O (l)

NaHCO3 (aq) + HBr (aq) NaBr (aq) + CO2 (g) + H2O(l)

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Sulfite ions produce SO2 and H2O

SrSO3 (s) + 2 HI (aq) SrI2 (aq) + SO2 (g) + H2O (l)

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Ammonium salts and soluble bases produce NH3 when solution

is warmed

• NH4Cl(aq) + NaOH (aq) --> NH3 (g) + H2O (l) + NaCl (aq)

• Theorically NH4OH is produced but is unstable and decomposes into ammonia and water.

• NH4OH (aq) NH3 (g) + H2O (l)

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• LOSING

• ELECTRONS

• OXIDATION

• O.N. increases in oxidation

• GAINING

• ELECTRONS

• REDUCTION

• O.N. decreases in reduction

LEO GER

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Oxidation-Reduction Reactions

• An oxidation occurs when an atom or ion loses electrons.

• A reduction occurs when an atom or ion gains electrons.

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Oxidation-Reduction Reactions

One cannot occur without the other.

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Oxidizing and Reducing Agents

• Oxidizing agents cause oxidation to occur. How?

• Reducing agents cause reduction to occur. How?

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A Summary of an Oxidation-Reduction Process

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September 28

• How to determine the Oxidation Number

of elements in compounds.

• How to predict if a reaction will ocurr using the reactivity series

• Single replacement reactions

• Balancing redox reactions

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Oxidation Numbers

To determine if an oxidation-reduction reaction has occurred, we assign an oxidation number to each element in a neutral compound or charged entity.

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Oxidation Numbers• Oxidation numbers are assigned by a series of

rules:1. If the atom is in its elemental form, the

oxidation number is zero. e.g., Cl2, H2, P4.2. For a monoatomic ion, the charge on the ion is

the oxidation state.3. Elements of Group I have always O.N. = +14. Elements of Group II have alwas O.N.= +2

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5. Oxidation number of O is usually –2. The peroxide ion, O2

2-, has oxygen with an oxidation number of –1 (H2O2, Na2O2).

• Oxygen with F- has O.N. = +2

6. Oxidation number of H is +1 when bonded to nonmetals and –1 when bonded to metals in metal Hydrides

7. The oxidation number of F is –1

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Oxidation Numbers

• The sum of the oxidation numbers in a neutral compound is 0.

• The sum of the oxidation numbers in a polyatomic ion is the charge on the ion.

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Examples: Determine the oxidation numbers of each element in each of the following:

1. H2SO4

2. KMnO4

3. NO3-

4. C2H6

5. CH3OH

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Displacement ReactionsSingle Replacement

(always redox reactions!)

• In displacement reactions, ions oxidize an element.

• The ions, then, are reduced.

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Displacement Reactions

In this reaction,

silver ions oxidize

copper metal.

Cu (s) + 2 Ag+ (aq) Cu2+ (aq) + 2 Ag (s)

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Displacement Reactions

The reverse reaction,

however, does not

occur.

Cu2+ (aq) + 2 Ag (s) Cu (s) + 2 Ag+ (aq) x

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Oxidation of Metals by Acids and Salts• Metals are oxidized by acids to form salts:

Mg(s) +2HCl(aq) MgCl2(aq) + H2(g)

• During the reaction, 2H+(aq) is reduced to H2(g).

• Metals can also be oxidized by other salts:Fe(s) +Ni2+(aq) Fe2+(aq) + Ni(s)

• Notice that the Fe is oxidized to Fe2+ and the Ni2+ is reduced to Ni.

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Activity Series• Some metals are easily oxidized whereas others

are not.• Activity series: a list of metals arranged in

decreasing ease of oxidation.• The higher the metal on the activity series, the

more active that metal.• Any metal can be oxidized by the ions of

elements below it.

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Examples: Write complete ionic and net ionic equations for the following:

1. Aluminum metal is added to an aqueous solution of copper chloride.

2. Zinc metal is added to a solution of hydrobromic acid.

3. Chromium metal is placed in a solution of potassium nitrate.

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Balancing by Half-Reaction Method

1. Write separate reduction, oxidation reactions.

2. For each half-reaction:

Balance elements (except H, O)

Balance O using H2O

Balance H using H+

Balance charge using electrons

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Balancing by Half-Reaction Method (continued)

3. If necessary, multiply by integer to equalize electron count.

4. Add half-reactions.

5. Check that elements and charges are balanced.

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Example

Balance the following oxidation reduction using the half-reaction method

MnO2 + Cl- Mn2+ + Cl2

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Half-Reaction Method - Balancing in Base

1. Balance as in acid.

2. Add OH that equals H+ ions (both sides!)

3. Form water by combining H+, OH.

4. Check elements and charges for balance.

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Example

Balance the following reaction which occurs in alkaline solution.

Ag (s) + CN-(aq) + O2(g) Ag(CN)2-(aq)

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September 29– Section 4.5Concentration of solutions

• Molarity Concept and problems• Find moles in a volume of solution• How to prepare a solution of a given

molarity• Conversions volume to moles, to mass

etc.• Expressing concentration of electrolytes• DILUTIONS

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NEXT WEEK LAB 1

• BOOKS ARE IN – NEED TO PURCHASE LAB BOOK BY WEDNESDAY –

• EQUATIONS BOOK ARE IN• STILL WAITING FOR TEXTBOOKS!

• HW – 61, 65, 67, 71 • TEST ON CH 1TO 4 ON MONDAY

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Molarity• Two solutions can contain the same

compounds but be quite different because the proportions of those compounds are different.

• Molarity is one way to measure the concentration of a solution.

moles of solute

volume of solution in litersMolarity (M) =

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Examples:

1. Calculate the molarity of a solution prepared by dissolving 3.89 g of sodium sulfate in enough water to prepare 250. mL of solution.

Step 1 Convert 3.89 g Na2SO4 to mol

FM = 142 g/mol 3.89g/142 g/mol= 0.027 mol

Step 2 Set up ratio

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1. How many grams of magnesium chlorate are required to prepare 100.0 mL of a 0.500 M solution?

Step 1 : Find the # mol in V of solution

Step 2: Convert mass to mol

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• How many g of KMnO4 are needed to prepare 100 ml of a 2M solution?

• Step 1: Find the # of mol

• Step 2: Convert mol to mass

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Mixing a Solution

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Dilution

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Dilution• We recognize that the number of moles are the

same in dilute and concentrated solutions.• So:

MdiluteVdilute = moles = MconcentratedVconcentrated

M1V1 = M2V2

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• Find the molarity of the solution if 1 ml of the solution are diluted to 500 ml

• Step 1: M V = M V

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Example: What volume of 18.0 M H2SO4 solution is required to prepare 1000.0 mL of 1.00 M H2SO4?

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October 1 -Section 4.6Solution Stoichiometry and

Chemical analysis

• Titration

• Problems

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Solution Stoichiometry and Chemical Analysis

• There are two different types of units: – laboratory units (macroscopic units: measure

in lab);– chemical units (microscopic units: relate to

moles).• Always convert the laboratory units into

chemical units first.– Grams are converted to moles using molar

mass.– Volume or molarity are converted into moles

using M = mol/L.

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•Use the stoichiometric coefficients to move between reactants and product.

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Titrations• Suppose we know the molarity of a NaOH

solution and we want to find the molarity of an HCl solution.

• We know:– molarity of NaOH, volume of HCl.

• What do we want?– Molarity of HCl.

• What do we do?– Take a known volume of the HCl solution,

measure the mL of NaOH required to react completely with the HCl.

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Titrations• What do we get?

– Volume of NaOH. We know molarity of the NaOH, we can calculate moles of NaOH.

• Next step?– We also know HCl + NaOH NaCl + H2O.

Therefore, we know moles of HCl.• Can we finish?

– Knowing mol(HCl) and volume of HCl (20.0 mL above), we can calculate the molarity.

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Examples:

1. Standardization of NaOH: A primary standard called potassium hydrogen phthalate is used to determine the exact molarity of a solution of base. KHP is a monoprotic weak acid (MW = 204.22). In one example, 0.4977 g of KHP is dissolved in 100.0 mL of water. This solution requires 14.86 mL of a solution of NaOH to neutralize it. What is the molarity of the NaOH solution?

2. Solid Acid: An unknown solid acid is analyzed by titration. 1.75 g of the acid is weighed and dissolved in 150.0 mL of distilled water. This solution is titrated with 0.250 M KOH, and requires 19.07 mL of the KOH solution to reach the endpoint. What is the molar mass of the solid, assuming it is monoprotic?

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3. Unknown Solution: A solution of the weak base, Na2SO4, is analyzed by titration with 0.1000 M HCl. A 10.00 mL sample of the base requires 38.04 mL of HCl to reach the endpoint. What is the molarity of the sodium sulfate solution?

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TitrationThe analytical technique in which one can calculate the concentration of a solute in a solution.

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Key Titration Terms

Titrant - solution of known concentration used in titration

Analyte - substance being analyzed

Equivalence point - enough titrant added to react exactly with the analyte

Endpoint - the indicator changes color so you can tell the equivalence point has been reached.

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Titration