AP Chemistry

Chapter 4:Aqueous Reactions & Solution

Stoichiometry

Solutions

• Homogeneous mixtures• Have 2 parts:– Solute• Smaller quantities• Dissolved in the solvent

– Solvent• Larger quantities• Dissolves the solute(s)• Often water

Electrolytic Properties

• Electrolytes– Aqueous solutions that contain ions– Have ionic compounds or strong acids or bases as

solute– Will conduct electricity

• Nonelectrolytes– Aqueous solutions that do not contain ions– Have molecular compounds as solute

Ionic Compounds in Water

Ions dissociate from each other in water

Ions attracted to one another to form a crystal

lattice

Ionic Compounds in WaterWater is polar

Negative end of wateris attracted to positive ion

Each ion will besurrounded bywater molecules:Process is called solvation

Water prevents ionsfrom recombiningto form a crystal

Ionic Compounds in Water

Molecular Compounds in Water

Molecules remain in tact

Ions not usually formed

Strong acids are exceptions

Strong vs. Weak ElectrolytesWeak Electrolytes:• Exist in solution

almost entirely as molecules (a small fraction may be ions)

• Molecular compounds fall into this category

Strong Electrolytes:•Exist in solution completely (or nearly completely) as ions

•Essentially all soluble ionic compounds & some molecular (some acids) fall into this category

Strong vs. Weak ElectrolytesStrong Electrolyte Example:

HCl(aq) H+(aq) + Cl-(aq)

have no tendency to recombine to form HCl

Weak Electrolyte Example:

HC2H3O2(aq) H+(aq) + C2H3O2-(aq)

reaction is significant in both directions is said to be in chemical equilibrium

Soluble Ionic Compounds are Strong Electrolytes

(metal & nonmetal or contain ammonium ion)

Strong vs. Weak ElectrolyteIS NOT the same thing as solubility….compounds can be very soluble and be a weak electrolyte and vice versa….

POINTS

TO

REMEMBER

Sample Exercise 4.1The diagram represents an aqueous solution of one of the following compounds: MgCl2, KCl, or K2SO4. Which solution does it best represent?

2-

2-

2-

2-+ +

+

+

++

+

+

Precipitation Reactions

Lead nitrate + Potassium Iodide

Precipitation Reactions

• Result in the formation of an insoluble compound

• A precipitate is an insoluble compound formed by a reaction in solution

Solubility Rules

HAS CANHalides

Except:Hg2

2+

AgPb All Alkali

metals

Sulfates

Except: Ba2+, Sr2+, Hg22+, Ag, and Pb

All AmmoniumcompoundsAll C2H3O2

- compounds

All Nitrates

Hydroxides and Sulfides are soluble if

bonded with Ca2+, Ba2+, or Sr2+, alkali

metals or ammonium

Precipitation ExampleWill a precipitate form when solutions of Mg(NO3)2 and NaOH are mixed?

Sample Exercise 4.2Classify the following ionic compounds as soluble or insoluble in water (a) sodium carbonate (b) lead sulfate

Metathesis Reactions

• Cation and anion appear to “change partners”• Precipitation reactions conform to this pattern

AY + BX AX + BY

Completing and Balancing Metathesis:

1. Use chemical formulas of reactants to determine the ions that are present

2. Write the chemical formulas for the products by combining the cation from one reactant with the anion form the other reactant

3. Balance the equation

* If all products are soluble, we say NO REACTION has occurred.

Sample Exercise 4.3

(a) Predict the identity of the precipitate that forms when solutions of BaCl2 and K2SO4 are mixed.

(b) Write the balanced equation for this reaction.

Ionic Equations

• Molecular equations show the chemical formulas for the reactants and the products

• Complete ionic equations show the ions in solution

Pb(NO3)2 (aq) + 2KI (aq) PbI2 (s) + 2KNO3 (aq)

Pb2+ (aq) +2NO3- (aq) + 2K+ (aq) + 2I- (aq) PbI2 (s) +2K+ (aq) +2NO3

- (aq)

Ionic Equations

• Spectator ions– appear in both the reactants and the products in

the equation– play no part in the reaction– can be omitted (cancelled out)

Ionic Equations

• Net ionic equations show only the ions and molecules that are directly involved in the reaction

Pb2+ (aq) +2NO3- (aq) + 2K+ (aq) + 2I- (aq) PbI2 (s) +2K+ (aq) +2NO3

- (aq)

Pb2+ (aq) + 2I- (aq) PbI2 (s)

NET IONIC EQUATION:

Reminders

Charge is

conserved: Sum

of charges of

reactants must

equal sum of

charges of

products

If every ion is a spectator, no reaction occurs

To write net ionic equations:

1. Write a balanced molecular equation for the reaction

2. Rewrite to show ions that form in solution when each soluble, strong electrolyte dissociates into its component ions (Only strong electrolytes written in ionic form)

3. Identify and cancel the spectator ions

Sample Exercise 4.4Write the net ionic equation for the precipitation reaction that occurs when solutions of calcium chloride and sodium carbonate are mixed.

Acid – Base Reactions• Acids:• Substances that ionize in solutions to from H+ ions• H+ is simply a proton• Acids often called proton donors• Monoprotic acids• Have only one H • Examples: HCl, HNO3

• Diprotic acids• Ionized in 2 steps• Examples:H2SO4

H2SO4(aq) H+(aq) + HSO4-(aq)

HSO4-(aq) H+(aq) + SO4

2-(aq)

Acid-Base Reactions

• Bases– Substances that accept protons (H+)– Produce hydroxide ions (OH-) when dissolved in water– Include ionic hydroxide compounds• NaOH, MgOH, etc.

– NH3(ammonia) is also a base…it accepts H+ from water leaving OH-

• NH3(aq) + H2O(l) NH4+(aq) + OH-(aq)

Only a small fraction will form NH4+ ammonia is a

weak electrolyte

Acid-Base Reactions

• Acids and bases that are strong electrolytes are called strong acids and bases

• Reactivity can be determined by strength of acid or on the strength of the anion– Example: HF is a weak acid, but is very reactive

because of the F- ion

STRONG ACIDSH2SO4

HNO3

HClO3

HClO4

HCl

HBr

HI

STRONG BASESGroup I Metal

Hydroxides

Ca, Sr, & BaHydroxides

Sample Exercise 4.5The following diagrams represent aqueous solutions of three acids (HX, HY, HZ) with water molecules omitted for clarity. Rank them from strongest to weakest.

+

-

-

+

HX HY HZ

+

+

+

+

+

+

+

+--

-

-

--

-

-

-

--+

+

+

Type of Compound

StrongElectrolyte

Weak Electrolyte Nonelectrolyte

Ionic All None None

Molecular Strong Acids

Weak Acids & Bases

All others

Electrolytic Behavior of Soluble Compounds

Sample Exercise 4.6Classify each of the following substances as strong electrolyte, weak electrolyte, or nonelectrolyte: CaCl2, HNO3, C2H5OH (ethanol), HCHO2 (formic acid), KOH

Acids vs. Bases

• Taste sour• Turn blue litmus paper red

Acids

• Taste bitter• Turn red litmus paper blue

Bases

NeutralizationReaction between acidic and basic solutions

If the base is a metal hydroxide

Products are a salt and water

Compound whose cation comesfrom a base and anion comesfrom an acid

Neutralization Example

HCl(aq) + NaOH(aq) NaCl(aq) + H2O(l)

H+(aq) + Cl-(aq) + Na+(aq) + OH-(aq) Na+(aq) + Cl-(aq) + H2O(l)

H+(aq) + OH-(aq) H2O(l)

Molecular Equation:

Complete Ionic Equation:

Net Ionic Equation:

Sample Exercise 4.7(a) Write a balanced molecular equation for the reaction between aqueous solutions of acetic acid (HC2H3O2) and barium hydroxide [Ba(OH)2]. (b) Write the net ionic equation for this reaction

Gas Forming Reactions• Acids will form gases when they react with– Sulfides…H2S (g) forms

– Carbonates & bicarbonates…CO2 (g) forms• First to form is the unstable carbonic acid (H2CO3) which

decomposes to form carbon dioxide gas and water vapor

• Example: Alka Seltzer: NaHCO3 reacts with stomach acid

• Sodium bicarbonate & sodium carbonate are used to clean up acid spills

Oxidation-Reduction Reactions

• Reactions in which electrons are transferred between substances

• Oxidation– Loss of electrons– Substance becomes more positive

• Reduction– Gain of electrons– Substance becomes more negative

Oxidation – Reduction Reactions

• Oxidation of one substance is ALWAYS accompanied by the reduction of another

Oxidation Numbers

• Assigned to each atom in a neutral substance or in a charged species

• This is the charge of the monatomic ions or the hypothetical charge assigned to the atom assuming the electrons are completely held by one atom or another

Oxidation Numbers

• In redox reactions, the ox. #’s change from products to reactants

• Oxidation = increase in oxidation number

• Reduction = decrease in oxidation number

Assigning Oxidation Numbers…the Rules

1. Oxidation number is 0 for atoms in elemental form H2 H = 0 S8 S = 0 P4 P = 0

2. Oxidation number = the charge of monatomic ionsTo distinguish between oxidation number and

actual charge: Ca2+ oxidation number = +2

3. Nonmetals (usually negative ox. nos.)a) Oxygen is usually -2 (both ionic & molecular)

Exception: in peroxide, each oxygen is -1b) Hydrogen is +1 when bonded with nonmetals and -1

when bonded with metalsc) Fluorine is ALWAYS -1d) Other halogens are usually -1

Exception: When bonded with oxygen, they are +

4. Sum of oxidation numbers in a compound is always 0: in a polyatomic ion the sum = the charge of the ion

Assigning Oxidation Numbers…the Rules

Sample Exercise 4.8Determine the oxidation number of sulfur in each of the following: (a) H2S (b) S8 (c) SCl2 (d) Na2SO3 (e) SO4

2-

Oxidation of Metals by Acids & Salts

• When a metal is oxidized, it appears to be “eaten away” as it reacts

• Called displacement reactions because the ion in solution is displaced (replaced) by the oxidation of an element

Oxidation of Metals by Acids and Salts• Many metals will react with acids to form hydrogen

gas and a salt

• For example: Mg(s) + 2 HCl(aq) MgCl2(aq) + H2(g)

ox. #’s 0 +1 -1 +2 -1 0

Mg is oxidizedH is reduced

Cl is a spectator

Oxidation of Metals by Acids & Salts

Fe (s) + Ni(NO3)2 (aq) Fe(NO3)2 (aq) + Ni (s)

One more time:

If something is oxidized….Something else must

be reduced!

Sample Exercise 4.9Write the balanced molecular and net ionic equations for the reaction of aluminum with hydrobromic acid.

The Activity Series

• A list of metals arranged in order of decreasing oxidation

• Can be used to predict reactions between metals and metals salts or acids

• Any metal on the list can be oxidized by ions of elements below it

• Most easily oxidized = most reactive is at the top of the list

• List is on page 143 of your textbook

Sample Exercise 4.10Will an aqueous solution of iron (II) chloride oxidize magnesium metal? If so, write the balanced molecular and net ionic equations for the reaction.

Concentrations of Solutions

• Designates the amount of solute dissolved in a given amount of solvent or solution

• Molarity– Quantitative expression of concentration– M = mol/L – Moles of solute/ Liter of solution

Sample Exercise 4.11Calculate the molarity of a solution made by dissolving 23.4 g of sodium sulfate in enough water to form 125 mL of solution.

Expression of Concentration

• Can be made in terms of the compound:– 1.0 M Na2SO4

• Can be made in terms of the ions:– 2.0 Na+ and 1.0 M SO4

2-

Sample Exercise 4.12What are the molar concentrations of each of the ions present in a 0.025 M aqueous solution of calcium nitrate?

Sample Exercise 4.13How many grams of Na2SO4 are required to make 0.350 L of 0.500 M Na2SO4?

Dilution

• Preparation of solutions of lower concentrations from solutions of higher concentrations by adding water

• The addition of water (solvent) does not change the # of moles of solute

• Equation:M1V1 = M2V2

Sample Exercise 4.14How many milliliters of 3.0 M H2SO4 are needed to make 450 milliliters of 0.10 M H2SO4?

Sample Exercise 4.15How many grams of Ca(OH)2 are needed to neutralize 25.0 mL of 0.100 M HNO3?

Titration

• Method used to determine the amount of solute in a solution

• Uses the addition of a solution of known concentration called a standard solution

• Can involve different types of reactions– Acid base– Redox– precipitation

Titration• Makes use of equivalence point– The point where stoichiometrically equivalent quantities

are brought together• Use indicators– Indicates when equivalence point is reached– Changes color in response to change in pH– Ex. Phenolphthalein is colorless in acid, but pink in a

base– End point of titration is when solution changes color– Must choose an indicator that’s end point coincides with

the equivalence point

Sample Exercise 4.16The quantity of Cl- in a municipal water supply is determined by titrating the sample with Ag+. The reaction taking place during the reaction is: Ag+(aq) + Cl-(aq) AgCl(s)The end point in this type of titration is marked by the color change of a special type of indicator. (a) How many grams of chloride ion are in a sample of the water if 20.2 mL of 0.100 M Ag+ is needed to react with all the chloride in the sample? (b) If the sample has a mass of 10.0 grams, what percent Cl- does it contain?

Sample Exercise 4.17One commercial method used to peel potatoes is to soak them in a solution of NaOH for a short time, remove them from the NaOH, and spray off the peel. The concentration of NaOH is normally in the range of 3 to 6 M. The NaOH is analyzed periodically . In one such analysis, 45.7 mL of 0.500 M H2SO4 is required to neutralize a 20.0 mL sample of NaOH solution. What is the concentration of the NaOH solution?

Integrative SampleA sample of 70.5 mg of potassium phosphate is added to 15.0 mL of 0.050 M silver nitrate, resulting in the formation of a precipitate. (a) Write the molecular equation for the reaction (b) What is the limiting reactant in the reaction? (c) Calculate the theoretical yield, in grams, of the precipitate that forms.

That’s all Folks!