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Chapter 4. Aqueous Reactions and Solution Stoichiometry Pg 105. Aqueous Solutions. -Aqueous Solutions are solutions that have water as the dissolving medium. -Many reactions contain substances that have been dissolved in water, making them aqueous solutions.

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Chapter 4

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Chapter 4

Chapter 4

Aqueous Reactions and Solution Stoichiometry

Pg 105

Aqueous solutions

Aqueous Solutions

-Aqueous Solutions are solutions that have water as the dissolving medium.

-Many reactions contain substances that have been dissolved in water, making them aqueous solutions.

3 main major chemical reaction types involving aqueous solutions

3 Main Major Chemical Reaction Types Involving Aqueous Solutions

  • Precipitation Reactions

  • Acid-Base Reactions

  • Redox Reactions

4 1 general properties of aqueous solutions

4.1 General Properties of Aqueous Solutions

  • Solutions are homogeneous mixtures

  • Usually has more solvent than solute.

  • Solute is the substance being dissolved in the solvent

Electrolytic properties

Electrolytic Properties

  • Pure water is a bad conductor

  • The presence of Ions in water makes it into a good conductor

  • Aqueous solution that conduct electricity such as NaCl(aq) or other ionic compounds are electrolyte.

  • Solutions that do not form ions like sucrose and other molecular compounds are nonelectrolytes.

Ionic compounds in water

Ionic Compounds in Water

  • Ionic compounds dissolve in water dissociating into component ions (ex. NaCl -> Na+&Cl-)

  • The polar nature of water makes it a very effective solvent

  • The polarity helps prevent anions and cations from rejoining.

Molecular compounds in water

Molecular Compounds in Water

  • Structure usually remains unchanged, they usually do not form ions

  • Acids and a few other compounds like ammonia react with water forming ions making an electrolyte.

  • Ex. HCl make H+ and Cl- ions

Strong weak electrolytes

Strong & Weak Electrolytes

  • Strong Electrolytes = Most ionic compounds and a few molecular compounds.

  • Weak Electrolytes = Molecular compounds that produce few ions when dissolved

  • If the chemical reaction goes both ways, breaking into ions, and recombining, than the substance is a weak electrolyte.

Hcl aq h cl one arrow means strong electrolyte

HCl(aq) --> H+ + Cl-One arrow means strong electrolyte

HC2H3O29(aq) <--> H+ + C2H3O2-

Double arrow means weak electrolyte

4 2 precipitation reactions

4.2 Precipitation Reactions

  • Precipitation Reaction = Reactions that result in the formation of an insoluble product.

  • Precipitate = Insoluble solid formed by a reaction in a solution

Solubility guidelines for ionic compounds

Solubility Guidelines for Ionic Compounds

  • Solubility = Amount of substance that can be dissolved in given amount of solvent

  • If less than .01 mol dissolves in a liter, substance is insoluble. In these substances intermolecular attraction is stronger than the waters polarity.

  • Table 4.1 pg 111 (Solubility Guidelines for Common Ionic Compounds in Water)

  • All ionic compounds with 1A elements or ammonia ions are soluble in water.

Is sodium carbonate soluble na 2 co 3

Is Sodium Carbonate Soluble(Na2CO3)

  • Yes. Carbonate is usually insoluble, but when paired with a 1A element, Sodium, the compound becomes soluble.

Exchange metathesis reactions

Exchange (Metathesis) Reactions

  • Exchange or Metathesis Reaction = AX+BY --> AY+BX

  • Precipitation and Acid Base Reactions conform to this pattern

What precipitate forms when bacl 2 and k 2 so 4 are mixed

What precipitate forms when BaCl2 and K2SO4 are mixed?

  • BaSO4, SO42- is soluble but Ba2+ is not

Ionic equations

Ionic Equations

  • Molecular Equation = complete chemical formulas of reactants and products

  • Complete Ionic Equation = All Soluble strong Electrolytes are shown as ions

  • Spectator ions = ions that are present in the same form on both product and reactant side. These are dropped out to form a Net Ionic Equation.

Steps to write a net ionic equation

Steps to Write a Net Ionic Equation

  • Write a balanced Molecular Equation

  • Rewrite to show ions that are formed during dissociation or ionization, only the strong electrolytes are written in ionic form

  • Cancel spectator ions on both sides

Write the net ionic equation for the mixing of cacl 2 and na 2 co 3

Write the net ionic equation for the mixing of CaCl2 and Na2CO3

  • CaCl2(aq) + Na2CO3(aq) --> CaCO3(S) + 2NaCl(aq)

  • Ca2++ 2Cl- + 2Na+ + CO32--->CaCO3(s) + 2Na+ + 2Cl-

  • Ca2+(aq) + CO32-(aq)-->CaCO3(s)

Acid base reactions

Acid-Base Reactions

  • Acids and Bases are common Electrolytes

  • Are some of the most common compounds we encounter



  • Substances that ionize in aqueous solutions upping H+ concentration

  • Protic refers to amount of H+ ions ionizing. Monoprotic = 1, Diprotic = 2.

  • Diprotic Acid ionization occurs in two steps, One hydrogen is separated at a time.



  • Substances that accept H+ ions, or increases OH- concentration.

  • Does not need to have an OH- ion, if accepts H+ like NH3 (ammonia is a weak electrolyte)

Strong and weak acids and bases

Strong and Weak Acids and Bases

  • Strong Acids and Bases are strong electrolytes that completely ionize in solutions

  • Weak Acids and Bases are electrolytes that partly ionize in solutions

  • Table 4.2 pg 115 (Common Strong Acids and Bases)

Identifying strong and weak electrolytes

Identifying Strong and Weak Electrolytes

  • Is the compound ionic, yes -> probably strong electrolyte

  • Not ionic, is it an acid

  • Yes, is an acid, if strong, is a strong electrolyte if weak, is a weak electrolyte.

  • Not an acid, is it NH3 or another molecular base, yes -> weak base, no -> probably nonelectrolyte

Classify hno 3 as a strong weak or non electrolyte

Classify HNO3 as a strong, weak, or non Electrolyte

  • Strong

  • HNO3 is a strong acid making it a strong electrolyte.

Neutralization reactions and salts

Neutralization Reactions and Salts

  • When acid and base react together, it is a neutralization reaction.

  • These reactions form a salt and a water.

  • Salt = any ionic compound whose cation comes form a base and anion comes from an acid.

Write the net ionic equation for hc 2 h 3 o 2 and ba oh 2

Write the net ionic equation for HC2H3O2 and Ba(OH)2

HC2H3O2(aq)+OH-(aq)--> H2O(l) + C2H3O2-(aq)

Acid base reaction with gas formation

Acid-Base Reaction with Gas Formation

  • The sulfide ion and carbonate ion react with acids to form gases

  • 2HCl(aq) + Na2S(aq) --> H2S(g)+2NaCl(aq)

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

4 4 oxidation reduction reactions

4.4 Oxidation - Reduction Reactions

Oxidation and reduction

Oxidation and Reduction

  • Metals undergoing erosion are losing electrons and forming cations

  • Loss of electrons is known as oxidization

  • The gain of electrons by a substance is called reduction

Oxidation numbers

Oxidation Numbers

  • Oxidization number is the actual charge of the of the atom.

  • In elemental form the Oxidization number is 0

  • Oxidization of monatomic ions equals the charge

  • Nonmetals are usually negative, Oxygen usually is -2, Hydrogen is +1, Fluorine is -1.

  • Sum of oxidation numbers in neutral compound is 0 or equal to the charge in a polyatomic ion.

Determine the oxidation stat of sulfur in h 2 s

Determine the oxidation stat of sulfur in H2S

  • -2, Hydrogen is always +1 2H = +2 so S = -2 so that sum of oxidation numbers = 0

Oxidation of metals by acids and salts

Oxidation of Metals by Acids and Salts

  • Displacement Reactions = ion is solution is displaced or replaced through the oxidation of an element.

  • A+BX-->AX+B

Write the net ionic equation for the reaction of aluminum and hydrobromic acid

Write the net ionic equation for the reaction of aluminum and hydrobromic acid

  • 2Al(S) + 6H+(aq)+6Br--->2Al3+(aq)+6Br-(aq)+3H2(g)

  • 2Al(s)+6H+(aq)-->2Al3++3H2(g)

The activity series

The Activity Series

  • Table 4.5 pg 124 Activity Series of Metals in Aqueous Solution

  • Is a table of metals arranged in order of decreasing ease of oxidation.

  • Alkali and Alkaline Earth Metals are at the top.

  • Metals on the list can be oxidized by any of the ions below them.

Can pb no 3 2 can oxidize zn cu or fe

Can Pb(NO3)2 can oxidize Zn, Cu, or Fe

  • Zn and Fe

  • Refer to table 4.5

4 5 concentrations of solutions

4.5 Concentrations Of Solutions

  • Concentration = Amount of solute dissolved in a given quantity of solvent or solution.



  • Molarity (M) = (moles of solute)/(Volume of Solution in Liters)

Chapter 4

Calculate the molarity of a solution made by dissolving 23.4 g of sodium sulfate in enough water to form 125ml of solution

  • 23.4g*(1mol Na2SO4/142g Na2SO4) = .165 mols

  • 125ml*(1L/1000ml)=.125

  • .165mols Na2SO4/.125L = 1.32M



  • Adding water to lower the concentration is called dilution

  • Mi*Vi=Mf*Vf

  • i = initial f = final M = Molarity V= Volume

How many ml of 3 0 m h 2 so 4 are required to make 450 ml of 10 m h 2 so 4

How many mL of 3.0 M H2SO4 are required to make 450 mL of .10 M H2SO4 ?

  • Vi = (MfVf)/Mi

  • ((.10M)(450mL))/3.0M = 15 mL

4 6 stoichiometry and chemical analysis

4.6 Stoichiometry and Chemical Analysis



  • Second solution of known concentration is called the standard solution

  • Combining the standard solution with a solution of unknown concentration to get a chemical reaction is called titration

  • Equivalence point is where equivalent quantities have been brought together indicators change the color helping us to find this point.

  • If molar ratio is 1 to 1 you may use the dilution equation.

  • If not, convert standard solution to mols, then use molar ratio to give you the mols of the unknown,then convert to grams.

Chapter 4

How many grams of chloride ion are in the sample of the water if 20.2 mL of .1 M Ag is required to react with all the chloride in the sample?

  • (20.2 mL solution) * (1L/1000mL solution) * (.1mol Ag+/L solution) = 2.02 * 10-3

  • 1mol Ag+ : 1mol Cl-

  • (2.02*10-3 mol Cl-) * (35.5g Cl-/1 mol Cl-) = 7.17 * 10-2 g Cl-

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