Chapter 4. Aqueous Reactions and Solution Stoichiometry

4.1 General Properties of Aqueous Solutions Read p. 122-128. Answer the Study Guide questions

1. Draw Water Molecule (p.125)

2. What is a solution?

3. Look at Figure 4.2 - What are the major differences between electrolytes and nonelectrolytes?

4. What occurs when ionic compounds dissolve?

5. Read Figure 4.3 – Write a summary about the 2 solutions.

6. How do strong electrolytes differ from weak electrolytes?

7. What do double arrows in a chemical reaction signify?

8. What is chemical equilibrium?

• A solution is a homogeneous mixture of two or more substances.

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• A solution is made when one substance (the solute) is dissolved in another (the solvent).• The solute is the substance that is present in the smallest amount.• Solutions in which water is the solvent are called aqueous solutions.

Electrolytic Properties• If a substance forms ions in solution, then the substance is an electrolyte and the solution conducts electricity. An example is NaCl.• If a substance does not form ions in solution, then the substance is a nonelectrolyte and the

solution does not conduct electricity. An example is glucose (sugar).

Ionic Compounds in Water – VERY IMPORTANT!!!• When an ionic compound dissolves in water, the ions are said to dissociate (ionize).

• This means that in solution, the solid no longer exists as a well-ordered arrangement of ions in contact with one another.

• Instead, each ion is surrounded by a shell of water molecules.• This tends to prevent cations and anions from recombining. • The positive ions have the oxygen atoms of water pointing towards the ion; negative ions

have the hydrogen atoms of water pointing towards the ion.

Molecular Compounds in Water – VERY IMPORTANT!!!• When a molecular compound (e.g. CH3OH ) dissolves in water, there are no ions formed.• Does not conduct electricity.• There an exceptions….. NH3(g) reacts with water to form NH4

+(aq) and OH– (aq).

Strong and Weak Electrolytes• Compounds whose aqueous solutions conduct electricity well are called strong electrolytes.

• • Example: NaClNaCl(aq) → Na+(aq) + Cl–(aq)

• The single arrow indicates that the Na+ and Cl– ions have no tendency to recombine to form NaCl molecules.

• In general, soluble ionic compounds are strong electrolytes.

• Compounds whose aqueous solutions conduct electricity poorly are called weak electrolytes • These substances exist as a mixture of ions and un-ionized molecules in solution. • Example: acetic acid, HC2H3O2

HC2H3O2(aq) → H+(aq) + C2H3O2–(aq)

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• The double arrow means that the reaction can go in both directions. • This balance produces a state of chemical equilibrium.

Types of Rxns:

1.Combination - Two or more elements or compounds combine to form a single product.

a. group IA or IIA metal may combine with a nonmetal to make a salt.

Ex. A piece of lithium metal is dropped into a container of nitrogen gas. Li (s)+ N2 (g) Li3N

b. Two compounds combine to form a single product.

Ex. Sulfur dioxide gas is passed over solid calcium oxide. SO2 + CaO CaSO3

Ex. The gases boron trifluoride and ammonia are mixed.

BF3 + NH3 H3NBF3

c. Nonmetallic oxides plus water yield acids.

Ex. Carbon dioxide is bubbled into water.

CO2(g) + H2O (l) H2CO3(aq)

d. Metallic oxide plus carbon dioxide yields a metallic carbonate. MO(s) + CO2 (g) MxCO3(aq)

e. A metallic oxide plus sulfur dioxide yields a metallic sulfite. MO(s) + SO2 (g) MxSO3(aq)

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f. A metallic oxide plus water yields a metallic hydroxide. MO(g) + H2O (l) MOH (aq)

2.Decomposition - Reactions where a compound breaks down into two or more elements or compounds. Heat, electrolysis, or a catalyst is usually necessary.

a. A compound may break down to produce two elements. Ex. Molten sodium chloride is electrolyzed.

NaCl Na+ + C12

b. A compound may break down to produce an element and a compound. Ex. A solution of hydrogen peroxide is decomposed. H2O2 H2O + O2

c. Metallic carbonates break down to yield metallic oxides and carbon dioxide, Ex. Solid magnesium carbonate is heated.

MgCO3 MgO + CO2

d. Metallic chlorates break down to yield metallic chlorides and oxygen. MClO3 MCl + O2

e. NH4OH decomposes NH4OH NH3 + H2O

f. Ammonium carbonate decomposes into ammonia, water and carbon dioxide. (NH4)2CO3 2NH3 + H2O + CO2

g. Sulfurous acid decomposes into water and sulfur dioxide. H2SO3 H2O + SO2

h. Carbonic acid decomposes into water and carbon dioxide. H2CO3 H2O + CO2

3. Combustion - Elements or compounds combine with oxygen.

a. Hydrocarbons (CxHy) or alcohols (CxHyOH) combine with oxygen to form carbon dioxide and water.

Ex. Ethanol is burned completely in air.

C2H5OH + O2 CO2 + H2O

4.2 Precipitation ReactionsRead p. 128-132. Answer the Study Guide questions

1. What are precipitate reactions and why do they occur?

2. How do you determine a substance’s solubility?

3. What does it mean if a substance is insoluble?

4. How are ionic equations different from molecular equations?

5. What occurs if every ion in a in a complete ionic equation is spectator?

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• Double Replacement (Exchange) Reactions, involve swapping ions in solution:AX(aq) + BY(aq) → AY(aq) + BX(s,l,g).

Rules for DR rxn: Both reactants are aqueous solutionsOne product is also an aqueous solutionThe other product must be either a precipitation(solid), liquid, or gas.

(A precipitate is an insoluble solid formed by a reaction in solution.) If ALL reactants and products are aqueous solutions = NR (NO REACTION)!!!!

Reactions that result in the formation of an insoluble product are also known as precipitation reactions. • Example: Pb(NO3)2(aq) + 2KI(aq) → PbI2(s) + 2KNO3(aq)

Solubility Rules for Ionic Compounds – Muy Importante!!!! • The solubility of a substance at a particular temperature is the amount of that substance that

can be dissolved in a given quantity of solvent at that temperature. • To determine whether if an ionic compound is aqueous, then you need to follow Solubility Rules:

ALWAYS Soluble – Aqueous Solutions (aq) Exceptions

1. Alkali metals (Group I) None

2. Ammonium (NH4+) None

3. Nitrates (NO3–) None

4. Acetates (C2H3O2–) None

5. Chlorates (ClO3- ) None

6. Perchlorates (ClO4- ) None

7. Hydrogen Carbonate (HCO3- ) None

8. Halides: Cl–, Br–, I– Ag+, Pb2+, Hg22+, Hg2+

9. Sulfates (SO42–) Ag+, Pb2+, Hg2

2+, Hg2+, Group II (Ca2+, Sr2+, Ba2+)

If it does not fit these rules above, assume it is INSOLUBLE (s, l, g). Insoluble Exceptions

Carbonates (CO3 -2) Group I, NH4

+

Chromate (CrO4 -2) Group I, (NH4

+), Group II (Ca2+, Ba2+, Sr2+)

Fluorides (F–) Group I, NH4+

Sulfites (SO3 -2) Group I, NH4

+

Phosphates (PO4 -3) Group I, NH4

+

Hydroxides (OH-) Group I, NH4+, Group II (Ca2+, Sr2+, Ba2+)

Oxides (O -2) Group I, NH4+, Group II (Ca2+, Sr2+, Ba2+)

Sulfides (S -2) Group I, NH4+, Group II (Ca2+, Sr2+, Ba2+)

Also, GASES, PURE LIQUIDS, and SOLIDS are non-electrolytes. Ionic Equations• The molecular equation lists all species in their molecular forms:

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

• The ionic equation lists all strong soluble electrolytes in the reaction as ions:Pb2+(aq) + 2NO3

–(aq) + 2K+(aq) + 2I–(aq) → PbI2(s) + 2K+(aq) + 2NO3–(aq)

Note that spectator ions, ions that are present in the solution but play no direct role in the reaction, are omitted in the net ionic equation.

• Strong electrolytes (aq) – ionic form. • Weak electrolytes and nonelectrolytes - molecular form (don’t dissociate).

• The net ionic equation lists only those ions which are not common on both sides of the reaction:

Pb2+(aq) + 2I–(aq) → PbI2(s)

Example Predict if a reaction takes place between the following solutions. If it does, write a net ionicequation for the reaction.(i) KOH(aq) + MgBr2(aq) →

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(ii) BaS(aq) + NiSO4(aq) →

4.3 Acid-Base ReactionsRead p. 132-138. Answer the Study Guide questions

1. What are the qualities of an acid?

2. Why are acids called proton donors?

3. What are the qualities of a base?

4. What is reactivity dependent upon?

5. How do you distinguish between strong and weak acids and bases?

6. Write out Table 4.3 of Summary of Electrolytes

7. What occurs when a solution of an acid and a solution of a base are mixed?

8. What are salts?

Acids – all aqueous!!• Acids are substances that are able to ionize in aqueous solution to form H+.

7 Strong Acids that are Strong Electrolytes: HCl, HBr, HI, HNO3, H2SO4, HClO3, HClO4

The rest of the acids in the world are weak electrolytes!!!!! (also they are all aqueous)

• Since H+ is a proton, we refer to acids as proton donors • Acids that ionize to form one H+ ion are called monoprotic acids.• Acids that ionize to form two H+ ions are called diprotic acids.

Bases• Bases are substances that accept H+ ions from acids. Bases are proton acceptors.• Hydroxide ions, OH–, react with the H+ ions to form water:

H+(aq) + OH–(aq) → H2O(l)

8 Strong Bases: LiOH, NaOH, KOH, RbOH, CsOH, Ca(OH)2, Sr(OH)2, Ba(OH)2

• Common bases are NH3 (ammonia), Draino, and milk of magnesia.• Compounds that do not contain OH– ions can also be bases.

• Proton transfer between NH3 (a weak base) and water (a weak acid) is an example of an acid–base reaction.

NH3(aq) + H2O(l) → NH4+(aq) + OH–(aq)

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

• They are completely ionized in solution.

HCl → H+ + Cl–

NaOH → Na+ + OH–

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• Weak acids and weak bases are weak electrolytes. (most acids and bases in the world are weak)

• Therefore, they are partially ionized in solution. HF → + F–

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Identifying Strong and Weak Electrolytes• Compounds can be classified as strong electrolytes, weak electrolytes, or nonelectrolytes

by looking at their solubility.• Strong electrolytes:

• Ionic compounds (metal + nonmetal)• Strong Acids and Bases.

• Weak electrolytes:• Molecular Compounds (nonmetals only) Weak acids and bases are weak electrolytes.

• Nonelectrolytes:• All other compounds….especially if they are CxHy molecules

Neutralization Reactions and Salts• A neutralization reaction occurs when an acid and a base react:

• HCl(aq) + NaOH(aq) → H2O(l) + NaCl(aq)• (acid) + (base) (water) + (salt)

• In general: an acid and a base react to form a salt and water.

• A salt is any ionic compound whose cation comes from a base and anion from an acid. • The other product, H2O, is a common weak electrolyte.

• A typical example of a neutralization reaction is:• the reaction between an acid and a metal hydroxide.

• Balanced equation:Mg(OH)2(s) + 2HCl(aq) → MgCl2(aq) + 2H2O(l)

• Net ionic equation:Mg(OH)2(s) + 2H+(aq) → Mg2+(aq) + 2H2O(l)

Acid-Base Reactions with Gas Formation• There are many bases besides OH– that react with H+ to form molecular compounds.

• Reaction of sulfides with acid gives rise to H2S(g). • Balanced equation:

Na2S(aq) + 2HCl(aq) → H2S(g) + 2NaCl(aq)• Net ionic equation:

2H+(aq) + S2–(aq) → H2S(g)

• Carbonates and hydrogen carbonates (or bicarbonates) will always form CO 2( g ) and water when treated with an acid.• Balanced equation:

HCl(aq) + NaHCO3(s) →NaCl(aq) + H2CO3(aq) → H2O(l) + CO2(g) + NaCl(aq)

• Net ionic equation:H+(aq) + HCO3

–(aq) → H2O(l) + CO2(g)

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4.4 Single Replacement & Oxidation-Reduction ReactionsRead p. 138-146. Answer the Study Guide questions

1. What occurs in redox reactions?

2. How is corrosion an example of a redox reaction?

3. What is the difference between reduction and oxidation?

4. How do you determine an atom’s oxidation number?

5. What are displacement reactions?

6. What is the purpose of the activity series?

Single Replacement. Also called Oxidation-Reduction (Redox) reaction • The general pattern is:

A + BX →AX + B• In the process the metal is oxidized and the H+ is reduced.

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

• The net ionic equation shows the redox chemistry well:Fe(s) + Ni2+(aq) → Fe2+(aq) + Ni(s)

• In this reaction Fe has been oxidized to Fe2+, while the Ni2+ has been reduced to Ni.

The Activity Series – Have to use for the Single Replacement Rxn!!!!!• We can list metals in order of decreasing ease of oxidation.

• The metals at the top of the activity series are called active metals.• A metal in the activity series can only be oxidized (replaced) by a metal ion below it.

Active metals replace less active metals or hydrogen from their compounds in aqueous solution. Use an activity series or a reduction potential table to determine activity. The more easily oxidized metal replaces the less easily oxidized metal. The metal with the most negative reduction potential will be the most active .

Cu(s) + 2AgNO3(aq) → Cu(NO3)2(aq) + 2Ag(s)or

Cu(s) + 2Ag+(aq) → Cu2+(aq) + 2Ag(s)

Activity Series - METALS

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Activity Series – NONMETALS (Halogens only)

Active nonmetals replace less active nonmetals from their compounds in aqueous solution. Each halogen will displace less electronegative (heavier) halogens from their binary salts.

a. Metallic oxides (basic anhydrides) plus water yield bases.

Ex. Solid sodium oxide is added to water.

Na2O + H2O Na+ + OH-

b. Metallic hydrides plus water yield metallic hydroxides and hydrogen gas.

Ex. Solid sodium hydride is added to water.

NaH + H2O Na+ + OH- + H2

Oxidation and Reduction (redox)

Reactions involve the transfer of electrons between reactants. • When a substances loses electrons, it undergoes oxidation:

Ca(s) + 2H+(aq) → Ca2+(aq) + H2(g)

• The neutral Ca has lost two electrons to become Ca2+. We say Ca has been oxidized.

• When a substance gains electrons, it undergoes reduction:

• The H+ ion gained electrons to become H2. We say H+ has been reduced.

In all redox reactions, one species is reduced at the same time as another is oxidized.

Oxidation Numbers

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• Electrons are not shown in chemical equations.• Oxidation numbers (or oxidation states) help up keep track of electrons during chemical reactions.

• Rules for Oxidation numbers: 1. For an atom in its elemental form, the oxidation number is always zero.2. For any monatomic ion, the oxidation number equals the charge on the ion.3. Nonmetals usually have negative oxidation numbers.

• The oxidation number of oxygen is usually –2. The major exception is in peroxides (containing the O2

2– ion). 4. The oxidation number of hydrogen is +1 when bonded to nonmetals and –1 when bonded

to metals. 5. Halogens have an oxidation number of –1. 6. The sum of the oxidation numbers in a polyatomic ion equals the charge of the ion. 7. The sum of the oxidation numbers of all atoms in a neutral compound is zero .

Oxidation = lose of electrons = increase in oxidation numberReduction = gain of electrons = decrease in oxidation number.

4.5 – 4.6 Concentrations of Solutions and Solution StoichiometryRead p. 146-153. Answer the Study Guide questions

1. How is concentration expressed?

2. How do you calculate molarity?

3. How do you interconvert molarity, moles and volume?

4. How are dilutions prepared in the laboratory? (READ FIGURE 4.15 & 4.16 – KNOW IT WELL HOW TO MAKE DILUTIONS. See next page of notes for picture)

5. What occurs to the number of moles when a solvent is added to a solution?

Ch 4.6 – Solution Stoichiometry and Chemical Analysis1. How are coefficients in a balanced equation used in stoichiometry?

2. Why are molar mass conversions not valid when working with a solution?

• The term concentration is used to indicate the amount of solute dissolved in a given quantity of solvent or solution.

Molarity• The amount (moles) of solute per liter of solution is the molarity (M) of the solution:

Expressing the Concentration of an Electrolyte• When an ionic compound dissolves, the relative concentrations of the ions in the solution depend on

the chemical formula of the compound. • Example: 1.0 M solution of NaCl:

• The solution is 1.0 M in Na+ ions and 1.0 M in Cl– ions. • Example: 1.0 M solution of Na2SO4:

• The solution is 2.0 M in Na+ ions and 1.0 M in SO42– ions.

Subscripts are very important!!!!

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Interconverting Molarity, Moles, and Volume• The definition of molarity contains three quantities: molarity, moles of solute, and liters of solution.

• If we know any two of these, we can calculate the third = STOICHIOMETRY!!!

Dilution – AP EXAM LAB QUESTION!!• A solution in concentrated form (stock solution) is mixed with solvent to obtain a solution of lower

solute concentration. • This process is called dilution.

• An alternate way of making a solution is to take a solution of known molarity and dilute it with more solvent. (diluting your Kool-Aid)

• Since the number of moles of solute remains the same in the concentrated and diluted forms of the solution, we can show:

MconcVconc = MdilVdil

• An alternate form of this equation is:M1V1 = M2V2

AP Exam Lab Question:Steps involved in the preparation of a standard solution. a) A weighed amount of a substance (the solute) is put into the volumetric flask, and a small quantity of water is added. (b) The solid is dissolved in the water by gently swirling the flask (with the stopper in place). (c) More water is added, until the level of the solution just reaches the mark etched on the neck of the flask (d).

SUMMARY of REACTIONS(1) Metal + oxygen --> most common oxide (COMBINATION)Examples:4Li(s) + O2(g) --> 2Li2O(s)2Cu(s) + O2(g) --> 2CuO(s)

(2) Alkali metal or alkaline earth metal + water --> metal hydroxide + hydrogen gas

Example:2Li(s) + 2H2O(l) --> 2LiOH(aq) + H2(g)

(3) Alkali metal oxide or alkaline earth metal oxide + water --> metal hydroxide (COMBINATION)

Example:Li2O(s) + H2O(l) --> LiOH(aq)

(4) Metal + soluble salt --> reduced metal + oxidized metal salt (depends on activity series) (SINGLE REPLACEMENT)

Example:Cu(s) + 2AgNO3(aq) --> 2Ag(s) + Cu(NO3)2(aq) [full ionic equation]Cu(s) + 2Ag+1(aq) --> 2Ag(s) + Cu+2(aq) [net ionic equation]

(5) Hydrates: (DECOMPOSTION)a. Heated hydrates release water

MgSO4•3H2O(s) --> MgSO4(s) + 3H2O(g)

(2) Metals + acid --> salt + hydrogen gas (SINGLE REPLACEMENT)

(remember that all acids are aqueous)(note that this is a reduction-oxidation)

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Example:Zn(s) + 2HCl(aq) --> ZnCl2(aq) + H2(g) [full ionic equation]Zn(s) + 2H+1(aq) --> Zn+2(aq) + H2(g) [net onic equation]

(3) Electricity + ionic compound --> elements (DECOMPOSTION)(note this is a reduction-oxidation)

Example:2NaCl(s) + electricity --> 2Na(s) +Cl2(g)

(4) Gas reactions: (DOUBLE REPLACEMENT)Any reaction that produces hydrogen carbonate will produce water + carbon dioxideAny reaction that produces hydrogen sulfite will produce water + sulfur dioxide

Examples:

Na2CO3(aq) + 2HCl(aq) --> 2NaCl(aq) + H2O(l) + CO2(g) [full ionic equation] CO3-2(aq) + 2H+1(aq) --> H2O(l) + CO2(g) [net ionic equation

Na2SO3(aq) + 2HCl(aq) --> 2NaCl(aq) + H2O(l) + SO2(g) [full ionic equation] SO3-2(aq) + 2H+1(aq) --> H2O(l) + SO2(g) [net ionic equation]

(5) Combustion: Hydrocarbon + oxygen gas --> water + carbon dioxide

Examples:2C8H18(l) + 25O2(g) --> 16CO2(g) + 18H2O(l)C6H12O6(s) + 6O2(g) --> 6CO2(g) + 6H2O(l)

(6) Precipitation reactions: PRODUCTS SOLIDS!!!

(7) Acids and bases: ONE OF THE PRODUCTS IS ALWAYS H20HCl(aq) + NaOH(aq) --> NaCl(aq) + H2O(l)H2SO4(aq) + Ba(OH)2(s) --> BaSO4(aq) + H2O(l)

(8) Reduction-oxidation: (ALSO CALLED SINGLE REPLACEMENT) These cover a wide variety. Example:2Al(s) + 3CuO(s) --> Al2O3(s) + 3Cu(s)

(9) Hydrogen peroxide decomposes: DECOMPOSTIONH2O2(aq) H2O(l) + O2(g)

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