chemical reactions in aqueous solutions

Prentice Hall © 2005General Chemistry 4th edition, Hill, Petrucci, McCreary, Perry

Chapter Four

1

Chemical Reactions in Aqueous Solutions

Chapter Four

Chapter OnePrentice Hall © 2005General Chemistry 4th edition, Hill, Petrucci, McCreary, Perry

Today…• Turn in:• Nothing

• Our Plan:• Notes – Synthesis & Decomposition• Begin Worksheet #1

• Homework (Write in Planner):• WS#1 Due Wednesday


Unit 1 Test Results

A 1B 6C 6D 3F 0

Average 78.47%

High Score 96%


Chapter Four

4

Quick Review•How do we know a

chemical reaction has occurred? What do we observe?


Chapter Four

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Evidence of a Chemical Rxn1. Color Change2. Precipitate (solid) forms3. Gas Evolved4. Heat/Light Given Off

–Endothermic – heat absorbed (gets cold)

–Exothermic – heat released (gets hot)• More about this in Unit 5 - Thermodynamics


Chapter Four

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Review of Basics• When writing out a balanced equation, it is

necessary to indicate what state the substance is in (s, l, g, or aq).

• For elements, you can look at the PT to determine their standard state.– Examples:

• Mercury• Fluorine • Iron


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Review of Basics• There are 7 diatomic elements.• They are only diatomic when

they are alone. They are all gases, except Br2, which is a liquid, as indicated on the PT.

• Remember the Super 7?–H2, N2, O2, F2, Cl2, Br2, I2


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Review of Basics• Soluble means the substance dissolves

in solution (water). –Label soluble substances aqueous (aq)

• Insoluble means the substance does not dissolve in solution.–Label insoluble substances solid (s)

• To determine if a substance is soluble or insoluble, use a solubility table.


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Solubility Chart


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Other Basic Tips• Acids and bases are aqueous• Water is almost always a liquid• Write water as HOH when it is

a reactant• Some common gases that aren’t

diatomic are CO2, CO, SO3, SO2, H2S, and NH3


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Other Basic Tips• If a reaction indicates that a catalyst was

present, it is not part of the actual reaction. Instead, you write the catalyst above the arrow.

• Catalysts are not used up in the reaction. They are just there to speed it up.

• If a reaction is heated, you draw a triangle above the arrow. Heat is used as a catalyst.


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Synthesis Reactions• Occur when two or more

reactants combine to form a single product.–In Chem 1 we called this type

“Dating”• There are several common types

of synthesis reactions.


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Synthesis ReactionsType 1: A metal combines with a nonmetal to form a binary salt.

Example: A piece of lithium metal is dropped into a container of nitrogen gas.

6Li (s) + N2 (g) → 2Li3N (aq)


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Synthesis ReactionsType 2: Metallic oxides and water form bases (metallic hydroxides).Example: Solid sodium oxide is added to water

Na2O (s) + HOH (l) → 2NaOH (aq)Example: Solid magnesium oxide is added to water.

MgO (s) + HOH (l) → Mg(OH)2 (aq)


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Synthesis ReactionsType 3: Nonmetallic oxides and water form acids. The nonmetal retains its oxidation number.Example: Carbon dioxide is bubbled in water.

CO2 (g) + H2O (l) → H2CO3 (aq)Example: Dinitrogen pentoxide is bubbled in water.

N2O5 (g) + H2O (l) → 2HNO3 (aq)


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Synthesis ReactionsType 4: Metallic oxides and nonmetallic oxides form salts.Example: Solid sodium oxide is added to carbon dioxide.

Na2O (s) + CO2 (g) → Na2CO3 (aq)Example: Solid calcium oxide is added to sulfur trioxide.

CaO (s) + SO3 (g) → CaSO4 (aq)


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Try It Out 1• Solid barium oxide is added to distilled

water (p. 142)


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Try It Out 2• Sulfur trioxide gas is added to excess water

(p. 146)


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Decomposition Reactions• Occur when a single reactant is

broken down into two or more products.– In Chem 1 we called this “The Break Up”

• There are several common types of decomposition reactions.


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Decomposition ReactionsType 1: Metallic carbonates decompose into metallic oxides and carbon dioxide

Example: A sample of magnesium carbonate is heated.

MgCO3 (s) → MgO (s) + CO2 (g)


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Decomposition ReactionsType 2: Metallic chlorates decompose into metallic chlorides and oxygen.

Example: A sample of magnesium chlorate is heated.Mg(ClO3)2 (aq) → MgCl2 (aq) + 3O2 (g)


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Decomposition ReactionsType 3: Ammonium carbonate decomposes into ammonia, water, and carbon dioxide.

Example: A sample of ammonium carbonate is heated.

(NH4)2CO3 (aq)→ 2NH3 (g) + H2O (l) + CO2 (g)


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Decomposition ReactionsType 4: Sulfurous acid decomposes into sulfur dioxide and water.

Example: A sample of ammonium carbonate is heated.

H2SO3 (aq) → H2O (l) + SO2 (g)


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Decomposition ReactionsType 5: Carbonic acid decomposes into carbon dioxide and water.

Example: A sample of carbonic acid is heated.

H2CO3 (aq) → H2O (l) + CO2 (g)


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Decomposition ReactionsType 6: A binary compound may break down to produce two elements.

Example: Molten sodium chloride is electrolyzed.

2NaCl (l) → 2Na (s) + Cl2 (g)


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Decomposition ReactionsType 7: Hydrogen peroxide decomposes into water and oxygen.

Example: 2H2O2 (aq) → 2H2O (l) + O2 (g)


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Decomposition ReactionsType 8: Ammonium hydroxide decomposes into ammonia and water.

Example: NH4OH (aq) → NH3 (g) + HOH (l)


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Try It Out 1• Solid calcium chlorate is heated in the

presence of a magnesium dioxide catalyst


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Try It Out 2• A sample of lithium carbonate is heated

strongly


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STOP•Complete Worksheet #1

by next class.


Today…• Turn in:• Get Out WS#1

• Our Plan:• Questions on WS#1• Quick Review• Notes – Single & Double Replacement• Begin WS#2

• Homework (Write in Planner):• WS#2 Due Friday


Quick Review• Write formulas for the reactants and predicted products for

the chemical reactions that follow: 1. Solid calcium carbonate is strongly heated in a test tube.

2. Solid lithium chlorate is heated in the presence of a manganese dioxide catalyst.

3. Excess chlorine gas is passed over hot iron filings.

32


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Single Replacement• Reactions that involve an element

replacing one part of a compound.• The products include the displaced

element and a new compound.–In Chem 1 we called this “Cheating”

• An element can only replace another element that is less active than itself.


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Single ReplacementGeneral Activity Series for Metals:

Most ActiveLiCaNaMgAlZnFePbH2

CuAgPtAu

Least Active


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Single ReplacementGeneral Activity Series for Nonmetals:

Most ActiveF2

Cl2

Br2

I2

Least Active


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Remember the Trend?• Or you can just look at the Periodic

Table–The most reactive metals are on the

bottom left (Francium)–The most reactive nonmetals are on the

top right (Fluorine)–Hydrogen is the tricky one.–Use the pink sheet for electronegativity

potentials otherwise.


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For metals, the more negative the reduction potential, the more reactive. For nonmetals, the more positive the reduction potential, the more reactive.


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Single ReplacementType 1: Active metals replace less active metals from their compounds in aqueous solution.

Example: Magnesium turnings are added to a solution of iron (III) chloride.

3Mg (s) + 2FeCl3 (aq)→ 2Fe (s) + 3MgCl2 (aq)


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Single ReplacementType 2: Active metals replace hydrogen in water.

Example: Sodium is added to water.2Na (s) + 2HOH (l)→ 2NaOH (aq) + H2 (g)


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Single ReplacementType 3: Active metals replace hydrogen in acids.

Example: Lithium is added to hydrochloric acid.

2Li (s) + 2HCl (aq) → 2LiCl (aq) + H2 (g)


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Single ReplacementType 4: Active nonmetals replace less active nonmetals from their compounds in aqueous solution.

Example: Chlorine gas is bubbled into a solution of potassium iodide.Cl2 (g) + 2KI (aq) → I2 (g) + 2KCl (aq)


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Single ReplacementType 5: If a less reactive element is combined with a more reactive element in compound form, their will be no resulting reaction.

Example: Chlorine gas is bubbled into a solution of potassium iodide.

Cl2 (g) + KF (aq) → No ReactionExample: Zinc is added to a solution of sodium chloride.

Zn (s) + NaCl (aq) → No Reaction


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Try It Out 1• A piece of aluminum metal is added to a

solution of gold nitrate. (p. 146)


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Try It Out 2• Liquid bromine is shaken with a potassium

iodide solution. (p. 152)


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Try It Out 3• Small chunks of potassium are added to

water. (p. 164)


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Try It Out 4• Small strips of platinum are placed in

hydrochloric acid.


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Double Replacement

• Reactions between two compounds in aqueous solution where the cations and anions appear to “switch partners”.

AX + BY → AY + BX• In Chem 1 we called this

“Swapping”


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Double Replacement

• All double replacement reactions (aka metathesis) must have a “driving force” or reason why the reaction will occur or “go to completion”.


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Double Replacement

• “Driving Force” for reactions:1. Formation of a precipitate2. Formation of a gas3. Formation of primarily

molecular species (nonelectrolytes, water, weak acids)


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Double Replacement• If one of these “driving forces” is NOT

present, then the reaction does not go to completion.

• This type of reaction is indicated by a double arrow


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• There are limits to the amount of a solute that will dissolve in a given amount of water.

• If the maximum concentration of solute is less than about 0.01 M, we refer to the solute as insoluble in water.

• When a chemical reaction forms such a solute, the insoluble solute comes out of solution and is called a precipitate.

Reactions that Form Precipitates


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Silver Iodide Precipitation

A solution containing silver ions and nitrate ions, when added to …

… a solution containing potassium ions and iodide ions,

forms …

… a precipitate of silver iodide.

What is the net ionic equation for the reaction that has

occurred here? (Hint: what species actually

reacted?)


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Double Replacement (Precipitate)

• In order to predict double replacement reactions yielding precipitates, one must memorize the solubility rules.

• I will let you use a solubility chart on your test!


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Solubility Rules


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Example 1• Predict and balance the following double

replacement reactions based on the solubility of the products. Use the abbreviations (aq) and (s) for the reactant and products. All reactants are aqueous.

1. Solutions of manganese (II) sulfate and ammonium sulfide are mixed.


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Example 22. Solutions of sodium iodide and lead (II)

nitrate are combined.


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Try it Out 1• Solutions of sodium carbonate and

iron (III) nitrate combine.


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Try it Out 2• Solutions of lead (II) acetate and

calcium chloride are combined.


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Double Replacement (Gas)

• Common Gases formed: H2S, CO2, SO2, NH3

• Reactions that produce three of the gases (CO2, SO2, and NH3) involve the initial formation of a substance that breaks down to give the gas and HOH.


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Double Replacement (Gas)Common Gases

H2S Any sulfide (salt of S2-) plus any acid form H2S (g) and a salt.

CO2 Any carbonate (salt of CO32-) plus any acid form CO2

(g), HOH, and a saltSO2 Any sulfite (salt of SO3

2-) plus any acid form SO2 (g), HOH, and a salt.

NH3 Any ammonium salt (salt of NH4+) plus any soluble

strong hydroxide react upon heating to form NH3 (g), HOH, and a salt.


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Double Replacement (Gas)• Example 1: The reaction of Na2SO3 and HCl

produces H2SO3:Na2SO3 (aq) + 2HCl (aq) → H2SO3 (aq) + 2NaCl (aq)

• Bubbling is observed in this reaction because the H2SO3 (sulfurous acid) is unstable and immediately decomposes to give HOH and SO2 gas:

H2SO3 (aq) → HOH (l) + SO2 (g)• The molecular equation for the overall or

complete reaction, therefore, is:Na2SO3 (aq) + 2HCl (aq) → HOH (l) + SO2 (g) + 2NaCl (aq)


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Double Replacement (Gas)• Example 2: A typical reaction of a carbonate

and an acid is:K2CO3 (aq) + 2HNO3 (aq) → HOH (l) + CO2 (g) + 2KNO3 (aq)

• Bubbling is also observed in this reaction. Theoretically H2CO3, carbonic acid, is formed, but the acid is unstable and immediately decomposes to form carbon dioxide gas and water according to the following equation:

H2CO3 (aq) → HOH (l) + CO2 (g)


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Double Replacement (Gas)• Example 3: Ammonium salts and soluble bases

react as follows (particularly when the solution is warmed):

NH4Cl (aq) + NaOH (aq) → NH3 (g) + HOH (l) + NaCl (aq)• The odor of ammonia gas is noted and moist blue

litmus paper held near the mouth of the container will turn blue. Theoretically NH4OH, ammonium hydroxide, is produced (also known as ammonia water). The compound is unstable and decomposes into ammonia gas and water:

NH4OH (aq) → NH3 (g) + HOH (l)


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Double Replacement (Gas)• Example 4: The odor of rotten

eggs and bubbling are noted when an acid is added to a sulfide. A typical reaction producing hydrogen sulfide gas is:

FeS (s) + 2HCl (aq) → FeCl2 (aq) + H2S (g)


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Try It Out 1• Dilute hydrochloric acid is added to a

solution of potassium sulfite. (p. 152)


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Try It Out 2• Concentrated hydrochloric acid is added to

solid manganese (II) sulfide. (p. 150)


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• Our Plan:• Questions on WS#2• Quick Review/Video• Notes – Aqueous Reactions, Net Ionic Equations, &

Math with Ion Concentration• Begin WS#3

• Homework (Write in Planner):• WS#3 Due Monday


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A Note about phases• A product can only be aqueous if there is an

aqueous or liquid reactant. That’s why some of the ones on the first WS that you labeled aqueous were solid.

• Examples:


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Quick Discussion of Weak Electrolytes

• One of the driving forces for a reaction is the production of weak electrolytes (water or a weak acid).

• For a list of strong acids see page 22 in booklet. Everything else is weak.

• The only thing to note is that if you form a weak acid, it is aqueous but you would have a forward arrow, not a double arrow.

• See examples on the board.


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Quick Review• Solutions of sodium fluoride and dilute

hydrochloric acid are combined. • Dilute acetic acid solution is added to

solid magnesium carbonate. • A saturated solution of calcium

hydroxide was added to a solution of magnesium chloride.


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Crash Course in Precipitation• http://

www.youtube.com/watch?v=IIu16dy3ThI

http://www.youtube.com/watch?v=IIu16dy3ThI

http://www.youtube.com/watch?v=IIu16dy3ThI


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• Why do some solutions conduct electricity?• An early hypothesis was that electricity

produced ions in solution, and those ions allowed the electricity to flow.

• Arrhenius’s theory:– Certain substances dissociate into cations and

anions when dissolved in water.– The ions already present in solution allow

electricity to flow.

Arrhenius’s Theory ofElectrolytic Dissociation


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• Electrolytes dissociate to produce ions.

Electrolytic Properties of Aqueous Solutions

The more the electrolyte dissociates, the more ions it produces.


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• A strong electrolyte dissociates completely.– A strong electrolyte is present in solution almost

exclusively as ions.– Strong electrolyte solutions are good conductors.

• A nonelectrolyte does not dissociate.– A nonelectrolyte is present in solution almost exclusively

as molecules.– Nonelectrolyte solutions do not conduct electricity.

• A weak electrolyte dissociates partially.– Weak electrolyte solutions are poor conductors.– Different weak electrolytes dissociate to different extents.

Types of Electrolytes


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• Strong electrolytes include:– Strong acids (HCl, HBr, HI, HNO3, H2SO4, HClO4)– Strong bases (IA and IIA hydroxides)– 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 of them are molecular)

Is it a strong electrolyte, a weak electrolyte, or a nonelectrolyte?

How do we tell whether an acid

(or base) is weak?


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Aqueous Solutions & Ionic Equations

• On the AP Exam, you do not have to write out complete molecular equations like we have been doing.

• Instead, you have to write ionic equations.


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Overall (Total) Ionic Equation• Formulas of the reactants and products are

written to show the predominant form of each substance as it exists in aqueous solution.

• Soluble salts, strong acids, and strong bases are written as separated ions. Everything else is written as a molecule.

• We will memorize strong acids and strong bases later in the year. For now, a list will be provided.


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Memorize Later…Strong Acids Strong Bases

HClO4 LiOH RbOHH2SO4 NaOH CsOH

HI KOH Mg(OH)2

HBr Ca(OH)2

HCl Sr(OH)2

HNO3 Ba(OH)2


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Overall Ionic Equation Example 1Cd(NO3)2 (aq) + Na2S (aq) → CdS (s) + 2NaNO3 (aq)


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Overall Ionic Equation Example 2CaCO3 (aq) + 2HCl (aq) → CaCl2 (aq) + HOH (l) + CO2 (g)


Chapter Four

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Try it Out(NH4)2S (aq) + 2LiOH (aq) → Li2S (aq) + 2NH3 (g) + 2HOH (l)


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Net Ionic Equation• Net ionic equations are written to show only

the species that react or undergo change in aqueous solutions.

• The net ionic equation is obtained by eliminating the spectator ions from the overall equation.

• All that is left are the ions that have changed chemically.


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Net Ionic Equation Example 1Cd(NO3)2 (aq) + Na2S (aq) → CdS (s) + 2NaNO3 (aq)


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Example 2• For each of these word equations, predict the product

and write an overall and net ionic equation.• Copper (I) nitrate is combined with silver chloride.


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Example 3• Magnesium hydroxide is combined with lead (II)

sulfate.


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Try It Out• Mercury (I) nitrate is combined with sodium sulfide.


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• “Trick” question …What is the concentration of Na2SO4 in a solution prepared by diluting 0.010 mol Na2SO4 to 1.00 L?

• The answer is:… zero …

• WHY??• And … how do we describe the concentration of

this solution?

Ion Concentrations in Solution


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• In 0.010 M Na2SO4:– two moles of Na+ ions are formed for each

mole of Na2SO4 in solution, so [Na+] = 0.020 M.

– one mole of SO42– ion is formed for each

mole of Na2SO4 in solution, so [SO42–] =

0.010 M.• An ion can have only one concentration

in a solution, even if the ion has two or more sources.

Calculating Ion Concentrations in Solution


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Example 4.1Calculate the molarity of each ion in an aqueous solution that is 0.00384 M Na2SO4 and 0.00202 M NaCl. In addition, calculate the total ion concentration of the solution.


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Try It Out!• Exercise 4.1A: Seawater is essentially 0.438 M NaCl and

0.0512 M MgCl2, together with several other minor solutes. What are the molarities of Na+1, Mg+2, and Cl-1 in seawater?


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Example 4.6One cup (about 240 g) of a certain clear chicken broth yields 4.302 g AgCl when excess AgNO3(aq) is added to it. Assuming that all the Cl– is derived from NaCl, what is the mass of NaCl in the sample of broth?


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Stop!• Complete Worksheet #4 by

Monday.–Let’s look at it together.


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• Our Plan:• Questions on WS#3• Quick Review• Notes – Acid/Base Reactions & Titrations• Begin WS#4

• Homework (Write in Planner):• WS#4 Due Wednesday


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Quick Review• Write the molecular and net ionic

equation for the following:1. Chlorine gas is bubbled into a

solution of potassium iodide.2. Aqueous solutions of potassium

chromate and silver nitrate react.


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Acid-Base Reactions• We will cover acid/base reactions in

great detail at the end of the school year. For now, we’re just going to cover very simple neutralization reactions.

• Neutralization is the transfer of PROTONS from an acid to a base.

• In neutralization, an acid and a base form a salt and water.


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Acid-Base Reactions• When writing net ionic equations, keep in mind

which acids are strong (written in ionic form) and which are weak (written in molecular form).

• Check the solubility rules of the salt produced. If it is soluble, it is written in ionic form; if it is insoluble it is written in molecular form.

• The salt produced always consists of a cation from the base and an anion from the acid.


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Remember, these are in your notes…

Strong Acids Strong BasesHClO4 LiOH RbOHH2SO4 NaOH CsOH

HI KOH Mg(OH)2

HBr Ca(OH)2

HCl Sr(OH)2

HNO3 Ba(OH)2


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Acid-Base Reactions• Example 1: Hydrogen sulfide

gas is bubbled through excess potassium hydroxide solution.

H2S (g) + 2KOH (aq) → K2S (aq) + 2HOH (l)


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Acid-Base Reactions• Polyprotic acids (more than one hydrogen like

H2SO4) can be tricky. If the base is in excess, all hydrogen ions will react with strong base to produce water. See Example 2.

• If however, this same reaction were described in terms of mixing equal numbers of moles, then the coefficients for both reactants would be one (the same number of H and OH must be given away). See Example 3.


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Acid-Base Reactions• Example 2: Dilute sulfuric acid is

reacted with excess sodium hydroxide.

H2SO4 (aq) + 2NaOH (aq) → Na2SO4 (aq) + 2HOH (l)


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Acid-Base Reactions• Example 3: Equal number of

moles of sulfuric acid and sodium hydroxide react.

H2SO4 (aq) + NaOH (aq) → NaHSO4 (aq) + HOH (l)


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Acid-Base Reactions• As the following example demonstrates, it

is important to take into account the quantity (concentration and amount) of each reactant.

• Example 4: Equal volumes of 0.1 M phosphoric acid and 0.2 M sodium hydroxide are reacted together.

H3PO4 (aq) + 2NaOH (aq) → Na2HPO4 (aq) + 2HOH (l)


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Try It Out 1• Equal molar volumes of sodium hydroxide

and hydrobromic acid are mixed together. Write the molecular equation and the net ionic equation.


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Try It Out 2• Equal volumes of 0.1 M phosphoric acid

and potassium hydroxide solutions are mixed.


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Titrations• A titration is a procedure where one

reactant is carefully added to another reactant until the two have combined in their exact stoichiometric proportions.

• There are different types of titrations, but no matter the type, both reactants are fully consumed at the end of the titration.


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Titrations• The purpose of a titration is to find the number

of moles, grams, concentration, or percentage of an unknown.

• The unknown is called the ANALYTE.• This is done by measuring the precise volume

and concentration of a known solution (TITRANT) needed to react completely with the analyte.

• Stoichiometry is then used to determine whatever quantity you’re looking for.


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Acid-Base Titrations• Acid-base titrations are the most common type• You conducted this type of titration in Chem 1

(remember the antacid lab?)• There are 3 things that you need for a titration:

– A way to accurately measure the titrant (BURET)

– A way to know your reaction is complete (INDICATOR or you can look at a TITRATION CURVE)

– A solution whose concentration you know (STANDARD SOLUTION)


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Acid-Base Titrations


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Acid-Base Titrations• Equivalence point – in an acid-base titration

it is the point when the titrant completely neutralizes the analyte.

• If you’re using an indicator, you choose one that changes color close to the neutralization point.

• If you are looking at a titration curve, it is the straight line on a graph of volume versus pH.


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Acid-Base Titrations• When the indicator changes color

in a titration, you have reached the endpoint.

• At the endpoint the titration is stopped and the volume of titrant is recorded.

• You want the equivalence point and endpoint to be the same!


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Titration Curve


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Calculation Example 1• Example 4.9: What volume (mL) of 0.210 M

NaOH is required to neutralize 20.00 mL of 0.103 M HCl in an acid-base titration?


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Calculation Example 2• Example 4.10: A 10.00 mL sample of an aqueous solution

of calcium hydroxide is neutralized by 23.30 mL of 0.020000 M HNO3 (aq). What is the molarity of the calcium hydroxide solution?


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Stop• Complete Worksheet #4 by next

class and complete the pre-lab for Investigation 4.


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• Our Plan:• Questions on WS#4• Investigation 4

• Homework (Write in Planner):• Lab Report Due Monday• Post results (Average concentration of

H3PO4) to Edmodo by Friday.


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Quick Review• Equal volumes of 0.1 M sulfuric acid and

0.2 M potassium hydroxide are mixed. Write the balanced molecular and net ionic equation. (p. 148)


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Today…• Turn in:• Questions on Lab Report so far?

• Our Plan:• Notes – Oxidation Numbers & Redox

Titrations• Begin WS#5 (10 minutes)• Sample Test (Pep Assembly Today)

• Homework (Write in Planner):• WS#5 Due Monday & Lab Report due Mon


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Redox Reactions• Balancing Redox Equations is

complicated. We will spend an entire week at the end of the year working on it.

• For now, you just have to identify substances that are oxidized and those that are reduced.


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• 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).• To remember, use the phrase LEO GER or

OIL RIG.

Reactions InvolvingOxidation and Reduction


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A Redox Reaction: Mg + Cu2+ → Mg2+ + Cu

… the products are Cu metal and Mg2+ ions.

Electrons are transferred from Mg metal to Cu2+

ions and …


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• 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 CO2 (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 CO2 has an oxidation number of +4 (Why?)

Oxidation Numbers


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1.For the atoms in a neutral species—an isolated atom, a molecule, or a formula unit—the sum of all the oxidation numbers is 0.

2.For the atoms in an ion, the sum of the oxidation numbers is equal to the charge on the ion.

3. In compounds, the group 1A metals all have an oxidation number of +1 and the group 2A metals all have an oxidation number of +2.

4. In compounds, the oxidation number of fluorine is –1.

Rules for Assigning Oxidation Numbers


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5. In compounds, hydrogen has an oxidation number of +1.

6. In most compounds, oxygen has an oxidation number of –2.

7. In binary compounds with metals, group 7A elements have an oxidation number of –1, group 6A elements have an oxidation number of –2, and group 5A elements have an oxidation number of –3.

8. Elements in their standard state (solid, liquid, or gas – look at PT) have an oxidation number of 0.

Rules for Assigning Oxidation Numbers


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• 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.

• Examples of redox reactions: displacement of an element by another element; combustion; incorporation of an element into a compound, etc.

Identifying Oxidation–Reduction Reactions


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LEO - GER


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Example 4.7What are the oxidation numbers assigned to the atoms of each element in

(a) KClO4 (b) Cr2O72– (c) CaH2 (d) Na2O2 (e)

Fe3O4


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Example 1• Which is oxidized and which is reduced in

the following equation:Fe2O3 (s) + 3CO (g) → 2Fe (s) + 3CO2 (g)


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Example 2: • Which is reduced/oxidized?

MnO2 + 4H+1 + 2Cl-1 → Mn+2 + 2H2O + Cl2


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• The maximum oxidation number of a nonmetal is equal to the group number.– For nitrogen, +5.– For sulfur, +6.– For chlorine, +7.

• The minimum oxidation number is equal to the group number – 8.

Oxidation Numbers of Nonmetals


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In the activity series, any metal above another can displace that other metal.

Activity Series of Some Metals

Mg metal can react with …

… Cu2+ ions to form Cu metal.

Will lead metal react with Fe3+ ions?

Will iron metal dissolve in an acid to

produce H2 gas?


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

• In foods and nutrition: redox reactions “burn” the foods we eat; antioxidants react with undesirable free radicals.

• In industry: to produce iron, steel, other metals, and consumer goods.

Applications of Oxidationand Reduction


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Redox Titrations• In a redox titration, one reactant (often the

titrant) is an oxidizing agent and the other reactant is a reducing agent.

• Permanganate ion, usually from KMnO4, is one of the most commonly used oxidizing agents in the chemical laboratory and makes an excellent titrant.

• The math is very similar to that used in acid-base titrations.


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Calculation Example 1• Example 4.12: A 0.2865 g sample of an iron ore is dissolved in

acid, and the iron is converted entirely to Fe+2 (aq). To titrate the resulting solution, 0.02645 L of 0.02250 M KMnO4 (aq) is required. What is the mass percent of iron in the ore?

5Fe+2 (aq) + MnO4-1 (aq) + 8 H+1 (aq) → 5Fe+3 (aq) + Mn+2 (aq) + 4H2O (l)


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Calculation Example 2• Example 4.12A: Suppose the titration in Example 4.12 was

carried out with 0.02250 M K2Cr2O7 (aq) rather than KMnO4 (aq). What volume of K2Cr2O7 (aq) would be required?

6Fe+2 (aq) + Cr2O7-2 (aq) + 14H+1 (aq) → 6Fe+3 (aq) + 2Cr+3 (aq) + 7H2O (l)


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Stop!

•Complete Worksheet #5 by next class period and finish your titration lab report!


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Today…• Turn in:• Get out WS#5 to Check• Put Lab Report in basket (rubric on top)

• Our Plan:• Questions on WS#5• Worksheet Race• Crash Course Video• Begin Report for Investigation 8

• Homework (Write in Planner):• Lab Report Due Friday


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Crash Course Water & Solutions

• Good introduction to redox lab:• http://

www.youtube.com/watch?v=AN4KifV12DA

http://www.youtube.com/watch?v=AN4KifV12DA




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Today…• Turn in:• Nothing

• Our Plan:• Investigation 8• Test Review

• Homework (Write in Planner):–Test Next Class–Breakfast Club Friday at 6:30 am!


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Today…• Turn in:• Lab Report – rubric on top

• Our Plan:• Questions on Test Review• Unit 2 Test

• Homework (Write in Planner):• Watch the Alkanes Notes online and

complete the booklet to p. 7


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