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Chapter 4

CHEMICAL REACTIONS


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I. Types of chemical reactions

A) There are millions of reactions. We can

remember them all so what to we do?

B) Fall into several categories.

C) Your text divides them into three types.

1. Precipitation reactions (metathesis

reactions - 2 compounds exchange parts -

generally take place in water solutions oftwo ionic solids and a solid (ppt) forms.)

The general form of the equation for

such a reaction is: AB + CD AD + CB


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2)2) Acid-Base reactions. An acid substance

react with a substance called a base. Such

reactions involve the transfer of a proton (H+)between reactants.

3) Oxidation-reduction reactions. These

involve the transfer of electrons betweenreactants.

D) Ionic Theory of Solutions and Solubility

Rules1) A solution is a homogeneous mixture of

particles less than one nanometer in diameter

distributed evenly throughout.


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3) We will focus on solids dissolved in water,

but there are other solvents than water, and

solutions other than liquids - gaseous (cleanair), solid solutions (alloys like steel, solder,

brass, bronze.

4) Solutes which dissolve are eitherelectrolytes or nonelectrolytes.

a) electrolytes are substances which when

dissolved in water result in a solution whichconducts electricity. Moving charged particles

must be there. These charged particles are

called ions.


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There are strong and weak electrolytes.

THE LIGHT BULBS !!!

b) nonelectrolytes are substances whichwhen dissolved in water do not conduct

electricity - uncharged molecules are the

particles in solution.


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c) Soluble Salts - substances which result

from the neutralization of an acid by a

base and are soluble in water are ofinterest - since they dissociate in water to

give separate ions

NaCl(aq) Na+(aq) + Cl-

(aq)

CaCl2(aq) Ca2+

(aq) + 2 Cl-

(aq)

C12H22O11(aq) C12H22O11(aq)

molecules stay intact - nonelectrolyte


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5) The formation of an insoluble solid (a

precipitate) drives a chemical reaction.

6) For these reactions we can write threekinds of equations.

MOLECULAR, IONIC, AND

NET IONIC EQUATIONSa) Molecular Equations-complete

formulas are written for all the reactants

and products, no ions are written.b) Ionic equations-all strongly soluble

electrolytes are written in their

dissociated (ionized) forms.


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c) Net Ionic equations-only involve those

chemical species which are involved in a

chemical reaction. All spectator ions areeliminated.

Spectator ions-those ions which do not

participate in the chemical reaction butare present in the reaction mixture.

d) Examples

Write the molecular, ionic, and net ionicequations for the reaction of an aqueous

solution of CaCl2 and an aqueous solution

of Na2CO3.


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1) the molecular equation is:

CaCl2(aq) + Na2CO3(aq) CaCO3(s)+2NaCl(aq)

2) the ionic equation is:

Ca2+(aq) + 2 Cl-(aq) + 2 Na

+(aq) + CO3

2-(aq)

CaCO3(s) + 2 Na+

(aq) + 2 Cl-(aq)

3) the net ionic equation is:

Ca2+(aq) + CO32-

(aq) CaCO3(s)


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D) PREDICTING METATHESIS

REACTIONS (PRECIPITATION

REACTIONS) USING SOLUBILITYRULES

The solubility rules are used to determine

whether precipitation reactions occur ornot.

1) you need to know whether the

predicted products are soluble or

insoluble in water. A precipitate is an

insoluble compound formed during a

chemical reaction in solution.


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9)The solubility rules are given on page

128 - table 4.1


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You do not have to memorize these rules,

I will give them to you on a quiz or exam.

You must know how to use the chart inorder to predict whether or not a

precipitate will occur and the formula of

the precipitate.

Examples:

If we add a solution of KCl to a AgNO3solution will a precipitate form? We first

write a predicted metathesis reaction.


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

Then we look at the solubility table to see

if any of the products are insoluble inwater.

We see that the table indicates that AgCl

is insoluble - most chlorides are solubleexcept for Ag+, Hg2+, Hg2

2+ and Pb2+

The molecular equation then becomes:

KCl(aq) + AgNO3(aq) KNO3(aq)+ AgCl(s)The net ionic equation is:

Cl-(aq) + Ag+(aq) AgCl(s)


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If we add a solution of NaNO3 to an

NH4Cl solution will a precipitate form?

The predicted equation then becomes:NaNO3(aq) + NH4Cl(aq) NaCl + NH4NO3

We see that the table indicates that both

compounds are soluble - most chlorides

are soluble, NaCl is soluble, and nitrates

are soluble as well as NH4

+ compounds, so

mixing these two solutions gives no

precipitates, no reaction results. After the

we would write N. R. for no reaction.


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What will result if we add a solution of

Pb(NO3)2 to a solution of KI?

II) ACID-BASE REACTIONS

A) There are three definitions of acids

and bases.1) The Arrhenius definition - which is

based on properties of these substances:

in water solution, acids taste sour, turn

blue litmus to red, react with active

metals to give off hydrogen gas,


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bases have a bitter taste, turn red litmus

to blue, feel soapy (slippery) and

neutralize acids.The result of these observations led

Arrhenius to the theory that acids ionize

in water to give H+

ions and bases giveOH- ions.

HCl(g) dissolves in water to give H+

(aq) and

Cl-(aq) ions. Therefore an Arrhenius acid.

NaOH(s) dissolves in water to give Na+

(aq)

ions and OH- (aq) ions. Therefore an

Arrhenius base.


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The neutralization reaction then

becomes:

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

Sodium hydroxide and Nitric acid

Potassium hydroxide and Sulfuric acid2) The Brnsted - Lowry Theory of acids

and bases.

A B-L acid is a species which donates aproton (H+ ion) to another species.


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A base is a species which accepts a proton

from another species. You must have a

reaction in order to name a species as aB-L acid or base.

HCl(aq) + H2O (l) H3O+

(aq) + Cl-(aq)

HCl is the B-L Acid, it donates the proton

to the H2O which is the B-L Base since it

accepts the proton (H+ion)


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NH3(g) + H2O(l) NH4+

(aq) + OH-(aq)

H2O is the B-L Acid, it donates the protonto the NH3 which is the B-L Base since it

accepts the proton (H+ion).

We will discuss the third theory of acidsand bases, the Lewis theory, later.

B) Strong and weak acids and bases

1) A strong acid is an acid that ionizes

completely in water; it is a strong

electrolyte.


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HCl(aq) + H2O (l) H3O+

(aq) + Cl-(aq)

At this time you should learn 6 strong

acids and bases:


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2) A weak acid is an acid that only

partially ionizes in water; it is a weak

electrolyte.The hydrogen cyanide molecule, HCN,

reacts with water to produce a small

percentage of ions in solution.

HCN(aq) + H2O(l) CN-(aq) + H3O

+(aq)

Table 4.1 lists some common weak acids,remember, if you memorize the strong

ones, all others are weak.


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3) A strong base is a base that is present

entirely as ions, one of which is OH-.


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4) Neutralization reactions involving

weak acids.


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Since KCN is a strong electrolyte, the net

ionic equation is

5) There are acid-base reactions which

give a gas as a product.


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The net ionic equation would be:

SO32-

(aq) + 2H+

(aq) H2O(l) + SO2(g)


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III) REDOX REACTIONS

A) When electrons are transferred from

one compound to another,an oxidation-reduction is said to occur.

B) The species which loses electrons is

said to be oxidized. The oxidationnumber of the species increases. This is

an oxidation reaction.

C) The species which gains electrons issaid to be reduced. The oxidation

number of the species decreases. This is

a reduction reaction.


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Oxidizing agents are themselves reduced.

They oxidize the other reactant.

Reducing agents are therefore oxidized.They reduce the other reactant.


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D) OXIDATION NUMBERS

Simply a bookkeeping method for

studying redox reactions.E) Rules for determining oxidation

numbers.

1. The oxidation number of any freeelement is zero, regardless of how

complex the allotrope may be.

2. The oxidation number for any simple,

monatomic ion is equal to the charge of

the ion.


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3. The sum of all the oxidation numbers

of the atoms in a molecule or polyatomic

ion must be equal to the charge of thesubstance.4. In its compounds, fluorine always has

an oxidation number of1.5. In its compounds, hydrogen has an

oxidation number of+1 except for

hydrides, where the oxidation number is-1.


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6. In its compounds, oxygen has an

oxidation number of-2 except for the

peroxides and superoxides, which haveoxidation numbers of -1 and -1/2

respectively.

7. Examples: Find the oxidation

numbers of all of the species present in

the following compounds.

KMnO4

Na2Cr2O7


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4 Fe(s) + 3 O2(g) 2 Fe2O3(s)

Feo Fe3+ + 3 e- (loss of electrons is

oxidation)

4 e + O2o

2 O2-

(gain of electrons isreduction)

Fe(s) is oxidized - it gives electrons - it is

the reducing agent.

O2(g) is reduced - it takes electrons - it is

the oxidizing agent.

F. An example:


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G. Most oxidation-reduction reactions

fall into one of the following simple

categories:

1. Combination

2. Decomposition


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4. Combustion

3. Displacement


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Balancing Redox equations on separate

sheets.

IV. Solutions

A) to describe the relative amounts ofsolute (that which is in the smaller

amount in the solution) and the solvent

(greater amount) the general term used isCONCENTRATION.


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B) Chemists like numbers, so we have

quantitative descriptions of

concentrations of solutions, one of whichis MOLARITY which MUST be related

to moles.

The important idea here is that when wemeasure out a certain volume of solution

we know how many moles of solute we

are getting.


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The number of grams of solute we are

pouring out of the container, is related to

the number ofPARTICLES OF SOLUTEwe have in our new container.

2) Examples: If we have 1.00 L of a 3.00

M solution of HCl, how many moles ofHCl are in the liter of solution?

How many grams of HCl are in 1.00 L

of 3.00 M HCl?


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If we have 500. mL of 6.00 M HCl

solution, how many moles do we have?

How many moles of HCl are in 25.0 mL of

6.00 M HCl solution?

3) Preparing solutions of a desired M

a) Calculate the amount of solute needed

and weigh out that amount.

b) Obtain the correct volumetric flask.


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c) Put some water into the flask, add the

solute and dissolve.

d) Fill the flask with solvent to thescratch on the flask and shake to

homogenize.

4) How would you prepare 250.00 mL of6.00 M aqueous solution of NaOH from

solid NaOH?

a) The calculation of the amount ofNaOH required is found as follows:


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b) Obtain the 250.00 mL volumetric flask.

c) Add a bit of water, add 60.0 g of NaOH,

dissolve.

d) Fill to scratch with water and shake to

homogenize.


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MAKE CERTAIN THAT YOU

REALIZE THAT YOU DID NOT

MEASURE THE AMOUNT OF WATERADDED.

C) Many times we start with

concentrated solutions to make dilute

solutions.

1) We must calculate the volume of the

concentrated solution required to make

the requested amount of dilute solution

and measure its volume.


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2) Obtain the correct volumetric flask.

3) put some water in the flask and a

measured amount of concentratedsolution.4) add solvent to the scratch and shake to

homogenize.D) To calculate the amount of

concentrated solution required, the major

concept to keep in mind is that the molesof solute obtained from the concentrated

solution are equal to the moles of solute in

the dilute solution.


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Upon dilution, the moles of solute do not

change only the amount of solvent

changes.

moles of solute in the conc. solution =

moles of solute in the dilute solution.

Mc X Lc or Mc X Vc


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Md X Ld or Md X Vd

Mc X Vc = Md X Vd

How would you prepare 250.0 mL of0.100M H2SO4 from 6.00 M H2SO4?


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1) first calculate the amount of

concentrated sulfuric acid (6.00 M) we

need.

2) obtain a 250.00 mL volumetric flask.


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3) put some water in the flask and add the

4.17 mL H2SO4 from 6.00 M bottle.

4) fill to the scratch with water andhomogenize.

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