Chapter 3 Chemical Reactions - Suffolk County Community ... · PDF fileconserved in a chemical reaction, chemical ... • In balancing a chemical equation, ... Ba(OH) 2 base H 3 PO

Jeffrey Mack

California State University,

Sacramento

Chapter 3 Chemical Reactions

Reactants: Zn + I2 Product: ZnI2

Chemical Reactions

Evidence of a chemical reaction:

• Gas Evolution

• Temperature Change

• Color Change

• Precipitation (insoluble species forms)

In general, a reaction involves a rearrangement

or change in oxidation state of atoms from

reactants to products.

Chemical Reactions

Chemical Equations show:

• the reactants and products in a reaction.

• the relative amounts in a reaction.

Example:

4 Al(s) + 3 O2(g) 2 Al2O3(s)

• The numbers in the front are called

stoichiometric coefficients

• The letters (s), (g), (l) and (aq) are the

physical states of compounds.

Chemical Equations

Notice the stoichiometric coefficients and the physical states of

the reactants and products.

Reaction of Phosphorus with Cl2

Notice the stoichiometric coefficients and the physical states of

the reactants and products.

Reaction of Iron with Cl2

4 Al(s) + 3 O2(g) 2 Al2O3(s)

This equation states that:

4 Al atoms + 3 O2 molecules

react to form 2 formula units

of Al2O3

or...

4 moles of Al + 3 moles of

O2 react to form 2 moles of

Al2O3

Chemical Equations

Law of the

Conservation of Matter

• Because the same

number of atoms are

present in a reaction at

the beginning and at

the end, the amount of

matter in a system

does not change.

2HgO(s) 2 Hg(l) + O2(g)

Chemical Equations

• Since matter is

conserved in a chemical

reaction, chemical

equations must be

balanced for mass!

• In other words, there

must be same number of

atoms of the each kind

on both sides of the

equatoin. Lavoisier, 1788

Chemical Equations

Steps in balancing a chemical reaction using coefficients:

1. Write the equation using the formulas of the reactants

and products. Include the physical states (s, l, g, aq

etc…)

2. Balance the compound with the most elements in the

formula first using integers as coefficients.

3. Balance elements on their own last.

4. Check to see that the sum of each individual elements

are equal on each side of the equation.

5. If the coefficients can be simplified by dividing though

with a whole number, do so.

Balancing Chemical Reactions

C2H6 + O2 CO2 + H2O

2 C’s & 6 H’s 2 O’s 1 C & 2 O’s 2 H’s & 1 O

balance last

Balancing Chemical Equations: Example

C2H6 + O2 CO2 + H2O

2 C’s & 6 H’s 2 O’s 1 C & 2 O’s 2 H’s & 1 O

balance last

3

balance H first

___C2H6 + O2 CO2 + ___ H2O

This side has an odd # of

O-atoms

This side will always have

an even # of O-atoms


C2H6 + O2 CO2 + H2O

2 C’s & 6 H’s 2 O’s 1 C & 2 O’s 2 H’s & 1 O

balance last

3

balance H first

___C2H6 + O2 CO2 + ___ H2O 2


C2H6 + O2 CO2 + H2O

2 C’s & 6 H’s 2 O’s 1 C & 2 O’s 2 H’s & 1 O

balance last

3

balance H first

___C2H6 + O2 CO2 + ___ H2O 2

balance C next

2C2H6 + O2 ___ CO2 + 6H2O 4


C2H6 + O2 CO2 + H2O

2 C’s & 6 H’s 2 O’s 1 C & 2 O’s 2 H’s & 1 O

balance last

3

balance H first

___C2H6 + O2 CO2 + ___ H2O 2

balance C next

2C2H6 + O2 ___ CO2 + 6H2O 4

balance O

2C2H6 + ____ O2 4CO2 + 6H2O 7


C2H6 + O2 CO2 + H2O

2 C’s & 6 H’s 2 O’s 1 C & 2 O’s 2 H’s & 1 O

balance last

3

balance H first

___C2H6 + O2 CO2 + ___ H2O 2

balance C next

2C2H6 + O2 ___ CO2 + 6H2O 4

balance O

2C2H6 + ____ O2 4CO2 + 6H2O 7

4 C’s 12 H’s 14 O’s 4 C’s 12 H’s 14 O’s


___ Al(s) + ___ Br2(l) ___ Al2Br6(s)

Balancing Equations

___C3H8(g) + ___ O2(g)

___ CO2(g) + _____ H2O(g)

___B4H10(g) + ___ O2(g)

___ B2O3(g) + ___ H2O(g)

Balancing Equations: Practice

• Solid magnesium hydroxide reacts with hydrochloric

acid to form aqueous magnesium chloride and

water.

• Write the balanced chemical equation for this

reaction.


_ Mg(OH)2(s) + _ HCl(aq)




water.


reaction.




water.


reaction.

_ Mg(OH)2(s) + _ HCl(aq) _ MgCl2(aq) + _ H2O(l)




water.


reaction.


Balance with a coefficient of ―2‖ in front of both HCl

and water.




water.


reaction.


Balance with a coefficient of ―2‖ in front of both HCl

and water.

Mg, Cl, O and H are now balanced.

• What Scientific Principles are used in the process of

balancing chemical equations?

• What symbols are used in chemical equations:

gasses: _____

liquids: _____

solids: _____

aqueous species in solution: _____

• What is the difference between P4 and 4P in an

eq.?

• In balancing a chemical equation, why are the

reactant and product subscripts not changed?

Chemical Equations: Review

When writing chemical reactions one starts with:

Reactants products

N2(g) + 3H2(g)

2NH3(g)

Some reactions can also run in reverse:

2NH3(g)

N2(g) + 3H2(g)

Under these conditions, the reaction can be written:

2 2 33H (g) N (g) 2NH (g)

Double arrows indicate ―Equilibrium‖.

Chemical Equilibrium

Chemical Equilibrium

Once equilibrium is achieved, reaction continues, but there

is no net change in amounts of products or reactants.

• Salts (ionic compounds): Composed of a

metal and non metal element(s).

• Acids: Arrhenius definition

Produce H+(aq) in water

Examples: HCl, HNO3, HC2H3O2

• Bases: Arrhenius definition

Produce OH (aq) in water

Examples: NaOH, Ba(OH)2, NH3

Classifying Compounds

• Molecular Compounds:

• Covalently bonded atoms, not acids, bases or

salts.

• Compounds like alcohols (C2H5OH) or table

sugar (C6H12O6)

• These never break up into ions.


• Classify the following as ionic, molecular,

acid or base.

Compound Type

Na2SO4

Ba(OH)2

H3PO4

CH4

P2O5

NH3

HCN



• Classify the following as ionic, molecular,

acid or base.

Compound Type

Na2SO4 ionic

Ba(OH)2 base

H3PO4 acid

CH4 molecular

P2O5 molecular

NH3 base

HCN acid

Aqueous Solutions:

There are three types of aqueous solutions:

Those with Strong Electrolytes

Water as the solvent

Solution = solute + solvent

That which is dissolved

(lesser amount)

That which is dissolves

(greater amount)

Those with Weak Electrolytes

& those with non-Electrolytes

Reactions in Aqueous Solutions

Many reactions involve ionic compounds, especially reactions in water — aqueous solutions.

KMnO4 in water K+(aq) + MnO4-(aq)

Reactions in Aqueous Solutions

Ionic Compounds (CuCl2) in Water

When ions are present in water,

the solutions conduct

electricity!

Ions in solution are called

ELECTROLYTES

Examples of Strong Electrolytes:

HCl (aq), CuCl2(aq) and NaCl

(aq) are strong electrolytes.

These dissociate completely (or

nearly so) into ions.

Strong Electrolytes conduct

electricity well.

Strong Electrolyte

HCl(aq), CuCl2(aq) and NaCl(aq) are strong electrolytes.

These dissociate completely (or nearly so) into ions.

Strong Electrolytes

Acetic acid ionizes only to a small

extent, it is a weak electrolyte.

Weak electrolytes exist in solution

under equilibrium conditions.

The small concentration of ions

conducts electricity poorly.

Weak electrolytes exit primarily in

their molecular form in water.

3 2 3 2CH CO H(aq) CH CO (aq) H (aq)

Weak Electrolytes

Weak electrolytic solutions are characterized by

equilibrium conditions in solution:

When acetic acid dissociates, it only partially

ionizes. +

2 3 2 2 3 2HC H O (aq) H (aq) + C H O (aq)

The majority species in solution is acetic acid in its

molecular form.

When writing a weak electrolyte in solution, one

NEVER breaks it up into the corresponding ions!

95% 5%

+2 3 2 2 3 2HC H O (aq) H (aq) + C H O (aq)×

Weak Electrolytes

Acetic acid ionizes only to a small extent, so it

is a weak electrolyte.

CH3CO2H(aq) CH3CO2-(aq) + H+(aq)

Weak Electrolytes

Some compounds dissolve in

water but do not conduct

electricity.

They are non-electrolytes.

Examples include:

• sugar

• ethanol

• ethylene glycol

Non-electrolytes do not

dissociate into ions!

Non-Electrolytes

Strong electrolytes: Characterized by ions only (cations &

anions) in solution (water).

Weak electrolytes: Characterized by ions (cations & anions)

& molecules in solution.

Non-electrolytes: Characterized by molecules in solution.

Conduct electricity well

Conduct electricity poorly

Do not conduct electricity

Species in Solution: Electrolytes

Solutes in Aqueous Solutions

How do we know if a compound will be soluble in

water?

For molecular compounds, the molecule must be

polar.

We will discuss polarity later, for now I will tell you

whether or not a molecular compound is polar…

For ionic compounds, the compound solubility is

governed by a set of SOLUBILITY RULES!

You must learn the basic rules on your own!!!

Solubility Rules

Water Solubility of Ionic Compounds

If one ion from the ―Soluble Compound‖ list is present in a compound, then the compound is water soluble.

Precipitation Reactions: A reaction where an

insoluble solid (precipitate) forms and drops out

of the solution.

Acid–base Neutralization: A reaction in which an

acid reacts with a base to yield water plus a salt.

Gas forming Reactions: A reaction where an

insoluble gas is formed.

Reduction and Oxidation Reactions (RedOx): A

reaction where electrons are transferred from

one reactant to another.

Types of Reactions in a Solution

REDOX REACTIONS

EXCHANGE

Acid-Base

Reactions

EXCHANGE

Gas-Forming

Reactions

EXCHANGE: Precipitation Reactions

REACTIONS

EXCHANGE REACTIONS

The anions exchange

places between cations.

A precipitate forms if one of

the products in insoluble.

Pb(NO3) 2(aq) + 2 KI(aq)

PbI2(s) + 2 KNO3 (aq)

Chemical Reactions in Water

The ―driving force‖ is the formation of

an insoluble solid called a precipitate.

Pb(NO3)2(aq) + 2 KI(aq)

2 KNO3(aq) + PbI2(s)

BaCl2(aq) + Na2SO4(aq)

BaSO4(s) + 2 NaCl(aq)

Precipitates are determined from the

solubility rules.

Precipitation Reactions

Which species is the precipitate?

Pb(NO3)2(?) + 2KI(?) 2KNO3(?) + PbI2(?)



Pb(NO3)2(?) + 2KI(?) 2KNO3(?) + PbI2(?)

From the solubility rules:

All nitrate salts are soluble, therefore:

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

Precipitation reactions


Pb(NO3)2(?) + 2KI(?) 2KNO3(?) + PbI2(?)




All potassium salts are soluble, therefore:

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



Pb(NO3)2(?) + 2KI(?) 2KNO3(?) + PbI2(?)




All potassium salts are soluble, therefore:

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

By the solubility rules:

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

PbI2 is the ppt.


Molecular Equation: all species listed as formula units or in

molecular form. reactants products

• Note all states of each reactant or product by: (s), (l), (g) or

(aq)

Ionic Equation: All soluble (aq) species present are listed as

ions.

• Leave all (s), (l) or (g) species as is. They do not dissociate

into ions

Net Ionic Equation:

• From the ionic equation, cancel out any species that appear

on either side of the equation.

• These are known as the ―spectator ions‖ and they are

never part of a net ionic equation!

Net Ionic Equations

Molecular Equation:


Writing Net Ionic Equations

Molecular Equation:


Total Ionic Equation:

Pb2+ (aq) + 2NO3– (aq) + 2K+(aq) + 2I–(aq)

2K+(aq) + 2NO3– (aq) + PbI2(s)


Molecular Equation:



Pb2+ (aq) + 2NO3– (aq) + 2K+(aq) + 2I–(aq)

2K+(aq) + 2NO3– (aq) + PbI2(s)

Never break up

any (s), (l) or (g)

or molecular

(aq) species!


Molecular Equation:



Pb2+ (aq) + 2NO3– (aq) + 2K+(aq) + 2I–(aq)

2K+(aq) + 2NO3– (aq) + PbI2(s)

Never break up

any (s), (l) or (g)

or molecular

(aq) species!

Cancel out the spectator ions to yield the net ionic equation:

Pb2+ (aq) + 2I–(aq) PbI2(s)


Arrhenius Definition:

• An acid is any substance that increases the

H+(aq) concentration in an aqueous solution.

HX(aq) H+(aq) + X–(aq)

• A base is any substance that increases the

OH–(aq) concentration in an aqueous

solution.

MOH(aq) M+(aq) + OH–(aq)

Acids & Bases

Brönsted-Lowry:

• An acid is any substance that donates H+(aq)

to another species in an aqueous solution.

HX(aq) + H2O(l) H3O+(aq) + X–(aq)

• A base is any substance that accepts an

H+(aq) in an aqueous solution.

H+(aq) + NH3(aq) NH4+(aq)

H3O+(aq) = H+(aq)

Acids and Bases

Acids

Strong acids are almost completely ionized in

water. (strong electrolytes)

Examples:

HX (aq) (X = Cl, Br & I) hydro ___ ic acid

HNO3 (aq) nitric acid

HClO4 (aq) perchloric acid

H2SO4 (aq)* sulfuric acid

* Only the 1st H is strong, sulfuric acid dissociates via:

H2SO4 (aq) H+ (aq) + HSO4– (aq)

Strong Acids

An acid: H3O+ in water

Acids

Weak Acids are incompletely ionized in water.

(weak electrolytes) Weak acids are governed by

dynamic equilibrium.

Examples:

HC2H3O2 (aq)

nitrous acid

hydrosulfuric acid

hydrogen sulfate ion

See you text and home work for more examples.

Weak acids are always written in their molecular form.

acetic acid (vinegar)

HNO2 (aq)

H2S (aq)

HSO4–(aq)

Weak Acids

2H O( )NaOH(s) Na (aq) ΟΗ aq

Bases: A base is a substance that produces OH– (aq) ions in

water by dissociation in water:

Strong bases are almost completely ionized in aqueous

solution. (Strong electrolytes)

Examples: Hydroxides of Group 1 (MOH(aq) where M = Li,

Na, K ect…) and Ca, Sr, Ba.*

*Ca(OH)2, Sr(OH)2 & Ba(OH)2 are slightly soluble, but that

which dissolves is present as ions only.

Strong Bases

Base: OH- in water

NaOH(aq) Na+(aq) + OH-(aq)

NaOH is a

strong base

Bases

NH3 acts as a base by reacting with water:

NH3(aq) + H2O(l)

Weak Bases:

Ammonia can also accept H+ from an acid:

NH3(aq) + H+(aq)

NH4+(aq) + OH –(aq)

NH4+(aq)

Weak Bases

Ammonia, NH3

K+(aq) + Br– (aq)

Salt + Water (usually)

HA (aq) + MOH(aq)

Acid + Base

Strong acid - Strong base neutralization: HBr(aq)/KOH(aq)

Molecular Equation:


HBr(aq) + KOH(aq) KBr (aq) + H2O(l)

H+ (aq) + Br– (aq) + K+(aq) + OH– (aq) + H2O(l)

Net Ionic equation:

H+ (aq) + OH– (aq) H2O (l)

/ / / /

MA(aq) + HOH(l)

Reactions of Acids & Bases: Acid-Base Neutralization

• The ―driving force‖ is the formation of water.

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

• Net ionic equation

OH-(aq) + H3O+(aq) 2 H2O(l)

• This applies to ALL reactions

of STRONG acids and bases.

Acid-Base Reactions


Reactions of weak acids and strong bases:

Molecular Equation:

HC2H3O2(aq) + NaOH(aq) NaC2H3O2(aq) + H2O(l)

HC2H3O2(aq) + Na+(aq) + OH–(aq) Na+(aq) + C2H3O2–(aq) + H2O(l)

Leave in

molecular

form

Net Ionic: HC2H3O2(aq) + OH–(aq) C2H3O2–(aq) + H2O(l)

/ /

Reactions of Acids & Bases: Acid-Base Neutralization

Nonmetal oxides can form acids in

aqueous solutions:

Examples:

CO2(aq) + H2O(s) H2CO3(aq)

SO3(aq) + H2O(s) H2SO4(aq)

Both gases come from the burning

of fossil fuels.

Non-Metal Acids

Metal oxides form bases in aqueous solution

CaO(s) + H2O(l) Ca(OH)2(aq)

CaO in water. Indicator

shows solution is basic.

Bases

Gas-Forming Reactions

Metal carbonate salts react with acids to the corresponding

metal salt, water and carbon dioxide gas.

2HCl(aq) + CaCO3(s) CaCl2(aq) + H2CO3(aq)

H2O(l) + CO2(g)

decomposes

Similarly:

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

acid base salt water

Neutralization!!!


Group I metals: Na, K, Cs etc.. react vigorously

with water

2K(s) + 2H2O(l) 2KOH(aq) + H2(g)

Metals & acid:

Some metals react vigorously with acid solutions:

Zn(s) + 2H+(aq) Zn2+(aq) + H2(g)


CaCO3(s) + H2SO4(aq) 2 CaSO4(s) + H2CO3(aq)

Carbonic acid is unstable and forms CO2 & H2O

H2CO3(aq) CO2 + water

(The antacid tablet contains citric acid + NaHCO3)


Thermite reaction:

Fe2O3(s) + 2Al(s)

2Fe(s) + Al2O3(s)

Oxidation-Reduction Reactions

REDOX = reduction & oxidation

O2(g) + 2 H2(g) 2 H2O(l)


• Oxidation involves a reactant atom or compound losing

electrons.

• Reduction involves a reactant atom or substance gaining

electrons.

• Neither process can occur alone… that is, there must be

an exchange of electrons in the process.

• The substance that is oxidized is the reducing agent

• The substance that is reduced is the oxidizing agent

Mg(s) + 2H+(aq) Mg2+(aq) + H2(g)

oxidized reduced

reducing

agent

oxidizing

agent


• Chemists use oxidation numbers to account for the transfer of electrons in a RedOx reaction.

• Oxidation numbers are the actual or apparent charge on atom when alone or combined in a compound.

1. The atoms of pure elements always have an

oxidation number of zero.

Examples: Mg(s)

Hg(l)

I2(s)

O2(g)

All have an

oxidation number

of zero (0)

Oxidation Numbers

2. If an atom is charged, then the charge is the

oxidation numbers .

Examples: Ion Oxidation Number

Mg2+(aq)

Cl (aq)

Sn4+(s)

22Hg (aq)

+2

1

+4

+2/2 = +1 for each Hg atom

Oxidation Numbers

3. In a compound, fluorine always has an oxidation numbers of 1.

4. Oxygen most often has an oxidation number of 2. » *When combined with fluorine, oxygen has a positive O.N.

» *In peroxide, the O.N. is 1.

5. In compounds, Cl, Br & I are 1 (Except with F and O present)

6. In compounds, H is +1, except as a hydride

(H : 1)

Oxidation Numbers

Examples:

compound Oxidation Numbers

HF(g) H = +1 F = 1

H2O(l) H = +1 O = 2

OF2(g) O = +2 F = 1

Na2O2(s) Na = +1 O = 1

HCl(g) H = +1 Cl = 1

NaH(l) Na = +1 H = 1

Oxidation Numbers

Most common oxidation numbers:

Oxidation Numbers

7. For neutral compounds, the sum of the oxidation numbers

equals zero.

For a poly atomic ion, the sum equals the charge.

Examples:

MgCl2

+2 + 2 × (−1) = 0

3 + 4 × (+1) = +1

4NH

Oxidation Numbers

Determine the oxidation number of iron in the

following compound:

Fe(OH)3

0 = 3 ( 1) ? +

Iron must have an oxidation number of +3!

Oxidation Numbers

In a RedOx reaction, the species oxidized and the

species reduced are identified by the changes in

oxidation numbers :

+ ++ ® + 22Ag (aq) Cu(s) 2Ag(s) Cu (aq)

Oxidation numbers:

+1 0

Oxidation numbers:

0 +2

Since silver goes from +1 to zero, it is reduced.

Since copper goes from zero to +2, it is oxidized.

The reaction is balanced for both mass and charge.

Recognizing a Redox Reaction

Practice:

Identify the species that is Oxidized and

Reduced by assigning oxidation numbers in the

following reaction.

2

4 2 7

3

3 2

3CH (g) Cr O (aq) 8H (aq)

3CH OH(l) 2Cr (aq) 4H O(l)

Answer:

Practice:



following reaction.

2

4 2 7

3

3 2



Answer:

• The carbon in methane (CH4) is oxidized ( 4 to 2)

2

4 2 7

3

3 2



Answer:

• Chromium in dichromate is reduced (+6 to +3)

Practice:



following reaction.

• The carbon in methane (CH4) is oxidized ( 4 to 2)

Redox Reactions

• Iron gains 3 electrons

(+3 to 0) oxidation

number change. It is

Reduced.

• Carbon loses 2

electrons (+2 to +4) it

is Oxidized.


REDOX = reduction & oxidation

Corrosion of aluminum

2 Al(s) + 3 Cu2+(aq) 2 Al3+(aq) + 3 Cu(s)

Redox Reactions

In all reactions if a

species is oxidized then

another species must

also been reduced

Cu(s) + 2 Ag+(aq) Cu2+(aq) + 2 Ag(s)

Redox Reactions

Cu(s) + 2 Ag+(aq) Cu2+(aq) + 2 Ag(s)

Redox Reactions

e e

• Two electrons leave copper.

• The silver ions accept them.

• The copper metal is oxidized to copper (II) ion.

• The silver ion is reduced to solid silver metal.

2Ag+(aq) + Cu(s) Cu2+(aq) + 2Ag(s)

Electron Transfer in a Redox Reaction

Manufacturing metals

Corrosion

Batteries

Fuels

Redox Reactions in Our World

Metal + halogen

2 Al + 3 Br2 Al2Br6

Examples of Redox Reactions

Metal (Mg) + Oxygen MgO

Nonmetal (P) + Oxygen P4O10


Metal + acid

Mg + HCl

Mg = reducing agent

H+ = oxidizing agent

Metal + acid

Cu + HNO3

Cu = reducing agent

HNO3 = oxidizing agent


• You have the following items available to you:

Deionized water, pH paper, test tubes various

metal nitrate salts, common acid and base

solutions.

• Suggest a simple test or set of tests for

identifying the unknown substances. Use

proper terminology and write balanced

chemical equations where applicable.

• Justify your answers thoroughly.

Reviewing What You’ve Learned

• How would you determine whether or not a

test tube containing a clear colorless solution

is water or sulfuric acid?

• Given a white powder that my be silver

chloride or sodium chloride.

• Whether a compound is silver nitrate or

sodium nitrate.3

Reviewing What You’ve Learned

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Chapter 3 Chemical Reactions - Suffolk County Community ... · PDF fileconserved in a chemical reaction, chemical ... • In balancing a chemical equation, ... Ba(OH) 2 base H 3 PO