chapter 4 types of chemical reactions -...

Chapter 4

Types of Chemical Reactions

Classifying Chemical Reactions by What Atoms Do

Classification of Reactions

+

Decomposition reaction

Single displacement reaction

+ +

Double displacement reaction

+ +

Synthesis reaction +

Classification of Reactions

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

2 H2 (g) + O2 (g) ---------> 2 H2O (g)

C2H4 (g) + H2O2 (aq) C2H6O2 (l)

Synthesis reaction 2 HgO (s) ---------> 2 Hg (l) + O2 (g)

CaCO3 (s) ---------> CaO (s) + CO2 (g)

2 NaCl (s) ---------> Cl2 (g) + 2 Na (l)

Decomposition reaction Cu (s) + 2 AgNO3 (aq) ---------> 2 Ag (s) + Cu(NO3)2 (aq)

2 Al (s) + Fe2O3 (s) ---------> Al2O3 (s) + 2 Fe (l)

Mg (s) + 2 HCl (aq) ---------> H2 (g) + MgCl2 (aq)

Single displacement reaction

Ba(NO3)2 (aq) + Na2SO4 (aq) ---------> BaSO4 (s) + 2 NaNO3 (aq)

PCl3 (l) + 3 AgF (s) ---------> PF3 (g) + 3 AgCl (s)

HCl (aq) + NaOH (aq) ---------> H2O(l) + NaCl (aq)

Double displacement reaction

Chemical Reactions Classified by

Reaction Type

Precipitation Reactions


Precipitation reactions are reactions in which a solid forms when we mix two solutions.

1) reactions between aqueous solutions of ionic compounds 2) produce an ionic compound that is insoluble in water 3) The insoluble product is called a precipitate.


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

No Precipitate Formation = No Reaction

KI(aq) + NaCl(aq) ➜ KCl(aq) + NaI(aq)

KI(aq)

NaCl(aq)

KCl(aq) + NaI(aq)

No precipitate forms, therefore, no reaction.

Process for Predicting the Products ofa Precipitation Reaction

1. Determine which ions are present in each aqueous reactant.

2. Determine formulas of possible products.

3. Determine solubility of each potential product in water.

4. If neither product will precipitate, write no reaction after the arrow.

5. If any of the possible products are insoluble, write their formulas as the products of the reaction using (s) after the formula to indicate solid. Write any soluble products with (aq) after the formula to indicate aqueous.

6. Balance the equation. Remember to only change coefficients, not subscripts

Solubility RulesCompounds Containing theFollowing Ions Are Mostly Soluble Exceptions

Li+, Na+, K+, NH4+ None

NO3-, C2H3O2

-, ClO4- None

Cl- , Br- , I- Ag+, Hg22+, Cu+, Pb2+

SO4 2- Ca2+, Sr2+, Ba2+,Pb2+

Compounds Containing theFollowing Ions Are Mostly Insoluble Exceptions

OH- Group I cations, Ca2+, Sr2+, Ba2+

S2- Group I cations, Ca2+, Sr2+, Ba2+, NH4+

CO32-, PO4

3- Group I cations, NH4+

Practice – Predict the products and balance the equation

K2CO3(aq) + NiCl2(aq) ➜

K2CO3(aq) + NiCl2(aq) ➜ 2 KCl (?) + NiCO3(?)

K2CO3(aq) + NiCl2(aq) ➜ 2 KCl (aq) + NiCO3(s)

K2CO3(aq) + NiCl2(aq) ➜ KCl (?) + NiCO3(?)


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

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

KCl(aq) + AgNO3(aq) ➜


Na2S(aq) + CaCl2(aq) ➜

Na2S(aq) + CaCl2(aq) ➜ NaCl(?) + CaS(?)

Na2S(aq) + CaCl2(aq) ➜ 2 NaCl(?) + CaS(?)

Na2S(aq) + CaCl2(aq) ➜ 2 NaCl(aq) + CaS(aq)

No Reaction !!!!!


(NH4)2SO4(aq) + Pb(C2H3O2)2(aq) ➜

(NH4)2SO4(aq) + Pb(C2H3O2)2(aq) ➜ NH4C2H3O2(?) + PbSO4(?)

(NH4)2SO4(aq) + Pb(C2H3O2)2(aq) ➜ 2 NH4C2H3O2(?) + PbSO4(?)

(NH4)2SO4(aq) + Pb(C2H3O2)2(aq) ➜ 2 NH4C2H3O2(aq) + PbSO4(s)

Ionic Equations

Equations that describe the material’s structure when dissolved are called complete ionic equations.

Aqueous strong electrolytes are written as ions.

Insoluble substances, weak electrolytes, and nonelectrolytes are written

as molecules.

Equations that describe the chemicals put into the water and the product molecules are called molecular equations.

2 KOH(aq) + Mg(NO3)2(aq) ➜ 2 KNO3(aq) + Mg(OH)2(s)

2K+(aq) + 2OH−(aq) + Mg2+(aq) + 2NO3−(aq) ➜ 2K+(aq) + 2NO3−(aq) + Mg(OH)2(s)

Ionic Equations

Ions that are both reactants and products are called spectator ions.

2 K+(aq) + 2 OH−(aq) + Mg2+(aq) + 2 NO3−(aq) ➜ 2 K+(aq) + 2 NO3−(aq) + Mg(OH)2(s)

An ionic equation in which the spectator ions are removed is called a net ionic equation.

2 OH−(aq) + Mg2+(aq) ➜ Mg(OH)2(s)

Write the ionic and net ionic equation

K2SO4(aq) + 2 AgNO3(aq) ➜ 2 KNO3(aq) + Ag2SO4(s)

2 Ag+(aq) + SO42−(aq) ➜ Ag2SO4(s)

2K+ (aq) + SO42-(aq) + 2Ag+ (aq) + 2NO3-(aq) ➜ 2K+ (aq) + 2NO3-(aq) + Ag2SO4(s)

Na2CO3(aq) + 2 HCl(aq) ➜ 2 NaCl(aq) + CO2(g) + H2O(l)

CO32−(aq) + 2 H+(aq) ➜ CO2(g) + H2O(l)

2Na+ (aq) + CO32-(aq) + 2H+ (aq) + 2Cl-(aq) ➜ 2Na+ (aq) + 2Cl-(aq) + CO2(g) + H2O(l)

Write the ionic and net ionic equation

Chapter 4 Acid/Base Reactions

Acids and Bases in Solution

Acids ionize in water to form H+ ions. (More precisely, the H+ from the acid molecule is donated

to a water molecule to form hydronium ion, H3O+)

Bases dissociate in water to form OH- ions. (Bases, such as NH3, that do not contain OH- ions, produce

OH- by pulling H+ off water molecules.)

In the reaction of an acid with a base, the H+ from the acid combines with the OH- from the base to make water.

The cation from the base combines with the anion from the acid to make a salt.

acid + base ➜ salt + water

Molecular Models of Selected Acids

Acid-Base ReactionsAlso called neutralization reactions because the acid and base neutralize each other’s properties

2 HNO3(aq) + Ca(OH)2(aq) ➜ Ca(NO3)2(aq) + 2 H2O(l)

Note that the cation from the base combines with the anion from the acid to make the water soluble salt.

H+(aq) + OH-(aq) ➜ H2O(l)

(as long as the salt that forms is soluble in water)

The net ionic equation for an acid-base reaction is

Common AcidsName Formula Uses Strength

Perchloric HClO4 explosives, catalysts Strong

Nitric HNO3 explosives, fertilizers, dyes, glues Strong

Sulfuric H2SO4 explosives, fertilizers, dyes, glue, batteries Strong

Hydrochloric HCl metal cleaning, food prep, ore refining, stomach acid Strong

Phosphoric H3PO4 fertilizers, plastics, food preservation Moderate

Chloric HClO3 explosives Moderate

Acetic HC2H3O2 plastics, food preservation, vinegar Weak

Hydrofluoric HF metal cleaning, glass etching Weak

Carbonic H2CO3 soda water, blood buffer Weak

Hypochlorous HClO sanitizer Weak

Boric H3BO3 eye wash Weak

Name Formula Common Name Uses Strength

Sodium Hydroxide NaOH Lye, Caustic Soda soap, plastic production, petroleum refining

Strong

Potassium Hydroxide

KOH Caustic Potash soap, cotton processing, electroplating

Strong

Calcium Hydroxide

Ca(OH)2 Slaked Lime cement Strong

Sodium Bicarbonate

NaHCO3 Baking Soda food preparation, antacids Weak

Magnesium Hydroxide

Mg(OH)2 Milk of Magnesia antacids Weak

Ammonium Hydroxide

NH4OH Ammonia Water fertilizers, detergents, explosives Weak

Common Bases

HCl(aq) + NaOH(aq) ➜ NaCl(aq) + H2O(l)

HCl(aq) NaOH(aq)

NaCl(aq) + H2O(l)

Write the molecular, ionic, and net-ionic equation for the acid-base reaction

HNO3(aq) + Ca(OH)2(aq) ➜

2H+ (aq) + 2NO3-(aq) + Ca2+ (aq) + 2OH-(aq) ➜ Ca2+ (aq) + 2NO3-(aq) + 2H2O(l)

HNO3(aq) + Ca(OH)2(aq) ➜ Ca(NO3)2(aq) + H2O(l)

2HNO3(aq) + Ca(OH)2(aq) ➜ Ca(NO3)2(aq) + 2H2O(l)

2H+(aq) + 2OH-(aq) ➜ 2H2O(l)

HCl(aq) + Ba(OH)2(aq) ➜


HCl(aq) + Ba(OH)2(aq) ➜ BaCl2(aq) + H2O(l)

2H+(aq) + 2OH-(aq) ➜ 2H2O(l)

2H+ (aq) + 2Cl-(aq) + Ba2+ (aq) + 2OH-(aq) ➜ Ba2+ (aq) + 2Cl-(aq) + 2H2O(l)

2HCl(aq) + Ba(OH)2(aq) ➜ BaCl2(aq) + 2H2O(l)

H2SO4(aq) + Sr(OH)2(aq) ➜

2H+ (aq) + SO42-(aq) + Sr2+ (aq) + 2OH-(aq) ➜ SrSO4 (s) + 2H2O(l)

2H+(aq) + SO42-(aq) + Sr2+ (aq) + 2OH-(aq) ➜ SrSO4 (s) + 2H2O(l)

H2SO4(aq) + Sr(OH)2(aq) ➜ SrSO4(s) + 2 H2O(l)


Chapter 4

Acid/Base Titrations

Titration

A solution’s concentration is determined by reacting it with another solution and using stoichiometry –

this process is called titration.

In the titration, the unknown solution is added to a known amount of another reactant until the reaction is just completed. At this point, called the endpoint,

the reactants are in their stoichiometric ratio.

The unknown solution is added slowly from an instrument called a burette.

Acid-Base Titrations

The difficulty is determining when there has been just enough of one solution (the titrant) added to complete the reaction.

In acid-base titrations, because both the reactant and product solutions are colorless, a chemical (indicator) is added that changes color when the solution undergoes large changes in

acidity/alkalinity.

At the endpoint of an acid-base titration, the number of moles of H+ equals the number of moles of OH- . This is the

equivalence point.

TitrationThe titrant is the base solution in the burette.

As the titrant is added to the flask, the H+ reacts with the OH– to form water. But there is still excess acid present so the color does not change.

At the titration’s endpoint, just enough base has been added to neutralize all the acid. At this point the indicator changes color.

Titration

The titration of 10.00 mL of HCl solution of unknown concentration requires 12.54 mL of 0.100 M NaOH solution to reach the end point.

What is the concentration of the unknown HCl solution?


L

mL

L of NaOH soln

mol of NaOH

L

mol

mol of HCl

mol

mol

mL of NaOH soln

mL of HCl soln

L of HCl soln Molarity =

mol of HCl

L of HCl soln

1.00 mol HCl1.00 mol NaOH

12.54 mL NaOH solution x x

x = mol HCl in the sample

0.100 mol NaOH1.000 L NaOH soln

0.001 L NaOH soln1.000 mL NaOH soln

The titration of 10.00 mL of HCl solution of unknown concentration requires 12.54 mL of 0.100 M NaOH solution to reach the end point.

What is the concentration of the unknown HCl solution?


1.25 x 10-3

10.00 mL HCl solution x = L HCl soln0.001 L HCl soln1.000 mL HCl soln

0.0100

Molarity of HCl solution = = 1.25 x 10-3 mol HCl0.0100 L HCl soln 0.125 M

What is the concentration of NaOH solution that requires 27.5 mL to titrate 50.0 mL of 0.1015 M H2SO4 ?

L

m

L of H2SO4 soln

mol of H2SO4

L

mol

mol of NaOH

mol

mol

mL of H2SO4 soln

mL of NaOH soln

H2SO4 (aq) + 2 NaOH (aq) ➜ Na2SO4 (aq) + 2 H2O (l)

L of NaOH soln

Molarity =mol of NaOH

L of NaOH soln

27.50 mL NaOH soln x = L NaOH soln

2.00 mol NaOH1.00 mol H2SO4

50.00 mL H2SO4 soln x x

x = mol NaOH in the sample

0.1015 mol H2SO41.000 L H2SO4 soln

0.001 L H2SO4 soln1.000 mL H2SO4 soln

0.1015

0.001 L NaOH soln1.000 mL NaOH soln

0.02750

Molarity of NaOH soln = = 0.1015 mol NaOH0.02750 L NaOH soln

0.3691 M

What is the concentration of NaOH solution that requires 27.50 mL to titrate 50.00 mL of 0.1015 M H2SO4 ?

H2SO4 (aq) + 2 NaOH (aq) ➜ Na2SO4 (aq) + 2 H2O (l)

Chapter 4

Gas-Evolving Reactions

Gas-Evolving Reactions

Some reactions form a gas directly from the ion exchange:

K2S(aq) + H2SO4(aq) ➜ K2SO4(aq) + H2S(g)

Other reactions form a gas by the decomposition of one of the ion exchange products into a gas and water.

K2SO3(aq) + H2SO4(aq) ➜ K2SO4(aq) + H2SO3(aq)

H2SO3 ➜ H2O(l) + SO2(g)

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

NaHCO3(aq)

HCl(aq)

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

Other reactions form a gas by the decomposition of one of the ion exchange products into a gas and water.

NaHCO3(aq) + HCl(aq) ➜ NaCl(aq) + H2CO3(aq)

H2CO3 ➜ H2O(l) + CO2(g)

Reactant ReactantExchange Product Gas Formed Example

Metal sulfide Metal hydrogensulfide Acid H2S H2S K2S (g) + HCl (aq)→

H2S (g) + KCl (aq)

Metal carbonate Metal hydrogencarbonate Acid H2CO3 CO2

K2CO3 (aq) + HCl (aq) → CO2 (g) + H2O (l) + KCl (aq)

Metal sulfite Metal hydrogensulfite Acid H2SO3 SO2

K2SO3 (aq) + HCl (aq) → SO2 (g) + H2O (l) + KCl (aq)

Ammonium salt Base NH4OH NH3KOH (aq) + NH4Cl (aq) → NH3 (g) + H2O (l) + KCl (aq)

Compounds that UndergoGas-Evolving Reactions

Practice – Predict the products and balance the equations

Na2CO3(aq) + 2 HNO3(aq) ➜

2 HCl(aq) + Na2SO3(aq) ➜

H2SO4(aq) + CaS(aq) ➜

2 NaNO3(aq) + H2O (l) + CO2(g)

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

CaSO4(aq) + H2S(aq)

“2 NaNO3(aq) + H2CO3”

“2 NaCl(aq) + H2SO3”

Chapter 4

Redox Reactions

Oxidation/Reduction Basic Definitionse-

X loses electrons Y gains electron

X is oxidized Y is reduced

X is the reducing agent Y is the oxidizing agent

X increases its oxidation number

Y decreases its oxidation number

Oxidation and Reduction - Symbolic Representation

Oxidation and Reduction at the Atomic Level

Oxidation/reduction reactions involve transferring electrons from one atom to another.

Also known as redox reactions

Many involve the reaction of a substance with O2(g).

4 Fe(s) + 3 O2(g) ➜ 2 Fe2O3(s)

Redox Reactions

Atoms in Elements-------> Ions in Compound

Combustion as Redox

2 H2(g) + O2(g) ➜ 2 H2O(g)

Redox without Combustion

2 Na(s) + Cl2(g) ➜ 2 NaCl(s)

Reactions of Metals with Nonmetals

Consider the following reactions:

4 Na(s) + O2(g) → 2 Na2O(s) 2 Na(s) + Cl2(g) → 2 NaCl(s)

The reactions involve a metal reacting with a nonmetal.

In addition, both reactions involve the conversion of free elements into ions.

Na2O = 2 Na+ + O2-

NaCl = Na+ + Cl-

Oxidation and Reduction

To convert a free element into an ion, the atoms must gain or lose electrons (of course, if one atom

loses electrons, another must accept them).

Atoms that lose electrons are being oxidized, atoms that gain electrons are being reduced.

2 Na(s) + Cl2(g) → 2 Na+Cl–(s) Na → Na+ + 1 e– oxidation Cl2 + 2 e– → 2 Cl– reduction

Chapter 4

Redox Reactions

Electron Bookkeeping

Electron BookkeepingFor reactions that are not metal + nonmetal, or do not involve O2, we need a method for determining how the

electrons are transferred.

Chemists assign a number to each element in a reaction called an oxidation state that allows them to determine

the electron flow in the reaction.

Even though they look like them, oxidation states are not ion charges!

Oxidation states are imaginary charges assigned based on a set of rules.

Ion charges are real, measurable charges.

Rules for Assigning Oxidation States (in order of priority)

1. Free elements have an oxidation state = 0.

In Na (s), Na = 0 ; In Cl2 (g), Cl2 = 0

2. Monatomic ions have an oxidation state equal to their charge.

In NaCl, Na = +1 and Cl = −1

3. (a) The sum of the oxidation states of all the atoms in a compound is 0.

Na = +1 and Cl = −1 in NaCl, (+1) + (−1) = 0


3. (b) The sum of the oxidation states of all the atoms in a polyatomic ion equals the charge on the ion.

In NO3–, N = +5 and O = −2 [3 x (2-) + 1 x (5+) = -1]

4. (a) Group I metals have an oxidation state of +1 in all their compounds.

(b) Group II metals have an oxidation state of +2 in all their compounds.


5. In their compounds, nonmetals have oxidation states according to the table below

Assign an oxidation state to each element in the following

Br2

K+

LiF

CO2

SO42−

Na2O2

Br = 0, (Rule 1)

K = +1, (Rule 2)

Li = +1, (Rule 4a) & F = −1, (Rule 5)

O = −2, (Rule 5) & C = +4, (Rule 3a)

O = −2, (Rule 5) & S = +6, (Rule 3b)

Na = +1, (Rule 4a) & O = −1 , (Rule 3a)

Determine the oxidation states of all the atoms in a propanoate polyatomic anion, C3H5O2–

There are no free elements or free ions in propanoate, so the first rule that applies is Rule 3b

(C3) + (H5) + (O2) = −1

Because all the atoms are nonmetals, the next rule we use is Rule 5, following the elements in order: H = +1 O = −2

(C3) + 5(+1) + 2(−2) = −1 (C3) = −2 C = −⅔

reduction

Oxidation and Reduction Another Definition

Oxidation occurs when an atom’s oxidation state increases during a reaction.

Reduction occurs when an atom’s oxidation state decreases during a reaction.

CH4 + 2 O2 → CO2 + 2 H2O-4 +4 0 -2

oxidation

Oxidation–Reduction

2 Na(s) + Cl2(g) → 2 Na+Cl–(s) Na is oxidized Cl is reduced

Na is the reducing agent Cl2 is the oxidizing agent

Oxidation and reduction must occur simultaneously.

If an atom “loses” electrons another atom must “take” them.

The reactant that reduces an element in another reactant is called the reducing agent.

The reactant that oxidizes an element in another reactant is called the oxidizing agent.

Assign oxidation states, determine the element oxidized and reduced, and determine the oxidizing agent and reducing

agent in the following reactions:

Sn4+ + Ca → Sn2+ + Ca2+

2 F2 + S → SF4

Sn4+ is being reduced; Sn4+ is the oxidizing agent. Ca is being oxidized; Ca is the reducing agent.

F is being reduced from F0 to F-;F2 is the oxidizing agent.

S is being oxidized from S0 to S+4;S is the reducing agent.

0

0 0 S +4

F -

reduction

Fe + MnO4− + 4 H+ → Fe3+ + MnO2 + 2 H2O

Assign oxidation states, determine the element oxidized and reduced, and determine the oxidizing agent and

reducing agent in the following reactions:

0 +3+7 +4

Fe is the reducing agent. MnO4− is the oxidizing agent.

oxidation

chapter 4 types of chemical reactions -...

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