Chemistry I Honors—Unit 6Chemical Equations, Reactions, & Redox

I. Chemical Equations Describe chemical reactions Starting substances are called

reactants Ending substances are called products All chemical reactions must follow the

Law of Conservation of Matter by being balanced

Objectives #1-3: Introduction to Chemical Reactions, Reaction Interpretation, and Balancing

II. Interpreting Chemical EquationsA. Symbols Symbol Meaning

+ and→ yields∆→

heat added

(s) solid(l) liquid(g) gas↑ gas produced↓ solid produced

elec

electricity added

(aq) aqueous solution

Fe→

catalyst required(element varies)

II. B. Writing Unbalanced Equations “Liquid hydrogen peroxide decomposes to form

water vapor and pxygen gas in the presence of the catalyst manganese (IV) oxide.”

“Solid calcium carbide (CaC2) reacts with water to form ethyne gas and aqueous calcium hydroxide.”

Objectives #1-3: Introduction to Chemical Reactions, Reaction Interpretation, and Balancing

“Ethyne gas (C2H2) reacts with oxygen in the the presence of a flame to produce carbon dioxide gas and water vapor.”

“Aqueous solutions of lead (II) nitrate and sodium iodine react to form lead (II) iodide and aqueous sodium nitrate.”

Basic Procedures: Be sure all formulas are correct before attempting to balance Never balance by changing subscripts Use coefficients to balance Type and number of atoms on each side of

reaction must balance Coefficients used must be in the lowest ratio possible

III. Balancing Chemical Equations

_____H2O2 _____ H2O + ______O2

___CaC2 + ___H2O ___C2H2 + __Ca(OH)2

___C2H2 + ____O2 ____CO2 + _____H2O

____Pb(NO3)2 + ___NaI ___NaNO3 + ___PbI2

Objectives #1-3: Introduction to Chemical Reactions, Reaction Interpretation, and Balancing

Objective #4: Assignment of Oxidation NumbersPart I: Oxidation vs. Reduction

Oxidation is the loss of electrons; during this process the charge of a species increases

Reduction is the gain of electrons; during this process the charge of a species decreases

“OIL RIG” or “LEO the lion goes GER”

Objective #4: Assignment of Oxidation Numbers Example I: Solid magnesium is reacted with

oxygen gas in the air to produce solid magnesium oxide

Equation: 0 0 +2 -2 Mg (s) + O2 (g) 2 MgO (s)

*What is the magnesium doing? Mg Mg+2 + 2 e-1

*What is the oxygen doing? O + 2e-1 O-2

Which element has been oxidized? Mg Which element has been reduced? O

Objective #4: Assignment of Oxidation Numbers Example II: Water is added to produce sufficient

heat to react solid forms of aluminum and iodine. The resulting reaction produces solid aluminum iodide.

Equation: 0 0 ∆ +3 -1 2 Al(s) + 3 I2 (s) 2 AlI3 (s)

*What is the aluminum doing? Al Al +3 + 3 e -1

*What is the iodine doing? I + e -1 I -1

*Which element has been oxidized? Al*Which element has been reduced? I

Objective #4: Assignment of Oxidation Numbers

In general, during REDOX reactions,

Metals tend to lose electrons and are oxidized

Nonmetals tend to gain electrons and are reduced

Objective #4: Assignment of Oxidation Numbers

Part II: Utilization of Oxidation Number RulesSee text p.232-233

The “Big 4”: Group I elements are +1 Group II elements are +2 H is usually +1 O is usually -2

Remember: 1) Elements are always neutral (zero)! 2) The total of the oxidation numbers in a compound must be neutral (zero)!!

He NaCl Na2Cr2O7

Ca(ClO3)2 OF Mg3(PO4)2

CrO4 -2

Oxidation Number Examples:

Demo Redox Reaction:

NaI(s) + H2SO4 (l) + MnO2 (s)

I2 (g) + MnSO4 (aq) + Na2SO4 (aq) + H2O(l)

Writing Half-Reactions (charges and atoms must balance to in order to be conserved! )

Examples:

K K+1 + _____ (__________) S + _______ S-2 (__________) Mg Mg+2 + _______ (__________) _____F-1 ______+ F2 (__________)

Objective #5: Balancing Redox Reactions

Objective #5: Balancing Redox Reactions

Key Steps:1.Write half-reactions for the oxidation

and reduction sections of the reaction.

2. Balance all elements except hydrogen and

oxygen.3. Balance oxygen by using water.4. Balance hydrogen by using hydrogen ions.

5. Balance charge by adding electrons to the side that is deficient in electrons.6. Equalize electrons lost and gained by multiplying each half-reaction by an appropriate factor.7. Add together half-reactions and cancel like species.8. Check that atoms and charges balance.

MnO4 -1 + Fe -2 Fe +3 + Mn-2

Example #1:

Cr2O7-2 + Cl-1 Cr+3 +

Cl2

Example #2:

Ce+4 + H3AsO3 Ce+3 + H3AsO4

Example #3:

Example #4:I2 + OCl -1 IO3

-1 + Cl -1

Examples: Copper + silver nitrate silver + copper (II)

nitrate

Element Oxidized:_______ “The Box:” Element Reduced: _______ O: OA: Oxidizing Agent:________ Reducing Agent: ________ R: RA:

Balancing Redox Reactions

Objective #6-8: Oxidizing and Reducing Agents

Examples—see packet

Objective #6-8: Oxidizing and Reducing Agents

Summary:The charge of the element oxidized goes upThe charge of the element reduced goes downThe item oxidized is the reducing agentThe item reduced is the oxidizing agentA species that is the source of BOTH oxidation and reduction is said to be disproportionate.

Objective #6-8: Oxidizing and Reducing Agents

Oxidizing and Reducing AbilityExample Demo: Cu + AgNO3 Cu(NO3)2 + Agassignment of oxidation numbers: 0 +1 +5 -2 +2 +5 -2 0 Cu + AgNO3 Cu(NO3)2 + AgCu has been oxidized and therefore Cu is the reducing agentAg has been reduced and therefore AgNO3 is the oxidizing agent

The more easily a species can lose electrons, the greater its ability to be a reducing agent and cause another species to gain electrons.

A species that loses electrons readily is unlikely to gain electrons and be reduced; such a species would not cause another species to lose electrons readily and therefore would act as a poor oxidizing agent

Objective #6-8: Oxidizing and Reducing Agents Example: Na + FeCl3 NaCl + Fe assignment of oxidation numbers: 0 +3 -1 +1 -1 0 Na + FeCl3 NaCl + Fe _____ is oxidized Na _____ is reduced Fe

*______ is the reducing agent and therefore would act as a ______ oxidizing agent

Na, poor*______ is the oxidizing agent and therefore

would act as a _______ reducing agentFeCl3, poor

A. Synthesis ReactionsGeneral formula: A + B AB

B. Decomposition ReactionsGeneral formula: AB A + B

C. Single-Displacement ReactionsGeneral formula: A + BC AC + B

D. Double Replacement Reactions General Formula: AB + CD AD + CB

E. Combustion ReactionsGeneral Formula: Hydrocarbon + O2 CO2 + H2O

Obj. 9 & 10—Types of Chem Rxns

Recall that oxidation-reduction reactions involve the transfer of electrons

A. Synthesis Reactions General formula: A + B AB Examples: Nonmetal + oxygen nonmetal oxide S + O2 SO3

N2 + O2 NO2

Objective #9: Oxidation-Reduction Reactions

Metal + oxygen metal oxide Rb + O2 Rb2O Mg + O2 MgO

Nonmetal + sulfur nonmetal sulfide C + S CS2

S + O2 SO3 (additional info needed)

Metal + sulfur metal sulfide Rb + S Rb2S Mg + S MgS

Metal + halogen metal halide Na + Cl2 NaCl Ca + I2 CaI2

Metal oxide + water metal hydroxide (base) Na2O + H2O NaOH MgO + H2O Mg(OH)2

Nonmetal oxide + water acid SO3 + H2O H2SO4 (add. info. needed)

SO2 + H2O H2SO3

B. Decomposition ReactionsGeneral formula: AB A + BExamples:Decomposition of binary compounds 2

elements H2O H2 + O2

NaCl Na + Cl2

Decomposition of metal carbonates carbon dioxide + metal oxide

BaCO3 BaO + CO2

Na2CO3 Na2O + CO2

Decomposition of metal hydroxides water + metal oxide

NaOH H2O + Na2O

Ca(OH)2 H2O + CaO

Decomposition of metal chlorates oxygen +metal chloride

KClO3 KCl + O2

Ca(ClO3)2 CaCl2 + O2

Decompostion of acids water + nonmetal oxide

H2SO4 H2O + SO2

C. Single-Displacement Reactions General formula: A + BC AC + BExamples:High metal + compound low metal + compound Fe + CuSO4 Cu + FeSO4

Cu + AgNO3 Ag + Cu(NO3)2

Active metal + water hydrogen + (low electronegativity) metal hydroxide Na + H2O H2 + NaOH

Ca + H2O H2 + Ca(OH)2

Metal + acid hydrogen + salt Zn + HCl ZnCl2 + H2

Mg + H3PO4 H2 + Mg3(PO4)2

High halogen + compound low halogen +

compound F2 + NaCl Cl2 + NaF

Br2 + NaI I2 + NaBr

An activity series is a vertical listing of elements in terms of their chemical reactivity; elements that are more reactive are listed at the top and less reactive elements are listed near the bottom (SEE RXN. PACKET!!)

A reactive element can readily transfer its valence electrons to another element

In general, for a single replacement reaction to go to completion, the lone element in the reaction must be higher on activity series that the element in the compound it is trying to displace.

Objective #11: Using an Activity Series

Remember, however, that an activity series should only be used as a general guide for predicting simple replacement reactions (see Table 3 on p.286)

Examples: Predict if the following reactions will occur: Zn + H2O --› (assume Zn is +2 if rx.

occurs) No Rx. Sn + O2 --› (assume Sn is +4 if rx.

occurs) Rx. Occurs; SnO2 will form

Cd + Pb(NO3)2 (assume Cd has a +2 charge if rx. occurs)

Rx. occurs ; Cd(NO3)2 + Pb

Cu + HCl (assume Cu has a charge of +2 if rx. occurs) No Rx.

D. Double Replacement Reactions General Formula: AB + CD AD + CB

Type I: Formation of a Precipitate (precipitation)

Ionic compound + ionic compound aqueous solution +

precipitate

Pb(NO3)2 + NaI NaNO3 + PbI2(s)

Na2S + Pb(NO3)2 PbS(s) + NaNO3

Type II: Formation of a GasIonic compound + ionic compound gas + aqueous solution +

water

NH4Cl + NaOH NH4OH + NaCl

NH3 + H2O

Na2SO3 + HCl H2SO3 + NaCl

SO2 + H2O

Type III: Formation of Water (acid-base) Acid + Base water + salt*

NaOH + HCl H2O + NaCl

Ca(OH)2 + HCl H2O + CaCl2

*SALT = an ionic compound that does NOT contain H+ or OH-

E. Combustion ReactionsExamples:Element + oxygen oxide Mg + O2 MgO

Na + O2 Na2O

Hydrocarbon + oxygen carbon dioxide + water

CH4 + O2 CO2 + H2O

C9H18 + O2 CO2 + H2O

(see example in lecture guide)

Practice in Predicting the Products of Chemical Reactions

Part I Dissociation of Ionic Compounds Dissociation process: The separation of ions

that occurs when an ionic compound is dissolved in water.

Examples: CaCl2(aq)

Ca+2(aq) + 2Cl-1(aq)

Al(NO3)3(aq) Al+3

(aq) + 3NO3-1

(aq)

Objectives #12: Compounds in Aqueous Solutions

Part II Predicting Precipitation Use of the solubility table in lecture guide Examples:

Part III: Writing Net Ionic Equations Net Reaction vs. Spectator IonsExamples:

Objective #12: Compounds in Aqueous Solutions