week 12: lectures 34 – 36 chemical reactionscourses.chem.psu.edu/chem110fall/lecture...

Week 12: Lectures 34 – 36

Lecture 34: W 11/9

Lecture 35: F 11/11

Lecture 36: M 11/14

Reading:

BLB Ch 3.1 – 3.2; 3.6 – 3.7; 4.2 – 4.4

Homework:

BLB 3: 1, 64; 4: 24, 39; Supp Rxns: 1 – 11

Reminder:

No Angel Quiz on Thur 11/10

ALEKS Objective 12 due on Tues 11/15

Jensen Office Hour: 501 Chemistry Building

Tuesdays and Thursdays 10:30 – 11:30 am

Late drop deadline: Friday 11/11 @11:59 pm

Final Exam: Monday Dec 12 2:30 – 4:20 pm !"#$%&'$("))$*+$,-*(.$!&/+0&*#$*$1*2,'3$

4,5',)6$7&8($98&($:+;,0$<=$>-*($?@A,#'0,$$

Jensen Chem 110 Chap 3 & 4 Page: 2

Chemical Reactions

Law of conservation of mass: total mass does not

change during a chemical reaction

• Mass of reactants MUST equal mass of products

# of atoms of each element on reactant side

= # of atoms of each element on product side

Example: Complete combustion of pentane

C5H12 + 8 O2 ! 5 CO2 + 6 H2O

If mass of reactants is 100 g, then the mass of

products must be ______ g.

Reactants: ____C, ____H, ____O atoms

Products: ____C, ____H, ____O atoms

___ mol C5H12 react completely with ____ mol O2,

produce ____ mol CO2 and ____ mol H2O.


Balancing Chemical Equations

When reactants and products are both given,

chemical reactions are balanced by changing

the _____________.

Four “Easy” Rules

1. Write the unbalanced molecular equation

correctly (molecules involved with correct

molecular formulas)

C6H6 + O2 ! CO2 + H2O

2. Balance the atoms of one element.

C6H6 + O2 ! CO2 + H2O

3. Balance atoms of remaining elements

C6H6 + O2 ! CO2 + H2O

4. Check you work! Make sure that you use the

smallest whole numbers.


Example: When the following equation

C4H8O2(l) + O2(g) ! CO2(g) + H2O(g)

is balanced with the smallest possible set of

integer coefficients, the coefficient of O2 is

A. 1

B. 2

C. 3

D. 5

E. 6

How many moles of water will be produced if

3 moles of C4H8O2 burn completely in air?


Patterns of Reactivity

1. Combination reactions (Chapter 3)

Elements react to form compounds

2 Mg (s) + O2 (g) ! 2 MgO (s)

Small compounds combine to form larger ones

MgO (s) + CO2 (g) ! MgCO3 (s)

2. Decomposition reactions (Chapter 3)

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

CaCO3 (s) ! CaO (s) + CO2 (g)

3. Complete Combustion reactions (Chapter 3)

CH4 (g) + 2 O2 (g) ! CO2 (g) + 2 H2O(g)

All hydrocarbons will produce ___________ when

they undergo complete combustion reactions.

4. Single displacement reactions (Chapter 4)

Zn(s) + CuSO4(aq) ! ZnSO4(aq) + Cu(s)


Patterns of Reactivity

5. Exchange reactions (Chapter 4) (Double Displacement or Metathesis Rxn)

Exchange reactions only occur if there is a driving force

a. Precipitation

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

b. Neutralization (weak or non-electrolyte)

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

c. Gas formation

2 HCl(aq) + Na2S(aq) ! H2S(g) # + 2 NaCl(aq)

How do you predict the phase of the products?

Use Solubility Rules (Table 4.1)

How do you know what is happening?

Use Net Ionic Equation


Table 4.1: Aqueous Solubility of Ionic

Compounds (Table provided on your data sheet)

Soluble ionic compounds Important exceptions

compounds containing NO3– None

C2H3O2– None

Cl– compounds of Ag+, Hg22+, and Pb2+

Br– compounds of Ag+, Hg22+, and Pb2+

I– compounds of Ag+, Hg22+, and Pb2+

SO42– compounds of Sr2+, Ba2+, Hg2

2+, and Pb2+

Insoluble ionic compounds Important exceptions

compounds containing S2– compounds of the alkali metal cations and

NH4+, Ca2+, Sr2+, and Ba2+

CO32– compounds of NH4

+,

the alkali metal cations

PO43– compounds of NH4

+,

the alkali metal cations

OH– compounds of the alkali metal cations and

NH4+, Ca2+, Sr2+, and Ba2+

Simplified Rules (in WATER)

1. Almost all ammonium and alkali metal salts are

soluble.

2. Most nitrates, acetates, chlorides, bromides, iodides,

and sulfates are soluble.

3. Most sulfides, carbonates, phosphates, and

hydroxides are insoluble.


Demo: Precipitation reactions

Use solubility rules to predict reactions based on

the reactants provided. ___________ drives the

reaction.

Mix MgCl2 and NaOH. What happens?

Mix KI and NaOH. What happens?

Mix AgNO3 and NaCl. What happens?

Mix AgNO3 and KI. What happens?


Net Ionic Equation: involves only the ions or

molecules directly involved in the reaction

1st

: Start with balanced molecular equation

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

2

nd: Dissociate all soluble strong electrolytes

(strong acids, strong bases, soluble ionic salts) to

get the Complete Ionic Equation.

Subscripts for ions in the chemical formula (in the

“Molecular” equation) become Coefficients for

those ions in the Complete Ionic Equation.

3rd

: Identify spectator ions (ions that appear on both

sides of equation)

4th

: Eliminate all spectator ions to get the Net Ionic

Equation (ions and molecules that directly involved

in the reaction)


Practice Examples:

1. Which ions are spectator ions in the reaction

represented by the following molecular equation?

2AgNO3(aq) + CaCl2(aq) !"2AgCl(s) + Ca(NO3)2(aq)

A. Ag+, Cl!, and Ca2+

B. Cl! and Ca2+

C. Ag+ and NO3!

D. Ca2+ and NO3!

E. Ca2+

2. Mixing solutions of K2SO4(aq) and BaCl2(aq)

produces an insoluble salt. What is the identity of

the spectator ions?

A. K+, SO42-, Ba2+, Cl-

B. K+, SO42-

C. K+, Cl-

D. Ba2+, Cl-

E. Ba2+, SO42-


Neutralization (Acid–Base) reactions

acid + base ! salt + water

• Acids: donate H+(aq)

HCl (aq) !

• Bases: raises concentration of OH–(aq) ions

KOH (aq) !

• Salts: ionic compounds;

replaces H+ of acid with positive ion

e.g., HCl becomes KCl

• Table 4.2—you"ve got this table memorized

by now, right??


Examples of Acid–Base Neutralization

1. Strong acid – strong base neutralization

HCl (aq) + KOH (aq) ! KCl (aq) + H2O (l)

# Complete ionic equation

# Spectator ions

# Net ionic equation

2. Weak acid – strong base neutralization

HF (aq) + KOH (aq) ! KF (aq) + H2O (l)

# Complete ionic equation

# Spectator ions

# Net ionic equation


Practice Example:

What is the net ionic equation for the

reaction between H3PO4 (aq) and KOH (aq)?

A. H3PO4 (aq) + KOH (aq)

! K3PO4 (aq) + H2O (l)

B. H3PO4 (aq) + 3 KOH aq)

! K3PO4 (aq) + 3 H2O(l)

C. H+ (aq) + OH— (aq) ! H2O (l)

D. H+(aq) + KOH (aq) ! K+ (aq) + H2O(l)

E. H3PO4 (aq) + 3 OH— (aq)

! PO43— (aq) + 3 H2O(l)


Driving Force: Gas Formation

• direct production of gas (e.g., H2, CO2, H2S)

2 HCl(aq) + Na2S(aq) ! H2S(g) # + 2 NaCl(aq)

• production of weak acid which decomposes to a

gas (e.g., H2CO3)

Example: Sodium Bicarbonate + HCl

Molecular equation:

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

H2CO3(aq) ! CO2(g)+H2O(l)

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

Complete ionic equation

Na+(aq) + HCO3

$(aq) + H+(aq) + Cl$(aq) !

Na+(aq) + Cl$(aq) + CO2(g) + H2O(l)

Net ionic equation:

HCO3$(aq) + H

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

Spectator ions:


Single Displacement Reactions

(Oxidation-reduction; aka, redox reactions)

Redox reactions: Reactions where electrons

are transferred from one reactant to another

• Reduction reaction: gaining e–

X2 + 2 e– ! 2 X

–

e.g. Cl2 + 2 e– ! 2 Cl

–

• Oxidation reaction: losing e–

M ! Mn+

+ n e–

e.g. Fe ! Fe2+

+ 2 e–

• Oxidation and reduction are always linked

• Must be balanced: atoms/electrons/charge


Rules for determining Oxidation Numbers

1. The oxidation number of an atom of a pure

element is ______.

e.g. oxidation number for Cl2 is ___, for Fe is ___.

2. The oxidation number of a monatomic ion equals

_____________.

e.g. oxidation number for Cl– is ___, for Fe

2+ is ___.

3. Some elements have the same oxidation number

in almost all their compounds, and can be used as

__________ to determine the oxidation numbers of

other atoms in the compound.

4. The sum of oxidation numbers in a neutral

compound is ________; The sum of oxidation

numbers in a polyatomic ion equals the __________

on the ion.


Example: What is the oxidation number of

Mn (Manganese) in MnO4–?

Practice example: What is the oxidation

state of S in H2SO4?

A. +2

B. +4

C. +6

D. -2

E. -4


Redox Reactions

• Oxidizing reagents: Elements or compounds

that oxidize the other reactant.

e.g.: O2, halogens, H2O2, HNO3, Cr2O7–, MnO4

–

• Reducing agents: Elements or compounds

that reduce the other reactant.

e.g.: H2, C, metals

Oxidation numbers always change in redox reactions!

Example: Balance the reaction between solid

lead (II) oxide and ammonia gas to produce

nitrogen gas, liquid water, and solid lead.


Examples of Single Displacement Reactions

1. Metal + Salt

Zn (s) + CuSO4 (aq) ! ZnSO4 (aq) + Cu (s)

Ionic equation:

Spectator Ions:

Net ionic equation:

What is the reducing agent (what is oxidized)?

What is the oxidizing agent (what is reduced)?

2. Metal + Acid

Zn (s) + 2 HCl (aq) ! ZnCl2 (aq) + H2 (g)

What is the reducing agent (oxidized)?

What is the oxidizing agent (reduced)?


Activity Series: predicts whether a certain metal

will be oxidized by an acid or a salt


Using the Activity Series

Remember: An element that is __________

in the activity series will be oxidized by the

ions of elements below it.

• Metals in the series will always be oxidized

by the ions of elements _______

Zn(s) + AgNO3(aq) !

• Metals _______ H2 in series (e.g., Mg, Zn) will

be oxidized by an acid (e.g., HCl) to form H2

Na(s) + H2O(l) !

• Metals toward bottom are unreactive (e.g.,

Ag, Pt, Au); that is, the elemental form is most

stable

Au(s) + H2O(g) !


Problem Solving with Chemical Reactions

Basic skills:

• Avogadro"s number and the definition of mole

• How to calculate formula weight (molar mass)

• How to do the following conversions:

gram % mole gram % molecules

• What is meant by:

empirical formula molecular formula

We use these along with balanced chemical

reactions to solve problems in chemistry

& Use balanced chemical equation to connect: moles reactants % moles products

& Use balanced chemical equation and

conservation of mass to connect: grams reactants % grams products

Example: C2H5OH + 3 O2 ! 2 CO2 + 3 H2O

moles 1 mol 3 mol 2 mol 3 mol

1 mol of C2H5OH reacts with 3 mol of O2;

produces 2 mol of CO2 and 3 mol of H2O


Steps to solve stoichiometry problems

1. Write the balanced chemical equation (or

process)

2. Make a table & fill in given information

Recognize:

• what you know already and what you are

being asked for

• what connections will take you from the

knowns to unknowns

3. Make connections between measured properties

and the balanced equation

Mass moles

Volume moles

(solutions)

P, V, T moles

(gases)

4. Fill in table until you are able to solve the problem

5. Make sure your answer is REASONABLE

concentration

molar mass

Ideal Gas Law


How much CO2 does your car produce?

The combustion of octane (C8H18) in the

presence of excess oxygen yields CO2 and

H2O. If 2.6 kg of octane is consumed, how

many kg of CO2 will it produce?


Limiting Reagents

& Reactant that is consumed completely;

& Determines the final amount of product;

& Must start with a balanced reaction.

When reactants mixed in unbalanced proportions,

some are left over (the ones in excess, unreacted)

Be sure to test all reactants!!!

• making a ham sandwich analogous to a chemical

reaction


Example: If 36.6 g of C2H5OH reacts with

63.8 g of O2 to form CO2 and H2O, how

many grams of CO2 will be produced?

A. 26.0g

B. 43.2g

C. 58.5 g

D. 70.4 g

E. 100.4g


Percent Yield

Theoretical yield: the yield of product that results

when the limiting reagent is completely consumed

Actual yield: the yield you actually get in the real

world

Percent yield:

% yield = actual yield

theoretical yield

!

"#$

%&'100

• Calculation is just one more step beyond a

standard stoichiometry calculation

• NOTE: if you get a % yield >100% something is

wrong, you"ve just created matter!??!!

Example continued:

If you obtained 50.0 g of CO2 from your

reaction, what is the percent yield?


Practice Example: The combustion reaction

between 1.0 mole of C3H8 (g) and 1.0 mole of

O2 (g) goes to completion:

C3H8(g) + 5 O2(g) ! 3 CO2(g) + 4 H2O(g)

Which of these statements are true?

i. All of the C3H8 (g) is used up.

ii. 3.0 moles of CO2 (g) is formed

iii. 0.8 moles of H2O (g) is formed

A. i only

B. ii only

C. i and ii only

D. iii only

E. i and iii only


Practice Example: Lithium and nitrogen react

to produce lithium nitride as follows:

6 Li (s) + N2 (g) ! 2 Li3N (s)

If 5.00 g of each reactant undergo a

reaction with a 80.5% yield, how many grams

of Li3N are obtained from the reaction?

A. 6.73 g

B. 1.67 g

C. 8.36 g

D. 2.08 g

E. 2.79 g


Scratch Paper:

week 12: lectures 34 – 36 chemical reactionscourses.chem.psu.edu/chem110fall/lecture...

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