Chapter 11Chemical Reactions

Hingham High SchoolMr. Dan Clune

All chemical reactions• Two parts:

• – what you start with

• what you end with• Reactants turn into the products.• Reactants ® Products

In a chemical reaction• The way atoms are joined is

changed.

• Atoms aren’t or .

In a chemical reaction• Can be described several ways:1. In a sentence

Copper with chlorine to copper (II) chloride.

2. In a word equation

Copper + chlorine ® copper (II) chloride

Symbols in equations-p.323

• () separates reactants from products

• Read “reacts ”• Plus (+) sign read “ ”

Cu + Cl2 CuCl2

Symbols used in equations

• (s) = • (g) = • (l) = • (aq) - dissolved in water,

an solution.

Cu(s) + Cl2(g) CuCl2(s)

Symbols used in equations

•­­↑ after product, indicates produced

•same as (g) - H2↑•¯­after product, indicates

produced •same as (s) - PbI2↓

Symbols used in equations

indicates reaction

shows that is supplied to the reaction

is - indicates a is supplied, in this case, platinum.

heat ,

Pt

What is a catalyst?

• up a reaction• Is NOT or

by the reaction.• Enzymes are biological or protein

catalysts.

Skeleton Equation

• Uses formulas to describe a reaction• doesn’t indicate how many.

•

Convert this to an equation

• Solid iron (III) sulfide reacts with gaseous hydrogen chloride to form iron (III) chloride and hydrogen sulfide gas.

Nitric acid dissolved in water reacts with solid sodium carbonate to form liquid water, carbon dioxide gas, sodium nitrate dissolved in water.

Now, read these:

• Fe(s) + O2(g) ® Fe2O3(s)

• Cu(s) + AgNO3(aq) ® Ag(s) + Cu(NO3)2(aq)

• NO2 (g) N2(g) + O2(g) Pt

Balancing Chemical Equations

Balanced Equation• Atoms can’t be created or destroyed• All the atoms we with

we must up with• A balanced equation has the

of each element on of the equation.

• C + O2 ® CO2

• This equation is already balanced• What if it isn’t?

C + OO® COO

• C + O2 ® CO• Need one more O in the .• Can’t change the ,

because it describes what it is (carbon monoxide in this example)

C + O® COO

• Must be used to make another • But where did the other C come from?

C +O

®C

OO

OC

• Must have started with two C•

C

+O

®C

OO

OC

C

Finding the number of atoms

• The subscript in front of an element is the number of atoms of that element / polyatomic.• Ex) CO2 B2(SO4)3

• C= O= B= SO4=

• A coefficient in front of the formula multiplies the amount of elements by the coefficient.• Ex) 3CO2 2B2(SO4)3

• C= O= B= SO4=

Finding the number of atoms

H2O H= O=

2H2O H= O=

B(NO3)2 B= NO3=

3B(NO3)2 B= NO3=

Pb3(PO4)4 Pb= PO4=

2Pb3(PO4)4 Pb= PO4=

Rules for balancing:

1. Determine the

for reactants and products.

2. Write a equation.

Rules for balancing:3. Count the of atoms

of each appearing on sides of the equation.

4. Balance the elements one at a time by adding (the numbers in front)-Save and until LAST!

Rules for balancing:

5. Check to make sure it is balanced.

6. Make sure the coefficients are in the possible ratio.

Don’t you ever…

• Never change a to balance an equation.• H2O is a different compound than H2O2

• Never put a coefficient in the of a formula

• 2 NaCl is okay, Na2Cl is not. X

Example

H2 + H2OO2 ®

R P

H

O

H2 + H2OO2 ® 22

AgNO3 + Cu ®­­­ Cu(NO3)2 + Ag

NO3

Cu

Ag

PR

Mg + N2 ® Mg3N2

R P

Mg

N

P + O2 ® P4O10

R P

P

O

Na + H2O ® H2 + NaOH

O

H

Na

PR

CH4 + O2 ® CO2 + H2O

O

H

C

PR

Section 8.2Types of Chemical Reactions

• OBJECTIVES:• Identify a reaction as combination,

decomposition, single-replacement, double-replacement, or combustion

Types of Reactions• 5 major types.• predict the products• predict whether or not they will happen at all

• How? We recognize them by their

#1 - Combination Reactions• Combine - put together• substances combine to

make one .

A + B ®­AB• Ca +O2 ®­CaO

• SO3 + H2O ® H2SO4

#1 - Combination Reactions

• We can predict the products if they are two elements.

Mg + N2 ®­

Write and balance

• Ca + Cl2 ®

• Fe + O2 ® iron (II) oxide

• Al + O2 ®• Remember that the first step is to write

the correct formulas• Then balance by using only

#2 - Decomposition Reactions• Decompose = fall apart• reactant falls apart into

elements or compounds.

AB A + B

#2 - Decomposition Reactions

•NaCl Na + Cl2

•CaCO3 CaO + CO2

•Note that is usually required to decompose

electricity

#2 - Decomposition Reactions• Binary compounds (made of 2 elements)

falls apart into its elements

• H2O

• HgO

electricity

#3 - Single Replacement Reactions• element

another (new dance partner)

• Reactants are an • Products will be a

element and different cmpd• Li + KCl ® K + LiCl

• F2 +2 LiCl ® 2LiF + Cl2

(Cations switched)

(Anions switched)

#3 Single Replacement Reactions• Metals replace other metals (and H)• K + AlN ®• Zn + HCl ® • Think of water as: HOH

• Metals replace first H, then combines w/ hydroxide (OH).

• Na + H2O ®

#3 Single Replacement Reactions• Sometimes, the reaction will happen:

Some chemicals are more “ ” than others• active replaces

active

The “Activity Series” of Metals• Lithium• Potassium• Calcium• Sodium• Magnesium• Aluminum• Zinc• Chromium• Iron• Nickel• Lead• Hydrogen• Bismuth• Copper• Mercury• Silver• Platinum• Gold

If the lone metal is the paired metal, replacement occur.

Ex) Li + NaCl ® Na + LiCl

If the lone metal is the paired metal, replacement will occur.

Ex) Na + LiCl ® Na + LiCl

Higher activity

Lower activity

The “Activity Series” of Halogens

Fluorine Chlorine Bromine Iodine

If the lone halogen is the paired

halogen, replacement occur.

2NaCl(s) + F2(g)

MgCl2(s) + Br2(g)

Higher Activity

Lower Activity

#3 Single Replacement Reactions Practice:

• Fe + CuSO4 ®

• Pb + KCl ®­­

• Al + HCl ®

• Lithium• Potassium• Calcium• Sodium• Magnesium• Aluminum• Zinc• Chromium• Iron• Nickel• Lead• Hydrogen• Bismuth• Copper• Mercury• Silver• Platinum• Gold

Higher activity

Lower activity

#4 - Double Replacement Reactions• things each

other.• Reactants must be two

compounds.

• NaOH + FeCl3 ®• positive ions change place

• NaOH + FeCl3 ®­­Fe+3 OH- + Na+1 Cl-1

=

Complete and balance:

• assume all of the following reactions actually take place:CaCl2 + NaOH ®

CuCl2 + K2S ®

KOH + Fe(NO3)3 ®

K2SO4 + BaF2 ®

How to recognize which type?• Look at the reactants:

El + El = Combination

Cpd = Decomposition

El + Cpd = Single replacement

Cpd + Cpd = Double replacement

Practice Examples:• H2 + O2 ®

• H2O ®­­­

• Zn + H2SO4 ®• HgO ® • KBr + Cl2 ®

• AgNO3 + NaCl ®

• Mg(OH)2 + H2SO3 ®

#5 – Combustion Reactions• Combustion means “ ”• Normally, a cpd composed of only C, H,

(sometimes O) is reacted with oxygen – called “burning”

• combustion, products are

Combustion Reaction Examples:

• C4H12 + O2 ®

• C6H12O6 + O2 ®

• C8H8 + O2 ®

SUMMARY: An equation...• Describes a rxn• Must be • only balance by changing

• special symbols to indicate physical state, catalyst or energy required, etc.

Reactions• 5 major types• We can tell what type they are by looking at

• Single Replacement happens based on the Series

Section 11.3p. 342Reactions in Aqueous Solution

Co(NO3)2

K2Cr2O7K2CrO4

NiCl2 CuSO4

KMnO4

Net Ionic Equations• Many reactions occur in water,

or solution• When dissolved in water, many ionic

cpds “dissociate”, or ,into cations & anions

• Now write ionic equation

Net Ionic Equations• Example (needs to be a double replacement reaction)

AgNO3 + NaCl AgCl + NaNO3

1. this is the full balanced equation2. next, write it as ionic equation by splitting

the cpds into their ions:Ag1+ + NO3

1- + Na1+ + Cl1-

AgCl(s) + Na1+ + NO31-

Solids do not split up.

Net Ionic Equations3. Crossing out ions that did not change

(called spectator ions)Ag1+ + Cl1- AgCl (s)

This is the

Predicting the Precipitate• Insoluble salt is a precipitate

• i.e. a solid• General solubility rules are found:

a) Table 11.3, p. 344 in textbook

Solubility Rules

BaCl2 + AgNO3 →

1. Break up into ions

2. Now write the net ionic equation.

NaCl + Ba(NO3)2 →

1. Break up into ions

2. Now write the net ionic equation.

PbCl2 (s) + Li2O

1. Break up into ions

2. Now write the net ionic equation.