
What is the difference between a chemical reaction and physical change?

When you watch a reaction occur, what are some hints that it is a chemical reaction?

Describing Chemical Reactions

Objectives




List three observations that suggest that a chemical reaction has taken place.
List three requirements for a correctly written chemical equation.
Write a word equation and a formula equation for a given reaction.
Balance a formula equation by inspection.

Chemical Reactions
 when a substance changes identity
 reactants- original
 products- resulting
 law of conservation of mass
 total mass of reactants = total mass of products

Chemical Reactions
 chemical equation

 represents identities and relative amounts of reactants and products in the chemical reaction uses symbols and formulas

Hints of Chemical Rxn
 heat or light
 can also happen with physical changes
 gas bubbles
 means a gas is being created as product
 precipitate
 solid is being created
 color change

Writing Chemical Equations

 most pure elements
 written as elemental symbol diatomic molecules



 molecule containing only 2 atoms some elements normally exist this way
H
2
, O
2
, N
2
, F
2
, Cl
2
, Br
2
, I
2 other exceptions
• sulfur: S
8
• phosphorus: P
4

Word Equations
 uses names instead of formulas
 helps you to write formula equation

Example

Description:
Solid sodium oxide is added to water at room temperature and forms sodium hydroxide.

Word Equation: sodium oxide + water  sodium hydroxide

Formula Equation:
Na
2
O + H
2
O  NaOH

Symbols Used in Equations yields reversible above arrow: or heat
MnO
2
25°C or Pt
2 atm heated catalyst specific T requirement specific P requirement after a formula:
(s) solid
(l) liquid
(aq) aqueous: dissolved in water
(g) gas



List three observations that suggest that a chemical reaction has taken place.

Acids you have to know!


2
4

3

3
4

2
3
2


Write the chemical equation from the following description:
Zinc metal is added to hydrochloric acid to create zinc chloride and hydrogen gas.

Aluminum reacts with oxygen to produce aluminum oxide
A.
B.
C.
Al + O  Al
2
O
3
Al + O
2
 Al
2
O
3
Al
3
+ O  Al
2
O
3

Aluminum reacts with oxygen to produce aluminum oxide
A.
B.
C.
Al + O  Al
2
O
3
Al + O
2
 Al
2
O
3
Al
3
+ O  Al
2
O
3

Phosphoric acid is produced through the reaction between tetraphosphorus decoxide and water
A. H
3
PO
4
B. H
3
PO
4
C. P
4
O
10

+ H
+ H
2
P
2
4
+ H
2
O  P
O
O  H
4
3
PO
4

Phosphoric acid is produced through the reaction between tetraphosphorus decoxide and water
A. H
3
PO
4
B. H
3
PO
4
C. P
4
O
10

+ H
+ H
2
P
2
4
+ H
2
O  P
O
O  H
4
3
PO
4

Iron(III)oxide reacts with carbon monoxide to produce iron and carbon dioxide
A.
B.
C.
FeO + CO  Fe + CO
2
Fe
2
O
3
+ CO  Fe + CO
2
Fe + CO  Fe
2
O
3
+ CO
2

Iron(III)oxide reacts with carbon monoxide to produce iron and carbon dioxide
A.
B.
C.
FeO + CO  Fe + CO
2
Fe
2
O
3
+ CO  Fe + CO
2
Fe + CO  Fe
2
O
3
+ CO
2

Coefficients
 whole numbers in front of formula
 distributes to numbers of atoms in formula
 specifies the relative number of moles and molecules involved in the reaction
 used to balance the equation


1.
2.
3.
4.
5.
6.
Balancing Equations
ONLY add/change coefficients-
NEVER subscripts!!!
balance one type of atom at a time balance polyatomic ions first balance atoms that appear only once second balance H and O last simplify if you can
Check at end!


Rules for writing and balancing equations – Pg.
327 in text.





Writing Equations
Write Word equations to help you organize reactants and products
Be sure to include symbols showing states of each reactant and product
Be sure to write the correct formula for each (crossing over for ionic compounds!)
Check your balancing of the equation when you are finished

Example 1

Description:

Aqueous iron III oxide reacts with hydrogen gas to produce iron metal and liquid water
Word Equation:
Iron III oxide + hydrogen gas  iron + water

Example 1

Formula Equation:
Fe
2
O
3 (aq)
+ H
2 (g)
 Fe
(s)
+ H
2
O
(l)

Balanced Formula Equation
Fe
2
O
3 (aq)
+ 3H
2 (g)
 2Fe
(s)
+ 3H
2
O
(l)

Example 2

Solid calcium metal reacts with water to form aqueous calcium hydroxide and hydrogen gas.
 calcium + water 


Ca
(s)
Ca
(s) calcium hydroxide + hydrogen
+ H
2
O
(l)
+ 2H
2
O
(l)
 Ca(OH)
2(aq)
 Ca(OH)
+ H
2(g)
2(aq)
+ H
2(g)

Example 3
 solid zinc metal reacts with aqueous copper (II) sulfate to produce solid copper metal and aqueous zinc sulfate
 zinc + copper (II) sulfate 


Zn
(s)
Zn
(s) copper + zinc sulfate
+ CuSO
4 (aq)
 Cu
(s)
+ CuSO
4 (aq)
 Cu
(s)
+ ZnSO
4 (aq)
+ ZnSO
4 (aq)

Example 4

Hydrogen peroxide in an aqueous solution decomposes to produce oxygen and water
 hydrogen peroxide  oxygen + water


H
2
O
2 (aq)
 O
2 (g)
2H
2
O
2 (aq)
 O
2 (g)
+ H
2
O
(l)
+ 2H
2
O
(l)

Example 5

Solid copper metal reacts with aqueous silver nitrate to produce solid silver metal and aqueous copper (II) nitrate
 copper + silver nitrate  silver + copper (II) nitrate


Cu
(s)
Cu
(s)
+ AgNO
3 (aq)
 Ag
(s)
+ 2AgNO
3 (aq)
 2Ag
(s)
+ Cu(NO
3
)
2 (aq)
+ Cu(NO
3
)
2 (aq)

Example 6

Carbon dioxide gas is bubbled through water containing solid barium carbonate, creating aqueous barium bicarbonate


 carbon dioxide + water + barium carbonate
 barium bicarbonate
CO
CO
2 (g)
2 (g)
+ H
+ H
2
2
O
O
(l)
(l)
+ BaCO
+ BaCO
3 (s)
3 (s)
 Ba(HCO
3
)
2 (aq)
 Ba(HCO
3
)
2 (aq)

Example 7


Acetic acid solution is added to a solution of magnesium bicarbonate to create water, carbon dioxide gas, and aqueous magnesium acetate.
acetic acid + magnesium bicarbonate  water + carbon dioxide + magnesium acetate


HCH
H
3
2
2HCH
2H
COO
O
(l)
(aq)
+ CO
3
COO
2
O
(l)
+ Mg(HCO
2 (g)
(aq)
+ Mg(HCO
+ 2CO
2 (g)
3
)
2 (aq)
+ Mg(CH
3
)

3
COO)
2 (aq)
+ Mg(CH
3

2 (aq)
COO)
2 (aq)


Write the balanced formula equation for:
Lithium metal is added to a solution of aluminum sulfate to make aqueous lithium sulfate and aluminum metal.

Types of Chemical Reactions

5 basic types discussed here
 not all reactions fall in these categories
 you should be able to:

 categorize a reaction predict the product(s)

1. Synthesis
 also called combination reaction
 reactants:

 more than one can be elements or compounds
 products: only one compound
A + X  AX where A is the cation and X is anion

1. Synthesis




Rubidium and sulfur
Rb
(s)
+ S
8 (s)
 Rb
Magnesium and oxygen
2
S
(s)
Mg
(s)
+ O
2 (g)
Sodium and chlorine
 MgO
Na
(s)
+ Cl
2 (g)
 NaCl
Magnesium and fluorine
Mg
(s)
+ F
2 (g)
(s)
(s)
 MgF
2 (s)

1. Synthesis
 calcium oxide and water
CaO
(s)
+ H
2
O
(l)
 Ca(OH)
2 (aq)
 sulfur dioxide and water
SO
2 (g)
+ H
2
O
(l)
 H
2
SO
3 (aq)
 calcium oxide and sulfur dioxide
CaO
(s)
+ SO
2 (g)
 CaSO
3 (s)

2. Decomposition
 opposite of synthesis
 usually require energy
 reactants: only one compound
 products: more than one
 usually elements but can be compounds
AX  A + X

2. Decomposition
 water

H
2
O
(l)
 H
2 (g) calcium carbonate

CaCO
3 (s)
 CaO calcium hydroxide
(s)
+ O
2 (g)
+ CO
2 (g)

Ca(OH)
2 (s) carbonic acid
 CaO
(s)
H
2
CO
3 (aq)
 CO
2 (g)
+ H
2
O
+ H
2
O
(l)
(l)

3. Single Replacement
 an element replaces a similar element in a compound
 reactants: 1 element & 1 compound
 products: 1 element & 1 compound
A + BX  B + AX
Y + AX  X + AY

3. Single Replacement
 zinc and hydrochloric acid
(aq)
 ZnCl
2 (aq)
+ H
2 (g)

Zn
(s)
+ HCl iron and water

Fe
(s)
+ H
2
O
(l)
 FeO
(aq)
+ H magnesium and lead (II) nitrate
2 (g)

Mg
(s)
+ Pb(NO
3
)
2 (aq)
 Mg(NO
3
)
3 (aq)
+ Pb chlorine and potassium bromide
Cl
2 (g)
+ KBr
(s)
 KCl
(s)
+ Br
2 (g)
(s)

4. Double Replacement
 two similar elements switch places
 reactants: 2 compounds
 products: 2 compounds
AX + BY  BX + AY

4. Double Replacement



 barium chloride and sodium sulfate
BaCl
2 (aq)
+ Na
2
SO
4 (aq)
 NaCl
(aq)
+ BaSO iron sulfide and hydrochloric acid
4 (s)
FeS
(aq)
+ HCl
(aq)
 FeCl
2 (aq)
+ H
2
S
(g) hydrochloric acid and sodium hydroxide
HCl
(aq)
+ NaOH
(aq)
 NaCl
(aq)
+ H
2
O potassium iodide and lead (II) nitrate
(l)
KI
(aq)
+ Pb(NO
3
)
2 (aq)
 KNO
3 (aq)
+ PbI
2 (s)

5. Combustion




Only responsible for one type releases energy in form of heat/light reactants: hydrocarbon + O
H
2
O and CO
2
2 as the only products
Ex: CH
4
+ O
2
 CO
2
+ H
2
O

Combustion
 propane and oxygen
C
3
H
8
(g) + O
2
(g)  CO
2
(g) + H
2
O(g)

Practice
Classify each of the following reactions one of the five basic types:

Na
2
O + H
2
O  NaOH
 synthesis

Zn

(s)
+ 2HCl
(aq)
 ZnCl
2 (aq)
+ H
2 (g) single replacement

Ca

(s)
+ 2H
2
O
(l)
 Ca(OH) single replacement
2 (aq)
+ H
2 (g)

Practice




2H

2
O
2 (aq)
 O
2 (g) decomposition
+ 2H
2
O
(l)
Cu

(s)
+ 2AgNO
3 (aq) single replacement
 2Ag
(s)
+Cu(NO
3
)
2 (aq)
C

2
H
4 (g)
+ O
2 (g) combustion
 CO
2 (g)
ZnO

(s)
+ C
(s)
 2Zn single replacement
(s)
+ H
2
O
+ CO
2 (g)
(g)

Practice

Na

2
O
(s)
+ 2CO
2 (g) synthesis
+ H
2
O
(l)
 NaHCO

Ca
(s)

+ H
2
O
(l)
 Ca(OH) single replacement
2 (aq)
+ H
2 (g)
3 (s)

KClO
3 (s)

 KCl decomposition
(s)
+ O
2 (g)

H
2
SO
4 (aq)
+ BaCl
2 (aq)
 HCl
 double replacement
(aq)
+ BaSO
4 (s)

Activity Series Pg. 333

Activity
 ability of an element to react
 easier it reacts, higher the activity
 activity series

 list of elements organized according to activities from highest to lowest

Activity Series
 metals

 greater activity, easier to lose electrons easier to become a cation
 nonmetals
 greater activity, easier to gain electrons
 easier to become an anion

Activity Series
 used to predict whether single replacement reactions will occur
 most active is on top
 an element can replace anything below it but not any above it

Practice
 zinc and hydrofluoric acid


Zn
(s)
+ HCl
(aq)
 ZnCl
2 (aq)
+ H
2 (g) calcium and lead (II) nitrate


Ca
(s)
+ Pb(NO
3
)
2 (aq)
 Ca(NO
3
)
2 (aq)
+ Pb copper and lithium sulfate
(s)


Cu
(s)
+ Li
2
SO
4 (aq)
 no reaction bromine and iron (II) chloride

Br
2 (l)
+ FeCl
2 (aq)
 no reaction