Transcript Honors Chapter 11 Reactions

 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?
Ch. 11 Chemical
Equations Reactions
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

can also happen with
physical changes
 gas

or light
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
H2, O2, N2, F2, Cl2, Br2, I2
other exceptions
• sulfur: S8
• phosphorus: P4
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:
Na2O + H2O  NaOH
Symbols Used in Equations
yields
reversible
above arrow:
or heat
MnO2 or Pt
25°C
2 atm
heated
catalyst
specific T
requirement
specific P
requirement
after a formula:
(s) solid
(l)
liquid
(aq) aqueous:
dissolved in
water
(g)
gas
Text
Pg. 323 Chart of
symbols used in chemical
equations
List three observations that suggest that
a chemical reaction has taken place.
Acids you have to know!
HCl
hydrochloric acid
H2SO4 sulfuric acid
HNO3 nitric acid
H3PO4 phosphoric acid
HC2H3O2 acetic acid
 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  Al2O3
Al + O2  Al2O3
Al3 + O  Al2O3
Aluminum reacts with
oxygen to produce aluminum
oxide
A.
B.
C.
Al + O  Al2O3
Al + O2  Al2O3
Al3 + O  Al2O3
Phosphoric acid is produced
through the reaction between
tetraphosphorus decoxide
and water
A. H3PO4  P4 + H2O
B. H3PO4 + H2O  P4
C. P4O10 + H2O  H3PO4
Phosphoric acid is produced
through the reaction between
tetraphosphorus decoxide
and water
A. H3PO4  P4 + H2O
B. H3PO4 + H2O  P4
C. P4O10 + H2O  H3PO4
Iron(III)oxide reacts with
carbon monoxide to produce
iron and carbon dioxide
A.
B.
C.
FeO + CO  Fe + CO2
Fe2O3 + CO  Fe + CO2
Fe + CO  Fe2O3 + CO2
Iron(III)oxide reacts with
carbon monoxide to produce
iron and carbon dioxide
A.
B.
C.
FeO + CO  Fe + CO2
Fe2O3 + CO  Fe + CO2
Fe + CO  Fe2O3 + CO2
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
Balancing Equations

1.
2.
3.
4.
5.
6.
ONLY add/change coefficientsNEVER 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:
Fe2O3 (aq) + H2 (g)  Fe
 Balanced
Fe2O3
(aq)
(s)
+ H2O
(l)
Formula Equation
+ 3H2
(g)
 2Fe
(s)
+ 3H2O
(l)
Example 2
 Solid
calcium metal reacts with water
to form aqueous calcium hydroxide
and hydrogen gas.
 calcium + water 
calcium hydroxide + hydrogen
 Ca(s) + H2O(l)  Ca(OH)2(aq) + H2(g)
 Ca(s) + 2H2O(l)  Ca(OH)2(aq) + H2(g)
Example 3
 solid
zinc metal reacts with aqueous
copper (II) sulfate to produce solid
copper metal and aqueous zinc sulfate
 zinc
 Zn(s)
 Zn(s)
+ copper (II) sulfate 
copper + zinc sulfate
+ CuSO4 (aq)  Cu (s) + ZnSO4 (aq)
+ CuSO4 (aq)  Cu(s) + ZnSO4 (aq)
Example 4
 Hydrogen
peroxide in an
aqueous solution
decomposes to produce
oxygen and water
 hydrogen peroxide 
oxygen + water
 H2O2 (aq)  O2 (g) + H2O (l)
 2H2O2 (aq)  O2 (g) + 2H2O (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)
+ AgNO3 (aq)  Ag

Cu
(s)
+ 2AgNO3 (aq)  2Ag
(s)
(s)
+ Cu(NO3)2 (aq)
+ Cu(NO3)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
 CO2 (g) + H2O (l) + BaCO3 (s)  Ba(HCO3)2 (aq)
 CO2 (g) + H2O (l) + BaCO3 (s)  Ba(HCO3)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
 HCH3COO (aq) + Mg(HCO3)2 (aq) 
H2O(l) + CO2 (g) + Mg(CH3COO)2 (aq)
 2HCH3COO (aq) + Mg(HCO3)2 (aq) 
2H2O(l) + 2CO2 (g) + Mg(CH3COO)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
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) + S8 (s)  Rb2S (s)
 Magnesium and oxygen
Mg (s) + O2 (g)  MgO (s)
 Sodium and chlorine
Na (s) + Cl2 (g)  NaCl (s)
 Magnesium and fluorine
Mg (s) + F2 (g)  MgF2 (s)
1. Synthesis
 calcium
oxide and water
CaO(s) + H2O(l)  Ca(OH)2 (aq)
 sulfur
dioxide and water
SO2 (g) + H2O (l)  H2SO3 (aq)
 calcium
oxide and sulfur dioxide
CaO (s) + SO2 (g)  CaSO3 (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
H2O (l)  H2 (g) + O2 (g)
 calcium carbonate
CaCO3 (s)  CaO (s) + CO2 (g)
 calcium hydroxide
Ca(OH)2 (s)  CaO (s) + H2O (l)
 carbonic acid
H2CO3 (aq)  CO2 (g) + H2O (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
Zn (s) + HCl (aq)  ZnCl2 (aq) + H2 (g)
 iron and water
Fe (s) + H2O (l)  FeO (aq) + H2 (g)
 magnesium and lead (II) nitrate
Mg (s) + Pb(NO3)2 (aq)  Mg(NO3)3 (aq) + Pb
 chlorine and potassium bromide
Cl2 (g) + KBr (s)  KCl (s) + Br2 (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
BaCl2 (aq) + Na2SO4 (aq)  NaCl (aq) + BaSO4 (s)
 iron sulfide and hydrochloric acid
FeS (aq) + HCl (aq)  FeCl2 (aq) + H2S (g)
 hydrochloric acid and sodium hydroxide
HCl (aq) + NaOH (aq)  NaCl (aq) + H2O (l)
 potassium iodide and lead (II) nitrate
KI (aq) + Pb(NO3)2 (aq)  KNO3 (aq) + PbI2 (s)
5. Combustion
 Only
responsible for one type
 releases energy in form of heat/light
 reactants: hydrocarbon + O2
 H2O and CO2 as the only products
Ex: CH4 + O2  CO2 + H2O
Combustion
 propane
and oxygen
C3H8(g) + O2(g)  CO2(g) + H2O(g)
Practice
Classify each of the following reactions
one of the five basic types:
 Na2O + H2O  NaOH

synthesis
 Zn (s)

(aq) 
ZnCl2 (aq) + H2 (g)
single replacement
 Ca (s)

+ 2HCl
+ 2H2O
(l)
 Ca(OH)2 (aq) + H2 (g)
single replacement
Practice
 2H2O2 (aq) 

decomposition
 Cu (s)

single replacement
(s)
+Cu(NO3)2 (aq)
+ O2 (g)  CO2 (g) + H2O
combustion
 ZnO (s)

(l)
+ 2AgNO3 (aq)  2Ag
 C2H4 (g)

O2 (g) + 2H2O
+C
(s)
 2Zn
single replacement
(s)
(g)
+ CO2 (g)
Practice
 Na2O (s)

+ H2O(l)  Ca(OH)2
 KCl
(aq) +
H2 (g)
(s)
+ O2 (g)
decomposition
 H2SO4 (aq) +

 NaHCO3 (s)
single replacement
 KClO3 (s)

(l)
synthesis
 Ca(s)

+ 2CO2 (g) + H2O
BaCl2 (aq)  HCl
double replacement
(aq) +
BaSO4 (s)
Activity Series
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

Zn
and hydrofluoric acid
(s)
+ HCl
 calcium

and lead (II) nitrate
and lithium sulfate
Cu (s) + Li2SO4 (aq)  no reaction
 bromine

ZnCl2 (aq) + H2 (g)
Ca (s) + Pb(NO3)2 (aq)  Ca(NO3)2 (aq) + Pb (s)
 copper

(aq) 
and iron (II) chloride
Br2 (l) + FeCl2 (aq)  no reaction