Download Chapter 10 - Chemical Reactions

Chapter 10 Chemical Reactions
Chemical changes and reactions between compounds
Chemical formula revisited:
types of elements and ratio making up compound
empirical formula : simplest whole number ratio of
elements in a compound
-general description of how to assemble
the compound
ionic-exists as separate ions
-empirical formula is the right description
-simple mixture of ions Na+ and Cl- ions
molecular (covalent)-exists as a particular
bonded entity
-empirical formula tells what it’s made of
-does not describe how it is bonded
molecular formula : identifies actual numbers of atoms
in a molecule-discrete units of atoms
-not appropriate for ionic – array of separate ions
-molecules described by molecular formula
-you know a formula is molecular if the numbers
are not the simplest whole number ratio
glucose: C6H12O6 molecular since divisible by 6
empirical – CH2O
Making compounds – cannot count atoms
need to know about masses of compounds
Formula weight (mass) –mass of one formula unit (amu)
sum of atomic weights for ALL atoms in chemical formula
general term for ionic and molecular
Molecular weight (mass) - mass of molecule
add up all atomic weights for each atom in molecule
specific to molecules only (covalent)
use atomic mass since it is an average
over all isotopes – naturally abundant species
Formula weight of NaCl:
Molecular weight of C12H22O12
Also want to be able to describe percent amount of
substances in compound
Identification of substances
percent composition: what percent by mass is
from each element
% composition = mass of the element in compound
X100%
total formula weight
CAN USE EMPIRICAL OR MOLECULAR
Find % comp of each element in NaCl:
Ex: find % comp of each element in C8H18
empirical easier C4H9
Try:
H3BO3
Ba(NO3)2
Chemical Equations
Describes what it takes for a chemical reaction
to occur (A recipe for producing compounds)
Which involves changes in compounds (one to
another), states of matter (liquid, solid, gas), and
exchanges in energy.
GENERAL CHEMICAL EQUATION
Reactants
Products
To quantitatively show how to make reaction succeed,
you have to describe the numbers of atoms involved
LAW OF CONSERVATION OF MASS
Matter is neither created nor destroyed
The products must balance the reactants
C + O2
CO2
C=1
O=2
C=1
O=2
NEW REACTION: synthesis of carbon monoxide
C + O2
CO
(unbalanced)
Unequal – mass of Reactants different than for products
2C + O2
2CO
Now mass of Reactants is equal with mass of Products
Balance equations by changing coefficients
NOTE: The formula CAN’T be changed
Be careful to take inventory of ATOMS in equation!
Ex: CH4 + 2O2
C=1
O=4
H=4
CO2 + 2H2O
C=1
O=4
H=4
Many times, Balancing equations is a trial & error process
Ex: Combustion of Gasoline (Octane)
2C8H18(g) + 25O2(g)
16CO2(g) + 18H2O(g)
However, you should be familiar with the rules which
describe balanced chemical reactions.
1. Number of Atoms of each element conserved in
reactants and products
2. Cannot change formula of reactants or products
3. Can only change coefficients to balance equation
Hints to help in balancing equations
1. Balance compounds with biggest numbers of atoms
first
2. Treat polyatomics as single units, especially when on
both sides of equation
3. Fractional coefficients are useful
Can also identify physical states of compounds in reaction
(g) – compound is gaseous
(l) – compound is liquid
(s) – compound is solid
(aq) – compound is an aqueous solution (in water)
() - temperature change (energy)
(or ) - solid precipitates out (or gas escapes)
Precipitation of calcium
bicarbonate from hard water
- changes in state
Ca(HCO3)2 (aq) + Na2CO3 (aq)
2NaHCO3 (aq)+ CaCO3
precipitate
HYDROCARBONS-combinations of hydrogen and carbon
-important energy source - combustion
gasoline, acetylene, propane
-react with O2 to form H2O and CO2 plus ENERGY
Example:
butane 2C4H10 + 13O2
8CO2 + 10H2O
(lighters)
+ energy
Carbohydrates: hydrocarbon with oxygen
plants store sun energy as carbs (photosythesis)
we eat and respire oxygen
energy + CO2
GLUCOSE
: C6H12O6(s) + 6O2(g)
6CO2(g) + 6H2O(g)
Enzymes - slowly release energy in plants and animals
free burning produces quick energy release and flame
Types of Chemical Reactions
Oxidation-Reduction (REDOX) Reaction - GENERAL
-electrons transferred from one atom to another
-one element oxidized (loses e-)
and the other is reduced (gains e-)
<balance of charge-losses cancel increases>
-oxygen often involved, but F, Cl & other
nonmetals do the same thing
OXIDIZING AGENT
- takes electrons from other substances
- takes e- away causing substances to be oxidized
- O2 in food and fuels gains e- to form octet
- Cl & F in bleach kill bacteria - pools
REDUCING AGENT
- provides e- to substance being reduced
- carbon: gives e- to form octet
- REDUCTION OF IRON ORE [iron(III) oxide]
2Fe O (s)+ 3C(s)
2 3crossover
Note reverse
To get charge for name
4Fe(s) + 3CO2
BUT reactions are normally defined in terms of the
effect on reactants and products: reaction type
3 types of redox eqs.: combine, decompose, replace
Ion exchange - not redox since no change in charge
ions just replaced
Reaction Types
Combination (Synthesis) Reaction
two or more substances combine to form one
X+Y
XY
lower energy
oxidation of metals- rust, burning
burning of Mg metal
2Mg(s) + O2(g)
2MgO(s)
rusting of iron
4Fe(s) + 3O2(g)
2Fe2O3(s)
burning non-metals
C(s) + O2(g)
CO2(g)
Decomposition Reaction
a compound is broken down into simpler
compounds or constituent elements
XY  X + Y requires energy
HYDROLYSIS - decompose water
with electricity - hydrogen
fuel cells
2H2O(l)
2H2(g) + O2(g)
electricity
decomposition of mercury(II) oxide
2HgO(s)
2Hg(s) + O2(g)
Replacement Reaction
One atom or polyatomic ion in a compound is replaced
XY + Z
XZ + Y or XY + A
AY + X
(-) part replaced
(+) part replaced
chemical activity - tendency of an element
to give up electrons
more chemically active if unable to keep electrons
MORE ACTIVE METALS GIVE UP ELECTRONS TO LESS ACTIVE
ACTIVITY SERIES - Fig. 10.12
shows most active elements on top
a metal will replace any metal it is above
has a larger chemical activity
2Al(s) +3CuCl2(aq)
2AlCl3(aq) + 3Cu(s)
Al more chem active - above Cu
- gives e- to Cu+2
- Cu precipitates, Al goes into solution
if less chem active, e- would
stay where they are (no reaction)
Alkali and Alkaline metals- very active
often replace other elements
2Na(s) + H2O(l)
2NaOH + H2
H takes e- from Na
makes Na+ and H2
Ion Exchange Reaction
Each ion of one compound is replaced by the ions of
a second compound
Ions mix together to form - a precipitate (insoluble)
- a gas
- or water
removes ions from the solution of ions
AX + BY
AY + BX
one of products must leave solution
(precipitate, gas or water)
Example:
3Ca(OH)2(aq) + Al2(SO4)3(aq)
3CaSO4(aq) +2Al(OH)3
Application:
Dissolved ions form insoluble aluminum hydroxide “net”
for water treatment- mixture traps suspended impurity
particles
Solubility Tables (Appendix B)
if the products are soluble --- no reaction
if one product is insoluble ---reaction occurs
precipitate or insoluble gas
Examples: identify the reaction type
2Al(s) + Fe2O3(s)
Al2O3(s) + 2Fe(s)
2Ag(s) + S(g)
Ag2S(s)
NaCl(aq) + AgNO3(aq)
NaNO3(aq) + AgCl
2KClO3(s)
2KCl(s) + 3O2
Summary of Chemical Equations
Tells how compounds combine together to form
new substances
what atoms are present before and after
(inventory)
2H2(g) + O2
2H2O
atomic description:
4 atoms H and 2 atoms O form 4 atoms H and 2 atoms O
MOLECULAR description:
2 molecules H2 and 1 molecule O2 forms 2 molecules of H2O
Also describes masses (formula weight):
reactant mass = 2x2.016amu + 32 amu = 36 amu
H2
O2
product mass = 2x18amu
=36 amu
H 2O
LAW OF CONSERVATION OF MASS
NOTE: 2g of H combines with 8 g of O
always the same mass ratio
remember amu based on weight of C-12
exactly 12 amu!
relationship between mass and number
historically misunderstood by Dalton
mis-measured mass of H and O to form H20
he thought water was HO
Led to failure of Daltons atomic theory
Gay-Lusaac and Avogadro
gases combine in whole number ratio
(at constant temperature and pressure-STP)
2H2(g) + O2
2H2O
2 volumes of H2 for each volume O2
gives one volume of water vapor
LAW OF COMBINING VOLUMES
equal volumes of gas at STP have equal
numbers of molecules
FIXED formula for water:
(twice as much H2) H2O
SHOWED H and O diatomic (H2&O2)
since two volumes of water produced
New interpretation of chemical equation:
coefficients in chemical equation give the volumes of gas
BUT we need to know how to get numbers of
molecules, and atoms by working with mass
mass is easier to determine (measure)
for high yield reactions (chemical co.
pharmacy, etc)
How do masses and numbers of atoms relate?
Each element has a averaged atomic weight based
on the mixtures of isotopes (natural abundance)
each weight a comparison to carbon-12:
Study weight mass relationship for C-12
the number of atoms in 12.0 g of
C-12 is 6.02x1023 experimentally
counted
Avogadro’s number - 6.02x1023
SI unit for counting numbers
Defines the mole - an amount of a substance that
contains Avogadro’s number
6.02x1023 H2O molecules
- 1 mole of water
- 1 mole of O atoms
- 2 moles of H atoms
- 3 moles total of atoms
All element weights based on a comparison to
carbon-12!
remember:
mass of one mole C-12=12 g
numerically equal to the atomic weight
Weight of one mole of any element is equivalent
Define weights based on large amounts of atoms
gram-atomic weight - the mass (in grams)
of one mole of an element
-numerically equal to atomic wt.
(in grams)
-represents 6.02x1023 atoms
- atomic weight since a mixture
of isotopes
- one gram-atomic weight of any
element contains the same number
of atoms
6.02x1023
gram-formula weight - the mass (in grams) of
one mole of the compound
-represents 6.02x1023 formula units
-numerically equal to formula wt.
(in grams)
- applies to ionic or covalent
gram-molecular weight - the mass (in grams) of
one mole of the molecules (covalent only)
-gram-formula weight of a
molecular compound
-represents 6.02x1023 molecules
-numerically equal to formula wt.
(in grams)
Can now go back and forth between number and mass
like conversion factors
Conversions between mass and number
Gram weights can be thought of as
conversion factors: grams per mole
Example: NaCl- formula weight = 58.44 amu
gram-formula wt.= 58.44 grams
58.44 g for every mole
conversion identity 58.44 g = 1 mole NaCl
24.7 g of NaCl is how many moles?
24.7 g 1 mole =0.423 moles
58.44 g
Example: 0.773 moles of C2H4 has what mass?
gram formula wt.=2x12.01g + 4x1.008g
= 28 g (per mole C2H4)
Convert:
28 g
0.773 moles 1 mole = 21.64 g of C2H4
How many molecules?
6.02x1023 molecules
0.773 moles
1 mole
= 4.65x1023 C2H4 molecules
How many C atoms: 2 C atoms per mole
2 C atoms
0.773 moles 1 mole C H =1.546 moles
2 4
Suggested problems for Chapter 10:
Parallel Exercises - # 1, 2, 3, 5
Additional problems:
1. Convert the following to numbers of moles
a) 6.7 g of iron
1 mole Fe
6.7g 55.8 g = 0.12 moles Fe
b) 45.3 g of Fe2(SO4)3
45.3 g 1 mole Fe2(SO4)3 =0.113 moles Fe2(SO4)3
c)
8.4x1024
399.9 g
molecules of C6H12O6
8.4x1024 molecules
d) 9.4x1023 Na+ ions
1 mole C6H12O6 =
6.02x1023 molecules
13.95 moles
of C6H12O6
2. Convert the following into masses
107.9 g
a) 0.55 moles of silver1 mole silver =59.34 g
16.16 g in a mole
c)
4.7x1022 molecules of CH4 6.02x1023 molecules =1.26 g
2 moles of AgCl - Answer: 286.8 g
d)
5.6x1024 Cl- ions - Answer: 330.23 g
b)
3. In two moles of BaF2:
a) how many moles of F are there? Answer: four
b) how many Ba+2 ions are there? Answer: two
c) what is the mass of F? Answer: 76 g
d) what is the mass percent of F? (Ch. 9) Answer: 21.7%