Download 1 mole = 6.02 X 10 23 Particles

The Mole Concept:
“How Chemist’s Measure”
and
Stoichiometry: “Balancing of
Chemical Rxns”
Ch# 3, Ch#7, Ch#9
I. Chemistry = Quantitative Science
A. Quantitative: Analyze/describe via definite
numeric results. Ex:___________________
B. Qualitative: Analyze via nonnumeric or descriptive
results (words) Ex:____________________
C. We measure the amt. of matter to do calculations
relating quantities of reactants to products to
chemical equations.
Yields
REACTANTS
PRODUCTS
D. We often use very small and very large numbers
in chemistry.
E. Scientific notation is a method to express these
numbers in a manageable fashion.
-Thus 0.000 000 1 cm can be written 1 x 10-7 cm.
F. Scientific notation expresses a number as the
product of two factors, the first falling between 1
and 10 and the second being a power of 10.
Format is
Mantissa
x
Base Power
Decimal part of
original number
6.02 x
Decimal
you moved
23
10
We just move the decimal point around.
602000000000000000000000
G. Converting into scientific notation:
1. Move decimal until there’s 1 digit to its left.
Places moved = exponent.
2. If number converting is:
a) Large # (>1)  positive exponent
b) Small # (<1)  negative exponent
3. Only include sig. figs.
2,500,080,000
0.00002205
H. Powers of 10 to know: Metric
•
•
•
•
•
•
0.000 000 001 or
0.000 001 or
0.001 or
0.01 or
0.1
1000 or
Symbols Used in Chemical Equations
“Yields”; indicates result of reaction
Used to indicate a reversible reaction
(s)
A reactant or product in the solid state;
also used to indicate a precipitate
Alternative to (s), but used only to indicate a precipitate
(l)
A reactant or product in the liquid state
(aq)
A reactant or product in an aqueous solution
(dissolved in water)
(g)
A reactant or product in the gaseous state
Additional Symbols Used in Chemical Equations
Alternative to (g), but used only to indicate a gaseous product
D
Reactants are heated
2 atm
pressure
Pressure at which reaction is carried out, in this case 2 atm
Pressure at which reaction is carried out exceeds normal
atmospheric pressure
0 oC
Temperature at which reaction is carried out, in this case 0 oC
MnO2
Formula of catalyst, in this case manganese (IV) oxide,
used to alter the rate of the reaction , NOT CONSUMED
Example: H2O2(aq)
MnO2
H2O (l) + O2 (g)
II. Measuring Matter
A. Measure by:
1)
2)
3)
4)
Counting (How many Ss are in this class?)
Weighing (Au sells for $25/gram)
Volume (The National Ave. for gas is $3.05/gal)
Sometimes by each (bananas 10¢/ea , by lb., by quart)
B. Chemists needed a counting unit: mole = Amt.
of Substance
III. The Mole (Counting)
A. Matter is made of Particles: (Representative)
1)
2)
3)
4)
Elements = Atoms
Ionic Cmpds = Formula unit (3D/ Cube)
Molecular Cmpds = Molecules
1 mole is the amount of substance that contains as
many particles (atoms or molecules) as there are in 12.0
g of C-12.
B. Count them using Avagadro’s Number:
1) Dozen (doz) eggs = 12 eggs
2) 1 mole = 6.02 X 1023 Particles
Pics of Particles/Pieces
3. Amedeo Avagadro (1766-1856) – Italian
scientist & lawyer who first tried to
experimentally determine the # of particles in
a mole.
– never knew his own number;
– it was named in his honor by a French scientist in
1909.
– its value was first estimated
by Josef Loschmidt, an
Austrian chemistry teacher,
in 1895.
How Big is a Mole?
One mole of marbles would cover the entire Earth
(oceans included) for a depth of two miles.
One mole of $1 bills stacked
one on top of another would
reach from the Sun to Pluto
and back 7.5 million times.
It would take light 9500 years to travel from the bottom
to the top of a stack of 1 mole of $1 bills.
Welcome to Mole Island
1 mole = 22.4 L
@ STP
1 mol = molar mass
1 mol =
6.02 x 1023 particles
Mass, Volume, Mole Relationship
IV. Particle Conversions:
A. Particles  Moles & Moles Particles
1. How many moles of Mg is 3.01X1022 Atoms of Mg?
Unknown =
Known =
Conv:
Soln: 3.01 x 1022 Atoms Mg X
1 Mol Mg
6.02 x 1023Atoms Mg
= 5.00 X 10-2 Mol Mg
2. How many moles are 1.20 X 1025 atoms of
phosphorus?
3. How many atoms are in 0.750 mol of zinc?
4. How many molecules are there in 4.00 mol of
glucose, C6H12O6?
Atoms, F. U.’s, or Molecules
B. Analogy: Have a doz. packages of pencils.
How many pencils do you have?
8 in a package = 8 x 12 = 96
C. To find the # of atoms in a Mol of a cmpd, must
determine the # of atoms in a representative
formula of that cmpd.
1) How many Fluoride ions are in 1.46 mol of aluminum
fluoride?
2) How many ammonium ions are in 0.036 mol
ammonium phosphate?
3) How many C atoms are in a mixture of 3.00 mol
C2H2 and 0.700 mol CO?
V. Mole (Mass/Weight)
A. Learned - mass of a single atom is given in
amu (atomic mass units).
B. Gram Atomic Mass (gam) – # of grams of an
element numerically equal to amu (we will
use the entire #).
Ex. C (gam) = 12.01115 g/mol (grams/mole)
H (gam) = 1.0079 g/mol
C. The gam = 1 mole = 6.02 x 1023 particles
for a mono-atomic element!!!
1 Mole of Particles
Do They look alike ? Do you think they will have the same mass if 1 mole?
Ex. What is the mass of 1 mol of each of these
monatomic elements?
a. Sodium
b. Arsenic
c. Uranium
D. What is the mass of 1 mole of a cmpd?
1) Formula of a cmpd tells us the # of atoms of each
element in a rep. particle of that cmpd.
2) Calculate the gram molecular mass (gmm) or gram
formula mass (gfm) by adding together the atomic
masses of the atoms making up a molecule or F.U.
3) Examples:
Sulfur Trioxide:
Ammonium Phosphate:
Aluminum Sulfate:
E. gam, gfm, & gmm are commonly referred to as
molar mass in chemistry.
(1 mol= mass from pd table= 6.02x1023 parts)
F. Mole  Mass (g) Conversions use Molar Mass
1) Molar mass of an element or cmpd is used to convert
grams of a substance into moles and vice versa.
2) Examples:
a) How many grams are in 7.20 mol of dinitrogen trioxide?
b) Find the number of moles in 922g of iron (III)
oxide.
c) What is the mass of 2.65 x 1024 molecules of
Carbon Tetrachloride?
d) A sample of chlorous acid has a mass of 3500.1
grams, how many representative particles are
present in this sample?
Each substance is a mole, so why do they look so different?
One mole each of a solid, a liquid, and a gas:
.
One mole of NaCl(the solid)
mass = 58.45 g
One mole Water, mass= 18.0 g
occupies a volume of 18.0 mL.
One mole of the O2 gas
has a mass of 32.0 g
& 22.4L volume
VI. Mole (Volume of A Gas)
A. Volumes of moles of different substances can
be different.
B. Vol of a mole of gas is much more predictable.
C. Vol of a gas is measured at STP (Standard
Temperature & Pressure)
1. Standard Temp = 0oC
2. Standard Pressure = 1 atm (atmosphere)
How do we measure pressure?
D. At STP, 1 mol of any gas occupies a vol of
22.4 L aka molar volume of a gas.
1 mol = 6.02 x 1023 P = Molar Mass ? = 22.4L
E. Problems:
1. Determine the volume, in liters, of 0.600 mol of
SO2 gas at STP.
2. Determine the number of moles in 333.6 mL of
He gas.
3. What is the volume at STP for 3.25 x 1024
molecules of CO2 gas?
Homework Problems:
1. What is the vol at STP of 0.960 mol CH4?
2. Determine the vol (in L) of 127.80 g of
Titanium (II) Oxide gas.
VII. Calculating Percent Composition
A. When a new cmpd is made in a laboratory,
we need to determine its formula.
B. So we use percent composition which is
simply finding the percent (%) by mass of
each element in a compound.
C. The % comp includes as many %’s as there
are elements in the cmpd.
For Example: K2CrO4
% comp: 40.3% K, 26.8% Cr, & 32.9% O
(∑=100%)
D. Percent Composition Formula: is the # of grams of
the element divided by the grams of the cmpd,
multiplied by 100%
Or
% Comp = grams of element
grams of cmpd
x 100%
We Can Use This Formula 2 Ways!!!!
Most Common Use:
E. Example Problems:
(given g amts < or > 1 mol  use #’s directly)
1. An 8.20 g piece of Mg combines completely w/
5.40 g of O to form a cmpd. What is the % comp of
this cmpd?
2. Calculate the % comp of ethane, C2H6.
(Given no #’s so you assume 1 mol & use molar mass)
VIII. Determining Formulas from %
Comp:
A. Empirical Formula: formula showing the
smallest whole number ratio of elements
present.
C6H12O6 = Molecular Formula
CH2O = Empirical Formula
B. Finding Empirical Formula:
1. Find % Composition:
-Unknown has 63% Mn & 37% O
2. Convert % to moles:(Assume 100g sample)
Mn: 63 g Mn x 1 mole Mn = 1.1 mole Mn
54.9 g Mn
O:
3. Convert Moles to Whole Numbers:
Method: divide all molar amounts by
the smallest molar amount
Mn: 1.1 mol Mn = 1 mol Mn
1.1
MnO2
O: 2.3 mol O = 2.1 mol O
1.1
Rule: round #’s to nearest whole # unless….
If answer is 0.4, 0.5, 0.6 then double ALL!
Practice Problem 1
Given the following data, find the correct empirical formula:
49.0% C, 2.70% H, 48.2% Cl
Practice Problem 2
Given the following data find the empirical formula:
N = 26.2%, H = 7.50%, Cl = 66.4%
C. Molecular Formula:
1. First find the empirical formula.
2. Then compare the formula mass for your
empirical formula with the formula mass for
the molecular formula.
3. You may have to multiply the subscripts
in your empirical formula by some factor.
Example #1:
Given: 38.7% C, 9.70% H, 51.6% O and a Molecular mass
of (gmm=62.0g) find the true molecular formula:
1.Empirical Formula:
2. gmm of E.F.:
3. Divide Masses:
4. Multiply E.F. by Factor:
Practice Problem #2
Given the following data, find the correct molecular formula:
24.3% C, 4.1% H, 71.6% Cl: gmm = 99.0g
Practice #3
Given the following data, find the correct molecular formula:
54.6% C, 9.00% H, 36.4% O and gmm: 176g
IX. Balancing Chemical
Equations
A. All equations we have written have been
correct in Qualitative sense but not
Quantitative.
B. Equation must be balanced so that they are
quantitatively correct.
--balanced equation: each side of the equation has
the same number of atoms of each element.
Law of Conservation of Mass!!!
Atoms can’t be created or destroyed just
REARRANGED!
C. Sometimes we write equations and it will
already be balanced:
C (s) + O2 (g)  CO2 (g)
D. Other equations we must correct or
balance the quantities of reactant atoms
to product atoms:
C (s) + O2 (g)  CO (g)
H2(g) + O2(g) ----->
H2O (l)
Pt
E. The numbers in front are called
STOICHIOMETRIC COEFFICIENTS
– These tell us the # of moles/subscripts =________
F. Guidelines For Balancing: Mostly Trial and Error
Do Not Change Subscripts!!!!
1. Count the number of atoms on each side
2. Polyatomic ion unchanged (consider it one unit and
balance first)
3. Even / odd (make all even)
4. Single Elements last
5. Make sure all coefficients are in simplest (lowest) ratio
AgNO3(aq) + Cu(s)  Cu(NO3)2 (aq) + Ag (s)
More Examples: (notice no physical states)
(NH4)3PO4 +
AlCl3
+
P +
O2 
Mg(OH)2 
Li2CO3
P2O5

Mg3(PO4)2 + NH4OH
Al2(CO3)3 +
LiCl
X. Mole Ratio is The KEY
A. Balanced equations show proportions:
Synthesis of Water – 2H2 + O2  2H2O
For Every 2 H molecules we need 1 O molecule!
B. Calculations involved in chemical rxns use
the proportions to find the quantity of
reactant and product.
- Assume all Rxns go to completion
C. Since Avogadro’s # relates molecules to
moles the coefficients also tells us the # of
moles needed for a rxn to progress
successfully.
D. Stoichiometry-the proportional relationship
between 2 or more substances during a
chemical rxn.
E. Converting Between Amounts in Moles
Unknown
Known
1. Identify the amount in moles you know.
2. Using coefficient from eqn., set up mole ratio with Known on Bottom
3. Multiply the original amount by the mole ratio and finish problem.
Example:
Consider the rxn for the commercial prep of
Ammonia.
N2 + 3H2  2NH3
How many moles of Hydrogen are needed
to prepare 312 moles of ammonia?
Example:
What mass of NH3 can be made from 1221 g H2
and excess N2?
Plan:
Use:
Fe2O3 + 2Al  2Fe + Al2O3
1. How many g of Al are needed for 135g Fe2O3?
2. How many g of Al2O3 can form when 23.6g Al
with excess Fe2O3?
3. How many g of Fe2O3 react with excess Al to
make 475g Fe?
4. How many g of Fe will form when 97.6g Al2O3
form?