Transcript Chapter 10 Chemical Quantities

Chapter 10
“Chemical Quantities”
You will need a
calculator for this
chapter!
Section 10.1 p. 287
The Mole: A Measurement of
Matter
How do we measure items?
 You can measure mass,
 volume,
 or count pieces
 We measure mass in grams
 We measure volume in liters
 We count pieces in MOLES
Other Ways to Measure
Amount




Pair: 1 pair of socks = 2 socks
Dozen: 1 dozen donuts = 12 donuts
Gross: 1 gross of pencils = 144 pencils
(12 dozen)
Ream: 1 ream of paper = 500 sheets
of paper
Guided Practice Problem p. 289
Practice Problem #2
pg. 289
• Assume 2.0 kg of apples is 1 dozen and
that each apple has 8 seeds. How
many apple seeds are in 14 kg of
apples? (work INDEPENDENTLY to solve)
What is the mole?
Not this kind of mole!
Moles (abbreviated mol)
 Derived from German word
molekül (molecule)
 SI measurement of an amount
 1 mole = 6.02 x 1023 of
representative particles, or…..
 # of carbon atoms in exactly 12
g of Carbon-12 isotope
 Called Avogadro’s number
What are Representative Particles?
(Table 10.1 p. 290)
 The smallest pieces of a substance:
1) molecular cmpd - molecule
2) ionic cmpd - formula unit (made
of ions)
3) element: is the atom
• Remember the 7 diatomic elements?
(made of molecules) BrINClHOF
Br2 I2 N2 Cl2 H2 O2 F2
Guided Practice Problem #3 p. 291
Mole Video 3:49
Quick Quiz
• How big is a mole?
6.02 x 1023
• If everyone in the world got a mole of
pennies, how much $ would every person
have?
1 trillion bucks
$1,000,000,000,000
• If you stacked a mole of paper how many
times would it go from the Earth to the
moon?
80 billion times
80,000,000,000
• How long would it take for every person in
the world to eat through a mole of
marshmellows?
40,000,000 years w/o a
bathroom break!
Consider these questions:
• How many oxygen atoms in
the following?
CaCO3 3 atoms of oxygen
Al2(SO4)3 12 (3 x 4) atoms of oxygen
• How many ions in the
following?
CaCl2 3 total ions (1 Ca ion and 2 Cl ions)
NaOH 2 total ions (1 Na ion and 1 OH ion)
Al2(SO4)3 5 total ions (2 Al + 3 SO ions)
2+
1-
1+
1-
3+
4
2-
Practice problems
The Mass of a Mole of an
Element
 Atomic mass of element (mass
of 1 atom) expressed in amu
- atomic masses - relative
masses based on mass of C-12
(12.0 amu)
- 1 amu is 1/12 mass of C-12
atom
Molar Mass….
 = mass of 1 mol of element in grams
(periodic table)
 12.01 grams C has same # particles
as 1.01 g H & 55.85 g Fe
 12.01 g C = 1 mol C
All contain
 1.01 g H = 1 mol H
6.02 x 1023
 55.85 g Fe = 1 mol Fe
atoms
Molar Mass Practice
Problems
What about compounds?
 1 mol of H2O molecules has 2 mol
of H atoms & 1 mol of O atoms (think
of a compound as a molar ratio)
 To find mass of 1 mol of a cmpd:
odetermine # moles of elements
present
oMultiply # times their mass (from
periodic table)
oadd up for total mass
Calculating Molar Mass
Calculate molar mass of
magnesium carbonate, MgCO3.
24.3 g
+
12.0 g
+ 3 x (16.00 g) =
84.3 g
So, 84.3 g = molar mass for MgCO3
Section 10.2
p. 297
Mole-Mass and MoleVolume Relationships
Molar Mass
 Molar mass - generic term for
mass of 1 mol of any substance
(expressed in grams/mol)
 Same as:
1) Gram Molecular Mass (for molecules)
2) Gram Formula Mass (ionic compounds)
3) Gram Atomic Mass (for elements)
o molar mass is more broad term than
these other specific masses
Examples
 Calculate the molar mass of:
= 78.05 g/mol
Na2S
N2O4
= 92.02 g/mol
C
= 12.01 g/mol
Ca(NO3)2 = 164.10 g/mol
C6H12O6 = 180.12 g/mol
(NH4)3PO4 = 149.12 g/mol
Molar Mass is…
 # of g in 1 mol of atoms,
formula units, or molecules
 Make conversion factors
from these
- To change btwn g of
cmpd and mol of cmpd
Using the Mole Roadmap
 How many moles is 5.69 g
of NaOH? 0.142 mol NaOH
The Mole-Volume Relationship
gases
- hard to determine mass
 how many moles of gas?
 2 things affect gas V:
 a) Temp & b) Pressure
 compare all gases at = temp &
pressure
Standard Temperature and Pressure
 0ºC & 1 atm pressure
- abbreviated “STP”
 At STP, 1 mol of any gas has
V of 22.4 L
- Called molar volume
 1 mol of any gas at STP = 22.4 L
Practice Examples
Mole Day
Celebrated on
October 23rd from 6:02
am until 6:02 pm
(6:02 on 10-23)
Density of a gas
 D = m / V (density = mass/volume)
- for gas units are: g / L
 find density of a gas at STP if
formula known
 You need: 1) mass and 2) volume
 Assume 1 mol, so mass is molar
mass (from periodic table)
 At STP, V = 22.4 L
Practice Examples
(D=m/V)
Another way:
 If given density, find molar mass of gas
 Assume 1 mol at STP, so V = 22.4 L
modify: D = m/V to show: m = D x V
 “m” will be mass of 1 mol, given 22.4 L
 What is molar mass of a gas with
density of 1.964 g/L? = 44.0 g/mol
 How about a density of 2.86 g/L?
64.0 g/mol
Summary
• all equal:
a) 1 mole
b) molar mass (in grams/mol)
c) 6.02 x 1023 representative
particles (atoms, molecules, or formula units)
d) 22.4 L of gas at STP
make conversion factors from
these 4 values (p.303)
Noticethis
all conversion
conversionsmap
mustinto
go
Copy
through
MOLE!
yourthe
notes!
Section 10.3
p. 305
Percent Composition
and Chemical Formulas
 All percent problems:
part
x 100 % = percent
whole
1) Find mass of each element,
2) Divide by total mass of cmpd; & x 100
%mass of element =
mass of element
mass of cmpd
x 100%
% composition from mass
 Calculate the percent composition
of a compound that is made of 29.0
grams of Ag with 4.30 grams of S.
29.0 g Ag
X 100 = 87.1 % Ag
33.3 g total
4.30 g S
X 100 = 12.9 % S
33.3 g total
Total = 100 %
% comp from the chemical formula
 If we know formula, assume
you have 1 mole,
 Subscripts used to calculate
mass of each element in 1
mole of cmpd
 sum of masses is molar mass
% Composition Examples
% composition as
conversion factor
 We can also use % as
conversion factor to calculate
# grams of element in cmpd
 Calculate % C in C3H8
 What is mass of C in 82.0 g sample
of propane (C3H8) 67.1 g C
% Composition
4:15
What is an Empirical Formula?
• Like ingredients for recipe –
double recipe, you double each
ingredient, but ratio of
ingredients stays same
• Empirical formula: lowest
whole number ratio of atoms
in cmpd
Calculating Empirical
 Find lowest whole number ratio
C6H12O6 = CH2O
CH4N = this is already the lowest ratio.
 A formula is not just ratio of atoms, it
is also ratio of moles
 1 molecule of CO2 = 1 atom of C
and 2 atoms of O
 1 mol of CO2 = 1 mol C and 2 mol O
Calculating Empirical
 get a ratio from % composition
1) Assume you have a 100 g sample
- the percentage become grams (75.1% = 75.1 grams)
2) Convert grams to moles.
3) Find lowest whole number ratio by
dividing each # of moles by
smallest value
Example
 Calculate empirical formula of
cmpd composed of 38.67 % C,
16.22 % H, and 45.11 %N.
CH5N
 Assume 100 g sample, so
 38.67 g C x
 16.22 g H x
 45.11 g N x
1mol C = 3.22 mole C
12.0 g C
1mol H
= 16.22 mole H
1.0 g H
1mol N = 3.22 mole N
14.0 g N
Now divide each value by the smallest value
Example
 The ratio is 3.22 mol C = 1 mol C
3.22 mol N
1 mol N
 The ratio is 16.22 mol H = 5 mol H
3.22 mol N
1 mol N
= C1H5N1
which is = CH5N
Practice Problem 36 p. 310
What is a Molecular Formula?
• Molecular formula: true # of atoms of
each element in formula of cmpd
• molecular cmpds only
• Example: molecular formula for
benzene is C6H6 (note that
everything is divisible by 6)
• Therefore, empirical formula =
lowest whole number ratio)
CH (the
Formulas (continued)
ionic compounds ALWAYS
empirical (cannot be reduced).
Examples:
NaCl
MgCl2
Al2(SO4)3
K2CO3
Formulas (continued)
Formulas for molecular compounds
MIGHT be empirical (lowest whole
number ratio).
Molecular:
H2O
C6H12O6
C12H22O11
H2O
CH2O
C12H22O11
(Correct formula)
Empirical:
(Lowest whole
number ratio)
Empirical to molecular
 Since empirical formula is lowest
ratio, the actual molecule weighs
more
Molar mass
=
Empirical formula mass
whole # to increase
each coefficient in
empirical formula
Empirical to molecular
practice problem
Empirical and Molecular
Formulas
3:29