Chem. 31 – 1/28 Lecture
Announcements
• Lab Adding Situation
– Sect. 2: 2 no shows (read names), so may be able to
add 3 students
– Sect. 4 appears to have zero wait list (may be able to
add if someone drops)
– However, if you are high on the list to be added,
check the labs again as space may open
• Homework and Quiz
– quiz next Wednesday
– corrected diagnostic quiz also due Wednesday
– quiz in lab next week - Wednesday and Thursday
Traditional vs. Modern Methods
Characteristic
Traditional
Modern
Equipment
Glassware and balances
(low cost)
Instruments
(high cost)
Precision
High
Moderate
Speed
slow
fast
Sensitivity
low
high
Selectivity
minimal
Good to great
Chapter 1 –
Measurements and Titrations
No measurement is valuable unless it is
given with units and some measure of
uncertainty
Units – Chapter 1
Uncertainty – Chapters 3 and 4
Units of Measure
• Most Basic –
Measure
Length
Mass
Time
Temperature
Amount
SI base units (important ones)
Unit_____
meter (m)
kilogram (kg) (only one with multiplier)
second (s)
Kelvin (K)
Mole (mol)
Units of Measure
• Directly Derived from Base Units
– Volume: cube volume = l3 so units = m3
l
– Density = m/V so kg/m3
– Pressure = force(kg·m/s2)/area(m2) =
kg/(s2·m)
Units of Measure
• Other metric units (not directly in SI units)
– Density (g/cm3)
– Pressure (Pascals or Pa = kg/(s2·m))
• Non-metric units (used commonly)
– For pressure 1 atmosphere (atm) = 101300
Pa
– English/Other system (not emphasized here)
Units of Measure
• Metric
Name
Kilo
Centi
Milli
Micro
Nano
Multipliers (ones you should know)
Abbreviation Multiplier
k
x103
c
x10-2
m
x10-3
m
x10-6
n
x10-9
Analytical chemists like small quantities. An instrument that can detect 1 fg
(1 x 10-15 g) is better than an instrument that can detect 1 pg (1 x 10-12 g)
Unit Conversion – Example Problem
• Convert the density of lead from g/cm3 to
kg m-3 if density = 11.7 g cm-3.
Concentration Units
• General form
mass or moles solute
mass or moles or volume of solution (or solvent)
• Note: sometimes, volume is required in the
denominator to be strictly considered
“concentration”, but for this class mass ratios
or mole ratios will be considered to be an
expression of concentration
Most Commonly Used
Concentration Units
• Molarity (M)
1 mol solute
1M 
L solution
• Mass Fraction (also valid for solids as mass analyte/mass sample)
g solute
Mass % 
 100
g solution
g solute
part per thousand by mass 
 1000
g solution
Most Commonly Used
Concentration Units
• Mass Fraction – continued
g solute
part per million (ppm) by mass 
 106
g solution
• Other Units:
Mass/volume units (e.g. mg/mL)
Conversion between Concentration
Units – Example Problem
What is the molarity of a H2O2 solution that
is listed as 25% H2O2 by mass and has a
density of 1.07 g mL-1? MW (H2O2) = 34.0
g mol-1
Preparing Solutions
• From Scratch (direct from solids)
Let’s say we want 100 mL of 0.100 M FeCl3.
What equipment do we need?
How do we make the solution?
Preparing Solutions
• From scratch
1. Solid to flask
2. Add liquid, dissolve
3. Fill to line
Preparing Solutions
• By Dilution
What if we need 20.0 mL of 0.00200 M FeCl3?
How do we make it?
If direct by weighing, mass FeCl3 = 0.0065 g (not
very accurate)
By Dilution from 0.100 M FeCl3
Preparing Solutions
• By Dilution
1. Pipet liquid
2. Add liquid, fill to mark
Stoichiometry
• Stoichiometry refers to ratios between moles of
reactants and products in chemical reactions
• The ratio of moles of reactants and products is
equal to the ratio of their stoichiometric
coefficients
Example:
aA + bB ↔ cC + dD
Moles A/moles B = a/b
Stoichiometry
•Example problem: How many moles of
H2O2 are needed to completely react with
25 mL of 0.80 M MnO4-?
Reaction:
5H2O2(aq) + 2MnO4- +6H+ ↔ 2Mn2+ + 5O2(g) + 8H2O(l)
Stoichiometry
• Remember: there are two (common) ways
to deliver a known amount (moles) of a
reagent:
– Mass (using formula weight)
– Volume (if molarity is known)
• Titrations = A practical way of using
stoichiometry with precise measure of
added volume