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The Role of Chemistry as a Prerequisite Course
Key facts and results:
Fact: The problem solving skills routinely utilized in the ‘high IQ’ and
related professions (such as nursing, business management, accounting,
etc.) are introduced, learnt and mastered during physical science
Result: Professional programs and subsequent employers insist that their
candidates have a background in one of the physical sciences – both for
specific (allied health, engineering) and general (your family room carpet)
Fact: Study within any of the ‘high IQ fields’ will increase cognitive
skills, but only the physical sciences do so via the study of fundamental,
everyday phenomena so are of broad relevance and interest (we all
interact with and benefit from the manipulation of matter on a daily basis
Result: Chemistry (and physics) may be considered to be the ‘gatekeepers’
of cognitive learning – chemistry in particular introduces, develops and
subsequently equips students with cognitive skills necessary to succeed
in their chosen careers
Take home message: While the direct relevance of chemistry to your chosen
course of study may at times seem tenuous, remember that the cognitive
skills developed during such programs of study are of significant importance
to your professional development and employability. In essence, this is why
you are here.
5 How Chemistry is Perceived & Skills Needed to Succeed in
How Chemistry is Perceived:
Discussion: How did your friends and family respond when you told them
you were taking a chemistry course this semester??
Study Skills Needed to Succeed in Chemistry:
Fact: As discussed above, chemistry is all about the student developing
and learning to apply problem solving skills - your study habits should
reflect this. Do NOT fall in to the trap of believing you can learn chemistry
simply by memorizing the information from your text – you must practice
applying this information, not just be familiar with it.
Result: Successful chemistry students typically spend most of their
independent study time working assigned problems, not just reading about
them. To learn chemistry you must do chemistry is a truism worth
remembering. An analogy would be this: you read all the books out there on
the subject of golf, but don’t get round to swinging a club – what do you
think happens when you tee off for the first time?
Fact: Chemistry relies on a cumulative method of learning, i.e. theories
learnt from week 1 onwards will be repeatedly applied all the way through
the course. Thus, it is important that the student does not let any ‘gaps’
in their knowledge develop. This fact exemplifies the differences in
philosophy between the sciences and arts, as art courses are often more
modular in nature. Example: I overhead a student tell another: “Yeah, I blew
off reading the first book in my English class, but read the second one and
got a ‘B’”. This method of study is not recommended in chemistry
6 Analogy: Building a tower
Result: Successful chemistry students typically have exemplary
attendance records. In some cases they may not be the ‘best’ students, but
guarantee themselves a better grade than more capable students, who in turn
typically may miss as few as one or two lecture sessions (this is especially
true with regard to 3 hr. class sessions).
Pictorial analogy of attendance vs cumulative knowledge
‘I missed a lab’ ‘I missed a lecture’ ‘I missed a couple of lectures’
Don’t ‘Swiss cheese’ or ‘torpedo’ your chances of passing the course
because of missed work
Take home message: Simply by attending class regularly and completing
the HWK assignments you essentially guarantee yourself a passing
grade for the course, while, due to the nature of the material, deviating
from this approach may ensure the opposite
Chemistry in action: Explaining what happens on your
The burning of a charcoal brick on your backyard grill
(MACRO) explained in terms of a balanced chemical
ANY large (MACRO) scale chemical process can be
described using a MICRO scale chemical equation
featuring individual atoms and/or molecules
A cartoon representation of the reaction of the pertinent atoms and
molecules; along with the Chemists’ description – a balanced chemical
equation illustrating a single microscopic event.
This process is repeated many billions of times (MICRO) for the burning of
a charcoal briquette (MACRO)
12 The Components of Matter
Reading: Ch 1 sections 1 - 5 Homework: Chapter 1: 37, 39, 41, 43, 45, 47, 49
= ‘important’ homework question
Review: What is matter?
Recall: “Chemistry is the study of matter and its properties, the changes
matter undergoes and the energy associated with those changes”
Recap: There are 3 stable states of matter – solid (s), liquid (l) and gas (g).
13 Specific macro- and microscopic physical properties define the three
states of matter
State of Matter Macroscopic Description Microscopic Description
(observation) (chemical model)
The state matter is in depends on the strength of the forces
(chemical bonds) between the individual microscopic
particles within the matter
Task: Rank the intermolecular forces present in steam, ice and water in
order of increasing strength. Use the included figures as a guide.
14 Changing between the 3 states of matter
Describe the relationship between the mpt. and bpt. of matter,
with regard to microscopic processes, occurring at these
Example: The boiling of water to make steam ( H O →( H O )
2 (l) 2 (g)
15 Physical and Chemical Properties – what’s the difference?
Analogy: We all posses ‘as is’ physical properties, or
characteristics, that define us. For example, Dr. Mills is 5’11”
and has green eyes.
As with people, each chemical also possesses a unique set of ‘as is’ physical
properties that define it. For example, water is a clear, colorless, tasteless
molecular material that has a fpt. of 0 C and a bpt. of 100 C.
Chemical Properties, in contrast, are a function of change (usually
associated with a chemical reaction). For example, Iron (Fe) reacts with
oxygen gas to form rust:
4 Fe (s) + 3 O (g) → 2 Fe O (s)
2 2 3
Task: Identify the flowing as either chemical or physical properties
Property Chemical or Physical
Diamond is the hardest known
Charcoal burns to make CO (g)
The statue of liberty turned ‘green’
Copper is a good conductor of
Sugar dissolves in water
Melting of ice
Think up two more chemical properties of your own
16 Elements and Compounds – the further classification of pure matter
Task: State which of the following are elements, and which are compounds.
When done, try to come up with a definition of what elements and
Material Chemical Formula Element or
Water H O (l)
Oxygen gas O (g)
Pure silver coin Ag (s)
Sugar crystals C H O (s)
6 12 6
Carbon dioxide gas CO (g)
17 Compounds and elements can have either ‘giant’ or molecular structures:
‘Giant’: Repeating lattice of particles – usually
strongly bound (high mpt.) solids.
Examples: sand (SiO ), diamond (C), table salt
Molecular: a collection of independent molecular
units (molecules will be discussed in more detail
later). Usually (low mpt) liquids or gasses at room
Definition: Molecule – a small, independent
particle of matter made up from 2 or more atoms
Examples: water (H O), carbon dioxide (CO ),
Nitrogen gas (N )
Think of molecules like cars on the expressway – each car
(molecule) is a separate, independent unit that contains a
number of passengers (atoms). The cars (molecules) are
free to move while the people (atoms) stay fixed inside.
‘Giant’ materials are like people (atoms) ‘locked’ in place
at a very crowded concert, the DMV waiting room
18 Review: A microscopic scale view of several materials is presented below.
Label each using elemental or compound and molecular or ‘giant’ tags
Water (H O (l)) Silicon (Si (s))
Steam (H O (g)) Sodium Chloride (NaCl)
Details: Ice is a solid (crystalline) form of water (a
molecular compound). How would you describe
the structure of ice? Can you think of other similar
More Details: Allotropes of an Element
C C C
(diamond) (graphite) 60
19 Pure Matter v Mixtures
Recap: Pure matter is classed as either an ELEMENT or a
Elements can have either Molecular or ‘giant’ structures.
Examples: N (g) (Nitrogen gas, molecular), Pb(s) (metallic
lead, a ‘giant’ structure)
Compounds can also have either Molecular or ‘giant’
structures. Examples: H O(l) (water, molecular), Fe O (s)
2 2 3
(‘rust’ (iron oxide), a ‘giant’ structure)
Recall: A molecule is an independent unit containing two or
more atoms. Remember the car / passenger analogy.
Molecules can exist as either elements or compounds
ANY combination of different types of pure matter ‘placed
together’ is defined as a mixture (eg. air, milk, pepsi).
Mixtures are NOT pure materials. eg. Pure gold (Au) vs
‘white’ gold (Au+ Ag), or water (H O) vs pepsi (H O +
Discussion: Air contains a number of different components – what are they?
How would you describe what air is made up from using words like element,
compound, gas, molecular etc.?
20 Task: Assign generic labels that describe to microscopic scale matter shown
on the slide (e.g. ‘gaseous atomic element’ etc.)
As viewed from a macroscopic perspective, mixtures are classified
as either HOMOGENEOUS or HETEROGENEOUS
Discussion: Can you think of something that is both a
homogeneous mixture and a solid?
A Bronze statue of
Examples of Alloys:
Classification of Matter Flowchart
(Dr. Mills really likes this slide – why? Hint: Recall the fundamental job
of a chemist)
22 Task: Use the ‘Classification of Matter’ flowchart (above) to classify the
1. The compressed gasses in a deep sea diver’s gas bottle (He(g) and
2. A ham and cheese omelet
3. An ice cube (made from pure water)
4. A ruby (Al O (s) with Cr impurities)
Extra Credit: Ask me about the separation of mixtures
assignment (based on background reading)
23 Units of Measurement
Reading: Ch. 1 sections 6 & 7 Homework: Chapter 1: 51, 55, 57, 59, 61, 65,
67, 69, 71, 75, 77, 81, 83
= ‘important’ homework question
Discussion: List some common units of measurement we use on a daily
basis. How did these units originate?
Quantity measured Familiar Unit
Question: What are the ‘metric’ (S.I.) versions of the everyday units listed
Quantity measured Fundamental S.I. Symbol
Unit (base unit)
Notes: SI base units are used to determine derived S.I. units, as discussed
below. Some S.I. base units feature a decimal prefix – which one(s)?
25 Derived S.I. Units
Insert appropriate S.I. base units into an equation that
defines the respective derived S.I. unit desired.
Area = length x length = m x m = m
the derived S.I. unit for area is m
Determine derived S.I. units for the following quantities
Quantity measured Math involving S.I. base units Derived S.I. unit
These are harder examples. To solve them start by inserting appropriate S.I. base units
into an equation that defines the quantity sought.
Discussion: Why do scientists prefer the S.I. system?
Is the S.I. unit of volume (m ) reasonable for everyday applications? Why?
What unit of volume do chemists prefer? Why?
1 dm = 1 L
More detail on the chemist’s volume unit
27 Scientific Notation and S.I. Prefixes
Fact: Chemical problem solving most often involves using either very large
or very small numbers (e.g. counting the number of molecules in a drop of
water, or quoting the mass of the water drop in kilograms)
Recall: How many individual H O (l) molecules are there
in a drop of water. Write this amount as a regular number:
Number H O (l) molecules in 1 drop water = _____________________________
Problem: How do we represent and manipulate such quantities in an ‘easier’
Overview Example: Consider the statement “eight million people live in
London”. How can this quantity be best expressed numerically?