Transcript Chemistry Module 1 About Chemistry Chemistry is the scientific study of matter, including its properties, its composition and its reactions. There are many branches of.

Chemistry
Module 1
About Chemistry
Chemistry is the scientific study of matter,
including its properties, its composition and
its reactions.
There are many branches of chemistry:
– Organic chemistry: --substances with carbon
– Inorganic chemistry: --substances without carbon
– Analytical chemistry: --composition / identification
– Physical chemistry: --theoretical basis of chemistry
– Biochemistry: --Substances in living things
(Optional Enrichment)
Chemistry evolved from Alchemy, the medieval
study of “magical” properties of materials
– Best known alchemist: Nicholas Flamel
rumoured to have found the secrets of the philosopher’s stone
and the elixir of life
About 1600 alchemy began to disappear, and be
replaced by the more systematic approach of
chemistry.
Early chemists include:
–
–
–
–
Robert Boyle (who worked on the gas laws)
Antoine Lavoisier (who found laws of fixed proportions)
John Dalton (who first described atoms)
Joseph Priestly (who discovered oxygen)
Review of Important Points from
Previous Science Courses
• Properties of matter
– Physical Properties:
• Properties that can be examined without reacting a material.
Examination may cause physical changes, such as change
of state or form.
– Chemical Properties
• Properties that can only be determined by reacting a material
with another material (which usually changes or “destroys” it)
– Characteristic Properties
• Properties that apply to a single material or a small group of
similar materials. They help us identify a material.
– Non-characteristic Properties
• Properties that are less helpful in identifying a material
because the apply to many different substances.
Pure substances vs. mixtures
• Pure substances are substances that are the
same throughout.
– Theoretically, all the particles in a pure substance are
the same.
• There are two types of pure substance
– Elements: usually composed of atoms*
– Compounds: usually composed of molecules*
• Most materials are mixtures. They contain two
or more types of particle mixed together.
– solutions, suspensions, colloids, emulsions and most
composite solid materials are mixtures.
*As we shall see, this is a slight over-simplification that ignores ionic compounds.
Important Physical Properties
of pure substances

• Density: The ratio of the mass of a material
to its volume.
• Melting point: The temperature at which a
pure substance will melt (for pure substances,
this is the same as freezing point)
• Boiling point: The temperature at which a
pure substance will boil (for pure substances
this is the same as the condensation point).

m
V
Classification:
• Everything in the world that has a mass
and takes up space is called “matter”
• Matter can be classified as:
All Matter
(solids, liquids, gases, plasma)
Pure Substances
Mixtures
Examples
separation Types
include:
HomoHeteroElementsof physical
Compounds
of chemical separation
Filtration,
distillation,include:
magnetic
geneousseparation,
geneous
Ex. goldevaporation,
Ex. water
chromatography, settling, decantation,
flotation,
sorting,
Mixtures
Mixtures
Electrolysis,
decomposition
screening.
And precipitation.
Solutions
colloids emulsions suspensions
Changes
• Physical Changes DO NOT alter the nature of
the substance, for example:
– Change of form (tearing, breaking, crushing)
– Change of state (melting, freezing, boiling)*
– Change of mixture (blending, dissolving)*
• The molecules do not change during a physical change.
• Chemical changes DO alter the substance.
–
–
–
–
Decomposition
-Combustion
Synthesis
-precipitation
Oxidation
-electrolysis
Single or double replacement
• The molecules become different in a chemical change
* note: sometimes attempting to cause a physical change may trigger a chemical change.
Summary of Lesson 1
– Chemistry is the study of matter, its properties,
compostition and reactions. Chemistry includes:
• Organic chemistry
• Analytical chemistry
• Biochemistry
 Inorganic chemistry
 Physical chemistry
– Matter has properties
• Physical properties
• Chemical properties
Characteristic properties
Non-characteristic
– Elements and Compounds are pure substances
– All other substances are mixtures
– Physical changes do not alter the composition
• Change of form: tearing, crushing, breaking
• Change of state: melting, freezing, boiling
• Change of mixture: dissolving
– Chemical changes do alter the composition
• Combustion, precipitation, decompostition etc.
• Element Song, Version 1
• Element Song, Version 2
• Element Song, Version 3
Assignment
• Read chapter 1 of Addison-Wesley
Chemistry (pp. 1-11)
• Answer the following questions in your
assignments book:
– Addison-Wesley Chemistry pp.17-18
• Questions # 9-20
Sample Answers
• 9. Chemistry is the branch of science that
studies matter, as well as the composition
of substances and changes they undergo.
• 10. Five divisions of chemistry include:
– Organic chemistry
– Inorganic chemistry
– Analytical chemistry
– Physical chemistry
– Biochemistry
• 11. A hypothesis is a descriptive model or
trial explanation, formed after observation.
A theory is a hypothesis that has been
thoroughly tested. A law is a statement
that summarizes the results of
observations.
• 12. Experiments are used to test a
hypothesis, or to gather more data to
make a better hypothesis.
• 13. a, b and e, or more completely:
– Matter: concrete, propanone vapour, air
– Not matter: heat, sound
• 14. some physical properties of a nail
– Mass
- volume
-length
– Density
- colour
-magnetism
– Diameter
- conductivity
-hardess
– Melting point
(pick four)
• 15. in which state of matter do each of the
following occur at room temperature?
– Diamond (solid)
– Oxygen (gas)
– Cooking oil (liquid)
Mercury (liquid)
Clay (solid)
neon (gas)
• 16.
–
–
–
–
A) incompressible
B) indefinite shape
C) definite volume
D) flows
solid, liquid
liquid, gas
solid, liquid
gas, liquid
• 17. how to physically separate:
– A) iron filings and salt could be separated by using a
magnet, or by dissolving the salt in water and filtering
off the iron filings
– B) Salt and water could be separated by evaporation
• 18. Physical properties that distinguish:
– A) water and rubbing alcohol: density, odor, boiling
point (2 of these)
– B) Gold and aluminum: density, colour, conductivity
– C) Helium and oxygen: density, solubility, diffusion
rate
• 19. A homogeneous mixture is uniform in
composition (ie. it appears to be the same
throughout). A heterogeneous mixture is
not uniform.
• 20. Some methods of separating mixtures
include evaporation, distillation, dissolution
and filtration.
Module 1, Lesson 2
This is an outline of today’s lesson, not the notes
•
•
•
•
•
•
•
•
States of Matter
Phases (optional material)
Symbols
Energy
Conservation of Energy
Identifying Chemical Reactions
Chemical equations
Conservation of Mass
States of matter
• Solid
– definite shape
– Incompressible
-definite volume
-does not flow
• Liquid
– Variable shape
– Incompressible
-definite volume
-fluid (can flow)
• Gas
– Variable shape
– Compressible
-variable volume
-fluid (can flow)
Exotic states
of matter:
(optional
enrichment)
Extreme
pressures
(Optional
enrichment)
Although liquids and solids are said to be incompressible under ordinary conditions, at
Plasma: At very high temperatures electrons separate from gases and they glow.
extreme pressures (thousands of atmospheres) they may actually compress slightly.
Superfluid:
At very
cold temperatures
helium
will flow
in ways
normal liquids all
don’t.
Some
scientists
theorize
that at extreme
pressures
(billions
of atmospheres)
matter
might compress into an exotic state nicknamed “neutronium”.
Phases (Optional enrichment)
• The term “phase” is sometimes used as a
synonym for “state”, but phases are more
general than states. Phases are portions of any
chemical system that have uniform composition
and properties.
• The most common phases are:
– Solid
-liquid
-gas
(just like states)
• But phases can also include:
– Solute
-gel
-crystal
– Colloid
-vapour
-etc.
(which technically speaking are not states of matter)
• A mixture can have several phases but appear
Another
difference
“state” and “phase” is that the
to exhibit
only between
one state
term
onlytwo
to pure
substances
– state
Oil onapplies
water has
phases,
but both(ie
arepure
liquid.
elements
or pureincompounds)
while
term but
phase
– Diamonds
graphite have
twothe
phases
bothcan
are
apply solid.
to portions of a mixture.
Chemical Symbols
• Each element has a symbol
• By now, you should know the symbols of
common elements, including:
•H
•F
• Cl
•I
He
Ne
Ar
Ni
Li
Na
K
Co
Be
Mg
Ca
Ag
B
Al
Br
Au
C
Si
Fe
Hg
N
P
Cu
Pb
O
S
Zn
Energy
• Energy is the ability to do work
• There are many types of energy:
– Heat, light, sound, electricity, chemical,
nuclear, thermal,
• But to a chemist, the two main divisions of
energy are:
– Kinetic: Energy of motion (active energy)
– Potential: Energy of position or composition.
(passive or hidden energy)
Law of Conservation of Energy
• “In any physical or chemical
process, energy is neither created
nor destroyed.”
• Energy can, however, be changed from
one form to another
– For example, from potential energy to kinetic
energy or vice-versa.
Chemical Reactions
• In a chemical change or “reaction” one or
more substances are changed into new
substances. We say that the composition
has changed.
• The materials we started with were called
reactants
• The new materials produced are called the
products.
Reactants  Products
For example:
Hydrogen + Oxygen  Water
( 2H2 + O2  2 H2O )
Hydrogen and oxygen are reactants
Water is the product.
Identifying Chemical Changes
• How do you identify if a change has been
chemical instead of physical?
• These are some of the indications
– Combustion: sudden release of heat or flames
– Precipitation: a solid separates from the
mixture of two solutions
– Effervescence: bubbles of gas forming in a
solution
– Colour change: a significant change of colour.
Law of Conservation of Mass
• “In any physical or chemical
process, mass (matter) is neither
created nor destroyed.”
• The mass of all the products must equal
the mass of all the reactants.
– Sometimes it is hard to show this, because
some products may escape the container.
Summary of Lesson 2
– Three important states of matter are:
• Solid: definite shape, definite volume, incompressible
• Liquid: indefinite shape, definite volume, incompressible
• Gas: indefinite shape, indefinite volume, compressible
– You should know symbols of common elements
– Energy is the ability to do work. It includes
• Kinetic energy: the energy of motion
• Potential energy: energy of position or composition
– Law of conservation of energy
• In reactions, Energy is neither created nor destroyed .
– Chemical reactions change substances
• Know what reactants & products are.
• Know how to identify a chemical change.
– Law of conservation of mass
• In a reaction, mass is neither created nor destroyed.
Assignment #2
• Read the rest of chapter 1 (pp. 11-16)
• Answer questions #21-29 from page 17 &
18 in your assignments folder.
• If you haven’t done questions #9-20, do them too.
• 21. Identify the following as homogeneous or
heterogeneous:
– A) milk: (arguable) Homogeneous or heterogeneous*
– why? Real milk, straight from the cow, separates into cream, water, and milk solids.
Skim milk and homogenized milk do not. Technically, milk is an emulsion. A mixture
between homogeneous and heterogeneous, but closer to heterogeneous.
– B) glass: homogeneous mixture
– C) Table sugar: homogeneous compound
– D) river water: (arguable) heterogeneous* mixture
(*at microscopic level. At the visible level, filtered river water looks homogeneous)
– E) cough syrup: homogeneous mixture
– F) Nitrogen: homogeneous pure element
*do not mark these two answers wrong, just add the opposing argument.
• 22. Two ways to distinguish a compound
from an element are:
– A compound can be broken down into
elements by decomposition.
– Compounds contain two or more different
types of atom
• 23. Identify the following element, compound
or mixture.
– A) milk: mixture (water, milkfat, milk solids)
– B) glass: mixture* (72% SiO2, 13%Na2O, 15% other)
• This one is very technical. Most people mistakenly classify glass
as a compound. One type of expensive glass (fused silica) is a
pure compound: 100% SiO2).
– C) Table sugar: compound (C12H22O11)
– D) river water: mixture (H2O, minerals, impurities)
– E) cough syrup: mixture (alcohol, water, medicine*)
• The medicine could be dextromethorphan, codeine, or
antihistamine, depending on the brand. Some also contain
sugar, flavour and colour.
– F) Nitrogen: element (N2)
• 24. The chemical symbols are:
– Copper: Cu
– Oxygen: O
– Phosphorus: P
Silver: Ag
Sodium: Na
Helium: He
• 25. The elements found in each are:
– NH4Cl:
– KMnO3:
– C2H7OH:
– CaI2:
Nitrogen, Hydrogen, Chlorine
Potassium, Manganese, Oxygen
Carbon, Hydrogen, Oxygen
Calcium, Iodine.
• 26. Kinetic energy is the energy of motion
(active energy), potential energy is the
energy of position or composition (hidden
energy).
• 27. Examples of types of energy (choose 5)
– Nuclear
– Radiant
– Thermal
-hydro
-chemical
-electrical -mechanical
-solar
-etc.
• 28. The law of conservation of energy
says that energy cannot be created or
destroyed during a chemical reaction.
• 29. Classify as physical or chemical
change:
– Bending wire: physical
– Burning coal: chemical
– Cooking steak: chemical
– Cutting grass: physical
Module 1 Lesson #3
• Overview of SI Metric system
– Prefixes
– Length
– Volume
– Mass
– Temperature
The SI metric system
• Resulted from an attempt to make a
sensible measurement system based on
powers of ten
• The metre was originally defined as
1/10000000 of the distance from the
equator to the north pole.
• All the other units were then derived from
the metre.
Metric Prefixes
•
•
•
•
•
•
•
•
•
•
•
•
•
•
•
•
•
•
•
•
•
•
•
•
Yotta
(100 zetta)
(10 zetta)
Zetta
(100 exa)
(10 exa)
Exa
(100 peta)
(10 peta)
Peta
(100 tera)
(10 tera)
Tera
(100 giga)
(10 giga)
Giga
(100 mega)
(10 mega)
Mega
(100 kilo)
(10 kilo)
Kilo
Hecta
Deca
1024
1023
1022
1021
1020
1019
1018
1017
1016
1015
1014
1013
1012
1011
1010
109
108
107
106
105
104
1000
100
10
Superclusters
Galaxy
nearby stars
Solar system
Inner planets
Earth/moon
East coast
Town
football field
Elephant
deci
centi
milli
(100 micro)
(10 micro)
micro
(100 nano)
(10 nano)
nano
(100 pico)
(10 pico)
pico
(100 femto)
(10 femto)
femto
(100 atto)
(10 atto)
atto
(100 zepto)
(10 zepto)
zepto
(100 yocto)
(10 yocto)
yocto
1/10
1/100
1/1000
10-4
10-5
10-6
10-7
10-8
10-9
10-10
10-11
10-12
10-13
10-14
10-15
10-16
10-17
10-18
10-19
10-20
10-21
10-22
10-23
10-24
hand
fingernail
sand
bacteria
molecule
atom
proton
electron?
quark??
strings???
Common metric units & prefixes
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•
•
•
•
•
•
•
•
•
•
•
•
MegaKiloHectaDeca-----DeciCentiMilliMicro-
M
k
prefixes (large)
h
da
metre, litre, gram, etc.
units
d
c
m
prefixes (small)
μ (or u)*
*if your keyboard does not support Greek letters
Length
• Unit of length is the metre (also spelled
meter)
• It can be divided into
– Decimetres
– Centimetres
– Millimetres
millimetre metre
centimetre
decimetre
Volume
• A cube 0.1m per side (a cubic decimetre)
is defined to have a volume of one litre
– 1 cubic decimetre = 1 Litre
– 1 cubic centimetre = 1 mL
– 1 cubic metre
= 1000 litres
= 1 kilolitre
• The symbol for litre can be L, l or curly l,
but in Canada the “L” is preferred.
Mass
• The mass of one litre of pure water at
standard conditions (4°C) is defined to be
one kilogram = 1000 g
• 1 litre of water = 1 kg
• 1 mL of water = 1 g
• 1 cubic metre of water = 1000 kg = 1 Mg = 1 tonne
Since it awkward to haul around a litre of distilled
water, and since the purity of local water is
questionable, a prototype kilogram was made of
platinum (IPK) and stored in the archives of France. It
is still used to calibrate balances around the world.
Kelvin
Celsius
Fahrenheit
473
200
392
453
180
356
433
160
320
413
140
284
393
120
248
373
100
212
353
80
176
333
60
140
Body temp.
Room temp.
Degrees Celsius (°C) A.KA. Centigrade Water freezes
313
40
99
293
20
68
273
0
32
Often used in Chemistry
253
-20
-4
233
-40
-40
213
-60
-76
193
-80
-112
173
-100
-148
153
-120
-184
133
-140
-220
113
-160
-256
93
-180
-292
73
-200
-328
53
-220
-364
33
-240
-400
13
-260
-436
0K
-273 C
-460 F
Temperature
• Degrees Fahrenheit (°F)
NOT to be used in Chemistry!
•
•
•
•
•
•
•
•
•
Freezing point
Room temp
Body temperature
Boiling point
Freezing point
Room temperature
Body temperature
Boiling point
32°F
68 °F
99 °F
212 °F
0 °C
20 °C
37 °C
100 °C
Water boils
Mercury freezes
• Kelvins (K), formerly: °K or Absolute °A )
The Best for Chemistry, especially with gas laws.
•
•
•
•
Freezing point
Room temperature
Body temperature
Boiling point
273 K
293 K
310 K
373 K
Absolute zero
Conversions
To convert :
Use these steps:
Example
°C to K
Add 273.15 (for simplicity, we often leave out the .15)
20°C = 293 K
K to °C
Subtract 273.15 (again, we can drop the .15)
300 K = 27 °C
°C to °F
multiply by 9, then divide by 5, then add 32
20°C = 68 °F
°F to °C
Subtract 32, then multiply by 5, then divide by 9
212°F = 100°C
°F to K
Subtract 32, then multiply by 5, then divide by 9
Then add 273
Or go to Google and type one of the following:
20 C in K
300 K in C
20 C in F
212 F in C
Module 1: Lesson #4
• Measurement
• Accuracy vs. Precision
• Significant Figures (Significant Digits)
– In measurement
– In calculations
Acceptable error of
several instruments
Measurement.
Thermometer ± 0.2°C
Balance ± 0.05g
Graduated cylinders:
• Measuring quantities is an
10 mL
important aspect of
50 mL
experimentation.
100 mL
• Instruments used for measuring
are seldom perfect. Each
instrument has an amount of
uncertainty or “error”
• Knowing the acceptable error
helps set the reliability of a
result.
± 0.1 mL
± 0.5 mL
± 1.0 mL
Accuracy vs. Precision
• Accuracy is how close an instrument’s
reading is to the actual correct value
• Precision is how well an instrument
reproduces a result
– An instrument that is inaccurate but precise
can often be adjusted to give better results.
– An instrument that is imprecise will have a
higher uncertainty or “error”.
– An instrument that is imprecise and
inaccurate should be discarded and replaced.
Significance
• It is misleading to write a result that
implies more precision than was
measured. To avoid excessive precision,
the concept of significance was
developed.
– Results should never be written with more
precision than the measurements that were
used to calculate them.
Example of Excess Precision
(discussion point)
• John wants to calculate the circumference
of a cylindrical water tank. He measures
the diameter as 2.55 m and then multiplies
the measurement by pi (3.1415926535)
• 2.55 x 3.1415926535 = 8.011061266425
m
8.01 m
• This is an extremely misleading number. His
measurement was nowhere near precise
enough to support this result. He must round
this off to a more reasonable result.
Your own measurements
• Make a judgement call of how accurate
your results are, based on your
instruments.
– For example, if your instrument allows you to
measure a value to the nearest tenth millilitre
(ie. Its acceptable error is  0.1mL) then you
can record values like:
• 3.9 mL or 4.0 mL
or 4.1 mL
– You measured to the nearest 0.1 mL
• Don’t write 4 mL
– it suggests that you were not precise enough
• Don’t write 4.00 mL
– it implies more precision than you actually measured
Interpreting Measurements
• If you see a measurement you may not
know how precise it is.
• We use the concept of significant digits to
determine its precision.
– The number of significant digits determines
the precision of the measurement and tells
how much you can safely round the results.
– Remember that a number with too many digits
is just as misleading as one with too few!
Rules of Significant Figures
(for measurements made by someone else)
• The digits 1 through 9 are ALWAYS treated as
significant in a written measurement.
• Zeros between significant digits are ALWAYS
significant.
• Leading zeros (in front of a number) are NEVER
significant.
• Trailing zeros ARE significant unless the person
who recorded them was estimating.
– If there is a decimal point they ARE significant
– If there is no decimal then you must use your
judgement. Was the person estimating or not? (see next slide)
Ambiguous Digits
If a “whole” number ends in zero, the situation is
ambiguous.
5280 ft
3 SD or 4 SD?
20 000 m 1 SD or 5 SD?
Unless we know who took the measurement, we
can’t tell if the trailing zeros are significant or not.
How do we handle this?
The textbook tells us to call these zeros significant.
Most other chemistry books say to treat them as
insignificant.
What’s a chemist to do? Use your judgement.
Using Your Judgement
• First look for clues.
• The word “about” or “approximately” in the
description of the measurement tells you to use the
lower number of significant digits.
• The word “exactly” or “precisely” tells you to count
all the digits.
• The context or type of source may help. Popular
scientific articles and newspapers usually round off
the number, so use the lower number of digits
• Textbooks and professional journals are usually
more accurate, so use all digits.
• If all the other measurements are very precise,
then assume the ambiguous measurement is too.
• If still in doubt, use all the digits.
Example
003.50270
6 significant digits
75000
5 significant digits (if accurate)
2 significant digits (if estimated)
0.001010
4 significant digits
All digits 1 through 9
are significant
All zeros between
significant digits are
significant
Leading zeros are never
significant
Trailing zeros with a
decimal point are
significant
Trailing zeros with no
decimal are sometimes
significant (use judgement)
Calculations with significant digits
You are the weakest link!
Good bye!
The result of a calculation
can have no more
significant digits than the
WORST measurement!
Multiplying and Dividing:
• Do the calculation, then round the answer so it
has the same number of significant digits as the
worst measurement.
2.514 cm x 3.1 cm = 7.7934  7.8 cm2
528 g ÷ 25 mL =21.12  21 g/mL
• However: if you multiply or divide by a number
that has no units (ie. An integer used to double
or triple a result) it does not reduce the number
of significant digits
23.4 g x 2 = 46.8 g  doesn’t change.
Adding and subtracting
• Make sure the units are the same before adding
or subtracting (convert metric units if necessary)
• Do the addition or subtraction, lining up decimals
• Round the answer to match the number with the
fewest decimals (or fewest significant digits if
there is no decimal marker)
23.45 cm
+
4.5123 cm
27.9623 cm  27.96 cm
Examples
• 5.3 cm x 4.33 cm = 22.949 cm2  23 cm2
2 significant digits
• 5.8798 mL ÷ 3 g = 1.95993 g/mL  2 g/mL
One significant digit
• 4.3576 m x 2 = 8.7152 m
5 significant digits
Not a
measurment
What about scientific notation?
• The digits are significant (following the
normal rules of significant figures)
• The 10 and exponents are not significant.
x Not Significant
Significant
• Example
–
6.02 x 1023
–
1.3200 x 10-7
Has 3 significant digits
Has 5 significant digits
What about exact numbers
• In the unlikely event that we have a
measurement of exactly 230 000 objects,
how should we represent it?
• Remember, someone might interpret it as having only
2 significant digits if they thought we were estimating.
– one way:
“230 000 exactly” (verbal description)
– A better way:
Write “exactly” if measured or “about” if not
“2.30000 x 105” (scientific notation)
Convert measurments to scientific notation
When in doubt, convert your
answers to scientific notation!
• On tests and assignments, assume that all
trailing zeros are significant, unless you
see the word “about” or “approximately” in
the question.
• It is my intention to never give you a
problem on a test or examination that has
an answer with less than 3 significant
figures.
Summary of Lesson #3
• Precision and accuracy are important
when reading instruments
– Knowing the acceptable error of an instrument
helps you know how much precision to record.
• A result of a calculation must never be
more precise than the worst measurement
used in the calculation.
• Rules of significant figures can help us
decide how to correctly round off the
results of a calculation.
(see Rules of Significant Figures earlier in this lesson)
• When in doubt, convert your answers to
scientific notation.
Assignment #3
• Do the sheet “Significant Figures”
Answers
• a) 165 283.78
• c) 165 280
• e) 165 000
b)
d)
f)
165 283.8
165 300
200 000
• A) 5280 feet
• C) 22.40 m
• E) 4000 kg
b)
d)
f)
007 A
23001 mm
0.000745 L
• a) 789.30 m (5)
• c) 0.04 V
(1)
• e) 0.4320 g (4)
b) 7400 mL (2)
d) 73.2469 cm (6)
f) 503 mm(3)
• a) 5
• c) 5
b) 3
d) 3
• A) 23 m2
• C) 31 cm
• E) 15 m2
•
•
•
•
Always
Always
Never
Sometimes
b) 200 V
d) 2.00 g/mL
f) 91.4 m
Module 1: Lesson #4
Problem Solving in Chemistry
• Problem solving techniques (self-review)
• Conversion factors (self-review)
• Dimensional analysis (lesson)
Read pages 49-52
• In your “notes” book, list the steps that are
suggested for solving chemistry problems
• Try the five problems on p.51-52. You
may check your own answers by looking
at the solutions (see p.719). This is for
your own practice.
Dimensional Analysis
AKA: unit analysis
• Dimensional analysis uses the units that
are part of the measurements to help
analyze and solve a problem.
– Adding and subtracting units
– Multiplying and dividing units
– Cancelling units
– Simplifying units
– Comparing the units of the answer to the
expected units can determine if the problem is
done correctly.
Rule 1: Adding and subtracting units
• You can only add or subtract
measurements that have the same unit
– Make sure that you have converted quantities to the
same units. If one measurement is in litres, and the
other in millilitres, you must change one of them.
2.5 L + 250 mL
or
 2500mL + 250 mL = 2750mL
 2.5 L + 0.25 L
= 2.75 L
Rule 2: Multiplying and Dividing
• Whenever you multiply units, place a dot
(•) between them:
• 10 N x 5 m = 50 N•m
(newton-metres)
• Whenever you divide units, place a slash
between them
• 50 g ÷ 10 mL = 5 g/mL
or
• 8 mol = 4 mol/L
2L
(grams per millilitre)
(moles per litre)
Rule 3: Cancelling Units
• You may cancel units if the same unit occurs in
a numerator as in a denominator.
• A unit that is immediately after a slash is in a
denominator: m/sec = m
sec
500 m
sec
_

60 sec
min

60 min
h

1 km
1000 m
= 1800 km/h
Rule 4
• Look for units that can be simplified:
• 5 A x 20 Ω = 100 A·Ω = 100 V
A few you may remember:
Amperes x Ohms =Volts
meters x metres = square metres
Ohms x Volts = Watts
A• Ω = V
m•m = m2
Ω•V = W
• Compare the units to what is expected.
Problem: Chili
•
•
•
•
You need 600 servings of chili
10 servings of chili needs 2 tsp. of chili powder
How much chili powder will you need?
This problem is easy to solve, but let’s show the
problem with dimensional analysis
600 servings
1

2 tsp
10 servings

1200
10
Tsp · servings
servings
= 120 tsp
Don’t Copy… but be aware.
• Dimensional analysis is not an end in
itself. You will never be tested on it directly
• It is a tool to help you solve problems and
give the correct units of your answer
• In tests and exams you must give units for
every answer. Not giving the correct unit
for an answer may cost up to 25% of the
value of the correct answer.
Module 1, Lesson #3
Models of the Atom
I am the very model of a modern
major element
(the Elements Song by Tom Leher)
Models of the Atom
• This section is mostly review from PSC416
with a few new bits added at the end.
• Since the time of Aristotle and Democritus
philosophers and scientists have tried to
determine what matter is made of.
• Since the particles of matter are too small
to see, we use models (pictures and other
representations) to try to understand them.
Dalton
• John Dalton was the first modern scientist
to propose that matter was made of tiny
particles.
– The philosopher Democritus had suggested
this two thousand years earlier, but had never
produced a model to explain.
• Dalton called these tiny particles “atoms”
Highlights of Dalton’s Model
• Remember– Dalton said:
– All matter is made of atoms
– Each different element has a different type of
atom.
• There are many different elements (we have now
named 109, have proven the existence of 112, and
have evidence of up to 118)
– Atoms of elements can combine to form
molecules of compounds.
Thompson & Rutherford
• J.J. Thompson discovered the existence of
electrons, particles smaller than, and
found inside atoms. He was the first to
suggest that the atom contained other
particles.
• Rutherford discovered that most of the
mass of an atom is concentrated in the
center. He stated that the atom has a
dense nucleus in the center.
Rutherford’s model
Bohr’s Discovery
• Niels Bohr studied the wavelength of light
given off by excited atoms, and
determined that electrons “orbit” the atom
in different energy levels or “shells”.
• By combining the idea of a central nucleus
(Rutherford) with the idea of orbiting
electrons (Bohr) we developed the BohrRutherford model of the atom
Bohr-Rutherford Model
Energy Level
or “Shell”
electron
Nucleus
(protons )
(in later versions
also neutrons)
Simplified Bohr-Rutherford
Heavy circle represents
The NUCLEUS
F
Semi-circle represents
First electron shell
9p+
10n0
• Shell maximums:
Semi-circle represents
2nd electron shell
2e-
7e-
Etc.
(2n2)
– 2e-, 8e-, 18e-, 32e-, 50e- …
• But… A shell does not have to be
completely filled! For example:
– 2e-, 8e-, 8e-, 2e- is the usual arrangement for
calcium, NOT: 2e-,8e-,10e-
Why?
(simplified answer)
• Atoms usually arrange themselves so that
most shells can have one of “magic
numbers”… 2,8,18, 32 etc,
• Ca: 2e, 8e, 10e
Y
Y
N
This is still an oversimplification,
but it is the best we can do unless
You learn the Aufbau diagram!
vs. Ca: 2e, 8e, 8e, 2e
Y
Y
Y
Y
This side wins, because it has more
Shells with a “magic number”
• Draw simplified Bohr-Rutherford diagrams
of:
–N
– Mg
– Cl
– Ca
Atom Overview
• Atoms consist of:
– Protons: positively charged particles located
in the nucleus with a mass ≈ 1 amu (1859/1860)
– Neutrons: neutral particles located in the
nucleus with a mass = 1 amu
– Electrons: negative particles, orbiting the
atom with a mass ≈ 0 amu ( 1/1860 amu)
Module 1, Lesson #7
• Advanced models of the atom
• Optional enrichment material.
(advanced classes)
Modern Model of the Atom
The Modern or
“Electron Cloud”
Model
Subshells (Orbitals)
(optional enrichment)
• Each shell has one or more subshells or
orbitals that look like clouds.
• Labelled s, p, d, or f based on their shape
– s orbitals are spherical
– p orbitals are “peanut” shaped
– d orbitals may* be doughnut shaped
– f orbitals are flower shaped.
• Each orbital can hold up to two rapidly
moving electrons.
*actually, some d orbitals look a lot like f orbitals.
Shell #
s
p
d
f
Shell 1 1 2eShell 2 1 2e- 3
6e-
Shell 3 1 2e- 3
6e-
510e-
h,i,j
Total
Orbitals
These orbitals are not
actually used.
Number of orbitals/electrons per shell
1 orbital 2e-
Shell 4 1
2e-
3
6e-
510e- 714e-
Shell 5 1
2e-
3
6e-
510e- 714e- 9
Shell 6 1
2e-
3
6e-
5
Shell 7 1 2e- 3
6e-
510e- 714e-
10e-
714e-
Total
Electrons
4 orbitals 8e9 orbitals 18e16 orbs
32e-
25 theory
16 typical
50 theory
32 typical
9,11
36 theory
16 typical
72 theory
32 typical
9,11,13
49 theory
16 typical
98 theory
32 typical
Aufbau Diagram
6d
5f
7s
6p
5d
4f
6s
5p
4d
5s
5th shell starts to fill
4th Shell continues filling
4p
(Ga to Kr)
3d
3rd shell finishes:10 more e(Sc to Zn)
4s
4th Shell starts to fill
3p
3rd shell starts filling: 8e-
3s
2p
2s
1s
(K and Ca)
(elements from Na to Ar)
2nd shell fills: 8e-
Shell one
(elements from Li to Ne)
First shell fills: 2e(H and He)
Modern Model
(optional enrichment)
nucleus
Electron Cloud
1st Shell
Contains one spherical s orbital with two e-
=2e-
One s orbital (2e-)+3 p orbitals (6e-)
=8e-
2nd Shell
3rd Shell
1 s orbital,
3 p orbital
and 5 d orbitals with 10 e-
Answers to Sheet
Element
Configuration
Diagram
N
(nitrogen)
1s2,
Mg
(magnesium)
1s2,
Sc
(scandium)
1s2,
Fe
(iron)
1s2,
2s2 2p6, 3s2 3p6 3d6, 4s2,
Mn
(manganese)
1s2,
2s2 2p6, 3s2 3p6 3d5,
Zr
(zirconium)
1s2, 2s2 2p6, 3s2 3p6 3d10,
4s2 4p6 4d2, 5s2
W
(tungsten)
1s2, 2s2 2p6, 3s2 3p6 3d10,
4s2, 4p6 4d10 4f14, 5s2 5p6 5d4, 6s2,
2s2 2p3
2electrons, 5 electrons
2s2 2p6,
2 electrons, 8 electrons ,
2s2 2p6,
1 electrons 8 electrons
3s2
2 electrons
3s2 3p6 3d1, 4s2,
9 electrons
2 electron
4s2