Introduction to Chemistry
“The world is full of obvious things. Which
nobody by any chance ever observe.”
-- Sherlock Holmes
• Careful observation is the foundation of
chemistry as an experimental science, leading
us to question what we have observed – how,
what why?
• The answers to these questions are sought in
experiments, which may be described as
observations made under controlled conditions
• Observation and experimentation – the twin
pillars of the scientific method
Scientific Method…
• The scientific way of knowing – often called the
scientific method – is sometimes presented as a
rigid sequence of events
• It is not however a rigid path – it is a process of
discovery!
• Discovery begins when we make observations
and then try to understand what we have
observed by asking key questions and
proposing possible answers
• This process of discovery begins as we design
and conduct experiments to test whether our
answers to these questions are valid!!
What are the steps of the Scientific
Method?
1. Make an observation(s)
2. Propose a hypothesis
3. Design and conduct a controlled
experiment
4. Analyze the results
5. Form conclusions
Controlled Experiment…
• Experiments should be designed so that the
effects of different variables on the behavior of a
substance can be studied independently
• A controlled experiment is when only one
variable at a time is changed
• There are two types of variables in an
experiment:
– Independent variable – the one that is deliberately
changed
– Dependent variable – the thing that changes because
of the independent variable
An Example…
• Students were trying to determine if the
amount of a Sodium chloride added to
calcium carbonate effected the amount of
heat given off
• Give 3 possible variables for this
experiment
• Give the independent and dependent
variables
What is chemistry?
• Chemistry is the study and investigation of
the structure, composition and property of
matter and the changes it undergoes
The properties of materials are always related
to their structure
Hence, structure determines properties
Measuring and Calculating in
Science…
•
Chemistry is a quantitative science
because it involves measuring and
calculating
 A measurement must have a number and a scale
(called a unit) to be meaningful
•
It can also be a qualitative science
because it can involve describing what is
happening in a reaction
What makes a measurement?
• In order to make measurements, we must
meet three requirements…
1.Know what we are trying to measure
2.Have a standard with which to compare
whatever we are measuring
3.Have a method for making comparisons
Exact Numbers…
• A number that is the result from a
definition or an exact count
• For example – there are 12 apples or π =
3.14
• All are significant
• Do not limit the number of sig figs
Uncertainty in Measurement…
• A measurement always has some degree
of uncertainty
• The amount of uncertainty depends on the
precision of the measuring device
• In science it is customary to report a
measurement by recording all certain
digits plus the first uncertain (estimated)
digit – these numbers are called the
significant figures of a measurement
Estimating Uncertainty in a
measurement…
• Remember all measurements are a result
of known values and one estimated
number
• When finding the uncertainty in a
measurement, we look at the estimated
number
• For example: 0.023
The 3 is the estimated number
It is in the 1000th place
The estimated uncertainty is written +0.001
This measurement has a very small uncertainty
Rules for Counting Significant
Figures…
1. Nonzero integers are always significant
figures
2. Zeros – there are three classes of zeros
a) Leading zeros precede all the nonzero digits
and are not significant figures
b) Captive zeros are between nonzero digits and
always count as significant figures
c) Trailing zeros are at the right end of the number
and are only significant if the number contains a
decimal point
Try a few…
1. Tell how many sig figs are in each
measurement and tell the uncertainty in
each measurement
a. 1508 cm
b. 300.0 ft
c. 20.003 g
d. 0.00705 L
Bell Ringer…
1. Tell the number of sig fig
– 94300
– 0.000400670
– 100000.
– 56.00
2. Tell the uncertainty of each measurement
– 3.45
– 6.0
– 12
– 4.725
Bell Ringer…
1. Each of the following are statements from
different labs – tell if they are quantitative
or qualitative
–
–
–
–
–
–
Bubbling
Heat given off
23.6 cm wide
A strong odor
pH of 5.4
273 K
Sig figs in Mathematical
Operations…
• To this point we have learned to count the
significant figures in a given number, but
we must also consider how uncertainty
accumulates as calculations are carried
out
Rules for Sig figs in Mathematical
Operations…
1. For multiplication or division the number of sig
figs in the answer is the same as the number
in the least precise measurement used in the
calculation
 4.56 x 1.4 = 6.38 → 6.4 (correct answer)
2. For addition or subtraction the number of sig
figs in the answer has the same number of
decimal places as the least precise
measurement used in the calculation
 12.11 + 18.0 + 1.013 = 31.123 → 31.1 (correct
answer)
Bell Ringer…
•
a.
b.
c.
d.
Carry out the following mathematical
operations and give each result with the
correct number of significant figures
1.05 x 10-3 / 6.135
21 – 13.8
20 X 23.00
14.75 + 34.25
Bell Ringer…
• The actual length of a certain plank is
26.782 cm. Which of the following
measurements is the most accurate? Are
the measurements precise?




26.5 cm
26.8 cm
26.202 cm
26.98 cm
Rules for Rounding…
•
•
1.
2.
In most calculations you will need to round numbers to
obtain the correct number of sig figs
When rounding, use only the first number to the right
of the last significant figure
In a series of calculations, carry the extra digits
through to the final result, then round
If the digit to be removed
a.
Is less than 5, then the preceding digit stays the same
1.33 → 1.3
b.
Is equal to or greater than 5, the preceding digit is increased by 1
1.36 → 1.4
Precision and Accuracy…
• Two terms often used to describe the
reliability of measurements are precision
and accuracy
Precision – the degree of agreement among
several measurements of the same quantity.
It also is known as the degree of
reproducibility of the measurement
Accuracy – the agreement of a particular
value with the true value
Bell Ringer…
• Decide if the following lab data is accurate
or precise or both
– 13.2 mL, 13.3 mL, 13.1 mL, 13.2 mL
– The actual value is 13.0 mL
• Make each of the following have 3 sig figs
– 34098
– 0.0003219
– 7154
– 76.78
Types of errors…
•
There are two types of errors in measurements
1. Random error (indeterminate error) – a
measurement that has an equal probability of being
high or low. This type of error occurs in estimating
the value of the last digit of a measurement
2. Systemic error (determinate error) – occurs in the
same directions each time. The measurement is
either always too high or too low
In groups…
a. There are 365 days/year, 24 hours/day, 12
months/year and 60 minutes/hr. Use this data
to determine how many minutes are in a
month.
b. Now use the following data to calculate the
number of minutes in a month: 24 hours/day,
60 minutes/hour, 7 days/week, and 4
weeks/month.
c. Why are these answers different? Which, if
any, is more correct and why?
Dimensional Analysis…
• It is often necessary to convert a given
result from one system of units to another
• The best way to do this is by a method
called unit factor method OR dimensional
analysis
Converting from One Unit to
Another…
• To convert from one unit to another, use
the equivalence statement that relates the
two units
• Derive the appropriate unit factor by
looking at the direction of the required
change (to cancel unwanted units)
• Multiply the quantity to be converted by
the unit factor to give the quantity with
desired units
Bell Ringer…
• A marathon race is 26 miles and 385
yards.
A. What is the distance in rods
B. What is the distance in meters
C. What is the distance in furlongs?
5.5 yards = 1 rod
40 rods = 1 furlong
8 furlongs = 1 mile
1 meter = 39.37 inches
1 yard = 36 inches
What if there is more than one unit
present?
• When more than one unit is present,
decide which unit you want to convert first
– Convert it first
– Then convert the second unit
***do not get confused!!!
EX:
How fast is a car going 35 miles/hour going in
yards/second?
1 mile = 1760 yards;1 hour = 60 minutes; 1 minute = 60 seconds
A few problems…
1. How many doughnuts can one purchase
for $123 if doughnuts cost $3.25/doz?
2. Convert 9.85 L to gal. 1.06 qt = 1.00 L
and 4 qt = 1 gal
3. A certain size of nail cost $1.25/lb. What
is the cost of 3.25 kg of these nails? 1kg
= 2.2 lb
Metric System Review…
• Scientists recognized that long ago a standard
system of units had to be adopted if
measurements were to be useful
• The system agreed upon in 1960 was the
International System or le Systeme International
(SI system)
• The SI system is based on the metric system
and units derived from the metric system
• Because fundamental units are not always
convenient, the SI system employs prefixes to
change the size of the unit
The Fundamental SI Units
Physical Quantity Name of Unit Abbreviation
Mass
gram
g
Length
Meter
m
Second
s
Temperature
Kelvin
K
Amount of Substance
Mole
mol
Electric current
Ampere
A
Luminous intensity
Candela
cd
Time
Derived units…
• Many SI units are combinations of
quantities
• These units are produced by multiplying or
dividing standard units
Derived SI units…
Quantity
Name of Unit
Abbreviation
Area (A)
square meter
m2
Length x width
Volume (V)
Cubic meter
m3
Length x width x
height
Density (D)
Kilograms per
cubic meter
kg/m3
Mass / volume
Molar mass (M)
Kilograms per
mole
kg/mol
Mass / amt of
substance
Concentration (c)
Moles per liter
M
Amt of
substance/
volume
Molar volume (Vm)
Cubic meters
per mole
m3/mol
Volume / amt of
substance
joule
J
Force x length
Energy (E)
Derivation
Dimensional Analysis with
metric units…
When converting with metric, always use
that value of the unit as compared to the
base unit
1.Convert 35.4 mm to m
2.Convert 2327.9 cg to kg
3.How many grams are in 53.24 dg?
Bell Ringer…
1. Why do we use the metric system?
2.Convert 35.4 mm to m
3.Convert 2327.9 cg to kg
4.How many grams are in 53.24 dg?
5. Convert the following:
a) How many inches are in 3.0 meters?
b) A baby weighs 8.5 lbs. How many grams is
that?
c) How many gallons of Coke would you drink if
you drank entire 2 liter?
• Science fiction often uses nautical
analogies to describe space travel. If the
starship U.S.S. Enterprise is traveling at
warp factor 1.71, what is its speed in
knots?
Warp 1.71 = 5.00 times the speed of light
The speed of light = 3.00 x 108 m/s
1 knot = 2000 yd/hr
Mass
• The measure of the resistance of an object
to a change in its state of motion OR the
amount of “stuff” in an object
• A scale is used to mass an object
Mass vs. weight…
• An important point concerning measurements is
the relationship between mass and weight
• Weight is the force gravity exerts on mass,
therefore weight varies with the strength of the
gravitational field
• Therefore if you went to the moon your weight
would change but not your mass
• Many times the terms mass and weight are
sometimes used interchangeably, although this
is incorrect!
Volume…
• The derived SI unit of volume is cubic meters (m3)
• Many times this unit is way too large to be a practical
way of expressing volume in a chemistry lab
• Instead, a smaller unit cubic centimeters (cm3) is used
• When dealing with the volumes of liquids and gases, the
non-SI unit liter (L) is often used
• Again the liter is often too large so the unit milliliter (mL)
is used
• This means 1 cm3 = 1mL
Review…
1. Round the following to 3 sig figs
– 96747210
– 91
– 0.0006589
2. How many sig figs are in each in #1?
3. What is the uncertainty of each measurement?
– 34.09
– 6.0222
– 12
4. What is the difference between precision and
accuracy?
Review…
1. Convert the following:
– How many grams are in 548.9 mg?
– How many feet are in 34.2 m?
– How many liters are in 2 gallon and 3.4
quarts?
It can be tricky with volume
conversions…
1. How many mL are in 14.65 kL?
2. How many L are in 48.6 cm3?
3. How many dm3 are in 29100 mL?
Bell Ringer…
1. A piece of metal has the mass of 3.45
kg. What is its mass in g?
2. A container has 2.3 L of gas in it? What
is its volume in mL?
3. A container has 750.00 mL of liquid in it.
What is its volume in m3?
What is Temperature?
• A measure of the AVERAGE kinetic
energy
• When looking at the different temperature
scales, all are talking about the same
height of mercury
Temperature Conversions…
• There are three systems used to measure
temperature
Degrees Fahrenheit (°F)
Degrees Celsius (°C)
Kelvin (K)
• Each has a different way of converting
between the values
How the equation for °F to °C was
derived…
• Notice -- 0°C = 32°F
and 100°C = 212°F
• If we subtract these
values then…
100°C = 180°F
* Find the value of 1°C
1°C = (180/100) °F
1°C = 9/5 °F
Converting…
1. Converting from °C to Kelvin
TC = TK – 273.15
TK = TC + 273.15
2. Converting from °C to °F
TF = TC x 9°F + 32°F
5°C
More Converting…
3. Converting °F to °C
TC = (TF - 32°F)5°C
9°F
Try these…
1. Normal body temperature is 98.6°F.
Convert this to the Celsius and Kelvin
scales.
2. Liquid nitrogen, which is often used as a
coolant for low-temperature experiments
has a boiling point of 77 K. What is this
temperature on the Fahrenheit scale?
Density…
•
•
•
•
•
Ratio of mass to volume
D = m/V
Unit usually: g/mL or g/cm3
Useful for predicting mass
Does not depend on the amount of
material in a compound
Density…
• A characteristic physical property of a
substance
• It does not depend on the size of the
sample because as the sample’s mass
increases, its volume increases
proportionally
• Density varies with temperature – V ↑ T↑
D↓
Bell Ringer…
• What is the density of an object that has a
mass of 14 g and a volume of 2 ml?
• What is the mass of an object with a
volume of 2 cm3 and a density of 1.5
g/mL?
• What is the volume of an object with the
mass of 20 kg and a density of 2.5 g/mL?
Bell Ringer…
• An empty container weighs 121.3 g. Filled with
carbon tetrachloride (density 1.53 g/cm3) the
container weighs 283.2 g. What is the volume of
the container?
• A 55.0 gal drum weighs 75 lbs when empty.
What will the total mass be when filled with
ethanol?
density of ethanol = 0.789 g/cm3
1 gal = 3.78 L
1 lb = 454 g
• In the opening scenes of the movie Raiders of the Lost
Ark, Indiana Jones tries to remove a gold idol from a
booby-trapped pedestal. He replaces the idol with a bag
of sand of approximately the same volume. (density of
gold = 19.32 g/mL; density or sand = 2 g/mL)
– Did he have a reasonable chance of not activating the
mass sensitive bobby-trap?
– In a later scene he and an unscrupulous guide play
catch with the idol. Assume the volume of the idol is
1.0 L. If it were solid gold, what mass would the idol
have?
• Is playing catch plausible?
It’s all that is Matter…
• Matter is anything that has mass and
takes up space
• All matter, regardless of form, has some
properties in common
Volume – the amount of 3-D space an object
occupies
Mass – a measure of the amount of matter in an
object
Basic Building Blocks of Matter…
• The most fundamental parts of matter are atoms
and molecules, which make up elements and
compounds
Atom – the smallest unit of an element that has all
the properties of that element
Element – a pure substance made of only one kind
of atom
Compound – a substance that is made from the
atoms of two or more elements that are
chemically bonded
Molecule – the smallest unit of an element or
compound that retains all the properties of that
element or compound
Classification of Matter…
• Matter is classified according to how it is
organized
• Matter is complex and has different levels
of organization:
Mixtures
Pure substances
Mixtures…
• Most of the matter around us consists of
mixtures of pure substances
• Mixture – consists of materials with
variable compositions
• Two types of mixtures
Homogeneous
heterogeneous
Homogeneous mixture…
• Homogeneous mixture – having visibly (to the
naked eye) indistinguishable parts
• Has one phase present
• It also called a solution
• There are two parts of a solution
1. the solute – the part that is dissolved
2. the solvent – the part that does the
dissolving
WATER IS THE UNIVERSAL SOLVENT!!
Alloy…
• A homogeneous mixture of metallic
elements with one solid phase
Alloy Problems…
Conversion of Total Mass to Mass of a
Component
• Manganese steel is very strong and finds
use as railroad rails. It is composed of
86.0% iron, 13.0% manganese, 1.0%
carbon. What is the mass of each of the
three elements in a 254-kg sample of
manganese steel?
One more try…
• A sample of brass is composed of 72%
copper and the remainder zinc. What
mass of brass can be made from 25-kg of
zinc?
Examples of solutions…
• Air – a gaseous mixture of various gases
• Brass – a solid mixture of various metals
• Iced Tea – a liquid mixture of various
materials
Heterogeneous mixtures…
• Heterogeneous mixture – having visibly
distinguishable parts
• Can usually be separated into two or more
homogeneous mixtures or pure
substances
• Has 2 or more phases present
Examples of heterogeneous
mixtures…
• Sand and water
• Iced tea and ice cubes
• Pepperoni pizza
Pure substances…
• Mixtures can be separated into pure
substances by physical methods
• Pure substance – contains materials with
a constant composition, such as
compounds and elements
• Have a definite composition and definite
unchanging properties (both chemical and
physical)
Classify the following as a mixture
or a pure substance
•
•
•
•
•
•
•
•
•
Salt water
Smog
Water
10 karat gold
Sugar
Diamond
Coffee
Chex mix
Chef Salad
Properties…
• We use two type of properties to describe
people
o Physical properties – what we look like
o Emotional/Personality – how we interact with other
people
We also use properties to describe matter
o Physical properties – what the matter looks like
o Chemical properties – how matter interacts with other
matter
Physical Properties…
• A physical property is a characteristic that
can be observed or measured without
changing the identity of the substance
• It describes a substance
Color
State at room temperature
Melting point
Boiling point
Density
Specific gravity
Physical Changes…
• A physical change is a change in a
substance that does not involve a change
in the identity of the substance
• A change of state is a physical change of
a substance from one state to another
Solid
Liquid
Gas
Plasma
Chemical Properties…
• A chemical property relates to a
substances ability to undergo changes that
transform it into different substances
• Chemical properties are easiest to see
when substances react to form new
substances which has different properties
than the original substances
Chemical change…
• A chemical change is a change in which one or
more substances are converted into different
substances
• It can also be called a chemical reaction
 Reactants – original substances
 Products – the new substances formed
• Chemical changes or reactions form products
whose properties differ greatly from the
reactants
• Chemical changes do not affect the total amount
of matter (and also the mass) present before
and after a reaction
Methods for Separating Mixtures…
• There are many different ways to separate
mixtures based on their physical properties
 Distillation – uses the volatility of the components
 Filtration – used when a mixture consists of a solid
and a liquid
 Chromatography – separates based on speed of
movement of the components of the mixture
 Solubility – uses the amount of solute that dissolves
in water at a given temperature
 Density
 Melting point
Energy Transfers…
• Physical and Chemical changes are
always accompanied by energy changes
• One way energy can be transferred is
through a temperature difference and is
called heat (q)
• The quantitative measurements in energy
changes are expressed in joules (J)
Energy…
• Energy is an important concept in chemistry and is a
property of all matter
• The ability to do work
• All objects possess energy
• Forms of energy:
1. Chemical energy – released as heat energy
2. Nuclear energy
3. Mechanical energy
4. Electrical Energy
5. Light energy
6. Radiant energy (not a property of objects) – the transfer of
energy through empty space
Categories of Energy…
•
There are two main categories of all
energy we see:
1. Potential energy – the energy stored in the
object
2. Kinetic energy – the energy due to motion
Laws of Matter and Energy…
•
•
•
•
Law of Conservation of Matter
Law of conservation of Energy
Law of Conservation of Mass-Energy
Mass and energy are interconvertible
Matter can change forms
Energy can change forms
Mass can change into energy
• However, never can mass or energy be
created or destroyed
Review Questions…
• What is the Law of Conservation of
Energy?
• What are the two main types of energy?
• Classify either as a chemical change or a
physical change
•
•
•
•
•
Digestion of a candy bar
Melting of ice
Formation of clouds
Growth of plants
Fading of dye in a cloth
A little more about the
Properties of Matter…
• Every substance, whether it is an element
or compound, has characteristics
properties
• These properties are used to distinguish
between substances and to separate them
Properties of Matter con’t…
• Properties can also be characteristics of
an entire group as with metals
• Properties can help identify unknown
substances – notice the plural – not just
one property can identify a substance
• Properties are either intensive or extensive
Intensive – does not depend on the amount of
matter present
Extensive – does depend on the amount of matter
present
States of Matter…
• Matter exists in three physical states:
 Gas – also known as vapor
* has no fixed volume or shape
* it takes the shape of the container by either
compression or expansion
 Liquid – has a definite volume but no specific
shape
 Solid – is rigid and has a fixed volume and a
fixed shape
• Neither liquids or solids are compressible to any
appreciable extent
Phase Change…
• The change of matter from one state to
another state
• It is a PHYSICAL change
• Energy changes ALWAYS accompany a
phase change between the three states
Changes in state…
• Melting/freezing – solid
liquid
• Sublimation/deposition – solid
gas
• Vaporization/condensation – liquid
gas
A Graphical Representation…
gas
liquid
solid
Some other terms associated with
phase changes…
• The melting point is the same temperature
as the freezing point differing only in the
direction from which the phase change is
approached
• Boiling point of a liquid is related to the
pressure – if you increase the pressure,
you increase the boiling point and visa
versa!
Compare and Contrast the Phases
of Matter…
Solid:
Similarities:
Liquid:
Gas:
Review Questions…
1. Describe the three phases of matter?
2. Classify the following as extensive or intensive
– color, mass, length, melting point, ductility
3. Classify the following as a chemical or physical
property
•
•
•
•
•
•
Reactivity
Odor
Rusting
Stability
Expansion
Porosity
Heat Energy
• Although there are different forms and
types of energy that are important to
chemistry, heat energy has the most
relevance to chemical changes
• The most obvious thing about heat energy
is that it causes changes in the
temperature of matter
• Just like matter, it too can be quantified –
how much heat?
How much heat?
• Heat can be lost or gained by a reaction
during a chemical change
Exothermic – when a chemical reaction
releases heat
Endothermic – when a chemical reaction
absorbs heat
Measuring Energy Changes…
• Experimentally, energy changes of
chemical reactions are measured in a
calorimeter
• To change the temperature of a
substance, heat must be added or
removed – some substances require little
heat to cause a change while others
require a great deal of heat
Specific Heat…
• The heat needed to raise the temperature
of one gram of a substance by one Celsius
degree is called the specific heat (Cp) of
a substance
• Every substance has its own specific heat
• Water has a specific heat of 4.18 J/g°C
while aluminum has a specific heat of
0.900 J/g°C
Calculating Specific Heat
• In calculating specific heat, some items must be
understood – the heat lost by one substance is
gained by another until equilibrium is reached
• The equation used:
q = (m)(ΔT) (Cp)
q – energy change
m – mass
ΔT – Tfinal – T initial
Cp – specific heat
Some examples…
How much heat is lost when a solid aluminum
ingot with mass 4100 g cools from 660.0 °C to
25 °C?
q=?
m = 4100g
ΔT = (660 – 25) = 635°C
Cp = 0.900 J/g°C
q = 4100g 635°C 0.900J = 2.35 x 106 J
g°C
Try these…
• How much heat is required to raise the
temperature of 54.5 g PCl3 from 18.6°C to
79.1°C? Cp = 0.874
• How much heat is required to raise the
temperature of 7.90 x 102 g H2O from
38.4°C to 85.4°C? Cp = 4.18 J/g°C