Chapter 1
An Introduction to Chemistry
Definition of chemistry
“Chemistry is the science of composition, structure,
properties and reactions of matter, especially atomic
and molecular systems.”
Broad field: includes everything in the whole universe, animate and inanimate.
Agriculture – find a wheat kernel that makes best bread
Astronomy – what matter exist on Mars?
Animal science – find the best animal feed
Geology – composition of rocks, water inclusion in lava
Medicine – MRI
Material science – catalysts for fuel cells
Molecular biology – protein structure and folding
Despite the common belief that a chemist mixes chemicals
just for the love of it – and gets unimaginable results – the truth is different:
a chemist tries to understand the nature and natural phenomena,
composition and structure of complex systems (like biomolecules) or to synthesize new substances in order to advance
science in general and contribute to the well-being of humanity.
Scientific method
- Collect data or facts relevant to the
problem and organize in data sets;
- Formulate prediction (hypothesis)
to explain the data;
- Plan and test the hypothesis with
additional experiments;
- Modify the hypothesis as needed.
Data
Data sets
Hypothesis – tentative explanation
Theory – well-established hypothesis;
explanation of general principles of certain
phenomena with facts to prove it.
- Einstein’s theory of relativity
Scientific laws – simple statements of
natural phenomena to which no
exceptions are known.
- Law of buoyancy: any
object submerged in a fluid
is acted on by a force
with magnitude equal to
the weight of the fluid
displaced by the object.
Theory
Aristotle believed that women
have less teeth than men.
Hypothesis
observation
hypothesis
experiment
theory
laws Bias- a
strong
preference
that inhibits
impartial
judgment.
Matter
… is anything that occupies space and has mass; or
… is anything that one can see or touch.
Matter is composed of tiny particles called atoms. It can
be in any of the three physical states: solid, liquid or gas.
The physical state in which the matter exists depends on
the nature of the matter (i.e. attractive forces in the atoms
or molecules of the matter), and external factors (P, T).
Physical transformations, or changes of physical states
are: freezing / melting (solid – liquid), vaporization /
condensation liquid-gas, and sublimation (solid-gas).
Pure substances are
particular form of matter
with definite and fixed
composition, and cannot
be separated by any
physical method. They are
divided into elements
(copper, gold aluminum)
and compounds (salt,
sugar, water).
Matter is divided into pure substances, and mixtures.
Mixtures can be:
homogeneous - uniform in appearance and with
same properties throughout, also called: solution, or
heterogeneous (two distinct phases with boundary
between them).
Mixtures contain two or more pure
substances with variable
composition; the substances can
be separated by a physical method
(e.g. by using a magnet, or by
boiling, filtration, etc.).
Physical state and properties
Different attractive forces in different matter determine
the transition point from one physical state to the next.
State
Solid
Liquid
Gas
Shape
Definite
Indefinite
Indefinite
Volume
Definite
Definite
Indefinite
Compress.
V. Slight
Slight
High
State of common materials
Solids
Most metals
(gold, copper, zinc)
Salt
Sugar
Sand
0
100 Temperature, oC
Liquids
Water
Alcohol
Vinegar
Blood
Oil
Gases
Air (oxygen,
nitrogen,…)
Acetylene
Chlorine
Noble gases
(He, Ne …)
Properties and changes of matter
Set of properties gives unique identity to the matter.
Physical properties can be determined without destroying the substance (state,
color, density, taste/odor, melting/boiling point).
Chemical properties describe the ability of the substance to form new substances,
either by decomposition or reaction with other substances (corrosiveness,
flammability, acidity, toxicity, etc.).
No two substances have identical physical and chemical properties.
Physical changes are reversible changes (e.g. changes in size, shape, state of
matter, density) and no new substances are formed.
Heating of Pt wire changes its appearance from silvery metallic to glowing yellowred. Upon cooling, original appearance is restored and no new compounds are
formed. Shattered glass looks different, but its (chemical) composition is the same.
Chemical changes are irreversible as new compounds are formed. Upon
heating, copper wire glows yellow-red like platinum, but appears black after it is
cooled. Copper (Cu) oxidizes to copper(II) oxide (CuO).
Physical change results in a
different form of the same
substance.
Chemical change results in
a completely different
substance(s).
Chapter 2
Measurements and significant figures
Measurement is a comparison of a property of your sample with that of the standard
Measurement is important – a small variation of a drug quantity may kill you!
Measurement never gives an exact value. Its numerical value always contains
all known figures the measuring tool provides, plus one that is estimated by
the measurer. These are called significant figures.
Certain
figures:
Measurement of a coin
Using
a
ruler
Estimated
2, 5
using a ruler marked in
marked in
figure: 5
centimeters gives the
millimeters makes
diameter to be about
the first decimal a
half-way between 2 and
certain figure.
3 cm, or about 2.5 cm,
Estimated
with the uncertainty in
figure: 0
the first decimal.
Certain figure: 2
This measurement gives the diameter
This measurement gives the diameter of
of 2.50 cm, or three significant figures.
2.5 cm. That is, two significant figures.
Uncertainty is +/- 0.01.
Uncertainty is +/- 0.1.
Accuracy: closeness of the
measurement to the true value.
Precision: closeness to one another of a series
of measurements made on the same object.
Using the first ruler we
found the diameter of
the coin to be 2.5 cm
(uncertainty +/- 0.1).
What is the radius of
the coin?
Radius = diameter / 2. Using your
calculator you find: 2.5 cm / 2 = 1.25 cm.
BUT: the second decimal place is beyond
the level of uncertainty. Thus, you must
round the number to a single decimal
place, i.e. 1.3 cm.
You can keep only significant figures – i.e. certain figures read by the
measuring tool, and an additional figure representing your estimate obtained by
reading between the smallest markings of your measuring tool.
Note that there are exact numbers: there is exactly 2 radii in a diameter!
Significant Figure Rules –
1. all non-zero digits are significant.
2. zero is NOT significant if:
– it is before the first nonzero digit or
- at the end of a number without decimal point.
Practice:
20.201
20.210
20.0002
120.
# of sig. figs.:
5
5
6
3
Leading (placeholder) zeroes only
give the location of the decimal
point and are NEVER significant.
0.06 inch has only one sig figs.
Decimal point tells you that the last zero is
an estimate (uncertainty +/- 1).
Without the decimal point, the number 120 can
be interpreted as 120 +/- 10, thus two sig. figs.
To avoid ambiguity, we use scientific notation.
1.2 x 102
Scientific notation
x=2
A = 1.2
1.2 x 100 = 120
Any number can
be presented
with scientific
notation
Scientific notation
A x 10x
A is a number equal or greater than 1 and
smaller than 10. A gives the # significant figures.
102 = 10 x 10 = 100
1000 = 1 x 103
100 = 1 x 102
10 = 1 x 101
1
= 1 100
0.1 = 1 x 10-1
0.01 = 1 x 10-2
0.001 = 1 x 10-3
10-2 =
1.23 x 102
2 places right
1.23 x 10-1
1 place left
0.123
1.23 x 100
doesn’t move
1.23
4.6 x 104
4 places right
4.6 x 10-4
4 places left
2 places left
1.2 x 10-1
1 place right
!
Conversion from scientific to regular notation:
When x is positive, you move the decimal
point x places to the right.
When x is negative, you move the
decimal point x places to the left.
Decimal point moves: Standard notation
1.20 x 102
1
102
x is an exponent
(or power) of 10.
Exponent x is an
integer (whole
number).
123.
no need for
decimal point
100 = 1 !
no decimal point!
Add 3 zeros after six
46,000
0.00046
120.
0.12
Add 3 zeros before four
Reverse movement of decimal
point when going from regular to
sci. notation
When doing math
Calculator often gives more digits than justified.
Rounding rule – use all digits you want to retain. Then check the digit
after the last one retained:
- If it is 4 or less, drop it and all the digits behind. Leave the retained number
unchanged.
- If it is 5 or more, drop it and all the digits behind. The last digit in the retained
number increases by one.
To avoid accumulating errors, always round the number at the end of calculation.
Significant figures rules for multiplication and division:
The result cannot have more significant figures than there are in the
measurement that has the smallest number of significant figures.
2.0 cm X 2 cm = 4 cm2
2 sig. figs. 1 sig. fig. 1 sig. fig.
2.0 cm X 2.0 cm = 4.0 cm2
2 sig. figs. 2 sig. fig. 2 sig. fig.
Significant figures rules for addition and subtraction:
The result can be no more certain than the least certain measurement in the
123 cm
series.
20.5 lb + 15 lb
35.5 lb
=
= 14.489…
+ 1.006 cm
2
2.45 ft x 1.00 ft
2.45 ft
+ 2.6 cm
= 14 lb/ft2
Calculator gives
126.606 cm
Uncertainty +/- 1 127 cm
Uncertainty +/- 1 Should be rounded to 36, or 2
significant figures... But then
the result would be 15 !
Number with a name - Unit of measure
A measurement makes no sense unless units are specified.
Say, you are to get to the airport in 3. Three what?
Minutes? Hours? Gallons of gas?
Remembering names of different units for the same
property, and the relationship between them can be
cumbersome....
12 inches in a foot, 3 feet in a yard, 22 yards in a chain,
10 chains in a furlong, 8 furlongs in a mile …
Still in doubt? Try yourself: how many feet in a yard?
8 fl. Oz in a cup, 4 cups in a quart, 4 quarts in a gallon,
31 gallons in a barrel…
… and if you make a mistake in conversion from one unit
into another, you are toast… like the princes feeding her
pet dragons. Even today!
NASA wasted $125 million for a Martian Lander that burned
up in the red planet’s atmosphere because of ambiguously
expressed retrorocket force. The company that built the
rockets expressed the force in US Customary units (in, ft,
lbs), whereas NASA assumed the data was given in metric
units (meters, kilograms, Newtons).
SI System
Whole world (except U.S.!) now use standardized system of units, called SI system.
SI system uses an unique name for a physical property and prefixes to describe
bigger or smaller units. The relationship between bigger & smaller units is in the
factor of 1000.
There is even relationship between units for different physical properties!
Base SI Units
Physical quantity
Length
Mass
Time
Temperature
Amount of subst.
Name
meter
kilogram
second
Kelvin
mole
Abbrev.
m
kg
s
K
mol
Derived SI Units
Physical quantity
Volume
Pressure
Energy
Electric charge
Name
Abbrev.
cubic meter m3
Pascal
Pa
joule
j
coulomb
C
A cube of 0.1 m length has a volume of 1 L; filled with water, its mass is 1 kg.
Because 1 m3 is a huge volume,
we commonly use non-SI units for
volume: Liter (L) and milliliter (mL).
1 m3 = 1000 L. 1 L = 1000 mL
1 L = (0.1 m)3 = 1 dm3 1 L = 1.057 qt.
1 m is 3.37 inches longer than a yard.
1 kg = 2.2046 lb
1 m = 1.0936 yd
Prefix - milli
SI System (continued)
Physical quantity: meter
1mm = 10-3 m
1km = 103 m
Other commonly used, non-SI units are
Pressure in: millimeters of mercury (mm Hg),
atmospheres (atm)
Temperature in: degrees of Celsius (oC)
1µm = 10-6 m
1mL = 10-3 L
1 MB = 106 B
1GB=109 B
Prefixes Used with SI Units
ten (101)
one hundred (102)
one thousand (103)
one million (106)
one billion (109)
one trillion (1012)
Rarely
used
Note: All prefixes for smaller units are lower case,
and those for commonly used larger are upper
case! Prefix kilo – is an exception.
Water boils: 373.15 K
Absolute zero: 0 K
Lowest possible
temperature
Note: Kelvin is
represented as
K, not oK!
100 oC
212 oF
0 oC
-273.15 oC
180 degrees
Water boils: 373.15 K
100 degrees
Meaning
one-trillionth (10-12)
one-billionth (10-9)
one-millionth (10-6)
one-thousandth (10-3)
one-hundredth (10-2)
one-tenth (10-1)
100 degrees
Greek Prefix
pico- (p)
nano- (n)
micro- (µ)
milli- (m)
centi- (c)
deci- (d)
unit
deca- (da)
hecta- (h)
kilo- (k)
mega- (M)
giga- (G)
tera- (T)
32 oF
-459.67 oF
K = oC + 273.15
oC = (5/9) (oF – 32)
oF = (9/5) oC + 32
Volume and Density
Density = mass of substance divided by the volume occupied by the mass: d = m /
V. By definition, the density of water is exactly 1 kg / L (hence the definition of L!)
Dimensionless specific gravity is the ratio of densities of a substance to that of water.
When immersed in water,
substances with density > 1 kg/L
(e.g. most metals) sink, and those
having density less 1 (ice, most
liquids, wood, etc) float.
Density depends on the way atoms
are packed in the crystal lattice.
Density
B>A
Volume of an
object with
orthogonal sides
can be caluclated
as Length x
Width x Height.
W
H
L
Volume of an
irregularly
shaped object
can be measured
by displacement
of water.
Density is an intensive property, i.e. does not depend on the amount of material.
The opposite is extensive property, one that depends on amount present.
Mass and volume are both extensive properties, but their ratio (density) is intensive.
Unit Analysis
Recall – density is the ratio of mass and volume.
d=m/V
What to do if you are given the mass and the density of mercury, and asked
to calculate the volume? (Suppose m = 2.00 g and d = 13.6 g/mL)
Hint:
Use above eq. to find V as a function of the other two quantities.
HOW?
Recall that equation stays unchanged if a mathematical
operation is applied to both sides of equal sign.
Multiply both sides with V:
Vx d=
Now divide both sides with d:
Vxd
d
Recall fractions:
=
x
V
m
=
d
2.00 g
=
V=
g
13.6
mL
Plug in the numbers:
1
2
3
4
m
solved for m
1
2
x
4
3
=
1x4
2x3
Check
the units:
V
Vxd=m
V =
m
d
solved for V
0.147 mL
g g x mL
g
1
g = g = g x 1 = mL
mL
mL
Unit Conversion
Conversion factors are used to introduce the desired unit and cancel the given one.
Example 1: how many minutes are in 13 hours?
exact number –
no need for
decimal point !
1 hr = 60 min
Use the identity relationship between the units.
Conversion
factors:
60 min
or
1 hr
There are two conversion
factors for each identity.
1 hr
60 min
Q: Which of the two conversion factors should I use?
A: The one that introduces the desired unit and cancels the given one!
Desired unit is in numerator,
the given one in denominator.
60 min
1 hr
Write down the given quantity:
The use of the other conversion
factor would not cancel hours:
Since 1 hr is identical to 60 min, the conversion
factor ratio is equal to ONE. When a number is
multiplied by one the number is unchanged.
13 hr
13 hr
60 min
x
x
1 hr
1 hr
60 min
= 780 min
= 0.22
hr2
min
not a unit for time!
Q: What to do if I cannot make one unit disappear immediately?
A: Try and try again!
Use multiple conversion factors; keep multiplying until you reach the desired
unit(s) and have all other “unwanted” units disappear.
Example 2:
1 week
How many seconds equals to (exactly) one week?
7 day
x
x
1 week
exact number
24 hour
x
1 day
3600 s
= 604,800. s
1 hour
1 week = 7 days
1 day = 24 hours
1 hour = 3600 s
Example 3: Convert 5.5 barrels into milliliters.
(Assume 5.5 barrel is a measurement.)
31 gal
5.5 barrel
x
3.7854 L
x
x
1 barrel
2 significant figures
1000 mL
1 gal
1L
exact number:
technically, the
decimal point is
not needed’; to
avoid confusion
about significant
figures it is better
to put it here.
= 586,737 mL
= 590,000 mL
1 barrel = 31 gal
1 gal = 3.7854 L
or 5.9 x 105 mL
1 L = 1000 mL
Homework:
Chapter 1: Review questions (p.10):
1, 6, 12, 13, 15
1. Explain the difference between:
(a) A hypothesis and a theory;
(b) A theory and a scientific law.
6. How many phases are present
in the graduated cylinder?
12. Is the system that contains
only one substance necessarily
homogeneous? Explain. (Hint:
ice in water!)
13. Is the system that contains two
or more substances necessarily
heterogeneous? Explain (Hint: salt
and sugar in water!)
15. Which of the following are pure substances?
a) table sugar
b) sand
c) gold
d) maple syrup
e) egg
Chapter 2: Paired exercises (p.39):
4 (a-d,f), 7, 9, 13, 18(b), 23, 41, 49, 55
4. State the abbreviation for each of the following units:
(a) milligram, (b) kilogram, (c) meter. (d) nanometer, (f)
microliter.
7. How many significant figures are in each of
the following numbers? (a) 0.025, (b) 22.4,
(c) 0.0404, (d) 5.50 x 103
9. Round each of the following numbers to three
significant digits: (a) 93.246, (b) 0.02857, (c)
4.644, (d) 34.250
13. Solve the following problems:
(a) 12.62 + 1.5 + 0.25 = ?
(b) (2.25 x 103) (4.80 x 104) = ?
(c) (452) (6.2) / (14.3) = ?
(d) (0.0394) (12.8) = ?
(e) (0.4278) / (59.6) = ?
(f) 10.4 + (3.75) (1.5 x 104) = ?
18. Solve equation for the variable x:
(b) 8.9 g/mL = 40.90 g / x (Hint: x is in mL!)
23. After you have worked out at a gym on a
stationary bike for 45 min, the distance gauge
indicates you have traveled 15.2 miles. What was
your rate in km/hour? (Hint: convert miles into km
and min into hr, then divide the two.)
41. A textbook is 27 cm long, 21 cm wide and 4.4
cm thick. What is the volume in (a) cubic
centimeters, (b) liters, (c) cubic inches? (Hint: (a)
multiply LxWxH; (b) since 1 cm3=1 mL, divide
result in (a) by 1000; (c) convert each dimension
into inches, then multiply them.)
49. The average temperature of Venus is 460 oC.
What is this temperature in oF?
55. Linseed oil has a density of 0.929 g/mL. How
many mL are in 15 g of oil? (Hint: V=m/d)