Introduction to
General Chemistry
Lecture 1
Suggested HW:
5, 7, 8, 11, 13, 15,
16, 27
Ch. 1.1- 1.5
What is Chemistry?
• Chemistry is the study of properties of substances and how
they react
• Chemical substances are composed of matter
– Matter is the physical material of the universe; anything
with mass that occupies space is matter
– Matter can take numerous forms
– Most matter is formed by unique arrangements of
elementary substances called elements
Elements, Compounds, and Molecules
• An element can easily be defined as a substance that can
not be broken down into simpler substances
• Millions of different materials in the world, all comprised
of some combination of only 118 elements
– Similar to how the alphabet combines 26 letters to
yield hundreds of thousands of words, elements bond
in unique arrangements to give different molecules
– Molecules agglomerate to yield compounds
Molecules Are Comprised of Uniquely Arranged Atoms.
Different Molecules Have Different Properties.
O
H
O
H
O
H
H
Small Molecular Differences Can Yield Vastly Different
in Terms of Biological Interactions
acetaldehyde
(hangover)
acetic acid
ethanol
BLINDNESS!!!
methanol
carbon dioxide
Small Molecular Differences Can Yield Vastly Different
in Terms of Biological Interactions
Relief of Morning Sickness
Severe Limb Defects
Small Molecular Differences Can Yield Vastly Different
in Terms of Biological Interactions
C6H12O6
C6H12O6
Technically, glucose and fructose are the same. So is high fructose corn
syrup really that bad for you???
Small Molecular Differences Can Yield Vastly Different
in Terms of Biological Interactions
The Properties of Molecules Differ Vastly from those of the
Atoms That Comprise Them
Na (sodium metal)
Cl2 (chlorine gas)
Na+Cl-
Different atomic arrangements can change physical
properties
Carbon (graphite vs diamond)
Spatial Dimensions of Compounds Can Alter Properties
Gold Nanoparticles
Bulk Gold
5 nm
50 nm
Spatial Dimensions of Compounds Can Alter Properties
2 nm
12 nm
CdSe quantum dots
Phases of Matter: Solids, Liquids and Gases
Solids
• Atoms tightly bound
• Fixed volume and shape (does not
conform to container)
• A chemical is denoted as solid by
labeling it with (s)
S(s)
Phases of Matter: Solids, Liquids and Gases
Liquids
• Atoms less tightly bound than
solids
• Has a definite volume, but not
definite shape (assumes the
shape of its container)
• Denoted by (L)
H2O (L)
Phases of Matter: Solids, Liquids and Gases
Gases
• Free atoms
• No shape, no definite volume
• Can be expanded or compressed
(like engine piston)
• Denoted by (g) ; ex. O2 (g)
Qualitative and Quantitative Analysis
• In chemistry, the scientific method is used to investigate
scientific phenomena & acquire new knowledge
• Empirical evidence is gathered which supports or refutes a
hypothesis
• Empirical evidence is either quantitative or qualitative
– Quantitative data is numerical, and results can be
measured
– Qualitative data is NOT numerical, but consists of
observations and descriptions
Quantitative and Qualitative Analysis
A+B
Quantitative data
• How much C is formed?
• How efficient is the
reaction?
• What is the rate of the
reaction?
C
Qualitative data
• What color is it?
• Is it solid, liquid, gas?
• How does it smell?
Units
• Quantitative measurements are represented by a:
NUMBER and a UNIT
• A unit is a standard against which a physical quantity is
compared physical quantity
– Temperature is measured in Co, Ko,or Fo
– Currency is measured in $USD
– Distance is measured in meters, miles, ft, etc.
– Time is reported in seconds, minutes, hr, etc.
• Internationally accepted system of measurements is called
the SI unit system
SI Unit System: The Units of Physical Science
Greek Prefixes
• Prefixes indicate powers of 10
– ex. k= 103; 5 kg = 5 x (103)g
A Review of Scientific Notation
Scientific notation indicates a factor (F) multiplied by a power (n) of 10
F x 10n (1 < F < 10)
• Important: All integers end with a decimal point, even though
it is not commonly written (1  1. )
• If no factor is shown, assume there is a 1. in front of powers
of 10:
102 = 1. x 102
10-7 = 1. x 10-7
• For every positive power of 10, shift the decimal 1 place to
the right, add a zero for each place
102 = 1. x 102 = 100.
105 = 1. x 105 = 100000.
A Review of Scientific Notation
• For all non integers, simply shift the decimal.
2.5 x 105 = 250000.
1.8773 x 108 = 187730000.
• For negative exponents, shift the decimal left. All values
less than 1 have negative exponents.
7.141 x 10-2 = .07141
3.867 x 10-7 = .0000003867
Convert to standard notation
• 3.4912 x 104
• 8.971 x 10-3
• 6.50 x 100
Convert to scientific notation
• 15
• 125.3
• 0.003003
Multiplying and Dividing Exponents (Review)
• When multiplying powers of 10, the product is the sum of the
powers
– 102 x 105 = 10 2+5 = 107
– (2.5 x 103) x (4 x 10-6) = (2.5 x 4) x (103+(-6)) = 10 x 10-3 = 1.0 x 10-2
• When dividing powers of 10, subtract
– 102 / 105 = 10 (2-5) = 10-3
– (6.6 x 1010) / (2.2 x 10-6) = 3.0 x (10 10-(-6)) = 3.0 x 1016
Group Work
Convert the following values to grams in proper scientific
notation.
– 421.4 kg
– 1170.1 mg
– 481 µg
Why Are Units Important? Example #1
• In 1999, NASA lost the $125M
Mars Orbiter System.
• One group of engineers failed to
communicate with another that
their calculated values were in
English units (feet, inches,
pounds), and not SI units.
• The satellite, which was
intended to monitor weather
patterns on Mars, descended
too far into the atmosphere and
melted.
Why Are Units Important? Example #2
• In 1983, an Air Canada Plane
ran out of fuel half way through
its scheduled flight. Why?
• Airline workers improperly
converted between liters and
gallons.
• Luckily, no one died.
Why Are Units Important? Example #3
• A case was reported in which a
nurse administered 0.5 g of a
sedative to a patient.
• The patient died soon after
• The patient should have only
received 0.5 grains (≈ 0.033 g)
but the units were not listed.
That was the equivalent of 8
doses!!
Derived SI Units: VOLUME
• Many measured properties have units that are combinations of
the fundamental SI units
• Volume: defines the quantity of space an object occupies; or the
capacity of fluid a container can hold
– expressed in units of (length)3 or Liters (L)
– 1 L is equal to the volume of fluid that a cube which is 10 cm on
each side can hold
V = (10 cm)3 = 1000 cm3
1 L = 1000 cm3
10
cm
1000 mL = 1000 cm3
10 cm
10 cm
mL = cm3
Derived SI Units: DENSITY
• All matter has mass, and must therefore occupy space. Density
correlates the mass of a substance to the volume of space it
occupies.
• Density = mass per unit volume (mass/volume). Different
materials have different densities.
Would a 20-gallon
filled with bricks
same mass as an
volume of feathers?
container
have the
equivalent
NO!
g
 feather  0.025 3
cm
g
 brick  1.90 3
cm
THE DENSITY OF WATER IS 𝟏
𝒈
𝒄𝒎𝟑
𝒐𝒓 𝟏
𝒈
𝒎𝑳
Group Work
• A cubic container that is 25 cm on each side is filled with
ethanol. The density of ethanol is 0.79 g/mL.
– What is the volume of ethanol in the cube in mL?
– What is the volume in L?
– What is the mass of ethanol in kg?
Give answers in scientific notation!!
Derived SI Units: ENERGY
• What is Energy?
– Energy is defined as the capacity to perform “work”
• How do we define work?
• Work is defined as the action of applying a force
acting over some distance. Work can not be done
if no energy is available.
• In SI units, we use the unit Joule (J) to represent energy.
𝑘𝑔 𝑚2
𝐽=
𝑠2
Conservation of Energy
Energy is never created or destroyed, merely converted between
forms and transferred from place to place. The total energy of
the universe is finite.
Forms of Energy
• Energy comes in many forms and can be converted from one
form to another. Some examples are given:
• Chemical Energy
– Energy stored in chemical bonds (e.g. gasoline, coal, etc.)
that can be released by chemical reaction, typically
combustion (fire)
• Heat Energy (thermal energy)
– Heat is defined as energy flow between bodies of matter
resulting from collisions of molecules or random motions
of electrons.
Forms of Energy
• Mass Energy
– Energy and mass are interchangeable. During a fusion
reaction (e.g. stars), mass is lost and energy is formed.
This mass appears as energy according to the following:
𝐄 = ∆𝐦𝐜 𝟐
where m is the change in mass (in kg), c is the speed of
light, and E is the energy released (J). This is the basis of
nuclear power.
• Kinetic Energy
– Energy of motion (e.g. a moving car). An object with mass
m, moving at a velocity V (meters/sec) has kinetic energy:
𝟏
𝐄𝐤 = 𝐦𝐕 𝟐
𝟐
Forms of Energy
• Potential Energy
– Potential energy corresponds to energy that is stored as a
result of the position of mass in a field.
• If a mass m is held at a height h (meters) above the
ground, assuming a gravitational acceleration of 9.8
m/s2 (g), its potential energy is:
𝐄𝐏 = 𝐦𝐠𝐡
– If the object is dropped, it loses potential energy. However,
it speeds up as it falls, so its kinetic energy increases
equally (conversion).
Forms of Energy
• Electrical Energy
– Energy resulting from electric current, the movement of
electrons through a conductive circuit. Electrical energy is
a type of potential energy. For a charge q (coulombs, C)
moving across a voltage, V
𝐄𝐞𝐥𝐞𝐜 = 𝐪𝐕
• Light/Radiation
– The energy of a wave of light is calculated as the product
of Planck’s constant, h (J s), and the wave frequency, ν
(1/s)
𝐄 = 𝐡𝐯
Power
• It is often necessary to express the rate of energy usage. This
is called power.
𝐞𝐧𝐞𝐫𝐠𝐲
𝐏𝐨𝐰𝐞𝐫 =
𝐭𝐢𝐦𝐞
• Typically, we speak in terms of energy per second. In SI units,
a joule per second (J/s) is known as a watt (W).
Temperature
• Temperature: a measure of the tendency of a substance to
lose or absorb heat. Temperature and heat are not the same.
• Heat always flows from bodies of higher temperature to those
of lower temperature
– The stove top is ‘hot’ because the surface is at a much
higher temperature than your hand, so heat flows rapidly
from the stove to your hand
– Ice feels ‘cold’ because it is at a lower temperature than
your body, so heat flows from your body to the ice, causing
it to melt
Temperature
• When performing calculations in chemistry, temperature
must always be converted to Kelvin (oK) units (unless
otherwise stated).
• The lowest possible temperature that can ever be reached
is 0oK, or absolute zero. At this temperature, all molecular
motion stops.
• To convert temperatures to the Kelvin scale:
oK
: oC + 273.15