Unit 1: Quantities and Units
Measurement:
Much of the chemistry we do will be quantitative, requiring measured, numbered
quantities. All measurements require a number and a unit associated with it to be
meaningful. All the units we use will be metric, based on the SI system. The SI
system, in place since 1960, has seven fundamental base units from which all other
units are derived. They are
Name
_
Mass
Length
Time
Temperature
Current
Luminous intensity
Amount of a substance
Quantity
(1 kg = 2.2 pounds)
(1 m = 1.0936 yd)
(3600s = 1Hour)
(1mol=6.023x1023molcules)
Table 5. SI prefixes
Factor Name
1024 yotta
1021 zetta
1018 exa
1015 peta
1012 tera
109
giga
6
10
mega
103
kilo
2
10
hecto
101
deka
Symbol
Y
Z
E
P
T
G
M
k
h
da
Factor
10-1
10-2
10-3
10-6
10-9
10-12
10-15
10-18
10-21
10-24
Name
deci
centi
milli
micro
nano
pico
femto
atto
zepto
yocto
Symbol
d
c
m
µ
n
p
f
a
z
y
The English system of
measurement grew out of the
creative way that people measured for
themselves. Familiar objects and
parts of the body were used as
measuring devices. For example,
people measured shorter distances on
the ground with their feet.
Symbol
(kg) kilogram or (g) gram
(m) meter
(s) second
(K) Kelvin
(A) Ampere
(cd) candela
(mol) mole
Conversion Tables
Metric System of Measurements
Length
10 millimeters = 1 centimeter
10 centimeters = 1 decimeter
10 decimeters = 1 meter
10 meters
= 1 decameter
10 decameters = 1 hectometer
10 hectometers = 1 kilometer
1000 meters
= 1 kilometer
Volume
1000 mm3
= 1 cm3
3
1000 cm
= 1 dm3
3
1000 dm
= 1 m3
1 million cm3 = 1 m3
Area
100 mm2
10 000 cm2
100 m2
100 acre
10 000 m2
100 hectares
1000000 m2
Capacity
10 mL = 1 cL
10 cL = 1 dL
10 dL = 1 L
1000 L = 1 m3
=
=
=
=
=
=
=
1
1
1
1
1
1
1
cm2
m2
are
hectare
hectare
km2
km2
Mass
1000 g = 1 kg
1000 kg = 1 ton
The U K (Imperial) System of Measurements
Length
12 inches = 1 foot
3 feet
= 1 yard
22 yards = 1 chain
10 chains = 1 furlong
8 furlongs = 1 mile
5280 feet = 1 mile
1760 yards = 1 mile
Volume
1728 in3 = 1 ft3
27 ft3
= 1 yd3
1L
= 1 dm3
1 L
= 0.01 m3
Mass (Avoirdupois)
437.5 grains = 1 ounce
16 ounces
= 1 pound (7000 grains)
14 pounds
= 1 stone
8 stones
= 1 hundredweight [cwt]
20 cwt
= 1 ton (2240 pounds)
20
3
8
20
Apothecaries'
minims
= 1
fl.scruples = 1
fl.drachms = 1
fl.ounces
= 1
Measures
fl.scruple
fl.drachm
fl.ounce
pint
144
9
4840
640
Area
sq. inches
sq. feet
sq. yards
acres
=
=
=
=
1
1
1
1
square foot
square yard
acre
square mile
Capacity
20 fluid ounces = 1 pint
4 gills
= 1 pint
2 pints
= 1 quart
4 quarts
= 1 gallon (8 pints)
Troy Weights
24 grains
= 1 pennyweight
20 pennyweights = 1 ounce (480 grains)
12 ounces
= 1 pound (5760 grains)
20
3
8
12
Apothecaries'
grains
= 1
scruples = 1
drachms = 1
ounces
= 1
Weights
scruple
drachm
ounce (480 grains)
pound (5760 grains)
Scientific Notation
Chemistry deals with very large and very small numbers.
Consider this calculation:
(0.000000000000000000000000000000663 x 30,000,000,000) ÷ 0.00000009116
Hopefully you can see, how awkward it is. Try keeping track of all those zeros! In
scientific notation, this problem is:
(6.63 x 10¯31 x 3.0 x 1010) ÷ 9.116 x 10¯8
It is now much more compact, it better represents significant figures, and it is easier to
manipulate mathematically. The trade-off, of course, is that you have to be able to read
scientific notation.
Format for Scientific Notation
1. Used to represent positive numbers only.
2. Every positive number X can be written as: (1 < N < 10) x 10some positive or negative integer
Where N represents the numerals of X with the decimal point after the first nonzero digit.
3. A decimal point is in standard position if it is behind the first non-zero digit. Let X be
any number and let N be that number with the decimal point moved to standard position.
 Rules to consider:
 The first number must be a number that is  l but < 10
 If the decimal point on any whole number or decimal is moved to the:
 Right, the exponent will be a negative sign
 Left, the exponent will be a positive sign
4. Some examples of number three:
4
 0.00087 becomes 8.7 x 10¯
0
0
 9.8 becomes 9.8 x 10 (the 10 is seldom written)
7
 23,000,000 becomes 2.3 x 10
5. Some more examples of number three:
7
 0.000000809 becomes 8.09 x 10¯
0
 4.56 becomes 4.56 x 10
11
 250,000,000,000 becomes 2.50 x 10
Example #1 - Convert 29,190,000,000 to scientific notation.
First Explanation
Step 1 - start at the decimal point of the original number and count the number of decimal
places you move, stopping to the right of the first non-zero digit. Remember that's the
first non-zero digit counting from the left.
Step 2 - The number of places you moved (10 in this example) will be the exponent. If
you moved to the left, it's a positive value. If you moved to the right, it's negative.
The answer is 2.919 x 1010.
Example 2 - Write 0.00000000459 in scientific notation.
Step 1 - Write all the significant digits down with the decimal point just to the right of the
first significant digit. Like this: 4.59. Please be aware that this process should ALWAYS
result in a value between 1 and 10.
Step 2 - Now count how many decimal places you would move from 4.59 to recover the
original number of 0.00000000459. The answer in this case would be 9 places to the
LEFT. That is the number 0.000000001. Be aware that this number in exponential
notation is 10¯9.
Step 3 - Write 4.59 times the other number, BUT, write the other number as a power of
10. The number of decimal places you counted gives the power of ten. In this example,
that power would be 9. The correct answer to this step is: 4.59 x 10¯ 9
Suppose the number to be converted looks something like scientific notation, but it really
is not. For example, look carefully at the example below. Notice that the number 428.5 is
not a number between 1 and 10. Although writing a number in this fashion is perfectly
OK, it is not in standard scientific notation. What would it look like when converted to
standard scientific notation?
Example #3 - Convert 428.5 x 109 to scientific notation.
Step 1 - convert the 428.5 to scientific notation. (The lesson up to this point has been
covering how to do just this step). Answer = 4.285 x 102.
Step 2 - write out the new number. Answer = 4.285 x 102 x 109.
Step 3 - combine the exponents according to the usual rules for exponents. Answer =
4.285 x 1011.
Example #4 - convert 208.8 x 10¯11 to scientific notation.
Step 1 - convert the 208.8 to scientific notation. Answer = 2.088 x 10 2.
Step 2 - write out the new number. Answer = 2.088 x 102 x 10¯11.
Step 3 - combine the exponents according to the usual rules for exponents. Answer =
2.088 x 10¯9.
1. When converting a number greater than one (the 428.5 and the 208.8 in the previous
examples), the resulting exponent will become more positive (11 is more positive than 9
while -9 is more positive than -11).
2. When converting a number less than one (the 0.000531 and the 0.00000306 in the
previous examples), the resulting exponent will always be more negative (10 is more
negative than 14 and -23 is more negative than -17).
If the decimal point is moved to the left, the exponent goes up in value (becomes more
positive).
If the decimal point is moved to the right, the exponent goes down in value (becomes
more negative).
Physical and Chemical Properties
I. Physical Properties
A physical property of a pure substance is anything that can be
observed without changing the identity of the substance. The
observations usually consist of some type of numerical
measurement, although sometimes there is a more qualitative (nonnumerical) description of the property. There are many physical
properties and each textbook will have a different list of examples.
Here are some of the more common ones:
melting point electrical conductivity color density
boiling point thermal conductivity
odor hardness
There are others which are not mentioned as often.
Examples include:
refractive index
atomic radius ductility
ionization energy allotropes
malleability
II. Chemical Properties
This one is more difficult. Here is one way to define "chemical
property:" characteristics which are exhibited as one substance is
chemically transformed into another. Or one could say that “a
chemical property describes the way a substance may change or
react to form other substances”.
Here are some examples.
(1) iron rusting. When iron (an element, symbol = Fe) rusts, it
combines in a complex fashion with oxygen to form a reddishcolored compound called ferric oxide (formula = Fe2O3). Not all
substances rust.
(2) glucose, mixed with yeast, ferments to make alcohol. Glucose
(C6H12O6) is a chemical compound which enzymes in yeast can use
to make ethyl alcohol (C2H5OH). Not all substances ferment.
(3) trinitrotoluene (TNT) reacts very, very fast when it is ignited.
Among other products, it makes LOTS of nitrogen gas and LOTS
of heat. Inside the proper container, it can cause an explosion. Not
all substances can make an explosion.
I. Physical Changes
A physical change is any change NOT involving a change in the
substance's chemical identity. Here are some examples:
(1) any phase change. Moving between solid, liquid and gas involves
only the amount of energy in the sample (this amount is the subject
of future lessons). There is no effect on the chemical identity of the
substance. For example, water remains water, no matter if it solid,
liquid or gas.
(2) grinding something into a powder. Or the reverse process of
making a bigger lump of stuff, say by melting lots of small pellets of
copper into one big piece.
(3) iron (and other metals) can be made to be magnetic. This change
in no way affects the chemical identity of the element. Iron that is
magnetized rusts just as easily as iron that is not magnetized.
Now would be a good time as any to list the names of the various
phase changes:
Change
Name of change
Solid to liquid melting, fusion
Liquid to gas
boiling, evaporation
Solid to gas
sublimation
Gas to solid
deposition
Gas to liquid
condensation, liquefaction
Liquid to solid freezing, solidification
II. Chemical Changes
A "chemical change" means that the reacting substances(s) are
changed into new substances. The actual atoms involved remain,
they are simply rearranged. The rearrangement is called a
chemical reaction. For example:
2H2O ---> 2H2 + O2
is a chemical reaction in which water is broken down into the
hydrogen and oxygen which make it up. Notice how the amounts of
hydrogen atoms (four) and oxygen atoms (two) do not change from
one side of the arrow to the other. However, the arrangement of
the atoms is different. Some chemical bonds (the one involved in
the water) have been broken and some new chemical bonds (the
one in hydrogen and oxygen) have been formed.
This is another way to define "chemical change:"
A process in which chemical bonds are broken and new ones are
made.
A process like grinding some salt crystals into a fine powder does
not involve the breaking of chemical bonds and the formation of
new ones, so it is a physical change.
Early Theories and Theorists of the Atom
1. Democritus - Greek Philosopher that proposed the concept of the
atom more than two thousand years ago.
 Matter is composed of empty space through which atoms move.
 Atoms are solid, homogeneous, indestructible, and indivisible.
 Different kinds of atoms have different sizes and shapes.
 The differing properties of matter are due to the size, shape, and
movement of atoms.
?What holds an atom together?
I have no idea??????
2. John Dalton - Adapted the theories of Democritus to develop his
own Atomic Theory.
 All matter is composed of extremely small particles called
atoms.
 All atoms of a given element are identical, having the same
size, mass, and chemical properties. Atoms of specific
element are different from those of any other element.
 Atoms cannot be created, divided into smaller particles, or
destroyed.
 Different atoms combine in simple whole number ratios to
form compounds.
 In a chemical reaction, atoms are separated, combined, or
rearranged.
Daltons 3 Main Laws:
Conservation of Mass
Definite Composition
Multiple Proportions
Joseph John Thomson: discovered the electron in a series of
experiments designed to study the nature of electric discharge in a
high-vacuum cathode-ray tube
Rays from the cathode (C) pass through a slit in the anode (A) and
through a slit in a grounded metal plug (B). An electrical voltage is
established between aluminum plates (D and E), and a scale pasted
on the outside of the end of the tube measures the deflection of the
rays.
Ernest Rutherford:
Rutherford's Gold Foil Experiment
> Over 98% of the
particles went straight
through.
> About 2% of the
particles went through
but were deflected by
large angles.
> About 0.01% of the
particles bounced off the
gold foil.
Subatomic Particles: the building blocks for atoms
 Nucleus - the center or core of an atom that contains
positively charged atoms.
 Protons - (p+) Positively charged subatomic particle Discovered
by Ernest Rutherford
 Electrons - (e-) Negatively charged subatomic particle that
surrounds the nucleus. Discovered by J.J. Thompson, Charged
determined by Robert Millikan. (Millikan's Oil-Drop experiment.
 Neutrons - (n) neutral, carries no electrical charge and are
contained in the nucleus with the protons. Discovered by James
Chadwick. Rutherford's coworker.
 Isotopes have different atomic masses, yet identical chemical and
physical properties
 Isotopes are common in nature and influence the measured
atomic masses of large assemblies of atoms
 Different isotopes of the same element have the same name (ex
Carbon-12, Carbon-13) Exception hydrogen, deuterium, tritium
Note: A proton's positive charge is equal to, but opposite to the
negative charge of an electron.
Particle Symbol
Location in
the Atom
Relative
Charge
Relative
Mass
Mass (g)
proton
p
nucleus
+1
1
1.673 × 10-24
neutron
n
nucleus
0
1
1.675 × 10-24
electrons
e-
electron
cloud
-1
~0
9.110 × 10-28
Chemistry Drill
Section Review
1. Define an atom in your own words.
2. Which statements in Dalton's original atomic theory are now
considered incorrect? Explain.
3. Compare and Contrast the atomic theories proposed by
Democritus and John Dalton.
4. Cite one contribution of each scientists toward the development
of the present atomic theory: a. Thompson b. Millikan c.
Rutherford
5. Compare and contrast the three types of subatomic particles in
terms of location in the atom, mass, and relative charge.
Answers:
1. An atom is the fundamental building block of nature. It is the
smallest component of an element that exhibits all the
characteristic properties of that element.
2. Dalton was wrong about atoms being indivisible (they are made
up of subatomic particles) and about all atoms of a given element
having identical properties (masses of isotopes differ).
3. Both believed: matter composed of extremely small particles
called atoms etc.
4. Thompson - provided strong support for the hypothesis that
cathode rays are negatively charged particles.
Millikan - work confirmed that an electron has the smallest
possible negative electric charge.
Rutherford - discovered the nucleus and described its contents.
5.
Particle Symbol
Location in the
Atom
Relative Charge
Relative Mass
Mass (g)
proton
p
nucleus
+1
1
1.673 × 10-24
neutron
n
nucleus
0
1
1.675 × 10-24
electrons
e-
electron cloud
-1
~0
9.110 × 10-28
How Atoms Differ
Atomic Weight (u): The average mass of an atom of an element,
usually expressed in atomic mass units. The terms mass and weight
are used interchangeably in this case. The atomic weight given on
the periodic table is a weighted average of isotopic masses found in
a typical terrestrial sample of the element.
 Atomic Number (Z) – is equal to number of Protons and
Electrons in the atom.
 # of Neutrons = Mass Number (MW) – Atomic Number.
Isotopes = Atoms of the same element having a different numbers
of neutrons due to different atomic masses. Same Atomic #!
Mass Number
3
1H
235
92U
Atomic Number
Ions = is an atom or group of atoms that has a positive or negative
charge. The atomic number still equals the number of protons,
but different number of electrons.
Na+1 = 11+ protons, 12 neutrons, & 10e' {-11 + (+1)} = -10e'
O-2 = 8+ protons, 8 neutrons, & 10e' {-8 + (-2)} = -10e'
Fe+3 = 26+ protons, 30 neutrons, & 23e' {-26 + (+3) = -23e'
Isotope Half Life
C-11
20.3 minutes
C-12
C-13
Stable
Stable
C-14
5730.0 years
How to Calculate
Average
Atomic Weight
C-15 an2.5
seconds
Example #1: Nitrogen
mass number exact weight percent abundance
14
14.003074
99.63
15
15.000108
0.37
(14.003074) (0.9963) + (15.000108) (0.0037) = 14.007
 Isotope Lab Activity
Chemistry Drill
Exercise 1: Complete the following chart.
Name
Sodium
Manganese
Lead
Symbol
Fe
Protons
Neutrons
29
4
Electrons
Atomic #
Mass #
Mercury
19
47
137
1. Which subatomic particle identifies an atom as that of a
particular element? How is this particle related to the atom's
atomic number?
2. What is an isotope? Give an example of an element with
isotopes.
3. Explain how the existences of isotopes are related to atomic
masses not being whole numbers.
Answers:
1. The Proton. The number of protons equals the atomic number.
2. Isotopes are atoms of the same element that have the same
number of protons but different number of neutrons. Carbon
12,13, & 14.
3. Atomic masses are not whole numbers because they represent
weighted averages of the masses of the isotopes of an element.
Quantum Theory of the Atom
The Periodic Table - Dimitri Mendeleev first invented and arranged
all the elements according to their unique properties.
1. Periods - are classified as the horizontal rows of the table.
2. Groups - are classified as the vertical rows of the table.
3. Chemical Symbols / Elements - are the shorthand ways to write
the name of the elements.
Rutherford's Nuclear Model
1. The atom contains a tiny dense center called the nucleus.
 the volume is about 1/10 trillionth the volume of the atom
2. The nucleus is essentially the entire mass of the atom.
3. The nucleus is positively charged & the amount of positive
charges of the nucleus balances the negative charge of the electrons.
4. The electrons move around in the empty space of the atom
surrounding the nucleus.
Thomson's Atomic Theory
1. An atom is breakable!
2. Atom has structure
3. Electrons suspended in a
positively charged electric
field.
 must have positive
charge to balance
negative charge of
electrons and make
the atom neutral
4. Mass of atom due to
electrons.
5. Atom mostly "empty"
space & compared size of
electron to size of atom
Niels Bohr Model
Atomic Structure: Size of the electron cloud – n values
n=1
n=2
n=3
n=4
all the way to n=7.
n = 1 contains a total of 2 electrons
n = 2 contains a total of 8 electrons
n = 3 contains a total of 18 electrons
n = 4 contains a total of 32 electrons (Max is 32)
Electron Configuration – The s, p, d, f orbitals. The s, p, d, f
orbitals contain the total amount of electrons that each shell
around the nucleus can hold.
 Atomic Sub-levels/Orbitals
s orbital contains 1 pair or 2 electrons.
p orbital contains 3 pairs or 6 electrons.
d orbital contains 5 pairs or 10 electrons.
f orbital contains 7 pairs or 14 electrons.
Writing the correct configuration:
Primary Quantum #
6
3p
Number of Electrons
(atomic number & orbital)
Sub-orbital
To predict a ground state electronic configuration:
 Aufbau principle - Lowest energy orbitals fill first
 Pauli exclusion principle - No 2 electrons can have the
same set of quantum numbers (maximum of 2 electrons
per orbital)
 Hund's rule - When filling degenerate orbitals preserve
the maximum multiplicity (maximum number of unpaired
electrons)
These rules often give the correct electron configuration for
an atom or ground state ion.
A guide to the order of orbital energies:
Order of increasing energy:
1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p 7s 5f
Practice Problems:
Al - Aluminum - Ca - Calcium 1s2 2s2 2p6 3s2 3p1 1s2 2s2 2p6 3s2 3p6 4s2
Ba - Barium - 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2
Sn-Tin- 1s2 2s 2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p2