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Chemistry
Chemistry is the science that studies the composition, structure and properties, and
interactions of matter
I. Brief History
Chemistry arose from attempts by people to transform metals into gold beginning about AD 100,
an effort that became known as alchemy Modern chemistry was established in the late 18th
century, as scientists began identifying and verifying through scientific experimentation the
elemental processes and interactions that create the gases, liquids, and solids that compose our
physical world. As the field of chemistry developed in the 19th and 20th centuries, chemists
learned how to create new substances that have many important applications in our lives.
Chemists, scientists who study chemistry, are more interested in the materials of which an object
is made than in its size, shape, or motion. Chemists ask questions such as what happens when
iron rusts, why iron rusts but tin does not, what happens when food is digested, why a solution of
salt conducts electricity but a solution of sugar does not, and why some chemical changes
proceed rapidly while others are slow. Chemists have learned to duplicate and produce large
quantities of many useful substances that occur in nature, and they have created substances
whose properties are unique.
Much of chemistry can be described as taking substances apart and putting the parts together
again in different ways. Using this approach, the chemical industry produces materials that are
vital to the industrialized world. Resources such as coal, petroleum, ores, plants, the sea, and the
air yield raw materials that are turned into metal alloys; detergents and dyes; paints, plastics, and
polymers; medicines and artificial implants; perfumes and flavors; fertilizers, herbicides, and
insecticides. Today, more synthetic detergent is used than soap; cotton and wool have been
displaced from many uses by artificial fibers; and wood, metal, and glass are often replaced by
plastics.
reference: http://encarta.msn.com/encyclopedia_762504460/Chemistry.html#s57
II. Chemistry as Quantitative Science
• The study of chemistry could involve both qualitative and quantitative aspect of materials
around us.
• The qualitative aspect involves describing materials such as color and the change they
under goes.
• The quantitative side concerned with measuring and calculating the characteristics of
materials. This quantitative aspect has played, and continues to play, an important role in
modern chemistry.
• Much of chemistry can be described as taking substances apart and putting the parts
together again in different ways. Using this approach, the chemical industry produces
materials that are vital to the industrialized world.
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Chemistry is an experimental science: base on the scientific method.
The scientific method is the use of carefully controlled experiments to answer scientific questions (the
combination is the combination of observation, experimentation, and the formulation of laws, hypothesis,
and theories.)
Observation (question): natural or experimental
Tentative explanation: hypthesis
Revise if experiments show
hypothesis is inadequate
Experimental designed to test hypothesis
Theory that amplifies hypothesis
and give prediction
Modify theory if
experiments show model is
inadequate
A hypothesis is a tentative explanation for an
observation or a phenomenon.. If a hypothesis
survives testing by experiments, it is often referred
to as a theory.
A theory is a model or way of lookingat nature that
can be used to explain natural laws and make further
predictions about nature phenomena.
An experiment is an observation of natural
phenomena carried out in a controlled manner so
that the result can be duplicated and rational
conclusions obtained.
A law is a concise statement or mathematical
equation about a basic relationship or
regularity of nature.
Experiments to test predication of theory
Theory estiblished unless later experiments
or observation show inadequacies of model
Quick-Review Question
III. Experimentation and Measure
Measurement is the comparison of a physical quantity to be measured with a unit of
measurement-that is a fixed standard of measurement. In 1960 the General Conference on
Weights and Measures adopted the International System of units (SI units) which is particular
choice of metric unit. This system has seven SI base units from which all others can be derived.
One of the advantages of any metric system is that it is a decimal system. Quantities differing
from the base unit by powers of ten are noted by the use of prefixes.
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Table 1 The seven fundamental units of measure
Physical quantity
Name of unit
Mass
kilogram
Length
meter
Temperature
Kelvin
Amount of substance
mole
Time
second
Electric current
ampere
Luminous intensity
candela
Table 2 Some SI Prefixes
Multiple
Prefix__
15
10
peta
1012
tera
9
10
giga
106
mega
103
kilo
2
10
hecto
10
deca
-1
10
deci
10-2
centi
-3
10
milli
10-6
micro
-9
10
nano
10-12
pico
-15
10
femto
Abbreviation
kg
m
K
mol
s
A
cd_______
Symbol__
P
T
G
M
k
h
da
d
c
m

n
p
f_______
Mass
Relation between mass and weight
Weight is the force of gravity on an object. It is directly proportional to mass.
Length
The meter (m) is the standard unit of length in the SI system.
1 m = 100 cm = 1,000 mm = 1,000,000 m
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Temperature
To establish a temperature scale, we arbitrarily set certain fixed points and temperature
increments called degree. Two commonly used fixed points are the temperature at which ice
melt and the temperature at which water boils. The Celsius scale is the temperature scale in
general scientific use. On this scale, the freezing point of water is 0C and the boiling point of
water at normal barometer pressure is 100C. However, the SI base unit of temperature is the
Kelvin (K), a unit on an absolute temperature scale. On this scale, the value zero is the lowest
possible temperature can be obtained theoretically which is 273.15C, sometime called absolute
zero.
0K = 273.15C.
A comparison of the Kelvin, Celsius, and Fahrenheit temperature scale.
Melting Point
Boiling point of water
E.g. The hottest place on record in North America is in Death Valley, California. A temperature
of 134F was reached there in 1913. What is this temperature reading in degree Celsius? in
Kelvin?
Derived Units
Area (m2), volume (m3), density (kg/m3), speed (m/s) etc are some of the derived quantities that
you are familiar with. They are derived quantities rather than fundamental quantities because
they can be expressed using one or more of the seven base units.
Volume, the amount of space occupied by an object, is measured in SI units by the cubic meter
(m3), defined as the amount of space occupied by a cubic 1 meter long on each edge.
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Density
Density is the intensity property that relates the mass of an object to its volume.
Because most substances change in volume when heated or cooled, densities are temperaturedependent for example, the density of water at 3.98C is 1.000g /mL and at 100C is 0.9584
g/mL as the volume expand. Although most substances expand when heated and contract when
cooled, water behaves differently. Water contract when cooled from 3.98C to 0C.
E.g. 2) What is the density of glass if a sample weighing 62.0 g has a volume of 12.4 cm3?
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IV. Working With Units: Dimensional Analysis
Dimensional analysis is the method of calculation in which one carries along the units for
quantities. Suppose we want to find the volume (V) of a cube, give l, the length of a side of the
cube. Because V = l3, if l = 5.00cm, we find that V = (5cm)3 = 125cm3. There is no guesswork
about the unit of volume here; it is cubic centimeter (cm3). Suppose, however, that we wish to
express the volume in liter (L), a metric unit that equals 103 cubic centimeters. We can write this
equality as
1L = 103 cm3.
If we divide both sides of the equality by the right-hand quantity, we get
Observe that units are treated in the same way as algebraic quantities. Note too that the righthand side now equals 1 and there are no units associated with it. Because it is always possible to
multiply any quantity by 1 without changing that quantity, we can multiply our previous
expression for volume by the factor 1 L/103 without changing the actual volume. We are
changing only the way in which we express this volume:
The ratio 1 L/ 103 cm3 is called a conversion factor because it is a factor equal to 1 that converts
a quantity expressed in one unit to one expressed in another unit. Note that the numbers in this
conversion factor are exact, because one liter equals exactly one thousand cubic centimeter.
Such exact conversion factors do not affect the number of significant figures in an arithmetic
result.
E.g. Convert 8.45 kg to milligrams
E.g. What is the density of a substance in g/mL, if a sample with a volume of 0.085 liters has a
mass of 1700 mg?
V. Uncertainties in Scientific Measurement
All measurements are subject to error.
Types of Errors:
1) Systematic error: arise because to some extent, measuring instruments have built-in, or
inherent, errors.
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2) Random errors: they arise from intrinsic limitation in the sensitivity of the instrument and
inability of observer to read a scientific instrument and give results that may be either too
high or too low.
Accuracy and Precision
In talking about the degree of uncertainty in a measurement, we use the words accuracy and
precision.
Accuracy refers to how close to the true value a given measurement is.
Precision refers to how well a number of independent measurements agree with one another.
There is no relationship between accuracy and precision, since an experiment can have small
random errors and still give inaccurate results due to large systematic errors.
E.g.
E.g. Imagine that you weigh a tennis ball whose true mass is 54.44178 g. Assume that you take
three independent measurements on each of three different types of balance to obtain data shown
in the following table
Measurement #
Bathroom scale1 Platform balance2 Analytical balance 3
1
0.0 kg
54.4g
54.4419g
2
0.0 kg
54.7g
54.4417g
3
0.1 kg
54.1g
54.4417g
average
0.03 kg
54.4g
54.4418g
VI. Significant Figures
To indicate the precision of a measured number (or result of a calculation with measured
numbers), we often use the concept of significant figures. Significant figures are those digits in
a measured number (or result of a calculation with measured numbers) that include all certain
digits plus a final one having some uncertainty. When we measured the rod, we obtained the
values 9.12 cm, 9.11 cm and 9.13 cm. We report the result as the average 9.12 cm. The first two
digits (9.1) are certain, and the next digit (2) is estimated, so has some uncertainty.
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The rules for determining significant figures (sig. fig.).
Number of significant figures refers to the number of digits reported for the value of a
measurement or calculated quantity, indicating the precision of the value. To count the number
of sig. fig. in a given measured quantity, we observed the following rules:
1 Zeros in the middle of a numbers are significant figures. E.g. 4023 mL
2 Zeros at the beginning of a number are not significant; they act only to locate the decimal point.
E.g. 0.00206L
3 Zeros at the end of a number and after decimal point are always significant. E. g. 2.200 g
4 Zeros at the end of number and before the decimal point may or may not be significant figures.
In such cases we must deduce the number of sig. fig. from the statement of the problem. E.g. The
statement “ 350,000 spectators lined the parade route” involves a number that probably has only
two sig. fig. because it is obvious that no one actually counted the spectators.
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that zeros are not sig. fig. if the zeros disappear when scientific notation is used.
Scientific Notation
• Scientific Notation was developed in order to easily represent numbers that are either
very large or very small.
• For example, instead of writing 0.0000000056, we write 5.6 x 10-9. So, how does this
work?
• We can think of 5.6 x 10-9 as the product of two numbers: 5.6 (the digit term) and 10-9
(the exponential term).
• A positive exponent shows that the decimal point is shifted that number of places to the
right. A negative exponent shows that the decimal point is shifted that number of places
to the left.
• In scientific notation, the digit term indicates the number of significant figures in the
number. The exponential term only places the decimal point.
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E.g.1) 46600000 = 4.66 x 10
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E.g.2) 0.00053 = 5.3 x 10-4
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E.g.(3) Write in scientific notation: 0.000467 and 32000000
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E.g.(4) Express 5.43 x 10-3 as a number.
Significant Figures in Numerical Calculations
1 Multiplication or Division
The result of multiplicand or division, may contain only as many sig. fig. as the least precisely
known quantity in the calculation.
If use scientific notion,
the digit terms are multiplied in the normal way and the exponents are added. The end
result is changed so that there is only one nonzero digit to the left of the decimal.
E.g.
(3.4 x 106)(4.2 x 103)
The digit terms are divided in the normal way and the exponents are subtracted. The
quotient is changed (if necessary) so that there is only one nonzero digit to the left of the
decimal.
E.g. (6.4 x 106)/(8.9 x 102)
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2 Addition and Subtraction:
The result of addition or subtraction must be expressed with the same digits beyond the
decimal point as the quantity carries the smallest number of such digits.
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When use scientific notion, all numbers are converted to the same power of 10, and the
digit terms are added or subtracted.
E.g. (4.215 x 10-2) + (3.2 x 10-4)
3 Exact numbers can be considered to have an unlimited number of sig. fig. There are
two situation when a quantity appearing in a calculation may be exact.
By definition such as 1 in = 2.54 cm
Counting such as six face on a cube or tow hydrogen atoms in a water molecule.
Powers of Exponentials:
The digit term is raised to the indicated power and the exponent is multiplied by the
number that indicates the power.
E.g. (2.4 x 104)3
Roots of Exponentials:
Change the exponent if necessary so that the number is divisible by the root. Remember
that taking the square root is the same as raising the number to the one-half power.
E.g.
VII. Rounding Numbers