Why should you study chemistry?
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How do we study chemistry?
Chemistry is an experimental science: base on the scientific method.
The scientific method is the combination is the combination of observation, experimentation, and the
formulation of laws, hypothesis, ans theories.
Observation: natural or experimental
Tentative explanation: hypthesis
A hypothesis is a tentative explanation of a natual
law. If a hypothesis survives testing by
experiments, it is often referred to as a theory.
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
Experiments to test predication of theory
Theory estiblished unless later experiments
or observation show inadequacies of model
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.
What are we studying in chemistry?
Chemistry is the science that studies the composition and
properties of matter.
Matter is anything that occupies space, displays a property known as
mass, and possesses inertia. Matter is the name scientists have
given to everything that you can touch, or see, or feel
Composition refers to the parts or components of a sample of matter
and their relative proportions.
Any characteristic that can be used to describe or identify matter is
called a property.
We classify matter (1) by its physical state as a solid, liquid, or gas, and
(2) by its chemical constitution as an element, compound, or
mixture.
Properties
Properties can be classified as
1) Intensive or extensive, depending on whether their value changes with the
size of the sample. Intensive properties, like temperature and melting point,
have values that do not depend on the amount of sample.
Extensive properties, like length and volume, have values that do depend on
the sample size.
2)
Physical or chemical.
~ Physical properties and physical change
A physical property, are those characteristics like temperature, color and
melting point, is one that a sample of matter displays without changing its
composition.
In physical change some of the physical properties of a sample may change,
but its composition remains unchanged.
E.g. Liquid water and ice (solid water) certainly different in many way but water
remains 11.9% hydrogen and 88.81% oxygen by mass.
~ Chemical properties and chemical change
A chemical property is the ability (or inability) of a sample of matter to
undergo a change in composition under stated condition.
In a chemical change (or chemical reaction), one of more kinds of
matter are converted to new kinds with different compositions.
E.g. the rust that occurs when a bicycle is left out in the rain is due to the
chemical combination of oxygen with iron to give the new substance iron
oxide. Rust is therefore a chemical properties of iron.
Table 1 Some Examples of Physical and Chemical Properties
Physical Properties
Temperature
Color
Melting point
Electrical conductivity
Chemical Properties
Amount
Odor
Solubility
Hardness
Rusting (of iron)
Combustion (of coal)
Tarnishing (of silver)
Hardening (of cement)
Chemistry in Action: cooking
• Cooking is chemistry in action, with the added benefit that
you can eat the results. As you follow the recipe below for
cheese making, make observations about the changes that
occur.
Quick review
• Problem 1: What is a physical change? List the physical changes
that occur when one makes cheese.
• Ans: Physical changes include:
1 Dissolving of salt in water
• 2 Evaporation of water from solution
• 3. Stirring to mix
• Problem 2: What is a chemical change? List the chemical
changes that occur when one makes cheese.
• Ans: changes include:
1 lactic acid bacteria convert the sugar in milk (lactose) to lactic
acid
• 2 special bacteria ferment the remaining lactose and produce
carbon dioxide bubbles in the cheese.
Classification of Matter
Matter
separate by
physical
process
Ionic
compound
Molecular
compound
no
yes
Substance
Mixture
chemically
decomposed
uniform
throughout
yes
may
be
Compound
no
Element
yes
Homogeneous
no
Heterogeneous
Classification of Matter by its chemical constitution
Matter
separate by
physical
process
no
Substance
A substance is a kind of
matter that cannot be
separated into other
kinds of matter by any
physical process.
yes
Mixture
A mixture can be separated
into its components by
appropriate physical means
such as filtration; distillation and
chromatography, etc.
Mixture
uniform
throughout
yes
Homogeneous
no
Heterogeneous
Homogeneous mixtures are uniform in composition and properties throughout
a given sample but variable from one sample to another. E.g. a solution of
sucrose (sugar) is uniformly sweet throughout a given sample, but the
sweetness of different sucrose solution samples may be different depend on the
proportion of sugar and water in the solutions.
In heterogeneous mixtures such as salad dressing the components separate
into distinct regions. The composition and physical properties vary from one
part of the mixture to another.
Substance
chemically
decomposed
Ionic
compound
Molecular
compound
yes
may
be
Compound
no
Element
An element is a substance
made up of only a single
type of atom.
Chemical compounds are substances, formed when atoms of two or more different
elements combine with one another. They retain their identity during physical
change, but they can be decomposed into its constituent elements by chemical
changes.
A molecule or formula unit (for ionic compound) is the smallest entity having the
same proportions of the constituent atoms as does the compound as a whole.
The composition and properties of an element or compound are uniform throughout
a given sample and from one sample to another.
State of matter
• Commonly, a given kind of matter exists in different physical
forms under different conditions. For example: water exists as ice
(solid), as liquid water and as steam (gaseous water).
• Solid: the form of matter characterized by rigidity; a solid is
relatively incompressible and has fixed shape and volume.
• Liquid: the form of matter that is a relatively incompressible fluid;
a liquid has a fixed volume but no fixed shape.
• Gas: the form of matter that is an easily compressible fluid; a given
quantity of gas will fit into a container of any size and shape.
The three form of matter-solid, liquid, and gar-are referred to as the
state of matter.
Chemistry and the Elements
Everything you see around you is formed from one or more of 115 presently
known elements. Only 92 of the 115 presently known elements occur naturally.
Chemical Symbol.
For simplicity, chemists referred to specific elements using one or twoletter abbreviations of the name of the element known as Chemical
symbols. The first letter is capitalized; E.g. carbon, C; neon Ne and silicon, Si.
Some elements have symbols based on their Latin name, E.g. Fe iron (ferrum),
Pb lead (plumbum), Na sodium (natrium for sodium carbonate), K potassium
(kalium for potassium carbonate). The symbol for tungsten, W, is based on the
German wolfram.
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
Measurement of Matter: SI (Metric) Units
~ is a decimal system. Quantities differing from the base unit by powers of
ten are noted by the use of prefixes.
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
Abbreviation
kg
m
K
mol
s
A
cd_______
Table 3 Some SI Prefixes
Multiple
Prefix__
Symbol__
1015
peta
P
1012
tera
T
109
giga
G
106
mega
M
103
kilo
k
102
hecto
h
10
deca
da
10-1
deci
d
10-2
centi
c
10-3
milli
m
You need to memorize the most
10-6
micro

common SI prefixes such as G,
10-9
nano
n
M, k, d, c, m, , n and p.
10-12
pico
p
10-15
femto
f_______
Fundamental Units
You are required to know all the conversion
Mass
~ describes the quantity of matter in an object. An object has a fixed mass
(m), which is independent of where or how the mass is measured (Stated more
scientifically, matter is anything that has mass). Mass is measured in SI units
by the kilogram.
1 kg =1,000 g = 1,000,000 mg = 1,000,000,000 g
Relation between mass and weight
Weight is the force of gravity on an object. It is directly proportional to mass.
W = gm
* Mass is a physical property that measures the amount of matter in an object,
whereas weight measures the pull of gravity on an object.
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
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 SI
temperature scale, called Kelvin scale, assigns a value of zero to the coldest
possible temperature, 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
Fahrenheit
32F
212F
Celsius
0C
100C
Kelvin
273 K
373 K
Celsius from Fahrenheit
(C) = 5/9[t (F) –32]
Kelvin from Celsius
T (K) = t (C) + 273.15
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.
Density
Density is the intensity properties that relates the mass of an object to its volume.
Density =
Mass (g)_______
Volume (mL or cm3)
Because most substances change in volume when heated or cooled, densities are
temperature-dependent. E.g. 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.
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
•
3
3
1 L
10
cm
=
=1
3
3
3
3
10 cm
10 cm
• Observe that units are treated in the same way as algebraic quantities. Note
too that the right-hand 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:
V = 125 cm3 x
1 L
103 cm3
= 125 x 10 -3 L = 0.125 L
converts cm3 to L
• 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.
• E.g. Convert 8.45 kg to milligrams
• 1 kg = 106mg
8.45kg =
106 mg
8.45kg x
= 8.45 x 106 mg
1 kg
• 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?
0.085L= 0.085Lx
1700mg = 1700mg
1000mL
= 85mL
1L
1g
1000mg
= 1.7g
d=
m
V
=
1.7 g
= 0.02 g/mL
85 mL
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.
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.
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.
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.
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.
In talking about the degree of uncertainty in a measurement, we use the
words accuracy and precision.
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.
Significant Figures
The rules for determining significant figures (sig. fig.).
1 Zeros in the middle of a numbers are significant figures. E.g. 4023 mL
has 4 significant figures.
2 Zeros at the beginning of a number are not significant; they act only to
locate the decimal point. E.g. 0.00206L
has 3 significant figures. The zeros to the left of 2 are not sig. fig.
3 Zeros at the end of a number and after decimal point are always
significant. E. g. 2.200 g
has 4 sig. fig.
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.
5 A useful rule of thumb to use for determining whether or not zeros are
significant figures is that zeros are not sig. fig. if the zeros disappear
when scientific notation is used.
E.g. 0.0197 = 1.97 x10-2 the zeros are not sig. fig.; 0.01090 = 1.090 x10-2
the first two zeros are not significant and the second two are.
Not significant
zero for cosmetic
purpose
Not significant: zeros
used only to locate
the decimal point
Significant:
all zeros between
nonzero numbers
0. 0 0 4 0 0 4 5 0 0
Significant:
all nonzero
integers
Significant: zeros at
the end of a number
and after decimal point
The number 0.004004500 has 7 sig. Fig.
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.
E.g. 14.79cm x 12.11cm x 5.05cm = 904cm
(4 sig.fig) (4 sig. fig) (3 sig. fig) (3 fig. sig)
If use scientific notion,
E.g.
(3.4 x 106)(4.2 x 103)
= (3.4)(4.2) x 10(6+3) = 14.28 x 109 = 1.4 x 1010
(to 2 significant figures)
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.
How to do calculations
• 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.
•
E.g.
3.18
•
+ 0.01315
•
3.19
• 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)
= (4.215 x 10-2) + (0.032 x 10-2) = 4.247 x 10-2
• 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.
Rounding Numbers
1) If the first digit you remove is less than 5, round down by drop it and
all following digits.
E.g. 5.664525  5.66 when rounded to three sig. fig.
2) If the first digit you remove is greater than 5 or 5 followed by nonzero,
round up by adding 1 to the digit on the left.
E.g. 5.664525  5.7 when rounded to two sig. fig.
3) What happens if there is a 5 or 5 following by zeros? There is an
arbitrary rule “round 5 to even”:
If the number before the 5 is odd, round up.
If the number before the 5 is even, let it be.
The justification for this is that in the course of a series of
many calculations, any rounding errors will be averaged out.
E.g. Round the following number to 2 significant figures:
(a) 2.35 x 102
(b) 2.45 x 102
(Answer: 2.4 x 102)
Summary of Key Concepts
Chemistry
uses
to be better
understand its
Matter
Experimentation
separate by
physical
process
Properties
involving
which
may be
no
Intensive or
Extensive
Measurement
Observation
yes
Physical or
Chemical
to generate
Substance
Mixture
may be
suggesting
futher
Hypotheses
leading to
maybe
Numbers
Theories
Compound
Element
Heterogeneous
Uncertainty
expressed using
may be
Ionic
compound
Homogeneous
Molecular
compound
which yields
Significant
figures
having
Acuracy
Units
using the
Precision
SI system