Download CHEM 101

Professor Abdul Muttaleb Jaber
Chemistry Department
Office: Room # 261F
Tel: 2611
E-mail: [email protected]
Office hours: S 9-11 am
U 10-11 am
M 9-11am
Chapter 1
Chemical Foundations
An overview
The scientific method
Units of measurements
Uncertainty in measurements
Significant figures and
calculations
Dimensional analysis
Temperature
Density
Classification of matter
1.1 Chemistry: an overview
Matter is composed of atoms
Atoms are found as individuals or
molecules
Atoms and molecules are connected by
electrons
The challenge of chemistry is to think
of the material the atomic level
100 different types of atoms form all
substances in the world
Matter is composed of various types of
atoms or molecules.
Water is composed of O and H; H2O
An electric spark causes a mixture of O2
and H2 to explode forming H2O.
One substance changes to another by
reorganizing the way atoms attached to
each other
Scientific method
It is a way of solving problems
It consists of the following steps:
– Observation- what is seen or measured
– Hypothesis- guess of why things behave
the way they do. (possible explanation
for an observation)
– Experiment- designed to test hypothesis
These steps would lead to new observations, and
the cycle goes on
Once a set of hypotheses agree with
observations, they are grouped into a theory
Scientific method
Thery is a set of tested hypothesis
that gives an overall explanation for
a natural phenomenon
Laws are summaries of observations
Often mathematical relationship
1.3 Units of measurements
Every measurement has two parts
Number
Scale (called a unit)
SI system (le Systeme International in
French) based on the metric system
Examples:
Prefix
20 grams
20 k g = 20 X103 g
20 m g = 20 X10-3 g
6.63   Joule seconds
Metric System
Fundamental SI base Units
Mass - kilogram (kg)
Length- meter (m)
Time - second (s)
Temperature- Kelvin (K)
Electric current- ampere (amp, A)
Amount of substance- mole (mol)
Prefixes used in SI units
giga-
G
1,000,000,000 109
106
mega - M
1,000,000
kilo decicentimillimicronano-
1,000
103
0.1
10-1
0.01
10-2
0.001
10-3
0.000001
10-6
0.000000001 10-9
k
d
c
m
m
n
Volume measurement: Liter
Liter is defined as the volume of 1 dm3
–1 dm3 =
– (10cm)3 =
– 1000 cm3 =
– 1000mL
Graduated
Cylinder
Pipet
Buret
Volumetric
Flask
Mass and Weight
Mass is measure of resistance to
change in motion
Weight is force of gravity.
Sometimes used interchangeably
Mass can’t change, weight can
Electronic
Analytical
Balance
Errors in Measurement
All scientific measurements are subject to
error.
These errors are reflected in the
observation that two successive
measurements of the same quantity are
different.
Types of Errors
Random Error (Indeterminate Error) measurement has an equal probability of
being high or low.
Systematic Error (Determinate Error) - Occurs
in the same direction each time (high or low),
often resulting from poor technique of
measurement or bad equipment.
You can have precision without accuracy
You can’t have accuracy without precision
(unless you’re really lucky).
Uncertainty in Measurements
Uncertainty in Measurement
A measurement always has
some degree of uncertainty.
Uncertainty has to be indicated in any
measurement.
Any measurement has certain digits and
one uncertain digit.
A digit that must be estimated is
called uncertain.
The number of certain digits + the
uncertain digit is called number of
significant figures.
Precision and Accuracy
Accuracy: Agreement of a particular
value to the true value (degree of
correctness)
– Measurements that are close to the
“correct” value are accurate.
Precision: The degree of agreement
among several measurements of the
same quantity (degree of
repeatability).
– Measurements that are close to
each other are precise.
Precision and Accuracy
1.5 Significant figures and calculations
•The number of digits reported in a measurement reflect the
accuracy of the measurement and the precision of the measuring
device.
# Sig figs. =?
# Sig figs. =?
•All the figures known with certainty plus one extra figure are
called significant figures.
Significant figures in computation
• For multiplication and division
• the results are reported to the least number of significant
figures
354.760 X 0.0004567 = 0.162018892
= 0.1620
• For addition and subtraction
• The results are reported to the least number of decimal places
345.672 – 34.56720 = 311.1048
= 311.105
Rules for Counting Significant Figures
Nonzero integers always count as significant
figures.
3456 has
4 sig figs.
Zeros
Leading zeros do not count as significant figures.
(Zeros before the nonzero digit)
0.0486 has
3 sig figs.
0.0003 has
one significant figure
Captive zeros always count as significant figures.
16.07 has
4 sig figs.
Zeros
Trailing zeros are significant only
if the number contains a decimal point.
9.300 has
4 sig figs.
Exact numbers have an infinite
number of significant figures.
1 inch = 2.54 cm, exactly
• Zeros at the end of a number before a
decimal place are ambiguous
10,300 g has: 3 or 4 or 5??
Scientific Notation
•Addition and Subtraction
(6.6 x
10-8)
(3.42 x
+ (4.0 x
-5
10 )
10-9)
– (2.5 x
=
-6
10 )
7 x 10-8
= 3.17 x 10-5
(Note that these answers have been
expressed in standard form)
Rules for Rounding Off
To get the correct number of significant digits
you need to round numbers
In a series of calculations get one extra digit then
round
If the digit to be removed
is less than 5, the preceding digit stays the same
is equal to or greater than 5, the preceding digit is
increased by 1
Don’t forget to add place-holding zeros if
necessary to keep value the same!!
13
Multiple computations
2.54 X 0.0028
0.0105 X 0.060
1) 11.3
2) 11
=
3) 0.041
Continuous calculator operation =
2.54 x 0.0028
 0.0105  0.060
 Here, the mathematical operation requires
that we apply the addition/ subtraction rule
first, then apply the multiplication/division
rule.
= 12
1.6 Dimensional Analysis
Using the units to solve
problems
Dimensional Analysis
Method of calculation utilizing a knowledge of
units.
Conversion factors are used to manipulate units:
The conversion factors are simple ratios:
m
100 cm
Conversion of m to cm :
;
100 cm
m
How many minutes are in 2.5 hours?
Initial unit
2.5 hr
Conversion
factor
2.5 hr x 60 min
1 hr
Final
unit
= 150 min
How many seconds are in 1.4 days?
Unit plan: days
hr
1.4 days x 24 hr
1 day
x
min
seconds
??
Exact numbers
1.4 day x 24 hr x 60 min x 60 sec
1 day
1 hr
1 min
= 1.2 x 105 sec
Multiple units
The speed limit is 65 mi/hr. What is
this in m/s?
– 1 mile = 1760 yds
– 1 meter = 1.094 yds
65 mi
hr
1760 yd
1m
1 hr 1 min
1 mi
1.094 yd 60 min 60 s
If you are running at a speed of 65 meters per
minute, how many seconds will it take for you to
walk a distance of 8450 feet?
Initial
8450 ft
x 12 in.
1 ft
x 1 min
65 m
x 2.54 cm
1 in.
x 60 sec
1 min
x 1m
100 cm
= 2400 sec
Units to a Power
How many m3 is 1500 cm3?
1500 cm3
1m
100 cm
3
1.7 Temperature
Define the three temperature scales:
Celsius , Fahrenheit and Kelvin
Perform conversion from one to
another.
Units of Temperature between
Boiling and Freezing
Fahrenheit
Water boils
212°F
180°
Water freezes 32°F
Celsius Kelvin
100°C
100°C
0°C
373 K
100K
273 K
K = oC + 273
5
C  F - 32
9
9
F  C   32
5
1.8 Density
Density is the mass of substance
per unit volume of the substance:
density =
mass
volume
Densities of Various Common Substances* at 20° C
Density Problem
An empty container weighs 121.3 g. When filled
with a liquid (density 1.53 g/cm3 ) the container
weighs 283.2 g. What is the volume of the
container?
Mass of the liquid  161.9 g
V  161.9 g
3
1cm
3
V  161.9 g 
 106cm
1.53g
1.9 Classification of Matter
Matter: Anything occupying space and having
mass.
Three States of Matter:
Solid: rigid - fixed volume and shape
Liquid: definite volume but assumes the
shape of its container
Gas: no fixed volume or shape - assumes
the shape of its container
Types of Mixtures
Matter could be pure (one component
only) or mixture
Mixtures have variable composition
(more than one component)
A homogeneous mixture has visibly
indistinguishable components usually
called a solution (for example, tap water)
A heterogeneous mixture has visibly
distinguishable components, clearly not
uniform (for example milk)
Organization of Matter
Physical and Chemical Changes
A physical change, is associated with a
change in the physical appearance but
not in chemical composition
– Ice melts: a solid is converted into a liquid.
When a substance changes its chemical
composition, it undergoes a chemical
change:
– H 2 + O2
pure water.
– In the flask containing water, there is no
oxygen or hydrogen left over.
Properties of Matter
Physical and Chemical Changes
Separation of Mixtures
Mixtures can be separated if their physical
properties are different.
Solids can be separated from liquids by
means of filtration.
The solid is collected in filter paper, and
the solution, called the filtrate, passes
through the filter paper and is collected in
a flask.
Filtration
Separation of Mixtures
Homogeneous liquid mixtures can be
separated by distillation.
Distillation requires the different liquids to
have different boiling points.
Thus, each component of the mixture is
boiled and collected.
The lowest boiling fraction is collected
first.
Distillation
Distillation is a physical change: No chemical
change occurs when salt water is distilled.
Separation of Mixtures
Chromatography can be used to separate
mixtures that have different abilities to
adhere to solid surfaces.
The greater the affinity the component
has for the surface (paper) the slower it
moves.
The greater affinity the component has
for the liquid, the faster it moves.
Chromatography can be used to separate
the different colors of inks in a pen.
Paper chromatography:
A Line of the mixture to
be separated is placed at
one end of a sheet
The Paper Acts
as a Wick to
Draw up the
Liquid
Component with
the weakest
attraction for the
paper travels
faster
Compounds and elements
Compound: A substance with a
constant composition that can be
broken down into elements by
chemical processes.
Element: A substance that
cannot be decomposed into
simpler substances by chemical
means.