Post on 21-Aug-2018
Introduction: Matter &
Measurement
CHEMISTRY The Central Science
Cpt. 1 and 2
What is Chemistry?
The central science
The study of the matter, its composition, properties, and the changes it undergoes.
The Molecular Perspective of Chemistry
• Matter is the physical material of the universe. It
has mass and occupies space.
• On the microscopic level, matter consists of atoms
and molecules.
• Atoms combine to form molecules.
• Molecules may consist of the same type of atoms
or different types of atoms.
The Molecular Perspective of Chemistry
Why Study Chemistry?
• Chemistry is central to our understanding of other
sciences.
• Chemistry is also encountered in everyday life.
States of Matter
• Matter can be a gas, a liquid, or a solid.
• Gases have no fixed shape or volume.
• Gases can be compressed to form liquids.
• Liquids have no shape, but they do have a defined volume. Liquids are fluid.
• Solids are rigid and have a definite shape and volume.
The Three States of Matter
• Elements and compounds are referred to as
Pure Substances. They contain one type of
particle.
• Mixtures are not considered pure
substances. They contain more than one
type of particle.
Pure
Substances
and Mixtures
Elements
• If a pure substance cannot be decomposed into
something else, then the substance is an element.
• There are 118 elements known.
• Each element is given a unique chemical symbol
(one or two letters).
• Elements are building blocks of matter.
Elements
Elements
• Chemical symbols with one letter have that letter
capitalized (e.g., H, B, C, N, etc.)
• Chemical symbols with two letters have only the
first letter capitalized (e.g., He, Be).
C Cu Na U
Compounds
• Most elements interact to form compounds.
• Law of Constant Composition (or Law of
Definite Proportions):
–The composition of a pure compound is always
the same.
Substances
• If water is decomposed, then there will always be
twice as much hydrogen gas formed as oxygen
gas.
• Pure substances that can be decomposed are
compounds.
Mixtures
• Heterogeneous mixtures are not uniform
throughout.
• Homogeneous mixtures are uniform throughout.
• Homogeneous mixtures are called solutions.
Physical and Chemical Properties
• Intensive physical properties do not depend on
how much of the substance is present. – Examples: density, temperature, color, melting point.
• Extensive physical properties depend on the
amount of substance present. – Examples: mass, volume, pressure.
Chemical Properties describe the chemical
behavior of a substance
-Example: Flammable, corrodes metal, basic
Physical and Chemical Changes
• When a substance undergoes a physical change, its
physical appearance changes. – Ice melts: a solid is converted into a liquid.
• Physical changes do not result in a change of
composition.
• When a substance undergoes a chemical change, it
changes its composition: – When pure hydrogen and pure oxygen react completely, they
form pure water.
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.
• Homogeneous liquid mixtures can be separated by
distillation.
Distillation
Separation of Mixtures
• Chromatography can be used to separate mixtures
that have different polarity.
• Chromatography can be used to separate the
different colors of inks in a pen.
Gas Chromatography
The Periodic Table
Periodic Table Basics
• Most elements are metals – left of zig-zag
• Non metals are on the right of zig-zag
• Borderline elements called metalloids
• Horizontal rows are called periods
• Vertical columns are called groups or
families. (similar properties – more later)
Metals and Nonmetals and
Their Ions
• Metals
– Good conductors of heat and electricity.
– Malleable and ductile.
– Moderate to high melting points.
• Nonmetals
– Nonconductors of heat and electricity.
– Brittle solids.
– Some are gases at room temperature.
C-2 Measures and metrics
SI Units
• There are two types of units: – fundamental (or base) units;
– derived units.
• There are 7 base units in the SI system.
• Powers of ten are used for convenience with
smaller or larger units in the SI system.
SI Units
SI Units
• Note the SI unit for length is the meter (m)
whereas the SI unit for mass is the kilogram (kg). – 1 kg weighs 2.2046 lb.
Temperature
There are three temperature scales:
• 1. Kelvin Scale – Used in science.
– Same temperature increment as Celsius scale.
– Lowest temperature possible (absolute zero) is zero Kelvin.
– Absolute zero: 0 K = -273 oC.
Temperature
• 2. Celsius Scale – Also used in science.
– Water freezes at 0 oC and boils at 100 oC.
– To convert: K = oC + 273.
• 3. Fahrenheit Scale – Not generally used in science.
– Water freezes at 32 oF and boils at 212 oF.
– To convert:
32-F
9
5C 32C
5
9F
Temperature
Scientific Notation
•Numbers written in scientific notation include a
numeral with one digit before the decimal point,
multiplied by some power of ten (6.022 x 1023)
•In scientific notation, all digits are significant.
•You should be able to convert from non-scientific
notation to scientific and vice-versa (on calc. as
well).
Examples
1. Convert to scientific notation:
a. 450 000 000
b. 0.000 000 047
2. Convert to standard notation:
a. 7.09 x 10-6
b. 3.39 x 105
Derived Units
• Derived units are obtained from the 7 base SI
units.
• Example:
m/s
seconds
meters
timeof units
distance of units velocityof Units
Volume
• The units for volume are
given by (units of
length)3. – SI unit for volume is 1 m3.
• Common: 1 mL=1 cm3.
• Other volume units: 1 L = 1 dm3 = 1000 cm3 = 1000mL.
Volume
Density
• Used to characterize substances.
• Defined as mass divided by volume:
• Units: g/cm3 or g/mL
• Originally based on mass (density was defined as
the mass/vol. of 1.00 g of pure water).
volume
massDensity
Examples:
1. An object with a mass of 17.95 g occupies a
volume of 11.8 mL. What is its density?
2. A sample with a density of 3.75 g/cm3 occupies a
volume of 10.44 cm3. What is the mass of the
sample?
3. A graduated cylinder is filled with 15.0 cm3 of
water. An object with a mass of 29.66 g causes
the total volume to increase to 23.4 mL. What is
the density of the sample?
1. 1.52g/mL 39.2 g 3.53g/ml
Dimensional Analysis
• Method of calculation utilizing a knowledge of
units. Conversion factors are simple ratios
(fractions):
unitgiven
unit desired factor Conversion aswritten
Using a Conversion Factor
• Example: convert length in meters to length in
centimeters:
• 3.25 meters x 100 cm = 325 cm
1 m
Using Two or More Conversion Factors
• Example to convert length in meters to length in
inches:
• 3.00 meters x 100 cm x 1 inch = 118 in
1 m 2.54 cm
Problem solving
• In dimensional analysis always ask three
questions:
• What data are we given?
• What quantity do we need?
• What conversion factors are available to take us
from what we are given to what we need?
Uncertainty in Measurement
• All scientific measures are subject to error.
• These errors are reflected in the number of figures
reported for the measurement.
Precision and Accuracy
• Measurements that are close to the “correct” value
are accurate.
• Measurements that are close to each other are
precise.
Precision and Accuracy
Significant Figures
• The number of digits reported in a measurement
reflect the accuracy of the measurement and the
precision of the measuring device.
• All the figures known with certainty plus one extra
figure (estimated digit) are called significant
figures.
• In any calculation, the results are reported to the
fewest significant figures (for multiplication and
division) or fewest decimal places (addition and
subtraction).
Significant Figures Rules:
1. Non-zero numbers are always significant.
2. Zeros between digits are always significant.
3. Initial zeros are not significant. (Example:
0.0003 has only 1 sf.)
4. Zeros at the end of the number are significant IF
there is a decimal point.(24.0 cm has 3 sf)
Examples: How many significant figures are in
each of the following?
1. 52300 m
2. 0.000487 kg
3. 29.0400 s
4. 507 people
5. 230,050 cm
6. 45.600 A
Reporting Uncertainties and Errors
in Measurements
• All instruments have limitations. When
recording a measurement, you should report
the measurement to the limits of the
instrument.
• The uncertainty (or tolerance) of an
instrument may be found on the instrument
( given as ±) , or you may be told what it is.
• Percent Error – used to compare data
(observed values) to accepted values in a
meaningful way
• │ Obs – Acc│ x 100 = % error
Acc
0
2
4
6
8
10
1 2 3 4 5 6
Series2
Direct Inverse
Proportions Proportions
0
5
10
15
20
25
30
1 2 3 4 5 6 7
Series2