Transcript Introduction

Chemistry
Introduction
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Definitions
Classification of Matter
Properties of Matter
Measurement and SI Units
Working with Numbers
Quit
Definitions
• Matter is anything that occupies space and
has mass.
• Chemistry is the study of matter and the
changes it undergoes
• A substance is matter that has a definite or
constant composition and distinct properties
Examples are water, silver, sugar, table salt,
etc.
Matter
Mixtures
Homogeneous
Mixtures
Separation by Physical
Methods
Heterogeneous
Mixtures
Pure
Substances
Compounds
Elements
Separation by Chemical Methods
Properties of Matter
Physical Property
Chemical Property
Extensive Property
Intensive Property
Physical Property
A physical property can be measured
and observed without changing the
composition of a substance.
Examples:
Boiling Point
Density
Conductivity
Chemical Property
A chemical property refers to the
ability of a substance to react with
other substances. In order to observe
this property a chemical change must
take place.
Examples:
Sugar ferments to form alcohol
Hydrogen burns in oxygen to create water.
Extensive Property
Measurable properties which depend
on the amount of substance present are
called extensive properties.
Examples:
Mass
Length
Volume
Intensive Property
Measurable or observable properties
which are independent of the amount
of substance present are called
intensive properties.
Examples:
Color
Density
Temperature
Measurement and SI Units
SI units are an international standard of units developed
in 1960 based on the decimal (base 10) system.
Base Quantity
Name of Unit
Symbol
Length
meter
m
Mass
Kilogram
kg
Time
second
s
Temperature
Amount of
Substance
kelvin
K
mole
mol
Length
Length measures the extent of an object.
Length can be used to determine derived units
such as area and volume.
Area = m x m = m2
Volume = m x m x m = m3
1 Liter (L) = 1dm3 (One cubic decimeter)
1 milliliter (mL) = 1cm3 (One cubic centimeter)
1 L = 1000 mL
Density d = m/V (mass per unit volume)
Mass
Mass is a measure of the quantity of matter inside of
a substance or object.
It should not be confused with the term weight,
which is a measure of the force that gravity exerts
on an object. They are related by the following
equation;
F = mg
where g is the acceleration due to gravity, m is the
mass and F is the force in Newtons
In chemistry, the smaller unit of mass grams (g) is
preferable to kilograms (kg). 1kg = 1000g
Temperature
Temperature measures the average kinetic energy of
the particles contained within a system or object.
Although Kelvin are the accepted SI unit, the Celsius
scale is often used. Both are based on the decimal
system. The Fahrenheit scale is seldom used for
scientific measurement.
Refer to the next frame for a comparison of
temperature scales and conversion factors.
Temperature Comparisons and
Conversions
373
100
212 Water Boils
oF
310
298
37
273
0
25
= 9/5 oC + 32
Body
o
o
98.6 Temperature C = ( F - 32)5/9
77 Room
K = oC + 273
Temperature
32 Water Freezes
Kelvin Celsius Fahrenheit
oC
oF
K
Working With Numbers
Scientific Notation
Significant Figures
Accuracy and Precision
Factor-Label Method of Solving
Problems
Scientific Notation
Allows representation of large or small numbers
accurately.
Removes possible ambiguity about significant figures.
Numbers are expressed follows;
N x 10n
where N is a number between 1 and 10 and n is an
integer exponent that is positive if the decimal point is
moved to the left to make N between 1 and 10, and
negative if it must be moved to the right.
Examples
1. The number 5,876.73 is expressed in scientific notation
as;
5.87673 x 103
2. The number .000034785 is expressed as;
3.4785 x 10-5
Addition and Subtraction
1. Write each number so that n has the same exponent
2. Add or subtract the N parts of the numbers
3. The exponent n remains the same
Example:
2.3x104 + 1.5x103 would be rewritten as 2.3x104 + .15x104
and the final answer would be 2.45 x 104.
Multiplication and Division
1. Multiply or divide the N parts of the numbers together
2. Add the exponents, n, if multiplying
3. Subtract exponents if dividing
Example:
3.0x103 x 4.0x104 = 12x107 = 1.2x108
Significant Figures
Significant figures refer to the meaningful digits in a
measured or calculated quantity
The last digit is understood to be uncertain when
significant figures are counted
Guidelines
• Any digit that is not zero is significant
•Zeros between nonzero digits are significant
•Zeros to the left of the first nonzero digit are not
significant
•If a number is greater than 1, then all the zeros
written to the right of the decimal point count as
significant figures
•For numbers that do not contain decimal points, the
trailing zeros (zeros after the last nonzero digit) may
or may not be significant. This is one reason why it is
important to use scientific notation
Calculations Involving Sig Figs
Addition and Subtraction: The number of digits to the right of
the decimal point in the final answer is determined by the lowest
number of significant figures to the right of the decimal in any of
the original numbers
Multiplication and Division: The number of significant figures in
the final answer is determined by the original number that has
the smallest number of significant figures
Exact numbers (from definitions or by counting) are considered
to have an infinite number of significant figures
For chain (multiple) calculations, carry the intermediate answers
to one extra decimal place and round the final answer to the
correct digits
Accuracy and Precision
Accuracy tells how close a measurement is to the true
value of the quantity that was measured
Precision refers to how closely two or more
measurements of the same quantity agree with one another
Precise and accurate
Precise but not
accurate
Neither precise nor
accurate
Dimensional Analysis
(Factor Label Method)
Allows accurate conversion between units of similar types
It utilizes the fact that equivalent quantities using different
units may be set up as a ratio to convert from one type of
unit to another
Algebraically, labels are treated exactly the same way as the
numbers they refer to
The unit you are converting to should always be placed in
the ratio such that the old units cancel out and the new unit
is in the desired position whether numerator or denominator
Examples
1in = 2.54cm therefore the ratio 1in/2.54cm or 2.54cm/1in may
by used to convert centimeters to inches or inches to
centimeters, respectively
100in x (2.54cm/1in) = 254cm
1km = 0.6215mi
10km x (0.6215mi/1km) = 6.215mi