Transcript Chapter 1 ppt

Antoine Lavoisier,
1743-1794
Marie Curie,
1867-1934
Joseph Priestly,
1733-1804
Dmitri Mendeleev,
1834-1907
John Dalton,
1766-1844
What is Matter?
Matter: Anything that occupies space and has mass
Physical States of Matter
Gas: Indefinite volume, indefinite shape,
particles far away from each other
Liquid: Definite volume, indefinite shape,
particles closer together than in gas
Solid: Definite volume, definite shape,
particles close to each other
Properties of Matter
Property: Characteristic of a substance
Physical Properties: Properties of matter that can be observed
without changing the composition or identity of a substance
Example: Size, physical state
Chemical Properties: Properties that matter demonstrates when
attempts are made to change it into new substances, as a result of
chemical reactions
Example: Burning, rusting
Changes in Matter
Physical Changes: Changes matter undergoes without changing
composition
Example: Melting ice; crushing rock
Chemical Changes: Changes matter undergoes that involve
changes in composition; a conversion of reactants to products
Example: Burning match; fruit ripening
Classifying Matter
Pure substance: Matter that has only 1 component; constant
composition and fixed properties
Example: water, sugar
•Element: Pure substance consisting of only 1 kind of atom
(homoatomic molecule)
Example: O2
•Compound: Pure substance consisting of 2 or more kinds of
atoms (heteroatomic molecules)
Example: CO2
Mixture: A combination of 2 or more pure substances, with each
retaining its own identity; variable composition and variable properties
Example: sugar-water
•Homogenous matter: Matter that has the same properties
throughout the sample
•Heterogenous matter: Matter with properties that differ
throughout the sample
Solution: A homogenous mixture of 2 or more substances (sugarwater, air)
Classification of Matter
Measurement Systems
Measurement: Determination of dimensions, capacity, quantity or
extent of something; represented by both a number and a unit
Examples: Mass, length, volume, energy, density, specific gravity,
temperature
Mass vs. Weight
Mass: A measurement of the amount of matter in an object
Weight: A measurement of the gravitational force acting on an object
English System Units: Inch, foot, pound, quart
The Metric System of Units
Each type of measurement has a base unit.
• Other units are related to the base unit by a power of 10.
• The prefix of the unit name indicates if the unit is larger
or smaller than the base unit.
Unit of Length
Meter = basic unit of length, approximately 1 yard
1 meter = 1.09 yards
Kilometer = 1000 larger than a meter
Centimeter = 1/100 of a meter
100 cm = 1 meter
Millimeter = 1/1000 of a meter
1000 mm = 1 meter
Unit of Mass
Gram: basic unit of mass
454 grams = 1 pound
Kilogram: 1000 times larger than a gram
1 Kg = 2.2 pounds
Milligram: 1/1000 of a gram
Unit of Volume
Liter: basic unit of volume
1 Liter = 1.06 quarts
1 Liter = 10 cm x 10 cm x 10 cm
1 liter = 1000 cm3
1 ml = 1 cm3 (1 cc)
Units of Temperature
Fahrenheit: -459oF (absolute zero) - 212oF (water boils)
Celsius: -273oC (absolute zero) - 100oC (water boils)
Kelvin: 0K (absolute zero) - 373 K (water boils)
Different Temperature Scales
Converting Celsius and Fahrenheit:
oC
= 5/9 (Fo - 32)
oF
= 9/5 (oC) +32
Converting Celsius and Kelvin:
oC
= K - 273
K = oC + 273
Scientific Notation and Significant Figures
Scientific notation: a shorthand way of representing very small
or very large numbers
Examples: 3 x 102, 2.5 x 10-4
The exponent is the number of places the decimal must be
moved from its original position in the number to its position
when the number is written in scientific notation
If the exponent is positive, move the decimal to the right of the
standard position
Example: 4.50 x 102  450
3.72 x 105  372,000
If the exponent is negative, move the decimal to the left of the
standard position
Example: 9.2 x 10-3  .0092
Practice with Scientific Notation
50,000 = 5.0 x 104
300 = 3.00 x 102
.00045 = 4.5 x 10-4
.0005 = 5 x 10-4
Significant Figures
Significant Figures: Numbers in a measurement that reflect the
certainty of the measurement, plus one number representing an estimate
Example: 3.27cm
Rules for Determining Significance:
All nonzero digits are significant
Zeroes between significant digits are significant
Example: 205 has 3 significant digits
1,006 has
4 sig. figs.
10,004 has 5 sig. figs.
Leading zeroes are not significant
Example: 0.025 has 2 significant digits
0.000459 has 3 significant digits
0.0000003645
4 sig. figs.
Trailing zeroes are significant only if there is a decimal point in the
number
Examples: 1.00 has 3 significant figures
2.0 has 2 significant digits
20 has 1 sig. fig.
1500
2 sig. figs.
1.500
4 sig. figs.
Calculations and Significant Figures
Answers obtained by calculations cannot contain more certainty
(significant figures) than the least certain measurement used in
the calculation
Multiplication/Division: The answers from these calculations must
contain the same number of significant figures as the quantity with
the fewest significant figures used in the calculation
Example: 4.95 x 12.10 = 59.895
Round to how many sig. figs.?
Final answer:
59.9
3
Addition/Subtraction: The answers from these calculations must
contain the same number of places to the right of the decimal point
as the quantity in the calculation that has the fewest number of places
to the right of the decimal
Example: 1.9 + 18.65 = 20.55
How many sig. figs. required? 1
Final answer: 20.6
Rounding Off
Rounding off: a way reducing the number of significant digits to
follow the above rules
Rules of Rounding Off:
Determine the appropriate number of significant figures; any and
all digits after this one will be dropped.
If the number to be dropped is 5 or greater, all the nonsignificant
figures are dropped and the last significant figure is increased by 1
If the number to be dropped is less than 5, all nonsignificant
figures are dropped and the last significant figure remains
unchanged
Example: 4.287 (with the appropriate number of sig. figs.
determined to be 2)
4.287 4.3
We only use significant figures when dealing with inexact numbers
Exact (counted) numbers: numbers determined by definition or
counting
Example: 60 minutes per hour, 12 items = 1 dozen
Inexact (measured) numbers: numbers determined by
measurement, by using a measuring device
Example: height = 1.5 meters, time elapsed = 2 minutes
Practice:
Classify each of the following as an exact or a inexact number.
Inexact
A.
A field is 100 meters long.
B.
There are 12 inches in 1 foot.
C.
The current temperature is 20o Celsius.
D.
There are 6 hats in the closet.
Exact
Exact
Inexact
Calculating Percentages
percent = “per hundred”
% = (part/total) x 100
Example: 50 students in a class, 10 are left-handed. What
percentage of students are lefties?
% lefties = (# lefties/total students) x 100
= 10/50 x 100
= .2 x 100
= 20%
Problem Solving Using Conversion Factors
Conversion factor: A term that converts a quantity in
one unit to a quantity in another unit.
original
quantity
x
conversion factor
=
desired
quantity
•Conversion factors are usually written as equalities.
2.21 lb = 1 kg
•To use them, they must be written as fractions.
2.21 lb
1 kg
or
1 kg
2.21 lb
•Units are treated like numbers
•Make sure all unwanted units cancel
Example: Convert 130 lb into kilograms
130 lb x
original
quantity
conversion factor
=
? kg
desired
quantity
2.21 lb
1 kg
130 lb
x
or
=
59 kg
1 kg
2.21 lb
• The bottom conversion factor has the original unit in the
denominator and the desired unit in the numerator.
• The unwanted unit lb cancels.
• The desired unit kg does not cancel.
Example
How many grams of aspirin are in a 325-mg
tablet?
Step [1]
Identify the original quantity and the desired
quantity, including units.
original quantity
325 mg
desired quantity
?g
Step [2] Write out the conversion factor(s) needed
to solve the problem.
1 g = 1000 mg
This can be written as two possible fractions:
1000 mg
1g
or
1g
1000 mg
Choose this factor to
cancel the unwanted
unit, mg.
Step [3]
Set up and solve the problem.
325 mg
3 sig. figures
Step [4]
x
1g
1000 mg
=
Unwanted unit
cancels.
0.325 g
3 sig. figures
Write the answer with the correct number
of significant figures.
Practice Using and Converting Units in Calculations
Practice: Convert 125m to yards.
125 x 1.09 yards /1 = 136.25 yards
•Determine appropriate amount of sig. figs. and round accordingly
Fewest sig. figs. in original problem is 3 (from 125), so final answer is
136 yards
Solving a Problem Using Two or More
Conversion Factors
Example: How many liters is in 1.0 pint?
1.0 pint
x conversion factor =
original quantity
?L
desired quantity
•Two conversion factors are needed:
2 pints = 1 quart
2 pt
1 qt
or
1 qt
2 pt
First, cancel pt.
1.06 quarts = 1 liter
1.06 qt
1L
or
1L
1.06 qt
Then, cancel qt.
Set up the problem and solve:
1.0 pt
2 sig. figures
x
1 qt
2 pt
x
1L =
1.06 qt
0.47 L
2 sig. figures
Density and Specific Gravity
Density: A physical property that relates the mass of
a substance to its volume.
density
=
To convert volume (mL)
to mass (g):
g
mL x
= g
mL
density
mass (g)
volume (mL)
To convert mass (g)
to volume (mL):
mL
g
x
= mL
g
inverse of density
Solving Problems with Density
Example: If the density of acetic acid is 1.05 g/mL, what is
the volume of 5.0 grams of acetic acid?
5.0 g
original quantity
? mL
x conversion factor =
desired quantity
•Density is the conversion factor, and can be
written two ways:
1.05 g
1 mL
1 mL
1.05 g
Choose the inverse density
to cancel the unwanted unit, g.
Set up and solve the problem:
5.0 g
x
1 mL
1.05 g
=
4.8 mL
Unwanted unit
cancels.
Write the final answer with the correct number of significant figures.
Specific Gravity
Specific gravity: A quantity that compares the density
of a substance with the density of water at the same temperature.
specific gravity
=
density of a substance (g/mL)
density of water (g/mL)
•The units of the numerator (g/mL) cancel the
units of the denominator (g/mL).
•The specific gravity of a substance is equal to its
density, but contains no units.