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BASICS OF CHEMISTRY
A. Chemistry is the study of the structure, properties, and composition
of substances, and the changes that substances undergo.
1. Definition of substance
A substance is a sample of matter having a uniform and definite composition
2. Definition of matter
Matter is anything that has mass and that takes up space.
B. Branches of chemistry
1. Analytical chemistry
Composition (qualitative)
and/or
Amounts of substances (quantitative)
2. Biochemistry
The chemistry of living organisms
3. Inorganic chemistry
The chemistry of substances not containing carbon
4. Organic chemistry
The chemistry of substances containing carbon
5. Physical chemistry
The theoretical basis of the chemical basis of the chemical
behavior of substances
C. Scientific method
1. Observation
Information obtained through the senses
Starts the search for the cause of the observed effect
2. Hypothesis
a. A tentative explanation for an observation
b. Links a cause to the observed effect
c. Generates one or more definite predictions
If-then: If one does this, then that will result
d. Is testable by making observations
Collecting data
astronomy
Making measurements
biology
Performing experiments
chemistry
e. Is measurable
Can be tested by making numerical measurements.
f. Is falsifiable
3. Experiment
a. A carefully controlled, repeatable procedure for
gathering data to test a hypothesis
b. Changes only one variable (the independent variable)
and measures the impact that this change has on
another variable (the dependent variable)
c. Produces quantifiable results
d. Is repeatable
e. Involves multiple trials
f. Controls variables
(1) A factor that can vary or change, and that can
affect the result of the experiment
(2) Independent variable
(a) The variable that is changed by the
researcher
(b) Is what is assumed to be the cause of
the observed effect
(c) Should be only one independent variable
(3) Dependent variable
(a) The variable that changes in response to
the change the researcher makes to the
independent variable
(b) Is what is assumed to be the effect
caused by changes in the independent
variable
(c) The new value of the dependent variable
is caused by, and depends on, the value
of the independent variable.
(4) Extraneous variables
(a) Variables other than the independent
variable that may influence the
dependent variable.
(b) When the researcher fails to control all
extraneous variables they will produces
effects that cannot be separated from the
effects of the independent variable the
experiment will be unrepeatable and the
data will be inconsistent.
(c) Extraneous variables must be controlled.
(5) Controlled variables
(a) Quantities that the researcher wants to
remain constant
(b) Quantities that the researcher must
observe as carefully as the dependent
variables
(c) The researcher expects that the
experiment will test only the effect of
changes in one variable (the independent
variable) all other factors must be
removed or otherwise accounted for.
(d) Two ways to control extraneous
variables
[1] Literally control them
[2] Use a control group
4. Analyze Results
a. Hypothesis is consistent with all observations?
Move on to drawing a conclusion
b. Hypothesis is not consistent with all observations?
(1) Determine how to modify the original
hypothesis and begin the process again.
(2) Or decide to reject the hypothesis and create
a new one
5. Draw your conclusions
a. State a law
A law summarizes all of the many measurements in a mathematical equation.
b. Propose a model
A model is a representation of a real structure, event, or class of events intended to facilitate a better understanding of abstract concepts or to
allow computer simulation of real world events.
c. Provide a theory
(1) Is built on much accumulated data.
(2) Purpose is to provide a framework for
understanding and explaining
(3) Falsifiable
(4) Refinable
D. Three of the basic laws of chemistry
1. The law of conservation of mass
Lavoisier 1789 ( the first to carry out quantitative
chemical experiments
The total mass remains the same during a chemical change, that is, the total mass of the reacting substances is equal to the total mass of the products formed.
Example:
2.73 g HgO heated strongly(2.53 g Hg + ? O
red-orange silver fluid gas
Answer: 0.20 g O
2. The law of constant composition also known as the law of
definite proportions
Proust 1808
A compound always contains the same elements in the
same proportion by mass, that is, the relative amounts of each element in a particular compound is always the same regardless of the source of the compound, how it is prepared, or the amount of the compound analyzed.
Examples:
Neither the amount analyzed nor the source of the compound makes a difference:
100.0 g MgS 43.13 g Mg = 43.13% Mg
source 1 56.87 g S = 56.87% S
10.00 g MgS 4.313 g Mg = 43.13% Mg
source 2 5.687 g S = 56.87% S
In both cases the composition is exactly the same.
How the compound is prepared makes no difference in its composition:
1000.0 g Mg + 56.87 g S ( 100.0 g MgS
(43.13% Mg + 56.87% S)
43.13 g Mg + 1000.0 g S(100.0 g MgS
(43.13% Mg + 56.87% S)
In both cases the composition is the same.
3. The law of multiple proportions
Dalton circa 1800
When two elements form more than one compound, the
different masses of one element that combine with the same mass of the other element are in the ratio of small whole numbers.
Example:
compound mass of H mass of O
H2O 2 g 16 g
H2O2 2 g 32 g
32 g O=216 g O1
INTRODUCTION TO MATTER
A. States of matter
1. Rigid state ( solid
Definite/fixed volume
Definite/fixed shape
2. Fluid states of matter
Fluid ( a substance that can flow
a. Incompressible state ( liquid
Definite/fixed volume
No specific shape ( takes the shape of its container
b. Compressible state ( gas
No specific volume ( fills its container
No specific shape ( takes the shape of its container
B. Changes in matter
1. Physical change
a. Definition
Processes that change a substances form or physical appearance but not its chemical identity or chemical composition
b. Examples:
Change of state
Dissolving
2. Chemical change
a. Definition
Processes that change one substance into a chemically different substance
b. Example:
Chemical reactions
C. Properties of matter
1. Intensive and extensive properties
a. Intensive properties.
(1) Definition
An intensive property does not depend on the amount of the sample.
(2) Examples
Temperature
Color
Melting or boiling point
b. Extensive properties
(1) Definition
An extensive property does depend on the amount of the sample.
(2) Examples
Heat
Mass
Volume
c. In some cases it may be useful to convert an extensive
property into an intensive property.
(1) Purpose
To convert a property that does depend on the amount of the sample into one that does not depend on the sample.
This allows the new intensive property to be used to identify a substance.
This also allows the new intensive property to be used in calculations with a wide range of amounts of the substance.
(2) Procedure
This can be done by dividing one extensive property by another extensive property, such as volume, mass, or moles.
(3) Examples
Density
=mass (extensive)volume (extensive)
= density (intensive)
Specific heat capacity
=heat capacity (extensive)mass(extensive)
= specific heat capacity
(intensive)
Molar mass
=mass (extensive)moles (extensive)
= molar mass (intensive)
1. Physical properties
a. Definition
Properties of a substance that can be observed or measured without changing the substances chemical identity or chemical composition
b. Examples:
physical state, density, melting point, solubility, electrical conductivity, magnetic
properties, etc.
2. Chemical properties
a. Definition
Properties that describe how a substance undergoes chemical reactions and forms new substances
b. Example
Flammability ( the tendency of a substance to burn in the presence of oxygen
D. Substances
1. Definition
A sample of matter having a uniform and definite composition
2. Key points
a. A substance cannot be separated into other kinds of
matter by any physical process.
b. A substance always has the same characteristic
properties regardless of its source.
3. There are two kinds of substances: elements and compounds
a. Elements
An element is a substance that cannot be separated
into simpler substances by chemical reactions
b. Compounds
A compound is a substance that can be separated into simpler substances (elements or other compounds) by chemical reactions.
E. Mixtures
1. Definition
A physical, but not chemical, combination of two or more substances
2. Key points
a. A mixture can be separated into two or more substances
by physical processes.
b. A mixture has a variable composition.
c. Most materials found in nature are mixtures.
d. In a mixture each substance retains its own chemical
identity and properties.
3. There are two kinds of mixtures: homogeneous and
heterogeneous
a. Homogeneous mixture
A mixture that is completely uniform in composition with its components evenly
distributed throughout the sample ( also
called a solution
b. Heterogeneous mixture
(1) Definitions
(a) A heterogeneous mixture is a mixture
that is not uniform in composition and
that has an uneven distribution of
components in two or more phases.
(b) A phase is any part of a system with
uniform composition and properties.
(2) Examples
(a) Italian salad dressing
Oil phase
Vinegar phase
(b) Mixture of sand and salt
Sand phase
Salt phase
4. Methods of separating mixtures
(involve physical methods taking advantage of differing physical properties)
a. Dissolving
Examples:
Sand and salt
Caffeine and coffee
b. Distillation
Examples:
Seawater solution
Alcohol and water solution
c. Fractional crystallization
Example:
An impure mixture of mostly A
(less soluble) with a small amount
of B (more soluble)
d. Chromatography
(used to separate liquid or gaseous solutions)
(1) Paper chromatography ( filter paper
(2) Thin layer chromatography ( thin layer of
silica gel on a glass plate
(3) Column chromatography ( replaced by high
pressure liquid chromatography
(4) Gas chromatography ( microliters of fluid
separated in a glass tube then passed through
a detector;
Example: blood test for drugs
THE INTERNATIONAL SYSTEM (SI) OF MEASUREMENT
A. Key points
1. Also called the metric system
2. Based on powers of ten
3. Uses a prefix plus a unit
B. Prefixes see Table 1.2 on page 9 or the handout SI System Prefixes
All the ones in blue plus giga and pico.
C. SI base units see Table 1.1 on page 8 or the handout SI Base Units
D. SI derived units
1. Volume
Liter symbolized L
Equal to one cubic decimeter
2. Density
g/cm3
d = m/V or ( = m/V
UNCERTAINTY IN MEASUREMENT
A. Precision and accuracy
1. Precision
a. Definition
Precision is defined as how closely individual measurements agree with each other.
b. Description
Precision has to do with reproducibility of measurements ( how close to each other a group
of measurements fall.
c. The phrase limits of precision refers to how minutely
a quantity can be measured by an instrument until there
is no agreement among sequential measurements of the
same quantity ( customarily, the first decimal place
which must be estimated.
2. Accuracy
a. Definition
Accuracy is defined as how closely the mean of a set of measurements agrees with the correct or true value.
b. Description
Accuracy has to do with reliability of measurements
( how close the measurements, as a group, comes to
the true value.
3. Comparing accuracy and precision
spread all overneither precise nor accurateclose together but far from bulls-eyeprecise
but not accuratespread out but averages in the bulls-eyeaccurate
but not preciseclose together in the bulls-eyeboth precise
and accurate
4. Types of errors
a. Random error
(1) Definition
An error that has an equal probability of being high or low
(2) Description
Occurs in estimating the value of the last digit of a measurement
(3) Also called indeterminate error
b. Systematic error
(1) Definition
An error that always occurs in the same
direction
(2) Description
Caused by a defect in analytical technique, an improperly functioning instrument, or an improper use of an instrument by the analyst
(3) Also called bias
5. Random errors and systematic errors and accuracy and precision
a. Random errors
Random errors can lead to poor precision.
However, random errors should not affect accuracy if enough measurements are made.
a. Systematic errors
Systematic errors can lead to poor accuracy.
Systematic errors can be very precise.
B. Significant digits
1. In measurements
a. Definition
All of the digits that can be known precisely plus a last digit that must be estimated.
b. In measurements the significant digits indicate the limits
of precision.
2. In calculations
In the result of a calculation the significant digits are all of the certain digits plus one uncertain digit.
C. Identifying significant digits
1. All nonzero digits are significant.
Example:
9876 4 significant digits
2. All captive zeros are significant.
(zeros between two nonzero digits)
Examples:
202 3 significant digits
1001 4 significant digits
3. No leading zeros are significant.
(zeros to the left of the leftmost nonzero digit)
Examples:
0.45 2 significant digits
.0051 2 significant digits
4. Some trailing zeros are significant.
a. When there is no decimal point the zeros are not
significant.
Example:
1700 2 significant digits
b. When there is a decimal point the zeros are significant.
Examples:
1700. 4 significant digits
83.0 3 significant digits
950.0 4 significant digits
5. In scientific notation, all of the digits in the number are
significant, but none of the digits in the exponent are significant.
Examples:
4 x 103 1 significant digit
6.29 x 10(41 3 significant digits
6.29 x 10(2 3 significant digits
6. Exact numbers have an infinite number of significant digits.
a. Exact due to counting
Example:
5 beakers
b. Exact due to defining
Example:
1 minute is defined to be 60 seconds
D. Using significant digits in calculations.
(Rule: An answer cannot be more precise than the least precise
measurement)
1. In addition and subtraction, locate the leftmost uncertain digit
and round to that place.
Examples:
12.52 g
49.0 g
+ 8.24 g
369.76 g
answer 369.8 g or 3.698 x 102 g
36,900 m
( 158 m
36,742 m
answer 36,700 m or 3.67 x 104 m
2. In multiplication and division count the significant digits and
round to the same number of digits as the measurement with
the fewest.
Examples:
7.55 cm x 0.34 cm = 2.567 cm2
(3) (2)
answer: 2.6 cm2
EMBED Equation.3
EMBED Equation = 0.291976190 g/cm3
answer: 0.29 g/cm3 or 2.9 x 10(1 g/cm3
REMEMBER: The good student EMBED Equation EMBED Equation out as
ASsuMD
(assumed)
Locate Uncertain Count Significant digits
Addition Subtraction Multiplication Division
3. Rules of rounding
If the digit to the right of the rightmost significant digit is:
a. Less than 5, then drop it
Example:
6.43 rounds to 6.4
b. More than 5, then round up
Example:
6.46 rounds to 6.5
c. 5, then round the rightmost significant digit to the
nearest even digit (this avoids a bias)
Another way to do the same thing is to look at the
digit to the left of the 5:
If the left-hand digit is even, leave it.
If the left-hand digit is odd, up it.
Even leavin; odd up
Examples:
6.45 rounds to 6.4
6.55 rounds to 6.6
TEMPERATURE CONVERSIONS
the 5, 9, and 32 are EXACT
the 273.15 is MEASURED
(C = 5/9 ((F ( 32)
(F = 9/5(C + 32
K = (C + 273.15
Convert 68.51(F to Celsius
(C = 5/9 (68.51 ( 32) = 20.28333333 = 20.28(C
Convert 24.08(C to Fahrenheit
(F = 9/5 (24.08) + 32 = 75.344 = 75.34(F
Convert 27.8(C to Kelvin
K = 27.8 + 273.15 = 300.95 = 301.0 K
DENSITY CALCULATIONS
25.0 mL of a liquid weighs 17.84 g. What is its density?
GivenFindV = 25.0 mL
m = 17.84 g d = ?
Equation:d=mV
Substitute into the equation
d=17.84 g= 0.7136 g/mL= 0.714 g/mL25.0 mL
Check: Are the units correct? Yes!
A liquid has a density of 3.10 g/mL. What is the volume of a sample weighing 88.50 g?
GivenFindd = 3.10 g/mL
m = 88.50 g V = ?
Equation:d=mV
After algebra:V=md
Substitute into the equation:
V=88.50 g = 28.548387g3.10 g/mLg/mL
V = 28.5 mL
Check: Are the units correct? Yes!
DIMENSIONAL ANALYSIS
also called THE FACTOR - LABEL METHOD
also called THE UNIT CANCELLATION
METHOD
A. Key points
1. Every measurement has a number and a unit.
2. A conversion factor enables us to change a measurement from
one unit to another
3. A conversion factor is equivalent to 1.
Examples:
1 dozen12 objects
12 inches1 foot
1 min60 s
B. Useful conversion factors
1. metric ( metric conversions
(exact)
Note: 1 mL = 1 cm3 (exactly)
2. English ( English conversions
(exact)
3. Helpful metric ( English conversions
(NOT exact)
a. Length
1 mi = 1.6093 km
1 in = 2.5400 cm
1 m = 39.370 in
b. Mass
1 kg = 2.2046 lb
c. Volume
1 L = 1.0567 qt
C. Conversion problems
1. The general approach to conversion problems
a. The first step is to draw a map.
b. The second step is to replace each arrow in your map
with a conversion factor that has the new units
cattycorner to the old units.
c. The third step is to set up this conversion factor in a
big, long line format.
2. One-step and two-step conversions
a. One step conversions
Convert 29.3 inches to feet
First step: Draw a map.
map: inches (( feet
Second step: Replace each arrow in your
map with a conversion factor
that has the new units
cattycorner to the old units.
inches (1 ft= feet12 in
Third Step: Set up this conversion factor
in a big, long line format.
Solution:
29.3 in1 ft12 in
= 2.4416667 ft
= 2.44 ft
b. Two-step conversions
Convert 3.77 inches to meters
map: inches (( cm (( m
solution:
3.77 in2.5400 cm1 m1 in100 cm
= 9.5758 x 10(2 m = 9.58 x 10(2 m
3. Conversions involving units that are ratios
a. General approach
(1) The first step is to draw a map so that the units
on the top are converted first, and then the units
on the bottom.
(2) The second step is to replace each arrow in your
map with a conversion factor that has the new
units cattycorner to the old units.
(3) The third step is to set up this conversion factor
in a big, long line format.
b. Example
Convert 176 m/s to km/hr
map:m((km((kmsshr
solution:
176 m1 km3600 ss1000 m1 hr
= 6.336 x 102 km/hr
= 6.34 x 102 km/hr
4. Conversions from cubic volume to volume
a. The general approach
(1) The first step is to draw a map.
(2) The second step is to replace each arrow in your
map with a conversion factor that has the new
units cattycorner to the old units.
(3) The third step involves using three linear
conversion factors to get from the original cubic
linear units to the new desired cubic linear units
or to the new unit.
(4) The fourth step is to set up this conversion factor
in a big, long line format.
b. Example
Convert 9.07 mm3 to mL
map: mm3 EMBED Equation.3 cm3 EMBED Equation.3 mL
* Note that because the unit is to the third power, we need to use the same conversion factor three times.
solution:
9.07 mm31 cm1 cm1 cm1 mL10 mm10 mm10 mm1 cm3
= 9.07 x 10(3 mL
PAGE
PAGE 11
Topic 1 Matter and Measurement
2006 Lloyd Crosby
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