Transcript 1 - Clayton State University

Daniel L. Reger
Scott R. Goode
David W. Ball
www.cengage.com/chemistry/reger
Chapter 1
Introduction to Chemistry
The Nature of Science and Chemistry
• Definitions
• Science: Study of the natural universe
• Specifically, knowledge acquired by experience
• Science is both an activity and the result of the
activity.
• Chemistry: the study of matter and its interactions
with other matter and with energy.
• Chemistry is how matter is
organized/reorganized/changed at the molecular level
Chemistry: The Central Science
• Chemistry is often called the central science because it is an
essential component of the natural and life sciences.
Chemistry and Astronomy
• Elemental composition of stars can be determined by
different wavelengths of visible light emitted.
• When starlight passes through a planets atmosphere,
certain frequencies of light disappear because they are
absorbed by compounds in the atmosphere.
http://cs.fit.edu/~wds/classes/cse5255/cse5255/davis/text.html
Chemistry and Geology
• Geochemistry: Study of the chemical composition of the
earth
• Chemical transformations in solids
• Ex. Polymorphism.
• How limestone becomes marble.
http://geology.com/rocks/limestone.shtml; http://www.italartworld.com/
Chemistry and Biology
• You reach a certain level in biology where processes can
only be understood in terms of chemistry.
• Chemistry in biology explains:
• Why your adrenaline levels increase when you are afraid or
excited
• Why a body fails to produce insulin (diabetes)
• Why cells become cancerous
• Neurotransmitter (e.g., dopamine, norepinephrine) imbalances
that can produce:
• Euphoria when you have a few beers or fall in love
• Depression
http://cs.fit.edu/~wds/classes/cse5255/cse5255/davis/text.html
The Scientific Method
Scientific method: investigations that are
guided by theory and earlier experiments.
• Hypothesis: a possible explanation for an
event.
• Law: a statement that summarizes a large
number of observations.
• Theory: an explanation of the laws of nature.
•
•
In the realm of science, theory has a much narrower
and more rigorous meaning than in general.
Matter, Mass and Weight
•
•
•
Matter: anything that has mass and occupies space.
Mass: the quantity of matter in an object.
Weight: the force of attraction between an object and other
objects.
Mass on
moon and
earth is the
same.
Weight on
moon and
earth is the
different.
Properties of Matter
Matter can be described by different properties
• Property: anything observed or measured about a
sample of matter.
•
Extensive property: depends on the size of the sample.
• mass, volume
• Intensive property: independent of sample size.
• density, color, melting or boiling point
• 2 samples with same intensive properties may be the
same material
•
Physical Properties and Changes
•
•
Physical properties: can be measured without changing the
composition of the sample.
• mass, density, color, MP, BP, solubility
Physical change: a change that occurs without changing the
composition of the material.
• freezing, melting, crushing a brick into powder
Physical Change
Chemical Change
Chemical Properties
•
Chemical properties: describe the reactivity of a
material.
– Flammability (whether something is ignitable, flash point <
100 °F)
– Combustibility (whether something will burn, flash point >
100 ° F)
• Iron rusts
•
Chemical change: at least part of the material is
changed into a different kind of matter.
•
•
Digestion of sugar is a chemical change
Acid/base neutralization
Classification of Matter - Substances
• Substances - Material that is chemically the same throughout.
• cannot be separated into component parts by physical methods
• Two types of substances
• Elements cannot be broken into simpler substances by
chemical methods
• Table 1.1 (p. 09): Memorize these!!!
• Compounds can be separated into simpler substances (or
elements) by chemical methods
• Always contain same elements in same proportions (H2O is
always 11.2% H and 88.8% O)
Classification of Matter - Mixtures
•
Mixture: matter that can be separated into simpler materials by
physical methods.
• Heterogeneous mixture: composition of the mixture changes
from one part to another.
• Chocolate chip cookies
• Italian dressing
• Homogeneous mixture or solution: composition of the mixture
is uniform throughout.
• Chocolate pudding
• Sugar dissolves in water
• Alloy: a solution of a metal and another material (usually another
metal).
What’s the Matter?
Accuracy and Precision
•
Modern chemistry is largely based on experimental measurements.
The confidence in measurements involves:
• Accuracy: agreement of a measurement with the true value.
• Precision: agreement among repeated measurements of the
same quantity.
accurate and
precise
accurate but
not precise
precise but
not accurate
neither accurate
nor precise
Significant Figures
• Significant figures are the numbers in a measurement that
represent the certainty of the measurement, plus one number
representing an estimate.
Q: When is a number NOT significant?
A: Look at the zeros
Leading zeros are NOT significant.
Confined zeros ARE significant.
Trailing zeros ARE significant, when decimal visible
But NOT significant if no decimal
Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7 th Edition, 2011
0.00123
0.00103
0.0012300
12300
Calculations with Significant Figures
Rules for Rounding
• If the first nonsignificant figure to drop from your answer is ≥ 5, all
nonsignificant figures dropped, last significant figure increased by 1.
• If the first nonsignificant figure to drop from your answer is < 5, all
nonsignificant figures dropped, last significant figure stays the same.
Exact Numbers
Numbers with no uncertainty, or are known values. Exact numbers do
not change.
Ex. 1 foot is always = 12 inches. It will never be = 12.5 inches.
Not used to determine sig figs
• Na = 1 mol = 6.02 x 1023
• π= 3.142
• 1m = 1000 mm
Significant Figures
How many significant figures are
present in each of the measured
quantities?
• 0.0012
106
2006
900.0
1.0012
0.001060
•
Significant Figures
•
How many significant figures are
present in each of the measured
quantities?
• 0.0012
2
106
3
2006
4
900.0
4
1.0012
5
0.001060
4
Significant Figures
Trailing zeros in numbers without a
decimal point may not be significant.
Avoid ambiguity by using scientific
notation.
• 100
1, 2 or 3
1 x 102
1
2
1.0 x 10
2
2
1.00 x 10
3
•
Test Your Skill
•Determine
the number of significant figures:
100.
100.0
30505
437,000
-3
125,904,000
4.80 x 10
-3
4.800 x 10
0.0048
Test Your Skill
•Determine
the number of significant figures
100.
100.0
30505
437,000
-3
125,904,000
4.80 x 10
-3
4.800 x 10
0.0048
•Answer:
100.
3
100.0
4
30505
5
437,000
3-6
-3
125,904,000 6-9 4.80 x 10
3
4.800 x 10-3 4
0.0048
2
Scientific Notation
• Scientific notation provides a convenient way to
express very large or very small numbers.
• Numbers written in scientific notation consist of a
product of two parts in the form M x 10n, where M is a
number between 1 and 10 (but not equal to 10) and n is
a positive or negative whole number.
• The number M is written
with the decimal in the
standard position.
Scientific Notation (continued)
• STANDARD DECIMAL POSITION
• The standard position for a decimal is to the right of
the first nonzero digit in the number M.
• SIGNIFICANCE OF THE EXPONENT n
• A positive n value indicates the number of places to
the right of the standard position that the original
decimal position is located.
• A negative n value indicates
the number of places to
the left of the standard
position that the original
decimal position is located.
Scientific ↔ Standard Notation
• Converting from scientific notation to standard numbers
1.1 x 102 = 1.1 x 10 x 10 = 1.1 x 100 = 110
Decimal 
1.1 x 10-2 = 1.1/ (10 x 10) = 1.1/100 = 0.011
 Decimal
Converting Exponents
Ex. 1
0.67 x 10-5
0.067 x 10-4
0.0067 x 10-3
0.00067 x 10-2
0.000067 x 10-1
0.0000067 x 100
0.00000067 x 101
0.000000067 x 102
Ex. 2
1.2 x 10-3
0.12 x 10-2
0.012 x 10-1
0.0012 x 100
0.00012 x 101
• When you move the decimal (l or r), the exponent will be
equal to number of places you moved the decimal.
Standard to Scientific Notation
60023.5
345.233
-345.233
0.00345
0.10345
1.42






6.00235 × 104
3.45233 × 102
-3.45233 × 102
3.45 × 10-3
1.0345 × 10-1
1.42 × 100
Calculations with Significant Figures
• Sum (addition) or difference (subtraction) must
contain the same number of places to the right
of the decimal (prd) as the quantity in the
calculation with the fewest number of places to
the right of the decimal (i.e., the least
accurate number).
Ex. 01
Ex. 02
5.325  5.5
 10.825  10.8
3 prd  1prd 
5.325  5.5
1 prd
 0.175  0.2
3 prd  1prd 
Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7 th Edition, 2011
1 prd
Calculations with Significant Figures
Product (multiplication) or quotient (division) must
have same number of sig figs as value with the
fewest number of sig figs (least accurate number).
Ex. 01
4.325  4.5  19.4625  19.
4 SF 2 SF 
2 SF
Ex. 02
4.325  4.5  0.961  0.96
4 SF  2 SF 
2 SF
Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7 th Edition, 2011
Mixed Operations
•
Determine accuracy in the same order as you
complete the mathematical operations, # of
significant digits are in red.
m
v
•
=
3
2.79 g
=
8.34 mL - 7.58 mL 0.76mL
2
3
3
density = 3.7 g/mL
2
3
2.79 g
Test Your Skill
•
Evaluate each expression to the correct number of
significant figures.
(a) 4.184 x 100.620 x (25.27 - 24.16)
(b)
(c)
8.925 - 8.904
x 100%
8.925
9.6 x 100.65
+ 4.026
8.321
Answers: (a) 467; (b) 0.24%; (c) 1.2 x 102
Test Your Skill
Calculate each to the correct number of
significant figures .
0.1654 + 2.07 - 2.114
8.27 x (4.987 - 4.962)
9.5 + 4.1 + 2.8 + 3.175
4
9.025 - 9.024
9.025
(4 is exact)
x 100%
Test Your Skill
Calculate each to the correct number of
significant figures .
0.1654 + 2.07 - 2.114
= 0.12
8.27 x (4.987 - 4.962)
= 0.21
9.5 + 4.1 + 2.8 + 3.175
4
9.025 - 9.024
9.025
= 4.89
(4 is exact)
x 100%
= 0.1%
Math Operations with Scientific Notation
Multiplication
Division
Addition/Subtraction
(a 104 )(b 102 )  (a  b)(104 2 )
(a  104 )  a  4  2
  (10 )
2
(b  10 )  b 
Convert numbers to the same exponents
(5.00 x 102) + (6.01 x 103) =
(0.500 x 103) + (6.01 x 103) =
(5.00 x 102) + (60.10 x 102) = (5.00 + 60.10) x 102 =
(65.10 x 102) = 6510
(6.01 x 103) - (5.00 x 102) =
(6.01 x 103) - (0.500 x 103) =
(60.10 x 102) - (5.00 x 102) = (60.10 - 5.00) x 102 =
(55.10 x 102) = 5510
Examples of Math Operations
Multiplication
a. (8.2 X 10-3)(1.1 X 10-2) = (8.2 X 1.1)(10(-3+(-2))) = 9.0 X 10-5
b. (2.7 X 102)(5.1 X 104) = (2.7 X 5.1)(102+4) = 13.77 X 106
Now change to Scientific Notation
1.4 X 107
Division
a. 3.1 X 10-3 = (3.1/1.2)(10-3-2) = 2.6 X 10-5
1.2 X 102
b. 7.9 X 104 = (7.9/3.6)(104-2) = 2.2 X 102
3.6 X 102
Adding/Subtracting
3.05 X 103 + 2.95 X 103 = (3.05 + 2.95)(103) = 6.0 X 103
Base Units in the SI
Quantity
Length
Mass
Time
Temperature
Amount
Electric current
Luminous intensity
Unit
Abbreviation
meter
m
kilogram
kg
second
s
kelvin
K
mole
mol
ampere
A
candela
cd
Common Prefixes Used With SI Units
Prefix
megakilocentimillimicronanopico-
Abbreviation
M
k
c
m
m
n
p
Meaning
106
103
10-2
10-3
10-6
10-9
-12
10
Unit Conversion Factors
Unit conversion factor: a fraction in which the
numerator is a quantity equal or equivalent to the
quantity in the denominator, but expressed in different
units
• The relationship 1 kg = 1000 g
• Generates two unit conversion factors:
•
 1kg 
 1000 g 

 and 

 1000 g 
 1kg 
What would be some other examples?
Conversion Among Derived Units
Volume is the product of three lengths.
• The standard unit of volume is the cubic meter
3
(m ).
100 cm = 1 m
(100 cm)3 = (1 m)3
106 cm3 = 1 m3
• Two important non-SI units of volume are the
liter and milliliter.
1 liter (L) = 1000 mL = 1000 cm3
1 mL = 1 cm3
•
Volume
Volumes can
be
expressed in
different
units
depending
on the size
of the object.
1 m3
contains
1000 L
1L
contains
1000 mL
1 mL = 1
cm3 or 1 cc
Using Unit Conversions
•
Express a volume of 1.250 L in mL,
3
3
cm , and m
1000 mL 

1.250 L  
  1,250 mL
 1L 
 1000 cm3 
3


1.250 L  
 1,250 cm

 1L

 1 m3 
-3 3
  1,250  10 m
1.250 L   6
3
 10 cm 
Factor Unit Method Examples
• A length of rope is measured to be 1834 cm. How many
meters is this?
• Solution:
• Write down known quantity (1834 cm).
• Set known quantity = units of the unknown quantity (meters).
• Use factor (100 cm = 1 m), to cancel units of known quantity
(cm) and generate units of the unknown quantity (m).
• Do the math.
1834 cm

m
 1m 
1834 cm
  18.34 m
 100 cm 
Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7 th Edition, 2011
Factor Unit Method Example
Q: If an arrow shot from a bow travels 30 yards in
1 second, many cm does it travel in 4 seconds?
•
•
•
•
•
Time = 4 s
Rate = 30 yards/sec
1 yard = 3 feet
1 foot = 12 in.
1 in = 2.54 cm
 30 yards  3 feet  12 in  2,54 cm 

4s


  ? cm
 1 s  1 yard  1 foot  1 in 
Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7 th Edition, 2011
Density
•
•
Density: mass per unit volume
m
d 
V
Density, in SI base units, is kg/m3
(kg m-3).
3
• Most commonly used density units are g/cm (g cm
3
or g/mL) for solids and liquids, and g/L for gases.
•
Conversions Between Equivalent Units
3
The density of Ti is 4.50 g/cm or
3
4.50 g = 1 cm
• Calculate the volume of 7.20 g Ti.
•
Conversions Between Equivalent Units
3
The density of Ti is 4.50 g/cm or
3
4.50 g = 1 cm .
• Calculate the volume of 7.20 g Ti.
•
 1 cm3 
3


 1.60 cm Ti
• 7.20 g Ti  

 4.50 g 
Percentage
• The word percentage means per one hundred. It is the
number of items in a group of 100 such items.
• PERCENTAGE CALCULATIONS
• Percentages are calculated using the equation:
part
%
 100
whole
• In this equation, part represents the number of
specific items included in the total number of items.
Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7 th Edition, 2011
Percentage Calculation
• A student counts the money she has left until pay day
and finds she has $36.48. Before payday, she has to
pay an outstanding bill of $15.67. What percentage of
her money must be used to pay the bill?
• Solution: Her total amount of money is $36.48, and the
part is what she has to pay or $15.67. The percentage
of her total is calculated as follows:
part
15.67
%
 100 
 100  42.96%
whole
36.48
Density Calculation
• A 20.00 mL sample of liquid is put into an empty beaker
that had a mass of 31.447 g. The beaker and contained
liquid were weighed and had a mass of 55.891 g.
Calculate the density of the liquid in g/mL.
• Solution: The mass of the liquid is the difference
between the mass of the beaker with contained liquid,
and the mass of the empty beaker or 55.891g -31.447 g
= 24.444 g. The density of the liquid is calculated as
follows:
m 24.444 g
g
d 
 1.222
v 20.00 mL
mL
Energy Calculations
Q: In order to lose 1 lb/week, you need to cut 500.0 Cal from your diet
each day, or use an equivalent number of joules by working each day.
How many equivalent joules would you have to spend on work to
achieve this each day? How many joules would you have to expend to
achieve this over 7 days?
To answer this, you need to first know the following:
• 1Cal = 1 kcal = 1000 scientific calories or 1 nutritional calorie
• 1 scientific calorie = 1 cal = 4.184 J
Temperature Scales
• The three most
commonly-used
temperature scales
are the Fahrenheit,
Celsius and Kelvin
scales.
• The Celsius and
Kelvin scales are
used in scientific
work.
Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7 th Edition, 2011
Temperature Conversions
• Readings on one temperature scale can be converted to
the other scales by using mathematical equations.
• Converting Fahrenheit to Celsius.
5 

C
F  32
9
• Converting Celsius to Fahrenheit.
9 

F
C  32
5


 
• Converting Kelvin to Celsius.

C  K  273
• Converting Celsius to Kelvin.
K   C  273
Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7 th Edition, 2011
Test Your Skill
•
Express 17.5°C in °F and in K.
Test Your Skill
•
Express 17.5°C in °F and in K.
•
Answer:
TF = 63.5° F; TK = 290.6 K
Useful Conversions/Constants
• Na = 1 mol = 6.02 x 1023
• π= (Circumference of circle/diameter
of circle) = 3.142
• 1 scientific calorie = 4.184 J
• 1Cal = 1 kcal = 1000 scientific
calories or 1 nutritional calorie
• 1 N = (kg)(m)/s2
• g = 9.81 m/s2 (earth’s gravity)
Chapter 1 Visual Summary