Download Lecture Notes 1 - Rutgers University

Chemical Principles 160:115
Prof. Luke A. Burke
http://camchem.rutgers.edu/~burke
Office: SCI 114B
Textbook: Chemistry, the Central Science, Brown, LeMay, and Bursten
9th or 10th Edition.
This course will cover chapters 1-13.
Office Hours: after Chem Prin class, and appointment
Lecture: MW: 8:30-11:30
Grading: 3 exams
3 x 20%
=
1 final (cumulative)
60%
40%
100%
If miss one exam (No make-ups):
2 exams
2 x 20%
1 final
(40+20%)
=
40%
60%
100%
There is no extra credit.Page xxv
Advice for learning and studying chemistry
• Keep up with your studying day to day. Don’t cram
Ch 1
Introduction: Matter and Measurement
1
• Focus your study.
• Keep good lecture notes.
• Skim topics in the text before they are covered in lecture.
Read the Intro and Summary first.
Then match the next day’s lecture notes to the text.
• After lecture, carefully re-read the topics covered in class.
Sample exercises (in the chapter)
Practice exercises (at end of chapter)
 Learn the language of chemistry
• Attempt all of the assigned end-of-chapter exercises
Spending more than twenty minutes on a single exercise is rarely effective
unless you know it is particularly challenging Chapter 1
Introduction:
Matter and Measurement
Chemistry is the study of the properties of materials and the changes that
materials undergo.
The Study of Chemistry
The Molecular Perspective of Chemistry
Matter is the physical material of the universe. It is anything that has mass and
occupies space.
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Introduction: Matter and Measurement
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Atoms are the almost infinitely small building blocks of matter.
Molecules combinations of two or more atoms with the atoms attached to one
another in a specific way.
O2 H2O
CO2
CH3 CH2 OH
HOCH2CH2OH
Macroscopic view versus microscopic view
Macroscopic is the scale of ordinary sized objects. Microscopic is the atomic
level. We explain macroscopic properties (e.g. boiling point) in terms of
microscopic properties (molecular structure)
Why Study Chemistry?
Classifications of Matter
States of Matter
Gas – atoms or molecules are far apart and moving very fast.
Liquid – particles are more closely packed and still moving relatively fast.
Solid – held tightly together, usually in definite arrangements, wiggle (vibrate)
only slowly in their fixed positions.
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Introduction: Matter and Measurement
3
Pure Substances and Mixtures
Pure substances – have fixed composition and distinct properties.
Elements – cannot be decomposed into simpler substances (have only one kind of
atom)
Compounds – composed of two or more elements (contain two or more kinds of
atoms)
Mixtures – combinations of two or more substances in which each substance
retains its own chemical identity and hence its own properties.
Heterogeneous – does not have the same composition, properties, and appearance
throughout. ex: sand
Homogeneous-constant composition throughout, also known as solutions.
Gaseous solutions air
Liquid solutions
gasoline
Solid solutions
brass
Separation of Mixtures
Solubility differences filtration
Volatility differences distillation
chromatography
Elements
112 elements
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Introduction: Matter and Measurement
4
Chemical symbols
H, O, N, C, Li, Cl, Br, Hg, Au
Periodic Table; elements with similar properties arranged in columns
Compounds – elements combine to form compounds. Compounds can be
decomposed to the elements.
The Law of Constant Composition: The elemental composition of a pure
substance is always the same.
Properties of Matter
Physical properties: color, odor, density, melting point, boiling point, hardness.
Chemical properties: reactivity; change in chemical composition.
Intensive properties: independent of amount of material present (melting point,
density).
Extensive properties: depend on amount of material present (mass, volume).
Physical and Chemical Changes
Physical changes do not effect the composition of the material. Most common is
change of state such as evaporation (vaporization: liquid to gas fusion: liquid to
solid sublimation: solid to gas)
Chemical changes (also known as reactions) a substance is changed into a
chemically different substance.
2 H2
Ch 1
 O2
 2 H2 O
Introduction: Matter and Measurement
5
Units of Measurement
Many properties of matter are quantitative.
Units must be specified (meters, grams, liters)
17.5 is meaningless without units (feet, inches, meters, gallons?)
17.5 cm specifies the length
The metric system is used scientific measurements.
SI Units (Système International d’Unitès)
Length and Mass
Unit of length is the meter.
Mass is a measure of the amount of material in an object, not the weight. Unit of
mass is the gram.
1 pound is about 2.2 kilograms.
The prefix kilo means one thousand.
Prefixes
Prefix
Abbreviation
Meaning
Example
Giga
G
109
Gigabyte
Mega
M
106
Megawatt
Ch 1
Introduction: Matter and Measurement
6
Kilo
k
103
Kilogram
Deci
d
10-1
Deciliter
Centi
c
10-2
Centimeter
Milli
m
10-3
Millimeter
Micro

10-6
micrometer
Nano
n
10-9
Nanometer
Pico
p
10-12
Picometer
Femto
f
10-15
femtosecond
Temperature
Fahrenheit
Celsius scale – based on freezing point (O °C) and boiling point (100 °C) of
water.
Kelvin scale – 0 K is –273.15 °C (absolute zero)
K  C  273.15
The Celsius and Kelvin scales have equal sized units.
5
C  F  32 
9
9
F  C  32
5
Ch 1
Introduction: Matter and Measurement
7
Derived SI Units
Speed = ms-1
force = ma = kg m s-2
(length)3
Volume
length cubed
SI: meter3
dm3 = 1 liter = 1000 ml = 1000 cm3
1 ml = 1 cm3
Density 
mass
volume

g
cm3
Density of water = 1.00 g/cm3
Uncertainty in Measurement
Exact numbers have defined numbers (c, e, π, R) or are integers that result from
counting numbers of objects (12 eggs is a dozen, 1000 g is a kg)
Ch 1
Introduction: Matter and Measurement
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Inexact numbers result from a measurement.
Equipment errors (calibration)
Human errors (how individuals take measurements)
Uncertainty always exists in measured quantities.
Precision and Accuracy
Precision is how closely individual measurements agree with one another
(reproducibility of results)
Accuracy is how closely individual measurements agree with the correct or
“true” value.
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Introduction: Matter and Measurement
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Significant Figures
Measured quantities are generally reported in such a way that only the last digit is
uncertain.
2.2405±0.0001 g
2.2405 g
All digits including the uncertain one are called significant figures. The number
above has 5 significant figures.
The number of significant figures indicates the exactness of a measurement.
Guidelines to significant figures in a measured quantity
1.
Nonzero digits are always significant
457 cm (3 sig. figs)
2.
Zeros between nonzero digits are always significant
1005 kg (4 sig. figs)
3.
2.5 g (2 sig. figs)
1.03 cm (3 sig. figs)
Zeros at the beginning of a number are never significant; they
merely indicate the position of the decimal point
0.02 g (1 sig. figs) 0.0026 cm (2 sig. figs)
4.
Zeros that fall both at the end of a number and after a decimal point
are always significant
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Introduction: Matter and Measurement
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0.0200 g (3 sig. figs)
5.
3.0 cm (2 sig. figs)
When a number ends in zero but contains no decimal point, the zeros
may or may not be significant
130 cm (2 or 3 sig. figs) 10,300 g (3, 4, or 5 sig. figs)
The use of exponential notation (Appendix A) avoids possible ambiguity seen in
the last example
Using exponential notation 10,300 g can be written
    g
(three significant figures)
    g
(four significant figures)
4
1.0300  10 g
(five significant figures)
Significant Figures in Calculations
In multiplication and division the result must be reported with the same number of
significant figures as the measurement with the fewest significant figures.
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Introduction: Matter and Measurement
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1.
If the leftmost digit to be removed is less than five, the
preceding number is left unchanged.
Rounding 7.248 to 2 sig. figs gives 7.2
2.
preceding
If the leftmost digit to be removed is 5 or greater, the
number is increased by 1.
Rounding 4.735 to three significant figures gives 4.74
Rounding 2.376 to two significant figures gives 2.4.
In addition and subtraction the result cannot have more digits to the right of
the decimal point than any of the original numbers.
20.4
1.322
Round off to 105
83
104.722
Dimensional Analysis
In dimensional analysis we carry units through all calculations.
A conversion factor is a fraction whose numerator and denominator are the same
quantity expressed in different units.
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Introduction: Matter and Measurement
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Example: 1 inch = 2.54 cm
2.54 cm
1in
and
1in
2.54 cm
How many centimeters are in 8.5 inches?
2.54 cm
Number of centimeters  8.5 in
 21.6 cm
1in
Given unit 
Desired unit
 Desired unit
Given unit
Using two or more conversion factors.
Convert 8.00 m into inches
100 cm  1 in 

  315 in
Number of inches  8.0 m 
 1 m 2.54 cm 
Conversions Involving Volume
What is the mass in grams of two cubic inches of gold?
The density of gold is 19.3 g/cm3.
19.3 g
1 cm 3
1 cm 3
19.3 g
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Introduction: Matter and Measurement
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3
2.54 cm  2.543 cm 3 16.39 cm 3

 

 1 in 
13 in 3
1 in 3
16.39 cm3 19.3 g 


Mass in grams  2.00 in 
3
3  633 g
 1 in
1 cm 

3

Summary of Dimensional Analysis
In using dimensional analysis to solve problems, we will always ask three
questions:
1.
What data are we given in the problem?
2.
What quantity do we wish to obtain in the problem?
3.
What conversion factors do we have available to take us
from the given quantity to the desired one?
Ch 1
Introduction: Matter and Measurement
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