Transcript Chem 105 Power Pt Lectures

Chem 140 Section A
Instructor: Ken Marr
Weekly Schedule
“Lecture”
“Lab”
9 -10, MWF in STB-2
8 -10 , Tu in STB-2
8 -10 , Th in STB-5
Chem 140 Section C
Instructor: Ken Marr
Weekly Schedule
“Lecture”
“Lab”
10 –11, MWF in STB-2
10-12 , Tu in STB-2
10-12 , Th in STB-5
Chem 140 Section E
Instructor: Ken Marr
Weekly Schedule
“Lecture”
“Lab”
1 - 2, MWF in STB-2
1 - 3 , Tu in STB-2
1 - 3 , Th in STB-5
Day 1 Activities

Introduction to Course
» Briefly Review Course Outline/Syllabus

Homework Assignments
» Reading: See Chem 140 schedule
» Lab: Do Prelab assignment for the “Measurement and
Density” Lab
» Stamped Assignment #1: Chapter 1 HW
– due Tues. 10/01/02 ……But start now!!!
– Note: 10/2 is a very special day for your instructor!!

Begin Chapter 1
» Alice 1 and 2
CHEMISTRY
The Study of
Matter and the
Changes that
Matter
Undergoes and
The Energy
Associated with
The Changes
Chemistry as the Central Science
Oceanography
Atmospheric
Sciences
Engineering
Physics
Economics
Medicine
Governments
Chemistry
People
Geology
Biology
Anthropology
Politics
Astronomy
Chapter# 1 : Keys to the Study of Chemistry
1.1 Some Fundamental Definitions
1.2 Chemical Arts and the Origins of
Modern Chemistry
1.3 The Scientific Approach: Developing a Model
1.4 Chemical Problem Solving
1.5 Measurement in Scientific Study
1.6 Uncertainty in Measurement:
Significant Figures
Measurement and Significant Figures

Measured Numbers are Never Exact...Why?
» Which Graduated Cylinder is the most precise?
» How is precision indicated when we record a
measurement?
The Number of Significant Figures in a
Measurement Depends Upon the
Measuring Device
Fig 1.14
3e
Significant Figures
 We
use significant figures to indicate the
maximum precision of a measurement
 Significant Figures
» The number of digits that are known with
certainty, plus one that is uncertain
» Significant figures are used only with
measured quantities.
» Some numbers are exact and do not have any
uncertainty......e.g...’s??
More Examples
Record the exact length in centimeters, cm
(T2c)
 Record the exact amounts for numbers 1-11
(T2d)

Sig. Fig. Rules to memorize.....
(See page 27-30 Silberberg 3e)
1.
All nonzero numbers are significant
e.g.
2.
3.
23.8 g,
2345 km, 11 mL, 5 inches
Zeros between nonzero digits are significant
i.e. Sandwiched zeros are significant
e.g. 509 m, 2001 mL, 2050.1 L
Zeros preceding the 1st nonzero digit are not
significant.......they serve only to locate the decimal
point
e.g. 0.083 m, 0.000306 L
–
Try converting these numbers to Scientific Notation to
prove this!
More Sig. Fig. Rules Involving Zeros
4.
Zeros at the end of number that include a
decimal point are significant
»
5.
0.800,
11.40,
10.00, 400.
Zeros at the end of a number without a
decimal point are not significant... The
Greenwater Rule!
»
40,
8800,
300,
– Use of underlining and decimal points
Examples of Significant Digits in Numbers
Number
0.0050 L
0.00012 kg
83.0001 L
0.006002 g
875,000 oz
30,000 kg
5.0000 m3
23,001.00 lbs
0.000108 g
1,470,000 L
- Sig digits
Number
-
1.34000 x 107 nm
87,000 L
78,002.3 ng
0.000007800 g
1.089 x 10 -6L
0.0000010048 oz
6.67000 kg
2.70879000 ml
1.0008000 kg
1,000,000,000 g
Sig digits
Rounding off Numbers

Rounding off is used to drop non-significant
numbers
» Rule 1
When the 1st digit after those you want to retain
is 4 of less, that digit and all others to the right
are dropped

Round off the following to 3 sig. figs.
» 105.29,
189.49999,
1.003, 100.3,
1001
Rounding off Numbers

Rule 2
When the 1st digit after those you want to retain
is 5 or greater, that digit and all others to the
right are dropped and the last digit retained is
increased by one

Round off the following to 4 sig. figs.
» 10.87519, 13.59800, 99.999,
1042.5
Sig. Figs. in Calculations

The Central Idea.....
» The result of a calculation based on
measurements can not be more precise
than the least precise measurement!

Some Rules to, yes, memorize......
Sig. Figs. in Multiplication and Division

“The Chain Rule”
» Your answer must contain the same
number of sig figs as the measurement
with the fewest sig figs.....Some e.g...’s...
(3.04) x (2.2) = 6.688 = ???
(2.00) / (0.3 ) = 6.666... = ???
(18.4) x (4.0) = 1.1117824 = ???
(66.2)
Sig. Figs. in Addition and Subtraction

“The Decimal Rule”
» The answer must have the same precision
as the least precise measurement...or...
– Your answer must be expressed to the same
number of decimal places as the
measurement with the fewest decimal
places.
 The number of sig figs are not
considered, only the number of decimal
places are considered!!!
» Some examples..
Sig. Figs. in Addition and
Subtraction

Examples.....
» 12.89 + 12.1 + 11.803 + 19 = 55.793 = ?
» 1786 - 130 = 1656 = ???
» 7331 + 0.495 = 7331.495 = ???
Scientific Notation

Scientific Notation
» Writing a number as a number between 1
and 10 times a power of 10
» WHY DO IT???

The Rules...
How to Write Numbers in Scientific Notation
1.
Move the decimal point in the original number
so that it is located after the first nonzero digit
» e.g. 5682  ????
2.
Multiply this number by the proper power of 10
» The power of 10 is equal to the number of places the
decimal point was moved.


POSITIVE IF MOVED TO THE LEFT
NEGATIVE IF MOVED TO THE RIGHT
Examples....

Express the following numbers in scientific
notation...
» 0.0421
» 150,000
» 5899

Express the following in “longhand”
-4
» 5.30 x 10
6
» 8.000 x 10
Meaning of Powers of 10
103
 102
 101
 100

=
=
=
=
10-3 =
10-2 =
10-1 =
Metric System
System of measure built around
standard or base units
 Uses factors of 10 to express larger or
smaller numbers of these units

Table 1. 2 (p. 17, 3e)
SI - Base Units
Physical Quantity
Unit Name
Abbreviation
Mass
Kilogram
kg
Length
meter
m
Time
second
s
Temperature
Kelvin
K
Electric current
ampere
A
Amount of substance mole
Luminous intensity
candela
mol
cd
Metric Base Units and their
Abbreviations
Length
 Mass
 Volume
 Temperature

» Prefixes are added to these base units for quantities
larger or smaller than the base unit
– Prefixes are a multiple of 10
Table 1.3
Prefix
tera
giga
Mega
Kilo
hecto
deka
----deci
centi
milli
micro
nano
pico
femto
Common Decimal Prefixes Used with SI Units.
Prefix
Symbol
Number
Word
Exponential
Notation
T
1,000,000,000,000
trillion
G
1,000,000,000
billion
M
1,000,000
million
k
1,000
thousand
h
100
hundred
da
10
ten
---1
one
d
0.1
tenth
c
0.01
hundredth
m
0.001
thousandth
millionth
n
0.000000001
billionth
p
0.000000000001
trillionth
f
0.000000000000001
quadrillionth
1012
109
106
103
102
101
100
10-1
10-2
10-3
10-6
10-9
10-12
10-15
Common Metric Prefixes

Memorize the Symbol, Numerical Value,
and Power of 10 Equivalent for.....
» kilo» centi» milli» micro» nano-
Common Prefix Applications

Length:
» km
» cm
» mm
» µm
» nm
1 km = ?
1 cm = ?
1 mm = ?
1 µm = ?
1 nm = ?
m
m
m
m
m
Common Prefix Applications

Mass
» kg
» mg
» µg
1 kg = ? g
1 mg = ? g
1 µg = ? g
Common Prefix Applications

Volume
» mL
» µL
1 mL = ? L
1 µL = ? L
Important Relationships

Length
»1m =
»1m =
»1m =
» 1 cm =
?? cm
?? mm
?? µm
?? mm
Important Relationships

Mass
» 1 g = ?? mg
» 1 kg = ?? g
» 1 kg = ?? lb..
Some Volume Relationships in SI Units
Fig. 1.9
Important Relationships

Volume
» 1 L = ?? mL
» 1 mL = ?? cm3
» 1 L = ?? cm3
» 1 L = 1.057 qt.
Solving Chemistry Problems
Develop a Plan  Carryout Plan  Check Answer
1.
Developing a Plan: Read the problem carefully!
• Clarify the know and unknown:


What information is given?
What are you trying to find?
• Think about how to solve the problem before you
start to juggle numbers



Suggest steps from the known to unknown
Determine principles involved and the relationships
needed
Use sample problems as a guides
• Map out the strategy you will follow
Solving Chemistry Problems (cont.)
2.
Solve the problem: Carry out your plan
»
»
»
»
»
Set up problem in a neat, organized, and logical way!
Unwanted units should cancel to give the desired
unit of measure
Make a rough estimate of the answer before using
your calculator
Round off to correct number of sig. figs.
Answer must have correct units
Solving Chemistry Problems (cont.)
3.
Check your answer
» Is it reasonable?
» Correct nits?
» Same “ballpark” as a rough estimate?
» Makes chemical sense?
Problem Solving: Some Examples
1.
How many hours would it take a pump to remove
the water from a flooded basement that is about 30
feet wide and 50 feet long with water at a depth of
about 2 feet? The pump has a capacity of 80 liters
per minute. See Table 1.4, Common SI-English
Equivalent Quantities, page 18 Silberberg 3e.
1062 min = 17.7 hours = 20 hours
Metric Conversion Factors

Be able to do conversions within the metric
system involving the common metric prefixes
» kilo» centi» milli» micro» nano– e.g. #32 on page 43
Metric - English Conversions
Given metric - English conversion factors, be
able to convert between these two systems
 You do not have to memorize metric to English
conversions factors

Measurement of Temperature

Heat vs. Temperature
» Temperature
(SI unit: Kelvin, K)
– A measure of how hot or cold an object is relative to
another object
– Also measured in degrees Celsius, oC
» Heat
(SI unit: joule, J)
– The energy transferred between objects at different
temperatures
– A form of energy associated with the motion of atoms
and molecules (the small particles of matter)
– Also measured in calories, cal
Application: Heat vs. Temperature

Which contains more heat...
» 1 mL of water at 90 oC or 1 liter of water at 90 oC ?
» 1 burning match or 10 burning matches?
Temperature Conversions
The boiling point of Liquid Nitrogen is - 195.8 oC, what is
the temperature in Kelvin and degrees Fahrenheit?
T (in K) = T (in oC) + 273.15
T (in K) = -195.8 + 273.15 = 77.35 K = 77.4 K
T (in oF) = 9/5 T (in oC) + 32
T (in oF) = 9/5 ( -195.8oC) +32 = -320.4 oF
The normal body temperature is 98.6oF, what is it in Kelvin
and degrees Celsius?
T (in oC) = [ T (in oF) - 32] 5/9
T (in oC) = [ 98.6oF - 32] 5/9 = 37.0 oC
T (in K) = T (in oC) + 273.15
T (in K) = 37.0 oC + 273.15 = 310.2
Density

Density = mass (g) / Volume (mL or cm3 or L)
» Physical characteristic of a substance
» Aids in identification of a substance
» Calculated by.....
– divide the mass of a substance by the volume occupied by that
mass
– Units


mass in grams
volume
» Solids and Liquids: mL or cm3
» Gasses: L
Density

Densities vary with temperature!
» Why??

Would you expect densities to increase
or decrease as the temperature
increases?
Density

Immiscible liquids and solids separate into
layers according to their densities
» List the order from top to bottom when the following
are mixed
– Hg (13.5525 g/mL)
– Carbon Tetrachloride (1.59525 g/mL)
– Mg (1.7425 g/mL)
– Water (1.004 g/mL)

What do the superscripts mean next to each density listed
above?
Calculations Involving Density

Be able to calculate the density, mass, or
volume of a substance
» Use the plug and chug method or use density as a
conversion factor

Practice makes perfect....
Specific Gravity

Compares the density of a liquid or solid to that
of water... Units???
»
Sp. Gravity = dsolid or liquid / dwater
– Usually use dwater @ 4oC = 1.000g/mL

Compares the density of a gas to that of air......
Units???
» Sp. Gravity = dgas/ dair
Mass vs. Weight

Mass
» Amount of matter in an object
» Independent of location
» Measure with a balance by comparison with
other known masses
Mass vs. Weight

Weight
» Measures earth’s gravitational attraction on
an object
» Measure with a scale
– measures force against a spring
» Depends on
– position relative to earth
– motion of object w.r.t. the earth
Scientific Approach: Developing a Model
Observations : Natural phenomena and measured events; universally
consistent ones can be stated as a natural law.
Hypothesis: Tentative proposal that explains observations.
Experiment: Procedure to test hypothesis; measures one variable
at a time.
Model (Theory): Set of conceptual assumptions that explains data
from accumulated experiments; predicts related
phenomena.
Further Experiment: Tests predictions based on model.