Transcript Chapter 3 Scientific Measurement - A

Chapter 3
Scientific Measurement
Chemistry 2
Measurements and Their
Uncertainty 3.1
1. Differentiate accuracy and precision.
2. How do you think scientist ensure measurements are accurate and precise?
Using and Expressing
Measurements 3.1
Quantity that has both a # AND a UNIT
 SI- International System of
Measurements

• Why SI and not IS?

Scientific Notation – using power of 10’s
to write small/large #s.
• # less than +/- 1 have a - exponent
• # bigger than +/- 1 have a + exponent
Scientific Notation 3.1

Multiplying numbers in scientific notation
1.Multiply number that appear before the
multiplication signs
2.Add the exponents

Dividing numbers in scientific notation
1.Divide the numbers that appear before the
exponential term
2.Subtract the exponents

Page R57 – Sample Problem MH-2
Accuracy, Precision, and Error 3.1
Accuracy, Precision, and Error 3.1



Accuracy – measurement of how CLOSE to
ACTUAL value
Precision – measurement of how close to a
series of measurement (COMPARING) or more
small measurement
Determining Error
• Error = experimental value – accepted value

Example: water boils at 100oC, in an experiment you find
water to boil at 99.1oC
• 99.1oC – 100oC = -.9oC Error
• Percent error =(lErrorl/Accepted Value) x 100


Error = absolute value  always a + answer
Example: (.9oC/100oC) x 100 = .9%
Significant Figures in Measurement 3.1

Measurement all the digits that are
known, plus a last digit that is estimated
• Example: have a scale that measures to the
tenth digit – you can measure up the nearest
hundredth digit

Show precision
Rules of Sig Figs

Every nonzero is significant
• 1.2583 = 5 sig figs

Zeros between nonzero digits are significant
• 10.001 = 5 sig figs

Zeros on the left of nonzero digits are not significant. They are
placeholders
• 0.0005 = 1 sig fig

Zeros at end of number to the right of the decimal point are
always significant
• 45.00 = 4 sig figs
• 0.0030 = 2 sig figs

Zeros at the rightmost end of a measurement that are to the left
of a decimal are not significant. However, if they are certain, they
are significant (normally use a decimal to show)
• 300 = 1
• 300. = 3
• Or 300 could be 3 IF 3.00 x102 is written

2 situations:
• Counting = Infinite sig figs
• Defined quantities: Infinite Sig Figs

Do not affect process of rounding an answer
Significant Figures in Calculations 1.3

Calculated answer cannot be more
precise than the lease precise
measurement from calculation
• Rounding
1st: How many sig figs should answer have?
 2nd: round to that many digits, counting from the
left
rd
 3 : if digit to right of last sig dig is less than 5 it
is dropped, if it is 5 or higher, last digit is
increased by 1

Addition and Subtraction of Sig
Figs 3.1

Round to the same number of DECIMAL
PLACES as the measurement with the
least number of DECIMAL PLACES
• 5.2 + 5.1111 = 10.3111 10.3
Multiplication and Division of Sig
Figs 3.1
Round the answer to the SAME
NUMBER sig figs as the measurement
with the LEAST NUMBER OF SIG FIGS
 3.4672 x 10 = 34.672  30

The International System of Units 3.2
Measuring with SI – 3.2

Metric system
• Simple & easy to use – multiples of10
Est. in France, 1795
 International adoption in 1960
 7 base units

Base Quantity
Base Unit
Symbol
Length
Meter
m
Mass
Kilogram
kg
Time
Second
s
Temperature
Kelvin
k
Amount of a substance
Mole
mol
Electric Current
Ampere
A
Luminous Intensity
Candela
cd
Units and Quantities 3.2
Kids Have Dropped Over Dead Converting Metrics
Kilo Hecto Deca
Origin deci Centi
Mili
Kg hg
dag(Dg) g
dg cg
mg
Giga = 1 000 000 000 = G
Mega = 1 000 000 = M
Micro = .000 001 = u
Nano = .000 000 001 = n
Conversion Factor – ratio expressed in one unit to
another unit

Units and Quantities 3.2
Length
• Linear measurement
• Meter (m)

Volume
• Liter (L) (10cm x 10 cm x 10 cm = 1000cm3 or 1L)
• Common: L, mL, cm3, microliter

Mass – measure of quantity of matter
• Kg
• Weight is different than mass


Weight – force that measures the pull on a given mass by gravity
Temperature
• Heat transfer: Hot  Cold
• Almost all substances expand with increase in temp and contact when it gets
cold (except water)
• K – freezing point of water 273.15K & boiling point 373.15K
• C – freezing point of water is 0 C &
boiling point is 100 C
• Absolute zero = 0K or -273.15 C
• K = C +273

Energy – capacity to do work or to produce heat
• Joule (J)
• James Prescott Joule – English Physicist
• Calorie (cal) – quantity of heat that raises the temp of 1 g of pure water by 1
C
• 1J = .2390cal or 1 cal = 4.184 J
Conversion Problems 3.3
Conversion Factors 3.3
$1 = 4 quarters = 10 dimes = 20 nickels
 1 m = 10 dm = 100 cm = 100 mm
 Ratio of equivalent measurements

• 100cm/1m

When Conversion factor is used,
numerical value changes, but actual size
does not
Dimensional Analysis 3.3
Way to analyze and solve problems using
units of measurement
 Different ways to solve every problem

Technology and Society

Scale Models – physical or conceptual
representation of object that is
proportional in size to the objects it
represents
• CAD programs
• Most trains are built in a scale of 1:87

Ratio means 1/87 the size of actual train
Who would use a scale model?
 What is the height of a building if it is
represented by a 1:30 m model using a
scale factor of 1:48?
 Answer:

• 62.4m
Density 3.4
Determining Density

Which is heavier, a lb of feathers or a lb
of lead?
• Lead has greater mass
• Larger volume of feathers needed
Density 3.4
Density = mass/volume
 Intensive property that depends only on
the composition of a substance, not the
size of the sample
 Density and Temperature

• Most substance: as temp. increases, volume
increases BUT mass remains the same…
• SO… Density must change as well
Careers in Chemistry

Analytical Chemist
• Spend most of time with measurement and
calculations
• Analyze composition
• Pharmaceuticals
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