Transcript Chapter 2 Measurements in Chemistry

Chapter 2
Measurements in Chemistry
Questions to be Answered
• What does a measurement involve?
• How do you make measurements
properly?
• How do we convert between
measurements of one unit to
measurements of a new unit?
Physical Quantities
• Physical quantities – measured physical
properties
– Number
– Unit
Proper Measurements
• Number
– Reflect the certainty to which the
measurement was made
• Unit
– Represent the type of measurement made
• Mass
• Volume
• Length
Measurement and Significant Figures
•
Number
– Certain digits - all digits
that can be stated as fact
• Read from smallest digit
– One Uncertain digit – the
first digit that is estimated
• No additional digits should
be recorded
– Ruler practice
– Uncertainty always exist in
the last digit of a number
• Balance example
Measurement and Significant Figures
• The total number of digits used to express
such a measurement is the number of
significant figures.
Measurement and Significant Figures
Instrument used directly impacts the certainty of
the measurement and hence the # of significant
figures that can be reported.
Scientific Notation
• Scientific notation - convenient way to write a
very small or a very large number.
– All digits listed in the number portion are significant
• Rules for conversion
– Move decimal so that it follows first non-zero digit
– Write all sig figs in number followed by (x 10)
– Raise the ten to the appropriate power
• Decimal moved left (+) the number of places moved
• Decimal moved right (-) the number of places moved
Physical Quantities
Physical Quantities
Measuring Mass
• Mass is a measure of the amount of
matter in an object. Mass does not depend
on location.
• Weight is a measure of the gravitational
force acting on an object. Weight depends
on location.
• Chemist measure grams or milligrams
Measuring Length and Volume
• Length has the SI unit of meter (m)
• Volume = length x width x height
– Units – m3
– Chemist tend to use milliliters (mL) or Liters
(L)
Measurement and Significant Figures
• When reading a measured value
– All nonzero digits are significant.
– Zero’s
• RULE 1. Zeros in the middle of a number are they are always
significant.
• RULE 2. Zeros at the beginning of a number are not
significant
• RULE 3. Zeros at the end of a number and after the decimal
point are significant.
• RULE 4. Zeros at the end of a number and before an implied
decimal point may or may not be significant. We cannot tell
whether they are part of the measurement or whether they
act only to locate the unwritten but implied decimal point.
– If a decimal point is shown the zero’s are significant
Problem
• Which measurement is expressed to 4
significant figures?
– A.
– B.
– C.
– D.
– E.
0.00423 kg
24.049 cm
1300 K
82,306 m
62.40 g
Performing Problems with
Measurements
• Why are significant figures important?
• How do we convert from one unit to
another?
Rounding Off Numbers
• Often when doing arithmetic on a
calculator, the answer is displayed with
several digits.
– Example - 13.6 / 28
• How many do you keep?
Rounding Off Numbers
•RULE 1. Multiplication or Division
– the answer cannot have more significant
figures than the original number with the
fewest.
Rounding Off Numbers
•RULE 2. Addition or Subtraction
– the answer cannot have more digits
after the decimal point than the original
number with the fewest.
Rounding Off Numbers
• Once you decide how many digits to
retain, the rules for rounding off numbers
are straightforward:
– If the first number dropped is:
• 4 or less – let it rest
• 5 or more – let it score
Problem
• The appropriate number of significant
figures in the result of 15.234 - 15.208 is
– A.
– B.
– C.
– D.
– E.
1
2
3
4
5
Problem
• Select the answer that expresses the result of
this calculation with the correct number of
significant figures.
–
–
–
–
–
A.
B.
C.
D.
E.
13.3568
13.357
13.36
13.4
13
Problem
• The result of (3.8621 × 1.5630) - 5.98 is
properly written as
– A.
– B.
– C.
– D.
– E.
0.06
0.056
0.0565
0.05646
0.056462
Converting a Quantity from One
Unit to Another
• Factor-Label Method:
(Starting quantity) x (Conversion factor) = Equivalent quantity
Converting a Quantity from One
Unit to Another
• What is a conversion unit
– Ratios, fractions, or two measured quantities
that are equivalent
– Equal 1
• The important item in these numbers are
UNITS
Example
• How many kilometers is 26.22 miles?
– STEP 1: Identify the information given.
– STEP 2: Identify the information needed to answer.
– STEP 3: Find the relationship(s) between the known
information and unknown answer, and plan a series of
steps, including conversion factors, for getting from
one to the other.
– STEP 4: Solve the problem.
– BALLPARK CHECK: Make a rough estimate to be
sure the value and the units of your calculated answer
are reasonable.
Problem
• The distance between carbon atoms in
ethylene is 134 picometers. Which of the
following expresses that distance in
meters?
– A.
– B.
– C.
– D.
– E.
1.34 × 10-13 m
1.34 × 10-12 m
1.34 × 10-10 m
1.34 × 10-7 m
1.34 × 10-6 m
Problem
• A dose of medication was prescribed to be
35 microliters. Which of the following
expresses that volume in centiliters?
– A.
– B.
– C.
– D.
– E.
3.5 × 105 cL
3.5 × 104 cL
3.5 cL
3.5 × 10-4 cL
3.5 × 10-3 cL
Problem
• The average distance between the Earth
and the Moon is 240,000 miles. Express
this distance in kilometers.
– A.
– B.
– C.
– D.
– E.
6.1 × 105 km
5.3 × 105 km
3.9 × l05 km
1.5 × 105 km
9.4 × 104 km
Problem
• The speed needed to escape the pull of
Earth's gravity is 11.3 km/s. What is this
speed in mi/h?
– A.
– B.
– C.
– D.
– E.
65,500 mi/h
25,300 mi/h
18,200 mi/h
1,090 mi/h
5.02 × 10-3 mi/h
Measuring Temperature
• 3 scales
– Fahrenheit
– Celsius
– Kelvin
Measuring Temperatures
• Converting Between Temperature Scales
–
–
oF
= (1.8 x oC) + 32
K = oC + 273.15
Problem
• Isopropyl alcohol, commonly known as
rubbing alcohol, boils at 82.4°C. What is
the boiling point in kelvins?
– A.
– B.
– C.
– D.
– E.
387.6 K
355.6 K
323.6 K
190.8 K
-190.8 K
Problem
• Acetic acid boils at 244.2°F. What is its
boiling point in degrees Celsius?
– A.
– B.
– C.
– D.
– E.
382.0°C
167.7°C
153.4°C
117.9°C
103.7°C
Units of Energy and Heat
• Energy: The capacity to do work or supply heat.
– SI units - Joule (J)
– calorie is another unit often used to measure energy.
• One calorie (cal) - the amount of heat necessary to raise the
temperature of 1 g of water by 1°C.
– Calorie – food calorie
• Energy equivalencies
– 4.184 J = 1 cal
– 1000 cal = 1 Cal
– 4.184 kJ = 1 Cal
Units of Energy and Heat
• A Snickers® candy bar contains 280 Calories, of
which the fat content accounts for 120 Calories.
What is the energy of the fat content, in kJ?
–
–
–
–
–
A.
B.
C.
D.
E.
5.0 × 10-1 kJ
29 kJ
5.0 × 102 kJ
1.2 × 103 kJ
5.0 × 105 kJ
Problem
• Natural gas, or methane, is an important fuel.
Combustion of one mole of methane releases
802.3 kilojoules of energy. How much energy
does that represent in kilocalories?
–
–
–
–
–
A.
B.
C.
D.
E.
1.918 × 10-1 kcal
1.918 × 102 kcal
3.360 × 103 kcal
1.918 × 105 kcal
3.360 × 106 kcal
Units of Heat and Energy
• Not all substances are created
equal.
– One calorie raises the
temperature of 1 g of water by
1°C but raises the
temperature of 1 g of iron by
10°C.
• The amount of heat needed to
raise the temperature of 1 g of
a substance by 1°C is called
the specific heat of the
substance (c).
• Specific heat is measured in
units of cal/gC or J/goC
Units of Heat and Energy
• q = mcΔT
– q = heat change
– m = mass
– c = specific heat
Units of Heat and Energy
• Calculate q when 28.6 g of water is heated
from 22.0°C to 78.3°C. (cwater = 4.184
J/goC)
– A.
– B.
– C.
– D.
– E.
0.385 kJ
1.61 kJ
6.74 kJ
9.37 kJ
1.61 × 103 kJ
Problem
• Ethylene glycol, used as a coolant in automotive
engines, has a specific heat capacity of 2.42
J/(goC). Calculate q when 3.65 kg of ethylene
glycol is cooled from 132°C to 85°C.
–
–
–
–
–
A.
B.
C.
D.
E.
-1900 kJ
-420 kJ
-99 kJ
-0.42 kJ
-4.2 × 10-6 kJ
Density
•Density relates the mass of an object to its
volume.
– Units
• grams per cubic centimeter (g/cm3) for solids
• grams per milliliter (g/mL) for liquids.
Density =
Mass (g)
Volume (mL or cm3)
Density
• If the gasoline in a full 20.0 gallon tank
weighs 116 lb, what is the density of
gasoline in g/mL
• How many grams does 1.2 L of water
weigh, if at room temperature water has a
density of 0.9970 g/cm3
Optional Homework
• Text - 2.44, 2.45, 2.46, 2.47, 2.48, 2.50,
2.52, 2.54, 2.56, 2.58, 2.62, 2.64, 2.66,
2.68, 2.70, 2.72, 2.74, 2.78, 2.88, 2.90,
2.96, 2.106
• Chapter 2 Homework - found online