Measurements in Chemistry

1
Chapter 2

• Measurements in Chemistry

2
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?

3
Physical Quantities

• Physical quantities measured physical
  properties
• Number
• Unit

4
Proper Measurements

• Number
• Reflect the certainty to which the measurement
  was made
• Unit
• Represent the type of measurement made
• Mass
• Volume
• Length

5
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

6
Measurement and Significant Figures

• The total number of digits used to express such a
  measurement is the number of significant figures.

7
Measurement and Significant Figures
Instrument used directly impacts the certainty of
the measurement and hence the of significant
figures that can be reported.
8
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

9
Physical Quantities
10
Physical Quantities
11
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

12
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)

13
Measurement and Significant Figures

• When reading a measured value
• All nonzero digits are significant.
• Zeros
• 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 zeros are
  significant

14
Problem

• Which measurement is expressed to 4 significant
  figures? 
• A.  0.00423 kg
• B.  24.049 cm
• C.  1300 K
• D.  82,306 m
• E.  62.40 g

15
Performing Problems with Measurements

• Why are significant figures important?
• How do we convert from one unit to another?

16
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?

17
Rounding Off Numbers

• RULE 1. Multiplication or Division
• the answer cannot have more significant figures
  than the original number with the fewest.

18
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.

19
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

20
Problem

• The appropriate number of significant figures in
  the result of 15.234 - 15.208 is 
• A.  1
• B.  2
• C.  3
• D.  4
• E.  5

21
Problem

• Select the answer that expresses the result of
  this calculation with the correct number of
  significant figures.   
• A.  13.3568
• B.  13.357
• C.  13.36
• D.  13.4
• E.  13

22
Problem

• The result of (3.8621 1.5630) - 5.98 is
  properly written as 
• A.  0.06
• B.  0.056
• C.  0.0565
• D.  0.05646
• E.  0.056462

23
Converting a Quantity from One Unit to Another

• Factor-Label Method

(Starting quantity) x (Conversion factor)
Equivalent quantity
24
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

25
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.

26
Problem

• The distance between carbon atoms in ethylene is
  134 picometers. Which of the following expresses
  that distance in meters? 
• A.  1.34 10-13 m
• B.  1.34 10-12 m
• C.  1.34 10-10 m
• D.  1.34 10-7 m
• E.  1.34 10-6 m

27
Problem

• A dose of medication was prescribed to be 35
  microliters. Which of the following expresses
  that volume in centiliters? 
• A.  3.5 105 cL
• B.  3.5 104 cL
• C.  3.5 cL
• D.  3.5 10-4 cL
• E.  3.5 10-3 cL

28
Problem

• The average distance between the Earth and the
  Moon is 240,000 miles. Express this distance in
  kilometers. 
• A.  6.1 105 km
• B.  5.3 105 km
• C.  3.9 l05 km
• D.  1.5 105 km
• E.  9.4 104 km

29
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.  65,500 mi/h
• B.  25,300 mi/h
• C.  18,200 mi/h
• D.  1,090 mi/h
• E.  5.02 10-3 mi/h

30
Measuring Temperature

• 3 scales
• Fahrenheit
• Celsius
• Kelvin

31
Measuring Temperatures

• Converting Between Temperature Scales
• oF (1.8 x oC) 32
• K oC 273.15

32
Problem

• Isopropyl alcohol, commonly known as rubbing
  alcohol, boils at 82.4C. What is the boiling
  point in kelvins? 
• A.  387.6 K
• B.  355.6 K
• C.  323.6 K
• D.  190.8 K
• E.  -190.8 K

33
Problem

• Acetic acid boils at 244.2F. What is its boiling
  point in degrees Celsius? 
• A.  382.0C
• B.  167.7C
• C.  153.4C
• D.  117.9C
• E.  103.7C

34
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 1C.
• Calorie food calorie
• Energy equivalencies
• 4.184 J 1 cal
• 1000 cal 1 Cal
• 4.184 kJ 1 Cal

35
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.  5.0 10-1 kJ
• B.  29 kJ
• C.  5.0 102 kJ
• D.  1.2 103 kJ
• E.  5.0 105 kJ

36
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.  1.918 10-1 kcal
• B.  1.918 102 kcal
• C.  3.360 103 kcal
• D.  1.918 105 kcal
• E.  3.360 106 kcal

37
Units of Heat and Energy

• Not all substances are created equal.
• One calorie raises the temperature of 1 g of
  water by 1C but raises the temperature of 1 g of
  iron by 10C.
• The amount of heat needed to raise the
  temperature of 1 g of a substance by 1C is
  called the specific heat of the substance (c).
• Specific heat is measured in units of cal/g?C or
  J/goC

38
Units of Heat and Energy

• q mc?T
• q heat change
• m mass
• c specific heat

39
Units of Heat and Energy

• Calculate q when 28.6 g of water is heated from
  22.0C to 78.3C. (cwater 4.184 J/goC)
• A.  0.385 kJ
• B.  1.61 kJ
• C.  6.74 kJ
• D.  9.37 kJ
• E.  1.61 103 kJ

40
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 132C to 85C. 
• A.  -1900 kJ
• B.  -420 kJ
• C.  -99 kJ
• D.  -0.42 kJ
• E.  -4.2 10-6 kJ

41
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.

Mass (g)
Density
Volume (mL or cm3)
42
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

43
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