1. Chemistry: Methods and Measurement
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Chapter 1
Chemistry:
Methods and Measurement
Denniston Topping Caret
5th Edition

2. 1.1 The Discovery Process Chemistry - The study of matter…
Matter - Anything that has mass and occupies space
A table
A piece paper
What about air?
Yes, it is matter

3. 1.1 The Discovery Process Chemistry: the study of matter
its chemical and physical properties
the chemical and physical changes it undergoes
the energy changes that accompany those processes
Energy - the ability to do work to accomplish some change
1.1 The Discovery Process

4. MAJOR AREAS OF CHEMISTRY
Biochemistry - the study of life at the molecular level
Organic chemistry - the study of matter containing carbon and hydrogen
Inorganic chemistry - the study of matter containing elements, not organic
Analytic chemistry - analyze matter to determine identity and composition
1.1 The Discovery Process

5. Physical chemistry - attempts to explain the way matter behaves
pharmaceutical industry
public health
1.1 The Discovery Process
CHEMISTRY
food science
medical practitioners
forensic sciences

6. 1.1 The Discovery Process THE SCIENTIFIC METHOD
The scientific method - a systematic approach to the discovery of new information
Characteristics of the scientific process
Observation
Formulation of a question
Pattern recognition
Developing theories
Experimentation
Summarizing information
1.1 The Discovery Process

7. 1.1 The Discovery Process

8. Models in Chemistry 1.1 The Discovery Process
To aid in understanding of a chemical unit or system
a model is often used
good models are based on everyday experience
Ball and stick methane model
color code balls
sticks show attractive forces holding atoms together
1.1 The Discovery Process

9. 1.2 Matter and Properties Properties - characteristics of matter
chemical vs. physical
Three states of matter
1. gas - particles widely separated, no definite shape or volume solid
2. liquid - particles closer together, definite volume but no definite shape
3. solid - particles are very close together, define shape and definite volume

10. Three States of Water
(a) Solid (b) Liquid (c) Gas

11. Comparison of the Three Physical States
1.2 Matter and Properties

12. Physical property - is observed without changing the composition or identity of a substance
Physical change - produces a recognizable difference in the appearance of a substance without causing any change in its composition or identity
conversion from one physical state to another
melting an ice cube
1.2 Matter and Properties

13. Separation by Physical Properties
Magnetic iron is separated from other nonmagnetic substances, such as sand. This property is used as a large-scale process in the recycling industry.

14. Chemical property - result in a change in composition and can be observed only through a chemical reaction
Chemical reaction (chemical change) - a process of rearranging, removing, replacing, or adding atoms to produce new substances
1.2 Matter and Properties
hydrogen + oxygen  water
reactants
products

15. Classify the following as either a chemical or physical property:
Color
Flammability
Hardness
Odor
Taste
1.2 Matter and Properties

16. Classify the following as either a chemical or physical change:
Boiling water becomes steam
Butter turns rancid
Burning of wood
Mountain snow pack melting in spring
Decay of leaves in winter
1.2 Matter and Properties

17. Intensive properties - a property of matter that is independent of the quantity of the substance
Density
Specific gravity
Extensive properties - a property of matter that depends on the quantity of the substance
Mass
Volume
1.2 Matter and Properties

18. Classification of Matter
1.2 Matter and Properties
Pure substance - a substance that has only one component
Mixture - a combination of two or more pure substances in which each substance retains its own identity, not undergoing a chemical reaction

19. Classification of Matter
1.2 Matter and Properties
Element - a pure substance that cannot be changed into a simpler form of matter by any chemical reaction
Compound - a substance resulting from the combination of two or more elements in a definite, reproducible way, in a fixed ratio

20. Classification of Matter
1.2 Matter and Properties
Mixture - a combination of two or more pure substances in which each substance retains its own identity
Homogeneous - uniform composition, particles well mixed, thoroughly intermingled
Heterogeneous – nonuniform composition, random placement

21. Classes of Matter
1.2 Matter and Properties

22. 1.3 Measurement in Chemistry
Data, Results, and Units
Data - each piece is an individual result of a single measurement or observation
mass of a sample
temperature of a solution
Results - the outcome of the experiment
Data and results may be identical, however usually related data are combined to generate a result
Units - the basic quantity of mass, volume or whatever quantity is being measured
A measurement is useless without its units

23. 1.3 Measurement in Chemistry
English and Metric Units
English system - a collection of functionally unrelated units
Difficult to convert from one unit to another
1 foot = 12 inches = 0.33 yard = 1/5280 miles
Metric System - composed of a set of units that are related to each other decimally, systematic
Units relate by powers of tens
1 meter = 10 decimeters = 100 centimeters = 1000 millimeters
1.3 Measurement in Chemistry

24. 1.3 Measurement in Chemistry
Basic Units of the Metric System
Mass gram g
Length meter m
Volume liter L
1.3 Measurement in Chemistry
Basic units are the units of a quantity without any metric prefix

25. 1.3 Measurement in Chemistry

26. 1.3 Measurement in Chemistry
UNIT CONVERSION
You must be able to convert between units
within the metric system
between the English system and metric system
The method used for conversion is called the Factor-Label Method or Dimensional Analysis
1.3 Measurement in Chemistry
!!!!!!!!!!! VERY IMPORTANT !!!!!!!!!!!

27. 1.3 Measurement in Chemistry
Let your units do the work for you by simply memorizing connections between units.
For example: How many donuts are in one dozen?
1.3 Measurement in Chemistry
We say: “Twelve donuts are in a dozen.”
Or: 12 donuts = 1 dozen donuts
What does any number divided by itself equal?
ONE!

28. 1.3 Measurement in Chemistry
This fraction is called a unit factor
1.3 Measurement in Chemistry
What does any number times one equal?
That number
Multiplication by a unit factor does not change the amount – only the unit

29. 1.3 Measurement in Chemistry
We use these two mathematical facts to use the factor label method
a number divided by itself = 1
any number times one gives that number back
Example: How many donuts are in 3.5 dozen?
You can probably do this in your head but try it using the Factor-Label Method.
1.3 Measurement in Chemistry

30. 1.3 Measurement in Chemistry
Start with the given information...
3.5 dozen
= 42 donuts
1.3 Measurement in Chemistry
Then set up your unit factor...
See that the units cancel...
Then multiply and divide all numbers...

31. Common English System Units
1.3 Measurement in Chemistry
Convert 12 gallons to units of quarts

32. Intersystem Conversion Units
1.3 Measurement in Chemistry
Convert 4.00 ounces to kilograms

33. 1.3 Measurement in Chemistry
Practice Unit Conversion
1. Convert 5.5 inches to millimeters
2. Convert 50.0 milliliters to pints
3. Convert 1.8 in2 to cm2
1.3 Measurement in Chemistry

34. 1.4 Significant Figures and Scientific Notation
Information-bearing digits or figures in a number are significant figures
The measuring devise used determines the number of significant figures a measurement has
The amount of uncertainty associated with a measurement is indicated by the number of digits or figures used to represent the information

35. 1.4 Significant Figures and Scientific Notation
Significant figures - all digits in a number representing data or results that are known with certainty plus one uncertain digit

36. Recognition of Significant Figures
All nonzero digits are significant
7.314 has four significant digits
The number of significant digits is independent of the position of the decimal point
73.14 also has four significant digits
Zeros located between nonzero digits are significant
has five significant digits
1.4 Significant Figures and Scientific Notation

37. Use of Zeros in Significant Figures
Zeros at the end of a number (trailing zeros) are significant if the number contains a decimal point.
4.70 has three significant digits
Trailing zeros are insignificant if the number does not contain a decimal point.
100 has one significant digit; 100. has three
Zeros to the left of the first nonzero integer are not significant.
has two significant digits
1.4 Significant Figures and Scientific Notation

38. 1.4 Significant Figures and Scientific Notation
Determining Significant
Figures
How many significant figures are in the following?
3.400
3004
300.
1.4 Significant Figures and Scientific Notation

39. 1.4 Significant Figures and Scientific Notation
Used to express very large or very small numbers easily and with the correct number of significant figures
Represents a number as a power of ten
Example:
4,300 = 4.3 x 1,000 = 4.3 x 103
1.4 Significant Figures and Scientific Notation

40. 1.4 Significant Figures and Scientific Notation
To convert a number greater than 1 to scientific notation, the original decimal point is moved x places to the left, and the resulting number is multiplied by 10x
The exponent x is a positive number equal to the number of places the decimal point moved
5340 = 5.34 x 104
What if you want to show the above number has four significant figures?
= x 104
1.4 Significant Figures and Scientific Notation

41. 1.4 Significant Figures and Scientific Notation
To convert a number less than 1 to scientific notation, the original decimal point is moved x places to the right, and the resulting number is multiplied by 10-x
The exponent x is a negative number equal to the number of places the decimal point moved
= 5.34 x 10-2
1.4 Significant Figures and Scientific Notation

42. 1.4 Significant Figures and Scientific Notation
Types of Uncertainty
Error - the difference between the true value and our estimation
Random
Systematic
Accuracy - the degree of agreement between the true value and the measured value
Precision - a measure of the agreement of replicate measurements
1.4 Significant Figures and Scientific Notation

43. Significant Figures in Calculation of Results
Rules for Addition and Subtraction
The result in a calculation cannot have greater significance than any of the quantities that produced the result
Consider:
liters
liters
liters
liters
1.4 Significant Figures and Scientific Notation
correct answer liters

44. Rules for Multiplication and Division
The answer can be no more precise than the least precise number from which the answer is derived
The least precise number is the one with the fewest significant figures
1.4 Significant Figures and Scientific Notation
Which number has the fewest significant figures? x 103 has only 2
The answer is therefore, 3.0 x 10-8

45. Exact and Inexact Numbers
Inexact numbers have uncertainty by definition
Exact numbers are a consequence of counting
A set of counted items (beakers on a shelf) has no uncertainty
Exact numbers by definition have an infinite number of significant figures
1.4 Significant Figures and Scientific Notation

46. Rules for Rounding Off Numbers
When the number to be dropped is less than 5 the preceding number is not changed
When the number to be dropped is 5 or larger, the preceding number is increased by one unit
Round the following number to 3 significant figures: x 104
1.4 Significant Figures and Scientific Notation
=3.35 x 104

47. How Many Significant Figures?
Round off each number to 3 significant figures:
61.40
6.171
1.4 Significant Figures and Scientific Notation

48. 1.5 Experimental Quantities
Mass - the quantity of matter in an object
not synonymous with weight
standard unit is the gram
Weight = mass x acceleration due to gravity
Mass must be measured on a balance (not a scale)

49. 1.5 Experimental Quantities
Units should be chosen to suit the quantity described
A dump truck is measured in tons
A person is measured in kg or pounds
A paperclip is measured in g or ounces
An atom?
For atoms, we use the atomic mass unit (amu)
1 amu = x g
1.5 Experimental Quantities

50. 1.5 Experimental Quantities
Length - the distance between two points
standard unit is the meter
long distances are measured in km
distances between atoms are measured in nm, 1 nm = 10-9 m
Volume - the space occupied by an object
standard unit is the liter
the liter is the volume occupied by 1000 grams of water at 4 oC
1 mL = 1/1000 L = 1 cm3
1.5 Experimental Quantities

51. 1.5 Experimental Quantities
The milliliter (mL) and the cubic centimeter (cm3) are equivalent
1.5 Experimental Quantities

52. 1.5 Experimental Quantities
Time
metric unit is the second
Temperature - the degree of “hotness” of an object
1.5 Experimental Quantities

53. Conversions Between Fahrenheit and Celsius
1.5 Experimental Quantities
1. Convert 75oC to oF
2. Convert -10oF to oC
1. Ans. 167 oF 2. Ans. -23oC

54. Kelvin Temperature Scale
The Kelvin scale is another temperature scale.
It is of particular importance because it is directly related to molecular motion.
As molecular speed increases, the Kelvin temperature proportionately increases.
1.5 Experimental Quantities
K = oC + 273

55. 1.5 Experimental Quantities
Energy
Energy - the ability to do work
kinetic energy - the energy of motion
potential energy - the energy of position (stored energy)
Energy is also categorized by form:
light
heat
electrical
mechanical
chemical
1.5 Experimental Quantities

56. Characteristics of Energy
Energy cannot be created or destroyed
Energy may be converted from one form to another
Energy conversion always occurs with less than 100% efficiency
All chemical reactions involve either a “gain” or “loss” of energy
1.5 Experimental Quantities

57. 1.5 Experimental Quantities
Units of Energy
Basic Units:
calorie or joule
1 calorie (cal) = joules (J)
A kilocalorie (kcal) also known as the large Calorie. This is the same Calorie as food Calories.
1 kcal = 1 Calorie = 1000 calories
1 calorie = the amount of heat energy required to increase the temperature of 1 gram of water 1oC.
1.5 Experimental Quantities

58. 1.5 Experimental Quantities
Concentration
Concentration:
the number of particles of a substance
the mass of those particles
that are contained in a specified volume
Often used to represent the mixtures of different substances
Concentration of oxygen in the air
Pollen counts
Proper dose of an antibiotic
1.5 Experimental Quantities

59. Density and Specific Gravity
the ratio of mass to volume
an extensive property
use to characterize a substance as each substance has a unique density
Units for density include:
g/mL
g/cm3
g/cc
1.5 Experimental Quantities

60. 1.5 Experimental Quantities
cork
1.5 Experimental Quantities
water
brass nut
liquid mercury

62. Calculating the Density of a Solid
2.00 cm3 of aluminum are found to weigh 5.40g. Calculate the density of aluminum in units of g/cm3.
Use the formula
Substitute our values
5.40 g
2.00 cm3
= g / cm3
1.5 Experimental Quantities

63. 1.5 Experimental Quantities
Density Calculations:
Air has a density of g/mL. What is the mass of 6.0-L sample of air?
Calculate the mass in grams of 10.0 mL if mercury (Hg) if the density of Hg is 13.6 g/mL.
Calculate the volume in milliliters, of a liquid that has a density of 1.20 g/mL and a mass of 5.00 grams.
1.5 Experimental Quantities

64. 1.5 Experimental Quantities
Specific Gravity
Values of density are often related to a standard
Specific gravity - the ratio of the density of the object in question to the density of pure water at 4oC
Specific gravity is a unitless term because the 2 units cancel
Often the health industry uses specific gravity to test urine and blood samples
1.5 Experimental Quantities