1. 1 Chemistry: Methods and Measurement GENERAL CHEMISTRY
General, Organic and Biochemistry 7th Edition
Copyright The McGraw-Hill Companies, Inc. Permission required for reproduction or display.

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. 1.1 The Discovery Process THE SCIENTIFIC METHOD
The scientific method - a systematic approach to the discovery of new information 2
Characteristics of the scientific process
Observation
Formulation of a question
Pattern recognition
Developing theories
Experimentation
Summarizing information
1.1 The Discovery Process

5. 1.1 The Discovery Process

6. Three States of Matter 1.2 Matter and Properties
gas - particles widely separated, no definite shape or volume
2. liquid - particles closer together, definite volume but no definite shape
3. solid - particles are very close together, define shape and definite volume 4
1.2 Matter and Properties

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

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

9. 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 5
1.2 Matter and Properties

10. 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.

11. 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 5
1.2 Matter and Properties
hydrogen + oxygen  water
reactants
products

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

13. 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

14. 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 is dependent on the quantity of the substance
Mass
Volume 6
1.2 Matter and Properties

15. 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 7

16. 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 7

17. 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 7

18. Classes of Matter
1.2 Matter and Properties

19. 1.3 Significant Figures and Scientific Notation
Information-bearing digits in a number are significant figures
The measuring device 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 8

20. 1.3 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

21. 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
1.3 Significant Figures and Scientific Notation

22. 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
1.3 Significant Figures and Scientific Notation

23. How many significant figures are in the following?
3.400
3004
300.
1.3 Significant Figures and Scientific Notation

24. 1.3 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  1,000 = 4.3  103
1.3 Significant Figures and Scientific Notation

25. 1.3 Significant Figures and Scientific Notation
For numbers greater than 1, 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  103
What if you want to show the above number has four significant figures?
=  103
1.3 Significant Figures and Scientific Notation

26. 1.3 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  10–2
1.3 Significant Figures and Scientific Notation

27. Convert the following to 2 SF using scientific notation…
0.0024
48.20
224
0.0180

28. 1.3 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
1.3 Significant Figures and Scientific Notation

29. Exact and Inexact Numbers
Inexact numbers have uncertainty by definition (any measured value)
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
1.3 Significant Figures and Scientific Notation

30. 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:  104
1.3 Significant Figures and Scientific Notation
=3.35  104

31. Round off each number to three significant figures:
61.40
6.171
1.3 Significant Figures and Scientific Notation

32. 1.4 Units and Unit Conversion
Units - the basic quantity of mass, volume or whatever is being measured
A measurement is useless without units 9
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 ten

33. 1.4 Units and Unit Conversion
Basic Units of the Metric System
Mass gram g
Length meter m
Volume liter L
1.4 Units and Unit Conversion
Basic units are the units of a quantity without any metric prefix.

34. 1.4 Units and Unit Conversion

36. 1.4 Units and 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 (FLM) or Dimensional Analysis 9
1.4 Units and Unit Conversion
!!!!!!!!!!! VERY IMPORTANT !!!!!!!!!!!

37. 1.4 Units and Unit Conversion
Let your units do the work for you by simply memorizing connections between units.
For example: How many donuts are in one dozen?
1.4 Units and Unit Conversion
We say: “Twelve donuts are in a dozen.”
Or: 12 donuts = 1 dozen donuts
What does any number divided by itself equal?
ONE!

38. 1.4 Units and Unit Conversion
This fraction is called a unit or conversion factor
1.4 Units and Unit Conversion
Multiplication by a unit factor does not change the amount – only the unit

39. 1.4 Units and Unit Conversion
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:
1.4 Units and Unit Conversion

40. 1.4 Units and Unit Conversion
Start with the given information...
3.5 dozen
= 42 donuts
1.4 Units and Unit Conversion
Then set up your unit factor...
See that the units cancel...
Then multiply and divide all numbers...

41. Common English System Units
1.4 Units and Unit Conversion
Convert 12 gallons to quarts
Convert 175 lbs to ounces
Convert 10,000 ft to miles

42. Intersystem Conversion Units
1.4 Units and Unit Conversion
Convert 4.00 ounces to kilograms
Convert 175 lbs to kgs
Convert 100 yards to meters

43. Practice Unit Conversions
Convert 12 gallons to quarts
Convert 10 cm to meters
Convert 360 feet to miles
Convert 60 miles/hr to m/s
5. Convert 1.8 in2 to cm2
1.4 Units and Unit Conversion

44. The speed of light is 299,792,458 m/s  convert this to miles per hour.
The speed of sound is m/s  convert this to miles per hour
(Use: 1 mile = 1.6 km; 1 km = 1000 m; 60 s = 1 min; 60 min = 1 hour)

45. 1.5 Experimental Quantities
Mass - the quantity of matter in an object
not synonymous with weight
standard unit is the gram
Weight = mass  acceleration due to gravity
Mass must be measured on a balance (not a scale)

46. 1.5 Experimental Quantities
Units should be chosen to best 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
For atoms, we use the atomic mass unit (amu)
1 amu =  g
1.5 Experimental Quantities

47. 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

48. 1.5 Experimental Quantities
The milliliter (mL) and the cubic centimeter (cm3) are equivalent
1.5 Experimental Quantities

49. 1.5 Experimental Quantities
Time
metric unit is the second
Temperature - the degree of “hotness” of an object 10
1.5 Experimental Quantities

50. Conversions Between Fahrenheit and Celsius
1.5 Experimental Quantities
Convert 75oC to oF
Convert -10oF to oC

51. 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
Convert 20ºC to K

52. Density and Specific Gravity
the ratio of mass to volume
an intensive property
use to characterize a substance as each substance has a unique density
Units for density include:
g/mL
g/cm3
g/cc 11
1.5 Experimental Quantities

53. 1.5 Experimental Quantities
cork
1.5 Experimental Quantities
water
brass nut
liquid mercury

55. 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.
1.5 Experimental Quantities

56. 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

57. 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