This document provides an overview of key concepts in chemistry, including: 1. Chemistry is the study of matter, its properties, and the changes it undergoes. The scientific method is used to systematically study matter through observation, questioning, experimentation and summarization. 2. Matter can exist in three physical states - solid, liquid, and gas. It also has physical and chemical properties and can undergo physical or chemical changes. 3. The metric system is the standard system of measurement in chemistry. It relates units decimally and is easier for conversion than the English system. Careful measurement and use of appropriate units is important in chemistry.

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Chapter 1
Chemistry:
Methods and Measurement
Deniston
Topping
Caret
7th 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

4. 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
1. Observation
2. Formulation of a question
3. Pattern recognition
4. Developing theories
5. Experimentation
6. Summarizing information

5. 1.1 The Discovery Process

6. 1.1 The Discovery Process
Models in Chemistry
• 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

7. 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

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

9. 1.2 Matter and Properties Comparison of the Three
Physical States

10. 1.2 Matter and Properties • 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

11. 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.

12. 1.2 Matter and Properties
• 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
hydrogen + oxygen  water
reactants products

13. 1.2 Matter and Properties Classify the following as either a
c
h
e
m
i
c
a
l
o
r
p
h

14. 1.2 Matter and Properties Classify the following as either a
chemical or physical change:
a. Boiling water becomes steam
b. Butter turns rancid
c. Burning of wood
d. Mountain snow pack melting in
spring
e. Decay of leaves in winter

15. 1.2 Matter and Properties
• 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

16. 1.2 Matter and Properties Classification of Matter
• 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

17. 1.2 Matter and Properties Classification of Matter
• 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

18. 1.2 Matter and Properties Classification of Matter
• 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

19. 1.2 Matter and Properties
Classes of Matter

20. 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

21. English and Metric Units
• English system - a collection of
1.3 Measurement in
functionally unrelated units
– Difficult to convert from one unit to
Chemistry
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

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

23. 1.3 Measurement in
Chemistry

24. UNIT CONVERSION
1.3 Measurement in
• You must be able to convert between
Chemistry
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
!!!!!!!!!!! VERY IMPORTANT !!!!!!!!!!!

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

26. 12 donuts
=1
1.3 Measurement in
1 dozen
Chemistry
• This fraction is called a unit factor
• What does any number times one
equal?
• That number
• Multiplication by a unit factor does
not change the amount – only the unit

27. • We use these two mathematical facts to
1.3 Measurement in
use the factor label method
– a number divided by itself = 1
Chemistry
– 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.

28. 1.3 Measurement in
Start with the given information...
12 donuts
3.5 dozen × = 42 donuts
Chemistry
1 dozen
Then set up your unit factor...
See that the units cancel...
Then multiply and divide all numbers...

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

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

31. 1.3 Measurement in
Chemistry
1. Convert 5.5 inches to millimeters
2. Convert 50.0 milliliters to pints
3. Convert 1.8 in2 to cm2

32. 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

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

34. and Scientific Notation Recognition of Significant Figures
1.4 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
• 60.052 has five significant digits

35. Use of Zeros in Significant
and Scientific Notation
1.4 Significant Figures
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.
• 0.0032 has two significant digits

36. and Scientific Notation
1.4 Significant Figures
How many significant figures are in
the following?
1. 3.400
2. 3004
3. 300.
4. 0.003040

37. and Scientific Notation Scientific Notation
1.4 Significant Figures
• 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

38. and Scientific Notation • To convert a number greater than 1 to
1.4 Significant Figures
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?
= 5.340 x 104

39. and Scientific Notation
1.4 Significant Figures
• 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
0.0534 = 5.34 x 10-2

40. and Scientific Notation Types of Uncertainty
1.4 Significant Figures
• 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

41. Significant Figures in Calculation of
and Scientific Notation Results
1.4 Significant Figures
Rules for Addition and Subtraction
• The result in a calculation cannot have greater
significance than any of the quantities that
produced the result
• Consider:
37.68 liters
6.71862 liters
108.428 liters
152.82662 liters
correct answer 152.83 liters

42. Rules for Multiplication and Division
and Scientific Notation
1.4 Significant Figures
• 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
4.2 × 103 (15.94)
= 2.9688692 ×10 −8 (on calculator)
2.255 ×10 − 4
Which number has the fewest
significant figures? 4.2 x 103 has only 2
The answer is therefore, 3.0 x 10-8

43. and Scientific Notation Exact and Inexact Numbers
1.4 Significant Figures
• 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

44. Rules for Rounding Off Numbers
and Scientific Notation
1.4 Significant Figures
• 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: 3.34966 x 104
=3.35 x 104

45. and Scientific Notation
1.4 Significant Figures
How Many Significant Figures?
Round off each number to 3 significant
figures:
1. 61.40
2. 6.171
3. 0.066494

46. 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)

47. 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 = 1.661 x 10-24 g

48. 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

49. 1.5 Experimental Quantities
The milliliter
(mL) and the
cubic centimeter
(cm3) are
equivalent

50. 1.5 Experimental Quantities • Time
- metric unit is the second
• Temperature - the degree of “hotness”
of an object

51. 1.5 Experimental Quantities Conversions Between Fahrenheit
and Celsius
o
F - 32
o
C=
1.8
o
F = 1.8 ×( C) + 32
o
1. Convert 75oC to oF
2. Convert -10oF to oC
1. Ans. 167 oF

52. 1.5 Experimental Quantities
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.
K = oC + 273

53. 1.5 Experimental Quantities
•
Energy

54. 1.5 Experimental Quantities 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

55. 1.5 Experimental Quantities Units of Energy
• Basic Units:
• calorie or joule
• 1 calorie (cal) = 4.184 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.

56. 1.5 Experimental Quantities
Density and Specific Gravity
• Density
– 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 mass m
d= =
• g/cc volume V

57. 1.5 Experimental Quantities
cork
water
brass nut
liquid mercury

59. 1.5 Experimental Quantities 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.
mass m
– Use the formula d= =
volume V
– Substitute our values
5.40 g
2.00 cm3
= 2.70 g / cm3

60. 1.5 Experimental Quantities
Air has a density of 0.0013 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.

61. 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
density of object (g/mL)
specific gravity =
density of water (g/mL)