Chemistry, The Central Science , 10th edition
Theodore L. Brown; H. Eugene LeMay, Jr.; and Bruce E. Bursten
John D. Bookstaver
St. Charles Community College
St. Peters, MO

2006, Prentice Hall
Matter
And
Measurement

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Matter
And
Measurement

Let the journey Begin…
• Chemistry the study of matter and the changes it undergoes.
• Matter – anything that has mass and takes up space
Matter
And
Measurement

Overview of
Matter
Atoms
Elements
Compounds
• Atoms are the building blocks of matter.
• Each element is made of the same kind of atom.
• A compound is made from atoms of two or more of elements that are chemically bonded together

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Matter
And
Measurement

States of Matter
• Solid - has definite shape and volume
• Liquid - definite volume; takes shape of its container
• Gas - no definite volume or shape; takes the shape of its container and is easily compressed/ expanded

Pure Substances:
• Matter that does not vary from sample to sample
• Have distinct or characteristic properties
• Have a constant composition
• All pure substances are either:
– Elements
– Compounds
Matter
And
Measurement

Elements:
• Made up of only one type of atom
• Cannot be decomposed into smaller substances

Compounds:
• Composed of two or more elements
• Ex: Water (H & O), salt (Na & Cl),
• Can be broken down into their elemental particles only by chemical means

Law of Constant Composition:
• A pure compound is always made of the exact same elemental composition (by mass)
• aka - Law of Definite Proportions
• Ex: Water (H
2
O) is always in the ratio of
1 oxygen : 2 hydrogen

Mixtures:
• Physical combination of many substances
• Composition can vary (variable)
• Each substance in a mixture retains its own chemical identity & properties
― allows the mixture to be separated by physical means into its component substances
• 2 types:
― Heterogeneous
― Homogenous

Mixtures: ( con’t)…
• Heterogeneous Mixture
― Not uniform in appearance, composition, or properties
― Appears as two or more distinct phases
― Ex: sand, chocolate chip cookies
• Homogenous Mixture (aka: solution)
― Uniform throughout
― appears as one phase
―Ex: air, saltwater, iced tea, Kool-Aid

Classification of Matter

Classification of Matter

Classification of Matter

Classification of Matter

Classification of Matter

Classification of Matter

Classification of Matter

Classification of Matter

Classification of Matter

Classification of Matter

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Properties of Matter
• Physical Properties :
– Observed without changing the substance into a different substance.
• Boiling point, density, mass, color, volume, etc.
• Chemical Properties:
– Only observed when the substance is changed into a different substance.
• Flammability, corrosiveness, reactivity with acid, etc.

Properties of Matter ( con’t)
• Intensive Properties:
– Independent of the amount/quantity of matter present.
– Only depends on the identity of the substance
• Density, boiling point, color, state of matter, etc.
• Extensive Properties:
– Dependent upon the amount of matter present.
• Mass, volume, energy, etc.

Changes of Matter
• Physical Changes:
– Changes in matter that do NOT change the composition of a substance.
• Changes of state, temperature, volume, etc.
• Chemical Changes (aka: reaction = rxn):
– Changes where the original substances
(reactants) are converted to all new & different substances (products).
• Combustion, oxidation, decomposition, etc.

Chemical vs. Physical
• Ask, “Can I somehow get the original substance back?”
– Yes = phys. change
– No = chem. change

Separation of Mixtures:
• Components of a mixture retain their individual properties
― Ex) magnetism, color, density
• A mixture can be physically separated based on these properties
― Each component remains chemically unchanged

Mixtures:
• Physical combination of 2 or more substances
• Composition is variable
• Each substance in a mixture retains its own chemical identity & properties
― allows the mixture to be separated by physical means into its component substances (by exploiting differences in properties)
• 2 types of mixtures:
― Heterogeneous = 2+ phases visible
― Homogenous = uniform (appears as 1 phase)

3 Methods to
Separate a Mixture:

Filtration :
• Separates solid substances from liquids (and solutions)
• Method is based on separation due to differences in the size of component particles in the mixture

Distillation :
Separates a homogeneous mixture on the basis of differences in boiling point.

Chromatography :
Separates on the basis of differences in the solubility of the components in the mixture

The Scientific Method
A systematic approach to solving problems
(textbook pg. 13)
Theory of
Evolution
Collect data
Analyze data
Tentative Generally explanation explains causes that can be of phenomena
TESTED by experimentation and can be used to make predictions

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Qualitative vs. Quantitative
• Qualitative:
– Observations which reflect a physical description of data
– Ex) color, odor, texture, state of matter
• Quantitative:
– Measurements that include numerical quantities & data
– Mass, volume, density, length, time

SI Units
• Système International d’Unités
• Base unit for each quantity

Derived Units:
• Two or more SI base units combined
Example: SI UNITS
Density = = volume kg m 3
• How would you determine the density of a liquid?
… A solid?

• Format : M x 10 n
– M is a # between 1-9
– n is an integer
• Insert an understood decimal point
2.5 x 10 9
• Decide where the decimal point needs to end up
• Count the # of places that the decimal point moved ( n )
• #s SMALLER than 1 = negative exponent, 
– 0.00000765 = 7.65 x 10 -6
• #s BIGGER than 1 = positive exponent, 
– 2, 340,000, 000 = 2.34 x 10 9

Temperature:
• Two Definitions:
1) A measure of the average kinetic energy of particles in a sample
2) A measure of how much heat is in an object

Temperature Conversions
• SI base unit = Kelvin
K =

C + 273.15

C = (

F − 32)
1.8

F = (1.8)(

C) + 32

Metric System
• Prefixes convert the base units into units that are appropriate for the item being measured.

• Prefixes convert the base units into units that are appropriate for the item being measured.

Metric Conversions
• The is a way to SHOW your work & units
Ex #1) Covert 945 mL to dL
945 mL x 1 dL = 9.45 dL
100 mL
Ex #2) Convert 0.54 km to nm
0.54 km x 1000 m x 10 9 nm = 5.4 x 10 11 nm
1 km 1 m

Density
• Density = mass volume
Volume
• Most common units:
L , mL, cm 3
• 1 mL = 1 cm 3
• 1 L = 1 dm 3

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Accuracy vs. Precision
• Accuracy - how close, a measurement is to the true or known value.
• Precision - several measurements with values very close to one other.

Percent Error Calculation
% error = experimental – theoretical x 100 theoretical
• Experimental = the value determined from lab data
• Theoretical = the standard, true, or accepted value

Exact vs. Inexact
• Exact #s:
– conversion factors & counting
Ex) There are exactly 5280 ft in 1 mile
“ “ “ 12 eggs in 1 dozen
“ “ “ 2.54 cm in 1inch
• Inexact #s:
– Measurements from equipment w/limitations
Ex) Mass using balance, volume using graduated cylinder

Uncertainty in Measurements
Different pieces of equipment have different uses and different degrees of accuracy.

Uncertainty in Measurements
What is the unit of volume here that can be measured with absolute certainty?
 43 mL … (last volume level marked on equipment)
All measurements must be reported with the last digit being uncertain (1 degree of uncertainty)
 Report the measurement as 43.0 mL (or 43.1 mL)

Significant Figures
• The term significant figures refers to digits that were “measured” or inexact #s .
• We pay attention to sig. figs. so we do not overstate the accuracy of our calculations

Rules for Sig. Figs
• All non-zero digits (1-9) are significant
• Zeroes between any non-zero digits are significant (ex: 101)
• Zeroes at the beginning of a number, before any non-zero digits, are
NEVER significant (ex: 0.082)
• Zeroes at the end of a number are significant ONLY if the number contains a decimal point (ex: 35.0 vs. 350 vs. 350. )

Practice with Sig. Figs
1) 0.0034567 g
2) 2.97 mL
3) 6.700 x 10 3 m
5 sig figs
3 sig figs
4 sig figs
6 sig figs 4) 598.300 K
5) 78,000 mg
6) 9020 cm
2 sig figs
3 sig figs

Calculations with
Significant Figures
• Addition or Subtraction:
– Round the answer off to the least number of decimal places
Ex: 20.42 g (2 dec. places)
+ 102.3 g (1 dec. place)
130.6964 g
The answer you report is = 130.7 g (1 dec. place)

Practice +/with Sig. Figs
1) 15.002 cm + 24.1104 cm 39.112 cm
3 dec. places
2) 22.35 kg – 0.154 kg 22.20 kg
2 dec. places
3) 100.0 g – 23.73 g 76.3 g
1 dec. places
4) 2.030 mL – 1.870 mL 0.160 mL
3 dec. places

Calculations with
Significant Figures
• Multiplication or Division:
– Round the answer off to the least number of significant figures x
32.3492 cm 2
The answer you report is = 32 cm 2 (2 sig. figs)

Practice x / ÷ with Sig. Figs
1) 75.246 g / 6.33 mL 11.9 g/mL
3 sig. figs
2) 16.00 cm x 2.5 cm x 3.66 cm 150 cm 3
2 sig. figs
3) 3.24 m x 7.00 m 22.7 m 2
3 sig. figs
4) 710 m / 3.0 s 240 m/s or 2.4 x 10 2 m/s
2 sig. figs

Calculations with
Significant Figures
• Mixed Operations
– Always follow PEMDAS
– Count sig. figs as you go to get it correct
Ex: (44.739 g – 36.21 g) = =
6.561 g/mL
1.3 mL 1.3 mL
The answer you report is = 6.6 g/mL (2 sig. figs)

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Matter
And
Measurement

Dimensional Analysis (D.A.)
• It’s all about converting & matching up units!
• D.A. provides a way to check your solution method by following the units or
“dimensions” of the problem
• When a problem is solved correctly, the unwanted units will cancel with one another.
This will leave behind only the desired units!

Dimensional Analysis
• Units are…
– multiplied together
– divided into one another
– or canceled completely
• The KEY is…
– Finding the correct conversion relationship
– Setting up the relationships to cancel out the units you don’t want

D.A. – Relationships as Ratios
• How many dozens are there in 48 eggs?
• HOW did you solve that problem?
• What are the relationships/conversions you needed to know?
– Relationships:
1 dozen eggs = 12 eggs
If we write that relationship as a ratio…
1 dozen - or 12 eggs .
12 eggs 1 dozen
• How did you know to multiply or divide?
– D.A. will figure that out for you!!!

D.A. - Problem Solving Steps
1) Identify the unknown (what are you solving for).
2) Write down what is given or known
3) Look for relationships between the known and unknown.
*Note: it may take more than one relationship to get from the known to the unknown*
4) Arrange the relationships into a series of ratios so that the given unit will cancel out and the desired unit will be left over.
5) Multiply numbers across the top of the expression and divide numbers on the bottom. given unit desired unit = desired unit given unit

D.A. - Intro Example
• How many large pizzas will I need for a birthday party with 22 guests if I allow each person to eat 4 slices?
Unknown/Solving for: # of pizzas
Given: 22 guests, 4 slices each
Relationships: 1 pizza = 8 slices, 1 guest = 4 slices
*** Must set up the relationships like ratios
(which are equal to 1)
1 pizza
8 slices
8 slices
1 pizza
1 guest
4 slices
4 slices
1 guest

Example Continued
22 guests x 4 slices x
1 guest
1 pizza =
8 slices
11 pizzas
Given in problem
All the unwanted units should cancel (top cancels w/matching bottom) leaving only the desired quantity and units

• On an island with 375 inhabitants, the people are growing very concerned with the declining number of turtles
• Your mission is to calculate how many turtles are on the island.
• Use the clues the inhabitants found to get an accurate count of the turtle population

Need-to-Know Conversions
Length
• 1 inch = 2.54 cm
• 1 mile = 5280 ft
• 1 ft = 12 in
Volume
• 1 cm 3 = 1 mL
• All the metric prefixes

Practice the following using dimensional analysis and sig figs:
1) How many decimeters are equal to
32.74 yards?
299.37456 dm (w/sig.fig = 299.4 dm)
2) How many milliliters are in 2.35 gal of water?
8892.4 mL (w/sig.fig = 8890 mL)

Derived Units as Ratios
• All derived units can be set-up as ratios for use in D.A.
– Density Example:
1.35 g/mL =
1.35 g
1 mL

Practice the following using dimensional analysis and sig figs:
3) The speed of light is 3.0 x10 8 m/s. What is the speed of light in mi/hr (mph)?
6.7108 x10 8 mi/hr (w/sig. figs = 6.7 x 10 8 mph)
4) What is the mass of gold, which has a density of 19.3 g/cm 3 and a volume of 2.0 in 3 ?
632.54 g (w/sig. fig = 630 g)