Matter and Measurement Unit 1
Dec 22, 2015
Introduction Matter Physical/Chemical Properties and Changes Extensive/Intensive Properties Scientific Notation Metric System SI units Conversions Density Significant Figures Uncertainty
Unit Overview
The study of matter and the changes it undergoes.
Major divisions Inorganic Compounds of elements
other than carbon
Organic Compounds of carbon
Biochemistry Compounds of living matter
Physical Theory and concepts
Analytical Methods of analysis
What is Chemistry?
Differing Views
We can explore the MACROSCOPIC world — what we can see —
to understand the PARTICULATE worlds we cannot see.
We write SYMBOLS to describe these worlds.
Matter has mass and occupies space
What is not matter?◦ Everything in the universe is either matter or
energy
What in the world isn’t matter?
Classification of Matter Matter
Pure Substance
Mixture
Element Compound Homogeneous
Heterogeneous
Iron CO2 Juice Trail
Mix
Pure Substances Element
◦ Cannot be converted to a simpler form by a chemical reaction.
◦ Example hydrogen and oxygen
Compound◦ Combination of two or more elements in a
definite, reproducible way.
◦ Example water - H2O
Mixtures
• A combination of two or more pure substances.
◦ Homogeneous - Looks the same throughout
◦ Heterogeneous - Does not look the same throughout
Which are homogeneous or heterogeneous?
◦ Blood Skittles “T-Bone” steak
◦ Orange Juice Vegetable Soup Salad Dressing
Qualitative and Quantitative Analysis
Qualitative analysis is data that is observed Colors, textures, smells, tastes, appearance, etc
Quantitative analysis is data that can be measured
Length, height, area, volume, mass, speed, time, temperature, humidity
Physical properties can be measured or observed
color density odor melting point taste boiling point
Chemical properties describe matter’s ability to change into another substance
ability to burn reactivity ability to decompose
Properties
Changes Physical changes do not change identity of
substance tearing melting freezing boiling grinding cutting
Chemical changes change identity of substance
burning reacting combusting
Extensive and Intensive Properties Extensive properties
Depend on the quantity of sample measured.
Example - mass and volume of a sample.
Intensive propertiesIndependent of the sample size.Properties that are often characteristic of the substance being measured.
Examples - density, melting and boiling points.
Scientific Notation
• Method to express really big or small numbers.
◦ Format is Mantissa x Base Power
We just move the decimal point around
Decimal part oforiginal number
Decimalsyou moved
Really Big Numbers If a number is larger than 1• The original decimal point is moved X places to
the left.
•The resulting number is multiplied by 10X.
•The exponent is the number of places you moved the decimal point.
1 2 3 0 0 0 0 0 0. = 1.23 x 108
Really Small Numbers If a number is smaller than 1• The original decimal point is moved X places to
the right.
•The resulting number is multiplied by 10-X.
•The exponent is the number of places you moved the decimal point.
0. 0 0 0 0 0 0 1 2 3 = 1.23 x 10-7
Using Your Calculator Most calculators use scientific notation
when the numbers get very large or small.
How scientific notation is displayed can vary.
It may use x10n
or may be displayed using an E.
They usually have an Exp or EE◦ This is to enter in the exponent. +
-1
/
x
0
2 3
4 5 6
7 8 9
.
CE
EE
log
ln
1/x
x2
cos tan
English units ◦ Still commonly used in the United States
Examples: pound, inch, foot, cup, pint
Why English system not used in chemistry ◦ Very confusing and difficult to keep track of the
conversions needed◦ Vary in size so you must memorize many
conversion factors
Measurements in Chemistry
Changing the prefix alters the size of a unit.
Metric prefixes
Prefix Symbol Factor
mega M 106 1 000 000
kilo k 103 1 000
hecto h 102 100
deka da 101 10
base - 100 1
deci d 10-1 0.1
centi c 10-2 0.01
milli m 10-3 0.001
micro µ 10-6 0.000001
nano n 10-9 0.000000001
pico p 10-12 0.000000000001
SI - System International - systematic subset of the metric system
Physical Quantity Name AbbreviationMass kilograms kgLength meters mTime seconds sTemperature Kelvin KAmount mole molElectric Current Ampere ALuminous Intensity candela cd
SI Units
Give you ability to convert between units
Problem solving technique (factor label method)
1. Write what you know2. Game plan3. Set up units4. Conversion factors5. Solve
Conversions
Convert 26 gallons to cups◦ Answer: 416 cups
Convert 18 miles to centimeters◦ Answer: 2.9×106 cm
Conversion practice
Allow you to convert between metric prefixes
Write what you know Set up units Bigger unit gets the “1” Smaller unit is 10x where “x” is how many places
apart the two units are
Example◦ Convert .25 kg to mg.
Answer: 2.5x105 mg
Metric Conversions
Ratio of mass to volume of matter
Common units are g / cm3 or g / mL.
Example: what is the density of 5.00 mL of a fluid if it has a mass of 5.23 grams?
d = mass / volume d = 5.23 g / 5.00 mL d = 1.05 g / mL
Example 2: What would be the mass of 1.00 liters of this sample?
Density
Density = Mass
Volume
cm3 = mL cm3 = mL
Measuring Mass Mass - the quantity of matter in an object. Weight - the effect of gravity on an object.
Since the Earth’s gravity is relatively constant, we can interconvert between weight and mass.
The SI unit of mass is the kilogram (kg). However, in the lab, the gram (g) is more commonly used.
Measuring Volume Volume - the amount of space that an object
occupies.
• The base metric unit is the liter (L).
• The common unit used in the lab is the milliliter (mL).
• One milliliter is exactly equal to one cm3.
• The derived SI unit for volume is the m3 which is too large for convenient use.
Significant Figures Method used to express accuracy and
precision.
You can’t report numbers better than the method used to measure them.
67.2 units = three significant figures
Significant Figures The number of significant digits is
independent of the decimal point.
255 25.5 2.55 0.255 0.0255
These numbersAll have three
significant figures!
Significant Figure Rules
Leading zeros are not significant.
Leading zeroLeading zero
Captive zeroCaptive zero
Trailing zeroTrailing zero
0.421 - three significant figures
4012 - four significant figures
114.20 - five significant figures
Zeroes between non-zeros are significant.
Trailing zeros are significant ONLY IF there is a decimal point in the number.
How many significant figures are in the following?
Significant Figures
1.0070 m 5 sig figs
17.10 kg 4 sig figs
100,890 L 5 sig figs
3.29 x 103 s 3 sig figs
0.0054 cm 2 sig figs
3,200,000 2 sig figs
Significant Figures and Calculations
123.45987 g+ 234.11 g 357.57 g
805.4 g- 721.67912 g 83.7 g
Addition and subtraction Report your answer with the same
number of digits to the right of the decimal point as the number having the fewest to start with.
Significant Figures and Calculations
Multiplication and division.Report your answer with the same
number of digits as the quantity have the smallest number of significant figures.
Example. Density of a rectangular solid.25.12 kg / [ (18.5 m) ( 0.2351 m) (2.1m) ]
= 2.8 kg / m3
(2.1 m - only has two significant figures)
After calculations, you may need to round off
◦ If the first insignificant digit is 5 or more, you round up
◦ If the first insignificant digit is 4 or less, you round down
Rounding Off
A properly written number in scientific notation always has the proper number of significant figures.
Scientific Notation and Significant Figures
0.00321 = 3.21 x 10-3
Three SignificantFigures
Three SignificantFigures
Measured and Exact Numbers
In science, all of our numbers are either measured or exact.
Exact - Infinite number of significant figures.1 foot = exactly 12 inches
Do not count toward significant figures
Measured - the tool used will tell you the level of significance and varies based on the tool.
When using a measuring tool, record the numbers that are certain and add a guess number
Significant Figures in Measurement
Types of Error Systematic• Errors in a single direction (high or low).
•Can be corrected by proper calibration or running controls and blanks.
Random•Errors in any direction.
•Can’t be corrected. Can only be accounted for by using statistics.
Uncertainty in Measurement
All measurements contain some uncertainty.• We make errors• Tools have limits
Accuracy How close to the true value
Precision How close to each other
Neither accurate nor precise
Precise but not accurate
Precise AND accurate