Introduction: Matter

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Introduction Matter Measurement

• AP Chemistry
• Chapter 1

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Chemistry

• What is chemistry?
• It is the study of the composition of matter and
  the changes that matter undergoes.
• What is matter?
• It is anything that takes up space and has mass.

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Elements, Compounds Mixtures

• A substance is matter that has a definite
  composition and constant properties.
• It can be an element or a compound.

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Elements, Compounds Mixtures

• An element is the simplest form of matter.
• It cannot be broken down further by chemical
  reactions.

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Elements, Compounds Mixtures

• A compound can be separated into simpler forms.
• It is a combination of two or more elements.

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Mixtures

• A mixture is a physical blend of two or more
  substances.
• 1. Heterogeneous Mixtures
• Not uniform in composition
• Properties indefinite vary
• Can be separated by physical methods

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Mixtures

• 2. Homogeneous Mixtures
• Completely uniform in composition
• Properties constant for a given sample
• Cannot be separated by physical methods (need
  distillation, chromatography, etc)
• Also called solutions.

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Separating mixtures

• Only a physical change- no new matter
• Filtration- separate solids from liquids with a
  barrier.
• Distillation- separate different liquids or
  solutions of a solid and a liquid using boiling
  points.
• Heat the mixture.
• Catch vapor and cool it to retrieve the liquid.
• Chromatography- different substances are
  attracted to paper or gel, so move at different
  speeds.

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Filtration
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Distillation
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Chromatography
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Physical Chemical Properties

• Physical property characteristics of a pure
  substance that we can observe without changing
  the substance the chemical composition of the
  substance does not change.

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Physical Chemical Properties

• Chemical property describes the chemical
  reaction of a pure substance with another
  substance chemical reaction is involved.

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Physical Chemical Properties

• Physical properties
• appearance
• odor
• melting point
• boiling point
• hardness
• density
• solubility
• conductivity

• Chemical properties
• reaction with oxygen (flammability)
• rxn with water
• rxn with acid
• Etc.

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Intensive Extensive Properties

• Intensive properties
• Do not depend on the amount of sample being
  examined
• temperature
• odor
• melting point
• boiling point
• hardness
• density

• Extensive properties
• Depend on the quantity of the sample
• mass
• volume
• Etc.

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Physical Chemical Changes

• Physical changes
• The composition of the substance doesnt change
• Phase changes (like liquid to gas)
• Evaporation, freezing, condensing, subliming,
  etc.
• Tearing or cutting the substance

• Chemical changes
• The substance is transformed into a chemically
  different substance
• All chemical reactions

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Signs of a Chemical Changes

1. permanent color change
2. gas produced (odor or bubbles)
3. precipitate (solid) produced
4. light given off
5. heat released (exothermic) or absorbed
   (endothermic)

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Making Measurements

• A measurement is a number with a unit.
• All measurements, MUST have units.

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Types of Units
Energy Joule J Pressure
Pascal Pa
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Prefixes

• giga- G 1,000,000,000 109
• mega - M 1,000,000 106
• kilo - k 1,000 103
• deci- d 0.1 10-1
• centi- c 0.01 10-2
• milli- m 0.001 10-3
• micro- m 0.000001 10-6
• nano- n 0.000000001 10-9
• pico- p 0.000000000001 10-12

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Measurements

• There are two types of measurements
• Qualitative data are words, such as color, heavy
  or hot.
• Quantitative measurements involve numbers
  (quantities), and depend on
• The reliability of the measuring instrument.
• The care with which it is read this is
  determined by YOU!

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Accuracy Precision

• Accuracy how close a measurement is to the true
  value.
• Precision how close the measurements are to
  each other (reproducibility).

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Precision and Accuracy
Precise, but not accurate
Neither accurate nor precise
Precise AND accurate
Our goal!
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Which are Precise? Accurate?
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Uncertainty in Measurements
Measurements are performed with instruments, and
no instrument can read to an infinite number of
decimal places

• Which of the balances below has the greatest
  uncertainty in measurement?

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2
3
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Uncertainty

• Basis for significant figures
• All measurements are uncertain to some degree
• Precision- how repeatable
• Accuracy- how correct - closeness to true value.
• Random error - equal chance of being high or low-
  addressed by averaging measurements - expected

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Uncertainty

• Systematic error- same direction each time
• Want to avoid this
• Bad equipment or bad technique.
• Better precision implies better accuracy.
• You can have precision without accuracy.
• You cant have accuracy without precision (unless
  youre really lucky).

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Significant Figures in Measurements

• Significant figures in a measurement include all
  of the digits that are known, plus one more digit
  that is estimated.
• Sig figs help to account for the uncertainty in a
  measurement.

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To how many significant figures can you measure
this pencil?
What is wrong with this ruler? What is it missing?
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Rules for Counting Significant Figures

• Non-zeros always count as significant figures
• 3456 has
• 4 significant figures

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Rules for Counting Significant Figures

• Zeros
• Leading zeroes do not count as significant
  figures
• 0.0486 has
• 3 significant figures

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Rules for Counting Significant Figures

• Zeros
• Captive zeroes always count as significant
  figures
• 16.07 has
• 4 significant figures

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Rules for Counting Significant Figures

• Zeros
• Trailing zeros are significant only if the number
  contains a written decimal point
• 9.300 has
• 4 significant figures

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Rules for Counting Significant Figures

• Two special situations have an unlimited
  (infinite) number of significant figures
• Counted items
• 23 people, or 36 desks
• Exactly defined quantities
• 60 minutes 1 hour

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Sig Fig Practice 1
How many significant figures in the following?
1.0070 m ?
5 sig figs
17.10 kg ?
4 sig figs
These all come from some measurements
100,890 L ?
5 sig figs
3.29 x 103 s ?
3 sig figs
0.0054 cm ?
2 sig figs
3,200,000 mL ?
2 sig figs
This is a counted value
3 cats ?
infinite
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Significant Figures in Calculations

• In general a calculated answer cannot be more
  accurate than the least accurate measurement from
  which it was calculated.
• Sometimes, calculated values need to be rounded
  off.

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Rounding Calculated Answers

• Rounding
• Decide how many significant figures are needed
• Round to that many digits, counting from the left
• Is the next digit less than 5? Drop it.
• Next digit 5 or greater? Increase by 1

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Rules for Significant Figures in Mathematical
Operations

• Addition and Subtraction
• The answer should be rounded to the same number
  of decimal places as the least number of decimal
  places in the problem.

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Rules for Significant Figures in Mathematical
Operations

• Addition and Subtraction The number of decimal
  places in the result equals the number of decimal
  places in the least accurate measurement.
• 6.8 11.934
• 18.734 ? 18.7 (3 sig figs)

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Sig Fig Practice 2
Calculation
Calculator says
Answer
10.24 m
3.24 m 7.0 m
10.2 m
100.0 g - 23.73 g
76.3 g
76.27 g
0.02 cm 2.371 cm
2.39 cm
2.391 cm
713.1 L - 3.872 L
709.228 L
709.2 L
1821 lb
1818 lb 3.37 lb
1821.37 lb
0.160 mL
0.16 mL
2.030 mL - 1.870 mL
Note the zero that has been added.
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Rounding Calculated Answers

• Multiplication and Division
• Round the answer to the same number of
  significant figures as the least number of
  significant figures in the problem.

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Rules for Significant Figures in Mathematical
Operations

• Multiplication and Division sig figs in the
  result equals the number in the least accurate
  measurement used in the calculation.
• 6.38 x 2.0
• 12.76 ? 13 (2 sig figs)

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Other Special Cases

• What if your answer has less significant figures
  than you are supposed to have?
• Calculator Example 100.00 / 5.00 20
• Add zeros!
• 20 is 1 sf
• 20. is 2 sf
• 20.0 is 3 sf

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Sig Fig Practice 3
Calculation
Calculator says
Answer
22.68 m2
3.24 m x 7.0 m
23 m2
100.0 g 23.7 cm3
4.22 g/cm3
4.219409283 g/cm3
0.02 cm x 2.371 cm
0.05 cm2
0.04742 cm2
710 m 3.0 s
236.6666667 m/s
240 m/s
5870 lbft
1818.2 lb x 3.23 ft
5872.786 lbft
2.9561 g/mL
2.96 g/mL
1.030 g x 2.87 mL
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Dimensional Analysis

• Using the units to solve problems

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Dimensional Analysis

• Use conversion factors to change the units
• Conversion factors 1
• 1 foot 12 inches (equivalence statement)
• 12 in 1 1 ft.
  
  1 ft. 12 in
• 2 conversion factors
• multiply by the one that will give you the
  correct units in your answer.

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Examples

• 11 yards 2 rod
• 40 rods 1 furlong
• 8 furlongs 1 mile
• The Kentucky Derby race is 1.25 miles. How long
  is the race in rods, furlongs, meters, and
  kilometers?
• A marathon race is 26 miles, 385 yards. What is
  this distance in rods and kilometers?

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Examples

• Science fiction often uses nautical analogies to
  describe space travel. If the starship U.S.S.
  Enterprise is traveling at warp factor 1.71, what
  is its speed in knots?
• Warp 1.71 5.00 times the speed of light
• speed of light 3.00 x 108 m/s
• 1 knot 2000 yd/h exactly

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Examples

• Because you never learned dimensional analysis,
  you have been working at a fast food restaurant
  for the past 35 years wrapping hamburgers. Each
  hour you wrap 184 hamburgers. You work 8 hours
  per day. You work 5 days a week. you get paid
  every 2 weeks with a salary of 840.34. How many
  hamburgers will you have to wrap to make your
  first one million dollars?

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• A senior was applying to college and wondered how
  many applications she needed to send. Her
  counselor explained that with the excellent grade
  she received in chemistry she would probably be
  accepted to one school out of every three to
  which she applied. She immediately realized that
  for each application she would have to write 3
  essays, and each essay would require 2 hours
  work. Of course writing essays is no simple
  matter. For each hour of serious essay writing,
  she would need to expend 500 calories which she
  could derive from her mother's apple pies. Every
  three times she cleaned her bedroom, her mother
  would made her an apple pie. How many times would
  she have to clean her room in order to gain
  acceptance to 10 colleges?

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Temperature and Density
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Temperature

• A measure of the average kinetic energy
• Different temperature scales, all are talking
  about the same height of mercury.
• We make measurements in lab using the Celsius
  scale, but most chemistry problems require you to
  change the temperature to Kelvin before using in
  an equation.

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Converting ºF to ºC and vice versa
Fahrenheit to Celsius (F - 32) x 5/9 C
Celsius to Fahrenheit (C 9/5) 32 F
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0ºC 32ºF
0ºC
32ºF
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0ºC 32ºF
100ºC 212ºF
100ºC
212ºF
0ºC
32ºF
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Converting oC to K and vice versa

• Celsius to Kelvin K oC 273.15
• Kelvin to Celsius oC K - 273.15

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Density

• Ratio of mass to volume
• D m/V
• Useful for identifying a compound
• Useful for predicting weight
• An intrinsic property- does depend on what the
  material is.

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Density Problem

• An empty container weighs 121.3 g. Filled with
  carbon tetrachloride (density 1.53 g/cm3 ) the
  container weighs 283.2 g. What is the volume of
  the container?

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Density Problem

• A 55.0 gal drum weighs 75.0 lbs. when empty. What
  will the total mass be when filled with ethanol?
  density 0.789 g/cm3
• 1 gal 3.78 L
• 1 lb 454 g