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Introduction Matter Measurement

- AP Chemistry
- Chapter 1

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.

Elements, Compounds Mixtures

- A substance is matter that has a definite

composition and constant properties. - It can be an element or a compound.

Elements, Compounds Mixtures

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

reactions.

Elements, Compounds Mixtures

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

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

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.

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.

Filtration

Distillation

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.

Physical Chemical Properties

- Chemical property describes the chemical

reaction of a pure substance with another

substance chemical reaction is involved.

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.

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.

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

Signs of a Chemical Changes

- permanent color change
- gas produced (odor or bubbles)
- precipitate (solid) produced
- light given off
- heat released (exothermic) or absorbed

(endothermic)

Making Measurements

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

Types of Units

Energy Joule J Pressure

Pascal Pa

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

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!

Accuracy Precision

- Accuracy how close a measurement is to the true

value. - Precision how close the measurements are to

each other (reproducibility).

Precision and Accuracy

Precise, but not accurate

Neither accurate nor precise

Precise AND accurate

Our goal!

Which are Precise? Accurate?

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?

1

2

3

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

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).

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.

To how many significant figures can you measure

this pencil?

What is wrong with this ruler? What is it missing?

Rules for Counting Significant Figures

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

Rules for Counting Significant Figures

- Zeros
- Leading zeroes do not count as significant

figures - 0.0486 has
- 3 significant figures

Rules for Counting Significant Figures

- Zeros
- Captive zeroes always count as significant

figures - 16.07 has
- 4 significant figures

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

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

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

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.

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

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.

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)

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.

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.

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)

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

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

Dimensional Analysis

- Using the units to solve problems

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.

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?

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

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?

- 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?

Temperature and Density

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.

Converting ºF to ºC and vice versa

Fahrenheit to Celsius (F - 32) x 5/9 C

Celsius to Fahrenheit (C 9/5) 32 F

0ºC 32ºF

0ºC

32ºF

0ºC 32ºF

100ºC 212ºF

100ºC

212ºF

0ºC

32ºF

Converting oC to K and vice versa

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

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.

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?

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