Introduction: Matter & Measurement

AP ChemistryChapter 1

ChemistryWhat is chemistry?What is chemistry?

It is the study of the composition composition of matterof matter and the changes that matter undergoes.

What is matter?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

Mixtures2. 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

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 PropertiesPhysical 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 PropertiesIntensive 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 ChangesPhysical changes• The composition of the

substance doesn’t 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

1. permanent color change2. gas produced (odor or

bubbles)3. precipitate (solid) produced 4. light given off 5. heat released (exothermic)

or absorbed (endothermic)

Making MeasurementsMaking Measurements

•A measurement is a A measurement is a number with a unit.number with a unit.

•All measurements, MUST All measurements, MUST have units.have units.

Types of UnitsTypes of Units

Energy Joule JPressure 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- 0.000001 10-6

• nano- n 0.000000001 10-9

• pico- p 0.000000000001 10-12

MeasurementsMeasurementsThere are two types of measurements: Qualitative data are words, such as

color, heavy or hot. Quantitative measurements involve

numbers (quantities), and depend on:

1. The reliability of the measuring instrument.

2. The care with which it is read – this is determined by YOU!

Accuracy & PrecisionAccuracy & Precision

AccuracyAccuracy – how close a – how close a measurement is to the true measurement is to the true value.value.

PrecisionPrecision – how close the – how close the measurements are to each measurements are to each other (reproducibility).other (reproducibility).

Precision and AccuracyPrecision and Accuracy

Neither accurate

nor precise

Precise, but not

accurate

Precise AND

accurate

Our goal!

Which are Precise? Which are Precise? Accurate?Accurate?

Uncertainty in MeasurementsUncertainty in MeasurementsMeasurements 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 can’t have accuracy without

precision (unless you’re really lucky).

Significant Figures in Significant Figures in MeasurementsMeasurements

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 Rules for Counting Significant FiguresSignificant Figures

Non-zerosNon-zeros always count as always count as significant figures:significant figures:

34563456 hashas

44 significant figuressignificant figures

Rules for Counting Rules for Counting Significant FiguresSignificant Figures

ZerosZerosLeading zeroes do not count Leading zeroes do not count as significant figures:as significant figures:

0.0486 0.0486 hashas

3 3 significant figuressignificant figures

Rules for Counting Rules for Counting Significant FiguresSignificant Figures

ZerosZerosCaptive zeroes always count Captive zeroes always count as significant figures:as significant figures:

16.0716.07 hashas

44 significant figures significant figures

Rules for Counting Rules for Counting Significant FiguresSignificant Figures

ZerosZerosTrailing zerosTrailing zeros are significant are significant

only if the number contains only if the number contains a written decimal point:a written decimal point:

9.300 9.300 hashas

4 4 significant figuressignificant figures

Rules for Counting Rules for Counting Significant FiguresSignificant Figures

Two special situationsTwo special situations have an have an unlimitedunlimited (infinite) number (infinite) number of significant figures:of significant figures:

1.1. Counted itemsCounted itemsa)a) 23 people, or 36 desks23 people, or 36 desks

2.2. Exactly defined quantitiesExactly defined quantitiesb)b) 60 minutes = 1 hour60 minutes = 1 hour

Sig Fig Practice #1Sig Fig Practice #1How many significant figures in the following?

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 mL 2 sig figs3 cats infinite

These all come from some measurements

This is a counted value

Significant Figures in Significant Figures in CalculationsCalculations

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 Rounding Calculated AnswersAnswers

RoundingRounding Decide Decide how manyhow many significant significant

figures are neededfigures are needed Round to that many digits, Round to that many digits,

counting from the counting from the leftleft Is the next digit less than 5? Is the next digit less than 5?

Drop it.Drop it. Next digit 5 or greater? Increase Next digit 5 or greater? Increase

by 1by 1

Rules for Significant Figures in Rules for Significant Figures in Mathematical OperationsMathematical Operations

Addition and SubtractionThe 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 Rules for Significant Figures in Mathematical OperationsMathematical Operations

• Addition and SubtractionAddition and Subtraction: The : The number of decimal places in the number of decimal places in the result equals the number of result equals the number of decimal places in the decimal places in the least least accurate accurate measurement.measurement.

6.8 + 11.934 =6.8 + 11.934 =18.734 18.734 18.7 18.7 ((3 sig figs3 sig figs))

Sig Fig Practice #2Sig Fig Practice #2

3.24 m + 7.0 m

Calculation Calculator says: Answer

10.24 m 10.2 m

100.0 g - 23.73 g 76.27 g 76.3 g

0.02 cm + 2.371 cm 2.391 cm 2.39 cm

713.1 L - 3.872 L 709.228 L 709.2 L

1818 lb + 3.37 lb 1821.37 lb 1821 lb

2.030 mL - 1.870 mL 0.16 mL 0.160 mL

*Note the zero that has been added.

Rounding Calculated Rounding Calculated AnswersAnswers

Multiplication and DivisionRound the answer to the same number of significant figures as the least number of significant figures in the problem.

Rules for Significant Figures in Rules for Significant Figures in Mathematical OperationsMathematical 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 CasesOther Special Cases

• What if your answer has less What if your answer has less significant figures than you are significant figures than you are supposed to have?supposed to have?– Calculator Example: 100.00 / 5.00 Calculator Example: 100.00 / 5.00

= 20= 20

• Add zeros!Add zeros!– 20 is 1 sf20 is 1 sf– 20. is 2 sf20. is 2 sf– 20.0 is 3 sf20.0 is 3 sf

Sig Fig Practice #3Sig Fig Practice #3

3.24 m x 7.0 m

Calculation Calculator says: Answer

22.68 m2 23 m2

100.0 g ÷ 23.7 cm3 4.219409283 g/cm3 4.22 g/cm3

0.02 cm x 2.371 cm 0.04742 cm2 0.05 cm2

710 m ÷ 3.0 s 236.6666667 m/s 240 m/s

1818.2 lb x 3.23 ft 5872.786 lb·ft 5870 lb·ft

1.030 g x 2.87 mL 2.9561 g/mL 2.96 g/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?

• 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?

Examples

• 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

100ºC 212ºF

100ºC = 212ºF

0ºC = 32ºF

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