Chemistry 100(02) Fall 2001Chemistry 100(02) Fall 2001

Dr. Upali Siriwardane

CTH 311 Phone 257-4941

Office Hours: M, Tu, W, Th, F 9:00-11:00 a.m.

Test 1 : Chapters 1, 2: September 26Test 2: Chapters 3, 4: October 24Test 3: Chapters 4, 5: November 14Make-up, Comprehensive, November 15

Chemistry 100(04) Fall 2001Chemistry 100(04) Fall 2001Dr. Upali Siriwardane

CTH 311 Phone 257-4941

Office Hours: 8:00-9:00, 11:00-12:00 M, W

Tu, Th, F 10:00-12:00 a.m.

Test 1 : Chapters 1, 2: October 2 Test 2: Chapters 3, 4: October 25Test 3: Chapters 4, 5: November 13Make-up, Comprehensive, November 15

Chemistry 100(05) Fall 2001Chemistry 100(05) Fall 2001

Dr. Upali Siriwardane

CTH 311 Phone 257-4941

Office Hours: M, Tu, W, Th, F 9:00-11:00 a.m.

Test 1 : Chapters 1, 2: October 2. Test 2: Chapters 3, 4: October 25Test 3: Chapters 4, 5: November 13Make-up, Comprehensive, November 15

KEY CONCEPTSKEY CONCEPTSWhat is chemistry?

Physical & chemical changes.

Physical & chemical properties.

Categories of matter. Separating Mixtures.Scientific Method. Scientific Measurement Observation Uncertainty.

Significant figure. Precision. Accuracy. Significant figures in calculations. Unit Conversions. Temperature Conversions. Unit conversion method.

Density Calculations.

What is chemistry?What is chemistry?• Chemistry deals with non-reversible

changes of matter.

• Chemistry explains using atoms and molecules.

• Chemical Concepts and Models improve your problem solving skills

• Chemistry is a Central Science

ChemistryChemistry• “The study of matter and the changes it undergoes.”

• Major divisionsMajor divisions

• InorganicInorganic Compounds of elements other than carbon

• OrganicOrganic Compounds of carbon

• BiochemistryBiochemistry Compounds of living matter

• PhysicalPhysical Theory and concepts

• AnalyticalAnalytical Methods of analysis

What is MatterWhat is Matter• Matter:

• Anything that has a mass and volume.

• Energy:

• Manifestations of matter.

• Matter and Energy is intertwined.

Classification of matterClassification of matter

Matter

Pure Substance Mixture

Element Compound Homogeneous

Heterogeneous

Iron Hemoglobin Plasma Blood

Mixtures Mixtures

HeterogeneousHeterogeneous HomogenousHomogenous

CompoundsCompounds

AtomsAtoms

ElectronsElectrons

NucleusNucleus

ProtonsProtonsNeutronsNeutrons

Pure SubstancesPure Substances

ElementsElements

Hierarchy of Matter

MixturesMixtures• A combination of two or more pure

substances.

– HomogeneousHomogeneous - Uniform composition

– HeterogeneousHeterogeneous - Non-uniform composition

• Which are homogeneous or heterogeneous?

– Blood Urine “T-Bone” steak

– Gasoline Twinkie Salad Dressing

How do you Separate How do you Separate Mixtures?Mixtures?

• Flotation: based on density

• Filtration: Solid- liquid

• Distillation- Liquid-liquid

• Magnetic Separation- Magnetic-

• Chromatography:

• 1) Paper 2) Column 3) Gas

Pure substancesPure substances

• ElementElement– Cannot be converted to a simpler form by a chemical

reaction.

– ExampleExample hydrogen and oxygen

• CompoundCompound– Combination of two or more elements in a definite,

reproducible way.

– ExampleExample water - H2O

• Both elements and compounds have characteristic properties such as color, boiling point and reactivity

Pure substancesPure substances

• The properties of a compound and the elements it is made of can differ greatly.

• Formula BP density Other

• Hydrogen H2 -253 0.90 Flammable

• Oxygen O2 -297 1.14 Supports combustion

• Water H2O 100 1000 Not flammable

Properties of SubstancesProperties of Substances• Physical propertiesPhysical properties:

• Physical properties are descriptions of matter such as color, density, viscosity, boiling point, and melting point.

• Chemical propertiesChemical properties:

• Chemical properties relates to the changes of substances making up the matter. For example, corrosiveness, Flammability

Extensive and intensive propertiesExtensive and intensive properties

• Extensive propertiesExtensive properties

• Depend on the quantity of sample measured.

• ExampleExample - mass and volume of a sample.

• Intensive propertiesIntensive properties

• Independent of the sample size.

• Properties that are often characteristic of the substance being measured.

• ExamplesExamples - density, melting and boiling points.

Physical propertiesPhysical properties

• Properties that do not involve substances changing into another substance.

• ExamplesExamples

– color density

– odor melting point

– taste boiling point

– feel compressibility

Chemical propertiesChemical properties• Properties that involve substances changing into

another substance.

• Chemical ReactionChemical Reaction - one or more substances are changed into other substances.

• Example A chemical property of wood is it’s ability to burn - combustioncombustion.

wood + oxygen carbon dioxide + water + heat

ReactantsReactants ProductsProducts The reactants and products are very different.

ExampleExample • Which are chemical or physical changes?

• Mulching leavesMulching leaves

• Milk turning sourMilk turning sour

• Making wineMaking wine

• Making ice waterMaking ice water

• Beer going flatBeer going flat

• Leaves changing colorLeaves changing color

Type of ChangesType of Changes• Physical changePhysical change:

• A change in the state of matter. It does not involve a change in the substances. E.g. melting of wax and water.

• Chemical changeChemical change: • A change involving at least one of the

substances making the matter. E.g. Electrolysis of water, formation of rust: reaction of iron and oxygen to from iron oxide.

Chemical verses Physical changeChemical verses Physical change

Sodium reacting Iodine changing with chlorine. from a solid to a gas

Scientific methodScientific method• All scientific studies follow the same

approach to examining a problem.

• The scientific method requires that we: •

• Make observations

• Apply logical, organized reasoning to observations made.

• Form a hypothesis.

• Reject or confirm that hypothesis through experiments.

Scientific MethodScientific Method. • A method common to all sciences has

• Four Basic Steps:

• a) Experiment

• b) Data or Results

• c) Hypothesis

• d) Further experiments to test hypothesis

Develop a theory

Do experiments

Do more experiments

Organize

Make hypothesis

Try newtests Did hypothesis work?

Yes

Scientific methodScientific methodMake observations

No

MeasurementMeasurement• Measurements or observations are made

• using our physical senses or using scientific instruments.

• 1) Qualitative measurements.

• Changes that cannot be expressed in terms of a number.

• 2) Quantitative measurements.

• expressed in terms of a number and an unit.

Units are importantUnits are important

• 45 000 has little meaning, just a number

• $45,000 has some meaning - money

• $45,000/yr more meaning - person’s salary

SI units• SI - System InternationalSystem International

– Systematic subset of the metric system.

• Only uses certain metric units.• Mass kilograms

• Length meters

• Time seconds

• Temperature kelvin

• Amount mole

• Other SI units are derived from SI base units.

Metric prefixesMetric prefixes• Changing the prefix alters the size of a unit.

PrefixPrefix Symbol Symbol Factor 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

Other UnitsOther Units• Derived UnitsDerived Units. Units consisting of more

than one one base unit. E.g. g/cm3

• English units.English units. – Still commonly used in the United States.

Weight ounce, pound, ton Length inch, foot, yard, mile Volume cup, pint, quart, gallon

• Not often used in scientific work.Not often used in scientific work. – Very confusing and difficult to keep track

of the conversions needed.

Converting unitsConverting units

• Factor label methodFactor label method

• Regardless of conversion, keeping track of units makes thing come out right

• Must use conversion factors• - The relationship between two units

• Canceling out units is a way of checking that your calculation is set up right!

• Other names used Unit Conversion Methoddimensional(Unit) Analysis

Example. Metric conversionExample. Metric conversion

How many milligrams are in a kilogram?

1 kg = 1000 g1 g = 1000 mg

1 kg x 1000 x 1000

= 1 000 000 mg

kg

g mg

g

ExampleExampleCreatinine is a substance found in blood. If

an analysis of blood serum sample detected 0.58 mg of creatinine, how many micrograms were present?

= 10-6 = micro = 10-6 = micro

0.580 mg = 580 g10-3 g1 mg( ) 1 g

10-6 g( )

Common conversion factorsCommon conversion factors

• English FactorFactor– 1 gallon = 4 quarts 4 qt/gal– 1 mile = 5280 feet 5280 ft/mile– 1 ton = 2000 pounds 2000 lb/ton

• Common English to Metric conversions FactorFactor– 1 liter = 1.057 quarts 1.057 qt/L– 1 kilogram = 2.2 pounds 2.2 lb/kg– 1 meter = 1.094 yards 1.094 yd/m– 1 inch = 2.54 cm 2.54 cm/inch

3.00 x 108 m 6.21 x 10-4 mile 31536000 s

1 s 1 m 1 yr

= 5.87 x 1012

mile/yr

.Speed of light is 3.00 x 108 m s-1 . Convert the speed of light to miles per year (1 mile = 1.61 km).

MeasurementMeasurement• Number PartNumber Part

Uncertainty in MeasurementUncertainty in Measurement

Significant FiguresSignificant Figures

• Exact MeasurementsExact Measurements

• Extensive and Intensive PropertiesExtensive and Intensive Properties

• DensityDensity

• Measuring Temperature and VolumeMeasuring Temperature and Volume

Uncertainty in MeasurementUncertainty in Measurement

• All measurements contain some uncertainty.

• We make errors

• Tools have limits

• Uncertainty is measured with

– AccuracyAccuracy How close to the true value

– PrecisionPrecision How close to each other

AccuracyAccuracy

Here the average valuewould give agood number but the numbersdon’t agree.

Large random error

How close our values agree with the true value.

PrecisionPrecision

Here the numbersare close togetherso we have goodprecision.

• Poor accuracy.

• Large systematic

error.

How well our values agree with each other.

Accuracy and precisionAccuracy and precision

• Predict the effect on accuracy and precisionPredict the effect on accuracy and precision..

• Instrument not ‘zeroed’ properly

• Reagents made at wrong concentration

• Temperature in room varies ‘wildly’

• Person running test is not properly trained

Types of errors

• Instrument not ‘zeroed’ properly– Reagents made at wrong concentration

– Temperature in room varies ‘wildly’– Person running test is not properly trained

Random

Systematic

Significant figuresSignificant 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

Certain Digits

UncertainDigit

• Leading zeros

• Trailing zeros in whole numbers (use scientific notion to avoid confusion.

• Exact numbers: unit definition has an unlimited number of sig. figs. 1 ft = 12 in

Not significantNot significant

SignificantSignificant

•Non-zero digits are always significant. •Any zeros between two significant digits •Trailing zeros in the decimal portion

ExamplesExamples

• 0.00341........3 sig. digs.

• 1.0040.........5 sig. digs.

• 0.00005........1 sig. digs.

• 65000.......… 2 sig. digs. 6.5 x 104

• 40300..........3 sig. digs.

• 200300.........4 sig. digs. 2.003 x 105

Significant figures: Significant figures: Rules for zerosRules for zeros

• Leading zeros are notare not significant.0.421 - three significant figures

Leading zeroLeading zero

Captive zeros areare significant.4012 - four significant figures

Trailing zeros areare significant.114.20 - five significant figures

Captive zeroCaptive zero

Trailing zeroTrailing zero

Significant figuresSignificant figures

• Zeros are what will give you a headache!Zeros are what will give you a headache!

• They are used/misused all of the time.

• ExampleExample

• The press might report that the federal deficit is three trillion dollars. What did they mean?

$3 x 1012

or$3,000,000,000,000.00

Significant figuresSignificant figures• In science, all of our numbers are either

measured or exact.

• ExactExact - Infinite number of significant figures.

• MeasuredMeasured - the tool used will tell you the level of significance. Varies based on the tool.

• ExampleExample

• Ruler with lines at 1/16” intervals.

• A balance might be able to measure to the nearest 0.1 grams.

• 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.

Scientific notationScientific notation

+

-1

/

x

0

2 3

4 5 6

7 8 9

.

CE

EE

log

ln

1/x

x2

cos tan

1.44939 E-2

ExamplesExamples• 378 000

• 3.78 x 10 5

• 8931.5

• 8.9315 x 10 3

• 0.000 593

• 5.93 x 10 - 4

• 0.000 000 4

• 4 x 10 - 7

Significant figuresand calculations

• An answer can’t have more significant figures than the quantities used to produce it.

– ExampleExample • How fast did you run if you• went 1.0 km in 3.0 minutes?

speed = 1.0 km / 3.0 min = 0.33 km / min

+

-1

/

x

0

2 3

4 5 6

7 8 9

.

CE

EE

log

ln

1/x

x2

cos tan

0.333333333

Significant figure in CalculationsSignificant figure in Calculations

Different rules apply in each case Different rules apply in each case

• Addition and subtraction

• In multiplication and division

• The root or power of a measurement

• Exact Numbers: Numbers coming from

definitions such as 12 in = 1 foot. They are not

considered in Sig. Fig. Calculations.

Significant figures and calculationsSignificant figures and calculations• Addition and subtractionAddition 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.

123.45987 g+ 234.11 g 357.57 g

805.4 g- 721.67912 g 83.7 g

Significant figures and calculationsSignificant figures and calculations• Multiplication and division.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.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)

Rounding off numbersRounding off numbers

• 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.

• If the digit 5 exactly rounded off to a even

a) 10.26 + 0.02 + 450 = 460.28 = 460 2 significant figures

ones place in 450 has the greatest uncertainty

b) 3.56 x 1.2 = 4.272 = 4.3 2 significant figuresc) 356.001 x 15.43

-------------------- = 45.58585 = 45.59 120.5

4 significant figuresc) (2.8 + 3.4) (2.35 - 4.1) = 6.2 x (-1.75) = 6.2 x (-1.75) = -10.85 = -11 2 significant figures

Temperature Conversions

• oF -- > oC ; C = 5/9 (F - 32)

• oC -- > oF ; F =9/5 C + 32

• oC -- > K ; K = C + 273.15• Human body temperature is 98.6 oF.

Convert this temperature to oC and K scale

• oC = 5/9 (98.6 - 32) = 5/9 (66.6) = 37.0

• oC--> K = 37.0 oC +273.15 = 310.2 K

Example. oF to K

If the temperature is 75.0 oF, what is it in K?

First convert to oC

Then convert to K

oC = (75.0oF - 32) 59

= 23.9

K = 23.9oC + 273

= 297

Measuring volumeMeasuring volume• VolumeVolume - the amount of space that an object

occupies.

• The base metric unit is the liter (L)liter (L).

• The common unit used in the lab is the milliliter (mL)milliliter (mL).

• One milliliter is exactly equal to one cmcm33.

• The derived SISI unit for volume is the mm33 which is too large for convenient use.

DensityDensity

• Density is an intensive property of a substance based on two extensive properties.

• Common units are g / cm3 or g / mL.

» g / cm3 g / cm3

• Air 0.0013 Bone 1.7 - 2.0• Water 1.0 Urine 1.01 - 1.03• Gold 19.3 Gasoline 0.66 - 0.69

Density = Mass

Volumecm3 = mL cm3 = mL

Example.Density calculation

What is the density of 5.00 mL of a fluid if ithas a mass of 5.23 grams?

d = mass / volume

d = 5.23 g / 5.00 mL

d = 1.05 g / mL

What would be the mass of 1.00 liters of thissample?

Density Calculations• Equation method:

• Density = mass ÷ volume; d = m/v

• Factor Label method:14.2 g -- > ? cm3

• conversion factor?• 2.70 g 1 cm3

• -------- or ------• 1 cm3 2.70 g

• 14.2 g x 1 cm3

• --------------------- = 5.26 cm3

• 2.70 g

Example.Density calculation

• What would be the mass of 1.00 liters of the fluid sample?

• The density was 1.05 g/mL.

» density = mass / volume

– so mass = volume x density

» mass = 1.00 L x 1000 x 1.05

» = 1.05 x 103 g

mlL

gmL

Specific gravity

• The density of a substance compared to a reference substance.

– Specific Gravity =

• Specific Gravity is unitless.

• Reference is commonly water at 4oC.

• At 4oC, density = specific gravity.

• Commonly used to test urine.

density of substancedensity of reference

Specific gravity measurement

Hydrometer

Float height willbe based onSpecific Gravity.