Unit 1: Introduction to Chemistry

Chemistry 1

Edmond North High School

Chapters: 1 & 2

Scientific Method / ProcessThe Scientific Method is a systematic

approach to problem solving.It is generally composed of the

following parts:QuestionHypothesisExperimentData AnalysisConclusion

Scientific Process

Question/Hypothesis

Experiment

Data AnalysisConclusion

Observation

MeasurementMeasurement: A quantitative observation consisting

of a numeric value and units.Three are 2 kinds of units: base and derived

Base Units: Units are mutually independent of each other Ex: 7 meters, 30 kg, 6 seconds

Derived Units: Units are obtained mathematically from base units Ex: 10 cm3, 0.9 g/mL, 22 meters/second

Metric SystemThe Metric System is a decimalized system of

measurement based on powers of 10Used internationally and in the scientific communityConsists of base units and prefixes

The United States, Liberia, and Myanmar (Burma) are the only three countries who do not use the metric system

Metric System - Base Units

Mass

TimeTemperature

Volume

Distance

grams (g)

meters (m)

liters (L)

Kelvin (K)seconds (s)

Accuracy & Precision

Accuracy: the degree of agreement between the true value and the measured value (bullseye)

Precision: the degree of agreement among several measured values (grouping)

Accuracy & PrecisionThree different groups of students measure the mass

of a medal, with a known value of 5.000 grams. Evaluate each group’s data for its accuracy and precision (low or high):

Trial 1 5.003 g

Trial 2 5.002 g

Trial 3 5.001 g

Trial 1 5.400 g

Trial 2 5.202 g

Trial 3 5.905 g

Trial 1 5.503 g

Trial 2 5.499 g

Trial 3 5.501 g

Group 2Group 1 Group 3

Accuracy ______Precision ______

Accuracy ______Precision ______

Accuracy ______Precision ______

high

high

low

low high

low

Uncertainty in MeasurementAll measuring instruments have

a degree of uncertainty. The more divisions a device has, the more accurate the measurement.

The last digit of a measurement is always estimated (uncertain).

How would you read the volume in this graduated cylinder if the markings showed every 5 mL instead of every 1 mL?

read from the bottom of the meniscus

Uncertainty in MeasurementIf we measure the length of the paw print with a

decimeter ruler, we know for a certainty that it is between 0 and 1 decimeters. We estimate the next digit:

0.3 decimetersestimated digit

Uncertainty in MeasurementIf we measure the length of the paw print with a

centimeter ruler, we know for a certainty that it is between 3 and 4 centimeters. We estimate the next digit:

3.5 centimetersestimated digit

Uncertainty in MeasurementIf we measure the length of the paw print with a

millimeter ruler, we know for a certainty that it is between 34 and 35 millimeters. We estimate the next digit:

34.5 millimetersestimated digit

Significant FiguresSignificant Figures are the digits in a measurement that

are known with some degree of certainty are called significant figures.The number of significant figures in a measurement = the

number of digits that are known + the estimated digitThe more significant figures after the decimal, the more

accurate the measurement.

Ex: This triple beam balance shows a mass of 62.41 grams. There are 4 significant figures in this measurement.

Significant Figure Rules

There are rules for counting significant figures related to:Non-zero integersZeroes

Leading zeroes Captive zeroes Trailing zeroes

Significant Figure Rules

Non-zero integers are always significant

3456 cm has

4 sig figs

Significant Figure Rules

Leading zeroes are never significant

0.0486 g has

3 sig figs

Significant Figure Rules

Captive zeroes are always significant

16.07 mL has

4 sig figs

Significant Figure Rules

Trailing zeroes are only significant if the number contains a decimal

9.30 m has

3 sig figs 930 m has

2 sig figs

Significant Figures Practice How many significant figures are in the following

measurements?

1. 8,675,309 km ____

2. 90,210 L ____

3. 0.07 mg ____

4. 0.2020 daL ____

5. 300.00 g ____

7

4

14

5

Significant Figure OperationsMultiplication & Division

Calculate “raw” answerRounded answer must contain no more significant

figures than the measurement with the least number of significant figures

Example: What is the density of a bar of gold with a mass of 87.82 g and a volume of 4.55 cm3? (Density = mass volume)

87.82 g 4.55 cm3 = 19.301099 g/cm3 19.3 g/cm3

4 sig figs 3 sig figs raw answer 3 sig figs

Significant Figure OperationsAddition & Subtraction

Calculate “raw” answerRounded answer can have no more digits to the right of

the decimal point than the measurement with the least number of digits to the right of the decimal point

Example: What is the difference in length between a professional shot put throw of 23.125 meters and an amateur shot put throw of 21.2 meters?

23.125 m 21.2 m = 1.925 m 1.9 m 3 digits 1 digit raw 1 digit

after decimal after decimal answer after decimal

Exact NumbersExact Numbers are different from

measurements because they have no uncertainty

Significant figures does not applyExamples

Conversions (3 feet = 1 yard)Counting Numbers (20 M&Ms)

Metric System - Base Units

Mass

TimeTemperature

Volume

Distance

grams (g)

meters (m)

liters (L)

Kelvin (K)seconds (s)

Metric System - PrefixesPrefix Symbol Numeric Representation

tera- T 1012 1,000,000,000,000

giga- G 109 1,000,000,000

mega- M 106 1,000,000

kilo- k 103 1,000

hecto- h 102 100

deka- da 101 10

(base unit) (none) 100 1

deci- d 10-1 0.1

centi- c 10-2 0.01

milli- m 10-3 0.001

micro- 10-6 0.000001

nano- n 10-9 0.000000001

pico- p 10-12 0.000000000001

Powers of Ten

From Quarks to Outer Space

Metric ConversionsExample: How many kilograms are in 75 decigrams?

Example: How centiliters are in 12 dekaliters?

Move decimal 4 places to the left: 75 dg = 0.0075 kg

Move decimal 3 places to the right: 12 daL = 12,000 cL

(start)

k h da d c m base x x

k h da d c m base x x

(start)

Metric Conversions Practice

Practice metric conversions:

1. _______ dg = 4.2 hg2. _______ dam = 6,055 mm3. _______ L = 1 cL4. _______ cK = 0.003 kK5. _______ s = 11,700,000 s

k h da d c m base x x

4,2000.6055

0.0130011.7

Remember

King Henry died by drinking chocolate milk.

Kilo Hecto Deka Base Deci, Centi, and Milli.

DensityDensity is the ratio of an

object’s mass and volumeThe formula for density is D =

m / vIn chemistry, the 2 most

common units of density will be g/mL and g/cm3

The four cubes to the right have the same volume (1 cm3), but different masses. How does this effect their densities?

DensityThe density of an object is an intensive physical

property, meaning it cannot be changed no matter the quantity.If the property changes with the amount of substance

present then it is an extensive property.The density of water is 1 g/mL

Objects that float < 1 g/mLObjects that sink > 1 g/mL

What happens to the density of an object when it is sawed I half?

Counting MatterWhat are some common ways we

count matter?

Dozen = 12

Ream = 500

Gross = 144

Mole = 6.02 x 1023

The MoleIn the same way a dozen is worth 12,

a mole is worth 6.02 x 1023

This number is called Avogadro’s number

Mole is abbreviated as “mol”Written in expanded form, that

number is: 602,000,000,000,000,000,000,000The mole is a large number because

particles are so small, it takes many of them make up an amount we can see and understand

We can use it to count anything!

1 dozen cookies = 12 cookies1 mole of cookies = 6.02 X 1023 cookies

1 gross cars = 144 cars1 mole of cars = 6.02 X 1023 cars

1 ream Al particles = 500 Al particles1 mole of Al particles = 6.02 X 1023 particles

The Mole

Scientific NotationScientific Notation is mathematical shorthand that

makes large and small numbers manageableIt is composed of three parts:

CoefficientBaseExponent

4.56 x 10-7Coefficient Base

Exponent

Scientific Notation RulesCoefficients

Coefficients must be greater than or equal to 1 and less than 10

Coefficients can be positive or negativeAll numbers in the coefficient are counted as significant

Which of the following numbers are written incorrectly?

22 x 105 9.5 x 102 10 x 108

7 x 10-3 0.3 x 10-9 -1.00 x 106

Scientific Notation RulesBases

Base is always a 10Bases are never counted as significant

ExponentsExponents are always integersExponents can be positive (big number) or negative (small number)

Ex: 1 x 103 = 1000 and 1 x 10-3 = 0.001Exponents are never counted as significant

Which of the following numbers are written incorrectly?

4.1 x 10-5 2.2 x 620 1.0 x 10-1

7 x 100 0.15 x 10-9 -7 x 106.3

Scientific NotationHow to enter the number 2.5 x 10-8

into the calculator:Enter the coefficient 2.5

Press 2nd, then EE

Enter the exponent -8

It should appear on your screen as 2.5E-8

Mole Calculations6.02 x 1023 particles = 1 mole

Ex: How many particles are in 3.00 moles of N2?

3.00 mol N2 6.02 x 1023 particles =

1 mol N2

Ex: How many moles of Na are in 1.10 x 1023 particles?

1.10 x 1023 particles Na 1 mol Na =

6.02 x 1023 particles Na

1.81 x 1024 particles

1.83 x 1022 moles

Classifying Matter

Pure SubstancesPure substances cannot be separated by physical means

Elements: cannot be chemically separated, listed on the periodic table

Compounds: can be chemically separated, made up of elements

carbon (C)

sulfur (S)

copper (Cu)

mercury (Hg)

salt (NaCl) water (H2O) sugar (C6H12O6) rust (Fe2O3)

Pure SubstancesParticle representations of…

Elements

Compounds

Periodic TableEach square on the

periodic table contains:Name of ElementSymbol

First letter is upper case, second letter must be lower case

Atomic Mass Number (usually with a

decimal) indicates the mass (g) of 1 mole of that element

Atomic Number

Types of Elements

Groups/Familes

Period/Series

The Mole

Molar Mass of AtomsHow do chemists “count” Avogadro’s number?By measuring mass (just like how they count

aluminum cans for recycling)The mass (think grams) of one mole of a substance

Atomic masses are based on the mass of carbonOne mole of carbon contains 6.02 x 1023 atoms of C

1 carbon atom is 12.011 amu 1 mole of carbon is 12.011 grams

Molar Mass of AtomsThe mass of 1 mole (in grams)Equal to the numerical value of the average atomic

mass (get from periodic table)Examples:

1 mole of C atoms = 12.0 g 1 mole of Mg atoms = 24.3 g 1 mole of Cu atoms = 63.5 g

PracticeFind the molar mass

1 mole of Au atoms =

1 mole of Sn atoms =

Chemical Formulas• The chemical formula describes a compound using

numbers and element symbols– Subscripts denote number of each element in a

compound• Ex: H2O has 2 hydrogens and 1 oxygen

• Ex: Al2(SO4)3 has 2 aluminums, 3 sulfurs, and 12 oxygens

– Coefficients act as a scalar (also called a multiplier)• Ex: 6H2O has 12 hydrogens and 6 oxygens

• Ex: 3Al2(SO4)3 has 6 aluminums, 9 sulfurs, and 36 oxygens

Naming Covalent CompoundsTwo words, with prefixesPrefixes tell you how many.

mono, di, tri, tetra, penta, hexa, septa, nona, decaFirst element whole name with the appropriate prefix,

except monoSecond element, -ide ending with appropriate prefix

Naming Covalent CompoundsCO2

CO CCl4

N2O4

Diflourine monoxideNitrogen trichlorideDiphosphorus pentoxide

Three Types of Ionic Compounds (Salts)We will learn how to write formulas and

nomenclature (naming system) for:Binary Salts

2 elements: a metal and a nonmetalTernary Salts

3 or more elements Includes a polyatomic ion

Salts with Multiple Oxidation Numbers Can be binary or ternary Includes a transition metal

Binary Ionic CompoundsA Binary Ionic Compound contains 2 elements: metal &

nonmetalWhen you write a formula, you add subscripts to the symbols

for the ions until the algebraic sum of the ions’ charges is zero.

Naming:Name the metalDrop the end of nonmetal and add “ide”

Example: MgCl2 = magnesium chloride

Molar Mass of CompoundsThe Molar Mass of a compound is the mass in grams

of 1 mole equal to the sum of the atomic massesExample: What is the mass of 1 mole of CaCl2?

1 mole Ca x 40.1 g/mol = 40.1 g Ca2 moles Cl x 35.5 g/mol = + 71.0 g Cl

111.1 g/mol CaCl2

Practice

Prozac, C17H18F3NO, is a widely used anti-depressant that inhibits the uptake of serotonin by the brain. Find its molar mass.

Calculations with Molar Mass Examples

How many grams are in 2 moles of Cu?2 moles of Cu = Xg1 mole of Cu 63.546 g of Cu

How many moles are in 100 g of H2O?

100 g of H2O = Xg

18.015 g of H2O 1 mole of H2O

Practice

The artificial sweetener aspartame (C14H18N2O5), commonly known as Nutra-Sweet, is used to sweeten diet foods, coffee and soft drinks. How many moles of aspartame are present in 225 g of aspartame?

Calculations with Particles/Mass

ExamplesHow many grams of CO2 are present in 3 particles of

CO2?

3 particles of CO2 = Xg

6.02 x 1023 particles of CO2 44.009 g of CO2

How many particles are present in 12.0 grams of NaCl?

12.0 g of NaCl = X particles

58.443 g of NaCl 6.02 x 1023 particles of NaCl

PracticeDiabetics do not properly

regulate insulin, which tells the body to take up glucose from the blood. What is the mass of 1.20 x 1024 molecules of glucose (C6H12O6)?

Diatomic ElementsSome elements are always found as a pair (2 together)

H2, O2, N2, F2, Cl2, Br2, I2

Try to remember HOF BrINCl

Example:How many atoms of O are present in 78.1 g of oxygen?

Practice:How many grams of Cl are present in 100 atoms of

chlorine?

78.1 g O2 = Xg 32.0 g O2 2(6.02 X 1023 atoms O)

Classifying Matter

MixturesMixtures are composed of two or more substances

that can be separated by physical means; contains elements and/or compoundsHeterogeneous Mixtures: not uniform throughoutHomogeneous Mixtures: uniform throughout

Particle representations of mixtures…

Heterogeneous MixturesHeterogeneous Mixture:

two or more substances physically combined; not uniform throughout

Ex: Granite, chex mix

SuspensionsA suspension is a

heterogeneous fluid containing solid particles that are sufficiently large for sedimentation.

The internal phase (solid) is dispersed throughout the external phase (fluid) through mechanicalSuspensions will

eventually settle.

ColloidsColloids are suspensions of particles larger than

individual ions or molecules, but too small to be settled out by gravity.

ColloidsThe Tyndall Effect is

scattering of light particles used to determine whether something is a solution or suspension

Click image to play clip

Homogeneous MixturesHomogeneous Mixture:

two or more substances physically combined; uniform throughoutKnow as a solution

composed of… Solute: substance being

dissolved (smaller amount) Solvent: substance that

does the dissolving (larger amount)

Examples: kool-aid, brass

Separating a MixtureSeparating a mixture -

components are separated without changing their physical identityManual SeparationMagnetismFiltrationEvaporationDistillationCentrifugingChromatography

Manual Separation

DecantingSeparates two liquids of

different densities Sifting

Separates two solids of different particle size

by pouringSorting

Separates two solids by picking

MagnetismSeparates metals (such as

iron) from a mixture

FiltrationSeparates solid

substances from liquids and solutions

EvaporationSeparates a dissolved

solid from its solvent

DistillationSeparates homogeneous

mixture with different boiling points (heat mixture and catch condensed vapor)

CentrifugingSeparates heavier particles

(bottom of tube) from lighter particles (top of tube) by spinning them at high speeds

ChromatographySeparates substances on the basis of their

differences in solubility in a solvent - different substances are attracted to paper or gel and move at different speeds

Properties of MatterA property is a characteristic that describe matter or

how it behavesPhysical – determined without a chemical change

Extensive – vary with amount of matter (Ex. Mass, volume, length, area)

Intensive – does not vary with amount (Ex. Density, color, odor, melting point, solubility)

Chemical – can only be determined by a chemical change Ex. Flammability, reactivity with acid, stability of a

compound (how easily it decomposes)

EnergyEnergy is the ability to do workHeat: Energy used to cause the temperature of an

object to rise.Temperature - measure of the average kinetic energy

of the particlesWe measure temperature with a thermometer.

Heat (Thermal) EnergyHeat is a form of energy that flows between two

samples of matter because of their difference in temperatureHeat flows from hot to coldIt can be absorbed or releasedMeasured in units of calories orJoules

Changes of MatterPhysical Changes – do not involve a change in chemical

identityEx: boiling, freezing, melting, dissolving, evaporating, and

crystallizing

Changes of MatterChemical Changes –

new substances are formed in the reactionEx: iron rusting, copper

oxidizing, wood burning, silver tarnishing

Evidence for Changes of Matter

Color ChangeGas produced without

heatingPrecipitate formedNew odor developsLarge amount of heat or

light produced

The End of Unit 1.

Be Prepared for Test.