Ch01 Chemistry and the Atomic Molecular View of Matter

Chapter 1: Chemistryand the

Atomic/MolecularView of Matter

Chemistry: The Molecular Nature of Matter, 6E

Jespersen/Brady/Hyslop

Jespersen/Brady/Hyslop Chemistry: The Molecular Nature of Matter, 6E

Why Study Chemistry? In every aspect of our modern life

Long life batteries Materials & miniaturization

Cell phones/pagers Laptops

Synthetic fibers Dyes CDs/DVDs—silicon wafers Medications DNA sequencing

Touches all areas of science2


Chemistry and the SciencesChemistry Study of matter & its transformations Seeks answers to fundamental

questions about: What makes up materials that compose our

world How composition affects properties of

substances How substances change when they interact

with each other = Chemical Reactions

3


Chemistry and the SciencesChemistry Seeks to understand:

Underlying structures of matter Forces that determine properties that we

observe Apply this knowledge to:

Create new materials not found in nature Understand fundamental biological

processes

4


Scientific Method Approach to gathering information &

formulating explanations. Scientists perform experiments in

laboratories under controlled conditions1.Make observations/collect data

Empirical fact Something we see, hear, taste, feel, or smell Something we can measure in laboratory Organize data so we can see relationships

5


Scientific Method2. Law or Scientific Law

Broad generalization Based on results of many experiments Only states what happens Doesn’t explain why they happen

3. Hypothesis Mental picture that explains observed laws Tentative explanation of data Make predictions Devise experiments to test Go back to laboratory & perform

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Scientific Method4. Theory

Tested explanation of how nature behaves

Devise further tests Depending on

results, may have to modify theory

Can never prove theory is absolutely correct

7


Scientific MethodEx. Study gases Discover Volume (V) of gas depends on

Pressure (P) Temperature (T) Amount (n)

Data Recorded observations of relationship between V, P, T &

n Law

R = constant Kinetic Theory of Gases

Explains gas behavior (Ch 11)

8

nRTPV


Atomic Theory Most significant theoretical model of

natureAtoms

Tiny submicroscopic particles Make up all chemical substances Make up everything in Macroscopic world Smallest particle that has all properties of

given element Composed of:

Electrons Neutrons Protons

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Matter & Its ClassificationsMatter

Anything that has mass & occupies spaceMass

How much matter given object has Measure of object’s momentum, or resistance

to change in motionWeight

Force with which object is attracted by gravityEx. Mass vs. Weight

Astronaut on moon & on earth Weight on moon = 1/6 weight on earth Same mass regardless of location

10


MatterChemical Reactions

Transformations that alter chemical compositions of substances

Decomposition Chemical reaction where 1 substance

broken down into 2 or more simpler substances

Ex.

11

Molten sodium chloride

Sodium metal

+ chlorine gas

Electric current


Elements Substances that can’t be decomposed into

simpler materials by chemical reactions Substances composed of only 1 type of atom Simplest forms of matter that we can work

with directly More complex substances composed of

elements in various combinations

12diamond = carbon gold sulfur


Chemical Symbols for ElementsChemical Symbol

One or two letter symbol for each element name First letter capitalized, second letter lower case

Ex. C = carbon S = sulfur Ca = calcium Ar = argon Br = bromine H = hydrogen Cl = chlorine O = oxygen

Used to represent elements in chemical formulas

Ex. Water = H2O Carbon dioxide = CO2

Most based on English name Some based on Latin or German names

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Chemical SymbolsEnglish Name

Chemical Symbol

Latin Name

Sodium Na NatriumPotassium K KaliumIron Fe FerrumCopper Cu CuprumSilver Ag ArgentumGold Au AurumMercury Hg HydrargyrumAntimony Sb StibiumTin Sn StanniumLead Pb PlumbumTungsten W Wolfram

(German) 14


Compound Formed from 2 or more atoms of

different elements Always combined in same fixed ratios

by mass Can be broken down into elements by

some chemical changes Ex. Water decomposed to elemental

hydrogen & oxygenMass of oxygen =

8 × mass of hydrogen

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Pure Substance vs. MixturePure substances

Elements and compounds Composition always same regardless of source

Mixture Can have variable compositions Made up of two or more substancesEx. CO2 in water—varying amounts of “fizz” in

soda 2 broad categories of mixtures:

Heterogeneous Homogeneous

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Homogeneous Mixtures Same properties throughout sample Solution

Thoroughly stirred homogeneous mixture Ex. Liquid solution

Sugar in water Gas solution

Air Contains nitrogen, oxygen, carbon

dioxide & other gases Solid solution

US 5¢ coin – Metal Alloy Contains copper & nickel metals

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Is honey a mixture?

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Heterogeneous Mixtures 2 or more regions of different properties Solution with multiple phases Separate layersEx.

Salad dressing Oil & vinegar

Ice & water Same composition 2 different physical states

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Physical Change No new substances formed Substance may change state or the

proportionsEx. Ice melting Sugar or salt dissolving Stirring iron filings & sulfur together

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Chemical Changeor Chemical Reaction

Formation of new substance or compound Involves changing chemical makeup of

substances New substance has different physical

properties Can’t be separated by physical means

Ex. Fool’s gold Compound containing sulfur & iron

No longer has same physical propertiesof free elements

Can’t be separated using magnet21


Learning Check:

Chemical

Physical

Magnesium burns when heated

Magnesium metal tarnishes in air

Magnesium metal melts at 922 K

Grape Kool-aid lightens when water is added

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For each of the following, determine if it represents a Chemical or Physical Change:

XX

XX


Classification of Matter

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Learning Check: ClassificationHot

CocoaIce(H2O

)

White Flour

Table Salt (NaCl)

Pure substanceElementCompoundMoleculeHeterogeneous

MixtureHomogeneous

Mixture

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XXX

X X

X

X


Law of Definite Proportions In given compound, elements always

combine in same proportions by mass. Ratio of masses of each element is fixed

for given compound Implication:

Each atom has fixed specific mass Ex. Fool’s gold, pyrite, iron (III) sulfide

Mass ratio always 1.00 g of Iron to 0.574 g of Sulfur

Ex. Water Mass ratio always: 8 g O to 1 g H

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Law of Conservation of Mass No detectable gain or loss of mass

occurs in chemical reactions. Mass is conserved.Implication: Reactions involve reorganization of

materials.

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Learning Check: Chemical Laws Magnesium burns in oxygen to form

magnesium oxide. If 16.88 g of Mg are consumed and 28.00 g of MgO are produced, what mass of oxygen was consumed?

28.00 g – 16.88 g = 11.12g O

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Learning Check: Chemical Laws In a sample of MgO, there are 16.89 g Mg and

11.11 g O. What mass of O would there be in a sample that contains 2.00 g of Mg?

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X = 1.32 g O


Dalton’s Atomic TheoryJohn Dalton Developed underlying theory to explain

Law of Conservation of Mass Law of Definite Proportions

Reasoned that if atoms exist, they have certain properties

Dalton’s Atomic Theory1.Matter consists of tiny particles called atoms.

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Dalton’s Atomic Theory (cont)2. Atoms are indestructible.

In chemical reactions, atoms rearrange but do not break apart.

3. In any sample of a pure element, all atoms are identical in mass & other properties.

4. Atoms of different elements differ in mass & other properties.

5. In given compound, constituent atoms are always present in same fixed numerical ratio.

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Proof Of Atoms Early 1980’s, use

Scanning Tunneling Microscope (STM)

Surface can be scanned for topographical information

Image for all matter shows spherical regions of matter Atoms

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STM of palladium


How Do We Visualize Atoms? Atoms represented by

spheres Different atoms have

different colors Standard scheme given in

Fig. 1.11 is represented on the right.

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Molecules Atoms combine to form more complex

substances Discrete particles Each composed of 2 or more atomsEx.

Molecular oxygen, O2

Carbon dioxide, CO2 Ammonia, NH3 Sucrose, C12H22O11

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Chemical Formulas Specify composition of substance Chemical symbols

Represent atoms of elements present Subscripts

Given after chemical symbol Represents relative numbers of each type of

atomEx.

Fe2O3 : iron & oxygen in 2:3 ratio

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Chemical FormulasFree Elements

Element not combined with another in compounds

Just use chemical symbol to representEx. Iron Fe Neon Ne

Sodium Na AluminumAl

Diatomic Molecule Molecules composed of 2 atoms each Many elements found in nature

Ex. Oxygen O2 Nitrogen N2

Hydrogen H2 Chlorine Cl235


Depicting Molecules Want to show:

Order in which atoms are attached to each other

3-dimensional shape of molecule Three ways of visualizing molecules:

1. Structural formula2. Ball-and-Stick model3. Space filling model

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1. Structural Formulas Use to show how atoms are attached

Atoms represented by chemical symbols Chemical bonds attaching atoms indicated

by lines

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H O H H C H

H

HH2Owater

CH4methane


3-D Representations of Molecules

Use touching spheres to indicate molecules Different colors indicate different elements Relative size of spheres reflects differing sizes

of atoms

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Hydrogen

molecule,

H2

Oxygen molecule

,O2

Nitrogen molecule

N2

Chlorine molecule,

Cl2


2. “Ball-and-Stick” Model Spheres = atoms Sticks = bonds

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Chloroform, CHCl3

Methane, CH4


3. “Space-Filling” Model Shows relative sizes of atoms Shows how atoms take up space in molecule

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Chloroform, CHCl3

Methane CH4

Water H2O


More Complicated Molecules Sometimes formulas contain

parentheses How do we translate into a structure?Ex. Urea, CO(NH2)2

Expands to CON2H4 Atoms in parentheses appear twice

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Ball-and-stick model Space-filling

model


Hydrates Crystals that contain water moleculesEx. plaster: CaSO4∙2H2O calcium sulfate dihydrate

Water is not tightly held Dehydration

Removal of water by heating Remaining solid is anhydrous (without water)

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Blue = CuSO4

•5H2OWhite = CuSO4


Counting Atoms1. Subscript following chemical symbol

indicates how many of that element are part of the formula

No subscript implies a subscript of 1.2. Quantity in parentheses is repeated a

number of times equal to the subscript that follows.

3. Raised dot in formula indicates that the substance is a hydrate

Number preceding H2O specifies how many water molecules are present.

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Counting AtomsEx. 1 (CH3)3COH Subscript 3 means 3 CH3 groups

So from(CH3)3, we get 3 × 1C = 3C 3 × 3H = 9H

#C = 3C + 1C = 4 C#H = 9H + 1H = 10 H#O = 1 OTotal # of atoms = 15 atoms

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Counting AtomsEx. 2 CoCl2 · 6H2O The dot 6H2O means you multiple both

H2 & O by 6 So there are:

#H 6 × 2 = 12 H#O 6 × 1 = 6 O#Co 1 × 1 = 1 Co#Cl 2 × 1 = 2 ClTotal # of atoms = 21 atoms

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Your Turn!

a. Na2CO3

b. (NH4)2SO4

c. Mg3(PO4)2

d. CuSO4∙5H2Oe. (C2H5)2N2H2

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a. ___Na, ___ C, ___ O

b. ___N, ___H, ___S, ___Oc. ___Mg, ___P, ___Od. ___Cu, ___S, ___O, ___He. ___C, ___H, ___N

32 12 8 1 43 2 81 1 9 10

Count the number of each type of atom in the chemical formula given below

4 12 2


Dalton’s Atomic Theory We now have the tools to explain this

theory & its consequences All molecules of compound are alike &

contain atoms in same numerical ratio.Ex. Water, H2O

Ratio of oxygen to hydrogen is 1 : 21 O atom : 2 H atoms in each moleculeO weighs 16 times as much as H1 H = 1 mass unit1 O = 16 mass units

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Atoms in Fixed Ratios by Mass

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For water in general: mass O = 8 mass H Regardless of amount of water present


Dalton’s Atomic TheorySuccesses: Explains Law of Conservation of Mass

Chemical reactions correspond to rearranging atoms.

Explains Law of Definite Proportions Given compound always has atoms of same

elements in same ratios. Predicted Law of Multiple Proportions

Not yet discovered Some elements combine to give 2 or

more compoundsEx. SO2 & SO3

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Law Of Multiple Proportions When 2 elements form more than one

compound, different masses of one element that combine with same mass of other element are always in ratio of small whole numbers. Atoms react as complete (whole)

particles. Chemical formulas

Indicate whole numbers of atoms Not fractions

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Using Law Of Multiple Proportions sulfur sulfur

dioxide trioxideMass S 32.06 g 32.06 gMass O 32.00 g 48.00 g

Use this data to prove law of multiple proportions

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Law of Multiple ProportionsCompound Sample

SizeMass of Sulfur

Mass of Oxygen

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23

00360048

2

3 gg

SOSO

.

. in O in O

Ratio of

Sulfur dioxide

64.06 g 32.06 g

Sulfur trioxide

80.06 g

32.06 g

32.06 g 48.00 g


Molecules Small and Large So far we’ve only discussed small

molecules Some are very large, especially those

found in nature Same principles apply to allEx. DNA - short segment

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How Do We Know Formulas? Hardly “out of the blue” Don’t know formula when compound 1st

isolated Formulas & structures backed by

extensive experimentation Use results of experiments to determine

Formula Chemical reactivity

Molecular Shape Can speculate once formula is known Determine from more experiments

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Visualizing Mixtures Look at mixtures at atomic/molecular level Different color spheres stand for 2

substances a. Homogeneous mixture/solution – uniform

mixingb. Heterogeneous mixture – 2 phases

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a. b.


Chemical Reactions When 1 or more substances

react to form 1 or more new substances

Ex. Reaction of methane, CH4, with oxygen, O2, to form carbon dioxide, CO2, & water, H2O.Reactants = CH4 & O2

Products = CO2 & H2O How to depict?

Words too long Pictures too awkward

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Chemical Equations Use chemical symbols & formulas to

represent reactants & products. Reactants on left hand side Products on right hand side Arrow () means “reacts to yield”

Ex. CH4 + 2O2 CO2 + 2H2O Coefficients

Numbers in front of formulas Indicate how many of each type of

molecule reacted or formed Equation reads “methane & oxygen

react to yield carbon dioxide & water”57


Conservation of Mass in Reactions Mass can neither be created nor destroyed

This means that there are the same number of each type of atom in reactants & in products of reaction If # of atoms same, then mass also same

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CH4 + 2O2 CO2 + 2H2O 4 H + 4O + C = 4 H + 4O

+ C


Balanced Chemical Equation

Ex. 2C4H10 + 13O2 8CO2 + 10H2O

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4 C & 10 H per

molecule2 O per molecul

e

2 H & 1 O per

molecule1 C & 2 O

per molecule

Subscripts Define identity of substances Must not change when equation

is balanced


Balanced Chemical EquationEx. 2C4H10 + 13O2 8CO2 + 10H2O

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2 molecules of C4H10

13 molecules of O2

10 molecules of C4H10

8 molecules of CO2Coefficients

Number in front of formulas Indicate number of molecules of each type Adjusted so # of each type of atom is

same on both sides of arrow Can change


Balanced Chemical Equations How do you determine if an equation is

balanced? Count atoms Same number of each type on both sides of

equation? If yes, then balanced If no, then unbalanced

Ex. 2C4H10 + 13O2 8CO2 + 10H2O Reactants Products2×4 = 8 C8×1 = 8 C2×10 = 20 H 10×2 = 20 H13×2 = 26 O (8×2)+(10×1)= 26 O

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Learning CheckFe(OH)3 + 2 HNO3 Fe(NO3)3 + 2 H2O

Not Balanced Only Fe has same number of atoms

on either side of arrow. 62

Reactants ProductsFe 1 1

3 + (2×3) = 9 (3×3) + 2 = 11O3 + 2 = 5 (2×2) = 4H

2 3N


Your Turn!How many atoms of each element appear on each side of the arrow in the following equation? 4NH3 + 3O2 → 2N2 + 6H2O

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Reactants ProductsN (4 × 1) = 4 (2 × 2) = 4O (3 × 2) = 6 (6 × 1) = 6H (4 × 3) = 12 (6 × 2) = 12


Your Turn!Count the number of atoms of each element on both sides of the arrow to determine whether the following equation is balanced.2(NH4)3PO4 + 3Ba(C2H3O2)2 → Ba3(PO4)2 + 6NH4C2H3O2

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Reactants ProductsN (2 × 3) = 6 (6 × 1) = 6H (2×3×4)+(3×3×2) =

42(6×4) + (6×3) = 42O (2×4) + (3×2×2) =

20(2×4) + (6×2) = 20P (2 × 1) = 2 (2 × 1) = 2

Ba (3 × 1) = 3 (3 × 1) = 3

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Ch01 Chemistry and the Atomic Molecular View of Matter