chemistry notes ch 3

8/3/2019 Chemistry Notes Ch 3

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Chapter 3 1

Chapter 3

Atoms: The BuildingBlocks of Matter


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Chapter 3 2

Section 3-1

The Atom: FromPhilosophical Idea toScientific Theory


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Chapter 3 3

Foundations of Atomic Theory

In 400 B.C. Democritis called natures

basic particle the atom

Aristotle, however, did not believe inatoms, but thought all matter was

continuous.

His ideas were accepted for about 2000

years.


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Chapter 3 4


In the late 1700s, during the

Enlightenment, science began to change

its ideas surrounding the structure ofmatter.

The modern definition of an element

came into general acceptance An element cannot be broken down into

simpler substances by ordinary means.


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Chapter 3 5


In the late 1790s, chemists began the

quantitative study of chemical

reactions. Law of conservation of Mass

Mass is neither destroyed nor created

during ordinary chemical or physicalreactions.


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Chapter 3 6


Law of Definite Proportions

The fact that a chemical compound

contains the same elements in exactly thesame proportions by mass regardless of

the size of the sample or source of the

compound. Example: NaCl is always NaCl


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Chapter 3 7


Law of Multiple Proportions

If two or more different compounds are

composed of the same two elements, thenthe ratio of the masses of the secondelement combined with a certain mass ofthe first element is always a ratio of small

whole numbers. Example: CO and CO2


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Chapter 3 8


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Chapter 3 9

Daltons Atomic Theory

1. All matter is composed of extremely

small particles called atoms.

2. Atoms of a given element areidentical in size, mass, and other

properties; atoms of different

elements differ in size, mass, andother properties.


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Chapter 3 10


3. Atoms cannot be subdivided, created,or destroyed.

4. Atoms of different elements combinein simple whole-number ratios toform chemical compounds.

5. In chemical reactions, atoms arecombined, separated, or rearranged.


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Chapter 3 11


Therefore, chemical reactionsinvolve the

Combination Separation or

Rearrangement of atoms


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Chapter 3 12

Modern Atomic Theory

Not all of Daltons theories arecorrect.

We know that atoms are divisibleAtoms of the same element can have

different masses


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Chapter 3 13

Section 3-2

TheS

tructure of theAtom


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Chapter 3 14

The Structure of the Atom

Dalton was incorrect that atomswere indivisible.

Atom: the smallest particle of anelement that retains the chemical

properties of that element.


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Chapter 3 15

The Structure of the Atom

The atom contains three types ofsubatomic particles

1. Protons: Positively chargedparticles, in nucleus

2. Neutrons: Neutral particles, in

nucleus3. Electrons: Negatively charged

surrounding the nucleus.


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Chapter 3 16

Discovery of the Electron

J. J. Thomson - English physicist.1897

Made a piece of equipmentcalled a cathode ray tube.

It is a vacuum tube - all the airhas been pumped out.


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Chapter 3

Thomsons Experiment

Voltage source

+-

Vacuum tube

Metal Disks


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Chapter 3

Thomsons Experiment

Voltage source

+-


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Chapter 3

Thomsons Experiment

Voltage source

+-


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Chapter 3

Thomsons ExperimentThomsons Experiment

Voltage source

+-


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Chapter 3

Passing an electric current makes aPassing an electric current makes abeam appear to move from thebeam appear to move from the

negative to the positive endnegative to the positive end


Voltage source

+-


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Chapter 3




Voltage source

+-


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Chapter 3




Voltage source

+-


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Chapter 3




Voltage source

+-


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Chapter 3

Voltage source

Thomsons Experiment

By adding an electric field


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Chapter 3

Voltage source


By adding an electric fieldBy adding an electric field

+

-


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Chapter 3

Voltage source



+

-


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Chapter 3

Voltage source



+

-


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Chapter 3

Voltage source



+

-


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Chapter 3

Voltage source



+

-


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Chapter 3

Voltage source


By adding an electric field he foundBy adding an electric field he found

that the moving pieces were negativethat the moving pieces were negative

+

-


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Chapter 3 32


Thompson determined that thecharge/mass ratio of the negative

particle was constant. Concluded that all cathode rays

are composed of identical

negatively charged particlescalled electrons.


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Chapter 3 33


The electron has a very largecharge compared with its mass

In 1909, Millikan measured themass of the electron as9.109 x 10-31 kg

Also confirmed the negativecharge of the electron


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Chapter 3

Other particles

Proton - positively charged pieces

1840 times heavier than the

electron by E. Goldstein

Neutron - no charge but the same

mass as a proton by J. Chadwick Where are the pieces?


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Chapter 3

Rutherfords experiment

Ernest Rutherford -Englishphysicist. (1910)

Believed in the plum pudding modelof the atom

Wanted to see how big they are.

Used radioactivity.


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Chapter 3


Alpha particles - positively chargedpieces- helium atoms minus

electrons Shot them at gold foil which can be

made a few atoms thick.


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Chapter 3


When an alpha particle hits a

fluorescent screen, it glows.

Heres what it looked like (page

72)


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Chapter 3

Lead

blockUranium

Gold Foil

Fluorescent

Screen


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Chapter 3

He Expected

The alpha particles to pass throughwithout changing direction very much.

Because? the positive charges were thought tobe spread out evenly. Alone they were

not enough to stop the alpha particles.


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Chapter 3

What he expected


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Chapter 3

Because


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Chapter 3

He thought the mass was evenly

distributed in the atom


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Chapter 3

Since he thought the

mass was evenly

distributed in the atom


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Chapter 3

What he got


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Chapter 3

How he explained it

+

Atom is mostly empty.

Small dense,

positive pieceat center.

Alpha particles

are deflected by

it if they get close

enough.


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Chapter 3

+


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Chapter 3

Subatomic particles p.74

Electron

Proton

Neutron

Name Symbol Charge

Relative

mass

Actual

mass (g)

e-

p+

n0

-1

+1

0

1/1840

1

1

9.11 x 10-28

1.67 x 10-24

1.67 x 10-24


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Chapter 3 48

Composition of the Atomic Nucleus

The short-range proton-neutron,proton-proton, and neutron-

neutron forces that hold thenuclear particles together arereferred to as nuclear forces.

Colossians 1:17b In Christ, allthings hold together.


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Chapter 3 49

Section 3-3

Counting Atoms


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Chapter 3 50

Objectives

Explain what isotopes are

Define atomic number and mass

number, and describe how theyapply to isotopes

Given the identity of a nuclide,

determine its number of protons,neutrons and electrons


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Chapter 3 51

Objectives

Define mole in terms ofAvogadros number and define

molar mass Solve problems involving mass in

grams, amount in moles, and

number of atoms of an element.


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Chapter 3 52

Atomic Number

Atoms of the same element have thesame numbers of protons

Atoms of different elements havedifferent numbers of protons.

The atomic number of an element is

the number of protons in the nucleusof each atom of that element.


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Chapter 3 53

Atomic Number

Elements are arranged in theperiodic table in order by atomic

number. The atomic number identifies the

element.


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Chapter 3 54

Isotopes

Isotope: Atoms of the sameelement that have different

masses. The isotopes of a particular

element have the same number of

protons and electrons but differennumbers of neutrons.


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Chapter 3 55

Mass Number

Mass Number: the total numberof protons an neutrons in the

nucleus of an isotope


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Chapter 3 56

Designating Isotopes

Isotopes are identified byspecifying their mass number.

Two methods for identifyingisotopes:

Hyphen notation: Uranium-235

Nuclear symbol


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Chapter 3 57


U23592AtomicNumber

Mass

Number


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Chapter 3 58


See sample problem 3-1 on page77.


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Chapter 3 59

Relative Atomic Masses

1 amu or atomic mass unit isexactly 1/12 the mass of a carbon-

12 atom, or The atomic mass of any nuclide is

determined by comparison with

the carbon-12 atom.

kg10660540.127

v


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Chapter 3 60

Average Atomic Masses of Elements

Most elements occur as a mixtureof isotopes

Average atomic mass: theweighted average of the atomicmasses of the naturally occurring

isotopes of an element.


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Chapter 3 61

Average Atomic Masses of Elements

Suppose you have a boxcontaining two sizes of

marbles. If 25% have massesof 2.00 g, and 75% havemasses of 3.00 g, how is the

weighted average calculated?


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Chapter 3 62

Relating mass to numbers of atoms

The mole

A mole is the amount of

substance that contains asmany particles as there are inexactly 12 g of carbon-12.


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Chapter 3 63


Avogadros Number

The number of particles in

exactly 1 mole of a puresubstance

2310022.6 v


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Chapter 3 64


Molar Mass

The mass of one mole of a pure

substance is called the molarmass of that substance

Usually written in units of g/mol

For elements, molar mass can befound in the periodic table


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Chapter 3 65

Gram/Mole Conversions

The decimal number on the periodictable is also the mass of 1 mole ofthose atoms in grams.

12.01 grams of C has the same number ofpieces as 1.008 grams of H and 55.85grams of iron.

We can write this as12.01 g C = 1 mole C


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Chapter 3 66


The decimal number on the periodictable is also the mass of 1 mole ofthose atoms in grams.

12.01 grams of C has the same number ofpieces as 1.008 grams of H and 55.85grams of iron.

We can write this as12.01 g C = 1 mole C


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Chapter 3 67


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Chapter 3 68


Work through sample problems on p.82-85 together

chemistry notes ch 3

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