H CHEM - WED, 9/7/16

Do Now

Be ready for notes.

Sigfig review problem

Homework

Error Analysis

Possibly atomic theory paragraph

Agenda

Atomic Theory

THE ATOM

DEFINITION TO START

Atom – smallest particle of an element that

retains its identity

They are tiny!

Electron microscope – allows us to observe

individual atoms

DEMOCRITUS THE PHILOSOPHER

~490 BC

First credited with proposing the existence of an

atom

Indivisible and indestructible

Shortcomings?

Did Aristotle agree?

ARISTOTLE

~380BC

Did not agree!

Earth, water, air and fire

THEN CAME JOHN DALTON

1766-1844

Experimental methods -> scientific theory

Dalton’s atomic theory

Matter is composed of indivisible atoms

Atoms of same element are identical

Combine in whole number ratios to form compounds

Rxns occur when atoms separate, bond, or rearrange.

Atoms of one element never become atoms of another

thru chemical rxns

JJ THOMSON’S EXPERIMENT -1897

Cathode Ray Experiment

Gas filled glass tube fitted with electrodes

Electricity -> cathode ray (travels from cathode to anode)

Conclusion: electrons - negatively charged subatomic particles

Further tests:

Mass to charge ratio

Different gases

Conclusion: Electrons are part of atoms of all elements.

THOMSON’S MODEL

Thomson – ‘plum pudding’ (1897)

Chocolate chips in cookie dough

MILIKAN - 1909

Oil Drop Experiment

Determined the charge of an electron

OTHER SUBATOMIC PARTICLES?

We know things are neutral…so where’s the positive

Actually discovery is disputed….

Goldstein: anode ray (1886)

Detected rays that contained positively charged particles

Rutherford: 1920

Coined proton after his work with H nuclei

Conclusion: protons – positively charged subatomic

particles

1840x mass of an electron

THOMSON’S MODEL TO RUTHERFORD’S

MODEL

Gold Foil Experiment (1911)

Alpha particles thru gold foil

Predictions: only slight deflection

Results:

Most: straight thru or slight deflection

Some: large deflection or ‘bounced’ back

toward the source

RUTHERFORD MODEL/NUCLEAR MODEL

Atom is mostly empty space

explains why the alpha particles could pass straight thru

All the positive charge and most of mass is located in a small region

explains the large deflections

Nucleus - protons and neutrons

Electrons are around the nucleus and account for most of the volume

Still not quite right!

IT’S ALL ABOUT COLOR…

In terms of atomic models, so far: Dalton (1803) = Tiny, solid particle

Thomson (1897) = “Plum Pudding” model –Electrons stuck on the outside of a big positive charge

Rutherford (1911) = Positively-charged nucleus with electrons moving around it

Rutherford’s model of the atom not quite right Could not explain chemical properties of elements

Could not explain color changes when metal is heated

BOHR MODEL OF THE ATOM-1913

Niels Bohr’s model of the atom

Electron found only on specific, circular paths around nucleus

Each orbit has fixed energy levelHypothesis: When electrons are excited (added

energy), jump into higher energy levels. When they moved back into lower energy levels - gave off light.

Electrons do not exist between levels (think of rungs on a ladder)

Electrons absorb and emit only certain quanta (amounts) of energy

Quantum of energy = fixed amount of energy required to move from one energy level to another energy level

Chapter 5

BOHR’S MODEL

Nucleus

Electron

Orbit

Energy Levels

BOHR’S PLANETARY MODEL OF THE ATOM

Electrons must have enough energy to keep moving around the nucleus

Electrons orbit nucleus in defined energy levels, just like planets orbit the sun

Each energy level assigned a principal quantum number n.

Lowest energy level called ground state (n=1)

Higher energy levels (n=2, 3, 4...) excited states

Model worked OK for hydrogen but not so good for other elements

Nucleus

n = 1n = 2

Ch

ap

ter 5

BOHR’S MODELIn

crea

sin

g e

ner

gy

Nucleus

First

Second

Third

Fourth

Fifth Further away from the

nucleus means more

energy.

There is no “in

between” energy

Energy Levels

Nucleus

Lowest energy level = ground state

Higher energy levels = excited states

Energy

Level 1

Energy

Level 2Energy

Level 3

Electron starts on lowest energy level (ground state)

Add energy to electron –moves to excited state

Energy levels are not evenly spaced

Nucleus

Electron returns to lower state –emits/gives off quantum of energy

Energy

Level 1

Energy

Level 2

Energy

Electron starts on lowest energy level (ground state)

Lowest energy level = ground state

Higher energy levels = excited states

Add energy to electron –moves to excited state

Energy levels are not evenly spaced

Energy

Level 3

Ch

ap

ter 5

Bohr used this

theory to explain

the lines in the

atomic emission

spectra for

hydrogen

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410 nm

434 nm

486 nm

656 nm

Each of these lines

corresponds to different

energy changes

So is the Bohr model correct?

For H, it does a wonderful job explaining things

For everything else….not so much

CHADWICK – THE LAST SUBATOMIC

PARTICLE

Chadwick (1932)

Chased evidence for a particle that Rutherford

predicted to exist

Neutrons – neutral subatomic particles

~same mass as proton

History lesson over….for now

WHAT CAN THE PERIODIC TABLE TELL US?

Atomic number – the number of protons in an

element

Elements are defined by their atomic number

Mass number – total number of neutrons and

protons

ISOTOPES!

Same number of protons

Different number of neutrons

Used to calculate the atomic weight of an element

Mass of an atom is tiny – so knowing the actual

mass is a bit impractical

ATOMIC MASS UNIT

Isotope Carbon 12 was assigned the mass of 12

atomic mass units.

So then, 1/12 of the mass of a carbon 12 atom is 1

amu.

So then helium 4 has 1/3 the mass of carbon 12.

Carbon has 6 protons and 6 neutrons…this

accounts for the bulk of the mass…so one proton

or 1 neutron has a mass of ~1 amu.

ATOMIC WEIGHT

Atomic weight – weighted average of the atomic

masses of the isotopes of an element in a

naturally occurring sample

Example problem….Grades!