Download CHEM Notes Unit 4 History of atomic theory Monday Oct 7 Greek

CHEM
Unit 4
Monday Oct 7
Greek philosophers
Leucippus and
Democritus
Aristotle
Skipping ahead – not on
the test…
Robert Boyle (about
1660)
Jump to around 1800,
give or take 20 years…
Antoine Lavoisier
Notes
History of atomic theory
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about 400 BC
Philosophers REASON, they don’t EXPERIMENT
Two groups of philosophers
o atomists
o nonatomists
 4 elements: earth, water, air, fire
 Leaders of the atomists
 If matter is divided enough times it reaches a point
where it is made of indivisible particles – atoms
 Atoms of different elements have different sizes and
shapes which make them behave differently
 Leader of the nonatomists
 Liked some things about the idea of atoms but had two
big problems – unanswered questions:
o What holds atoms together?
o What’s in between atoms? Aristotle believed
there could not be a location where there is
nothing.
Both sets of ideas were known and studied for the next 2000
years. HOWEVER, there weren’t many schools, not many
people could even read, and most people who studied matter
were ALCHEMISTS – studied matter more as if it was magic –
they weren’t scientists.
 one of the first true scientists as opposed to alchemists
(mostly)
 Wrote book – The Sceptical Chymist
 Two important things in this book
o Defined element as matter that is made of
fundamental, simplest particles that cannot be
simplified any more (elements can’t be broken
down by any chemical or physical change)
o Said science was only valuable when every
scientist shares not only his experimental results
but the process he used to get them and his
numerical results
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Farther of modern chemistry
Wrote the first true chemistry textbook – Elements of
Chemistry
Identified the law of conservation of mass – mass can’t be
created or destroyed; in any chemical or physical change,
the mass you start with, you end with
CHEM
Unit 4
Notes
History of atomic theory
Lavoisier burns a
diamond with a
magnifying glass
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Joseph Proust
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Two ways to
calculate %
composition:
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Identified the law of constant composition (definite
proportion)
Any sample of the same compound will always contain the
same elements in the same proportion by mass
The percent by mass (mass percent) does not change.
Mass percent can be calculated two ways: using the
compound’s formula and the PT OR measuring the
masses of elements in the compound in lab. You get the
same answers.
Use data collected in an experiment.
o From ex. above: % comp. of MgO
𝑚𝑎𝑠𝑠 𝑀𝑔
o %𝑀𝑔 = 𝑚𝑎𝑠𝑠 𝑀𝑔𝑂 × 100
o %𝑀𝑔 =
o %𝑂 =
o %𝑂 =

5.00 𝑔
8.29 𝑔
𝑚𝑎𝑠𝑠 𝑂
× 100 = 60.3% 𝑀𝑔
𝑚𝑎𝑠𝑠 𝑀𝑔𝑂
3.29 𝑔
8.29 𝑔
× 100
× 100 = 39.7 % 𝑂
Use masses from PT
o MgO contains 1 atom Mg and 1 atom O per
molecule
o Mass of 1 atom Mg = 24.305 u
o Mass of 1 atom O = 15.9994 u
o Mass of 1 molecule MgO =
mass 1 atom Mg + 1 atom O =
24.305 u + 15.9994 u = 40.304 u
𝑚𝑎𝑠𝑠 𝑀𝑔
o %𝑀𝑔 = 𝑚𝑎𝑠𝑠 𝑀𝑔𝑂 × 100
o %𝑀𝑔 =
o %𝑂 =
o %𝑂 =
24.305 𝑢
40.304 𝑢
𝑚𝑎𝑠𝑠 𝑂
𝑚𝑎𝑠𝑠 𝑀𝑔𝑂
15.9994 𝑢
40.304 𝑢
× 100 = 60.3% 𝑀𝑔
× 100
× 100 = 39.7 % 𝑂
CHEM
Unit 4
John Dalton
Math of law of multiple
proportions (you don’t
have to DO it, but you do
need to understand it)
Dalton’s atomic theory
(*essay question on test)
Notes
History of atomic theory
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Identified the law of multiple proportions
If you have two different compounds that contain exactly
the same elements, AND you take a sample of each
element so that both elements contain exactly the same
mass of element A, and then you look at the ratio of the
mass of element B in compound 1 to the mass of element
B in compound 2, the ratio is ALWAYS a simple wholenumber ratio.
 First person to measure and publish a table of atomic
weights – got the idea from his law of multiple proportions.
 Developed the first complete atomic theory – ideas about
atoms that EXPLAINED how matter behaved.
Copper and oxygen form two compounds:
 copper (I) oxide, Cu2O
 copper (II) oxide, CuO
If you take the amount of each compound that contains 10.00
g copper:
 The sample of Cu2O has a mass of 11.26 g. 10.00 g of
this is copper; 1.26 g is oxygen.
 The sample of CuO has a mass of 12.52 g. 10.00 g of this
is copper; 2.52 g is oxygen.
Compare the masses of oxygen in each sample as a ratio:
1.26 g O : 2.52 g O
Simplify the ratio by dividing both numbers by the smaller one:
1.26 𝑔 𝑂 2.52 𝑔 𝑂
∶
1.26 𝑔 𝑂 1.26 𝑔 𝑂
1∶2
The ratio is a simple whole-number ratio.
 All matter is made of tiny invisible indivisible particles
called ATOMS. (definition of element – Boyle)
 All atoms of the same element are identical and different
from atoms of every other element. (definition of element –
Boyle; idea of atomic weight – Dalton)
 Atoms combine in simple whole-number ratios to form
compounds. (law of constant composition – Proust)
 If atoms of the same element can combine in more than
one ratio, they form more than one compound. (law of
multiple proportions – Dalton)
 During chemical reactions, atoms break attachments,
rearrange and form new attachments to different atoms but
they are not created, destroyed or transformed into atoms
of a different element. (law of conservation of mass –
Lavoisier)
CHEM
Unit 4
Dalton’s model of
the atom
October 8, 2013
You-tube videos:
Brian Cox,
J. J. Thomson
Notes
History of atomic theory
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Billiard ball model
Atoms are solid spheres
Atoms of different elements have different masses.
Atoms have little hooks on them that let them attach to other
atoms.
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Cathode ray tube
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Beam of particles he shot went between two charged plates.
The beam bent toward positive plate and away from negative
plate.
Opposite charges attract so particles in beam are negative
Atoms are neutral
Neutral atom gives off (contains) negative particles – very, very
small
Called them electrons
Atom left behind became positive after electrons left
called plum pudding model (like blueberry muffin)
Blueberries are electrons – negative
Bread is rest of atom - positive
measured exact mass and charge of an electron
gold foil experiment
Shot alpha particles
(small, positive
particles) at gold foil
Most went straight
through
A few bent sideways
a few bounced
backward
WHY?
Straight through 
atom is mostly empty
Bent sideways  pushed away by something positive (like
charges repel)
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Thomson’s model
Robert Millikan
Ernest Rutherford
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CHEM
Unit 4
Notes
History of atomic theory
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Rutherford’s
model – nuclear
model
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Henry Moseley
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Protons
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James Chadwick
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Bounced back – hit something with a lot of mass and positively
charged and bounced off
Atom is mostly empty space
Center is small, positive and massive – very dense – called
nucleus
Atoms are outside in large, mostly empty region called electron
cloud.
Discovered the atomic number – amount of positive charge in the
nucleus.
Predicted the atomic number matched a number of positive
particles but died prior to proving it experimentally (it was later
proved by Rutherford)
Wrote periodic law – the properties of elements vary in a
predictable repeating way based on atomic number
Idea proposed by Moseley
Experiments proving done by Rutherford due to Moseley’s death
in WWI
Discovered the neutron
(how is really complicated, so don’t worry about it)