Transcript Atoms – Building Blocks of Matter Notes

Atoms – Building Blocks of Matter
Notes - Chapter 3
FIRST CHEMICAL REACTION
I. The Atom: From Idea to Theory
A. 400 BC Democritus VS Aristotle
Democritus, an ancient Greek and student of
Aristotle, proposed the 1st atomic theory he said
that the world is composed of 2 things: void
(empty space) and matter. No one supported
him and he had NO experimental evidence to
support his idea.
Aristotle proposed that matter was composed of
one continually flowing substance called hyle.
This idea was widely supported and accepted
until the late 1700’s and he too had NO
experimental evidence to support his idea.
What Accounts for Matter’s Secrets?
The Greeks came up with at least 3 different theories
ARISTOTLE – 4 ELEMENTS: AIR, EARTH, FIRE, WATER
ALCHEMISTS ACCOMPLISHMENTS – NOT GOLD
B. Late 1700’s Isaac Newton and
Robert Boyle
It was not until the late 1700’s that anyone dared to
question Aristotle’s wisdom. They suggested that
Aristotle was incorrect but did not have their own theory
to submit.
At this time chemist did believe, based on experiments,
that there were different elements and that an element
was a substance that could not be broken down by
chemical means. Chemist knew that some substances
could transform into different or new substances, they
called this a chemical reaction.
C. 1790’s - Basic laws that were
established:
Chemist also new, via improved balances, that
when a chemical reaction occurred in a closed
space that the mass of the material before the
change equaled the mass of the material after
the change. Now known as the Law of
Conservation of Mass.
Another realization was that substances always
contained their elements in the same proportions
by mass. For example: for any sample of
sodium chloride, the mass of the sample is
always 39.34% Na and 60.66% Cl. Now known
as the Law of Definite Proportions.
It was also known that elements combined to form
more than one compound. Example: carbon
monoxide and carbon dioxide. This is the Law
of Multiple Proportions.
D. 1803 John Dalton
British chemist who was the first to have a theory
about matter being composed of atoms and how
atoms might look and behave. Dalton proposed
an explanation for the Law of Conservation of
Mass, Law of Definite Proportions, and Law of
Multiple Proportions. He reasoned that elements
were composed of atoms and that only whole
numbers of atoms can combine to form
compounds. He conceived on the atom as a
solid billiard ball. Here is a summary of his
theory:
JOHN DALTON (1766 - 1844) REVIVES ATOMIC THEORY OF MATTER
1. All matter is composed of atoms.
2. Atoms of the same elements are exactly the
same and atoms of different elements are
different.
3. Atoms cannot be created, destroyed, or
subdivided.
4. Atoms of different elements combine in whole
number ratios to form compounds.
5. In chemical reactions, atoms are combined,
separated, or rearranged.
Democritus’s idea, because Dalton was
able to relate atoms to the measurable
property of mass, turned into a scientific
theory!!
The only aspect of Daltons’ Theory that is now
known to be incorrect is the fact that atoms can
be subdivided (into p+, e-, n). And that atoms of
the same element can have deterrent masses
(these are called isotopes).
II. The Structure of Atoms
Atom – smallest particle of an element that retains
the chemical properties of that element. All
atoms consist of 2 regions – the nucleus (p+ & n)
and surrounding the nucleus is the electron
cloud – a region occupied by the negatively
charged particles called electrons. How do we
know this?!
1. Discovery of Electron 1897 (by
J.J. Thomson and Robert Millikan)
1st subatomic particle to be discovered – Thompson
was working with electricity and magnetic fields. He
was taking various gases and sending an electric
current through the gas. When he did this he
noticed that a glow was emitted. (What he was
doing, he believed, was separating the electron
from the nucleus of the gas atoms – this caused the
glow!) Thompson went on to prove that the glow
was actually a stream of negatively charged
particles – called electrons. Symbol e-, charge –1,
and mass of 0.00055amu (atomic mass unit, 1amu
= 1.66 X10 –27 kg)
Plum Pudding Model – Thompson proposed
that the atom had negative electrons scattered
throughout a positively charged area (proton
area).
2. Protons 1919 (discovered by
Rutherford/J.J. Thompson)
Both Rutherford and Thompson knew that
positively charged particles (protons) must exist
(because an atom is neutral, if there is a
negative charged electron then there has to be a
positively charged proton to make it neutral.)
They worked together to prove they existed.
Proton symbol: +p, charge +1, mass 1.008 amu.
3. Discovery of the Atomic Nucleus
1911
Subatomic ParticleMassElectron 0.00055amuProton1.008
amu.Neutron1.008 amu.
Discovered by Rutherford during his famous gold-foil
experiment and realized that the main part of the atom’s
mass is in the nucleus, and that it is positively charged.
Summary of his experiment:
-Bombarded a thin piece of gold foil with positive alpha
particles
-Most went through as though nothing was there
-Few (1 in 8000) ricochet back toward the source
-Few were deflected off to the side
Rutherford’s Conclusion: the positive alpha particles
had to have hit something else that was positively
charged to cause the ricochet effect. The
“something” was very small and dense because
only a few hit it, therefore the atom must have a
small positively charged nucleus, surrounded by
mostly empty space (because most particles went
through the gold foil.) New model of atom:
Electron 0.00055amu
Proton 1.008 amu.
Neutron 1.008 amu.
4. Neutrons 1932 (Proved by
Chadwick)
New something else existed in an atom because of
the mass of the atom. Neutron is an electrically
neutral particle, symbol n, mass equal that of
protons.
5. Nuclear Forces – the +p and n stay close to
each other due to these short-range forces that
hold the +p and n together.
Current Model of Atom:
III. Counting Atoms
Reading the periodic table
11
atomic number
Na
symbol
22.990
average atomic mass (in amu’s)
Sodium
name of element
23
mass number (the average atomic mass
rounded to the nearest whole number)
1. Atomic Number
the number of protons in the nucleus. The
atomic number identifies the element!!!!!!!!!
Because atoms are neutral they contain the
same number of electrons as protons.
(Therefore the atomic number is the number of
electrons as well.)
2. Atomic Mass
– mass of 1 atom (measured in amu’s)
3. Mass Number
– the average atomic mass rounded to the
nearest whole number, therefore it is the total
number of protons and neutrons in an atom’s
nucleus.
Practice: How many protons are in each
of the following? neutrons? electrons?
Symbol At #
Mass # p+
Be
4
9
Ne
10
20
Na
11
23
n
e-
Symbol
4. Isotopes
– atoms with the same number of protons (atomic
number is the same) but different numbers of
neutrons (mass number is different). Usually
isotopes are referred to by their name (of symbol)
and their mass number. Every element on the
chart has at least 2 isotopes and some elements
have as many as 25 isotopes.
Example: The isotopes of hydrogen have
separate names rather than being called
hydrogen-1, hydrogen-2, etc. Their names are
protium (H-1), deuterium (H-2), and tritium (H-3).
Diagram of protium, deuterium, and tritium:
Name
Protium (H-1)
Deuterium (H-2)
Tritium (H-3)
p+ e- n
Mass #
Symbol
Practice: carbon-14, carbon-13, carbon12
Name
p+ e- n
Mass #
Symbol
Most elements occur naturally as mixtures of
isotopes, as indicated in Table 3-4 of textbook.
The percentage of each isotope in the naturally
occurring element on Earth is nearly always the
same, no matter where the element is found.
The percentage at which each of an element’s
isotopes occurs in nature is taken into account
when calculating the element’s average atomic
mass (which appears on the periodic table).
Nuclide
-a general term for any isotope of any element
5. Relative atomic masses
a.m.u.- atomic mass unit; One amu is exactly
1/12 the mass of a carbon-12 atom. So the
atomic mass of any nuclide is determined by
comparing it with the mass of the carbon-12
atom. The hydrogen-1 atom has an atomic mass
of about 1/12 that of the carbon-12 atom, or 1
amu. 1 amu = 1.66X10-27kg
6. Average atomic mass
It is the weighted average of the masses of all the isotopes
of that element. A weighted average reflects both the
mass and the abundance of the isotopes as they occur in
nature.
Ex:ample:
isotope
Atomic mass abundance (%)
protium
1
99.985
deuterium
2
0.015
tritium
3
negligible
The average atomic mass of hydrogen is 1.0079. Multiply
each mass number by the percent abundance and add
them up.
Practice: Element Z has 2 natural isotopes. The
isotope with a mass number of 15 has a relative
abundance of 30%. The isotope with a mass
number of 16 has a relative abundance of 70%.
Estimate the average atomic mass for this
element.
IV. Relating Mass to Numbers
of Atoms
1. The Mole (can be abbreviated mol, but NOT
m, which is the abbreviation for meter!) - the SI
unit for amount of substance. A mole is the
amount of a substance that contains as many
particles as there are atoms in exactly 12 grams
of carbon-12.
2. Avogadro’s Number
-the number of particles in exactly one mole of a
pure substance. This number was determined
experimentally and its value is 6.022 X 1023,
which means that 12 g of carbon-12 contains
6.022 x 1023 carbon-12 atoms.
3. Using the Mole and Avogadro’s
Number
A mole can be thought of as a counting unit
just like a dozen (12), gross (144), pair (2),
ream (500), mole (6.022X1023 ).
A. How many is a mole?
If every person living on Earth (6 billion people)
worked to count out one mole of oranges (or anything
else), and if each person counted continually at a rate
of one orange per second, it would take about 4
million years for all the oranges to be counted!
If we had a mole of sand it would cover the earth 7
times over! If you had a mole of dollar bills, you
could spend a million dollars every minute of your life
and never spend it all!
Since the mole is so large, we use it to count
very tiny things – like atoms. Because the mole
is so large, (and we now know that we cannot
count out a mole of anything), how do we know
when we have a mole of anything?
We determine the mass and relate that to the
number of atoms present.
4. Molar Mass
– The mass of one mole of a pure substance.
The pure substance can be an element or a
compound.
The atomic mass is the mass of 1 atom of that
element measured in amu’s.
The atomic mass is also equal to 1 mole of
atoms measured in grams it is called the
molar mass!!!! What a coincidence!!!!
Mass of 1 atom of Pb = 207.2 amu
Mass of 1 mole of Pb atoms = 207.2 g
Mass of 1 atom of N = 14.01 amu
Mass of 1 mole of N atoms = 14.01 g
Mass of 1 atom of Ba = 137.33 amu
Mass of 1 mole of Ba atoms = 137.33 g
Mass of 1 atom of Al = 26.98 amu
Mass of 1 mole of Al atoms = 26.98 g
Let’s prove it: Determine the mass,
in grams, of 6.022X1023 atoms of
aluminum. Use 1amu = 1.66X1027kg.
With this information we can write some
new conversion ratio’s!!
1 mole = 6.022X1023 atoms OR molecules OR
formula units
1 mole Al = 26.98 grams
1 atom Al = 26.98 amu
V. Mole Problems – When in
doubt go to the mole!
The MOLE has been defined as 6.022 x l023 atoms of a
pure element or the molar mass of a substance
expressed in grams. It can also be defined as 6.022 x
l023 molecules of a compound or diatomic molecule
(O2, N2, H2, etc)
THE ONLY THING HARD ABOUT UNDERSTANDING
THE DEFINITION OF A MOLE IS THAT YOU
UNDERSTAND THAT THE VALUE OF A MOLE IS
DIFFERENT FOR EVERY DIFFERENT ELEMENT
AND COMPOUND.
1. Gram/Mole conversions-how to convert moles to
grams or grams to moles.
Example: 120 g Ca x 1 mole Ca = 2.99 mole Ca
40.1 g Ca
Practice:
How many grams of sodium are in 5.00 moles of
sodium?
How many grams of magnesium are in 0.250 moles of
magnesium?
How many moles of lead, Pb, are in 210. g of lead?
How many moles of nitrogen are in 44.0 g of nitrogen?
2. Conversions with Avogadro’s Number
Example: How many atoms of silver, Ag, are in
4.25 moles of Ag?
4.25 moles Ag X 6.022 x 1023 atoms Ag =
1 mole Ag
Practice:
How many atoms of Pb are in 3.80 moles of Pb?
How many moles of Na are in 8.24 x 1024 atoms
of Na?
Two-step conversions:
Ex: How many atoms of sodium, Na, are in 5.25 g of
Na?
5.25 g Na
x 1 mole Na
23 g Na
x 6.022 x 1023 atoms Na =
1 mole Na
Practice:
How many atoms of potassium, K, are in 3.99 g of
K?
How many g of He are in 3.03 x 1021 atoms of He?
3. How many atoms of Li are in 0.755 g of Li?
3. Molar Mass for compounds
How to find the molar mass:
Write a CORRECT formula for the compound
(we’ll do this later)
Look up the atomic mass of each element in the
compound
Multiply the atomic mass by the subscripts, if any.
Add all masses of elements together and use the
unit, g/mol
Example: find the molar mass of NaCl.
Na=23.0 g/mol
Cl=35.5 g/mol
58.5 g/mol
Example: find the molar mass of calcium phosphate,
Ca3(PO4)2.
Ca = 40.1 x 3 = 120.3
P = 31.0 x 2 = 62.0
0 = 16.00 x 8 = 128.00
310.3 g/mol
Practice:
Find the molar mass of ammonium sulfate,
(NH4)2SO4
Find the molar mass of Cl2O7
Find the molar mass of NaCl
Hydrates - Some compounds trap water
inside their crystal structure and are known
as hydrates. You will not be able to predict
which compounds will form hydrates. All
you have to do is to be able to name them
and find their molar masses (including the
water).
CuSO4 · 5H20 is an example of a hydrate.
This says that one formula unit of cupric
sulfate will trap 5 molecules of water inside
its crystal.
Hydrates are named by naming the ionic compound
by the regular rules and then adding (as a second
word) a prefix indicating the number of water
molecules. You will use the word “hydrate” to
indicate water. The above compound would be
called cupric sulfate pentahydrate.
To find the formula mass of a hydrate, simply find
the mass of the ionic compound by itself and then
ADD the mass of water molecule(s) to that mass.
Practice: What is the formula mass of barium chloride
dehydrate, BaCl2 · 2H2O?
What is the formula mass of aluminum sulfate
octahydrate, Al2(SO4)3 · 8H2O?