Download 6.022 X 10 23 atoms - Fort Thomas Independent Schools

Chapter 3
Atoms: The Building Blocks of
Matter
Chapter 3
The Atom: From Philosophical Idea To
Scientific Theory
The Atom: From Philosophical Idea To
Scientific Theory
 Remember the idea that matter is made of
atoms dates back to 400 BC.
 Not proven experimentally until 1700’s
The Greeks
History of the Atom
 In 400 B.C the Greeks tried
to understand matter
(chemicals) and broke them
down into earth, wind,
fire, and air.
~
~
Greek Model
Democritus
 Greek philosopher
 Idea of ‘democracy’
 Idea of ‘atomos’
 Atomos = ‘indivisible’
 ‘Atom’ is derived
 No experiments to support idea
Democritus’s model of atom
No protons, electrons, or neutrons
Solid and INDESTRUCTABLE
Alchemy
 After that chemistry was
ruled by alchemy.
 They believed that that
could take any cheap
metals and turn them
into gold.
 Alchemists were almost
like magicians.
 elixirs, physical
immortality
Alchemy
Alchemical symbols for substances…
..
.
......
.
.....
GOLD
SILVER
COPPER
IRON
SAND
transmutation: changing one substance into another
D
In ordinary chemistry, we cannot transmute elements.
Contributions
of alchemists:
Information about elements
- the elements mercury, sulfur, and antimony were discovered
- properties of some elements
Develop lab apparatus / procedures / experimental techniques
- alchemists learned how to prepare acids.
- developed several alloys
- new glassware
Timeline
Greeks
(Democratus ~450 BC)
Discontinuous
theory of matter
Issac Newton
(1642 - 1727)
ALCHEMY
400 BC
300 AD
1000
2000
American
Independence
(1776)
Foundations of Atomic Theory
Observations and chemical reactions led to
the following scientific laws that describe
how compounds are formed.
Foundations of Atomic Theory
Law of Conservation of Mass (Lavoisier)
Mass is neither destroyed nor created during ordinary chemical
reactions or physical changes.
Law of Definite Proportions (Proust)
The fact that a chemical compound contains the same elements
in exactly the same proportions by mass regardless of the size
of the sample or source of the compound.
Law of Multiple Proportions (Dalton)
If two or more different compounds are composed of the
same two elements, then the ratio of the masses of the
second element combined with a fixed mass of the first
element is always a ratio of small whole numbers.
Conservation of Mass
2 H2 + O2
2 H2O
H
H
H2
O
H
O2
+
H2
H
O
H2O
O
H 2O
H
H
O
H
H
4 atoms hydrogen
2 atoms oxygen
4 atoms hydrogen
2 atoms oxygen
Legos are Similar to Atoms
H
H2
H
H
O
+
H2
H
H
O2
H
O
H 2O
H
O
O
H
H 2O
Legos can be taken apart and built into many different things.
Atoms can be rearranged into different substances.
Conservation of Mass
High
voltage
electrodes
Before reaction
After reaction
glass
chamber
O2
High
voltage
H2O
O2
5.0 g H2
H2
80
0 g H2
g O2
300 g (mass
of chamber)
+
385 g total
45
? g H2O
40 g O2
300 g (mass
of chamber)
+
385 g total
Dorin, Demmin, Gabel, Chemistry The Study of Matter , 3rd Edition, 1990, page 204
Law of Definite Proportions
Joseph Louis Proust (1754 – 1826)
 Each compound has a specific ratio of elements
 It is a ratio by mass
 Water is always 16 grams of oxygen for every 2
grams of hydrogen
Law of Definite Proportions
Whether synthesized in the laboratory or obtained from
various natural sources, copper carbonate always has
the same composition.
Analysis of this compound led Proust to formulate
the law of definite proportions.
+
103 g of
copper carbonate
53 g of
copper
+
40 g of oxygen
10 g of carbon
Law of Multiple Proportions
John Dalton (1766 – 1844)
If two elements form more than one compound,
the ratio of the second element that combines
with the first element in each is a simple whole
number.
H2O
H2O2
water
hydrogen peroxide
Ratio of oxygen is 1:2 (an exact ratio)
Dalton’s Atomic Theory
 English chemist in the early
1800’s
 Dalton stated that elements
consisted of tiny particles
called atoms
 He also called the elements
pure substances because all
atoms of an element were
identical and that in
particular they had the
same mass.
Dalton’s Atomic Theory
1. All matter consists of extremely small particles that are indivisible
and indestructible called atoms.
2. Atoms of a given element are identical in their physical and
chemical properties. Atoms of different elements have different
physical and chemical properties.
3. Atoms of different elements combine in simple, whole number
ratios to form chemical compounds.
4.
Atoms cannot be subdivided, created or destroyed.
5. In chemical reactions, atoms are combined, separated, or
arranged.
Although some exceptions have been discovered, the theory still stands
today, and has been expanded and modified.
Daltons’ Models of Atoms
Carbon dioxide, CO2
Water, H2O
Methane, CH4
Structure of Atoms
 Scientist began to wonder what an atom was like.
 Was it solid throughout with no internal structure or was it
made up of smaller, subatomic particles?
 It was not until the late 1800’s that evidence became available
that atoms were composed of smaller parts.
Chapter 3
The Structure of the Atom
Structure of Atoms
 The cathode ray tube led to the discovery of electrons,
small, negatively charged particles at the turn of the
century.
 J. J. Thomson - English physicist
 Made a piece of equipment called a cathode ray tube.
 It is a vacuum tube - all the air has been pumped out.
A Cathode Ray Tube
Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 58
Thomson’s Experiment
-
voltage
source
+
vacuum tube
metal disks
Thomson’s Experiment
ON
-
OFF
voltage
source
+
Passing an electric current makes a beam appear
to move from the negative to the positive end
Thomson’s Experiment
ON
-
OFF
voltage
source
+
Thomson’s Experiment
ON
-
OFF
voltage
source
+
+
By adding a magnetic field…
he found that the moving pieces were negative.
J.J. Thomson
 He proved that ALL atoms
of any element must
contain these negative
particles.
 He knew that atoms did not
have a net negative charge
and so there must be
balancing the negative
charge.
J.J. Thomson
Plum Pudding Model
 In 1910 proposed the
Plum Pudding model
 Negative electrons
were embedded into a
positively charged
spherical cloud.
Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 56
Spherical cloud of
Positive charge
Electrons
Television Picture Tube
Blue beam
Green beam
Red beam
Glass window
Shadow mask
Fluorescent screen
Electron
gun
Electron
beam
Deflecting
electromagnets
Red beam
Green beam
Blue beam
Shadow mask
Fluorescent
screen with
phosphor dots
Ernest Rutherford (1871-1937)
 Learned physics in J.J.
Thomson’ lab.
 Noticed that ‘alpha’
particles were
sometime deflected by
something in the air.
 Gold-foil experiment
Rutherford ‘Scattering’
 In 1909 Rutherford undertook a series of experiments
 He fired a (alpha) particles at a very thin sample of gold foil
 According to the Thomson model the a particles would only
be slightly deflected
 Rutherford discovered that they were deflected through large
angles and could even be reflected straight back to the source
Lead collimator
Gold foil
a particle
source
q
Rutherford’s Apparatus
beam of alpha particles
radioactive
substance
circular ZnS - coated
fluorescent screen
gold foil
Dorin, Demmin, Gabel, Chemistry The Study of Matter , 3rd Edition, 1990, page 120
What He Expected
 The alpha particles to pass through without
changing direction (very much).
 Because…
 The positive charges were spread out evenly.
Alone they were not enough to stop the alpha
particles
California WEB
What he expected…
California WEB
What he got…
Density and the Atom
 Since most of the particles went through, the atom
was mostly empty.
 Because the alpha rays were deflected so much, the
positive pieces it was striking were heavy.
 Small volume and big mass = big density
 This small dense positive area is the nucleus
California WEB
Rutherford’s Experiment
Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 56
The Rutherford Atom
n+
Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 57
Size of an atom
 Atoms are incredibly tiny.
 Measured in picometers (10-12 meters)
 Hydrogen atom, 32 pm radius
 Nucleus tiny compared to atom
 Radius of the nucleus near 10-15 m.
 Density near 1014 g/cm
 If the atom was the size of a stadium, the nucleus would be
the size of a marble.
California WEB
Atoms
All atoms have similar structure
 protons and neutrons cluster together to form a
nucleus, or central core
 electrons orbit the space surrounding the
nucleus
 then things change depending on the element
Chapter 3
Counting Atoms
Atoms
Each element has a characteristic number of
protons
 Si- silicon 14 protons
 H- hydrogen 1 ptoton
 Ag- silver 47 protons
↓
Atomic number
Atomic Particles
 Protons are constant for an element, but electrons and
neutrons can vary
 When electrons vary, the charge of the atom changes
 When neutrons vary, you have a different isotope of
the atom
 Isotope- one of 2 or more atoms having the same number of
protons but different numbers of neutrons
Isotopes
Hydrogen
3 isotopes
protium
deuterium
tritium
1 proton
1 proton
1 proton
0 neutrons
1 neutron
2 neutrons
 isotopes have very similar chemical properties
Deuterium or
hydrogen-2
Isotopes
Contain the symbol of the element, the mass number and
the atomic number
# protons
+ # neutrons
mass number
# protons
Mass
number
Atomic
number
X
Isotopes
Atomic Number = number of protons
# of protons determines kind of atom/element
Atomic Number = number of electrons in a neutral atom
Mass Number = the number of protons + neutrons
California WEB
Mass Number
 mass # = protons + neutrons
• always a whole number
Neutron
+
• NOT on the
Periodic Table!
Electrons
Nucleus
+
+
+
+
+
Nucleus
Carbon-12
Neutrons 6
Protons
6
Electrons 6
Proton
Isotopes
 Find the
 number of protons
 number of neutrons
 number of electrons
 Atomic number
 Mass number
=9
= 10
=9
=9
= 19
+
19
9
F
Isotopes
Find the
– number of protons = 35
– number of neutrons = 45
– number of electrons = 35
– Atomic number = 35
– Mass number = 80
80
35
Br
Isotopes
Find the
 number of protons
 number of neutrons
 number of electrons
 Atomic number
 Mass number
23
11
Na
Sodium atom
Isotopes
Find the
 number of protons
 number of neutrons
 number of electrons
= 11
 Atomic number
= 12
 Mass number
= 10
= 11
= 23
23
11
1+
Na
Sodium ion
Isotopes
If an element has an atomic number of
23 and a mass number of 51 what is
the
– number of protons = 23
– number of neutrons = 28
– number of electrons = 21
– Complete symbol
51
23
V
+2
Isotopes
If an element has 60 protons and 144
neutrons what is the
– Atomic number = 60
= 204
– Mass number
– number of electrons = 60
– Complete symbol
204
60
Nd
Can atoms be counted or measured?
 atomic mass- the mass of an atom in atomic mass
units
 atomic mass unit- 1/12 of the mass of the
carbon-12 isotope
 Atoms too small to measure in grams. Created
new unit. Carbon-12 atom was assigned a value
of 12 atomic mass units (amu).
56
•22.990 is the
average atomic mass
of all the isotopes
•a weighted average
57
16
Hg
Isotopes
200.59
The percent natural abundances
for mercury isotopes are:
Hg-196
Hg-198
Hg-199
Hg-200
Hg-201
Hg-202
Hg-204
0.146%
10.02%
16.84%
23.13%
13.22%
29.80%
6.85%
(0.00146)(196) + (0.1002)(198) + (0.1684)(199) + (0.2313)(200) + (0.1322)(201) + (0.2980)(202) + (0.0685)(204) = x
0.28616 + 19.8396 + 33.5116 + 46.2600 + 26.5722 + 60.1960 + 13.974 = x
x = 200.63956 amu
Example: Si
92.21% of Si atoms in nature have mass of 27.98 amu
4.70% of Si atoms in nature have mass of 28.98 amu
3.09% of Si atoms in nature have mass of 29.97 amu
What is the average of these masses, taking into
consideration how often they are found in nature
(abundance)?
(92.21)(27.98) + (4.70)(28.98) + (3.09)(29.97)
100
= 28.09 amu
59
Example: Cl
2 isotopes occur naturally, chlorine-35, which has a natural
abundance of 75% and a mass of 34.969 amu and chlorine-37
which occurs only 25% of the time and has a mass of 36.966
amu. What is the average atomic mass of chlorine?
chlorine-35 75% 34.969 amu
chlorine-37 25% 36.966 amu
(75)(34.969) + (25)(36.966) = 35.453 amu
100
60
Mole
 discussing tiny particles and tiny amounts
 cannot create and work with in lab
 chemists derived a new unit as a bridge
between the microscopic and macroscopic
worlds
 fundamental SI unit used to measure the
amount of a substance
61
 collection of 6.022137 X 1023 particles or usually 6.022 X
1023 (Avogadro’s number)
 1 mole of any substance contains 6.022 X 1023 particles
 1 mole of oxygen contains 6.022 X 1023 atoms
 1 mole of water contains 6.022 X 1023 molecules
Just like a dozen is always 12 pieces!!!
62
molar mass- the mass in grams of 1 mole of a given
substance
molar mass = average atomic mass in grams
C → 1 mol = 6.022 X 1023 C atoms = 12.011 g of C
Fe → 1 mol = 6.022 X 1023 Fe atoms = 55.85 g of Fe
Mo → 1 mol = 6.022 X 1023 Mo atoms = 95.94 g of Mo
H2O → 1 mol = 6.022 X 1023 H2O molecules =
2(1.007 g H atoms) + 15.999 g O atoms = 18.013 g of H2O
Ca(OH)2 → 1 mol = 6.022 X 1023 Ca(OH)2 molecules =
40.08 g + 2(1.007 g) + 2(15.999 g) = 74.092 g of Ca(OH)2
So…
63
1 mol = 6.022 X 1023 atoms for any
substance
 can be written as conversion factor
1 mol
or 6.022 X 1023
6.022 X 1023
1 mol
64
How many atoms in 2.5 mols of Si?
2.5 mol Si
1 mol = 6.022 X 1023 atoms
? = # of atoms
2.5 mol Si • 6.022 X 1023 atoms
1 mol
= 1.5 X 1024 atoms Si
65
How many mols is 9.7 X 1024 atoms of Cu?
9.7 X 1024 atoms
1 mol = 6.022 X 1023 atoms
9.7 X 1024 atoms •
= 16 mol Cu
66
1 mol
6.022 X 1023 atoms
Determine the mass in grams of 3.5 mols of
Cu.
3.5 mol Cu
molar mass of Cu = 63.55g, which means
1 mol = 63.55g of Cu
mass of Cu = ?
3.5 mol Cu • 63.55 g Cu = 222 g Cu
1 mol
67
What is the mass of a single atom of Si?
molar mass of Si = 28.09g
1 mol = 6.022 X 1023 atoms
mass of 1 atom = ?
28.09g Si •
1 mol
1 mol
6.022 X 1023 atoms
= 4.665 X 10-23 g/atom
68