Download Unit 2 Atomic Structure 2015

Unit B: Atomic Structure &
Nuclear Chemistry
Atoms: Building Blocks of Elements
• An atom is the
smallest
particle of an
element that
has the
characteristics
of that element.
The Atom
Proton
Electron
Neutron
Subatomic Particles
• Proton+ has a positive charge and is located
in the nucleus of an atom
• Neutron0 has no charge and is also located in
the nucleus
• Electron- has a negative charge and travels at
the speed of light around the outside of the
atom
The Nucleus
• Is positively charged
because it contains
protons (p+) and
neutrons (n0)
• Is much smaller than
the whole atom yet
contains most of its
mass
This is what the atomic model would
look like drawn to scale
History of the Atom
• In ~400 BC,
Democritus
proposed that
everything in the
universe was
composed of small,
indivisible objects
called “atomos”
John Dalton
• In the 1800’s, John
Dalton performed a
series of experiments
combining elements
that showed that
matter actually did
consist of small
particles, but he
couldn’t predict their
structure
• He used this to come up
with his Atomic Theory
Dalton’s Atomic Theory
1) chemical elements are made of atoms. 2)
the atoms of an element are identical in their
masses 3) atoms of different elements have
different masses 4) atoms only combine in
small, whole number ratios (there are no ½
atoms) 5) atoms can be neither created nor
destroyed
Evaluate
• Which aspects of Dalton’s Atomic Theory were
correct? Which were not?
J.J. Thomson
• In 1897, another English
scientist, J.J. Thomson,
discovered the electron
• Based upon his
experiments, he
proposed the “Plum
Pudding” model of the
atom, with negative
electrons floating in a
sea of positive matter
Ernest Rutherford
• In 1911, Ernest Rutherford
performed experiments that led
him to discover that electrons
orbit a positively charged nucleus
• He later discovered it contained
protons and neutrons
His hypothesis was incorrect…and he
learned something!
Rutherford “planetary” model
Niels Bohr
• Danish scientist Niels
Bohr came up with the
final model of the atom
• He discovered that
electrons can only orbit
the nucleus at certain
specific distances, which
he called “orbitals” or
“energy levels”
Energy Levels
• Bohr discovered that different orbitals
required different amounts of energy for an
electron to reach them
• Electrons could not exist “between” levels, but
they could leap from one level to the next if
charged with enough energy
• Each level could only hold a certain number of
electrons (2, 8, 18, etc)
The Bohr Model
Present – Quantum Mechanical Model
(Heisenberg/Schrodinger)
• Bohr and other scientists
(Heisenberg) updated the
model in light of quantum
mechanics
• Rather than treating
electrons as having
specific locations and
circular orbits, they are
represented by
probability clouds where
electrons are most likely
to be found
Not THAT Heisenberg…
This one.
History of Atomic Theory Summary
Scientist/Philosopher
Discovery/Contribution
Democritus
Atoms are “uncuttable”
John Dalton
Atomic Theory; elements are
composed of one type of atom
JJ Thomson
Electron
Ernest Rutherford
Nucleus
Neils Bohr
Energy Levels
Werner Heisenberg/Schrodinger Electron Clouds
This is what they look like
Elements
• Atoms of different
elements differ in
their number of
protons
• All known 118
elements in the
universe are found
on the Periodic
Table of Elements
Chemical Symbols
• Each element has
its own unique
symbol which
always begins
with a CAPITAL
letter
Atomic Number
• Is the number of
protons in an
atom
• In a neutral atom,
it is also the
number of
electrons
Atomic Mass
• Is the average weight
of an atom based
upon the number of
protons and
neutrons in the
nucleus
• Measured in atomic
mass units (amu)
• Tells you the mass
number of MOST
atoms of an element
Mass Number
• Is the total number
of protons and
neutrons in an atom.
• To find the typical
mass number of an
element, round its
atomic mass to the
nearest whole
number
• The mass number for
Hydrogen is 1 amu
Simplified for notes
• Atomic # = p+
• e- = p+(if neutral)
• Mass # = p+ + no
• # of neutrons =
mass # - atomic #
Examples
•
•
•
•
•
•
Element name: Carbon
Atomic Number: 6
Mass Number: 12
# Protons: 6
# Electrons: 6
# Neutrons: 12-6 = 6
Another...
• Element name:
Gold
• Atomic Number: 79
• Mass Number: 197
• # Protons: 79
• # Electrons: 79
• # Neutrons: 197-79
= 118
Isotopes
• -are atoms of an element that have the same number
of protons but different numbers of neutrons
• E.g.:
• -Carbon-12 has 6 protons, 6 neutrons
• -Carbon 13 has 6 protons, 7 neutrons
Representing Isotopes
This is the isotope needed for nuclear
fission
NUCLEAR CHEMISTRY
Marie Curie
• In 1896, Marie Curie
demonstrated that the
fogging of film plates
exposed to uranium
was caused by rays
emitted by the atoms
• She called this
radioactivity and
became the first woman
to receive a Nobel Prize
in 1903
Radioactive Isotopes
• Are atoms with an
imbalance of protons to
neutrons and an
unstable nucleus
• To become stable, they
slowly give off excess
particles in the form of
radiation
• Any atom that is giving
off radiation is said to
be radioactive
Radiation
• The penetrating rays
and particles emitted by
a radioactive source is
known as radiation
• Radiation that is
harmful to the body is
known as ionizing
radiation
• It can take 3 forms:
alpha, beta, and gamma
Alpha Radiation (α)
• Consists of Helium
nuclei that have been
emitted from a
radioactive source
• Because it is just the
nucleus (no electrons),
these particles have a
positive charge
Beta Radiation (β)
• Consists of fast moving
electrons formed by the
decomposition of a
neutron
• The neutron breaks into
a proton, which stays in
the nucleus, and an
electron which is
emitted
Gamma Radiation
• Is high-energy
electromagnetic
radiation (light) given
off by a radioisotope
alongside alpha or beta
particles
• Most damaging type of
radiation
Gamma ray bursts like this one could
wipe out all life on Earth at any moment.
Property
Alpha Radiation Beta Radiation
Gamma
Radiation
Particle
Helium Nucleus
Electron
High-energy
electromagnetic
radiation
Symbol
α, 42He
β, 0-1e
γ
Charge
Positive (+2)
Negative (-1)
0
Common Source Radium-226
Carbon-14
Cobalt-60
Penetrating
Power
Low (0.05 mm
body tissue)
Moderate (4 mm Very High
body tissue)
(penetrates
body easily)
Shielding
Paper, clothing
Metal foil
Lead, concrete
Radiation Penetration
Half-life
• Radioactive isotopes
decay at very specific,
unique rates
• The time required for
one-half of the nuclei of
a sample to decay is
called its half-life
Half Life Problems
• Each time a sample goes through 1 half life,
half of the sample decays and half remains.
• Ex: Let’s say a sample has a half life of 5 years
and we start with 8g…
8g remain
Year 0 (0 half
lives)
4g remain
2g remain
1g remains
Year 5 (1 half
life)
Year 10 (2
half lives)
Year 15 (3
half lives)
Note: will the sample ever reach 0?
No!
Graphing Half Lives
E=mc2
• In the early 1900’s,
Albert Einstein showed
that energy and matter
were equivalent with
his famous equation,
E=mc2
• This lead later scientists
to hypothesize that
tremendous amounts of
energy could be
produced by tiny
amounts of matter
Nuclear Fission
• The splitting of an
atomic nucleus is
known as nuclear
fission
• The split atoms
then release more
neutrons, leading
to a chain reaction
Nuclear Chain Reactions
Controlled (nuclear power plant)
Uncontrolled (atomic bomb)
Nuclear Fusion
• Occurs when nuclei
combine to produce a
bigger nucleus
• For example, 2
Hydrogen nuclei can
combine to form
Helium
• The sun is powered by
nuclear fusion
Tsar Bomba (Hydrogen bomb) – the
biggest bomb ever
A nuclear fission reaction sets off a second nuclear reaction, this one an
uncontrolled fusion reaction, resulting in incredible amounts of energy
COUNTING ATOMS
Avg Atomic Mass
• Is the weighted average
of all of the isotopes of
an element
• To find the mass #,
multiply the %
abundance of each
isotope by the isotope’s
mass and then add the
products together
Avg Atomic Mass of Carbon
• 0.989 x 12 = 11.868
• 0.011 x 13 = 0.143
• 0.000001 x 14 =
0.000014
Atomic Mass =
11.868
+ 0.143
+ 0.000014
= 12.01 amu
Avogadro's Number (the Mole)
• A unit used for counting atoms
• Similar to a dozen, except instead of 12,
it’s 602 billion trillion
602,000,000,000,000,000,000,000
• 6.02 X 1023 (in scientific notation)
• This number is named in honor of Amedeo
_________ (1776 – 1856), who studied
quantities of gases and discovered that no
matter what the gas was, there were the
same number of molecules present in a
given volume (Avogadro’s Number)
The mole defined…
• The mole is defined as the number of carbon12 atoms in exactly 12g of carbon
• This number turns out to be Avogadro’s
Number, 6.02 x 1023 (scientists have actually
calculated this down to 8 significant figures,
but we will only hold you to the 3 )
Just How Big is a Mole?
• Enough soft drink cans to cover the
surface of the earth to a depth of
over 200 miles.
• If you had Avogadro's number of
unpopped popcorn kernels, and
spread them across the United
States of America, the country would
be covered in popcorn to a depth of
over 9 miles.
• If we were able to count atoms at the
rate of 10 million per second, it
would take about 2 billion years to
count the atoms in one mole.
• In other words: a LOT.
Everybody Has Avogadro’s Number!
But Where Did it Come From?
• It was NOT just picked!
It was MEASURED.
• One of the better
methods of measuring
this number was the
Millikan Oil Drop
Experiment
• Has been consistently
supported by repeated
experiments
Why do scientists use the mole?
• Used for counting
extremely large
numbers, such as
the number of
atoms in 18g of
water
• Similar to why
bakers use dozens
Counting units
Suppose we invented a new collection unit
called a rapp. One rapp contains 8 objects.
1. How many paper clips in 1 rapp?
a) 1
b) 4
c) 8
2. How many oranges in 2.0 rapp?
a) 4
b) 8
c) 16
3. How many rapps contain 40 gummy bears?
a) 5
b) 10
c) 20
The Mole
• 1 dozen cookies = 12 cookies
• 1 mole of cookies = 6.02 X 1023 cookies
• 1 dozen cars = 12 cars
• 1 mole of cars = 6.02 X 1023 cars
• 1 dozen Al atoms = 12 Al atoms
• 1 mole of Al atoms = 6.02 X 1023 atoms
Note that the NUMBER is always the same,
but the MASS is very different! (a mole of
cars will probably weigh more than a mole of
cookies!)
Mole is abbreviated mol (gee, that’s a lot
quicker to write, huh?)
A Mole of Particles
Contains 6.02 x 1023 particles
1 mole C
= 6.02 x 1023 C atoms
1 mole H2O
= 6.02 x 1023 H2O molecules
1 mole NaCl
= 6.02 x 1023 NaCl “molecules
(technically, ionic compounds are ratios not
molecules so they are called formula units)
6.02 x 1023 Na+ ions and
6.02 x 1023 Cl– ions
Avogadro’s Number as
Conversion Factor
6.02 x 1023 particles
1 mole
or
1 mole
6.02 x 1023 particles
Note that a particle could be anything since the mole is a
counting unit
Converting from moles to particles
• If I have 2 moles of pennies, how many
pennies do I have?
2 moles X 6.02 x 1023 pennies = 12.04 x 1023
of pennies mole pennies
pennies
= 1.20 x 1024
pennies
Particles to moles
• If I have 1.8 x 1024 pencils, how many
moles of pencils do I have?
1.8 x 1024 pencils x 1 mole of pencils
6.02 x 1023 pencils
= 3 moles
of pencils
Learning Check
1. Number of atoms in 0.500 mole of Al
a) 500 Al atoms
b) 6.02 x 1023 Al atoms
c) 3.01 x 1023 Al atoms
2.Number of moles of S in 1.8 x 1024 S atoms
a) 1.0 mole S atoms
b) 3.0 mole S atoms
c) 1.1 x 1048 mole S atoms
Part 2: Molar Mass
• Explain what molar mass is
• Distinguish between gram molecular mass and
gram formula mass
• Find the molar mass of elements and
compounds
Molar Mass
• The mass (in grams) of 6.02 x 1023 particles (1
mole) of a substance
• Relative in value to the avg atomic mass:
1 C atom
=
1 mole of C atoms
1 Mg atom
atoms
12 amu
= 12 g
= 24 amu 1 mole of Mg
= 24 g
Same number, different unit!
The Unit
• We express molar mass using the
unit g/mol(grams per mole)
• Similar to density (g/mL)
Atomic Mass and Molar Mass
Atomic Mass
• the mass of one ATOM
• Measured in amu (atomic
mass units)
Molar Mass
• The mass of one MOLE of
atoms (6.02 x 1023 atoms)
• Measured in grams (g) per
mole
The atomic mass and molar mass are the same
number…the only thing that is different is the unit!
Formula/Atomic/Molecular Mass
• Gram atomic mass (gam) is the mass of 1
mole of atoms
• Gram Molecular Mass (gmm) is the mass of
1 mole of molecules (covalent only!)
• Gram Formula Mass (gfm) is the mass of 1
mole of ionic compounds
These are all examples of the same thing:
molar mass!
Learning Check!
Find the molar mass
(usually we round to the hundredths place)
A. 1 mole of Br atoms = 79.90 g/mole
B. 1 mole of Sn atoms = 118.71 g/mole
Molar Mass of Molecules and
Compounds
Mass in grams of 1 mole equal numerically to
the sum of the atomic masses
1 mole of CaCl2 = 110.98 g/mol
1 mole Ca x 40.08 g/mol
+ 2 moles Cl x 35.45 g/mol
1 mole of N2O4
= 110.98 g/mol CaCl2
= 92.02 g/mol
Learning Check!
A. Molar Mass of K2O = ? Grams/mole
94. 20 g/mol
B. Molar Mass of antacid Al(OH)3 = ?
Grams/mole
78.01 g/mol
Challenge
Prozac, C17H18F3NO, is a widely used
antidepressant that inhibits the uptake of
serotonin by the brain. Find its molar
mass.
Pt 3: Calculations with Molar Mass
molar mass
Grams
Moles
• Moles to Grams/Grams to Moles
• Particles to Grams/Grams to Particles
• Volume
Converting Moles and Grams
Aluminum is often used for the structure
of light-weight bicycle frames. How
many grams of Al are in 3.00 moles of
Al?
3.00 moles Al
? g Al
1. Molar mass of Al
1 mole Al = 26.98 g Al
2. Conversion factors for Al
26.98g Al
1 mol Al
or
1 mol Al
26.98 g Al
3. Setup 3.00 moles Al x
26.98 g Al
1 mole Al
Answer
= 80.94 g Al
Example 2
After leaving his bike out all night, Mr.
Pinson collected a 36g sample of Iron
(II) Oxide from his wheel. How many
moles of rust were formed?
1. First, find the molar mass of FeO
Fe 1 55.85
O 1 16.00
= 71.85 g/mol
2. Write the known quantity
36g FeO x 1 mole FeO
71.85g FeO
3. Multiply by the
conversion factor.
= 0.50 Moles FeO