Download Atomic Theory Notes (Chap 3,18

Chapter 3 and 18:
Modern Atomic Theory
I. Atomic Structure
A. The atom is mostly empty
space with a solid core
composed of protons and
neutrons surrounded by a
series of “clouds” or “energy
levels” containing electrons.
1. Three fundamental
particles compose atoms:
protons, electrons, and
neutrons.
2. Atoms by definition are
electrically neutral (# of
protons which are positive
= # of electrons which are
negative)
3. Rutherford Gold Foil
experiment (early 1900’s)
Topic
2
Atomic
Atomic Structure:
Structure: Basic
Basic Concepts
Concepts
The Gold Foil Experiment
Topic
2
Atomic
Atomic Structure:
Structure: Basic
Basic Concepts
Concepts
The Nuclear Model of the Atom
• The new model of the atom as pictured by
Rutherford’s group in 1911 is shown below.
4. Cathode ray tube –
used to discover the
electron and proton
(Thomson)
B. Atomic number vs atomic
mass
1. Each element (atom)
has a unique number of
protons and electrons that
define it. This number is
called the atomic number.
2. Each element is also
characterized by an
average atomic mass that
neglects the number of
electrons because they are
so small in mass. This
number is called the
atomic mass.
3. Atomic mass (rounded
to the nearest whole
number
-atomic number
# of neutrons
4. Atoms are arranged in
the Periodic Table in
order of increasing atomic
number.
C. Nucleus = central core,
contains most of the mass of
the atom (protons and
neutrons) vs energy levels
(electrons)
1. protons = positive
charge (+1), mass of 1
amu (1.673 x 10-24 g),
nucleus
2. neutrons = no net
electric charge (0), mass
of 1 amu (1.675 x 10-24 g),
nucleus
3. electrons = negative
charge
(-1), mass of 0.0005 amu
(9.1 x
10-28 g), energy
levels
4. protons are made of 2
“up” and 1 “down” quark
5. neutrons are made of 2
“down” and 1 “up” quark
6. gluons are smaller
particles transferred between
quarks
7. pions are smaller
particles transferred
between protons and
neutrons
8. quarks, pions, gluons,
and other particles
(neutrinos, etc.) were
discovered by blowing
apart the nucleus of
atoms in particle
accelerators
9. forces inside the
nucleus: strong nuclear
force holds particles
together whereas
electrostatic repulsion
pushes them apart
10. Electrons fill up energy
levels, sublevels, and
orbitals in order
1st level = 2
2nd level = 8
3rd level = 18
4th level = 32
5th – 7th level = 32
D. Ions are atoms with a net
electrical charge
1. atoms become ions by
gaining or losing electrons
to satisfy the octet rule
(Rule of 8) to become
more “stable.” Atoms
with 8 electrons in their
outermost energy level
are stable.
2. Positive ions (cations)
form when atoms lose
electrons.
3. Negative ions (anions)
form when atoms gain
electrons.
4. Each ion is assigned a
charge, oxidation
number, or valence. (+1,
+2, +3, +4, -1, -2, -3, -4)
written as superscripts.
5. Determining the ionic
charge:
6. Writing ionic
equations:
7. Periodic Table and
Ionic Charge:
E. Isotopes: atoms with a
different number of neutrons
1. Writing isotopes using
the “mass number”:
2. Determining the most
common isotope:
F. “Particle” nature of
matter
1. Democritus (400 BC) –
Substances around us are
made of tiny “indivisible”
particles = “atomos”
2. Law of Definite
Proportions – two
samples of the same
compound have the same
proportions by mass
Example: NaCl
60.7% - Na; 39.3% - Cl
3. Law of Conservation of
Mass: mass of reactants
must equal the mass of
the products
Example: 2H2 + 1O2 
2H2O
4. Law of Multiple
Proportions: If two or
more different
compounds are composed
of the same two elements,
the ratio of the masses of
the second element to the
first is always a simple
whole number.
NO
NO2
1.14g of O
1.00g of N
2.28g of O
1.00g of N
5. Dalton’s “concept” of
matter and atoms
a. All matter is
composed of atoms
which can not be
subdivided, created, or
destroyed
b. Atoms of the same
element are identical
c. Atoms of different
elements are different
d. Atoms of different
elements combine in
simple whole number
ratios
e. In chemical
reactions atoms are
rearranged but neither
created nor destroyed
II. Electrons and Their Properties
A. Electrons travel around
the nucleus at great speeds in
energy levels, sublevels, and
orbitals. Atoms can contain
as many as 7 energy levels.
B. The greater the distance
from the nucleus, the higher
the energy level of the
electron.
C. Simplified “Bohr”
diagrams of atomic structure
D. Electrons have both the
properties of matter as well
as energy (waves) – De
Broglie
E. Electromagnetic Energy
and “Waves”
1. amplitude
2. wavelength
3. frequency
4. speed
5. Electromagnetic
spectrum
6. wavelength x
frequency = speed of light
F. Quantum Theory of the
Atom
1. atoms absorb or emit
specific amounts of
energy or quanta
(Planck); E = hv
2. photoelectric effect –
each photon of light
carries a specific amount
of energy given by the
equation E = hv (Einstein)
3. dual nature of light:
possesses the properties of
energy (waves) and
matter (particles)
4. when excited and
viewed through a
spectroscope, atoms
produce various colored
“line spectra” of specific
colors, wavelengths, and
energies. This reflects the
“jumps” electrons make
between energy levels and
the “ground” state and
therefore the specific
“energies” of electrons.
5. electrons
behave as matter “waves”
– DeBroglie
6. We can not know the
exact position nor
momentum of an electron
simultaneously –
Heisenberg uncertainty
principle
G. Quantum Mechanical
Model of the Atom –
Schrodinger and the “Wave”
Equation
1. Electrons in an atom
can be described by a
“wave” equation using 4
quantum numbers that
describe the properties of
electrons
2. 1st quantum number –
describes the average
distance from the nucleus
or principal energy level
(1-7) (n)
3. 2nd quantum number –
describes the sublevel
(s,p,d, and f) and shape of
the cloud (spherical,
barbel, cloverleaf, etc) (l)
4. 3rd quantum number –
describes the orientation
in space and orbital (m)
s
p
d
f
sublevels – 1 orbital
sublevels – 3 orbitals
sublevels – 5 orbitals
sublevels – 7 orbitals
Orbital = space occupied
by a pair of electrons
5. 4th quantum number –
describes the electron spin
rotation (s)
positive spin ↑ - clockwise
negative spin ↓ - counterclockwise
6. writing electron
configurations:
7. writing orbital filling
diagrams:
8. writing electron dot
diagrams:
9. guiding principles
a.
Aufbau principle
b. Pauli exclusion
principle
c.
Hund’s rule
III. Counting Atoms Using the
Mole
A. Mole = SI unit for the
quanitity of particles
(abbreviation for “molecule”)
1. One mole of any kind of
particles (atoms,
molecules, ions, or
formula units) contains
6.02 x 10 23 particles
(similar to a dozen: 1
dozen = 12 objects)
2. particles:
Fe = atoms
H2O = molecules
(covalent bonds)
NaCl = formula units
(ionic bonds)
Na+, Cl- = ions
3. One particle has a mass
in amu’s equivalent to the
mass of one mole of
particles in grams
1 Fe atom = 55.85 amu
(atomic mass)
1 mole of Fe atoms = 55.85g
(molar mass)
1 H2O molecule = 18.02amu
(molecular mass)
1 mole of H2O molecules =
18.02g (molar mass)
4. Calculating the
molecular/molar mass
a. molecular mass of
CH4 =
b. molar mass of Ag =
c. formula mass of
CuCl2 =
B. We can use the mole to
“count” particles indirectly
by massing them.
1. Why do we need to
“count” particles?
(because chemicals react
with each other in specific
proportions based upon
the number of particles or
moles)
2. 3CuCl2 + 2Al  3Cu +
2AlCl3
3. 2H2 + 1O2  2H2O
C. Mole calculations
1. Changing grams to
moles
2. Changing moles to
grams
3. Changing grams to
particles
4. Changing particles to
grams
5. Changing moles to
particles
6. Changing particles to
moles
7. Changing to volume of
a gas at STP (Standard
temperature and pressure
= 1 atm or 101.3kPa, 0
degrees Celsius or 273K)
a. 1 mol of any gas
occupies 22.4L at STP
(molar volume)
b. Sample calculations
8. Mole “Road Map”
IV. Nuclear stability and
radioactivity (Chapter 18)
A. Nuclear Stability
1. Not all combinations of
protons and neutrons
make a stable nucleus
(one that remains in tact).
a.
nucleons =
b. nuclides =
2. The strong nuclear
force holds the nucleus
together by overcoming
the “repulsion” force of
like charges.
3. Atoms that contain
more or less neutrons
than the number of
protons become less stable
and begin to emit
radiation.
a. small atoms with an
equal number of
protons and neutrons
are stable!!!
b. Larger atoms with
an unequal number of
protons and neutrons
are unstable
c. Uranium – unstable
and radioactive, last
“naturally occurring”
element
4. Forms of radioactivity:
a.
alpha particles
(stopped by a piece of
paper or your skin)
b.
beta particles
(stopped by several sheets of
aluminum foil)
c. gamma radiation
(stopped by a thick sheet of
lead!)
5. Writing nuclear decay
equations:
B. Nuclear reactions (Fission
vs Fusion) – Chapter 18
1. Fission = splitting
heavy atoms to produce
smaller ones
2. Fusion = combining
smaller atoms to make
heavier ones
3. E = mc2
a. nucleons  nucleus +
energy
b. when a nucleus is
formed from separate
nucleons, the mass of
the nucleus is less than
the total mass of the
nucleons that compose it
= mass defect, the
nucleus is at a lower
energy state than the
individual nucleons
composing it
c. nuclear binding energy
= energy released when
nucleons come together
d. in nuclear reactions,
mass may be converted
into energy
e. even a very small
amount of mass
converted into energy
releases huge amounts
of energy
f. 1 g of U-235 releases the
same amount of energy
in fission as burning 7
500 000g of coal. (16,500
lbs)