Download Structure of the Atom - Dr. Vernon-

General Chemistry
Unit 3 Note Packet – Atomic Structure And The Periodic Table
Atom – the smallest part of matter that retains the characteristics
of that type of matter
 considered to be the building blocks of matter
 atoms consist of a positively charged center, or nucleus,
surrounded by negatively charged particles
 nucleus consists of protons and neutrons
 negatively charged particles are called electrons existing outside
the nucleus in mathematical “regions of probability”
particle
mass
(g)
mass
(amu)
relative
charge
location in atom
proton
neutron
electron
Types of Atoms  elements
 each different type of atom is an element
 elements (different types of atoms) are distinguished by the #
of protons
 every atom of the same element has the same # of protons (ex:
every atom of hydrogen has one proton, every atom of boron has
5 protons, etc.)
 the # of protons = atomic number
 elements are organized on the periodic table according to atomic
number
 each element has a chemical symbol
 the symbol is the first letter of an element’s name capitalized and
may have another letter from the name in lower case (examples: C
for carbon and Al for aluminum)
 some symbols are derived from Latin names (example: Argentum
is Latin for silver; the symbol for silver is Ag)
 # protons = # electrons in a neutral atom
 # protons ≠ # electrons in a charged (positive or negative) atom
(called an ion)
 # neutrons can vary (not the same for every atom of an element)
Ions  the result of a neutral atom gaining or losing an electron
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cation  positive ion
anion  negative ion
charge of an ion = number of protons – number of electrons
Isotopes atoms that have the same number of protons but
different numbers of neutrons
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for example…..hydrogen has three isotopes:
hydrogen 1 the atomic mass is “1” which means there is only one
proton in the nucleus
hydrogen 2 (deuterium) the atomic mass is “2” which means
there is one proton and one neutron in the nucleus
hydrogen 3 (tritium) the atomic mass is “3” which means there is
one proton and two neutrons in the nucleus
mass number = # protons plus neutrons
atomic mass or atomic weight takes into account all the different
isotopes
atomic mass/weight is an AVERAGE of all the different isotopes
an example calculation for atomic mass is show below:
A sample of silver is 51.35%
atomic mass/weight?
107
Ag and 48.65%
108
Ag. What is its average
Answer: 0.5135 x 107 = 54.9445
0.4865 x 108 = 52.542
Total
= 107.4865 amu = average atomic mass/weight
(107.49 amu)
MODELS OF THE ATOM (1900 to present)
Rutherford proposed an atom with a nucleus and negative
particles outside of the nucleus.
The next major step forward on the model of the atom came
after some important breakthroughs in our understanding of
electromagnetic energy (light!).
 light behaves like a wave, but also like a stream of extremely
tiny, fast-moving particles (wave-particle duality)
 light is a form of electromagnetic radiation
 electromagnetic radiation can be described in terms of:
 amplitude
– height of wave from origin to crest
 wavelength ()
– distance between successive crests
(meters)
 frequency ()
– how fast the wave oscillates; the
number of cycles per second
(high frequency = high energy)
 speed of light (c) – 3.00 x 108 meters per second
c = 
Notice: if  is large, then  must be small since
c is constant; the reverse is also true
(therefore, long  light is low energy
while short  light is high energy)
 visible spectrum of light is just part of the electromagnetic
spectrum – the part detectable by the human eye
 the energy absorbed by and given off by objects was
thought to be continuous
 Max Planck, in 1900, proposed that energy is only available in
discrete packets, or quanta (singular is quantum); the size of
the quantum is related to the frequency of the radiation by
a simple equation
Bohr Model of the Atom (1913)
 Bohr applied Planck’s idea of quantized energy to the model of
the atom.
 Electrons move around the nucleus in specific orbits, or energy
levels. These energy levels are said to be quantized.
 An atom has several orbits, each representing a specific
energy level. The energy levels are like steps of a ladder.
Only specific energy values (steps) can exist – there are none
in between.
 When an atom is not excited (ground state), its electrons are
in orbits close to the nucleus. The electrons are at the lowest
energy level.
 If an atom gains energy (excited state), an electron is
displaced farther away from the nucleus to one of the higher
energy levels.
An atom emits energy when an electron falls from a higher
energy level to a lower energy level in one sudden drop, or
transition. The energy released is electromagnetic energy.
 The frequency () of the radiation emitted depends on the
difference between the higher and lower energy levels
involved in the transition. Remember that a specific frequency
correlates with a specific wavelength (and color!) through
c = .
Bohr model of the atom
Bohr used this model to explain the discontinuous line spectrum
of excited hydrogen atoms.
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Electron is close to the nucleus of the atom  PE is low
Electron is further away from the nucleus  higher PE
Absorption of energy increases the PE of the electron as it
moves further away from the nucleus.
Thus, an electron of an excited atom absorbs energy in the
form of heat or electricity when it moves to a higher energy
level. When the electron returns to ground state, it emits the
energy it absorbed. The packet of energy emitted
corresponds to a specific color on the line spectrum.
Problems with Bohr’s Model:
1. it only explained the line spectrum for hydrogen
2. he couldn’t explain why the electrons, which he pictured as
circling the nucleus like tiny planets, did not fall into the
nucleus every time they emitted light
Wave/Particle Paradox:
 Light sometimes acts as a particle and sometimes acts as a
wave.
 Thomas Young noticed that when electrons are forced through
a narrow slit, a pattern of wave interference emerged. He was
familiar with Rutherford’s experiment that showed electrons
are particles and needed to reconcile the two points of view.
 This led to the development of the Quantum Mechanical Model
of the atom.
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The Quantum Mechanical Model is based on complex
mathematical equations developed by Erwin Schrodinger that
describe the wave nature of the atom and the locations of the
electrons.
Bohr Model vs. Quantum Theory
Bohr
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nuclear model
electrons orbit the nucleus like
planets around the sun
electrons are treated as a
particle
electrons’ energy is quantized
Quantum Theory
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nuclear model
electrons occupy orbitals which
are probability spaces described
by a mathematical equation
electrons are treated as a wave
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electrons’ energy is quantized
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QUANTUM THEORY
based partly on Heisenberg’s Uncertainty Principle
 the position and the momentum of a moving object cannot
simultaneously be measured and known exactly
 there is an inherent limitation to knowing both where a particle is at a
particular moment and how it is moving in order to predict where it will
be in the future
an electron is in an orbital (electron cloud) – probability space where an
electron can be found a certain percentage of the time as defined
by Schrodinger’s equations
Schrodinger’s equations
 Schrodinger’s equations are like your class schedule on file in the
office
 the equations tell us where we are most likely to find an electron
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your schedule tells us where we are most likely to find you during, for
example, 2nd period
there is no guarantee you will be in your 2nd period classroom
there is no guarantee that the electron will be in particular location at
a particular time
there is information on your schedule that help us find you
there is information derived from Schrodinger’s equations (once all
the mathematics has been done) that describes where a particular
electron is likely to be found
FOUR QUANTUM NUMBERS – information to describe the location
of an electron
Principal quantum number (n)
 values = 1,2,3,4,5,6,7
 describes the approximate distance of the electron from the
nucleus and therefore its energy
2
 maximum capacity for an energy level = 2n
Sublevel quantum number (l)
 values = 0,1,2,3…n-1
if n = 1
l = 0 (there is only one sublevel possible)
if n = 2
l = 0,1 (there are two sublevels)
if n = 3
l = 0,1,2 (there are three sublevels)
if n = 4
l = 0,1,2,3 (there are four sublevels)
 describes the shape of the electron cloud
if l = 0 then the orbital
shape is spherical
(called the s sublevel)
if l = 1 then the orbital
shape is like a dumb-bell
(called the p sublevel)
insert diagram
if l = 2 then the orbitals are shaped is like… (called the d sublevel)
if l = 3 then the orbitals are shaped like… (called the f sublevel)
Magnetic quantum number (lm)
 values depend on l
 describes the orientation of the electron cloud in space
for l = 0, lm = 0 (no way to describe the orientation of a sphere)
for l = 1, lm = -1,0,1 (dumb-bell is along the x, y, or z axis)
for l = 2, lm = -2,-1,0,1,2
Spin quantum number (ls)
 values = + ½ or – ½
 describes the two possible orientations of the spin axis of an electron
 electrons act as though they are spinning on an axis like Earth spins on
its axis generates an electric field
 for two electrons to occupy the same orbital, they must have opposite
spins
DESCRIBING THE LOCATION OF ELECTRONS IN AN ATOM:
Aufbau Principle:
Electrons are added one at a time to the lowest energy orbitals available
until all the electrons of the atom have been accounted for.
Pauli Exclusion Principle:
No two electrons in an atom can have the same set of four quantum
numbers. Taken together, the four quantum numbers describe the state
of a particular electron. For example, an electron may have these
quantum numbers: 1,0,0,-1/2 (corresponding to n,l,lm, and ls). Think of
these as a social security number for an individual, or a street address
for a house. No two are identical.
Hund’s Rule:
Electrons occupy equal-energy orbitals so that a maximum number of
unpaired electrons results. In the 2p orbital with three equal energy
orbitals, for example, electrons go into 2p orbitals one by one until all
three orbitals have one electron. The next three electrons are then
paired with the previous three unpaired electrons.