Download Chapter 6 - Atomic Theory

Chapter 6
Atomic Theory
Greek Idea
• Democritus
– Matter is made up of
indivisible particles
• Dalton
– One type of atom for
each element
Thomson’s Model
• Discovered electrons
• Atoms were made of
positive “stuff”
• Negative electron
embedded inside the atom
• “Plum-Pudding” model
– [Blueberry muffin model ]
Rutherford’s Model
• Discovered dense
positive piece at the
center of the atom
– Nucleus
– Electrons moved
around
– Mostly empty space
Bohr’s Model
– Why don’t the electrons fall into the
nucleus?
– Electrons move like planets around the sun
• In circular orbits at different levels
– Amounts of energy separate one level from
another
Bohr’s Model
Nucleus
Electron
Orbit
Energy Levels
Bohr’s Model
Energy Levels
Increasing energy
Fifth
Fourth
Third
Second
First
Nucleus
• Further away
from the
nucleus means
more energy
• There is no “in
between”
energy
Bohr’s Model
Energy Levels
The Bohr model shows electrons
circling around the nucleus in
definite orbits or paths. The
electrons orbits the nucleus much
like planets circle the sun.
Electrons move from one orbit to
another. The further away the
orbit from the nucleus means
more energy. The electrons can
not move “in between” orbits;
electrons can only exist in an
orbit at definite energy levels.
The energy absorbed or released
when electrons change energy
levels is in the form of
electromagnetic radiation (light).
Energy
– When electrons “jump” orbits
(move up or down) energy is
needed or released respectively
• Energy is in the form of light
Energy
needed
2
– It was through the study of light
that several advances in science
were made
• Physicists determined the properties
of light
• The Quantum Mechanical model
was produced
1
Energy
released
Electromagnetic Radiation
– Light is just one component of the
Electromagnetic Spectrum
• Composed of waves of different energies
– Waves travel through empty space as well as
through air and other substances
– Electromagnetic radiation has a dual "personality“
• Acts like waves and particles (photons)
• The photons with the highest energy correspond to the
shortest wavelengths
Electromagnetic Spectrum
Electromagnetic Radiation
Parts of the Wave
Crest
Wavelength
Origin
Amplitude
Trough
Parts of a Wave
– Origin
Crest
Wavelength
• the base-line of energy
– Crest
Origin
• Highest point on a wave
Amplitude
– Trough
• Lowest point on a wave
Trough
– Amplitude
• Distance from origin to crest
– Wavelength
• Distance from crest to crest (abbreviated “l” - lambda)
– Frequency
• The number of waves passing a given point per second
• Units are cycles/sec or hertz (Hz)
Frequency and the Spectrum
– Frequency and wavelength are inversely related
n=C
l
Frequency = speed of light / wavelength
– Different frequencies of light have different energy levels
–E = hν (Energy = Planck’s constant x frequency)
• High frequency light has high energy (violet light)
• Low frequency light has low energy (red light)
– White light is made up of all the colors of the visible
spectrum, and all frequencies therein
– The whole range of colours is called a continuous
spectrum; one colour leads into the next with no break
And now back to the atom …
– Electrons occupy the lowest energy levels, making the
atom stable (low energy content; ground state)
– When electrons interact with energy (photon) they may
absorb it and move away from the nucleus
• Referred to as an electron transition
– The electrons are no longer in the ground state; they are in
an excited state
– The electrons can return to the ground state by releasing
quanta of energy
• Energy released is of a definite quantity
• The energy of electron transitions are quantized (fixed), quantum of
energy is released
Catch and Release
Quantized energy is
absorbed
– Electron is excited
into a higher energy
level
Quantized energy is
emitted (released)
– Electron jumps to a
lower energy level
Bohr Model
Emission Spectra
– Light of different frequencies have different characteristic
colours
• Low energy, low frequency light is seen as red light; high frequency,
high energy light is seen as violet light
– All possible jumps from one energy level to another can
occur at the same time because of the number of atoms
present, each giving off its characteristic colour
– The colour that we see is a mixture of colours
• Can be separated by using a diffraction grating
– Breaks up the predominant colour into bright lines of specific colours
representing electron transitions
Emission Spectra
Hydrogen Emission Spectrum
Iron Emission Spectrum
The Quantum Mechanical Model
– Bohr’s model of the atom introduced the concept
of quantum energy levels, but couldn’t explain how
electrons are arranged in atoms
– Louis de Broglie (1892-1987)
• Proposed that if waves can have particle-like behaviour,
particles of matter can behave like waves [under
appropriate conditions]
• Suggested that as an electron moves about the
nucleus, an appropriate wavelength is associated with it
– Just a few years later, experimental evidence
supported de Broglie …
The Quantum Mechanical Model
– Werner Heisenberg (1901-1976)
• Limit to how precisely we can know both location and
momentum of any object [only important with subatomic
particles]
– Impossible to know both exact momentum and exact location
of electron at any point in time
» Heisenberg Uncertainty Principle
– Erwin Schröndinger (1887-1961)
• Proposed an equation which leads to a series of wave
functions
– Related the probability of finding the electrons to a particular
volume of space (3D)
» Schröndinger’s wave equation
The Plot for atomic orbital
• The orbital is classified by its shape.
• S orbital a sphere.
• P orbital is a dumbbell shape.
Atomic Orbitals
– Represents the likelihood of finding an electron at
a particular point
• Densest near the nucleus; less dense with increasing
distance
– Indicates most probable location around the nucleus
– Scientists arbitrarily draw the orbital surface to contain 90% of
the total probability distribution
– To assign relative sizes and energies to orbitals,
the quantum mechanical model assigns Principal
Quantum Numbers (n)
• “n” specifies the atom’s major energy levels [Principal
Energy Levels]
Atomic Orbitals
– Within each energy level the complex math of
Schrödinger’s equation describes several shapes
• Each type of shape is identified as a sublevel containing a specific
atomic orbital
– Energy level 1 (n=1) has one sublevel that contains one
“s” orbital
• Every succeeding energy level contains an s sublevel with an s
orbital
• The s orbital is spherically shaped
– Each s sublevel is identified by its principal quantum
number n
• Called 1s, 2s, 3s, …
– Each s sublevel has one s type orbital, and can hold 2
electrons
Atomic Orbitals
– Energy level 2 (n=2) has 2 sublevels, the s and p sublevel
(The sublevel name comes from the type of orbital it contains.)
• Every succeeding energy level contains a p sublevel with three p
orbitals
– The p orbital is shaped like a dumbbell and labelled
depending on its orientation
• One p orbital is oriented along the x axis (px)
• One p orbital is oriented along the y axis (py)
• One p orbital is oriented along the z axis (pz)
– Each p orbital can hold 2 electrons, so the p sublevel can
hold a maximum of six electrons
Atomic Orbitals
– The third energy level (n=3) has three sublevels
• s sublevel with one s orbital
• p sublevel with three p orbitals
• d sublevel with five d orbitals
– The d orbitals are more complex in shape and are
oriented along planes not axes
– Each of the five orbitals can hold two electrons so
the d sublevel can hold a maximum of ten electrons
Atomic Orbitals
– The fourth energy level (n=4) has four sublevels
• s sublevel with one s orbital
• p sublevel with three p orbitals
• d sublevel with five d orbitals
• f sublevel with seven f orbitals
– The f orbital are even more complex in shape
– Each of the seven orbitals can hold two electrons
so the f sublevel can hold a maximum of fourteen
electrons
Orbitals Summary
– Atomic orbital pictures
– Atomic orbital density representations
Principal
quantum number
(n)
Number of
sublevels
Number of
orbitals
Total number of
electrons
n
n
n2
2n2
1
1 (s)
1
2
2
2 (s, p)
4
8
3
3 (s, p, d)
9
18
4
4 (s, p, d, f)
16
32
Electron Configuration
– The Electron Configuration is the arrangement of
the electrons within an atom
• The atomic orbitals do not fill up in a neat order; the
energy levels overlap
– Large complex orbital shapes cause the electron to be, on
average, further from the nucleus than the simple orbital shapes
from greater energy levels
– Arrangements follow specific principles:
• aufbau principle
– Electrons enter the lowest available energy level or sublevel
first
– This causes difficulties because of the overlap of orbitals of
different energies
Electron Configuration
– Pauli Exclusion Principle
• At most 2 electrons per orbital
– Electrons have the same charge and will repel each other
– The repulsion can be reduced if the electrons have different
spins (paired spins) leading to opposite magnetic force fields
– Hund’s Rule
• When electrons occupy orbitals of equal energy
(orbitals within the same sublevel) they don’t pair up until
the sublevel is half full
Increasing energy
7s
6s
5s
7p
6p
5p
4p
4s
3p
3s
2p
2s
1s
6d
5d
4d
3d
5f
4f
n=4
With more sublevels the energy level
n=3
n=2
n=1
gets wider (leads to an overlap).
Electron Configuration
7s
6s
5s
4s
3s
2s
1s
Electron Configuration
7p
6d
5f
6p
5d
4f
5p
4d
←
4p
3d
←
3p
←
2p
←
←
←
Start
←
←
Determining Electron Configuration
– Let’s determine the electron configuration for phosphorus
• Need to account for 15 electrons
– The first two electrons go into the 1s orbital (1s2)
• Remember electrons have opposite spins
– The next two electrons go into the 2s orbital (2s2)
– The next six electrons go to the 2p sublevel (2p6)
• 3 paired spins.
– Next is the 3s orbital with 2 electrons (3s2)
• Total is now 12, only 3 more to go
– The next three electrons go 3px1, 3py1, 3pz1
– The Electron Configuration for phosphorus is:
1s2 2s2 2p6 3s2 3px1, 3py1, 3pz1
1s2 2s2 2p6 3s2 3p3 (not showing the individual p orbitals)
Determining Electron Configuration
– Try for:
• Na
•C
• Cl
– Start off with “How many
electrons to place?”
– Add electrons to the orbitals
starting at 1s and following
through the table
Electron Configuration
7p
6d
5f
←
7s
6p
5d
4f
←
6s
5p
4d
←
5s
4p
3d
←
4s
3p
←
3s
2p
←
2s
←
1s
←
Start
Valence Shell and Valence Electrons
– The highest numbered energy level is given a
special name: the valence shell
• The electrons in this valence shell are referred to as
valence electrons
– Only the outer s, and p electrons
Valence Shell and Valence Electrons
As electrons fill the sublevels, we
see:
Electron Configuration
7p
6d
5f
←
←
• 1s2 2s2 2p6 3s2 all fill up first
7s
6p
5d
4f
6s
5p
4d
←
• 3p6 4s2  now we’re in the 4th
level, so we start in a new valence
shell
5s
4p
3d
←
4s
3p
←
3s
2p
←
2s
←
1s
←
• 3d10 4p6 5s2  again, a new shell
• The d and f electrons never “see”
valence because a new level is
started before they’re filled
Start
Valence Shell and Valence Electrons
– The highest numbered energy level is given a
special name: the valence shell
• The electrons in this valence shell are referred to as
valence electrons
– Only the outer s, and p electrons
–
Identify each element, state the valence shell
and the number of valence electrons for:
1s2
2s2
2p6
a.
b. 1s2 2s2 2p6 3s2 3p6 4s1
c. 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p2
d. 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2
Ne
K
Ge
Sr
Exceptions to the Electron
Configuration
– Orbitals fill in order of lowest to highest energy
• The most stable sublevel or energy level is “full”
• The next most stable is a “half full”
– Adding electrons can change the energy of the orbital or sublevel
– “More than half filled” sublevels or “less than half” filled
sublevels are higher in energy than those that are exactly
half full
• The most stable is a full sublevel (least energy)
– Sometimes electrons will be promoted/demoted to help the
atom stay in the lowest energy state (most stable) possible
Electron Exceptions
– Copper has 29 electrons so we expect
1s22s22p63s23p64s23d9
– The electron configuration for the copper is also
1s22s22p63s23p64s13d10
• One 4s electron has been promoted to a 3d orbital,
leaving a valence shell with only 1 electron and a full 3d
– Copper exists in both states
• 1s22s22p63s23p64s23d9
• 1s22s22p63s23p64s13d10
Copper Ions
– Because copper can exist in both states, it can
form 2 different cations
•Cu2+
– 1s22s22p63s23p64s03d9
4s2 electrons are lost
• Cu+
– 1s22s22p63s23p64s03d10
4s1 electron is lost
– Many transition metals can form multiple cations
Electron Exceptions
– Iron with 26 electrons has an electron
configuration of 1s22s22p63s23p64s23d6
– When reacting, it can lose the 2 valence 4s2
electrons creating a 2+ cation
• 1s22s22p63s23p64s03d6
– Iron can also lose one 3d electron in addition, to
yield a half full 3d sublevel 3d5, creating a 3+ cation
• 1s22s22p63s23p64s03d5
– most stable
Silver (Ag)
– Electron configuration should be
1s22s22p63s23p64s23d104p65s24d9
• It’s actually 1s22s22p63s23p64s23d104p65s14d10
– Results in a lower overall energy, and higher
stability
– Full d sublevel; partially filled s sublevel
E
5p
5s
4d
4f