Download - Chem 101 week 8 ch7
Chapter 7Electronic Structure
of Atoms
Recall:
• The concept of atoms proposed by Dalton explained many important observations:
- such as why compounds always have the same composition and how chemical rxns. occur• Once chemists were convinced of the
existence of atoms…..they naturally began to ask what atoms looked like?
The studies of Thomson, Rutherford and Chadwick
• Lead to our picture of the atom: - which includes a solid dense nucleus
containing protons and neutrons about which we find the negatively charged
electrons
In this Chapter we will look at the atomic structure in more detail.
• In particular we will develop a picture of the electron arrangements in atoms
• This picture will allow us to account for the chemistry of the elements
• With this knowledge we will revisit the arrangements of the elements in the periodic table and see why there are striking differences in the chemical properties of elements in the groups
Electromagnetic Radiation and Energy
• The sun is a central source of energy
• Energy from the sun travels through space in the form of electromagnetic radiation
Forms of Electromagnetic Radiation
• Visible light is a form of electromagnetic radiation• Microwaves are another form of electromagnetic
radiation• X-rays are yet another form of electromagnetic
radiation***all of these forms of energy are different, but they are
all forms of electromagnetic radiation
Electromagnetic Radiation
• All forms of electromagnetic radiation exhibit wave-like behavior and travel at the speed of light in a vacuum (airless space)
speed of light = 3.8 x 1010cm/s or 186,000 mi/s
Waves
• Waves have 3 primary characteristics:1. Wavelength- λ (Greek letter lambda) is the distance between 2
consecutive peaks or troughs in a wave2. Frequency- υ (Greek letter nu) it indicates how many waves
pass a given point per second3. Speed- indicates how fast a given peak moves through space
• Electromagnetic radiation from the sun is divided into various classes (forms) according to λ (wavelength)
• Energy from the sun reaches the earth in the form of visible and ultra-violet light
• Hot coals emit infrared radiation• Microwave ovens use microwave radiation to heat food**** all have different wavelengths
Visible Light
• Also known as white light• When passed through a prism
it separates into a continuous range of colors with one gradually blending into another called the continuous spectrum(rainbow)
• Separation is due to different speeds in the prisim
• Each color of light has a specific wavelength
Continuous Spectrum
• Violet has the shortest wavelength• Red has the longest wavelength
Wavelength and energy
• Wavelength and energy have an inverse relationship, as shown below
• h is Planck’s constant (6.626 10-34 J·s)
• c is the speed of light
hc
EE 1
Energy and Wavelength
• Red light with the longest wavelength has the lowest energy
• Purple light with the shortest wavelength has the highest energy
The Nature of Energy
• The wave nature of light does not explain how an object can glow when its temperature increases.
• Max Planck explained it by assuming that energy comes in packets called quanta.
Photoelectric Effect
A freshly polished, negatively charged zinc plate looses its charge if it is exposed to ultraviolet light. This phenomenon is called the photoelectric effect.
The Nature of Energy
• Einstein used this assumption to explain the photoelectric effect.
• He concluded that energy is proportional to frequency:
E = hwhere h is Planck’s constant, 6.63 10−34 J-s.
The Nature of Energy
• Therefore, if one knows the wavelength of light, one can calculate the energy in one photon, or packet, of that light:
c = E = h
The Nature of Energy
Another mystery involved the emission spectra observed from energy emitted by atoms and molecules.
Heating an Element
• When an element is heated strongly to the point that the element changes phase to a gas, the gaseous atoms emit light like the sun
• One might think that the light produced would result in a continuous spectrum like light emitted from the sun…………instead…….
When an element is heated..
• Only definite or discrete colors are produced• Since colors are discrete (definite) and the colors
correspond to energies• The energy being emitted by the atoms of the
element is also discrete
An Elements Spectrum
• Is unique to that element• Are called Atomic Emission Spectra or Line
Spectra• Are the basis of a fireworks display
The Nature of Energy
• Niels Bohr adopted Planck’s assumption and explained these phenomena in this way:
1. Electrons in an atom can only occupy certain orbits (corresponding to certain energies).
The Nature of Energy
• Niels Bohr adopted Planck’s assumption and explained these phenomena in this way:
2. Electrons in permitted orbits have specific, “allowed” energies; these energies will not be radiated from the atom.
The Nature of Energy
• Niels Bohr adopted Planck’s assumption and explained these phenomena in this way:
3. Energy is only absorbed or emitted in such a way as to move an electron from one “allowed” energy state to another; the energy is defined by
E = h
The Nature of Energy
The energy absorbed or emitted from the process of electron promotion or demotion can be calculated by the equation:
E = −RH ( )1nf
2
1ni
2-
where RH is the Rydberg constant, 2.18 10−18 J, and ni and nf are the initial and final energy levels of the electron.
Electronic Structure
Good Points• Electrons in Quantized
Energy Levels• Maximum # electrons in
each n is 2n2
• Sublevels (s,p,d,f) and # electrons they hold
Bad Points• Electrons are placed in
orbits about nucleus• Only explains emission
spectra of H2
• Does not address all interactions
• Treats electron as particle
The Wave Nature of Matter
• Louis de Broglie postulated that if light can have material properties, matter should exhibit wave properties.
• He demonstrated that the relationship between mass and wavelength was
=h
mv
The Uncertainty Principle
• Heisenberg showed that the more precisely the momentum of a particle is known, the less precisely is its position known:
• In many cases, our uncertainty of the whereabouts of an electron is greater than the size of the atom itself!
(x) (mv) h4
Dual Nature of Electron
Previous Concept; A Substance is Either Matter or Energy
• Matter; Definite Mass and Position Made of Particles
• Energy; Massless and Delocalized Position not Specificed Wave-like
Dual Nature of Electron
• Electron is both “particle-like” and “wave-like” at the same time.
• Previous model only considered “particle-like” nature of the electron
Orbitals Replace Orbits
• Bohr Orbits- Both electron position and energy known with certainty
• Orbitals – Regions of space where an electrons of a given energy will most likely be found
Quantum TheoryOrbitals Replace Orbits
Orbits Orbitals
Quantum Mechanics
• Erwin Schrödinger developed a mathematical treatment into which both the wave and particle nature of matter could be incorporated.
• It is known as quantum mechanics.
Schrodinger Wave Equation ()
Describes size/shape/orientation of orbitals
7.5
• Wave Equation is based on…
1. Dual Nature of Electron (Electron both particle and wave-like at the same time.)
2. Heisenberg Uncertainty Principle(Orbitals describe a region in space an electron will most likely be.)
Wave Equation ()
• Wave Equation describes the size, shape, and orientation of the orbital the electron (of a given energy) is in. There are four variables in the function
-n; Energy and size of orbital– l; Shape of orbital– ml; Orientation of orbital
– ms; Electron Spin
(n, l, ml, ms)
Quantum Mechanics
• The square of the wave equation, 2, gives a probability density map of where an electron has a certain statistical likelihood of being at any given instant in time.
1. Each electron has a unique set of 4 Quantum Numbers
2. Each orbital described by the Quantum Numbers can hold a maximum of 2 electrons.
Principal Quantum Number, n
• The principal quantum number, n, describes the energy level on which the orbital resides.
• The values of n are integers ≥ 0.• n= 1, 2, 3, 4, ….
distance of e- from the nucleus
Azimuthal Quantum Number, l
• This quantum number defines the shape of the orbital.
• Allowed values of l are integers ranging from 0 to n − 1.
• We use letter designations to communicate the different values of l and, therefore, the shapes and types of orbitals.
Azimuthal Quantum Number, l
Value of l 0 1 2 3
Type of orbital s p d f
Magnetic Quantum Number, ml
• Describes the three-dimensional orientation of the orbital in space.
• Values are integers ranging from -l to l: −l ≤ ml ≤ l.
• Therefore, on any given energy level, there can be up to 1 s orbital, 3 p orbitals, 5 d orbitals, 7 f orbitals, etc.
Magnetic Quantum Number, ml
• Orbitals with the same value of n form a shell.• Different orbital types within a shell are subshells.
s Orbitals
• Value of l = 0.• Spherical in shape.• Radius of sphere
increases with increasing value of n.
s Orbitals
Observing a graph of probabilities of finding an electron versus distance from the nucleus, we see that s orbitals possess n−1 nodes, or regions where there is 0 probability of finding an electron.
p Orbitals
• Value of l = 1.• Have two lobes with a node between them.
d Orbitals
• Value of l is 2.• Four of the five
orbitals have 4 lobes; the other resembles a p orbital with a doughnut around the center.
f-orbitals
Orbital Shapes
Orbital Type Shape Name
s Spherical
p Dumbbell
d Complex
f More complex
Energies of Orbitals
• For a one-electron hydrogen atom, orbitals on the same energy level have the same energy.
• That is, they are degenerate.
Energies of Orbitals
• As the number of electrons increases, though, so does the repulsion between them.
• Therefore, in many-electron atoms, orbitals on the same energy level are no longer degenerate.
Spin Quantum Number, ms
• In the 1920s, it was discovered that two electrons in the same orbital do not have exactly the same energy.
• The “spin” of an electron describes its magnetic field, which affects its energy.
Spin Quantum Number, ms
• This led to a fourth quantum number, the spin quantum number, ms.
• The spin quantum number has only 2 allowed values: +1/2 and −1/2.
Pauli Exclusion Principle
• No two electrons in the same atom can have exactly the same energy.
• For example, no two electrons in the same atom can have identical sets of quantum numbers.
How many 2p orbitals are there in an atom?
2p
n=2
l = 1
If l = 1, then ml = -1, 0, or +1
3 orbitals
How many electrons can be placed in the 3d subshell?
3d
n=3
l = 2
If l = 2, then ml = -2, -1, 0, +1, or +2
5 orbitals which can hold a total of 10 e-
7.6
Three Manners to Convey How Electrons are Arranged
1. Electron Configuration ; List Orbitals and Number of Electrons in Each(1s22s22p63s2…)
2. Quantum Numbers (2,0,0,+1/2)3. Orbital Diagrams; List Orbitals and show
location of electrons and their spin
1s 2s 2p
Electron Configurations
• Distribution of all electrons in an atom
• Consist of – Number denoting the
energy level
Electron Configurations
• Distribution of all electrons in an atom
• Consist of – Number denoting the
energy level– Letter denoting the type of
orbital
Electron Configurations
• Distribution of all electrons in an atom.
• Consist of – Number denoting the
energy level.– Letter denoting the type of
orbital.– Superscript denoting the
number of electrons in those orbitals.
Writing Atomic ElectronConfigurations
Two ways of writing configurations:• One is called the spdf notation.spdf notation for H, atomic number = 1
1 s1
no. ofelectrons
value of lvalue of n
Writing Atomic ElectronConfigurations
• The spdf notation can be shortened using the noble gas notation.
spdf notation for K, atomic number = 191s22s22p63s23p64s1 OR [Ar]4s1
core e- valence e-
Orbital Diagrams
• Each box represents one orbital.
• Half-arrows represent the electrons.
• The direction of the arrow represents the spin of the electron.
Hund’s Rule
“For degenerate orbitals, the lowest energy is attained when the number of electrons with the same spin is maximized.”
7.7
Orbital Diagrams
1s 2s 2p
Carbon; 6 electrons
Electron Configuration; 1s22s22p2
Orbital Diagram
Paramagneticunpaired electrons
2p
Diamagneticall electrons paired
2p7.8
Periodic Table
• We fill orbitals in increasing order of energy.
• Different blocks on the periodic table, then correspond to different types of orbitals.
Order of orbitals (filling) in multi-electron atom
1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d < 5p < 6s7.7
Some Anomalies
• This occurs because the 4s and 3d orbitals are very close in energy.
• Cu29 is also an anomaly with an expected config. of
[Ar] 4s2 3d9 and an actual config. of
[Ar] 4s1 3d10
• These anomalies occur in f-block atoms, as well.
Ion Configurations
• To form cations from elements remove e-’sfrom the subshell with the highest n.P [Ne] 3s2 3p3 - 3e- → P3+ [Ne] 3s2 3p0
• For transition metals, remove ns electronsand then (n - 1) electrons.Fe [Ar] 4s2 3d6
loses 2 electrons → Fe2+ [Ar] 4s0 3d6
Na+: [Ne] Al3+: [Ne] F-: 1s22s22p6 or [Ne]
O2-: 1s22s22p6 or [Ne] N3-: 1s22s22p6 or [Ne]
Na+, Al3+, F-, O2-, and N3- are all isoelectronic with Ne
What neutral atom is isoelectronic with H- ?
H-: 1s2 same electron configuration as He
8.2