V.Montgomery & R.Smith - 2.1 Atomic Structure complete OK
Atomic Structure
From Indivisible to Quantum Mechanical Model of the Atom
Classical Model
Democritus Dalton Thomson Rutherford
Democritus
Circa 400 BC Greek philosopher Suggested that all matter is composed of tiny, indivisible particles, called atoms Dalton’s Atomic Theory (1808)
1. All matter is made of tiny indivisible particles called atoms.
2. Atoms of the same element are identical. The
atoms of any one element are different from those of any other element.Atoms of different elements can combine with 3. one another in simple whole number ratios to form compounds.
4. Chemical reactions occur when atoms are
separated, joined, or rearranged;however,
atoms of one element are not changed into atoms of another by a chemical reaction.J.J. Thomson (1897)
Determined the charge to mass ratio for electrons Applied electric and magnetic fields to cathode rays “Plum pudding” model of the atom Rutherford’s Gold Foil Experiment (1910)
Alpha particles (positively charged helium ions) from a radioactive source was directed toward a very thin gold foil.
A fluorescent screen was placed behind the Au foil to detect the scattering of alpha () particles.
Rutherford’s Gold Foil Experiment (Observations)
Most of the -particles passed through the foil.
Many of the -particles deflected at various angles.
Surprisingly, a few particles were deflected back from the Au foil.
Rutherford’s Gold Foil Experiment (Conclusions)
Rutherford concluded that most of the mass of an atom is concentrated in a core, called the atomic nucleus. The nucleus is positively charged. Most of the volume of the atom is empty space.
Shortfalls of Rutherford’s Model
Did not explain where the atom’s negatively charged electrons are located in the space
surrounding its positively charged nucleus.
We know oppositely charged particles attract each other What prevents the negative electrons from being drawn into the positive nucleus?Bohr Model (1913)
Niels Bohr (1885-1962), Danish scientist working with Rutherford Proposed that electrons must have enough energy to keep them in constant motion around the nucleus Analogous to the motion of the planets orbiting the sun
Planetary Model
The planets are attracted to the sun by gravitational force, they move with enough energy to remain in stable orbits around the sun.
Electrons have energy of motion that enables them to overcome the attraction for the positive nucleus
Think about satellites…
We launch a satellite into space with enough energy to orbit the earth The amount of energy it is given, determines how high it will orbit We use energy from a rocket to boost our satellite, what energy do we give electrons to boost them?
Electronic Structure of Atom
Waves-particle duality Photoelectric effect Planck’s constant Bohr model de Broglie equation Radiant Energy
Radiation the emission of energy in various forms
A.K.A. Electromagnetic Radiation Radiant Energy travels in the form of waves that have both electrical and magnetic impulses
Electromagnetic Radiation radiation that consists of wave-like electric and
magnetic fields in space, including light, microwaves, radio signals, and x-rays Electromagnetic waves can travel through empty space, at the speed of 8 light (c=3.00x10 m/s) or about
300million m/s!!! Waves
Waves transfer energy from one place to Think about the damage done by waves during strong hurricanes. another
Think about placing a tennis ball in your bath tub, if you create waves at one it, that energy is transferred to the ball at the other = bobbing
Electromagnetic waves have the same characteristics as other waves Wave Characteristics
Wavelength, (lambda) distance between successive points
10m 2m Wave Characteristics
Frequency, (nu) the number of complete wave cycles to pass a given
point per unit of time; Cycles per t=0 t=5 second t=0 t=5
Units for Frequency
1/s -1 s hertz, Hz Because all electromagnetic waves travel at the speed of light, wavelength is determined by frequency Low frequency = long wavelengths High frequency = short wavelengths Waves
Amplitude maximum height of a wave Waves
Node points of zero amplitude
Electromagnetic Spectrum
Radio & TV, microwaves, UV, infrared, visible light = all are examples of electromagnetic radiation (and radiant energy) Electromagnetic spectrum: entire range of electromagnetic radiation
Electromagnetic Spectrum
10 24 10 20 10 18 10 16 10 14 10 12 10 10 10 8 10 6 Gamma Xrays UV Microwaves FM AM IR Visible Light Frequency Hz 10 -16 10 -9 10 -8 10 -6 10 -3 10 10 2 Wavelength m
Notes
Higher-frequency electromagnetic waves have higher energy than lower- frequency electromagnetic waves All forms of electromagnetic energy interact with matter, and the ability of these different waves to penetrate matter is a measure of the energy of the waves
What is your favorite radio station?
Radio stations are identified by their frequency in MHz.
We know all electromagnetic radiation(which includes radio waves) travel at the speed of light. What is the wavelength of your favorite station? Velocity of a Wave
Velocity of a wave (m/s) = wavelength (m) x frequency (1/s) c = 8 c= speed of light = 3.00x10 m/s
My favorite radio station is 105.9 Jamming Oldies!!! What is the wavelength of this FM station? Wavelength of FM
c = 8 c= speed of light = 3.00x10 m/s 8 Hz = 105.9MHz or 1.059x10 8 m/s = 2.83m
= c/ =3.00x10 8 1.059x10 1/s What does the electromagnetic spectrum have to do with electrons?
It’s all related to energy – energy of motion(of electrons) and energy of light
States of Electrons
When current is passed through a gas at a
low pressure, the potential energy (energy
due to position) of some of the gas atoms
increases.Ground State: the lowest energy state of an atom Excited State: a state in which the atom has a higher potential energy than it had in its ground state
Neon Signs
When an excited atom returns to its ground state it gives off the energy it gained in the form of electromagnetic radiation! The glow of neon signs,is an example of this process
White Light
White light is composed of all of the colors of the spectrum = ROY G BIV When white light is passed through a prism, the light is separated into a spectrum, of all the colors What are rainbows?
Line-emission Spectrum
When an electric current is passed through a vacuum tube containing H 2 gas at low pressure, and emission of a pinkish glow is observed.
What do you think happens when that pink glow is passed through a prism?
Hydrogen’s Emission Spectrum
The pink light consisted of just a few specific frequencies, not the whole range of colors as with white light Scientists had expected to see a continuous range of frequencies of electromagnetic radiation, because the hydrogen atoms were excited by whatever amount of energy was added to them.
Lead to a new theory of the atom Bohr’s Model of Hydrogen Atom Hydrogen did not produce a continuous spectrum New model was needed:
Electrons can circle the nucleus only in allowed paths or orbits When an e- is in one of these orbits, the atom has a fixed, definite energy
e- and hydrogen atom are in its lowest energy state when it is in the orbit closest to the nucleus
Bohr Model Continued…
Orbits are separated by empty space, where e- cannot exist
Energy of e- increases as it moves to orbits farther and farther from the nucleus (Similar to a person climbing a ladder)
Bohr Model and Hydrogen Spectrum
While in orbit, e- can neither gain or lose energy But, e- can gain energy equal to the difference between higher and lower orbitals, and therefore move to the higher orbital (Absorption) When e- falls from higher state to lower state, energy is emitted (Emission)
Bohr’s Calculations
Based on the wavelengths of hydrogen’s line-emission spectrum, Bohr calculated the energies that an e- would have in the allowed energy levels for the hydrogen atom
Photoelectric Effect
An observed phenomenon, early 1900s When light was shone on a metal, electrons were emitted from that metal Light was known to be a form of energy, capable of knocking loose an electron from a metal Therefore, light of any frequency could supply enough energy to eject an electron.
Photoelectric Effect pg. 93
Light strikes the surface of a metal (cathode), and e- are ejected.
These ejected e- move from the cathode to the anode, and current flows in the cell.
A minimum frequency of light is used. If the
frequency is above the minimum and the
intensity of the light is increased, more e- are ejected.Photoelectric Effect
Observed: For a given metal, no electrons were emitted if the light’s frequency was below a certain minimum, no matter how long the light was shone Why does the light have to be of a minimum frequency? Explanation….
Max Planck studied the emission of light by hot objects Proposed: objects emit energy in small, specific amounts = quanta
(Differs from wave theory which would say objects emit electromagnetic radiation continuously)
Quantum: is the minimum quantity of energy that can be lost or gained by an atom. Planck’s Equation
E = Planck’s constant x frequency radiation of radiation E = h
- -34 • h = Planck’s constant = 6.626 x 10 J s When an object emits radiation, there must be a minimum quantity of energy that can be emitted at any given time. Einstein Expands Planck’s Theory
- -34 • h = Planck’s constant = 6.626 x 10 J s c = 8 c= speed of light = 3.00x10 m/s 8 m/s)/(6.85x10-7m)
- h ( ә Ψ )+ (ә Ψ )+( ә Ψ ) +Vψ =Eψ 2 2) 2) 2 8(π) m (әx (әy (әz ) Scary??? Definitions
Theorized that electromagnetic radiation had a dual wave-particle nature! Behaves like waves and particles Think of light as particles that each carry one quantum of energy =
photons Photons
Photons: a particle of electromagnetic radiation having zero mass and carrying a quantum of energy E photon
= h Back to Photoelectric Effect
Einstein concluded:
Electromagnetic radiation is absorbed by matter only in whole numbers of photons In order for an e- to be ejected, the e- must be struck by a single photon with minimum frequency Example of Planck’s Equation
CD players use lasers that emit red light with a of 685 nm. Calculate the energy of one photon.
Different metals require different minimum frequencies to exhibit photoelectric effect Answer E photon = h
= (3.00x10 14 1/s =4.37x10 -34 14 •
E = (6.626 x 10 J s)(4.37x10 1/s) photon
-19
E = 2.90 x 10 J photonWave Nature of Electrons
We know electrons behave as particles In 1925, Louis de Broglie suggested that electrons might also display wave properties de Broglie’s Equation
A free e- of mass (m) moving with a velocity (v) should have an associated wavelength: = h/mv Linked particle properties (m and v) with a wave property ()
Example of de Broglie’s Equation
Calculate the wavelength associated -28 with an e- of mass 9.109x10 g traveling at 40.0% the speed of light. 2 2
1 J = 1 kg m /s Answer
C=(3.00x10 8 m/s)(.40)=1.2x10 8 m/s = h/mv
= (6.626 x 10 -34 J • s) =6.06x10 -12 m
(9.11x10 -31 kg)(1.2x10 8 m/s) Remember 1J = 1(kg)(m) 2 /s 2 Wave-Particle Duality
de Broglie’s experiments suggested that e- has wave-like properties.
Thomson’s experiments suggested that
e- has particle-like properties
measured charge-to-mass ratio Quantum mechanical model Ö
Schr dinger Heisenberg Pauli Hund
Where are the e- in the atom?
e- have a dual wave-particle nature If e- act like waves and particles at the same time, where are they in the atom? First consider a theory by German theoretical physicist, Werner Heisenberg.
Heisenberg’s Idea
e- are detected by their interactions with photons Photons have about the same energy as
e- Any attempt to locate a specific e- with a photon knocks the e- off its course ALWAYS a basic uncertainty in trying to locate an e-
Heisenberg’s Uncertainty Principle
Impossible to determine both the position and the momentum of an e- in an atom simultaneously with great certainty.
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Schr dinger’s Wave Equation
An equation that treated electrons in atoms as waves Only waves of specific energies, and therefore frequencies, provided solutions to the equation Quantization of e- energies was a natural outcome
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Schr dinger’s Wave Equation
Solutions are known as wave functions Wave functions give ONLY the probability of finding and e- at a given place around the nucleus
e- not in neat orbits, but exist in regions called orbitals
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Schr dinger’s Wave Equation
Here is the equation Don’t memorize this or write it down
It is a differential equation, and we need
calculus to solve it 2 2 2Probability likelihood Orbital wave function; region in space where the probability of finding an electron is high
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Schr dinger’s Wave Equation states that orbitals have quantized energies But there are other characteristics to describe orbitals besides energy Quantum Numbers
Definition: specify the properties of atomic orbitals and the properties of electrons in orbitals There are four quantum numbers The first three are results from
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Schr dinger’s Wave Equation Quantum Numbers (1)
Principal Quantum Number, n Quantum Numbers
Principal Quantum Number, n Values of n = 1,2,3,… Positive integers only! Indicates the main energy level occupied
by the electron Quantum Numbers
Principal Quantum Number, n
Values of n = 1,2,3,… Describes the energy level, orbital size Quantum Numbers
Principal Quantum Number, n
Values of n = 1,2,3,… Describes the energy level, orbital size
As n increases, orbital size increases. Principle Quantum Number Energy n = 1 n=2 n=3 n=4 n=5 n=6
Principle Quantum Number
More than one e- can have the same n value These e- are said to be in the same e- shell The total number of orbitals that exist 2 in a given shell = n Quantum Numbers (2)
Angular momentum quantum number,
l Quantum Numbers
Angular momentum quantum number,
l
Values of
l = n-1, 0 Quantum Numbers
Angular momentum quantum number,
l
Values of
l = n-1, 0
Describes the orbital shape Quantum Numbers
Angular momentum quantum number,
l
Values of l = n-1, 0 Describes the orbital shape Indicates the number of sublevel (subshells)
st (except for the 1 main energy level, orbitals of diferent shapes are known as sublevels or subshells) Orbital Shapes
For a specific main energy level, the number of orbital shapes possible is equal to n.
Values of l = n-1, 0 Ex. Orbital which n=2, can have one of two shapes corresponding to l = 0 or l=1
Depending on its value of l , an orbital is assigned a letter. Orbital Shapes
Angular magnetic quantum number,
l
If s .
l = 0, then the orbital is labeled s is spherical. Orbital Shapes
If p .
l = 1, then the orbital is labeled
“dumbbell” shape Orbital Shapes
If d .
l = 2, the orbital is labeled
“double dumbbell” or four-leaf clover Orbital Shapes
If
l = 3, then the orbital is labeled
f .Energy Level and Orbitals
n=1, only s orbitals n=2, s and p orbitals n=3, s, p, and d orbitals n=4, s,p,d and f orbitals
Remember: l = n-1 Atomic Orbitals
Atomic Orbitals are designated by the principal quantum number followed by letter of their subshell st Ex. 1s = s orbital in 1 main energy level th Ex. 4d = d sublevel in 4 main energy level Quantum Numbers (3)
Magnetic Quantum Number, m
l Quantum Numbers
Magnetic Quantum Number, m
l
Values of
m l
= +
l…0…-l Quantum Numbers
Magnetic Quantum Number, m
l m
Values of = +
l…0…-l l
Describes the orientation of the orbital
Atomic orbitals can have the same shape but different orientations
Magnetic Quantum Number
s orbitals are spherical, only one orientation, so m=0 p orbitals, 3-D orientation, so m= -1, 0 or 1 (x, y, z) d orbitals, 5 orientations, m= -2,-1, 0, 1 or 2 Quantum Numbers (4)
Electron Spin Quantum Number,m s Quantum Numbers
Electron Spin Quantum Number,m s Values of m = +1/2 or –1/2 s
e- spin in only 1 or 2 directions A single orbital can hold a maximum of 2 e-, which must have opposite spins
Electron Configurations
Electron Configurations: arragenment of
e- in an atom There is a distinct electron configuration for each atom
There are 3 rules to writing electron configurations: Pauli Exclusion Principle
No 2 e- in an atom can have the same set of four quantum numbers ( n , l , m l ,
m ). Therefore, no atomic orbital can s contain more than 2 e-.
Aufbau Principle
Aufbau Principle: an e- occupies the lowest energy orbital that can receive it.
Aufbau order:
Hund’s Rule
Hund’s Rule: orbitals of equal energy are each occupied by one e- before any orbital is occupied by a second e-, and all e- in singly occupied orbitals must have the same spin
Electron Configuration
The total of the superscripts must equal the atomic number (number of electrons) of that atom. The last symbol listed is the symbol for the differentiating electron.
Differentiating Electron
The differentiating electron is the electron
that is added which makes the configuration
different from that of the preceding element.
The “last” electron. 1 H 1s 2 He 1s 2 1 Li 1s , 2s 2 2 Be 1s , 2s 2 2 1 B 1s , 2s , 2pOrbital Diagrams
These diagrams are based on the electron configuration.
In orbital diagrams:
Each orbital (the space in an atom that will hold a pair of electrons) is shown.
The opposite spins of the electron pair is
indicated.Orbital Diagram Rules
1. Represent each electron by an arrow
2. The direction of the arrow represents the electron spin Draw an up arrow to show the first electron
3. in each orbital.
4. Hund’s Rule: Distribute the electrons among
the orbitals within sublevels so as to give
the most unshared pairs.Put one electron in each orbital of a sublevel before the second electron appears.
Half filled sublevels are more stable than partially full sublevels.
Orbital Diagram Examples
H _ 1s Li _ 1s 2s B __ __ 1s 2s 2p N _ 1s 2s 2p
Dot Diagram of Valence Electrons
When two atom collide, and a reaction takes
place, only the outer electrons interact.
These outer electrons are referred to as the
valence electrons.
Because of the overlaying of the sublevels in
the larger atoms, there are never more than
eight valence electrons.Rules for Dot Diagrams
P orbital
x. .
S sublevel P orbital y electrons
:Xy:
. .
P orbital z
Rules for Dot Diagrams
Remember: the maximum number of valence electrons is 8.
Only s and p sublevel electrons will ever be valence electrons.
Put the dots that represent the s and p electrons around the symbol.
Use the same rule (Hund’s rule) as you fill the designated orbitals.
Examples of Dot Diagrams
H He Li Be
Examples of Dot Diagrams
C N O Xe
Summary
Both dot diagrams and orbital diagrams will be
use full to use when we begin our study of atomic bonding. We have been dealing with valence electrons since our initial studies of the ions.The number of valence electrons can be determined by reading the column number. Al = 3 valence electrons Br = 7 valence electrons All transitions metals have 2 valence electrons.