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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.

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atomic structure

Atomic Structure

From Indivisible to Quantum Mechanical Model of the Atom

V.Montgomery & R.Smith

classical model
Classical Model
  • Democritus
  • Dalton
  • Thomson
  • Rutherford

V.Montgomery & R.Smith

democritus
Democritus
  • Circa 400 BC
  • Greek philosopher
  • Suggested that all matter is composed of tiny, indivisible particles, called atoms

V.Montgomery & R.Smith

dalton s atomic theory 1808
Dalton’s Atomic Theory (1808)
  • All matter is made of tiny indivisible particles called atoms.
  • 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 one another in simple whole number ratios to form compounds.
  • 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.

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j j thomson 1897
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

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rutherford s gold foil experiment 1910
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.

V.Montgomery & R.Smith

rutherford s gold foil experiment observations
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.

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rutherford s gold foil experiment conclusions
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.

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shortfalls of rutherford s model
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?

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bohr model 1913
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

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planetary model
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

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think about satellites
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?

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electronic structure of atom
Electronic Structure of Atom
  • Waves-particle duality
  • Photoelectric effect
  • Planck’s constant
  • Bohr model
  • de Broglie equation

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radiant energy
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

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slide16
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 light (c=3.00x108m/s) or about 300million m/s!!!

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waves
Waves

Waves transfer energy from one place to another

Think about the damage done by waves during strong hurricanes.

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

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wave characteristics
Wave Characteristics

Wavelength,  (lambda)  distance between

successive points

2m

10m

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wave characteristics19
Wave Characteristics
  • Frequency,  (nu)  the number of complete wave cycles to pass a given point per unit of time; Cycles per second

t=5

t=0

t=0

t=5

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units for frequency
Units for Frequency
  • 1/s
  • s-1
  • 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

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waves21
Waves
  • Amplitude maximum height of a wave

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waves22
Waves
  • Node points of zero amplitude

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electromagnetic spectrum
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

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electromagnetic spectrum24
Electromagnetic Spectrum

Frequency Hz

1024 1020

1018 1016 1014

1012 1010 108 106

Gamma Xrays UV

Microwaves FM AM

IR

10-16 10-9 10-8 10-6 10-3 100 102

Wavelength m

Visible Light

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notes
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

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what is your favorite radio station
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?

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velocity of a wave
Velocity of a Wave
  • Velocity of a wave (m/s) = wavelength (m) x frequency (1/s)
  • c = 
  • c= speed of light = 3.00x108 m/s
  • My favorite radio station is 105.9 Jamming Oldies!!!
  • What is the wavelength of this FM station?

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wavelength of fm
Wavelength of FM
  • c = 
  • c= speed of light = 3.00x108 m/s
  •  = 105.9MHz or 1.059x108Hz
  •  = c/ =3.00x108 m/s = 2.83m

1.059x1081/s

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what does the electromagnetic spectrum have to do with electrons
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

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states of electrons
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

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neon signs
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

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white light
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?

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line emission spectrum
Line-emission Spectrum
  • When an electric current is passed through a vacuum tube containing H2 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?

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hydrogen s emission spectrum
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

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bohr s model of hydrogen 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

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bohr model continued
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)

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bohr model and hydrogen spectrum
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)

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bohr s calculations
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

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photoelectric effect
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.

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photoelectric effect pg 93
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.

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photoelectric effect41
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?

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explanation
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.

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planck s equation
Planck’s Equation
  • E radiation = Planck’s constant x frequency of radiation
  • E = h
  • h = Planck’s constant = 6.626 x 10-34 J•s
  • When an object emits radiation, there must be a minimum quantity of energy that can be emitted at any given time.

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einstein expands planck s theory
Einstein Expands Planck’s Theory
  • 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

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photons
Photons
  • Photons: a particle of electromagnetic radiation having zero mass and carrying a quantum of energy
  • Ephoton = h

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back to photoelectric effect
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

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example of planck s equation
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

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answer
Answer
  • Ephoton = h
  • h = Planck’s constant = 6.626 x 10-34 J•s
  • c = 
  • c= speed of light = 3.00x108 m/s
  • = (3.00x108 m/s)/(6.85x10-7m)
  • =4.37x10141/s
  • Ephoton= (6.626 x 10-34 J•s)(4.37x10141/s)

Ephoton= 2.90 x 10-19J

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wave nature of electrons
Wave Nature of Electrons
  • We know electrons behave as particles
  • In 1925, Louis de Broglie suggested that electrons might also display wave properties

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de broglie s equation
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 ()

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example of de broglie s equation
Example of de Broglie’s Equation
  • Calculate the wavelength associated with an e- of mass 9.109x10-28 g traveling at 40.0% the speed of light.
  • 1 J = 1 kg m2/s2

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answer52
Answer
  • C=(3.00x108m/s)(.40)=1.2x108m/s
  •  = h/mv
  •  = (6.626 x 10-34 J•s) =6.06x10-12m

(9.11x10-31kg)(1.2x108m/s)

Remember 1J = 1(kg)(m)2/s2

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wave particle duality
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

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quantum mechanical model
Quantum mechanical model
  • SchrÖdinger
  • Heisenberg
  • Pauli
  • Hund

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where are the e in the atom
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.

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heisenberg s idea
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-

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heisenberg s uncertainty principle
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
SchrÖdinger’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 equation59
SchrÖdinger’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 equation60
SchrÖdinger’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

-h (ә2 Ψ )+ (ә2Ψ )+( ә2Ψ ) +Vψ =Eψ

8(π)2m (әx2) (әy2) (әz2 )

Scary???

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definitions
Definitions
  • Probability  likelihood
  • Orbital  wave function; region in space where the probability of finding an electron is high
  • SchrÖdinger’s Wave Equation states that orbitals have quantized energies
  • But there are other characteristics to describe orbitals besides energy

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quantum numbers
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 SchrÖdinger’s Wave Equation

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quantum numbers 1
Quantum Numbers (1)
  • Principal Quantum Number, n

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quantum numbers64
Quantum Numbers
  • Principal Quantum Number, n
    • Values of n = 1,2,3,… 
    • Positive integers only!
    • Indicates the main energy level occupied by the electron

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quantum numbers65
Quantum Numbers
  • Principal Quantum Number, n
    • Values of n = 1,2,3,… 
    • Describes the energy level, orbital size

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quantum numbers66
Quantum Numbers
  • Principal Quantum Number, n
    • Values of n = 1,2,3,… 
    • Describes the energy level, orbital size
    • As n increases, orbital size increases.

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principle quantum number
Principle Quantum Number

n=6

n=5

n=4

n=3

Energy

n=2

n = 1

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principle quantum number68
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 in a given shell = n2

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quantum numbers 2
Quantum Numbers (2)
  • Angular momentum quantum number, l

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quantum numbers70
Quantum Numbers
  • Angular momentum quantum number, l
    • Values of l = n-1, 0

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quantum numbers71
Quantum Numbers
  • Angular momentum quantum number, l
    • Values of l = n-1, 0
    • Describes the orbital shape

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quantum numbers72
Quantum Numbers
  • Angular momentum quantum number, l
    • Values of l = n-1, 0
    • Describes the orbital shape
    • Indicates the number of sublevel (subshells)

(except for the 1st main energy level, orbitals of different shapes are known as sublevels or subshells)

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orbital shapes
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.

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orbital shapes74
Orbital Shapes
  • Angular magnetic quantum number, l
  • If l = 0, then the orbital is labeled s.
  • s is spherical.

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orbital shapes75
Orbital Shapes
  • If l = 1, then the orbital is labeled p.
  • “dumbbell” shape

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orbital shapes76
Orbital Shapes
  • If l = 2, the orbital is labeled d.
  • “double dumbbell” or four-leaf clover

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orbital shapes77
Orbital Shapes
  • If l = 3, then the orbital is labeled f.

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energy level and orbitals
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

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atomic orbitals
Atomic Orbitals
  • Atomic Orbitals are designated by the principal quantum number followed by letter of their subshell
    • Ex. 1s = s orbital in 1st main energy level
    • Ex. 4d = d sublevel in 4th main energy level

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quantum numbers 3
Quantum Numbers (3)
  • Magnetic Quantum Number, ml

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quantum numbers81
Quantum Numbers
  • Magnetic Quantum Number, ml
    • Values of ml = +l…0…-l

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quantum numbers82
Quantum Numbers
  • Magnetic Quantum Number, ml
    • Values of ml = +l…0…-l
    • Describes the orientation of the orbital
      • Atomic orbitals can have the same shape but different orientations

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magnetic quantum number
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

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quantum numbers 4
Quantum Numbers (4)
  • Electron Spin Quantum Number,ms

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quantum numbers85
Quantum Numbers
  • Electron Spin Quantum Number,ms
    • Values of ms = +1/2 or –1/2
    • e- spin in only 1 or 2 directions
    • A single orbital can hold a maximum of 2 e-, which must have opposite spins

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electron configurations
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:

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pauli exclusion principle
Pauli Exclusion Principle
  • No 2 e- in an atom can have the same set of four quantum numbers (n, l, ml, ms ). Therefore, no atomic orbital can contain more than 2 e-.

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aufbau principle
Aufbau Principle
  • Aufbau Principle: an e- occupies the lowest energy orbital that can receive it.
  • Aufbau order:

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hund s rule
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

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electron configuration
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.

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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.
  • H 1s1

He 1s2

Li 1s2, 2s1

Be 1s2, 2s2

B 1s2, 2s2, 2p1

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orbital diagrams
Orbital 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.

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orbital diagram rules
Orbital Diagram Rules
  • Represent each electron by an arrow
  • The direction of the arrow represents the electron spin
  • Draw an up arrow to show the first electron in each orbital.
  • 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.

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orbital diagram examples
Orbital Diagram Examples
  • H _

1s

  • Li _

1s 2s

  • B  __ __

1s 2s 2p

  • N   _

1s 2s 2p

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dot diagram of valence electrons
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.

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rules for dot diagrams
Rules for Dot Diagrams

:Xy:

. .

Px orbital

S sublevel electrons

. .

Py orbital

Pz orbital

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rules for dot diagrams97
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.

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examples of dot diagrams
Examples of Dot Diagrams
  • H
  • He
  • Li
  • Be

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examples of dot diagrams99
Examples of Dot Diagrams
  • C
  • N
  • O
  • Xe

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summary
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.

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