copyright©2000 by houghton mifflin company. all rights reserved. 1 honors chemistry chapter 5

Copyright©2000 by Houghton Mifflin Company. All rights

reserved.

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Honors Chemistry

Chapter 5


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Max Planck (Early 1900’s)• Studied Radiation emitted by solid bodies

heated to incandescence.

• Thought of the day: Matter could absorb or emit any quantity of energy.

• He could not explain his results based on this!!!

• So, he proposed that energy can be gained or lost only in whole number intervals of h.


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Planck’s Constant

E = change in energy, in J

h = Planck’s constant, 6.626 1034 J s

= frequency, in s1

= wavelength, in m

E hhc

= =

Transfer of energy is quantized, and can Transfer of energy is quantized, and can only occur in discrete units, called only occur in discrete units, called quantaquanta..


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Energy is Quantized

• Discrete units of h.

• Each small “packet” of energy is called a Quantum.

• Energy is transferred in Whole Quanta.


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Surprise!

Energy seems to have particle-like properties!

Before - energy always assumed to be continuous.


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Albert Einstein• Proposed that electromagnetic Radiation

is itself Quantized, that is,

It can be viewed as a stream of Particles called Photons.

• Energy of a photon:E = h = h c/


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Energy and Mass

• Also, Einstein proposed that

Energy has mass

• E = mc2

• E = energy

• m = mass

• c = speed of light


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Dual Nature of Light

Electromagnetic radiation exhibits:

1) Wave Properties

2) Particulate Properties


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Figure 7.4

Dual Nature of Light


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Light thought to be purely wavelikewas found to have particulate properties.

Matter

thought to be purely particulateDoes it exhibit wave properties?


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Wavelength and Massde Broglie’s Equation (1923): Allows one to de Broglie’s Equation (1923): Allows one to calculate an apparent wavelength for a calculate an apparent wavelength for a particle.particle.


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• Continuous spectrum: Contains all the wavelengths of light.

• Line (discrete) spectrum: Contains only some of the wavelengths of light.


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Figure 7.6

A Continuous Spectrum (a)

and

A Hydrogen Line Spectrum (b)


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Sample of H2 gas (H—H)

• Introduce a high energy spark• H2 molecules absorb energy• Some of the H—H bonds break• Resulting H atoms are “EXCITED”,

i.e. contain excess energy.• They will eventually “relax” & will release

excess energy by emitting light of various wavelengths.

LINE SPECTRUM


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Line Spectrum

• Line spectrum results because only certain energies are allowed for the electron in H atom.

• That is, energy of electron in H atom is QUANTIZED.

E = h ν = h c/ λ


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Figure 7.7A Change between Two Discrete Energy Levels


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If any energy were allowed then we would see a Continuous Spectrum (a)

and

When only certain energies are possiblewe see only a discrete Line Spectrum (b)


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Niels Bohr (1913)

• Developed Quantum Model for the Hydrogen Atom.

• The Electron in a Hydrogen Atom moves around the nucleus only in certain allowed circular orbits.


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Figure 7.8

Electronic Transitions in the Bohr Model for the Hydrogen Atom

♦ He calculated the radii for theallowed circular orbits.

♦ Only certain electron energiesallowed .

♦ Energy levels consistent with the Hydrogen line-emission spectrum.


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The Bohr Model

• Ground State: The lowest possible energy state for an atom (n = 1).


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TWO IMPORTANT POINTS

• Bohr Model correctly fits Quantized Energy Levels of the H-atom.

Postulates only certain allowed circular orbits.

• As electron is brought closer to the nucleus, Energy is released from the system.


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Bohr’s Model

• Appeared promising.

• Calculations worked well for hydrogen.

• Didn’t work when applied to other atoms.

• Something fundamentally incorrect.

• Important for its introduction of the concept of Quantization of energy in atoms.


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Quantum Mechanical Model of the Atom

• Totally different approach was needed.

• Three physicists: Heisenberg, de Broglie, & Schrodinger.

• Emphasizes the wave properties of an electron.


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For the Electron in a Hydrogen Atom

• Electron bound to the nucleus.

• Similar situation of only certain allowable “Electron Waves.”

• Modeled by Schrödinger


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Solution of Schrodinger Equation for the Hydrogen Atom

• Atomic Orbitals:space that encloses 90% of the total

electron probability.

Wave function for an electron in the Hydrogen atom.

• Each electron described by 4 different quantum numbers.


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Section 7.6Quantum Numbers (QN)

• Four different quantum numbers.

• Three (n, l, ml) specify the wave function that gives the probability of finding the electron at various pts. in space.

• Fourth (ms) specifies the electron”spin”.


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Quantum Numbers (QN)

• 1. Principal QN (n = 1, 2, 3, . . .)

- related to size and energy of the orbital.

- Shell Number

- larger n, then higher energy


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2. Angular Momentum Quantum Number

• (l = 0 to n 1)

- relates to shape of the orbital.

- Subshell

l = 0 s subshell

l = 1 p subshell

l = 2 d subshell

l = 3 f subshell


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3. Magnetic Quantum Number

(ml = l to l)

- relates to orientation of the orbital in space relative to other orbitals.


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SHAPES----------

Two Representations of the Hydrogen 1s, 2s, and 3s Orbitals

Node – Area of Zero Probability.


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Figure 7.14

The Boundary Surface Representations of All Three 2p Orbitals

No 1p orbital.In 2p, two lobes separated by a node at the nucleus.Labeled according to orientation.


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Figure 7.16

The Boundary Surfaces of All of the 3d Orbitals

No 1d or 2d orbitals.

Five 3d orbital


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Figure 7.17

Representation of the 4f Orbitals in Terms of Their Boundary Surfaces


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4. Electron Spin Quantum Number

(ms = +1/2, 1/2)

- relates to the spin states of the electrons.


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Electron Spin & Pauli Exclusion Principle

• In a given atom, no two electrons can have the same set of four quantum numbers (n, l, ml, ms).

• Therefore, an orbital can hold only two electrons, and they must have opposite spins.


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Aufbau Principle

• As protons are added one by one to the nucleus to build up the elements, electrons are similarly added to these hydrogen-like orbitals.


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Hund’s Rule

• The lowest energy configuration for an atom is the one having the maximum number of unpaired electrons allowed by the Pauli principle in a particular set of degenerate orbitals.


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Greatest triumph of quantum mechanical model

Is its ABILITY

to account for thearrangement of elements in thePeriodic Table.


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Write the Electronic Configuration for the first 18 elements.

• Write the

--- full electronic configuration and

--- the noble gas configuration and

--- the orbital diagram.


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Figure 7.26The Orbitals Being Filled for Elements in Various Parts of the Periodic Table


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Figure 7.24

The Electron Configurations in the Type of Orbital Occupied Last for the First 18 Elements


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Figure 7.25

Electron Configurations for Potassium Through Krypton


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• Know exceptions of Cu & Cr


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Broad Periodic Table Classifications

• Representative Elements (main group): filling s and p orbitals (Na, Al, Ne, O)

• Transition Elements: filling d orbitals (Fe, Co, Ni)

• Lanthanide and Actinide Series (inner transition elements): filling 4f and 5f orbitals (Eu, Am, Es)

copyright©2000 by houghton mifflin company. all rights reserved. 1 honors chemistry chapter 5

Documents

houghton mifflin company

energy of electron

excess energy

quantity of energy

transfer of energy

h c slide

discrete energy levels

quantized energy levels