Presentation is loading. Please wait.

Presentation is loading. Please wait.

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 1 Honors Chemistry Chapter 5.

Similar presentations


Presentation on theme: "Copyright©2000 by Houghton Mifflin Company. All rights reserved. 1 Honors Chemistry Chapter 5."— Presentation transcript:

1 Copyright©2000 by Houghton Mifflin Company. All rights reserved. 1 Honors Chemistry Chapter 5

2 Copyright©2000 by Houghton Mifflin Company. All rights reserved. 2 Max Planck (Early 1900s) 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.

3 Copyright©2000 by Houghton Mifflin Company. All rights reserved. 3 Plancks Constant E = change in energy, in J h = Plancks constant, J s = frequency, in s 1 = wavelength, in m Transfer of energy is quantized, and can only occur in discrete units, called quanta.

4 Copyright©2000 by Houghton Mifflin Company. All rights reserved. 4 Energy is Quantized Discrete units of h. Each small packet of energy is called a Quantum. Energy is transferred in Whole Quanta.

5 Copyright©2000 by Houghton Mifflin Company. All rights reserved. 5 Surprise! Energy seems to have particle-like properties! Before - energy always assumed to be continuous.

6 Copyright©2000 by Houghton Mifflin Company. All rights reserved. 6 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/

7 Copyright©2000 by Houghton Mifflin Company. All rights reserved. 7 Energy and Mass Also, Einstein proposed that Energy has mass E = mc 2 E = energy m = mass c = speed of light

8 Copyright©2000 by Houghton Mifflin Company. All rights reserved. 8 Dual Nature of Light Electromagnetic radiation exhibits: 1) Wave Properties 2) Particulate Properties

9 Copyright©2000 by Houghton Mifflin Company. All rights reserved. 9 Figure 7.4 Dual Nature of Light

10 Copyright©2000 by Houghton Mifflin Company. All rights reserved. 10 Light thought to be purely wavelike was found to have particulate properties. Matter thought to be purely particulate Does it exhibit wave properties?

11 Copyright©2000 by Houghton Mifflin Company. All rights reserved. 11 Wavelength and Mass de Broglies Equation (1923): Allows one to calculate an apparent wavelength for a particle.

12 Copyright©2000 by Houghton Mifflin Company. All rights reserved. 12 Continuous spectrum: Contains all the wavelengths of light. Line (discrete) spectrum: Contains only some of the wavelengths of light.

13 Copyright©2000 by Houghton Mifflin Company. All rights reserved. 13 Figure 7.6 A Continuous Spectrum (a) and A Hydrogen Line Spectrum (b)

14 Copyright©2000 by Houghton Mifflin Company. All rights reserved. 14 Sample of H 2 gas (HH) Introduce a high energy spark H 2 molecules absorb energy Some of the HH 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

15 Copyright©2000 by Houghton Mifflin Company. All rights reserved. 15 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

16 Copyright©2000 by Houghton Mifflin Company. All rights reserved. 16 Figure 7.7 A Change between Two Discrete Energy Levels

17 Copyright©2000 by Houghton Mifflin Company. All rights reserved. 17 If any energy were allowed then we would see a Continuous Spectrum (a) and When only certain energies are possible we see only a discrete Line Spectrum (b)

18 Copyright©2000 by Houghton Mifflin Company. All rights reserved. 18 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.

19 Copyright©2000 by Houghton Mifflin Company. All rights reserved. 19 Figure 7.8 Electronic Transitions in the Bohr Model for the Hydrogen Atom He calculated the radii for the allowed circular orbits. Only certain electron energies allowed. Energy levels consistent with the Hydrogen line-emission spectrum.

20 Copyright©2000 by Houghton Mifflin Company. All rights reserved. 20 The Bohr Model Ground State: The lowest possible energy state for an atom (n = 1).

21 Copyright©2000 by Houghton Mifflin Company. All rights reserved. 21 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.

22 Copyright©2000 by Houghton Mifflin Company. All rights reserved. 22 Bohrs Model Appeared promising. Calculations worked well for hydrogen. Didnt work when applied to other atoms. Something fundamentally incorrect. Important for its introduction of the concept of Quantization of energy in atoms.

23 Copyright©2000 by Houghton Mifflin Company. All rights reserved. 23 Quantum Mechanical Model of the Atom Totally different approach was needed. Three physicists: Heisenberg, de Broglie, & Schrodinger. Emphasizes the wave properties of an electron.

24 Copyright©2000 by Houghton Mifflin Company. All rights reserved. 24 For the Electron in a Hydrogen Atom Electron bound to the nucleus. Similar situation of only certain allowable Electron Waves. Modeled by Schrödinger

25 Copyright©2000 by Houghton Mifflin Company. All rights reserved. 25 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.

26 Copyright©2000 by Houghton Mifflin Company. All rights reserved. 26 Section 7.6 Quantum Numbers (QN) Four different quantum numbers. Three (n, l, m l ) specify the wave function that gives the probability of finding the electron at various pts. in space. Fourth (m s ) specifies the electronspin.

27 Copyright©2000 by Houghton Mifflin Company. All rights reserved. 27 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

28 Copyright©2000 by Houghton Mifflin Company. All rights reserved Angular Momentum Quantum Number ( l = 0 to n 1) - relates to shape of the orbital. - Subshell l = 0 s subshell l = 1p subshell l = 2 d subshell l = 3f subshell

29 Copyright©2000 by Houghton Mifflin Company. All rights reserved Magnetic Quantum Number (m l = l to l ) - relates to orientation of the orbital in space relative to other orbitals.

30 Copyright©2000 by Houghton Mifflin Company. All rights reserved. 30 SHAPES Two Representations of the Hydrogen 1s, 2s, and 3s Orbitals Node – Area of Zero Probability.

31 Copyright©2000 by Houghton Mifflin Company. All rights reserved. 31 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.

32 Copyright©2000 by Houghton Mifflin Company. All rights reserved. 32 Figure 7.16 The Boundary Surfaces of All of the 3d Orbitals No 1d or 2d orbitals. Five 3d orbital

33 Copyright©2000 by Houghton Mifflin Company. All rights reserved. 33 Figure 7.17 Representation of the 4f Orbitals in Terms of Their Boundary Surfaces

34 Copyright©2000 by Houghton Mifflin Company. All rights reserved Electron Spin Quantum Number (m s = + 1 / 2, 1 / 2 ) - relates to the spin states of the electrons.

35 Copyright©2000 by Houghton Mifflin Company. All rights reserved. 35 Electron Spin & Pauli Exclusion Principle In a given atom, no two electrons can have the same set of four quantum numbers (n, l, m l, m s ). Therefore, an orbital can hold only two electrons, and they must have opposite spins.

36 Copyright©2000 by Houghton Mifflin Company. All rights reserved. 36 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.

37 Copyright©2000 by Houghton Mifflin Company. All rights reserved. 37 Hunds 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.

38 Copyright©2000 by Houghton Mifflin Company. All rights reserved. 38 Greatest triumph of quantum mechanical model Is its ABILITY to account for the arrangement of elements in the Periodic Table.

39 Copyright©2000 by Houghton Mifflin Company. All rights reserved. 39 Write the Electronic Configuration for the first 18 elements. Write the --- full electronic configuration and --- the noble gas configuration and --- the orbital diagram.

40 Copyright©2000 by Houghton Mifflin Company. All rights reserved. 40 Figure 7.26 The Orbitals Being Filled for Elements in Various Parts of the Periodic Table

41 Copyright©2000 by Houghton Mifflin Company. All rights reserved. 41 Figure 7.24 The Electron Configurations in the Type of Orbital Occupied Last for the First 18 Elements

42 Copyright©2000 by Houghton Mifflin Company. All rights reserved. 42 Figure 7.25 Electron Configurations for Potassium Through Krypton

43 Copyright©2000 by Houghton Mifflin Company. All rights reserved. 43 Know exceptions of Cu & Cr

44 Copyright©2000 by Houghton Mifflin Company. All rights reserved. 44 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)


Download ppt "Copyright©2000 by Houghton Mifflin Company. All rights reserved. 1 Honors Chemistry Chapter 5."

Similar presentations


Ads by Google