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Electronic Structure of Atoms Chapter 6 BLB 12 th
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Electronic Structure of Atoms Chapter 6 BLB 12 th.

Dec 26, 2015

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Page 1: Electronic Structure of Atoms Chapter 6 BLB 12 th.

Electronic Structure of Atoms

Chapter 6 BLB 12th

Page 2: Electronic Structure of Atoms Chapter 6 BLB 12 th.

The Periodic Table: the key to electronic structure

Page 3: Electronic Structure of Atoms Chapter 6 BLB 12 th.

Expectations Work with wavelength, frequency, and energy

of electromagnetic radiation. Know the order (energy and wavelength) of

the regions in the electromagnetic spectrum. Interpret line spectra of elements (lab). Understand electronic structure.

Quantum numbers Orbitals Electron configurations

Page 4: Electronic Structure of Atoms Chapter 6 BLB 12 th.

6.1 The Wave Nature of Light Electromagnetic Radiation

A form of energy Light, heat, microwaves, radio waves Speed of light in a vacuum: c = 2.9979 x 108 m/s

Wave characteristics Wavelength (λ) in m or nm Frequency (ν) in s-1 or Hz Velocity (c) in m/s Amplitude – height of a wave Node – point where amplitude equals zero

Page 5: Electronic Structure of Atoms Chapter 6 BLB 12 th.

Electromagnetic Spectrum

Page 6: Electronic Structure of Atoms Chapter 6 BLB 12 th.

Electromagnetic Spectrum

Page 7: Electronic Structure of Atoms Chapter 6 BLB 12 th.

Relating Wavelength and Frequency

All light travels at the same velocity:

As λ ↑, ν ↓. Electromagnetic spectrum (Fig. 6.4, p. 209)

Arranged by wavelength and frequency Visible portion: 350 – 760 nm Know order of spectrum!

νλc

Page 8: Electronic Structure of Atoms Chapter 6 BLB 12 th.

Calculate the wavelength (in nm) for light with a frequency of 5.50 x 1014 s-1.

Page 9: Electronic Structure of Atoms Chapter 6 BLB 12 th.

6.2 Quantized Energy and Photons Matter (atoms) and energy (light) were thought to

be unrelated until 1900. Matter – consists of particles with mass and position Energy (waves) – massless with uncertain position

3 problems:1. Emission of light from hot objects (blackbody radiation)

2. Emission of electrons from a metal surfaces on which light shines (photoelectric effect)

3. Emission of light from electronically excited gas atoms (emission spectra)

Page 10: Electronic Structure of Atoms Chapter 6 BLB 12 th.

What do these have in common?Color and intensity of light are temperature dependent.

Different colors of light are produced by each gaseous element.

Page 11: Electronic Structure of Atoms Chapter 6 BLB 12 th.

Quantized Energy Max Planck (1900) – postulated that energy of

matter is quantized, that is, occurs only in certain discrete units of energy Equantum = hν h = 6.626 x 10-34 J·s

(Planck’s constant) E = nhν n = an integer

Quantized – restricted to certain quantities Quantum – fixed amount

Page 12: Electronic Structure of Atoms Chapter 6 BLB 12 th.

Quantized Energy Einstein (1905) – explained that electro-

magnetic radiation is quantized Assumed light consists of tiny energy packets

called photons that behave like particles!?! Ephoton = hν

Some types of light have more energy than other types.

Energy ⇐ ? ⇒ Matter

Page 13: Electronic Structure of Atoms Chapter 6 BLB 12 th.

Photoelectric Effect

Page 14: Electronic Structure of Atoms Chapter 6 BLB 12 th.

Molybdenum requires a photon with a frequency of 4.41 x 1015 s-1 to emit electrons. Calculate (a) the energy of one photon and (b) of one mole of photons.

Page 15: Electronic Structure of Atoms Chapter 6 BLB 12 th.

6.3 Line Spectra and the Bohr Model

Spectrum – radiation separated into different wavelengths Continuous – light of all

wavelengths Line – contains only

specific wavelengths

Different gases produce different line spectra.

Page 16: Electronic Structure of Atoms Chapter 6 BLB 12 th.

Continuous Spectrum

Page 17: Electronic Structure of Atoms Chapter 6 BLB 12 th.

Atomic Line Spectra

Page 18: Electronic Structure of Atoms Chapter 6 BLB 12 th.

Rydberg Equation

2217 11

m10096776.1λ

1

if nn

ni and nf are integers; ni > nf for emission

Used to calculate the wavelengths of the lines in the line spectrum of hydrogen

↑ RH

Page 19: Electronic Structure of Atoms Chapter 6 BLB 12 th.

Bohr Model of the AtomPostulates:

1. Orbits have certain radii, which correspond to certain energy levels.

2. An electron in an orbit has a specified energy.

3. Energy is only emitted and absorbed by an electron as it moves from one energy state to another.

Page 20: Electronic Structure of Atoms Chapter 6 BLB 12 th.

Energy States of the Hydrogen Atom Bohr calculated the energy of each orbit.

n is an integer (1…∞); the principal quantum number

Ground state (n = 1) – lowest energy state Excited state – higher energy state Energy values are negative, indicating stability

from the electron-nucleus attraction.

218 1

J)1018.2(n

E

Page 21: Electronic Structure of Atoms Chapter 6 BLB 12 th.

Electronic Transitions

Ultraviolet (Lyman) ni → nf = 1

Visible (Balmer) ni → nf = 2

Infrared (Paschen) ni → nf = 3

Page 22: Electronic Structure of Atoms Chapter 6 BLB 12 th.

2217 11

m10097.1λ

1

if nn

Electronic Transitions:

hc

Eand

hchE

nnE

if

λ

1

λ

11J)1018.2(

2218

Page 23: Electronic Structure of Atoms Chapter 6 BLB 12 th.

e- transition practiceCalculate the wavelength, energy, and frequency for an electron transition from n = 5 to n = 3.

Page 24: Electronic Structure of Atoms Chapter 6 BLB 12 th.

Limitations of the Bohr Model

Offered an explanation of the hydrogen atom, but failed for other atoms.

The electron does not orbit about the nucleus.

But,

1. Electrons do exist in energy levels.

2. Energy is involved in moving an electron between energy levels.

Page 25: Electronic Structure of Atoms Chapter 6 BLB 12 th.

6.4 The Wave Behavior of Matter Louis de Broglie (1923) – discovered the

relationship between a particle’s mass and wavelength Diffraction of x-rays and electrons Tiny particles like electrons have wave-like

properties!?!

Sample Exercise 6.5: λ = 1.22 x 10-10 mmv

λh

v – velocity (m/s)

m – mass (kg)

Page 26: Electronic Structure of Atoms Chapter 6 BLB 12 th.

Heisenberg Uncertainty Principle

It is impossible to know both the position and momentum of a particle at a given time.

Heisenberg may have been here!

Enter quantum mechanics, a way to deal with both the wavelike and particle-like behavior of the electron.

4(mv)x

h

Page 27: Electronic Structure of Atoms Chapter 6 BLB 12 th.

6.5 Quantum Mechanics and Atomic Orbitals

Quantum (or wave) mechanics Schrödinger solved an equation.

Treated electron as wave Solved for energy of the wave Mathematical solution gives the size and

shape of a wave function or orbital.

Page 28: Electronic Structure of Atoms Chapter 6 BLB 12 th.

Schrödinger solved an equation. Predicts the probability of finding an

electron (electron density) Node – where probability of finding an

electron is zero Each equation solution uses four variables

called quantum numbers.

6.5 Quantum Mechanics and Atomic Orbitals

Page 29: Electronic Structure of Atoms Chapter 6 BLB 12 th.

So what are these orbitals anyway?

The Periodic Table gives us the answer.

A rule first:

Each orbital can hold only one pair of electrons – with spins of +½ and −½

↿⇂

Page 30: Electronic Structure of Atoms Chapter 6 BLB 12 th.

The Periodic Table: the key to electronic structure

Page 31: Electronic Structure of Atoms Chapter 6 BLB 12 th.

The Periodic Table: the key to electronic structure

Page 32: Electronic Structure of Atoms Chapter 6 BLB 12 th.

The Periodic Table

Page 33: Electronic Structure of Atoms Chapter 6 BLB 12 th.
Page 34: Electronic Structure of Atoms Chapter 6 BLB 12 th.

6.6 Representations of Orbitals

Electron density map

Page 35: Electronic Structure of Atoms Chapter 6 BLB 12 th.

Radial Probability Plot

Page 36: Electronic Structure of Atoms Chapter 6 BLB 12 th.

The 1s, 2s, and 3s orbitals

< <

Page 37: Electronic Structure of Atoms Chapter 6 BLB 12 th.

The p orbitals

Page 38: Electronic Structure of Atoms Chapter 6 BLB 12 th.

The d orbitals

Page 39: Electronic Structure of Atoms Chapter 6 BLB 12 th.

Quantum Numbers (p. 220, 227)Quantum # Values Purpose Describes

Principal (n) 1, 2, 3, … ∞ size and energy

shell

Angular (ℓ) 0, 1, 2, 3, … (n−1)s, p, d, f

shape subshell

Magnetic (mℓ) −ℓ, … 0 …, +ℓ orientation in space

orbital

Spin magnetic (ms) +½, −½ spin of electron

electron

Each set of four quantum numbers defines individual electrons.

Page 40: Electronic Structure of Atoms Chapter 6 BLB 12 th.

Orbital energy levels in the hydrogen atom

n = 4 shell

Page 41: Electronic Structure of Atoms Chapter 6 BLB 12 th.

Quantum Numbers

Page 42: Electronic Structure of Atoms Chapter 6 BLB 12 th.

Quantum # Practice

Page 43: Electronic Structure of Atoms Chapter 6 BLB 12 th.

6.7 Many-Electron AtomsRules of Orbital Filling

1. Pauli Exclusion Principle – In a given atom, no two electrons can have the same set of four quantum numbers, i.e. only two electrons per orbital.

2. Aufbau Principle – Lowest energy orbitals are filled first.

3. Hund’s Rule – For degenerate orbitals, the lowest energy is attained with maximum number of unpaired electrons.

Page 44: Electronic Structure of Atoms Chapter 6 BLB 12 th.

Orbital energy levels in many-electron atoms

Page 45: Electronic Structure of Atoms Chapter 6 BLB 12 th.

6.8 Electron Configurations Shows the number of electrons and type

of orbitals present in an atom or ion Orbital diagram – use boxes or lines Spectroscopic notation – 1s2 2s2 2p6 etc. Core electrons – electrons in filled shells Valence electrons – electron(s) in unfilled

shells

Page 46: Electronic Structure of Atoms Chapter 6 BLB 12 th.
Page 47: Electronic Structure of Atoms Chapter 6 BLB 12 th.

6.9 Electron Configurations and The Periodic Table

Remember: The Periodic Table is the answer. Use it!

Exceptions (p. 237):

24Cr

28Cu

Page 48: Electronic Structure of Atoms Chapter 6 BLB 12 th.

Valence Electron configurations. p. 236

Page 49: Electronic Structure of Atoms Chapter 6 BLB 12 th.

Electron Configurations of Ions (pp. 262-263)

Cations: Remove electron(s) from the orbital(s) with

the highest n and highest l. For transition metals the s electrons are

removed first.

Anions: Add electron(s) to the empty or partially filled

orbitals with the lowest value of n.

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