Upload
tracy-stephens
View
235
Download
1
Tags:
Embed Size (px)
Citation preview
Chapter 2
Atomic Structure
and Periodicity
Chapter 2
Table of Contents
Return to TOC
Copyright © Cengage Learning. All rights reserved 2
2.1 Electromagnetic Radiation
2.2 The Nature of Matter
2.3 The Atomic Spectrum of Hydrogen
2.4 The Bohr Model
2.5 The Quantum Mechanical Model of the Atom
2.6 Quantum Numbers
2.7 Orbital Shapes and Energies
2.8 Electron Spin and the Pauli Principle
2.9 Polyelectronic Atoms
2.10 The History of the Periodic Table
2.11 The Aufbau Principle and the Periodic Table
2.12 Periodic Trends in Atomic Properties
2.13 The Properties of a Group: The Alkali Metals
Section 2.1
Electromagnetic Radiation
Return to TOC
Copyright © Cengage Learning. All rights reserved 3
Different Colored Fireworks
Section 2.1
Electromagnetic Radiation
Return to TOC
Properties of Light
Eletromagnetic radiation
(The way that energy travels through space)
(a) ex: Sun light, microwave. X-ray, radiant heat
(b) Wavelike behavior:
= cwavelength ( ) : m
frequency (v): s-1 (= hertz, Hz)
velocity ( c ) : m/s
Section 2.1
Electromagnetic Radiation
Return to TOC
Copyright © Cengage Learning. All rights reserved 5
Classification of Electromagnetic Radiation
Section 2.2
The Nature of Matter
Return to TOC
Copyright © Cengage Learning. All rights reserved 6
Pickle Light
Section 2.2
The Nature of Matter
Return to TOC
Classic theory Quantum theory
matter : particle energy is same mass as matter !
energy: continuous , wavelike
Section 2.2
The Nature of Matter
Return to TOC
3 paradoxes :
(1) Blackbody radiation radiation depend on Temp
Plank : energy is quantized (quanta)
only certain values allowed
(2) Photoelectric effectEinstein : light has particulate behavior
photon
(3) Atomic line spectraBohr : energy of atoms is quantized
photon emitted when electron change orbit
Section 2.2
The Nature of Matter
Return to TOC
Planck’s eqn (1900):
Observation : solid body (metal) is heat
T : 750℃ T > 1200℃metal → dull red → brighter → brilliant white lightClassical physics : atoms & molecules could emit or
absorb any arbitrary amount of E
continuePlanck proposal : energy , like matter , is discontinuous
quantum of energy & the energy E=nh n : positive integer
h : 6.6210-34 JS An atom can emit only certain amounts of energy
E = h , 2h , 3h , is not continuous but quantized
Section 2.2
The Nature of Matter
Return to TOC
Einstein & the photoelectric effect (1905
(1) Photons : particles of light
> 0 photon current
< 0 no e - ejected
classical : energy associate intensityweak blue light &
intense red light
Section 2.2
The Nature of Matter
Return to TOC
Einstein & the photoelectric effect (1905)
(2) Ephoton = h = hc/
E = mc2
a) energy is quantized
light wave
photon (particle)
mass Speed of light
energy
b) light : dual nature
Section 2.3
The Atomic Spectrum of Hydrogen
Return to TOC
Section 2.4
The Bohr Model
Return to TOC
Bohr’s postulation (for hydrogen atom)a) e moving around the nucleus in a circular orbit
Planetary model
b) only a limited number of orbits with certain E are allow orbits are quantized
c) E of electrons in orbit its distance from nucleusE = -2.178 10-18 (Z2 / n2)
d) Electrons can pass from one allowed orbit to another.Fig. 2.9
Section 2.4
The Bohr Model
Return to TOC
lower E → higher E ni < nf
E > 0 absorption spectrum
higher E → lower E ni > nf
E < 0 emission spectrum (fire works)
Niel Bohr had tied the unseen (interior of the atom)
to the seen (spectrum)
But the model is only good for one e atom:
H , He+ , Li2+
Section 2.4
The Bohr Model
Return to TOC
p. 70, Fig. 2-9
Section 2.5
The Quantum Mechanical Model of the Atom
Return to TOC
wave mechanics or quantum mechanics
(A) Louis de Broglie (1982 – 1987)
light wave
photon
How about matter ?
matters have both
wave & particle behavior
2r = nmr = n(h/2)
h = ──
m
wave properties
particle properties
Section 2.5
The Quantum Mechanical Model of the Atom
Return to TOC
(B) Schrödinger’s model of H atom & wave function
() = f (x, y, z)
(1) Ĥ = E an electron in an atom could be described by equation for wave motion
wave function ()
characterize the e as a matter wave.
Section 2.5
The Quantum Mechanical Model of the Atom
Return to TOC
Copyright © Cengage Learning. All rights reserved 18
(2) Schrödinger’s theory choose to define the E of e precisely, i.e. can only describe the probability of electron.
E is quantized : only certain are allowed& each with allowed E.
electron density = probability of finding the e =
orbitals : specific wave functions for a given e. () The matter waves for the allow E states.
orbits : Bohr’s model , was a path supposedly followed by the electron.
Section 2.5
The Quantum Mechanical Model of the Atom
Return to TOC
(C) The uncertainty principle (1927)
Heisenberg: It’s impossible to know simultaneously both the momentum & the position of a
particle at a given time with certainty.
only probability of finding an e with a given energy a given space.
(X)(P)
(X)(mν) h/4
h/2 ( h= h/2)
Section 2.5
The Quantum Mechanical Model of the Atom
Return to TOC
Copyright © Cengage Learning. All rights reserved 20
Probability Distribution for the 1s Wave Function
Section 2.5
The Quantum Mechanical Model of the Atom
Return to TOC
Copyright © Cengage Learning. All rights reserved 21
Radial Probability Distribution
Section 2.6
Quantum Numbers
Return to TOC
Quantum Numbers
3 quantum numbers are required to describe the distribution of e in atoms
(1)The principle quantum number (n)(a) n = 1, 2, 3,……..∞ (shell)
(b) related to the size & energy of the orbital.
(c) the bigger the n, the larger the orbital, the less stable the orbital.
Section 2.6
Quantum Numbers
Return to TOC
(2) The angular momentum quantum number ( l )(a) l = 0,1, 2, 3,……., n - 1 (subshells)
(b) tell the orbital shapes or types.
(c)
(3) The magnetic quantum number ( ml )a) ml = l , ﹣ - l +1 , …, 0 , … , ( l-1 ) , l
( 2l + 1 ) integral values
b) relates to the orientation of the orbital in space
l 0 1 2 3 4name of orbital s p d f g
Section 2.6
Quantum Numbers
Return to TOC
Copyright © Cengage Learning. All rights reserved 24
Quantum Numbers for the First Four Levels of Orbitals in the Hydrogen Atom
Section 2.7
Orbital Shapes and Energies
Return to TOC
Copyright © Cengage Learning. All rights reserved 25
Two Representations of the Hydrogen 1s, 2s, and 3s Orbitals
Section 2.7
Orbital Shapes and Energies
Return to TOC
Copyright © Cengage Learning. All rights reserved 26
The Boundary Surface Representations of All Three 2p Orbitals
Section 2.7
Orbital Shapes and Energies
Return to TOC
Copyright © Cengage Learning. All rights reserved 27
The Boundary Surfaces of All of the 3d Orbitals
Section 2.7
Orbital Shapes and Energies
Return to TOC
Copyright © Cengage Learning. All rights reserved 28
Representation of the 4f Orbitals in Terms of Their Boundary Surfaces
Section 2.8
Electron Spin and the Pauli Principle
Return to TOC
(1) n, l, ml : define the orbital for an electron
(2) for muti-electron atom : we need one more quantum number : electron spin ( ms )
ms = + ½ , - ½
(3) the Pauli exclusion principle
no two electrons in an atom can have the same set of four quantum numbers (n, l, ml, ms)
no atomic orbital can contain more than two electrons (opposite spins)
He : 1s2 , (n, l, ml, ms) = (1, 0, 0, ½)
(1, 0, 0, - ½)
Section 2.8
Electron Spin and the Pauli Principle
Return to TOC
(4) Paramagnetic : can be attracted by a magnetic fied
atoms contain upaired e.
Diamagnetic : e spin are paired with partners
magnetic effects cancel out.
odd / even e ?
Section 2.9
Polyelectronic Atoms
Return to TOC
(1) For polyelectronic atoms, we need electron configuration to understand electrons behavior.
(2) electron configuration : how the electrons are
distributes among the
various atomic orbitals.
(3) order of subshell E - depend on n & l
a) E ↑ with “ n + l ” value ↑
b) if same value of ( n + l ) , then lower n lower E
Section 2.9
Polyelectronic Atoms
Return to TOC
(4) Effective nuclear charge (Zeff)
a) polyelectronic atoms with two type of interactionsnucleus - electron attraction, Zeff ↑
electron - electron repulsions, Z eff ↓
b) atomic E has a value∵ ,stronger attractions, lower Ebut repulsions, higher E
Section 2.12
Periodic Trends in Atomic Properties
Return to TOC
(1) ground state electron configuration
(2) Hund’s rule :the most stable arrangement of e in subshell (p, d, f) is that with the maximum number of unpaired e.
H 1S1
1S
n
l
# of e in the orbital (subshell)
Section 2.11
The Aufbau Principle and the Periodic Table
Return to TOC
Copyright © Cengage Learning. All rights reserved 34
The Orbitals Being Filled for Elements in Various Parts of the Periodic Table
Section 2.12
Periodic Trends in Atomic Properties
Return to TOC
Copyright © Cengage Learning. All rights reserved 35
The Values of First Ionization Energy for the Elements in the First Six Periods
Section 2.12
Periodic Trends in Atomic Properties
Return to TOC
Copyright © Cengage Learning. All rights reserved 36
Atomic Radii for Selected Atoms
Section 2.13
The Properties of a Group: The Alkali Metals
Return to TOC
Copyright © Cengage Learning. All rights reserved 37
Special Names for Groups in the Periodic Table