Chapter 2 Atomic Structure and Periodicity. Chapter 2 Table of Contents Return to TOC Copyright ©...

Chapter 2

Atomic Structure

and Periodicity

Chapter 2

Table of Contents

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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

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Different Colored Fireworks

Section 2.1

Electromagnetic Radiation

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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

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Classification of Electromagnetic Radiation

Section 2.2

The Nature of Matter

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Pickle Light

Section 2.2

The Nature of Matter

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Classic theory Quantum theory

matter : particle energy is same mass as matter !

energy: continuous , wavelike

Section 2.2

The Nature of Matter

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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

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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

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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

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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

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Section 2.4

The Bohr Model

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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

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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

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p. 70, Fig. 2-9

Section 2.5

The Quantum Mechanical Model of the Atom

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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

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(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

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(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

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(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

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Probability Distribution for the 1s Wave Function

Section 2.5

The Quantum Mechanical Model of the Atom

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Radial Probability Distribution

Section 2.6

Quantum Numbers

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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

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(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

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Quantum Numbers for the First Four Levels of Orbitals in the Hydrogen Atom

Section 2.7

Orbital Shapes and Energies

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

Section 2.7

Orbital Shapes and Energies

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The Boundary Surface Representations of All Three 2p Orbitals

Section 2.7

Orbital Shapes and Energies

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The Boundary Surfaces of All of the 3d Orbitals

Section 2.7

Orbital Shapes and Energies

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

Section 2.8

Electron Spin and the Pauli Principle

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(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

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(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

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(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

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(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

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(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

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

Section 2.12

Periodic Trends in Atomic Properties

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The Values of First Ionization Energy for the Elements in the First Six Periods

Section 2.12

Periodic Trends in Atomic Properties

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Atomic Radii for Selected Atoms

Section 2.13

The Properties of a Group: The Alkali Metals

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Special Names for Groups in the Periodic Table