Transcript Ch.4 The Electronic Structure of Atoms

Chapter 4
The Electronic Structure of Atoms
1
4.1
The Electromagnetic Spectrum
4.2
Deduction of Electronic Structure
from Ionization Enthalpies
4.3
The Wave-mechanical Model of
the Atom
4.4
Atomic Orbitals
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Chapter 4 The electronic structure of atoms (SB p.90)
The Electronic Structure of Atoms
Niels Bohr
Bohr’s Model of H atom
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New Way Chemistry for Hong Kong A-Level Book 1
Chapter 4 The electronic structure of atoms (SB p.90)
The Electronic Structure of Atoms
Niels Bohr
Bohr’s Model of H atom
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4.1 The Electromagnetic Spectrum (SB p.91)
The Electromagnetic Spectrum


4
c
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c

4.1 The Electromagnetic Spectrum (SB p.92)
Continuous spectrum of white light
Fig.4-5(a)
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4.1 The Electromagnetic Spectrum (SB p.93)
Line Spectrum of hydrogen
Fig.4-5(b)
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4.1 The Electromagnetic Spectrum (SB p.93)
The Emission Spectrum of Atomic Hydrogen
UV
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Visible
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IR
4.1 The Electromagnetic Spectrum (SB p.94)
Interpretation of the Atomic Hydrogen Spectrum
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4.1 The Electromagnetic Spectrum (SB p.94)
Interpretation of the Atomic Hydrogen Spectrum
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4.1 The Electromagnetic Spectrum (SB p.94)
Interpretation of the Atomic Hydrogen Spectrum
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4.1 The Electromagnetic Spectrum (SB p.95)
Bohr proposed for a hydrogen atom:
1. An electron in an atom can only exist in certain
states characterized by definite energy levels
(called quantum).
2. Different orbits have different energy levels. An
orbit with higher energy is further away from the
nucleus.
3.When an electron jumps from a higher energy level
(of energy E1) to a lower energy level (of energy
E2), the energy emitted is related to the frequency
of light recorded in the emission spectrum by:
E = E1 - E2 = h
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4.1 The Electromagnetic Spectrum (SB p.96)
How can we know the
energy levels are getting
closer and closer
together?
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4.1 The Electromagnetic Spectrum (SB p.97)
E = E1 - E2 = h
Planck ’s
constant
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Frequency of light emitted
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4.1 The Electromagnetic Spectrum (SB p.97)
Emission spectrum of hydrogen
dark background
(photographic
plate)
bright background
Absorption spectrum of hydrogen (photographic plate)
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bright lines
dark lines
4.1 The Electromagnetic Spectrum (SB p.97)
Production of the Absorption Spectrum
Absorption spectrum of hydrogen
bright background
(photographic plate)
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dark lines
4.1 The Electromagnetic Spectrum (SB p.97)
Convergence Limits and Ionization
What line in the H
spectrum corresponds to
this electron transition
(n= ∞  n=1)?
Last line in the Lyman Series
For n=∞  n=1:
H (g)
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H+(g) + e-
4.1 The Electromagnetic Spectrum (SB p.99)
The Uniqueness of Atomic Emission Spectra
No two elements have identical atomic spectra
atomic spectra can be used to identify unknown elements.
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4.2 Deduction of Electronic Structure from Ionization Enthalpies (p.100)
Ionization Enthalpy
Ionization enthalpy (ionization energy) of an
atom is the energy required to remove one mole
of electrons from one mole of its gaseous atoms
to form one mole of gaseous positive ions.
The first ionization enthalpy
M(g)  M+(g) + e-
H = 1st I.E.
The second ionization enthalpy
M+(g)  M2+(g) + e18
H = 2nd I.E.
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4.2 Deduction of Electronic Structure from Ionization Enthalpies (p.101)
Evidence of Shells
 shells
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4.2 Deduction of Electronic Structure from Ionization Enthalpies (p.102)
Evidence of Sub-shells
2,8
2,5
 subshells
2,7
2,2
2,4
2,6
2,3
2,1
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4.3 The Wave-mechanical Model of the Atom (p.104)
Bohr’s Atomic Model and its Limitations
Bohr considered the electron in the H atom (a oneelectron system) moves around the nucleus in circular
orbits.
Basing on classical mechanics, Bohr calculated values
of frequencies of light emitted for electron transitions
between such ‘orbits’.
The calculated values for the
frequencies of light matched with the
data in the emission spectrum of H.
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4.3 The Wave-mechanical Model of the Atom (p.104)
Bohr’s Atomic Model and its Limitations
Bohr tried to apply similar models to atoms of other
elements (many-electron system), e.g. Na atom.
Basing on classical mechanics, Bohr calculated values
of frequencies of light emitted for electron transitions
between such ‘orbits’.
The calculated values for the frequencies
of light did NOT match with the data in
the emission spectra of the elements.
 The electron orbits in atoms may
NOT be simple circular path.
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4.3 The Wave-mechanical Model of the Atom (p.104)
Wave Nature of Electrons
A beam of electrons shows diffraction phenomenon
Electrons possess wave properties
(as well as particle properties).
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4.3 The Wave-mechanical Model of the Atom (p.104)
Wave Nature of Electrons
Schrödinger used complex differential
equations/wave fucntions to describe
the wave nature of the electrons
inside atoms (wave mechanic model).
The solutions to the differential
equations describes the orbitals of the
electrons inside the concerned atom.
An orbital is a region of space having
a high probability of finding the
electron.
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4.3 The Wave-mechanical Model of the Atom (p.104)
Quantum Numbers
The solutions of the
wave functions are
the orbitals -which are
themselves
equations
describing the
electrons.
Electrons in orbitals are specified
with a set of numbers called
Quantum Numbers:
1. Principal quantum number (n)
n = 1, 2, 3, 4, …...
2. Subsidiary quantum number (l)
l = 0, 1, 2, 3…, n-1
s p d f
3. Magnetic quantum number (m)
m = -l, …, 0, …l
4. Spin quantum number (s)
s= +½, -½
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4.3 The Wave-mechanical Model of the Atom (p.105)
Principal
quantum
number
(n)
1
Subsidiary
quantum
number (l)
2
0
1
0
1
2
0
1
2
3
3
4
26
0
Number of Symbol of Maximum
orbitals (2l+1) orbitals number of
electrons
held
1
1s
2
1
3
1
3
5
1
3
5
7
2s
2p
3s
3p
3d
4s
4p
4d
4f
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2
6 8
5
6 18
10
2
6
10 32
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4.3 The Wave-mechanical Model of the Atom (p.105)
3d
4s
3p
3s
2p
2s
1s
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Each orbital can
accommodate 2
electrons with
opposite spin.
4.4 Atomic Orbitals (p. 107)
The s Orbitals
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4.4 Atomic Orbitals (p.107)
The s Orbitals
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4.4 Atomic Orbitals (p.109)
The p Orbitals
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The END
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