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Ch.4 The Electronic Structure of Atoms

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Chapter 4 The Electronic Structure of Atoms ... 4.2 Deduction of Electronic Structure from Ionization Enthalpies 4.3 The Wave-mechanical Model of the Atom – PowerPoint PPT presentation

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Title: Ch.4 The Electronic Structure of Atoms


1
Chapter 4
The Electronic Structure of Atoms
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
2
The Electronic Structure of Atoms
Chapter 4 The electronic structure of atoms (SB
p.90)
Niels Bohr
Bohrs Model of H atom
3
The Electronic Structure of Atoms
Chapter 4 The electronic structure of atoms (SB
p.90)
Niels Bohr
Bohrs Model of H atom
4
The Electromagnetic Spectrum
4.1 The Electromagnetic Spectrum (SB p.91)
5
Continuous spectrum of white light
4.1 The Electromagnetic Spectrum (SB p.92)
Fig.4-5(a)
6
Line Spectrum of hydrogen
4.1 The Electromagnetic Spectrum (SB p.93)
Fig.4-5(b)
7
The Emission Spectrum of Atomic Hydrogen
4.1 The Electromagnetic Spectrum (SB p.93)
UV
Visible
IR
8
Interpretation of the Atomic Hydrogen Spectrum
4.1 The Electromagnetic Spectrum (SB p.94)
9
Interpretation of the Atomic Hydrogen Spectrum
4.1 The Electromagnetic Spectrum (SB p.94)
10
Interpretation of the Atomic Hydrogen Spectrum
4.1 The Electromagnetic Spectrum (SB p.94)
11
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?
12
4.1 The Electromagnetic Spectrum (SB p.96)
How can we know the energy levels are getting
closer and closer together?
13
4.1 The Electromagnetic Spectrum (SB p.97)
?E E1 - E2 h?
14

4.1 The Electromagnetic Spectrum (SB p.97)
Emission spectrum of hydrogen
Absorption spectrum of hydrogen
15
Production of the Absorption Spectrum
4.1 The Electromagnetic Spectrum (SB p.97)
Absorption spectrum of hydrogen
16
Convergence Limits and Ionization
4.1 The Electromagnetic Spectrum (SB p.97)
What line in the H spectrum corresponds to this
electron transition (n 8 ? n1)?
Last line in the Lyman Series
For n8 ? n1
17
The Uniqueness of Atomic Emission Spectra
4.1 The Electromagnetic Spectrum (SB p.99)
No two elements have identical atomic spectra
?atomic spectra can be used to identify unknown
elements.
18
Ionization Enthalpy
4.2 Deduction of Electronic Structure from
Ionization Enthalpies (p.100)
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)
e- ?H 2nd I.E.
19
Evidence of Shells
4.2 Deduction of Electronic Structure from
Ionization Enthalpies (p.101)
? shells
20
Evidence of Sub-shells
4.2 Deduction of Electronic Structure from
Ionization Enthalpies (p.102)
? subshells
21
Bohrs Atomic Model and its Limitations
4.3 The Wave-mechanical Model of the Atom (p.104)
Bohr considered the electron in the H atom (a
one-electron 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.
22
Bohrs Atomic Model and its Limitations
4.3 The Wave-mechanical Model of the Atom (p.104)
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.
23
Wave Nature of Electrons
4.3 The Wave-mechanical Model of the Atom (p.104)
A beam of electrons shows diffraction phenomenon
  • Electrons possess wave properties
  • (as well as particle properties).

24
Wave Nature of Electrons
4.3 The Wave-mechanical Model of the Atom (p.104)
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.
25
Quantum Numbers
4.3 The Wave-mechanical Model of the Atom (p.104)
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
The solutions of the wave functions are the
orbitals -- which are themselves equations
describing the electrons.
26
4.3 The Wave-mechanical Model of the Atom (p.105)
Principal quantum number (n) Subsidiary quantum number (l) Number of orbitals (2l1) Symbol of orbitals Maximum number of electrons held
1 0 1 1s 2
2 0 1 1 3 2s 2p 2 6
3 0 1 2 1 3 5 3s 3p 3d 5 6 10
4 0 1 2 3 1 3 5 7 4s 4p 4d 4f 2 6 10 14
8
18
32
27
4.3 The Wave-mechanical Model of the Atom (p.105)
3d
4s
3p
3s
2p
2s
Each orbital can accommodate 2 electrons with
opposite spin.
1s
28
The s Orbitals
4.4 Atomic Orbitals (p. 107)
29
The s Orbitals
4.4 Atomic Orbitals (p.107)
30
The p Orbitals
4.4 Atomic Orbitals (p.109)
31
The END
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