New Way Chemistry for Hong Kong A-Level Book 11 Chapter 4 The Electronic Structure of Atoms 4.1The Electromagnetic Spectrum 4.2 Deduction of Electronic

New Way Chemistry for Hong Kong A-Level Book 11

Chapter 4Chapter 4The Electronic Structure of AtomsThe Electronic Structure of Atoms

4.14.1 The Electromagnetic SpectrumThe Electromagnetic Spectrum

4.24.2 Deduction of Electronic Structure Deduction of Electronic Structure from Ionization Enthalpies from Ionization Enthalpies

4.34.3 The Wave-mechanical Model of The Wave-mechanical Model of the Atom the Atom

4.44.4 Atomic Orbitals Atomic Orbitals


The Electronic Structure of Atoms

Niels Bohr

Bohr’s Model of H atom

Chapter 4 The electronic structure of atoms (SB p.90)


The Electronic Structure of Atoms

Niels Bohr

Bohr’s Model of H atom

Chapter 4 The electronic structure of atoms (SB p.90)


The Electromagnetic Spectrum

c

4.1 The Electromagnetic Spectrum (SB p.91)

c


Continuous spectrum of white light

Fig.4-5(a)



Line Spectrum of hydrogen4.1 The Electromagnetic Spectrum (SB p.93)

Fig.4-5(b)


The Emission Spectrum of Atomic Hydrogen

UV Visible IR



Interpretation of the Atomic Hydrogen Spectrum



Interpretation of the Atomic Hydrogen Spectrum4.1 The Electromagnetic Spectrum (SB p.94)


Interpretation of the Atomic Hydrogen Spectrum4.1 The Electromagnetic Spectrum (SB p.94)


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




How can we know the energy levels are getting closer and closer together?



E = E1 - E2 = h

Planck ’s constant

Frequency of light emitted


Emission spectrum of hydrogen

Absorption spectrum of hydrogen

dark background(photographic plate)

bright lines

bright background(photographic plate)

dark lines



Production of the Absorption Spectrum

Absorption spectrum of hydrogen


bright background(photographic plate)

dark lines


Convergence Limits and Ionization

What line in the H spectrum corresponds to this electron transition (n= ∞ n=1)?

What line in the H spectrum corresponds to this electron transition (n= ∞ n=1)?

Last line in the Lyman SeriesLast line in the Lyman Series

For n=∞ n=1:


H (g) H+(g) + e-


The Uniqueness of Atomic Emission Spectra

No two elements have identical atomic spectraatomic spectra can be used to identify unknown elements.



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.

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.

4.2 Deduction of Electronic Structure from Ionization Enthalpies (p.100)


Evidence of Shells

shells



Evidence of Sub-shells


2,1

2,2

2,3

2,4

2,5

2,6

2,7

2,8 subshells


Bohr’s Atomic Model and its LimitationsBohr 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.

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.



Wave Nature of Electrons

A beam of electrons shows diffraction phenomenonElectrons possess wave properties

(as well as particle properties).



Wave Nature of ElectronsSchrödinger 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.



Quantum NumbersElectrons 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.

The solutions of the wave functions are the orbitals -- which are themselves equations describing the electrons.




Principal quantum number

(n)

Subsidiary quantum

number (l)

Number of orbitals (2l+1)

Symbol of orbitals

Maximum number of electrons

held1 0 1 1s 2

2 01

13

2s2p

26

3 012

135

3s3p3d

56

104 0

123

1357

4s4p4d4f

26

1014

8

18

32


Each orbital can accommodate 2 electrons with opposite spin.

1s

2s

2p

3s

3p

3d

4s



The s Orbitals

4.4 Atomic Orbitals (p. 107)


The s Orbitals4.4 Atomic Orbitals (p.107)


The p Orbitals

4.4 Atomic Orbitals (p.109)


The END

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New Way Chemistry for Hong Kong A-Level Book 11 Chapter 4 The Electronic Structure of Atoms 4.1The Electromagnetic Spectrum 4.2 Deduction of Electronic