sign in sign up
<<
Home , Atomic orbital

The Electronic Structure of

1 The Electronic Structure of the y From Classical to Quantum Theory (7.1) y The (7.2) y Bohr’s Theory of the Atom (7.3) y The Dual of the (7.4) y (7.5) y Quantum Numbers(7.6) y Atomic Orbitals (7.7) y Electron Configurations (7.8) y The Building-Up Principle (7.9) 7.1 Light as and black body radiation y What properties are used to describe a ? y How are frequency, wavelength and speed of light related?

Figure 7.1, p. 214 7.1 Light as waves and black body radiation y What properties are used to describe a wave? y How are frequency, wavelength and speed of light related? y What is electromagnetic radiation? 7.1 Light as waves and black body radiation

Maxwell (1873), proposed that visible light consists of electromagnetic waves.

Electromagnetic radiation is the emission and transmission of in the form of Speed of light (c) in vacuum = 3.00 x 108 m/s electromagnetic

Electromagnetic radiation Electromagnetic waves. All electromagnetic radiation λ x ν = c Figure 7.2, p. 215 7.1 Light as waves and black body radiation y What is black-body radiation and how was this important in describing light? Electromagnetic radiation Electromagnetic

Figure 7.3, p. 216 7.2 The photoelectric effect – the duality of light

y What is the photoelectric effect? y How is this explained in terms of light as ? y In terms of particles, what is intensity? y In terms of particles, how does increasing the frequency of the light affect the particles? 7.2 Photoelectric Effect

If light is not composed of particles, then increasing the intensity of light (even of low frequency) should be able to eject an electron.

But it was observed that ejected are dependent on the frequency of the light and not the intensity. Above a certain frequency (called the threshold frequency), as the intensity increases, the number of electron ejected increases.

Figure 7.4, p. 218 7.2 Photoelectric Effect

When the has a frequency ≥ the energy holding the electron in the (the ), the electron is ejected. Any additional energy of the photon (above this) can be given to the electron:

Ephoton = BE + KE

KE(ejected electron) = Ephoton –BE (or W)

Figure 7.4, p. 218 7.3 Bohr’s theory of the and atomic orbits

y What is the atomic for an element? y What does the color of the emission lines tell us about the frequency or wavelength of the emitted photon? y Why is it not continuous? y Describe the atomic emission spectrum for hydrogen 7.3 Bohr’s theory of the hydrogen atom and atomic orbits Atomic Emission Spectrum of Hydrogen Atomic Emission

Figure 7.5, p. 220 7.3 Bohr’s theory of the hydrogen atom and atomic orbits

y What is the atomic emission spectrum for an element? y What does the color of the emission lines tell us about the frequency or wavelength of the emitted photon? y Why is it not continuous? y When is energy absorbed and emitted in Bohr’s model of the atom? 7.3 Bohr’s theory of the hydrogen atom and atomic orbits

y Key Definitions: ◦ or ground level ◦ or excited level Atomic Emission Spectrum of Hydrogen Atomic Emission

Figure 7.7, p. 222 Atomic Emission Spectrum of Hydrogen Figure 7.9, p. 223 7.3 Bohr’s theory of the hydrogen atom and atomic orbits

What is the change in energy for the lowest energy line in the Balmer series for atomic hydrogen?

What is the frequency and wavelength of the corresponding ? Atomic Emission Spectrum of Hydrogen Atomic Emission

Table 7.1, p. 223

7.4 Dual nature of the electron

y Why are electrons limited to these orbits? y How can electrons be modeled so this makes sense? y Does this new model make sense (based on empirical data)? Standing Waves 7.4 Dual nature oftheelectron Figure 7.10, p. 225 Figure 7.11, p. 225 7.5 Quantum mechanics

y How do we model electrons now? y What does this mean for the model of the atom? y Is it accurate to “dot” electrons in an atomic model? y What is Heisenberg’s ? y What do quantum numbers tell us? Schrödinger Wave Equation solution for multi-electron systems. for thehydrogen atom. Must approximate its Schr ( the e andwavedescribed both theparticle nature of In 1926Schr 7.5 1. of e energy 2. offindinge ö - dinger’s equation canonly besolved exactly Quantum mechanics ö dinger wrote an equationthat Ψ - ) describes: with agiven - in avolume ofspace Ψ Atomic Orbitals

What is a boundary surface diagram? Figure 7.15, p. 231 Quantum numbers (7.6) and atomic orbitals (7.7) As we consider each , we will answer: y What is the symbol for the quantum number? y What is the rule for which quantum numbers are allowed? y What does the quantum number tell us about the atomic orbital or the electrons within these atomic orbitals? y What differentiates the different of atomic orbitals (within one atom)? Atomic Orbitals

l = 0

Figure 7.16 p. 231 Atomic Orbitals

l = 1

Figure 7.17, p. 231 Atomic Orbitals

l = 2

Figure 7.18, p. 232 Atomic Orbitals – Practice orbitals ( Quantum numbers ( r hr n tmcobtl with Are there any atomicorbitals How many have How many have 1node? are totally spherical? how many atomicorbitals For anatomwithelectrons in How many? 7.7 ) l = 2and m 7.6 l = 2? n ) andatomic = 1, 2, 3, and4, m l = 3? 7.8 Electron configurations

y What are electron configurations? y How are these determined and written? 7.8 Electron configurations

Electron configuration number of electrons in the orbital or subshell 1s1 principal quantum number, n quantum number, ℓ 7.8 Electron configurations

y What are electron configurations? y How are these determined and written? y How are these modeled in an orbital diagram? 7.8 Electron configurations

Electron configuration number of electrons in the orbital or subshell 1s1 principal quantum angular momentum number, n quantum number, ℓ

Orbital diagram H 1s1 Energy of orbitals in a single electron atom

Energy only depends on n

n=3

n=2

1 E = -R n H ( n 2 )

n=1

Figure 7.19, p. 233 7.8 Electron configurations Shielding EffectShielding –Atoms Multi-Electron Figure 7.24, p. 237 Energy of orbitals in a multi-electron atom

Energy depends on n and l

n=3 l = 2

n=3 l = 1 n=3 l = 0

n=2 l = 1 n=2 l = 0

n=1 l = 0

Figure 7.19, p. 234 7.8 Electron configurations

y What are electron configurations? y How are these determined and written? y How are these modeled in an orbital diagram? y What is the Pauli exclusion principle? y What is Hund’s rule? y What is diamagnetic vs. paramagnetic? Paramagnetic Diamagnetic (attracted by a magnet) (slightly repelled by a magnet) unpaired electrons all electrons paired

2p 2p Figure 7.22, p. 236 Table 7.3, p. 242 7.9 The building-up principle

y What is the ? y What is a core? y How can we use the to write electron configurations? 7.9 The building-up principle the Periodic Table the Periodic Electron ConfigurationsElectron and

Figure 7.25, p. 243 7.9 The building-up principle

Of the ground state atoms of C, O, and Ne, which are paramagnetic and which are diamagnetic?

Of the 3 and 4 elements in the ground state, which are paramagnetic and which are diagmagnetic?

Is the silver paramagnetic or diamagnetic? Electron ConfigurationsElectron – Practice Chapter 7 – Practice Chapter 7 – Practice