The Electronic Structure of Atoms 1 The Electronic Structure of the Atom y From Classical Physics to Quantum Theory (7.1) y The Photoelectric Effect (7.2) y Bohr’s Theory of the Hydrogen Atom (7.3) y The Dual Nature of the Electron (7.4) y Quantum Mechanics (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 waves and black body radiation y What properties are used to describe a wave? 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 energy 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 particles? 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 electrons 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 photon has a frequency ≥ the energy holding the electron in the metal (the binding energy), 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 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 Describe the atomic emission spectrum for hydrogen 7.3 Bohr’s theory of the hydrogen atom and atomic orbits Atomic Emission Spectrum of Hydrogen 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: ◦ Ground state or ground level ◦ Excited state or excited level Atomic Emission Spectrum of Hydrogen Figure 7.7, p. 222 Atomic Emission Spectrum of Hydrogen Figure 7.9, p. 223 p. 7.9, Figure 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 photons? Atomic Emission Spectrum of Hydrogen 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)? 7.4 Dual nature of the electron Figure 7.10, p. 225 Standing Waves 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 uncertainty principle? y What do quantum numbers tell us? 7.5 Quantum mechanics In 1926 Schrödinger wrote an equation that described both the particle and wave nature of the e- Wave function (Ψ) describes: 1. energy of e- with a given Ψ 2. probability of finding e- in a volume of space Schrödinger Wave Equation Schrödinger Wave Schrödinger’s equation can only be solved exactly for the hydrogen atom. Must approximate its solution for multi-electron systems. 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 quantum number, 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 energies 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 Quantum numbers (7.6) and atomic orbitals (7.7) For an atom with electrons in n = 1, 2, 3, and 4, how many atomic orbitals are totally spherical? How many have 1 node? How many have l = 2 and ml = 2? Are there any atomic orbitals with ml = 3? Atomic Orbitals – Practice How many? 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 angular momentum 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 principal quantum number n n=3 n=2 1 E = -R n H ( n 2 ) n=1 Figure 7.19, p. 233 7.8 Electron configurations Shielding Effect –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 Aufbau principle? y What is a noble gas core? y How can we use the periodic table 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 period 3 and 4 elements in the ground state, which are paramagnetic and which are diagmagnetic? Is the silver ion paramagnetic or diamagnetic? Electron ConfigurationsElectron – Practice Chapter 7 – Practice Chapter 7 – Practice.