Topic E: Atomic Structure
- Page ID
- 170001
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WHAT YOU SHOULD BE ABLE TO DO WHEN YOU HAVE FINISHED THIS TOPIC:
Electromagnetic radiation
- Interconvert frequency and wavelength for electromagnetic radiation, and relate each of these to photon energy in J (per photon) and in kJ/mol. (\(λ\nu = c\) and \(E_{photon}= h\nu\))
Electron energy levels and atomic emission spectra
- Recognize that electrons can only exist in certain allowed energy states.
- Recognize that in light emission and absorption, the photon energy equals the energy difference between two electron energy states. ( Ephoton = |ΔEelectron| )
- Understand the measurement and meaning of ionization energy.
Special case: Bohr model (electronic structure of a one-electon atom or ion)
- Use the Bohr energy equation to calculate the energies of the quantum levels in a hydrogen atom and in one-electron ions.
- Calculate the energy, wavelength or frequency of light that corresponds to any electron transition in a one-electron atom or ion (using the Rydberg equation or the Bohr energy equation.)
General case: Quantum Mechanical model (modern electronic structure of an atom)
- Understand that an energy state for a one-electron atom can be described by a wave function Ψ (atomic orbital); and describe the general shape of s, p and d orbitals.
- Understand the relationship between atomic wave functions Ψ (atomic orbitals) and the quantum numbers \(n\), \(\ell\), and \(m_\ell\).
- Know and apply the rules that govern the allowed values of the quantum numbers \(n\), \(\ell\), and \(m_\ell\).
- Understand the physical meaning of \(Ψ^2\) and \(Ψ^2ΔV\).
- Understand that one-electron wave functions (atomic orbitals) can be used to “build-up” a total wave function describing a multi-electron atom (i.e. electron configurations.)
- Understand how orbital energy for one-electron atoms depends only on quantum number n and how orbital energy for multi-electron atoms depends on the quantum numbers n, l, and ml.
- Understand how the quantum number ms describes electron spin, and the significance of the Pauli Principle and Hund’s Rule in multi-electron atoms.
- Predict the electron configurations for all of the elements in the periodic table, and identify a configuration as diamagnetic or paramagnetic.
- Relate the electron configurations of the elements to the following properties, with a focus on the periodic nature of these properties: Ionization energy, Atomic radius, Ionic radius, Electronegativity, and Electron affinity.
- 7: Electronic Structure of Atoms
- In this chapter, we describe how electrons are arranged in atoms and how the spatial arrangements of electrons are related to their energies. We also explain how knowing the arrangement of electrons in an atom enables chemists to predict and explain the chemistry of an element. As you study the material presented in this chapter, you will discover how the shape of the periodic table reflects the electronic arrangements of elements.
- 8: Periodic Properties of the Elements
- In this chapter, we explore the relationship between the electron configurations of the elements, as reflected in their arrangement in the periodic table, and their physical and chemical properties. In particular, we focus on the similarities between elements in the same column and on the trends in properties that are observed across horizontal rows or down vertical columns.