Electronic Structure and Periodic Properties of Elements
Electronic Structure and Periodic Properties of Elements CHEM 1127Q 001
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This 10 page Class Notes was uploaded by Caitrín Hall on Thursday April 14, 2016. The Class Notes belongs to CHEM 1127Q 001 at University of Connecticut taught by Fatma Selampinar (TC), Joseph Depasquale (PI) in Spring 2016. Since its upload, it has received 11 views. For similar materials see General Chemistry in Chemistry at University of Connecticut.
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Date Created: 04/14/16
Chapter 6: Electronic Structure and Periodic Properties of Elements 6.1 Electromagnetic Energy Electromagnetic radiation – light is the visible part of a vast spectrum of electromagnetic waves; kilometers (10 m) to picometers (10 -12m) Light behaves like a wave and a particle (wave-particle duality) Waves A wave is an oscillation o periodic movement that can transport energy from one point in space to another Kinetic energy is transferred through matter while the matter remains in place Travel through a vacuum at a constant speed of 2.998 x 10 m/s = speed of light (c) Wavelength (λ) – the distance between two consecutive peaks or troughs in a wave Frequency (v) – wave cycles that pass a specified point in space in a specified amount of time; measured in hertz (Hz) or seconds inverse (s ) Amplitude – corresponds to the magnitude of the wave’s displacement; one-half the height between the peaks and troughs; related to intensity (brightness or loudness) Electromagnetic spectrum – the range of all types of electromagnetic radiation o The human eye sees only between 400 and 700 nm c = 2.998 × 10 ms −1 = λν Short wavelength means high frequency and energy Long wavelength means low frequency and energy Continuous spectrum – electromagnetic radiation given off in an unbroken series of wavelengths Blackbody – idealized perfect absorber of all incident electromagnetic radiation; emit electromagnetic radiation in continuous spectra called blackbody radiation o Led to the discovery that as temperature increases, atoms vibrate at higher frequencies with lower wavelengths; their vibrations are the source of emitted electromagnetic radiation o Planck found that by restricting the vibrational energies to discrete values for each frequency (quantized energies), he could derive an expression for blackbody radiation that correctly had the intensity dropping rapidly for the short wavelengths in the UV region E = nhv -34 h = Planck’s constant = 6.626 x 10 joule seconds (J s) o Quantization – only discrete values from a more general set of continuous values of some property are observed; hence why n = 1, 2, 3… The Photoelectric Effect Electrons can be ejected from the clean surface of a metal when light having a frequency greater than some threshold frequency was shone on it KE of ejected electrons does not depend on brightness of light, but increases with increasing frequency of the light # ejected electrons depends on brightness because greater # incoming photons greater likelihood of collision with electrons of metal Einstein realized that Planck’s concept of quantized energies could be applied to the light in the photoelectric effect if the light striking the metal is viewed as a photon rather than a wave Photons – smallest possible packet of electromagnetic radiation, a particle of light whose energy depends on its frequency Planck’s formula: E = hv OR E = (hc)/ λ Line Spectra Line spectrum – electromagnetic radiation emitted at discrete wavelengths by a specific atom (or atoms) in an excited state Light can occur as line spectra having very narrow line widths interspersed throughout the spectral regional Exciting a gas at low partial pressure using an electrical current (heating it) produces line spectra (ex: fluorescent light bulbs, neon signs) Each emission line consists of a single wavelength of light, which implies that the light emitted by a gas consists of a discrete set of energies Each element has a unique energy shell system 6.2 The Bohr Model Preceding Bohr was the planetary model of electrons orbiting atoms Bohr assumed that the electron orbiting the nucleus would not normally emit any radiation (the stationary state hypothesis), but it would emit or absorb a photon if it moved to a different orbit (energy level) The energy absorbed or emitted reflects differences in the orbital energies according to this equation: ∣ΔE∣ = ∣E −E ∣ = hν = hc/ λ f i o h = Planck’s constant o Eiand Efare the initial and final orbital energies Bohr assumed that only discreet values of angular momentum, energy, and orbit radius could occur Expression for quantized energies: En= -k/(n2) , n=1,2,3, ... o k is a constant comprising fundamental constants such as electron mass and charge and Planck’s constant OR 2 2 En= -(R H Z )/(n ) o R (Hhe Rydberg constant) = 2.179 x 10 -18joules o Z = nuclear charge of an atom (ex: +1 for H, +2 for He, +3 for Li…) Bohr’s model o Models the hydrogen atom only but does not account for electron-electron interactions in atoms with more than one electron o Introduce important features of all models An electron in its lowest energy orbit is in its ground electric state If the atom receives energy from an outside source, an electron can move to an orbit of a higher n value; excited electronic state When an electron moves from an excited state to a less excited state, the difference is the energy emitted as a photon Energies of electrons are quantized, described by quantum numbers; integer #s having only specific allowed value Electron’s energy increases with increasing distance from the nucleus Discrete energies (lines) in the spectra result from quantized energies 6.3 Development of Quantum Theory Why do electrons orbit at only fixed distances defined by a single quantum number? Why did Bohr’s model work for only hydrogen and one-electron ions? Behavior in the Macroscopic World Louis de Broglie extended wave-particle duality of light to electrons and predicted that a particle with mass m and velocity v (with linear momentum p) exhibit the behavior of a wave with wavelength λ λ = h/mv = h/p “de Broglie wavelength” o Characteristic of particles and other bodies rather than electromagnetic radiation; v = velocity NOT frequency!! Heisenberg uncertainty principle: It is fundamentally impossible to determine simultaneously and exactly both the momentum and the position of a particle; consequence of wave-particle duality The Quantum-Mechanical Model of an Atom Erwin Schrödinger thought of electrons in terms of three-dimensional stationary waves, or wavefunctions o Schrödinger equation: Hψ = E ψ 2 o The square of the magnitude of a wavefunction ∣ψ∣ describes probability of quantum particle being near a certain location in space (idea of orbital) Understanding Quantum Theory of Electrons in Atoms Atomic orbital – general region in an atom within which an electron is most probable to reside Principle quantum number – defines the location of an energy level; same concept as n in the Bohr model; greater distance from nucleus greater energy; defines the energy of an electron in hydrogen or an ion with one electron and the orbital in which discrete energy levels of electrons in multi-electron atoms/ions are located Angular momentum quantum number (l) – an integer that defines shape of orbital and takes on values of n – 1; orbitals with the same l value form a subshell o l = 0 s orbitals o l = 1 p orbitals o l = 2 d orbitals o l = 3 f orbitals Magnetic quantum number (m) – specifils orientation of an orbital in space; ranges in value from – l to + l Spin quantum number (m ) – intrsnsic electron rotation; + or – 1/2 The Pauli Exclusion Principle No two electrons in the same atom have exactly the same set of all four quantum #s Two electrons in the same orbital must have opposite spins 6.4 Electronic Structure of Atoms (Electron Configurations) Orbital Energies and Atomic Structure Electron configuration – the arrangement of electrons in the orbitals of an atom 1. The # principal quantum shell, n 2. The letter that designates the orbital type (subshell, l) 3. A superscript number that designates # electrons in that subshell The Aufbau Principle – each added electron (across a period) occupies the subshell of lowest energy available Hund’s Rule – the lowest-energy configuration for an electrons within a set of degenerate (same energy) orbitals is having max # unpaired electrons Orbital diagrams – pictorial representations of the electron configuration Electron Configurations and the Periodic Table These classifications determine which orbitals are counted in the valence shell 1. Main group (representative) elements – the last electron added enters and s or p orbital in the outermost shell; valence electrons are those with highest n; completely filled d orbitals count as core 2. Transition elements or transition metals – metallic elements in which the last electron added enters d orbital; valence electrons include ns & (n – 1) d electrons 3. Inner transition elements – metallic elements in which the last electron added occupies f orbital; valence shells consist of (n – 2)f, (n – 1)d, and ns subshells Electron Configuration of Ions For main group elements, electrons added last are the first to be removed For transition and inner transition metals, electrons in the s orbital are easier to remove than d or f electrons highest ns electrons are lost, then (n – 1)d or (n – 2)f Added electrons fill in the order predicted by the Aufbau principle 6.5 Periodic Variations in Element Properties Across a period from left to right: add a proton to the nucleus and an electron to the valence shell with each successive element Down the elements in a group, # electrons in the valence shell remains constant, but principal quantum # increases by one each time Radius of atoms and ions, ionization energies, and electron affinities vary Variation in Covalent Radius Covalent radius – ½ the distance between the nuclei of two identical atoms when joined by a covalent bond (possible because atoms within molecules retain identity) Size increases from top to bottom and decreases from left to right o Effective nuclear charge, Z eff– the pull exerted on a specific electron by the nucleus, taking into account electron-electron repulsions o Z effncreases from left to right across a period; stronger pull experience by electrons on the right side of the periodic table draws them closer to the nucleus, making the radii smaller o Shielding is determined by the probability of another electron being between the electron of interest and the nucleus, as well as by the electron-electron repulsions Zeff= Z – shielding Variation in Ionic Radii Cation is smaller than its atom because when electrons are removed from the outer valence shell, the remaining core electrons experience a greater Z eff Down the groups of the periodic table, cations of successive elements with the same charge have larger radii Anion is larger than its atom because addition of one or more electrons to the valence shell increases repulsion among electrons and decreases Z effper electron Isoelectric atoms and ions have the same electron configuration Variation in Ionization Energies The amount of energy required to remove the most loosely bound electron from a gaseous atom in its ground state is called its first ionization energy (IE ) 1 IE1= energy required to form a cation with a +1 charge Endothermic process (IE values are always positive) As atomic radius increases, ionization energy decreases (theoretically) IE1increases with increasing Zacross a period and decreases with increasing Z down a group Exceptions: o IE 1roup 2 > IE gro1p 3; IE group15 > IE group 61 o Because it is easier to remove electrons paired electrons from higher orbitals Variation in Electron Affinities Electron affinity [EA] – the energy change for the process of adding an electron to a gaseous atom to for an anion Endo/exothermic depending on element It becomes easier to add an electron across a series of atoms as Z eff increases From left to right across a period, EAs tend to become more negative From top to bottom of each group, EA is less clear
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