CHEM101 NOTES Properties of Light • Found that the wave theory could not work for some experiments e.g. the photovoltaic effect • This is because the classic EM view of light could not account for some of the observations • When then intensity changed, the energies would not vary. This did not fit the classical theory. • When the frequency of the radiation was changed, photons were ejected with more energy - classical theory states that this should only happen with a change of intensity. • focal light consists of packets of energy called photons • these photons needed to have a certain threshold of energy in order to eject photons of light Problems with the Wave Theory of Light (Photoelectric Effect) • when you heat up elementary substances (such as argon, neon, hydrogen, etc) by passing through an electrical discharge they emit only certain frequencies of light - unequally spaced on the frequency spectrum • Balmer’s equation calculated this but only for the hydrogen spectrum Emission Spectra • Each orbit has a very specific, quantised energy Bohr Model of the Hydrogen Atom • The lower the energy the more stable the element • when electrons are energised they jump between energy states/orbital shells • in order to return to ground state they will eject a photon of light • depending which energy state they are jumping between a different colour/wavelength photon is emitted. Bohr’s Postulates 1. Electrons revolve around the nucleus in circular path, which are known as "ORBITS" or "ENERGY LEVEL”. 2. Energy of an electron in one of its allowed orbits is fixed.As long as an electron remains in one of its allowed orbit, it can not absorb or radiate energy. 3. If an electron jumps from lower energy level to a higher energy level, it absorbs a definite amount of energy. 4. If an electron jumps form higher energy level to a lower energy level, it radiates a definite amount of energy. 5. Energy released or absorbed by an electron is equal to the difference of energy of two energy levels. 6. Spectrum of light emitted from an electron is a "LINE SPECTRUM”. Page "1 of "11 Problems with Bohr’s model • Bohr model could only successfully explain the hydrogen spectrum. • could not explain why the intensity of the spectra lines were NOT all equal. • With better equipment and careful observation, it was found that there were previously undiscovered spectral lines. These were named Hyperfine lines and they accompanied the other more visible lines. • The Zeeman effect: It was found that, when hydrogen gas was excited in a magnetic field, the produced emission spectrum was split. • Stationary states - Although Bohr stated that electrons were in stationary states, he could not explain why. • The Bohr model was a mixture of quantum and classical physics Uncertainty Principle • The more you know about an electron’s velocity, the less you know about its position and vice versa. • Electrons dont occupy orbits around the nucleas, more they are in electron clouds. • *** watch video: Quantum Mechanics - The Uncertainty Principle (the science channel) Quantum Mechanics • Principal Quantum Number, ’n’ (energy level on which the orbital resides) • Angular Momentum Quantum Number, ‘I’ (shape of an orbital) • ***N AND L CAN NEVER HAVE THE SAME VALUE • S orbital: value of l = 0 • P orbital: Value of l = 1 • D orbital: Value of l = 2 Electron Density Probability Functions Page "2 of "11 Electron Configuration • Key to understanding chemical properties and reactivity • Can also explain the physical properties of substances Periodic Table • The periodic table places elements of similar reactivity in the same column • We fill orbitals (with electrons) in an increasing order of energy - Aufbau Principle • 1s • 2s • 2p • Different blocks on the periodic table correspond to different types of orbitals Rules for writing electron configuration • Shorthand method of describing the distribution of all electrons in an atom • Electrons occupy the orbital with the lowest energy • No more than two electrons can occupy a single orbital (Pauli exclusion principle) • Each configuration consists of a series of terms Terms in Electron Configurations • Each term consists of 3 parts: - Number denoting the principle quantum number for the orbital - Letter denoting the type of orbital - Superscript denoting the number of electrons in these orbitals Orbital Diagrams • These are an alternative way of showing electronic configurations for atoms • Each box represents one orbital • Half arrows represent the electrons • Since the value of ms for two electrons in the same orbital have different signs, the arrows are places in an anti-parallel fashion Page "3 of "11 Condensed / Short Hand Electron Configurations • Determine which (if any) noble gas precedes the atom or ion you are interested in • Write [He], [Ne], [Ar] etc to indicate full electron configuration for that noble gas • Add terms corresponding to the remaining (valence) electrons Unexpected Electronic Configurations • Some elements do not have the expected electronic configurations • e.g. Cr is [Ar]4s13d5 instead of [Ar]4s23d4 • e.g. Cu is [Ar]4s13d10 instead of [Ar]4s^2 3d^9 • This can occur due to: - The similarity in energy between the 4s and 3d orbitals - Because half-filled and fully-filled orbital sub-shells are especially stable Electronic Configuration and Reactivity - Noble Gases • All have outer electron sub-shells that are completely full • This is a very stable arrangement • Therefore these elements are extremely stable and unreactive Why are the Alkali Metals so Reactive? • All have one electron in the outermost (highest energy) orbital • By losing that electron in a chemical reaction they can achieve a much more stable electronic configuration Valence Electrons • Are those in the outermost shell Page "4 of "11 • Atoms and ions with a complete octet of electrons int their valence shell are very stable and unreactive • Elements in the same column of the periodic table have the same number of valence electrons as well as similar chemical properties Orbitals for the Hydrogen Atom • For the one-electron hydrogen atom, all orbitals with the same principal quantum number, n, have the same energy. • That is, they are Degenerate Orbitals for Multi-Electron Systems • As the number of electrons increases, so does the repulsion between them • Therefore, in many electron atoms, orbitals with the same principle quantum number, n, are not always degenerate • For a given value of n the energies of the orbital sub-shells increase in the following order: • s < p < d Electron Spin • In the 1920’s, it was discovered that two electrons in the same orbital do not have exactly the same energy • The ‘spin’ of an electron describes its magnetic field, which affects its energy Spin Quantum Number, ms • This led to a fourth quantum number, the spin quantum number, Ms • The spin quantum number has only 2 allowed values: +1/2 and -1/2 Page "5 of "11 Pauli Exclusion Principle • No two electrons in the same atom can have exactly the same set of four quantum numbers • Therefore a single orbital can accommodate a maximum of two electrons Effective Nuclear Charge • Is the net positive charge experienced by an outer shell electron in a multi- electron atom. • Is dependent on the number of protons in the nucleus and the number of inner shell electrons. • Is the key to understanding trends in physical properties and reactivity across the Periodic Table. Trends in Zeff across the Periodic Table • As you move across any row (period) in the Periodic Table Zeff increases. • What effect would this have on properties such as the size of atoms? Size of Atoms Page "6 of "11 Periodic Trends in Atomic Radii • With each row, atomic radius tends to decrease on moving from left to right. • This is due to increasing Zeff, which draws the valence electrons closer to the nucleus. • Within each column, atomic radius tends to increase on moving from top to bottom. • This is because the valence electrons are in a higher principal quantum shell, and therefore on average further from the nucleus. Cations • Cations are smaller than their parent atoms. This is because: - One or more electrons have been removed from the outermost electron shell. - Repulsions between the remaining electrons are reduced, resulting in the remaining electrons being more strongly attracted to the nucleus. Anions • Anions are larger than their parent atoms. This is because: - Addition of electrons results in greater electron-electron repulsions, forcing the electrons to stay further apart from each other. Size of Ions in a Group •Both cations and anions increase in size as you go down a column (group). •This is because the outermost electrons belong to a higher principal quantum number shell. Isoelectronic Series • Isoelectronic species are neutral or charged atoms or molecules with the same electron configuration. • For example, Li+, Be2+ and B3+ all have the electron configuration 1s2. Page "7 of "11 • Do the ionic radii of the above ions show any trend • Ionic size decreases with increasing effective nuclear charge. • O2- and F- are also isoelectronic. • With what neutral atom are they both isoelectronic? • Do the ionic radii of the above ions show any trend • Again, ionic size decreases with increasing effective nuclear charge. Ionisation Energy • Is the amount of energy required to remove an electron from the ground state of a gaseous atom or ion. • First ionization energy is the amount of energy required to remove the first electron. • Na(g) → Na+(g) + e- • Second ionization energy is the amount of energy required to remove the second electron. • Na+(g) → Na2+(g) + e- Trends in Successive Ionisation Energy • For any atom is requires more energy to remove each successive electron • For the same reasons cations are smaller than the corresponding neutral atoms. • When all valence electrons have been removed, the ionization energy increases dramatically.