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Chapter 8

CHM120-Chapter 8 Notes.docx

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Department
Chemistry
Course
CHM120H5
Professor
Judith C Poe
Semester
Winter

Description
Chapter 8: Electron Configuration & Chemical Periodicity The Original Periodic Table Mendleev arranged the 65 elements known at the time into a table and summarized their behavior in the periodic law: when arranged by atomic mass, the elements exhibit a periodic recurrence of similar properties. The Modern Periodic Table Arranges the table by atomic number (number of protons) not by atomic mass. 8.1: Characteristics of Many-Electron Atoms A. The Electron-Spin Quantum Number  The three quantum numbers, n, l and m describe the size (energy), shape and orientation respectively.  An additional quantum number describes the spin, which is a property of an electron not the orbital.  When a beam of atoms that have one or more lone pairs of electrons passes through a magnetic field, it splits into two beams. Each electron behaves like a spinning charge and generates a tiny magnetic field, which can have one of two values of spin.  The two electrons have opposing directions, so half of the electrons are attracted by the large external magnetic field while the other half are repelled by it.  The spin quantum number (m ) has two possible values: +½ or –½.  Each electron in an atom is described completely by a set of four quantum numbers: the first three describe its orbital, and the fourth describes its spin. B. The Exclusion Principle  Pauli Exclusion Principle states that: no two electrons in the same atom can have the same four quantum numbers.  The major consequence is that an atomic orbital can hold a maximum of two electrons and they must have opposing spins. C. Electrostatic Effects and Energy-Level Splitting  Electrostatic Effects: attraction of opposite charges and repulsion of like charges  They play a major role in determining the energy states of many-electron atoms.  The Hydrogen atom is only effected by the n value whereas other elements are also affected by electron-electron repulsion.  These additional interactions give rise to the splitting of energy levels into sublevels of differing energies: the energy of an orbital in a many-electron atom depends mostly on its n value (size) and to a lesser extent on its l value (shape).  This energy difference arises from three factors: 1. Nuclear Attraction 2. Electron Repulsions 3. Orbital Shape  These three factors lead to shielding and penetration that occur in all elements but H.  The Effect of Nuclear Charge (Z) on Sublevel Energy. o Higher charges interact more strongly than lower charges (according to Coulomb’s Law) o Thus, a higher nuclear charge increase nucleus-electron attraction and thus lowers sublevel energy. o Least Stable = Highest Energy sublevel, least energy required to remove it.  Sheilding: The Effect of Electron Repulsion on Sublevel Energy o each electron feels not only the attraction to the nucleus but also repulsion from other electrons o repulsions counteract nuclear attraction, making each electron easier to remove by helping to push it away o each electron “shields” the other electrons from the nuclear charge somewhat o Shielding (also called screening) reduces the full nuclear charge to an effective nuclear charge eff), which is the nuclear charge an electron actually experiences, and this lower nuclear charge makes the electron easier to remove.  Shielding by other electrons in a given energy level: electrons in the same energy level shield each other somewhat.  ex. He and He : both have a 2+ nuclear charge (two protons) but He has 2e in the 1s + - sublevel and He only has one, thus it takes less than ½ as much energy to remove an e from He -  Shielding by Electrons in in inner energy levels: b/c inner e spend all their time b/w the outer electrons and nucleus, they cause a much lower Z effhan do e in the same level. This makes it easier to remove e from outer levels.  Penetration: The Effect of Orbital Shape on Sublevel Energy o Penetration has two effects:  it increases the nuclear attraction for a 2s electron over that for a 2p electrons  it decreases the shielding of a 2s electron by the 1s electron  Splitting of Levels into Sublevels o penetration and the resulting effects on shielding cause an energy level to split into sublevels of differing energy. o the lower the l value of a sublevel, the more its electrons penetrate, and so the greater their attraction to the nucleus o thus, for a given n value, a lower l value indicates a more stable (lower energy) sublevel  Order of sublevel energies: s < p < d < f 8.2: The Quantum-Mechanical Model and The Periodic Table A. Building the Periodic Table  the ground state electron configuration = the lowest-energy distribution of electrons in the sublevels of its atoms.  Aufbau Principle: electrons are always placed in the lowest energy sublevel available.  The Pauli Exclusion Principle states that each orbital may contain a maximum of 2 electrons, which must have opposite spins  There are two common ways to indicate the distribution of electrons: o Electron configuration is indicated by a shorthand notation which consists of the principle energy level (n value), the letter designation of the sublevel (l value), and the number of electrons in the sublevel o Orbital Diagram: Orbital diagrams make use of a box, circle, or line for each orbital in the energy level. An arrow is used to represent an electron and its spin.  Hund’s Rule: specifies that when orbitals of equal energy are available, the lowest energy electron configuration has the maximum number of unpaired electrons with parallel spins.  Elements in the same group have similar outer electron configurations and similar patterns of reactivity  Partial orbital diagrams: shows only the sublevels being filled  Condensed Electron Configuration: A condensed electron configuration has the element symbol of the previous 2 1 noble gas in square brackets. Al has the condensed configuration [Ne]3s 3p  Elements in the same group of the periodic table have the same outer electron configuration.  Elements in the same group of the periodic table exhibit similar chemical behavior.  Similar outer electron configurations correlate with similar chemical behavior.  Sublevels are filled in order of increasing energy, which leads to chemical properties that recur periodically B. Building Up the Transition Series:  Effects of Shielding and Penetration on sublevel energy: the 3d sublevel is filled in period 4, but the 4s level is filled first, this switch in filling order is due to shielding and penetration effects. o Based on the 3d radial probability distribution, a 3d electron spends most of the time outside the filled inner n = 1, and n=2 levels, so it is shielded very effectively from nuclear charge. o The outermost 4s electron penetrates close to the nucleus part of the time, so it is subject to greater attraction. o Thus, the 4s orbital is slightly lower in energy than the 3d and fills first o In any period the ns sublevel fills before the (n-1)d sublevel  Filling the 4s and 3d sublevels: o filling the 3d orbitals proceeds at one electron a time with the exception of Cr and Cu  Stability of Half-Filled and filled sublevels o half-filled and full sublevels are the most stable 2 4 1 5 o Cr – instead of [Ar]4s 3d as expected has [Ar]4s 3d , an electron from the s sublevel moves to the d sublevel to increase stability C. Categories of Electrons  Inner (core) electrons are those an atom has in common with the pervious noble gas and any completed transition series. The fill all the lower energy levels of an atom  Outer electrons are those in the highest energy level (highest n value), they spend most of their time farthest from the nucleus.  Valence Electrons: are those involved in forming compounds o For main group elements, the valence electrons are the outer electrons o For transition elements, the valence electrons include the outer electrons and any (n-1)d electrons. Group and Period Numbers  among the main group elements the A number equals the number of outer electrons  the period number is the n value of the highest energy level 2 2  for any energy level, n = the number of orbitals, and 2n = maximum number of electrons D. Intervening Series: Transition and Inner Transition Metals  Transition Series: Periods 4,5,6 and 7 have 3d,4d,5d and 6d orbitals respectively. the order of filling is ns, (n-1)d, np  Inner Transition Series: f orbitals, the f orbitals have l=3, so the possibll m values are [-3,-2,-1,0,+1,+2,+3] o the period 6 inner transition series are called the lanthanides o the period 7 inner transition series are called the actinides o the filling order is: ns, (n-1)d, (n-2)f , np o 8.3: Trends in Three Atomic Properties A. Trends in Atomic Size Atomic size varies from substance to substance 1. Metallic Radius: it is one-half of the shortest distance b/w nuclei of adjacent, individual atoms in a crystal of the element (ex. Aluminum) 2. Covalent Radius: used for elements occurring as molecules, it is one-half of the shortest distance b/w nuclei of bonded atoms (ex. Chlorine) Main Group Elements Among the main group(1-8) elements atomic size varies within both groups and periods as a result of two opposing influences. 1. Changes in n: as the principal quantum number ,n, increases, the probability that outer electrons spend most of their time farther from the nucleus also increases, thus atomic size increases (down the group). As we move down a group each member has one more level of inner electrons that shields the outer electrons, thus atomic radius increases down a group. 2. Changes in Zeff : As the effective nuclear charge increase, outer electrons are pulled closer to the nucleus; thus atomic size decreases (across a period). Across a period, electrons are added to the same outer level, so the shielding by inner electrons doesn’t change. Thus, atomic radius decreases across a group Transition Elements 1. Down a transition group n increases but shielding by an additional level of inner electrons results in only a small change in size from period 4 to 5 and none from 5 to 6. 2. Across a transition period, atomic size shrinks through the first two or three elements because of the incr
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