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Copy of Chapter 8: Electron Configuration and Chemical Periodicity

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Ryan Walker

on 5 March 2013

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Transcript of Copy of Chapter 8: Electron Configuration and Chemical Periodicity

The First Periodic Table Arranged the 65 known elements based on atomic mass
Found repeating patterns of chemical properties he termed "periodic law" Dmitri Mendeleev Electron Configuration and Chemical Periodicity KEY QUESTION:
How does the "electron configuration" of an element relate to its chemical and physical properties? ELECTRON CONFIGURATION the distribution of electrons within the orbitals of an atom The Modern Periodic Table Consists of many more elements than were known in Mendeleev's time
Arranged in order of ATOMIC NUMBER rather than atomic mass Quantum Mechanical Model of Hydrogen Quantum Numbers the same quantum numbers apply to many-electron atoms
n - size/energy level
l - shape of orbital
ml - orientation of orbital
one additional quantum number is needed to describe individual electrons within orbitals, called SPIN ms - spin of electron A beam of H-atoms passed through a magnetic field splits into two beams
each electron acts like a spinning charge creating a tiny magnetic field
Each H-atom (single electron) can have one of two values of SPIN, +1/2 or -1/2 +1/2 or -1/2 0, 1, 2, 3... 0 to (n-1) -l to +l Electron Spin spin is an intrinsic property of an electron
you can think of electrons like tops that can spin one way or the other Quantum Numbers We can now use a set of four quantum numbers to describe any single electron in an atom's ground state electron configuration For the lone electron in our friend hydrogen: n = 1
l = 0
ml = 0
ms = +1/2 by convention, the "first" electron in an orbital gets assigned +1/2 The Pauli Exclusion Principle No two electrons in the same atom can have the same set of four quantum numbers Helium is the next element on the table, and the first to have more than one electron The first electron in He has the same quantum numbers as the electron in H The second electron in He can't have the same 4 quantum numbers (by the Pauli exclusion principle) Every electron in an atom has its own unique identity, given by a unique set of quantum numbers The second He electron has a spin quantum number of -1/2 SO? Each orbital can hold a maximum of 2 electrons with opposing spins The 1s orbital of He is "filled" and the electrons have "paired spins" Coulomb's Law Electrostatic Effects In hydrogen, the energy state is dependent only on the energy level, n In atoms with multiple electrons, the energies arise not only from attractions with the nucleus, but also from repulsions between electrons There is only one ever-present attraction between the lone electron and the nucleus Energy Level Splitting hydrogen others Lithium Lithium has 3 electrons In the ground state, the first two fill the 1s orbital, the third must go into the next-lowest-energy orbital 2s is lower in energy than 2p, so the 3rd electron in Li is found in 2s Split energy levels result from 3 factors leading to 2 phenomena Factors Phenomena nuclear charge electron repulsions orbital shape shielding
penetration Nuclear Charge Increasing nuclear charge increases nucleus-electron attractions increasing stability lowering energy Electron Repulsion Additional electrons means additional repulsive forces between them counteracting nuclear attractions decreasing Zeff increasing energy Zeff the amount of charge an electron actually "feels" Orbital Shape Quantum Mechanics and The Periodic Table Quantum mechanics provides the theoretical foundation for the experimentally-derived periodic table We will build a periodic table using electron configurations

Pay special attention to the recurring pattern of electron configurations (relating to periodic law) The Aufbau Principle "to build up" We will start with H and, for every new element, add one electron to the lowest-energy orbital available

This will result in the ground-state electron configuration for each element along the way H and He The quantum numbers designating a particular element refer to the LAST electron added Two common ways to designate electrons in orbitals electron configuration orbital diagram Period 2 Hund's Rule We place one electron in EACH ORBITAL OF A SUBLEVEL before pairing any of them up Example 8.1 Period 3
Example 8.2 Link between QM and Periodic Law Orbitals are filled in order of increasing energy, leading to outer electron configurations that recur periodically

The resulting electron configurations determine chemical properties, also recurring periodically 3d: The First d-orbital 3d is heavily shielded, but 4s penetrates near the nucleus Result: 4s is slightly lower in energy than 3d In general, the ns sublevel fills before the (n-1)d sublevel Period 4 This leads to the conclusion that half-filled and filled sublevels are more stable than other configurations Electron Configurations Elements within a group exhibit similar chemical properties because they have similar outer electron configurations Categories of Electrons 1. Inner (core) electrons Any electrons included in the previous noble gas AND any completed transition series 2. Outer electrons 3. Valence electrons All electrons in the HIGHEST ENERGY LEVEL For main-group elements, count the outer electrons For transition elements, the electrons in the unfilled d sublevel [(n-1)d electrons] are counted as well Key Info in the Periodic Table For main-group elements, the group # gives the number of valence electrons The period number gives the highest n-value in the configuration Transition and Inner Transition Elements Periods 4-7 incorporate the d-block transition elements Periods 6 and 7 include the f-block inner transition elements After lanthium begins the 4f orbitals Trends All physical and chemical behavior of the elements is ultimately based on the electron configurations of their atoms Three periodic properties that we will investigate: Atomic size
Ionization energy
Electron affinity Atomic Size Two common definitions: Metallic radius Covalent radius half the distance between the nuclei of adjacent atoms in a crystal of that element half the distance between the nuclei of identical covalently bonded atoms Two influences Changes in n Changes in Zeff Radius generally increases in a group from top to bottom Radius generally decreases in a period from left to right Ionization Energy The energy required to remove one mole of electrons from 1 mole of gaseous atoms of an element Removing an electron always requires energy, so IE is always positive You can think of it as a measure of "how hard it is to remove an electron" Many atoms may lose more than one electron, so there are 1st, 2nd, 3rd IE's and so on Removing a second electron involves removing it from a positive ion, so the energy required will be greater IE Trends IE generally decreases down a group IE generally increases across a period Successive IE's For a given element, there is a point where there is a LARGE jump between successive IE's This jump appears after the valence electrons have been removed and we are reaching into the inner electrons It is MUCH harder to remove a core electron than an outer electron Electron Affinity The energy given off or absorbed with the addition of 1 mole of electrons to 1 mole of atoms You can think of it as a measure of "how much an atom wants to become an anion" Many atoms may gain additional electrons, so there are 1st, 2nd, 3rd EA's and so on The first EA is usually negative, successive EA's are positive EA Trends Factors other than Zeff and atomic size affect EA, so its trend is spotty and less reliable EA generally becomes more negative farther up and to the right IE vs EA The relationship between trends in IE and EA lead to 3 important understandings 1. Reactive nonmetals (groups 6 and 7) 2. Reactive metals (groups 1 and 2) 3. Noble gases (group 8) High IE's, very negative EA's - tend to form ANIONS Low IE's, slightly negative EA's - tend to form CATIONS High IE's, positive EA's - DO NOT REACT METALS Shiny solids, moderate to high melting points, good thermal/electrical conductors, can be formed into wires and sheets

In reactions with nonmetals, tend to give up electrons to form cations

In reactions with metals, form metallic bonds NONMETALS Dull, brittle solids or gases, low melting points, poor thermal/electrical conductivity

In reactions with metals, tend to gain electrons to form anions

In reactions with nonmetals, form covalent bonds
METALLOIDS Have properties existing somewhere between those of metals and nonmetals Predicting Ion Formation For elements at the beginning or end of a period, the of ion a particular element forms is related to the low reactivity of noble gases isoelectronic having the same electron configuration These elements tend to ionize so as to become isoelectronic with the nearest noble gas Large Main-Group Metal Ions pseudo-noble gas configuration This (n-1)d configuration is called a Stability of a filled n = 4 energy level Stability of filled 5s and 4d sublevels Transition Metal Ions Transition metals typically form more than one cation by losing all of the ns and some of the (n-1)d electrons Because the FILLED 3d orbitals are more stable, when Period 4 transition metals form ions, the 4s electrons are removed first-- "first in, first out" Ion Formation Summary Main group s- and p-block metals: Transition metals: Nonmetals:
Remove any np-electrons, then ns-electrons Remove ns electrons before (n-1)d electrons Add electrons to np sublevels to obtain noble gas configuration Magnetic Properties Electron spin creates a tiny magnetic field
For paired electrons, magnetic fields cancel out Species with unpaired electrons exhibit paramagnetism and are attracted by an external magnetic field

Species with only paired electrons exhibit diamagnetism and are very slightly repelled by an external magnetic field Ionic Radius Cations are smaller than their parent atoms Anions are larger than their parent atoms Ionic Size Trend Down a group, ionic size increases

Across a period, ionic size decreases, then increases drastically, then decreases again For isoelectronic species, ionic size decreases with increasing positive charge Example 8.5
Example 8.6 Henry Moseley Studied x-ray spectra of elements

Was able to determine atomic numbers from data and correlate that to a new ordering system
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