Chapter 8 (Essentials of General Chemistry, 2 nd Edition) (Ebbing and Gammon)

Transcription

1 Chapter 8 (Essentials of General Chemistry, 2 nd Edition) (Ebbing and Gammon) Electron Configuration and Periodicity Electron Spin spin quantum number (m s ) -describes the spin orientation of an electron electron spin - beam of hydrogn atoms is split into two since each electron in each atom acts as tiny magnet with only two possible orientations - electrons acts as a ball of spinning charge and like a circulating electric charge, would create a magnetic field - resulting directions of spin magnetism correspond to m s = +1/2 and m s = 1/2 Karen Hattenhauer (Fall 2007) 2 Electron Spin Electron spin magnetism Spin quantum numbers m s +1/2 and 1/2 Karen Hattenhauer (Fall 2007) 3 1

2 Electron Configurations: Orbital Diagrams electron configuration - a particular distribution of electrons among the available subshells - notation for configuration lists the subshell symbols, one after the other, with a subscript giving the number of electrons in that subshell eg. boron 1s 2 2s 2 2p 1 orbital diagram - a diagram to show how the orbitals of a subshell are occupied by electrons 1s 2s 2p Karen Hattenhauer (Fall 2007) 4 Pauli Exclusion Principle - no two electrons in an atom can have the same four quantum numbers restated: an orbital can hold at most two electrons, and then only if the electrons have opposite spins Note: each subshell holds a maximum of twice as many electrons as the number of orbitals in the subshell Karen Hattenhauer (Fall 2007) 5 The Pauli Exclusion Principle The maximum number of electrons and their orbital diagrams are: Sub shell s (l = 0) p (l = 1) d (l =2) f (l =3) Number of Orbitals Maximum Number of Electrons Karen Hattenhauer (Fall 2007) 6 2

3 Building-Up Principle (Aufbau Principle) ground state - the configuration with the lowest energy level of the atom - chemical properties of an atom are related primarily to the electron configuration of its ground state excited state - associated with energy levels other than the lowest Aufbau Principle - a scheme used to reproduce the electron configurations of atoms by successfully filling subshells with electrons Karen Hattenhauer (Fall 2007) 7 - following this principle, electron configuration of an atom may be obtained by successively filling subshells in the following order: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f - this order can be easily obtained by using the periodic table as a template - filling orbitals of lowest energy first, usually gives lowest total energy (ground state) of the atom - recall: energy of orbital depends only on quantum numbers n and l - orbitals with same n and l but different ml (different orbitals of same subshell) have same energy Karen Hattenhauer (Fall 2007) 8 - energy depends primarily on n, increasing with its value - energies of orbitals with same n increase with the l quantum number - that is, 3p orbital has slightly greater energy than 3s orbital - exception: when subshells have nearly same energy, building- up order is not strictly determined by order of energies - ground-state configurations are determined by total energies of atoms which depend not only on energies of subshells but also on energies of interaction among different subshells - note that for elements with Z=21 or greater, the energy of the 3d subshell is lower than the energy of the 4s subshell (opposite to the building-up order) Karen Hattenhauer (Fall 2007) 9 3

4 Order for Filling Atomic Subshells 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 6f Karen Hattenhauer (Fall 2007) 10 Orbital Energy Levels in Multi-electron Systems Energy 4s 3s 3p 2p 3d 2s 1s Karen Hattenhauer (Fall 2007) 11 Electron Configuration and the Periodic Table noble gases - relatively unreactive - configurations in which p subshell has just filled (neon, argon, krypton; also He with filled 1s subshell) alkaline earth metals (Group IIA) - noble-gas core plus two out electrons with an ns 2 configuration Beryllium Magnesium Calcium 1s 2 2s 2 or [He]2s 2 1s 2 2s 2 2p 6 3s 2 or [He]3s 2 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 or [He]4s 2 noble-gas core - an inner-shell configuration corresponding to one of the noble gases Karen Hattenhauer (Fall 2007) 12 4

5 Group IIA Boron 1s 2 2s 2 2p 1 or [He]2s 2 2p 1 Aluminum 1s 2 2s 2 2p 6 3s 2 3p 1 or [He]3s 2 3p 1 Gallium 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 1 or [He]3d 10 4s 2 4p 1 - boron and aluminum noble-gas core plus three electrons with configuration ns 2 np 1 - gallium pseudo-noble-gas core plus ns 2 np 1 pseudo-noble-gas core - noble-gas core together with (n-1)d 10 electrons - these electrons are usually not involved in chemical reactions Karen Hattenhauer (Fall 2007) 13 valence electron - an electron in an atom outside the noble-gas or pseudo-noble gas core - primarily involved in chemical reactions - similarities among configurations of valence electrons (valence-shell configurations) account for similarities of the chemical properties among groups of elements - figure 8.5 (p.231) note similarity in electron configuration within any group (column) Karen Hattenhauer (Fall 2007) 14 Valence-shell Configurations Figure 8.5: Periodic Table Karen Hattenhauer (Fall 2007) 15 5

6 main-group (or representative) elements - have valence-shell configurations nsansb as the outer s or p subshell is being filled d-block transition elements (transition elements) - a d subshell is being filled f-block transition elements (inner-transition elements) - an f subshell is being filled Exceptions to the Building-up Principle Element Predicted Experimental Cr [Ar]3d 4 4s 2 [Ar]3d 5 4s 1 Cu [Ar]3d 9 4s 2 [Ar]3d 10 4s 1 Karen Hattenhauer (Fall 2007) 16 Electron Configuration Using the Periodic Table Configuration of the outer (valence-shell) electrons - for main-group elements is ns a np b - n = principal quantum number of outer shell and period for the element - total valence electrons (a+b) comes from group number Karen Hattenhauer (Fall 2007) 17 Orbital Diagrams and Hund s Rule Consider possible arrangements of electrons in the ground-state configuration of C (1s 2 2s 2 2p 2 ): Diagram 1: 1s 2s 2p Diagram 2: Diagram 3: - these diagrams represent different states of the carbon atom Karen Hattenhauer (Fall 2007) 18 6

7 Hund s Rule Friedrich Hund (1927) - the lowest energy arrangement of electrons in a subshell is obtained by putting electrons into separate orbitals of the subshell with the same spin before pairing electrons Magnetic Properties of Atoms - the magnetic properties of an atom are best observed by subjecting it to the field of a strong magnet paramagnetic substance - a substance that is weakly attracted by a magnetic field, and this attraction is generally the result of unpaired electrons diamagnetic substance - a substance that is not attracted by a magnetic field or is very slightly repelled by such a field - means that the substance has only paired electrons Karen Hattenhauer (Fall 2007) 19 Some Periodic Properties periodic law - states that when the elements are arranged by atomic number, their physical and chemical properties vary periodically Physical properties: a.) atomic radius b.) ionization energy c.) electron affinity Karen Hattenhauer (Fall 2007) 20 Atomic Radii Atomic radius (or size of an atom) - measured as covalent radii obtained from measurements of distances between the nuclei of atoms in the chemical bonds of molecular substances Karen Hattenhauer (Fall 2007) 21 7

8 Observed variations in Atomic Size 1.) decrease in atomic radii from left to right across the periodic table - b/c as atomic number increases, the effective nuclear charge also increases Karen Hattenhauer (Fall 2007) 22 effective nuclear charge - the positive charge that an electron experiences from the nucleus, equal to the nuclear charge but reduced by any shielding or screening from any intervening electron distribution Consider Li: - nuclear charge is +3 but effective nuclear charge is +1 Karen Hattenhauer (Fall 2007) 23 2.) atomic radii increase going down a group of the periodic table Atomic Radii of the Elements (pm) Karen Hattenhauer (Fall 2007) 24 8

9 Ionization Energy first ionization energy (first ionization potential) - the minimum energy needed to remove the highestenergy (that is, the outermost) electron from the neutral atom in the gaseous state Consider Li: Li (1s 2 2s 1 ) Li + (1s 2 ) + e - I.E. = 520 kj/mol General trends: 1.) increasing ionization energy with atomic number in a given period 2.) decreasing ionization energy down any column of main-group elements Karen Hattenhauer (Fall 2007) 25 - electrons of an atom can removed successively with ionization energies increasing as more electrons are removed Karen Hattenhauer (Fall 2007) 26 Electron Affinity electron affinity - the energy change for the process of adding an electron to a neutral atom in the gaseous state to form a negative ion Cl ([Ne]3s 2 3p 5 ) + e - Cl - ([Ne]3s 2 3p 6 ) E.A.= -349kJ/mol - if the negative ion is stable, the energy change for its formation is a negative number (energy is released) - large negative numbers indicate a very stable negative ion is formed (as is the case with Cl - ) - small negative number indicate that a less stable ion is formed Karen Hattenhauer (Fall 2007) 27 9

10 General trend: - broadly speaking, more negative electron affinities from left to right in any period Karen Hattenhauer (Fall 2007) 28 Periodicity in Main-Group Elements - chemical and physical properties of main-group elements clearly display periodic character - move left to right in any row, metallic character of the elements decreases - proceed down a column, elements tend to increase in metallic character - low ionization energy tend to be metallic - high ionization energy tend to be nonmetallic Karen Hattenhauer (Fall 2007) 29 - basic-acidic behavior of oxides of elements is good indicator of metallic-nonmetallic character basic oxide - an oxide that reacts with acids - metal oxides are basic acidic oxide - an oxide that reacts with bases - nonmetal oxides are acidic amphoteric oxides - an oxide that has both basic and acidic properties Karen Hattenhauer (Fall 2007) 30 10

11 Hydrogen (1s 1 ) - electron configuration places is in Group IA - properties different and better to consider it belonging to its own group Group IA Elements; Alkali Metals (ns 1 ) - reactivities increase as you move down a group - form basic oxides of the formula R 2 O Group IIA Elements; Alkaline Earth Metals (ns 2 ) - reactivities increase as you move down a group; but in general are less reactive than the alkali metals - form basic oxides of the formula RO Karen Hattenhauer (Fall 2007) 31 Group IIIA Elements (ns 2 np 1 ) - shows significant increase in metallic character down the group - oxides have general formula R 2 O 3 (include acidic oxide and amphoteric oxides) Note: indicative of increase in metallic character Group IVA Elements (ns 2 np 2 ) - distinct change in metallic character down the group - oxides have general formula RO 2 (includes acid and amphoteric oxides) Karen Hattenhauer (Fall 2007) 32 Group VA Elements (ns 2 np 3 ) - distinct change in metallic character down the group - oxides have general empirical formulas R 2 O 3 and R 2 O 5 with molecular formulas of R 4 O 6 and R 4 O 10 (includes acid, amphoteric and basic oxides) Group VIA Elements; the Chalcagens (ns 2 np 4 ) - distinct change in metallic character down the group - oxides have general formula RO 2 and RO 3 (includes acid and amphoteric oxides) Karen Hattenhauer (Fall 2007) 33 11

12 Group VIIA Elements; the Halogens (ns 2 np 5 ) - reactive nonmetals with general formula, X 2 - each halogen forms several compounds with oxygen; generally unstable, acidic oxides Group VIIIA Elements; the Noble Gases (ns 2 np 6 ) - exist as gases consisting of uncombined atoms - thought to be inert until 1960s - now, compounds of xenon, krypton and radon are known (although still relatively unreactive) Karen Hattenhauer (Fall 2007) 34 12

13 This document was created with Win2PDF available at The unregistered version of Win2PDF is for evaluation or non-commercial use only. This page will not be added after purchasing Win2PDF.