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Chapter 8 Electron Configurations, Atomic Properties, and the Periodic Table General Chemistry: An Integrated Approach Hill, Petrucci, 4 th Edition Mark.

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Presentation on theme: "Chapter 8 Electron Configurations, Atomic Properties, and the Periodic Table General Chemistry: An Integrated Approach Hill, Petrucci, 4 th Edition Mark."— Presentation transcript:

1 Chapter 8 Electron Configurations, Atomic Properties, and the Periodic Table General Chemistry: An Integrated Approach Hill, Petrucci, 4 th Edition Mark P. Heitz State University of New York at Brockport © 2005, Prentice Hall, Inc.

2 Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table 2 Multielectron Atoms Electrons are attracted to the nucleus while simultaneously repelling one another EOS In the hydrogen atom, all subshells of a principal shell are at the same energy level recall E n = –B/n 2

3 Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table 3 Multielectron Atoms The increasing energy order of subshells is generally: s < p < d < f EOS In a multielectron atom the various subshells of a principal shell are at different energy levels, but all orbitals within a subshell are at the same energy level

4 Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table 4 Multielectron Atoms In higher numbered principal shells of a multielectron atom, some subshells of different principal shells have nearly identical energies EOS Orbital energies are lower in multielectron atoms than in the hydrogen atom

5 Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table 5 Electron Configurations Electron configuration describes the distribution of electrons among the various orbitals in the atom The spdf notation uses numbers to designate a principal shell and the letters to identify a subshell; a superscript number indicates the number of electrons in a designated subshell EOS

6 Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table 6 Formula Mass EOS Each box has arrows representing electron spins; opposing spins are paired together An orbital diagram uses boxes to represent orbitals within subshells and arrows to represent electrons:

7 Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table 7 Rules for Electron Configurations Electrons occupy the lowest available energy orbitals Pauli exclusion principle – no two electrons in the same atom may have the same four quantum numbers EOS Orbitals hold a maximum of two electrons spins must be opposed

8 Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table 8 Rules for Electron Configurations For orbitals of identical energy, electrons enter empty orbitals whenever possible – Hund’s rule Electrons in half-filled orbitals have parallel spins EOS

9 Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table 9 Rules for Electron Configurations EOS Capacities of shells (n) and subshells (l)

10 Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table 10 Rules for Electron Configurations Subshell filling order... Each subshell must be filled before moving to the next level EOS 1s 2 2s 2 2p 6 3s 2 3p 6... Illustration

11 Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table 11 The Aufbau Principle A hypothetical building up of an atom from the one that precedes it in atomic number (Z = 1) H 1s 1 (Z = 2) He 1s 2 (Z = 3) Li 1s 2 2s 1 EOS (Z = 3) Li 1s 2 2s 1  [He]2s 1 Abbreviated electron configuration

12 Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table 12 The Aufbau Principle... [He]2p 2 [He]2p 3 [He]2p 4 [He]2p 5 EOS [He]2p 6

13 Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table 13 Main Group and Transition Elements Elements in which the orbitals being filled in the aufbau process are either s or p orbitals of the outermost shell are called main group elements “A” group designation on the periodic table The first 20 elements are all main group elements In transition elements, the subshell being filled in the aufbau process is in an inner principal shell EOS Fourth period transition elements have n = 4 as their outermost shell as the 3d subshell fills

14 Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table 14 Main Group and Transition Elements EOS Completely filled and half- filled sublevels are more energetically favorable configurations

15 Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table 15 Periodic Relationships The valence shell is the outermost occupied shell The period number = principal quantum number, n, of the electrons in the valence shell EOS

16 Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table 16 Periodic Relationships EOS For main group elements the number of valence shell electrons is the same as the periodic table “A” group number

17 Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table 17 Periodic Relationships We can deduce the general form of electron configurations directly from the periodic table EOS

18 Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table 18 Valence Electrons and Core Electrons Valence electrons are those with the highest principal quantum number EOS Sulfur has six valence electrons

19 Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table 19 Valence Electrons and Core Electrons Electrons in inner shells are called core electrons EOS Sulfur has 10 core electrons

20 Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table 20 Electron Configurations of Ions Anions: gain e – to complete the valence shell Example: EOS -

21 Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table 21 Electron Configurations of Ions Cations: lose e – to attain a complete valence shell Example: (Z = 11) Na EOS (Z = 11) Na +

22 Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table 22 Electron Configurations of Ions Cations formed from transition metals lose e – from the highest principal energy level (n) EOS

23 Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table 23 Magnetic Properties Diamagnetism is the weak repulsion associated with paired electrons Paramagnetism is the attraction associated with unpaired electrons EOS Ferromagnetism is the exceptionally strong attractions of a magnetic field for iron and a few other substances

24 Periodic Trends Use must justify the trend across the period, you cannot simply state the trend. A trend is an observation, not an explanation! You should state the trend in your answer, but you must also go further by explaining what causes the observed trend! Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table 24

25 Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table 25 Periodic Atomic Properties of the Elements Periodic law states that certain sets of physical and chemical properties recur at regular intervals when the elements are arranged according to increasing atomic number EOS Consider atomic radii: distance between the nuclei of two atoms The distance between the nucleus and the outer edge of the electron cloud

26 Atomic Radii Atomic Radii decrease as atomic numbers increase in an given period (going across). –A proton and electron are added so the effective nuclear charge increases because each proton has more of an effect than each additional electron As that attraction between the nucleus and electrons increases, and the atomic radius decreases Atomic Radii increase Sgoing down –In going from top to bottom of a group, the valence electrons are assigned to orbitals with increasingly higher values of n (prin. Quantum number) The underlying electrons requires some space, so the electrons of the outer shell must be further (Your are adding energy levels) 26

27 Atomic Radii Z eff effective nuclear charge: the nuclear charge experienced by a particular electron in a multielectron atom –Increases the attraction of the nucleus and pulls the electron cloud closer to the nucleus resulting in a smaller atomic radius Atomic radii of transtion metals trend a little differently Exceptions in atomic radii also exist in the lanthanide and actinide series because of how the f subshells are uniquely filled by electrons 27

28 Z eff & Shielding –The order in which electrons are assigned to subshells in an atom, as well as other properties are because of Z eff –Shielding: electrons closer to the nucleus screen or shield the effect of nuclear charge on valence electrons the number of shielding electrons increases when you reach the end of the periodic table and go on to the next period. Shielding increases in steps as you start a new period or go down a group Video Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table 28

29 Transition Metal Atomic Trends From left to right across a period, the radii initially decrease, then size remains almost the same, then slightly increases toward the end. The small increase in atomic radii is because of the d subshell is filled with electrons and thus the ele-eletron repulsions cause the size to increase 29

30 Atomic Radii Properties The increased number of energy levels (n) increases the distance over which the nucleus must pull and therefore reduces the attraction for electrons Full energy levels provide shielding between the nucleus and valence electrons, so you see an increase in shielding as the level gets full Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table 30 EOS Illustration

31 Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table 31 Ionic Radii EOS The ionic radius of each ion is the portion of the distance between the nuclei occupied by that ion If the size of an atom is determined by the outermost electrons, what happens if you remove or add an electron?

32 32 Ionic Radii Cations are smaller than the atoms from which they are formed – the nucleus attracts the remaining electrons more strongly EOS Anions are larger than the atoms from which they are formed – the greater number of electrons repel more strongly Think of the proton/electron ratio, -as electrons are lost, the ratio of p+/e- increases and so the electrons are held closer vv.

33 Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table 33 Isoelectronic Configurations Isoelectronic species are elements that all have the same number of electrons For isoelectronic species, the greater the nuclear charge, the smaller the species EOS Effective nuclear charge

34 Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table 34 Atomic and Ionic Radii EOS

35 Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table 35 Ionization Energy Ionization energy is the energy required to remove an electron from a ground state atom in the gaseous state to remove an electron, energy must be supplied to overcome the attraction of the nuclear charge (endothermic, always +) Continual removal of electrons increases ionization energy greatly B  B + + e – I = 801 kJ mol –1 B +  B +2 + e – I = 2427 kJ mol –1 B +2  B +3 + e – I = 3660 kJ mol –1 B +3  B +4 + e – I = 25,025 kJ mol –1 EOS B +4  B +5 + e – I = 32,822 kJ mol –1 Illustration

36 Ionization energy First ionization energy- energy is increased with each successive removal because the electron is being removed from an increasingly positive ion –The remaining electrons are held more tightly –Notice the large jump at the 3 rd level for Mg. –There is a large increase as you remove electrons from lower (inner) energy subshells 36

37 Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table 37 First Ionization Energies EOS Illustration

38 Ionization energy Ionization energy increases as atomic number increases in any given period –Z eff increases the attraction of the nucleus and holds the electrons more tightly Exceptions: group II to III, IE drops because the p electrons do not penetrate the nuclear region as well as s electrons so aren’t as tightly held Drop in IE also occurs between V & VI because of increased repulsion created by the first pairing of electrons, that is stronger than the increase in Z eff, lowering the energy required to remove the electron Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table 38

39 Ionization energy Ionization energy decreases as atomic number increases down a column or group –The increased number of energy levels (n) increases the distance over which the nucleus must pull, reducing the attraction for electrons –A full energy level provides some shielding between the nucleus and valence electrons Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table 39

40 Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table 40 Electron Affinity Electron affinity is the energy change that occurs when an electron is added to a gaseous atom -How much an atom ‘likes’ Electrons (+ or -) -the more negative it is the higher the EA (energy is flowing out of the system) EOS Electron affinities are expressed as negative because the process is exothermic Illustration

41 Electronegativity A measure of the attraction of an atom for the pair of outer shell electrons in a covalent bond with another atom –Pattern is same as electron affinity for same reasons –Both are attraction nucleus has for electrons, one in forming an ion (EA) and one in forming a molecule (EN) –Fluorine is the most electronegative. The closer it is to fluorine, the more electronegative it is. 41

42 Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table 42 Metals, Nonmetals, and Metalloids Metals have a small number of electrons in their valence shells and tend to form positive ions Except for hydrogen and helium, all s-block elements are metals EOS All d- and f-block elements are metals

43 Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table 43 Metals, Nonmetals, and Metalloids Atoms of a nonmetal generally have larger numbers of electrons in their valence shell than do metals, and many tend to form negative ions Nonmetals are all p- block elements and include hydrogen and helium EOS Metalloids have properties of both metals and nonmetals

44 Metals Metals react by losing electrons –A loosely held electron will result in a more reactive metal –This is tied directly to ionization energy –With an increased # of energy levels (n), comes increased distance from the nuclear attraction and thus a more loosely held electron available for reactions Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table 44

45 Non-metals Non-metals tend to gain electrons, a strong nuclear attraction will result in a more reactive non-metal This means that an atom with the highest Z eff and the least number of energy levels should be the most reactive nonmetal (F) because its nucleus exerts the strong pull Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table 45

46 Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table 46 A Summary of Periodic Trends EOS

47 Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table 47 The Inert (Noble) Gases The six noble gases, He, Ne, Ar, Kr, Xe, and Rn, rarely enter into chemical reactions EOS All have complete octets... = stability!

48 Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table 48 “Periodic” Behavior of Elements Flame tests: elements with low first ionization energies are excited in a flame EOS Atoms emit energy when electrons fall from higher to lower energy states FlameTests

49 Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table 49 “Periodic” Behavior of Elements halogens (Group 7A) are good oxidizing agents EOS When Cl 2 is bubbled in a solution containing iodide ions, chlorine oxidizes I – to I 2

50 Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table 50 The s-Block Metals as Reducing Agents EOS Recall activity series... H + is reduced by these metals 2 K + 2 H 2 O  2 K OH – + H 2

51 Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table 51 Acidic, Basic, and Amphoteric Oxides Acidic oxides are oxides that produce acids by reacting the oxide with water e.g., SO 3 + H 2 O  H 2 SO 4 EOS

52 Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table 52 Acidic, Basic, and Amphoteric Oxides Basic oxides are oxides that produce bases by reacting with water e.g., MgO + H 2 O  Mg(OH) 2 EOS

53 Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table 53 Acidic, Basic, and Amphoteric Oxides Oxides that can react with either acids or bases are amphoteric oxides e.g., Al 2 O 3 EOS Behavior of Oxides

54 Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table 54 Summary of Concepts The wave-mechanical treatment of the hydrogen atom can be extended to multielectron atoms, but with two differences Electron configuration is the distribution of electrons in orbitals among the subshells and principal subshells EOS There are two types of electron configuration notation: spdf and orbital

55 Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table 55 Summary of Concepts The aufbau principle describes a process of hypothetically building up an atom from the atom of the preceding atomic number Elements in similar electron configurations fall in the same group of the periodic table An atom with all the electrons paired is diamagnetic; an atom with one or more unpaired electrons is paramagnetic EOS Certain atomic properties, such as atomic radius, ionic radius, ionization energy, and electron affinity, vary periodically with increasing atomic number

56 Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table 56 Summary of Concepts The regions of the periodic table ascribed to metals, nonmetals, metalloids, and the noble gases are related to the value of atomic properties EOS


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