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  1. PRINCIPLES OF CHEMISTRY I CHEM 1211 CHAPTER 8 DR. AUGUSTINE OFORI AGYEMAN Assistant professor of chemistry Department of natural sciences Clayton state university


  3. CLASSIFICATION OF THE ELEMENTS - Elements in a given group have similar chemical properties because the outer-shell electron arrangements are similar Group IIA elements Be: 1s22s2 Mg: 1s22s22p63s2 Ca: 1s22s22p63s23p64s2 Sr: 1s22s22p63s23p64s23d104p65s2

  4. CLASSIFICATION OF THE ELEMENTS - The last electron in an element’s electron configuration causes the difference in the electron configuration of the preceding element and is referred to as the distinguishing electron Homework Write notes (one page) on the different classifications of the elements based on electronic properties. Briefly describe the s-area, p-area, d-area, and the f-area.

  5. CLASSIFICATION OF THE ELEMENTS • Elements can be classified as • Metalsor Nonmetals • Based on physical properties • Elements can also be classified as • Noble-gas, Representative, • Transition, or Inner Transition • - Based on electron configuration

  6. CLASSIFICATION OF THE ELEMENTS Noble-gas Elements - Group VIIIA (18) elements on the periodic table (far right column) - Gases at room temperature - Little tendency to form chemical compounds - Electron configuration ends in p6 - Completes p subshell (except Helium) - Nonmetals

  7. CLASSIFICATION OF THE ELEMENTS Representative Elements - Elements in the s-area (Groups IA and IIA) first five columns of the p-area (Groups IIIA, IVA, VA, VIA, and VIIA) - Metals and nonmetals

  8. CLASSIFICATION OF THE ELEMENTS Transition Elements - Elements in the d-area of the periodic table - Groups IIIB (3), IVB (4), VB (5), VIB (6), VIIB (7), VIIIB (8, 9, 10), IB (11), and IIB (12) - Distinguishing electron in a d subshell - Metals

  9. CLASSIFICATION OF THE ELEMENTS Inner Transition Elements - Elements in the f-area of the periodic table - The two-row block of elements below the main table - Distinguishing electron in an f subshell - Metals

  10. VALENCE ELECTRONS - Electrons in the highest principal quantum number of an atom, and any electrons in an unfilled subshell from a lower shell CORE ELECTRONS - Electrons in filled subshells that have lower principal quantum numbers

  11. VALENCE ELECTRONS - Representative elements in the same group of the periodic table have the same number of valence electrons - The number of valence electrons for representative elements is the same as the group number (with A) in the periodic table - The maximum number of valence electrons for any given element is eight

  12. VALENCE ELECTRONS - Not all electrons in a given atom participate in bonding - Only valence electrons are available for bonding - For representative and noble-gas elements these electrons are always found in the s or p subshells

  13. VALENCE ELECTRONS - The electron configuration can be used to determine the number of valence electrons of an atom C: 1s22s22p2 O: 1s22s22p4 Na: 1s22s22p63s1 As: 1s22s22p63s23p64s23d104p3

  14. ELECTRON CONFIGURATION OF IONS Cl: [Ne]3s23p5 Cl-: [Ne]3s23p6 Na: [Ne]3s1 Na+: [Ne] F : 1s22s22p5 F- : 1s22s22p6 Co: [Ar]3d74s2 Co3+: [Ar]3d6

  15. EFFECTIVE NECLEAR CHARGE • Negatively charged electrons are attracted to the positively • charged nucleus • The force of attraction between an electron and the nucleus • - Depends on the magnitude of the net nuclear charge acting • on the electrons • - Depends on the average distance between the nucleus and • the electron • (Coulombs law)

  16. EFFECTIVE NECLEAR CHARGE The force of attraction - Increases with increasing nuclear charge - Decreases with increasing average distance between electrons and the nucleus - Electrons also experience repulsion by other electrons in the atom Zeff = Z – S Zeff = effective nuclear charge Z = actual nuclear charge (number of protons in the nucleus)(> Zeff) S = screening constant (represents number of core electrons)

  17. EFFECTIVE NECLEAR CHARGE Atomic number of Na = 11 Number of valence electrons = 1 Number of core electrons = 10 As simplified from this model Z = 11 S = 10 Zeff = 11-10 = +1 In actual fact Zeff in Na is about +2.5

  18. EFFECTIVE NECLEAR CHARGE In general Zeff in s orbital > Zeff in p orbital > Zeff in d orbital > Zeff in f orbital - This is the result of the trend in energy levels ns < np < nd < df - Zeff increases across the periods (from left to right) in the periodic table (Z increases but S remains the same) - Zeff changes slightly down the groups in the periodic table (essentially the same in a given group)

  19. SIZES OF ATOMS - Atomic radius tends to decrease across the periods (from left to right) in the periodic table - Due to increase in effective nuclear charge which draws valence electrons closer to the nucleus - Atomic radius tends to increase down the groups (from top to bottom) of the periodic table - Due to increase in principal quantum number of the outer electrons (number of shells)

  20. SIZES OF ATOMS - Cations are smaller than their parent atoms - Decrease in the number of electrons decreases electron-electron repulsions - Anions are larger than their parent atoms - Increase in the number of electrons increases electron-electron repulsions - Ionic size increases down the group of the periodic table for ions carrying the same charge

  21. ISOELECTRONIC SERIES - A group of atoms and ions containing the same number of electrons Due to the same number of electrons - Ionic radius decreases with increasing nuclear charge (electrons are more strongly attracted to the nucleus) increasing nuclear charge O2- , F- , Na+ , Mg2+ , Al3+ S2-, Cl-, K+ , Ca2+ , Ga3+ Se2- , Br- , Rb+ , Sr2+ , In3+ decreasing ionic radius

  22. IONIZATION ENERGY • The energy required to remove an electron from • a gaseous atom or ion • X(g) → X+(g) + e- • E is positive • - The atom or ion is assumed to be in its ground state • - The highest energy electron is always removed first • Units: kJ/mol • 96.485 kJ/mol = 1 eV

  23. IONIZATION ENERGY Ionization Energies of Magnesium (Mg) Mg(g) → Mg+(g) + e- I1 = 738 kJ/mol Mg+(g) → Mg2+(g) + e- I2 = 1450 kJ/mol Mg2+(g) → Mg3+(g) + e- I3 = 7734 kJ/mol I1 < I2 < I3

  24. IONIZATION ENERGY First Ionization Energy (I1) - Energy required to remove the highest energy electron of an atom - The first electron is removed from a neutral atom - The second electron is removed from a positive ion (more difficult) - Increase in positive charge binds electrons more tightly - Large jump is observed in going from removal of valence electrons to removal of core electrons

  25. IONIZATION ENERGY - First ionization energy increases across the period of the periodic table (from left to right) - Electrons added in the same principal quantum number do not completely shield increasing nuclear charge - First ionization energy decreases down the group of the periodic table (from top to down) - As n increases, the size of orbital increases (distance from nucleus increases) and electrons are easier to remove

  26. IONIZATION ENERGY - Discontinuities are due to electron repulsions and shielding (Be to B, N to O) - Representative elements show a larger range of values of I1 than the transition-metal elements - Smaller atoms have higher ionization energies

  27. PSEUDO-NOBLE-GAS CONFIGURATION - Several cations with the pseudo-noble-gas electron configuration are more stable [noble gas](n-1)d10 In: [Kr]5s24d105p1 In+: [Kr]5s24d10 Pseudo = In3+: [Kr]4d10 Sn: [Kr]5s24d105p2 Sn2+: [Kr]5s24d10 Pseudo = Sn4+: [Kr]4d10

  28. ELECTRON AFFINITY - The energy change that occurs when an electron is added to a gaseous atom - A measure of the attraction of the atom for the added electron - Energy is released when an electron is added to most atoms X(g) + e- → X-(g) E is negative Units: kJ/mol

  29. ELECTRON AFFINITY Cl(g) + e-→ Cl-(g) E = -349 kJ/mol - The greater the attraction between an atom and an added electron the more negative the atom’s electron affinity - Electron affinities for noble gases are positive values (E > 0) - Halogens have the most negative electron affinities - Electron affinity changes slightly down the group of the periodic table

  30. METALS • - Refer to chapter 2 for properties of metals • - Metals tend to have low ionization energies • - Metals form positive ions relatively easily • - Metals lose electrons (oxidize) when they • undergo chemical reactions • Charge on metals • Alkali metals: 1+ • Alkaline earth metals: 2+

  31. METALS - The charge on transition metals do not follow any obvious pattern - Transition metals are able to form more than one positive ion - Compounds of metals and nonmetals are ionic 2Na(s) + Cl2(g) → 2NaCl(s) (contains Na+ and Cl- ions)

  32. METALS Most Metal Oxides are Basic - Dissolve in water to form metal hydroxides Na2O(s) + H2O(l) → 2NaOH(aq) O2-(aq) + H2O(l) → 2OH-(aq) (net ionic equation) - React with acid to form salt and water NiO(s) + 2HNO3(aq) → Ni(NO3)2(aq) + H2O(l)

  33. NONMETALS - Refer to chapter 2 for properties of nonmetals - Nonmetals tend to have high electron affinities - Nonmetals tend to gain electrons when they react with metals - Compounds composed of only nonmetals are generally molecular substances

  34. NONMETALS Most Nonmetal Oxides are Acidic CO2(g) + H2O(l) → H2CO3(aq) (acidity of rainwater) - Dissolve in basic solutions to form salt and water CO2(g) + 2NaOH(aq) → Na2CO3(aq) + H2O(l)

  35. ALKALI METALS (GROUP IA) - Alkali means ‘ashes” Relatively abundant in the - earth’s crust (Na, K) - sea water - human bodies - Have low densities and melting points - Very reactive and readily lose an electron to form 1+ ions

  36. ALKALI METALS (GROUP 1A) - Form hydrides with hydrogen and sulfides with sulfur 2M(s) + H2(g) → 2MH(s) 2M(s) + S(s) → M2S(s) - React vigorously with water to produce hydrogen gas and alkali metal hydroxide (very exothermic and may explode) 2M(s) + 2H2O(l) → 2MOH(aq) + H2(g)

  37. ALKALI METALS (GROUP 1A) - Can react with oxygen to form oxides, peroxides, and superoxides Li forms only oxides (O2-) 4Li(s) + O2(g) → 2Li2O(s) (lithium oxide) Na can form peroxides (O22-) 2Na(s) + O2(g) → Na2O2(s) (sodium peroxide) K, Rb, and Cs can form peroxides (O22-) and superoxides (O2-) K(s) + O2(g) → KO2(s) (potassium superoxide)

  38. ALKALINE EARTH METALS (GROUP 2A) Compared to alkali metals - harder - more dense - higher melting points - less reactive than the respective adjacent alkali metal - Tend to lose two outer s electrons to from 2+ ions - Give off characteristic colors when heated in a flame (salts used in fireworks)

  39. ALKALINE EARTH METALS (GROUP 2A) - Reactivity increases down the group - Beryllium does not react with water - Magnesium reacts slowly with water - Calcium and elements below react readily with water Ca(s) + 2H2O(l) → Ca(OH)2 + H2(g)

  40. HYDROGEN - First element in the periodic table (1s1 electron configuration) - Nonmetal - Can be metallic under extreme pressures - Colorless diatomic gas - Has very high ionizaton energy - More than double those of alkali metals - Due to absence of nuclear shielding of the 1s electron

  41. HYDROGEN - Does not easily lose its valence electron - Share with nonmetals to form molecular compounds - Can lose its electron to form a cation (H+) - Can gain electron to form the hydride ion (H-)

  42. CHALCOGENS (GROUP 6A) THE OXYGEN GROUP - Properties change from nonmetallic to metallic down the group Nonmetallic properties: oxygen, sulfur, selenium Metallic properties: tellurium and below - Oxygen is a colorless gas at room temperature - The other group members are solids - Oxygen exists as O2 (oxygen gas) and O3 (ozone) - Allotropes (different forms of the same element in the same state)

  43. CHALCOGENS (GROUP 6A) THE OXYGEN GROUP - O2 can produce O3 in lightning storms 3O2(g) → 2O3(g) Ho = 284.6 kJ - Sulfur also has several allotropic forms - The most common is S8 (yellow solid)

  44. THE HALOGENS (GROUP 7A) - All halogens are nonmetals - Melting and boiling points increase with increasing atomic number - Consist of diatomic molecules (F2, Cl2, Br2, and I2) - Form colored gases

  45. THE HALOGENS (GROUP 7A) At room temperature - Fluorine and chlorine are gases - Bromine is a liquid - Iodine is a solid - Have highly negative electron affinities - Tend to gain electrons to form 1- ions - Reactivity decreases down the group

  46. THE HALOGENS (GROUP 7A) - React readily with most metals to form ionic halides - React with hydrogen to form gaseous hydrogen halides H2(g) + X2 → 2HX(g) - Hydrogen halides dissolve in water to form acids [HCl(aq)] - Fluorine is very reactive (dangerous) - Chlorine is the most industrially useful

  47. THE NOBLE GASES (GROUP 8A) - Nonmetals - Monatomic - Gases at room temperature - Have completely filled s and p subshells - Have high first ionization energies - Have stable electron configuration - Unreactive