8.1 Early Periodic Tables CHAPTER 8. Modern Periodic Table. Mendeleev s 1871 Table

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1 8.1 Early Periodic Tables CHAPTER 8 Periodic Relationships Among the Elements 1772: de Morveau table of chemically simple substances 1803: Dalton atomic theory, simple table of atomic masses 1817: Döbreiner's triads 3 elements w/ regularly varying properties: S Se Te 1865: Newlands "law of octaves", about 55 elements Early tables were based on mass number (A) or combining weight Modern Periodic Table 20 elements Published about : Mendeleev and Meyer "properties of the elements are a periodic function of their atomic weights;" 63-element table : Rayleigh et al. discover noble gases with no tendency to form ions 1913: Moseley X-ray emission spectra vary with atomic number (Z) 1940: McMillan and Abelson first transuranium element, neptunium Mendeleev s 1871 Table 65 elements Arranged by chemical similarity Predicted missing elements 8.2 Periodicity Example: Chalcogens, group 6A (16) O = [He]2s 2 2p 4 S = [Ne]3s 2 3p 4 Se = [Ar]4s 2 3d 10 4p 4 Te = [Kr]5s 2 4d 10 5p 4 Similarities Nonmetals (Te is metalloid) Form 2- ions Form oxides EO 2

2 Valence electrons Noble Gases (18) completely filled subshells Similarity of the electron configurations causes the chalcogens to behave similarly Electrons after the noble gas core are called outer electrons or valence electrons Valence electrons are involved in chemical bonding Underlying (n-1)d 10 electrons don t have a major effect on chemistry Lanthanides, actinides incompletely filled f subshells (3) Transition metals - atom or cation w/ incompletely filled d subshell Representative elements incompletely filled s or p subshells (1, 2, 13-17) Elements in Chemical Equations Metals, metalloids and non-metals with complex 3-D network structures use empirical formula (element symbol) Fe, Si, C (except molecular forms like C 60 ) Non-metals that exist as discrete molecules use molecular formula H, O, F, Br, I, N, Cl are diatomic H 2, O 2 etc. Other common non-metals: P 4, S 8 e - Configurations of Anions Additional electrons go into next available orbital (Aufbau) S 2- S [Ne] 3s 2 3p 4 S 2- [Ne] 3s 2 3p 6 Usually formed from non-metals Anion usually has noble gas configuration Charge usually Group # - 18 (or 8) e - Configurations of Cations Start with configuration neutral element Remove electrons from subshell with the highest n, p before s Usually formed from metals Main group cations Often have noble gas configuration Often charge = Group # (-10) Transition metals can have several cations, with noble gas configuration + d e - s Cation Examples Sr 2+ Sr [Kr] 5s 2 Sr 2+ [Kr] Ti 4+ Ti [Ar] 4s 2 3d 2 Ti 4+ [Ar] Mn 2+ Mn [Ar] 4s 2 3d 5 Mn 2+ [Ar] 3d 5

3 Isoelectronic Atoms or ions with same e - configuration F - [He] 2s 2 2p 6 O 2- [He] 2s 2 2p 6 Name a cation isoelectronic with F -. Na + or Mg 2+ Transition Metal Cations Electron configuration of Ti (Z = 22)? Ti [Ar] 4s 2 3d 2 Electron configuration of Cr (Z = 24)? Cr [Ar] 4s 1 3d 5 Electron configuration of Cr 2+? Cr2+ [Ar] 3d 4 Are Ti and Cr 2+ isoelectronic? No; same number of e - but different configuration. 8.3 Periodicity of Physical Properties Shielding electrons reduce electrostatic attraction between nucleus and outer electrons Effective nuclear charge (Z eff ) Nuclear charge felt by an electron Z eff = Z - σ (σ is the shielding or screening constant, 0 < σ < Z) Degrees of shielding e - s are shielded very effectively by e - s in subshells closer to the nucleus (lower n) e - s are shielded slightly by other e - s in the same shell (same n) e - s are not shielded by e - s further from the nucleus (higher n) Shielding of outer e - in Li and Be Shielding in multi-electron atoms This 1s e - spends very little time between the other 1s e - and nucleus. e e - e - Li Z eff = 3-2 = 1 e e- e - e - Be Z eff = 4-2 = 2 Ineffective shielding: Each 1s e - feels nearly full nuclear charge This 1s e - spends most of its time between this 2p e - and nucleus. Effective shielding: 1s e - feels nearly full nuclear charge; 2p e - feels diminished nuclear charge

4 Defining Atomic Radius Low Z eff High Z eff n = 1 1/2 distance between nuclei of two adjacent, identical atoms (a) In metal or 3-dimensional network, e.g., diamond (b) In covalent, diatomic molecule n = 6 Variation of Atomic Radius with Z Ionic Radii: Cations Top to bottom: increasing n d increasing r Left to right: increasing Z eff d decreasing r Li vs. Li + Each has nucleus with +3 charge Li has 3 e -, highest is 2s 1 Li + has 2 e -, higher is 1s 2 Li = 152 pm, Li + = 60 pm Less interelectronic repulsion d smaller cloud Ionic Radii: Anions Reaction of Li with F to form LiF Cl vs. Cl - Each has nucleus with +17 charge Cl has 17 e -, highest is 3p 5 Cl - has 18 e -, highest is 3p 6 Cl = 99 pm, Cl - = 181 pm Greater interelectronic repulsion d larger cloud Remove valence electron from Li to form Li + reduce interelectronic repulsion Li+ is smaller Add another electron to F to form F - increase interelectronic repulsion F - is larger

5 Radii of Isoelectronic Series Relative sizes are determined by number of electrons and protons. S 2-, Cl -, Ar, K +, Ca 2+ All have 18 e - Nuclear charge (Z) increases S (16), Cl (17), Ar (18), K (19), Ca (20) Higher Z, greater attraction for valence e - s Size S 2 - > Cl - > (Ar?) > K + > Ca 2+ Radii (pm) 184 > 181 > (98?) > 133 > 99 Why is Ar out of line? Ionic radii are inferred from distances between ions in crystals; Ar forms no bonds, radius measured in gas. 8.4 Ionization Energy Minimum energy required to remove an electron from a ground-state, gaseous atom Always positive (requires energy) Measures how tightly e - is held in atom Energy associated with: X(g) d X + (g) + e - Ionization Energy 1st ionization energy (I 1 ) E required to remove first electron X(g) d X + (g) + e - 2nd ionization energy (I 2 ) E required to remove second electron X + (g) d X 2+ (g) + e - Which requires more energy? 2nd ionization is always larger harder to remove negative charge from a cation Table 8.3 Ionization Energies (kj/mole) Z Element First Second Third 1 H He Li Be Variation of I 1 with Z Top to bottom: increasing n d increasing r d weaker attraction d decreasing I 1 Left to right: increasing Z eff d increasing I 1

6 Inconsistent Variation of I 1 Across Period Be: [He]2s 2 B: [He]2s 2 2p 1 Outer e- of B is in 2p, less strongly held than 2s, lowers I 1 N: [He]2s 2 2p 3 O: [He]2s 2 2p 4 2 outer e- of O are in same 2p, inter-e - repulsion lowers I Electron Affinity Energy change when an e - adds to a gas-phase, ground-state atom Energy associated with this reaction: X(g) + e - d X - (g) By convention, positive EA means that energy is released, e - addition is favorable! Caution: Thinkwell uses the opposite sign convention! Mostly positive, a few zero or negative Electron Affinity Trends Left to right: increasing Z eff d increasing EA Top to bottom: not much variation (same valence e - configuration) EAs often lower for metals than nonmetals Alkali metal (Grp. 1) EAs are positive! added e- gives favorable, filled ns 2 Be and Mg (Grp. 2) EAs are negative added e- goes into higher energy p subshell Nitrogen (Grp. 15) EA is 0 added e - goes into already half -filled 2p 3 interelectronic repulsion disfavors addition Halogen EAs (Grp. 17) are largest added e - produces full ns 2 np 6 configuration Noble gas EAs are negative added e - goes into next higher n shell 8.6 Periodic Properties of the Representative Elements Ionization energy and electron affinity are related for most atoms Metals usually lose e- relatively easily (low IE) but release little E when e - adds (low EA) Nonmetals usually lose e- with difficulty (high IE) but release much E when e - adds (high EA) Combination indicates overall tendency to hold e - similar to electronegativity, Chap. 9 Min Min Z eff Periodic trends Max Ionization E (I 1 ) Max Atomic radius Min Max Electron affinity Min Max

7 Charge Density Ratio of ionic charge to size is a useful measure of ionic behavior Cations with a high positive charge density tend to draw electrons toward themselves Anions that are large and/ or highly negative tend to be strongly distorted by cations None Slight Large Distortion (polarization) - Diagonal Relationship Chemical resemblance between an element and the element one down and to the right Diagonal relationships mainly result from similarity in charge density Charge density (approx., C mm -3 ) Li + 98 Na + 24 Be Mg B Al C 4+ Si Hydrogen Unique behavior because of small size 1s 1 electron configuration H loses 1 e - to form H + (proton) e.g., H 2 (g) + Cl 2 (g) d 2 HCl (g) H + bonds to solvent, e.g. H 3 O + H gains 1 e - to form H - (1s 2 ) w/ Grp. 1 or 2 metals e.g. 2K(s) + H 2 (g) d 2KH (s) Group 1 (1A): Alkali Metals Li, Na, K, Rb, Cs, (Fr) ns 1 electron configurations often lose 1 e- to form M + react with water 2M(s) + 2 H 2 O(l) d 2MOH(aq) + H 2 (g) react with oxygen (O 2 ) form Li 2 O (oxide), Na 2 O 2 (peroxide) and MO 2 (superoxide) depending on conditions Group 2 (2A): Alkaline Earths Be, Mg, Ca, Sr, Ba, Ra ns 2 electron configurations often lose 2 e- to form M 2+ react with water less vigorously than Grp. 1 M(s) + 2 H 2 O(l) d M(OH) 2 (aq) + H 2 (g) react with oxygen (O 2 ) on heating Be and Mg require heat typically form MO (oxide) Group 13 (3A) B (metalloid), Al, Ga, In, Tl (metals) ns 2 ((n-1)d 10 ) np 1 electron configurations often lose 3 e- to form M 3+ esp. B tends to form molecular rather than ionic compounds Ga, In, Tl also form M + (ns 2 ) inert lone pair effect Al forms protective coating of Al 2 O 3 resists water and oxygen

8 Group 14 (4A) C (nonmetal), Si, Ge (metalloids), Sn, Pb (metals) C allotropes diamond, graphite and fullerenes ns 2 ((n-1)d 10 ) np 2 electron configurations formally lose 4 e- to form M 4+ tend to form molecular rather than ionic compounds, e.g., CO (g), SiO 2 (s) Sn, Pb also form M 2+ (ns 2 ) inert lone pair effect Group 15 (5A): Pnicogens N, P (nonmetal), As (metalloid), Sb, Bi (metals) ns 2 ((n-1)d 10 ) np 3 electron configurations elements N 2 (g), P 4 (s) etc. N adds 3 e- to form N 3 - [Ne] also form M (III) and M (V) compounds many molecular rather than ionic compounds, e.g., NO 2 (g), P 4 O 10 (s), AsH 3 (g) HNO 3 (nitric) and H 3 PO 4 (phosphoric) acids Group 16 (6A): Chalcogens O, S, Se (nonmetals), Te, Po (metalloids) ns 2 ((n-1)d 10 ) np 4 electron configurations elements O 2 (g), O 3 (g), S 8 (s) etc. O, S, Se, Te add 2 e- to form E 2 - [Noble gas] form many M (II), (IV) and (VI) compounds many molecular rather than ionic compounds, e.g., SO 2 (g), SF 6 (g), H 2 Se(g) H 2 SO 4 = sulfuric acid Group 17 (7A): Halogens F, Cl, Br, I, At (nonmetals) ns 2 ((n-1)d 10 ) np 5 electron configurations elements F 2 (g), Cl 2 (g), Br 2 (l), I 2 (s) all add 1 e - to form X - [Noble gas] M + X - compounds are usually ionic X(I), X(III), X(V) and X(VII) compounds are usually molecular; e.g., ICl 3, Cl 2 O 7 HClO x (x = 0, 1, 2, 3, 4) acids Group 18 (8A): Noble Gases He, Ne, Ar, Kr, Xe, Rn (nonmetals) ns 2 ((n-1)d 10 ) np 6 electron configurations elements all monatomic gases very unreactive Kr and Xe form a few compounds with F and O, e.g. KrF 2, XeO 3 Group 1 (Ia) vs. 11 (IB) K = [Ar] 4s1 I 1 = 419 kj/mol r(k) = 227 pm r(k+) = 133 pm K is very reactive! K + is not very polarizable Cu = [Ar] 3d 10 4s1 I 1 = 745 kj/mol r(cu) = 128 pm r(cu+) = 96 pm Cu is unreactive! Cu+ is very polarizable Why? Inefficient shielding of 4s e - by underlying 3d 10 e - s low d-e - density at nucleus d e - s are easily polarized

9 Oxide Trends NaO MgO Al 2 O 3 SiO 2 P 4 O 10 Cl 2 O 7 Ionic assssssssssssd Molecular 3-D lattice asssssssd Discrete units High-melting solid asssssssd Gas Basic assd Amphoteric* assssd Acidic * Displays both acidic and basic properties

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