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Chemical Periodicity 化學週期性

5. Chemical Periodicity 化學週期性. The reaction of H 2 O and Li (left) to produce LiOH and H 2(g) is much slower than the analogous reaction between H 2 O and Na (right). Chapter Goals. 5-1 More About the Periodic Table Periodic Properties of the Elements 元素的週期特質 5-2 Atomic Radii 原子半徑

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Chemical Periodicity 化學週期性

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  1. 5 Chemical Periodicity 化學週期性 The reaction of H2O and Li (left) to produce LiOH and H2(g) is much slower than the analogous reaction between H2O and Na (right).

  2. Chapter Goals 5-1 More About the Periodic Table Periodic Properties of the Elements元素的週期特質 5-2 Atomic Radii 原子半徑 5-3 Ionization Energy游離能 5-4 Electron Affinity電子親和力 5-5 Ionic Radii 離子半徑 5-6 Electronegativity 電負性 5-7 Oxidation States Chemical Reactions and Periodicity 5-7 Hydrogen & the Hydrides • Hydrogen • Reactions of Hydrogen and the Hydrides 5-8 Oxygen & the Oxides • Oxygen and Ozone • Reactions of Oxygen and the Oxides • Combustion Reactions • Combustion of Fossil Fuels and Air Pollution

  3. More About the Periodic Table • Establish a classification scheme of the elements based on their electron configurations. • Noble Gases 鈍氣 Group 8A • All of them have completely filled electron shells. • Since they have similar electronic structures, their chemical reactions are similar. ns2np6 He 1s2 Ne [He] 2s2 2p6 Ar [Ne] 3s2 3p6 Kr [Ar] 4s2 3d10 4p6 Xe [Kr] 5s2 4d10 5p6 Rn [Xe] 6s2 4f14 5d10 6p6

  4. More About the Periodic Table • Representative Elements典型元素 • Are the elements in A groups on periodic chart. • These elements will have their “last” electron in an outer s or p orbital. • These elements have fairly regular variations in their properties.

  5. More About the Periodic Table • d-Transition Elements過渡金屬元素 • Elements on periodic chart in B groups. • Sometimes called transition metals. • Each metal has d electrons. • ns (n-1)d configurations • These elements make the transition from metals to nonmetals. • Exhibit smaller variations from row-to-row than the representative elements. First transition series (4s and 3d orbital occupy): 21Sc through 30Zn Second transition series (5s and 4d orbital occupy): 39Sc through 48Cd Third transition series (6s and 5d orbital occupy): 52La and 72Hf to 80Hg Fourth transition series (6s and 5d orbital occupy): 89Ac and 104Rf to 112

  6. More About the Periodic Table • f - transition metals • Sometimes called inner transition metals. • Electrons are being added to f orbitals. • Outermost electrons have the greatest influence on the chemical properties of elements. • Adding an electron to an sorp orbital •  dramatic change in the physical and chemical properties • Adding an electron to a dorf orbital •  much smaller effect on properties

  7. Periodic Propertiesof the Elements Atomic Radii 原子半徑 Ionization Energy游離能 Electron Affinity電子親和力 Ionic Radii 離子半徑 Electronegativity 電負性

  8. Atomic Radii原子半徑 relative sizes Atomic radii decrease 同族元素原子序增加  原子半徑愈大 同週期典型元素大致隨原子序增加 →原子半徑愈小 Å = 10-10m Atomic radii increase

  9. Atomic Radii • The reason the atomic radii decrease across a period is due to shielding or screening effect遮罩效應. • Effective nuclear charge有效核電荷, Zeff, experienced by an electron is less than the actual nuclear charge, Z. • 最外層電子受原子的吸引力被內層電子對外層電子的排斥力給抵銷即稱遮罩效應 • Moving across a period, each element has an increased nuclear charge and the electrons are going into the same shell (2s and 2p or 3s and 3p, etc.). • Consequently, the outer electrons feel a stronger effective nuclear charge. • For Li, Zeff ~ +1 For Be, Zeff ~ +2

  10. 有效核電荷 原子核質子數－內層電子數總和 鎂的有效核電荷約為12 – 10 = 2

  11. 原子半徑隨原子序數的增加呈現週期性變化。這與原子有效核電荷的週期性變化相關原子半徑隨原子序數的增加呈現週期性變化。這與原子有效核電荷的週期性變化相關 • 因為有效核電荷愈大, 對外層電子的吸引力愈大, 原子半徑就愈小 • 各週期的主族從左到右, 電子層數不變, 有效核電荷增加明顯, 原子半徑的逐漸減少也就比較明顯 • 長週期中的過渡元素原子半徑先是緩慢縮小然後略有增大 • 內過渡元素，有效核電荷變化不大，原子半徑幾乎不變 • 稀有氣體原子半徑突然增大，因為它是van der Waals半徑 • 同一主族從上到下，由於電子層數增加，使遮罩效應明顯加大，所以原子半徑遞增。 • 凡得瓦半徑 為兩個相鄰非鍵結原子間距離的一半

  12. Atomic Radii • Example 5-1: Arrange these elements based on their atomic radii. • Se, S, O, Te • P, Cl, S, Si • Ga, F, S, As • Cs, F, K, Cl 6A O < S < Se < Te Period 3 Cl < S < P < Si F < S < As < Ga F < Cl < K < Cs Atomic radii decrease Atomic radii increase

  13. Ionization Energy 游離能 • First ionization energy (IE1) 第一游離能 • The minimum amount of energy required to remove the most loosely bound electron from an isolated gaseous atom to form a 1+ ion 處於基態的氣態原子失去一個電子，成為＋1 價氣態陽離子，所需吸收的能量 • Symbolically: Atom(g) + energy  ion+(g) + e- Mg(g) + 738kJ/mol  Mg+ + e-

  14. Ionization Energy • Second ionization energy (IE2)第二游離能 • The amount of energy required to remove the second electron from a gaseous 1+ ion. • Symbolically: ion+ + energy  ion2+ + e- • Mg+ + 1451 kJ/molMg2+ + e- • Atoms can have 3rd (IE3), 4th (IE4), etc. ionization energies.

  15. Ionization Energy Periodic trends for Ionization Energy: • IE2 > IE1 It always takes moreenergy to remove a second electron from an ion than from a neutral atom. • IE1 generally increases moving from IA elements to VIIIA elements. Important exceptions at Be & Mg, N & P, etc. due to filled and half-filled subshells. • IE1 generally decreases moving down a family. IE1 for Li > IE1 for Na, etc. 隨原子序增大而遞減 increase 隨原子序增大漸增趨勢 Decrease

  16. 同一族元素的游離能 • 隨原子序的增加而變小 • 例如第1族元素的第一游離能大小順序為：Li > Na > K > Rb > Cs > Fr • 同族元素的原子序愈大，原子半徑愈大，最外殼層電子受原子核的引力愈小，故電子愈容易移除。

  17. 游離能的週期性 • 同一週期之各元素的第一游離能，隨原子序的增加，呈現鋸齒狀的增加 • 原子半徑愈小，且有效核電荷愈大，所需的IE愈大 • 同一週期第一游離 • 2族> 3族，鎂(738 kJ/mol)>鋁(578 kJ/mol) • 15(5A) >16(6A)，磷(1012 kJ/mol)>硫(1000 kJ/mol) • 同一週期過渡元素，因原子半徑及有效核電荷彼此差異不大，故其第一游離能彼此差異相對較小

  18. 游離能的規則 同一週期元素的第一游離能從左至右漸增, 但在中間有些起伏IIA > IIIA, 而VA > VIA: • IIA(ns2) > IIIA(ns2np1)是因為np能階較ns高, 因此IIIA原子的游離能較小 • VA (ns2np3) > VIA(ns2np4) • 因為VA 族的價電子組態為全半填滿, 而VIA族有一個p軌域填入兩個電子, 將增加電子的斥力, 因此較易移去,故VA > VIA

  19. First Ionization Energies of Some Elements Elements with low ionization energies (IE) easily lose to form cations (positive charge)

  20. Ionization Energy • Example 5-2: Arrange these elements based on their first ionization energies. • Sr, Be, Ca, Mg • Al, Cl, Na, P • B, O, Be, N • Na, Mg, Al, Si Sr < Ca < Mg < Be Na < Al < P < Cl B < Be < O < N Na < Al < Mg < Si increase Decrease

  21. Ionization Energy • First, second, third, etc. ionization energies exhibit periodicity as well • Look at the following table of ionization energies versus third row elements • Notice that the energy increases enormously when an electron is removed from a completed electron shell Al: 1s22s22p63s23p1 Na: 1s22s22p63s1

  22. Ionization Energy • The reason Na forms Na+ and not Na2+ is that the energy difference between IE1 and IE2 is so large. • Requires more than 9 times more energy to remove the second electron than the first one. • The same trend is persistent throughout the series. • Thus Mg forms Mg2+ and not Mg3+. • Al forms Al3+.

  23. Ionization Energy • Example 5-3: What charge ion would be expected for an element that has these ionization energies? Notice that the largest increase in ionization energies occurs between IE7 and IE8. Thus this element would form a 1- ion

  24. Electron Affinity電子親和力 • Electron affinity(EA) is the amount of energy absorbed when an electron is added to an isolated gaseous atom to form an ion with a 1- charge. 加一個電子給中性氣體原子而形成負離子時所需的能量即稱電子親合力 • Sign conventions for electron affinity. • If electron affinity > 0 energy is absorbed • If electron affinity < 0 energy is released 對於放熱愈多的反應，我們稱其電子親和力愈大，表示該原子愈易獲得電子 • Electron affinity is a measure of an atom’s ability to form negative ions • Symbolically: atom(g) + e- + EA ion-(g)

  25. Electron Affinity Two examples of electron affinity values: Mg(g) + e- + 231 kJ/mol Mg-(g) EA = +231 kJ/mol • Br(g) + e-  Br-(g) + 323 kJ/mol • EA = -323 kJ/mol 電子親和力大 Most elements have no affinity for an additional electron and thus have an electron affinity equal to zero He(g) + e-He-(g) EA=0 kJ/mol

  26. 電子親和力的性質 • 電子親和力的負值愈大, 表示該原子接受電子的傾向愈強, 所形成的陰離子也愈加穩定 • 若電子親和力為正值, 則所形成的陰離子較不穩定

  27. Electron Affinity • General periodic trend for electron affinity is • the values become more negative from left to right across a period on the periodic chart • the values become more negative from bottom to top up a row on the periodic chart • Measuring electron affinity values is a difficult experiment

  28. Electron Affinity 沒有規律性

  29. Electron Affinity 2A 5A

  30. Electron Affinity • Example 5-4: Arrange these elements based on their electron affinities. • Al, Mg, Si, Na Si < Al < Na < Mg

  31. Ionic Radii 離子半徑 • Cations (positive ions) are always smallerthan their respective neutral atoms • Anions (negative ions) are always larger than their neutral atoms.

  32. 核電荷增加 半徑減少 Ionic Radii • Cation (positive ions) radii decrease from left to right across a period • Increasing nuclear charge attracts the electrons and decreases the radius • Anion (negative ions) radii decrease from left to right across a period • Increasing electron numbers in highly charged ions cause the electrons to repel and increase the ionic radius 原子序增加 半徑增加

  33. Ionic Radii Isoelectronic species: have the same number of electron

  34. Ionic Radii • Example 5-5: Arrange these elements based on their ionic radii. • Ga, K, Ca • Cl, Se, Br, S K1+ > Ca2+ > Ga3+ Cl1- < S2- < Br1- < Se2-

  35. Electronegativity (EN) 電負度；陰電性 • Electronegativity is a measure of the relative tendency of an atom to attract electrons to itself when chemically combined with another element. • Electronegativity is measured on the Pauling scale. • Fluorine is the most electronegative element. • Cesium (Cs) and francium (Fr) are the least electronegative elements. 電負性是指元素的原子在分子中吸引電子的能力的相對大小，電負性大，原子在分子中吸引電子的能力強，反之就弱。

  36. Electronegativity 電負度；陰電性 • 金屬元素的電負度較小，而非金屬元素的電負度較大 • Elements with high electronegativity (nonmetals) often gain electrons to form anions. • Elements with low electronegativity (metals) often lose electrons to form cations. • For the representative elements, electronegativities usually increase from left to right across periods and decrease from top to bottom within groups.

  37. 惰性氣體無電負度 • 同族由上而下遞減，同週期由左而右遞增

  38. Electronegativity

  39. Electronegativity • Example 6-11: Arrange these elements based on their electronegativity. • Se, Ge, Br, As • Be, Mg, Ca, Ba Ge < As < Se < Br Ba < Ca < Mg < Be

  40. Periodic Trends • It is important that you understand and know the periodic trends described in the previous sections. • They will be used extensively in Chapter 7 to understand and predict bonding patterns.

  41. 週期表各元素性質之變化趨勢

  42. Oxidation Numbers氧化數 • The transfer of electrons from one species to another are called oxidation-reduction reactions 氧化還原反應, redox reaction. • Oxidation numberor oxidation state of an element in a simple binary ionic compound is the number of electrons gained or lost by an atom of that element when it forms the compound. • 金屬與非金屬反應形成離子化合物，涉及一個或更多個電子從金屬(形成一個陽離子) 轉移到非金屬 (形成一個陰離子)。 • 涉及電子轉移的反應稱為氧化－還原反應 (oxidation-reduction reaction)。例如：金屬鎂與氧的反應。

  43. Oxidation Numbers氧化數 • Guidelines for assigning oxidation numbers. • The oxidation number of any free, uncombined element is zero. Such as: H2, O2, P4, S8 • The oxidation number of an element in a simple (monatomic) ion is the charge on the ion. • In the formula for any compound, the sum of the oxidation numbers of all elements in the compound is zero. • In a polyatomic ion, the sum of the oxidation numbers of the constituent elements is equal to the charge on the ion.

  44. Oxidation Numbers • Fluorine, F,has an oxidation number of –1 in its compounds. • Hydrogen, H, has an oxidation number of +1 unless it is combined with metals, where it has the oxidation number -1. • Examples – LiH, BaH2 • Oxygen usually has the oxidation number -2. Exceptions: • In peroxides O has oxidation number of –1. • Examples - H2O2, CaO2, Na2O2 • In OF2, O has oxidation number of +2.

  45. Oxidation Numbers • Use the periodic table to help with assigning oxidation numbers of other elements. • IA metals have oxidation numbers of +1. • IIA metals have oxidation numbers of +2. • IIIA metals have oxidation numbers of +3. • There are a few rare exceptions. • VA elements have oxidation numbers of –3 in binary compounds with H, metals or NH4+. • VIA elements below O have oxidation numbers of –2 in binary compounds with H, metals or NH4+. • Summary in Table 4-10(Table 5-4 p193).

  46. Oxidation Numbers

  47. Oxidation Numbers • Example5-5: Assign oxidation numbers to each element in the following compounds: • NaNO3 Na = +1 (Rule 8) O = -2 (Rule 7) N = +5 • Calculate using rule 3 • +1 + 3(-2) + x = 0 • x = +5 • K2Sn(OH)6 • K = +1 (Rule 8) • O = -2 (Rule 7) • H = +1 (Rule 6) • Sn = +5 • Calculate using rule 3 • 2(+1) + 6(-2) + 6(+1) + x = 0 • x = +5

  48. Oxidation Numbers • HCO3- • O = -2 (Rule 7) • H = +1 (Rule 6) • C = +4 • Calculate using rule 4. • +1 + 3(-2) + x = -1 • x = +4 • NO2- O = -2 (Rule 7) N = +3 • Calculate using rule 4. • 2(-2) + x = -1 • x = +3

  49. Example 5-6 Oxidation Numbers Determine the oxidation numbers of nitrogen in the following species (a) N2O4, (b)NH3, (c)HNO3, (d) NO3-, (e)N2 • (a) N2O4 • O = -2 • 2x +(-2)4 = 0 • x = +4 • N=+4 • (b) NH3 • H = +1 • x +(+1)3 = 0 • x = -3 • N=-3 • (c) HNO3 • H = +1 O=-2 • 1+x+(+2)3 = 0 • x = +5 • N=+5 • (d) NO3- • O = -2 • x +(-2)3 = -1 • x = +5 • N=+5 (b) N2 The oxidation number of any free element is zero Exercise 42

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