chemical periodicity 化學週期性

55Chemical PeriodicityChemical Periodicity

化學週期性化學週期性

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).

Chapter GoalsChapter Goals5-1 More About the Periodic TablePeriodic 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 Periodicity5-7 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

2

More About the Periodic TableMore 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

3

More About the Periodic TableMore 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.

4

More About the Periodic TableMore 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.

5

First transition series (4s and 3d orbital occupy): 21Sc through 30ZnSecond transition series (5s and 4d orbital occupy): 39Sc through 48CdThird transition series (6s and 5d orbital occupy): 52La and 72Hf to 80HgFourth transition series (6s and 5d orbital occupy): 89Ac and 104Rf to 112

More About the Periodic TableMore About the Periodic Table• f - transition metals

– Sometimes called inner transition metals.

– Electrons are being added to f orbitals.

6

Outermost electrons have the greatest influence on the chemical properties of elements.

• Adding an electron to an s or p orbital dramatic change in the physical and chemical properties• Adding an electron to a d or f orbital much smaller effect on properties

Periodic PropertiesPeriodic Propertiesof the Elementsof 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

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8

Atomic RadiiAtomic Radii原子半徑原子半徑

relative sizesA

tom

ic r

ad

ii in

cre

ase

Atomic radii decrease Atomic radii decrease

Å = 10-10m

同週期典型元素大致隨原子序增加 →原子半徑愈小

同族元素原子序增加原子半徑愈大

Atomic RadiiAtomic 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

9

有效核電荷有效核電荷

鎂的有效核電荷約為 12 – 10 = 2

原子核質子數－內層電子數總和

•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。

11

凡得瓦半徑為兩個相鄰非鍵結原子間距離的一半

Atomic RadiiAtomic 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

6AO < S < Se < Te

Period 3Cl < S < P < Si

F < S < As < Ga

F < Cl < K < Cs 12

Atom

ic ra

dii i

ncre

ase

Atomic radii decrease

Ionization Energy Ionization Energy 游離能游離能• 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-

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Mg(g) + 738kJ/mol Mg+ + e-

Ionization EnergyIonization 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-

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Mg+ + 1451 kJ/mol Mg2+ + e-

•Atoms can have 3rd (IE3), 4th (IE4), etc. ionization energies.

Ionization EnergyIonization EnergyPeriodic trends for Ionization Energy:

1. IE2 > IE1

It always takes more energy to remove a second electron from an ion than from a neutral atom.

2. 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.

3. IE1 generally decreases moving down a family.IE1 for Li > IE1 for Na, etc.

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Dec

reas

eD

ecre

ase

increase increase 隨原子序增大漸增趨勢

隨原子序增大而遞減

同一族元素的游離能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。

游離能的週期性游離能的週期性 •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

游離能的規則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

First Ionization Energies First Ionization Energies of Some Elementsof Some Elements

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Elements with low ionization energies (IE) easily lose to form cations (positive charge)

Ionization EnergyIonization 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 < BeNa < Al < P < ClB < Be < O < NNa < Al < Mg < Si

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Dec

reas

e

increase

Ionization EnergyIonization 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

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Group and element

IANa

IIAMg

IIIAAl

IVASi

IE1 (kJ/mol) 496 738 578 786IE2 (kJ/mol) 4562 1451 1817 1577IE3 (kJ/mol) 6912 7733 2745 3232IE4 (kJ/mol) 9540 10,550 11,580 4356

Na: 1s22s22p63s1 Al: 1s22s22p63s23p1

Ionization EnergyIonization 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+.

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Ionization EnergyIonization Energy• Example 5-3: What charge ion would be

expected for an element that has these ionization energies?

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IE1 (kJ/mol) 1680

IE2 (kJ/mol) 3370

IE3 (kJ/mol) 6050

IE4 (kJ/mol) 8410

IE5 (kJ/mol) 11020

IE6 (kJ/mol) 15160

IE7 (kJ/mol) 17870

IE8 (kJ/mol) 92040

Notice that the largest increase in ionization energies occurs between IE7 and IE8. Thus this element would form a 1- ion

Electron AffinityElectron Affinity 電子親和力電子親和力• 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 對於放熱愈多放熱愈多的反應，我們稱其電子親和力愈大電子親和力愈大，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:

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atom(g) + e- + EA ion-(g)

Electron AffinityElectron Affinity

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Mg(g) + e- + 231 kJ/mol Mg-(g)

EA = +231 kJ/molBr(g) + e- Br-

(g) + 323 kJ/molEA = -323 kJ/mol

Two examples of electron affinity values:

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

電子親和力大

電子親和力的性質電子親和力的性質•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


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• 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


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沒沒有有規規律律性性


30

2A 5A


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

Si < Al < Na < Mg

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Ionic Radii Ionic Radii 離子半徑離子半徑

32

• Cations (positive ions) are always smaller than their respective neutral atoms

• Anions (negative ions) are always larger than their neutral atoms.

Ionic RadiiIonic 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

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原子序增加

半徑增加

核電荷增加

半徑減少

Ionic RadiiIonic Radii

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Isoelectronic species: have the same number of electron

Ionic RadiiIonic 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-

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Electronegativity (EN) Electronegativity (EN) 電負度；陰電性電負度；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.

36

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。

Electronegativity Electronegativity 電負度；陰電性電負度；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

38

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

ElectronegativityElectronegativity

39

ElectronegativityElectronegativity• Example 6-11: Arrange these elements based on

their electronegativity.– Se, Ge, Br, As– Be, Mg, Ca, Ba

Ge < As < Se < BrBa < Ca < Mg < Be

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Periodic TrendsPeriodic 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週期表各元素性質之變化趨勢

Oxidation NumbersOxidation Numbers 氧化數氧化數

43

• The transfer of electrons from one species to another are called oxidation-reduction reactions 氧化還原反應 , redox reaction.

• Oxidation number or 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。

Oxidation NumbersOxidation Numbers 氧化數氧化數• Guidelines for assigning oxidation numbers.1. The oxidation number of any free, uncombined

element is zero. Such as: H2, O2, P4, S82. The oxidation number of an element in a simple

(monatomic) ion is the charge on the ion.3. In the formula for any compound, the sum of the

oxidation numbers of all elements in the compound is zero.

4. In a polyatomic ion, the sum of the oxidation numbers of the constituent elements is equal to the charge on the ion.

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Oxidation NumbersOxidation Numbers

5. Fluorine, F, has an oxidation number of –1 in its compounds.

6. Hydrogen, H, has an oxidation number of +1 unless it is combined with metals, where it has the oxidation number -1.

–Examples – LiH, BaH27. 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.

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Oxidation NumbersOxidation Numbers8. Use the periodic table to help with assigning

oxidation numbers of other elements.a.IA metals have oxidation numbers of +1.b.IIA metals have oxidation numbers of +2.c. IIIA metals have oxidation numbers of +3.

•There are a few rare exceptions.d.VA elements have oxidation numbers of –3 in

binary compounds with H, metals or NH4+.

e.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).

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Oxidation NumbersOxidation Numbers

47

Oxidation NumbersOxidation 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

48

•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

Oxidation NumbersOxidation Numbers• NO2

-

O = -2 (Rule 7)N = +3

– Calculate using rule 4.– 2(-2) + x = -1– x = +3

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• HCO3-

O = -2 (Rule 7)H = +1 (Rule 6)C = +4

– Calculate using rule 4.– +1 + 3(-2) + x = -1– x = +4

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Example 5-6 Oxidation NumbersDetermine the oxidation numbers of nitrogen in the following species (a) N2O4, (b)NH3, (c)HNO3, (d) NO3

-, (e)N2

Example 5-6 Oxidation NumbersDetermine the oxidation numbers of nitrogen in the following species (a) N2O4, (b)NH3, (c)HNO3, (d) NO3

-, (e)N2

Exercise 42

(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

Chemical Reactions & PeriodicityChemical Reactions & Periodicity• In the next sections periodicity will be applied to

the chemical reactions of hydrogen, oxygen, and their compounds.

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Hydrogen and the HydridesHydrogen and the Hydrides

• Element hydrogen is colorless, ordorless, tasteless diatomic gas with the lowest molecular weight and density of any known substance

• Hydrogen gas, H2, can be made in the laboratory by– Displacement (and redox) reaction 3Fe(s)+4H2O(g)Fe3O4(s)+4H2(g)

Zn(s)+2HCl(aq)ZnCl2(aq)+H2(g)– Electrolysis of water 2H2O(l) 2H2(g)+O2(g)– Water gas reaction C(s)+H2O(g)CO(g)+H2(g)

– Steam cracking CH4 (g)+H2O(g) CO(g)+3H2(g)

蒸氣裂解52

In cock steam Water gas 水媒氣反應可當燃料

Nisynthesis gas•To produce organic chemicals like Methanol...

Hydrogen

• Hydrogen reacts with metals and other nonmetals to form binary compound called hydrides 氫化物 . – Ionic hydrides — all basic

• Hydride ions, H- , formed when hydrogen gains one electron per atom from an active metal

• All basic because the react with water to form hydroxide ions. LiH(s) +H2O(l) LiOH(s)+H2(g)

– Molecular hydrides• Hydrogen share s electrons with an atom of another nonmetal• Many are acidic; their aqueous contain hydrogen cation (H+),

such as HF, HCl, HBr, HI, H2S, H2Se, H2Te– The ionic or molecular character of the binary compounds

of hydrogen depends on the position of the other elements in the periodic table

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Hydrogen and the HydridesHydrogen and the Hydrides

54

2Li(l) + H2(g) 2LiH(s)Ca(l) + H2(g) CaH2(s)

Ionic hydrides

Molecular hydrides H2(g) + F2(g) 2HF(g) hydrogen fluoride2H2(g) + O2(g) 2H2O(g)

N2(g) + 3H2(g) 2NH3(g)

1A: 2M(l) + H2(g) 2(M+,H-)(s)2A: M(l) + H2(g) (M2+,2H-)(s)

Reactions of Hydrogen andReactions of Hydrogen andthe Hydridesthe Hydrides

• The ionic hydrides produced in the two previous reactions are basic.– The H- reacts with water to produce H2 and OH-.

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H- + H2O H2 + OH-

•For example, the reaction of LiH with water proceeds in this fashion.

LiH(s) + H2O(l) H2(g) + OH-(aq) + Li+

(aq)


• Hydrogen reacts with nonmetals to produce covalent binary compounds.

• One example is the haloacids 鹵化酸類 produced by the reaction of hydrogen with the halogens 鹵素 .

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H2 + F2 2 HFH2 + Br2 2 HBr

H2 + X2 2 HX•For example, the reactions of F2 and Br2 with H2 are:


• Combustion of H2: highly exothermic 2 H2(g) + O2(g) 2 H2O(l) + energy

• Hydrogen reacts with oxygen and other VIA elements to produce several common binary covalent compounds.– Examples of this reaction include the production of

H2O, H2S, H2Se, H2Te.

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2 H2 + O2 2 H2O8 H2 + S8 8 H2S


• The hydrides of Group VIIA and VIA hydrides are acidic.

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HCl H+(aq)+ Cl-

(aq) (a strong acid)

H2S H+(aq)+ HS-

(aq) (a weak acid)


• There is an important periodic trend evident in the ionic or covalent character of hydrides.

1.1.Metal hydridesMetal hydrides are ionic compounds and form basic aqueous solutions.

2.2.Nonmetal hydridesNonmetal hydrides are covalent compounds and form acidic aqueous solutions.

59

60

Example 5-7, 5-8 Predicting Products of ReactionsPredict the products of the reactions involving the reactions shown. Write a balanced formula unit equation for each. and predict the ionic or molecular character of the products.(a) H2(g) + I2(g) (b) K(l) +H2(g) (c) NaOH(s) +H2O (l)(excess)

(a) H2(g) + I2(g) 2HI(g)

Molecular (b) K(l) +H2(g) 2KH(s)

Ionic (c) NaH(s) +H2O (l) NaOH(aq) +H2(g) Ionic Molecular

Example 5-7, 5-8 Predicting Products of ReactionsPredict the products of the reactions involving the reactions shown. Write a balanced formula unit equation for each. and predict the ionic or molecular character of the products.(a) H2(g) + I2(g) (b) K(l) +H2(g) (c) NaOH(s) +H2O (l)(excess)

(a) H2(g) + I2(g) 2HI(g)

Molecular (b) K(l) +H2(g) 2KH(s)

Ionic (c) NaH(s) +H2O (l) NaOH(aq) +H2(g) Ionic Molecular

Exercise 54,55

Oxygen and the OxidesOxygen and the Oxides

• Joseph Priestley discovered oxygen in 1774 using this reaction (decomposition) :

61

2 KClO3 (s) 2 KCl (s) + 3 O2(g)

•A common laboratory preparation method for oxygen is:

•Commercially, oxygen is obtained from the fractional distillation of liquid air.

Oxygen and Ozone

2 HgO(s) 2 Hg() + O2(g)

•O2•an odorless and colorless gas.•make up 21% by volume of dry air•only slightly soluble in water

Oxygen and the OxidesOxygen and the Oxides• Ozone (O3) is an allotropic form 同素異構物 of

oxygen which has two resonance structures.

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2 O3(g) 3 O2(g)

in presence of UV

• Ozone (O3) is an unstable, pale blue gas, its density is about 1.5 times of O2. it is a very strong oxidizing agent.

•Ozone is an excellent UV light absorber in the earth’s atmosphere.

Reactions of Oxygen andReactions of Oxygen andthe Oxidesthe Oxides

•Oxygen is an extremely reactive element.– O2 reacts with most metals to produce normal oxides

having an oxidation number of –2.

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4 Li(s) + O2(g) 2 Li2O(s) lithium oxide

2 Na(s) + O2(g) Na2O2(s) sodium peroxide

However, oxygen reacts with sodium to produce a peroxide having an oxidation number of –1.


• Oxygen reacts with K, Rb, and Cs to produce superoxides having an oxidation number of -1/2.

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K(s) + O2(g) KO2(s) potassium superoxide

2 M(s) + O2(g) 2 MO(s)

2 Sr(s) + O2(g) 2 SrO(s)

Oxygen reacts with IIA metals to give normal oxides.

Class Contains Ions

Oxidation No. of Oxygen

Normal oxides O2- -2

Peroxides O22- -1

superoxides O2- -1/2

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4Li(s) + O2(g) 2Li2O(s) lithium oxide

2Na(s) + O2(g) 2Na2O2(g) sodium peroxideK(s) + O2(g) KO2 (s) potassium superoxide


• At high oxygen pressures the IIA metals can form peroxides.

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Ca(s) + O2(g) CaO2(s)

2 Mn(s) + O2(g) 2 MnO(s)

4 Mn(s) + 3 O2(g) 2 Mn2O3(s)

Metals that have variable oxidation states, such as the d-transition metals, can form variable oxides.

For example, in limited oxygen:

In excess oxygen:

+0 +2

+0 +3


• Oxygen reacts with nonmetals to form covalent nonmetal oxides.

• For example, carbon reactions with oxygen:– In limited oxygen

– In excess oxygen C(s) + O2(g) CO2(g)

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2 C(s) + O2(g) 2 CO(g)

+2

+4


• Phosphorous reacts similarly to carbon forming two different oxides depending on the oxygen amounts:– In limited oxygen

P4(s) + 3 O2(g) P4O6(s)– In excess oxygen

P4(s) + 5 O2(g) P4O10(s)

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+3

+5

+0

+0

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Purple: amphoteric oxide 兩性氧化物is one that shows some acidic and some basic properties


• Similarly to the nonmetal hydrides, nonmetal oxides are acidicacidic.– Sometimes nonmetal oxides are called acidic anhydrides

(acidic oxides).• Form acid with no change in oxidation state of nonmetal.

– They react with water to produce ternary acids. 三元酸• For example: CO2(g) + H2O () H2CO3(aq)

Cl2O7(s) + H2O () 2 HClO4(aq)

As2O5(s) + 6 H2O() 4 H3AsO4(aq)

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+4 +4

+7 +7

+5 +5


• Similarly to the hydrides, metal oxides are basicbasic.– These are called basic anhydrides.– They react with water to produce ionic metal

hydroxides (bases) Li2O(s) + H2O() 2 LiOH(aq)

CaO(s) + H2O () Ca(OH)2(aq)

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Metal oxides are usually ionicionic and basicbasic. Nonmetal oxides are usually covalentcovalent and acidicacidic. An important periodic trend.


• Nonmetal oxides react with metal oxides to produce salts.

Li2O(s) + SO2(g) Li2SO3(s)

Cl2O7(s) + MgO(s) Mg(ClO4)2(s)

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+1 +4 +1+4

Combustion ReactionsCombustion Reactions• Combustion reactions are exothermic redox

reactions – Some of them are extremely exothermic.

• One example of extremely exothermic reactions is the combustion of hydrocarbons.– Examples are butane and pentane combustion. 2 C4H10(g) + 13 O2(g) 8 CO2(g) + 10 H2O(g)

C5H12(g) + 8 O2(g) 5 CO2(g) + 6 H2O(g)

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-2 -2+0

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Example 6-8, 6-9 Acidic Character of OxidesArrange the following oxides in order of increasing molecular (acidic) character: SO3, Cl2O7, CaO, PbO2 and arrange these oxides in order of increasing basicity.

increasing nonmetallic characterCa < Pb < S < Cl2A 4A 6A 7A

increasing molecular characterCaO < PbO2 < SO3 < Cl2O7

increasing basic character Cl2O7 < SO3< PbO2 <CaO

Example 6-8, 6-9 Acidic Character of OxidesArrange the following oxides in order of increasing molecular (acidic) character: SO3, Cl2O7, CaO, PbO2 and arrange these oxides in order of increasing basicity.

increasing nonmetallic characterCa < Pb < S < Cl2A 4A 6A 7A

increasing molecular characterCaO < PbO2 < SO3 < Cl2O7

increasing basic character Cl2O7 < SO3< PbO2 <CaO

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Example 6-10 Predicting Reaction ProductsPredict the products of the following pairs of reactants.Write a balanced equation for each reaction.(a) Cl2O7(l) + H2O(l)(b) As4(s)+O2(g)(excess) (c) Mg (s)+O2(g)

(a) Cl2O7(l) + H2O(l) 2HClO4(aq) (b) As4(s)+5O2(g) 2As2O5(s) (c) 2Mg (s)+O2(g) 2MgO2(s)

Example 6-10 Predicting Reaction ProductsPredict the products of the following pairs of reactants.Write a balanced equation for each reaction.(a) Cl2O7(l) + H2O(l)(b) As4(s)+O2(g)(excess) (c) Mg (s)+O2(g)

(a) Cl2O7(l) + H2O(l) 2HClO4(aq) (b) As4(s)+5O2(g) 2As2O5(s) (c) 2Mg (s)+O2(g) 2MgO2(s)

Fossil Fuel ContaminantsFossil Fuel Contaminants• When fossil fuels are burned, they frequently have

contaminants in them.• Sulfur contaminants in coal are a major source of air

pollution.– Sulfur combusts in air. S8(g) + 8 O2(g) 8 SO2(g) most harmful pollutant

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2 SO2(g) + O2(g) 2 SO3(g)

SO3(g) + H2O() H2SO4(aq)

Next, a slow air oxidation of sulfur dioxide occurs.

Sulfur trioxide is a nonmetal oxide, i.e. an acid anhydride.

Main cause of acid rain

Fossil Fuel ContaminantsFossil Fuel Contaminants• Nitrogen from air can also be a source of significant

air pollution.• This combustion reaction occurs in a car’s cylinders

during combustion of gasoline. N2(g) + O2(g) 2 NO(g)

• After the engine exhaust is released, a slow oxidation of NO in air occurs.

2 NO(g) + O2(g) 2 NO2(g)

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Fossil Fuel ContaminantsFossil Fuel Contaminants

• NO2 is the haze that we call smog 煙霧 (=smoke and fog)

– Causes a brown haze in air.• NO2 is also an acid anhydride.

– It reacts with water to form acid rain and, unfortunately, the NO is recycled to form more acid rain.

3 NO2(g) + H2O() 2 HNO3(aq) + NO(g)

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8080

Dec

reas

e

increase

chemical periodicity 化學週期性

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