Chapter 2 Chemical Bonding PDF

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Chapter 2 Chemical Bonding The interaction between atoms that leads to a rearrangement of the electrons to a more stable state is what we define as chemical bonding. All atoms, except those of the noble gases*, readily engage in chemical bonding either with atoms of their own kind (in elements) or with atoms of a different kind (in compounds). *) This statement is not absolutely true. Some of the heavier noble gases such as xenon and krypton can be persuaded to form compounds with very strong oxidizers. Compounds of helium and neon are presently not known. Three types (extremes) of chemical bonds exist: –

Ionic bond - electrostatic forces hold together oppositely charged ions.

–

Covalent bond - two atoms sharing electron pair(s). Electrons are mutually attracted by the nuclei of adjacent atoms.

–

Metallic bond -is a complicated type of bond. At the simples level metal atoms are envisaged of being metal cations in a “sea” of electrons. The electrons are shared between all the ions simultaneously. A better description of metals is obtained from the band theory.

Ionic bond Ionic bonds form between oppositely charged ions, generally, but not always, between metals and nonmetals. Metals versus Nonmetals Metals are conductors of heat and electricity, have a lustre (shiny) appearance, are ductile (can be drawn in to wirers) and are malleable (can be hammered out into thin foils). With the exception of hydrogen, which is a non-metal, metals are located to the left in the periodic table. Non-metals are generally non conductors, do not have a metallic lustre and are brittle. They are located to the right in the periodic table. The elements near the “step ladder” have both, metallic and non-metallic properties and are called metalloids or semi metals.

1

2

3

4

5

6

7

1 1 H 3 Li 11 Na 19 K 37 Rb 55 Cs 87 Fr

2

4 Be 12 Mg 20 Ca 38 Sr 56 Ba 88 Ra

* **

57 La 89 Ac

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4

5

21 22 23 Sc Ti V 39 40 41 Y Zr Nb 57* 72 73 La Hf Ta 89** 104 105 Ac Rf Db 58 Ce 90 Th

59 Pr 91 Pa

60 Nd 92 U

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7

8

9

10

11

12

13

14

15

16

17

6 C 14 Si 32 Ge 50 Sn 82 Pb

7 N 15 P 33 As 51 Sb 83 Bi

8 O 16 S 34 Se 52 Te 84 Po

9 F 17 Cl 35 Br 53 I 85 At

69 Tm 101 Md

70 71 Yb Lu 102 103 No Lr

24 25 Cr Mn 42 43 Mo Tc 74 75 W Re 106 107 Sg Bh

26 Fe 44 Ru 76 Os 108 Hs

27 Co 45 Rh 77 Ir 109 Mt

28 Ni 46 Pd 78 Pt

29 Cu 47 Ag 79 Au

30 Zn 48 Cd 80 Hg

5 B 13 Al 31 Ga 49 In 81 Tl

61 Pm 93 Np

63 Eu 95 Am

64 Gd 96 Cm

65 Tb 97 Bk

66 Dy 98 Cf

67 Ho 99 Es

68 Er 100 Fm

62 Sm 94 Pu

To understand chemical bonding trends in three atomic properties must be considered: electron affinity, ionization energy and atomic radius

Atomic radius Atomic radius is difficult to define – there is a small probability of finding an electron at an infinite distance from the nucleus. Often atomic radius is defined in terms of internuclear distance in a compound. Methods of representing Size covalent radii: half of the internuclear distance in a homonuclear X-X single bond. ionic radii:

The ionic radius, rion, is the size of an ion (due to changes in effective nuclear charge; for cations, rion < rcov; whilst for anions, rion > rcov). These values are estimated from electron density maps determined from crystallographic data.

van der Waals Radii:

half of the distance of closest approach of two non-bonded atoms. The forces holding the atoms together are weak London Dispersion forces (more to come…).

18 2 He 10 Ne 18 Ar 36 Kr 54 Xe 86 Rn

Source: petrucci Periodic Trends in Atomic Radii Trends in atomic radii can be rationalized if we consider the shielding or screening effect of the core electrons.

Effective nuclear charge: Zeff = Z – S Z = atomic number, S = number of inner electrons that screen an outer electron The more electronic shells in an atom, the larger the atom. Atomic radius increases from top to bottom in a group. The atomic radius decreases from left to right through a period of elements.

Class exercise: Arrange the following atoms in order of increasing atomic radius: Mg, N, Rb, Be

Source Petrucci Ions Ions are charged atoms or groups of atoms •

Ionic Radius: Cations are smaller than the atoms from which they are formed. Isoelectronic cations have the same number of electrons in identical configurations. Na+ and Mg2+ both have closed shells (both have 10 electrons but the magnesium nucleus has more protons to attract the 10 electrons then the sodium nucleus). For isoelectronic cations, the more positive the ionic charge, the smaller the ionic radius.

•

Anions are larger than the atoms from which they are formed. For isoelectronic anions, the more negative the charge, the larger the ionic radius.

Source Petrucci

Class exercise: Arrange the following ions in order of increasing ionic radius: Al3+, Cl-, Na+, F-

Electron affinity and ionization energy: As we have seen an atom can gain one electron or more. This fundamental atomic property is known as electron affinity and is usually exothermic. H + e- → H- + energy An atom can also lose one electron (or more). This fundamental atomic property is called ionization or ionization energy. Atoms are stable species. It requires energy to release an electron from an atom so ionization is an endothermic process and ionization energies are positive H + energy → H+ + eIt is important to understand that electron gain and ionization are truly atomic properties. They are properties of isolated gas phase atom (not bulk phase atoms as in a chunk of copper). Metals tend to lose electrons and form cations →

Mg+(g) + e-

I1 = 738 kJ/mol

Mg+(g) →

Mg2+ (g)+ e-

I2 = 1451 kJ/mol

Mg(g)

I1 is called the first ionization energy, I2 the 2nd ionization energy. Trends of ionization energy within a group in the periodic table: As you go down a group (chemical family) the ionization energy decreases. It is not difficult to understand the rationale behind that. With every new row (period) in the periodic table a new electronic shell is being filled with electrons (atoms get bigger!) The further away the electron is from the nucleus, the smaller is its attractive interaction with the nucleus and the easier it is to remove it. Chemically speaking, the easier it is to remove the electron the greater the reactivity of the element toward an electron accepting element (an oxidizer). The table below lists the first ionization energies of the alkaline metals (group 1 in the periodic table) Element Lithium Sodium Potassium Rubidium Cesium

Radius in pm 152 186 227 248 265

I1 in kJ/mol 520.2 495.8 418.8 403.0 375.7

Trends of ionization energy within a period: As you go from the left to the right ionization energy increases. Electrons enter the same shell within a perioid. This leads to an increase in repulsion between the electrons in the outermost electronic shell. At the same time as the number of electrons increases so does the number of protons in the nucleus and the greater the positive charge of the nucleus the

greater the attraction towards the electrons. It is this increased attraction that dominates over electron repulsion. Thus ionization energies generally increase from the left to the right within a period. We will take a closer look at the exceptions later. This nucleus having a “tighter grip” on the outermost electrons also means that atomic radius decreases from the left to the right in a period. Strictly speaking this is only true for gas phase atoms. In the bulk phase elements show different bonding arrangements that might stabilize the atoms and make them less prone to react e.g. the reactivity of the 3rd row elements toward O2 decreases from Na to Si but hugely increases at P! Why? The table below lists ionization energies of the 3rd row elements in KJ/mol

I1 I2 I3 I4 I5 I6 I7

Na 496 4562

Mg 738 1451 7733

Al 578 1817 2745 11580

Si 787 1577 3232 4356 16090

P 1012 1903 2912 4957 6274 21270

S 1000 2251 3361 4564 7013 8496 27110

Cl 1251 2297 3822 5158 6542 9362 11020

Ar 1521 2666 3931 5771 7238 8781 12000

Consider the reactivity of the following elements Na and K as well as Na and Mg!

There are two more points worth noting in this table. First there is a huge increase in ionization energy after the last valence electron has vacated the valence shell and a core electron from the inner electronic shell is removed e.g. it takes 496 kJ /mol to remove one electron from sodium but it takes 4562 kJ/mol to remove a second electron which is why no compounds are known that contain Na2+ ions but essentially all compounds of sodium contain Na+ ions. Second, by the time you get to the middle of the 3rd period (e.g. Si) ionization energies become quite large and it is less feasible for atoms to lose electrons. Nonmetals which are located to the right in the periodic table do not lose but tend to gain electrons and form anions. We called this atomic property electron affinity and the following trend is observed within the periodic table: Electron affinity generally increases from the left to the right within a period and decreases from the top the bottom in a group. Chemically speaking this means the element most eager to gain an electron is the most reactive non metal element which is fluorine. Fluorine reacts with most elements in the periodic table. Class exercise: Arrange the following atoms in order of increasing ionization energy: Si, F, Na, O

Arrange the following atoms in order of increasing electron affinity: Cl, As, P, Ba

Octet Rule In taking a closer look at the number of electrons that metals lose and nonmetals gain it becomes apparent that many try to obtain the same electron count as their nearest noble gas, especially the elements close to the right or the left. So in compounds group 1 elements such as sodium or potassium always lose just one electron, never two or more. In doing so they obtain the same number of electrons as the nearest noble gas. Similarly in compounds with metals group 17 elements such as fluorine or bromine always gain just one electron never two or more to achieve a noble gas electronic configuration (more to come).

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1 2 +/1 H 1+ 2+ Li Be 1+ 2+ Na Mg 1+ K 1+ Rb 1+ Cs

2+ Ca 2+ Sr

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3+ Al 3+ Sc

4+ Ti

2,3,4, 5+

3, 6+

2,4, 7+

2,3, 6+

V

Cr

Mn

Fe

-1 to +7

2+ Co

2+ Ni

3,4,6, 8+

1,2+ Cu 1+ Ag

2+ Zn 2+ Cd

2+ 3,4+ 2,3,4+ 1,3+ 1,2+ Ba Re Os Ir Pt Au Hg Common oxidation states/oxidation numbers of selected elements

15

16 17

3N 3P

2O 2S

2- 1Se Br 2- 1Te I

Ga 2,4+

Sn 1,3+

2,4+

3,5+

Tl

Pb

Bi

Since metals tend to lose electrons and non-metals tend to gain electrons combining a metal and a nometal leads to a chemical reaction in which electrons are completely transferred from the metal to the non-metal resulting in metal cations and non-metal anions that feel a strong electrostatic attraction. This might be represented by the equation below, which uses Lewis dot formulas or Lewis structures including the number of valence electrons (electrons in the outermost electronic shell) as dots around the element symbol.

Li

+

F

Li

+

1F 1Cl

F

Compounds between metals and non-metals are called salts (e.g. sodium chloride, magnesium bromide) or ionic compounds (e.g. magnesium oxide, copper sulfide). Many important materials are ionic including ordinary table salt, ceramics, many ores and gypsum. There is a fundamental difference between an atom and an ion. Sodium ions cannot be isolated and placed in a jar on a shelf like sodium atoms which come in form of bars or chunks of metal under

mineral oil in a tin can. They are associated with a counter ion (chloride, fluoride, sulfate) to preserve overall electrical neutrality. A sodium ion is a stable species abundant in large quantity in the human body. Sodium metal is a reactive element that vigorously reacts with water to produce hydrogen and sodium hydroxide, a strong base. Class exercise Write balanced chemical equations for the reaction of:

sodium with bromine

magnesium with chlorine

aluminum with sulfur

strontium with fluorine

hydrogen with calcium

nitrogen with lithium

Remember that electrical neutrality must be preserved in the final product and that the charges of the ions cancel properly.

Energetic of ionic bonding As we have seen it requires 496 kJ/mol to remove electrons from sodium (first ionization energy of sodium). We get 349 kJ/mol back by giving the electrons to chlorine (electron gain is exothermic). But these numbers don’t explain why the reaction of sodium metal and chlorine gas to form sodium chloride is so exothermic! What is as of yet unaccounted for is the electrostatic attraction between the newly formed ions. Negative anions are attracted to positive cations. The result is an ionic bond and a three-dimensional crystal lattice of anions and cations. The binding energy might be calculated from the following experimental data (Born Haber Cycle:

Na+(g) + e- + Cl(g)

Na+(g) + Cl-(g) Na(g) + Cl(g)

+496 kJ

Na(g) + 0.5Cl2(g)

+122 kJ

Na(s) + 0.5Cl2(g)

+107 kJ

Start

-787 kJ

-411 kJ

NaCl(s) End

-349 kJ

Step 1 (Start) Sublime one mole of sodium atoms Na(s) → Na(g)

+107 kJ/mol

Step 2 Dissociate half a mole of chlorine molecules 0.5 Cl2(g) → Cl(g)

+122 kJ/mol

Step 3 Ionize one mole of sodium atoms Na(g) → Na+(g)

+496 kJ/mol

Step 4 Convert one mole of chlorine atoms to one mole of chlorine ions Cl(g) + e- → Cl-(g)

-349 kJ/mol

Step 5 Combine one mole of chloride ions with one mol of sodium ions to form one mole of sodium chloride Na+(g) + Cl-(g) → NaCl(s)

(lattice energy unknown)

Step 6 (End) Measure the heat of formation for the overall reaction Na(s) + 0.5 Cl2(g) → NaCl(s)

-411 kJ/mol

The energy change between reactants and products is independent of the path that the reaction takes (more to come in thermochemistry). The heat of the reaction that we obtain by directly combining the elements should thus be equal to the energy change if we take the reactants clockwise through the Born-Haber cycle because we end up with the same product in the same state. -411 kJ/mol = 107 kJ/mol + 122 kJ/mol + 496 kJ/mol +(-349 kJ/mol) +x kJ/mol x = -787 kJ/mol As you can see it is the hugely negative lattice energy that is largely responsible for the exothermic reaction between the elements sodium and chlorine that releases 411 kJ/mol. Relationship between Bonding and Properties Now why is an understanding of chemical bonding important? There is a relationship between structure and properties of compounds (materials!) If you know and understand the structure of a material, the bonding at the atomic level, you can understand and often predict its physical properties. This is why chemists have been so successful in engineering materials with properties tailored to specific applications. Ionic compounds have very largely negative lattice energies. This tells us that they are very stable with respect to the elements from which they have formed and explains why ionic compounds are usually hard materials (ceramics, “rock salt”). Another property of ionic compounds is their very high melting

and boiling points which is related directly to the high lattice energy. Although anions and cations can and do form ion pairs in solution and in the gas phase in the solid state they form crystalline solids which can be described by a three dimensional crystal lattice. The lattice energy is defined as the energy given off when oppositely charged isolated gaseous ions come together to form one mole of ionic solid. The greater the lattice energy the harder it is to vaporize an ionic compound. Alumina and magnesia have such high lattice energies that they make excellent materials for furnace bricks or crucibles. Zirconium oxide on top of having high melting and boiling points has exceptional fracture toughness along with very good chemical resistance. Another ceramic, zirconium nitride is a cement like refractory material. A coating of zirconium nitride is commercially applied on parts that are subject to high wear and tear, corrosive environments or to improve refraction. Ta4HfC5 is presently holding the record with the highest melting point of any compound at 4215 °C.

Class exercise Given below are calculated lattice energies as well as measured melting points, boiling points and aqueous solubility of selected ionic compounds from Handbook of Chemistry and Physical 90th Edition 2009. Rationalize the trends in lattice energies, melting and boiling points.

Compound LiF NaCl KBr RbI

Compound Na2O MgO Al2O3 ZrN

Lattice energy (kJ/mol) 1,049 787 691 632

Melting point (°C) 848 801 734 656

Boiling Point (°C) 1673 1465 1435 1300

Solubility (g/100mL H2O) 0.13 36 67 165

Lattice energy (kJ/mol) 2481 3795 15916 7633

Melting point (°C) 1134 2825 2054 2952

Boiling Point (°C) decomposition 3600 2977 _____

Solubility (g/100mL H2O) decomposition insoluble insoluble insoluble

Picture source: Wikipedia The sodium chloride structure in which each sodium ion is surrounded by six chloride ions and each chloride ion is surrounded by six sodium ions in octahedral geometry is adopted by many other compounds which have a stoichiometry of 1:1 and where anions and cations don’t differ to greatly in ionic radius.

Covalent bond For some elements it is not feasible to gain the number of electrons required to reach an octet. e.g. carbon would have to gain 4 electrons to have the same number of electrons as neon. Carbon (and many nonmetals) can share valence electrons so that they are surrounded by 8 electrons (hydrogen by 2 electrons). This sharing of electrons is called covalent bonding.

It is not immediately obvious why the sharing of electrons leads to a more stable state. Electrons after all are negatively charged particles and as such repel another. Let’s consider the simplest case, a dihydrogen molecule. Hydrogen possesses only one electron. To obtain the same number of electrons as its nearest noble gas (helium) it would have to gain one electron. In reactions with active metals this is exactly what hydrogen do...