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This chapter explores the three basic types of chemical bonds - ionic, covalent, and metallic. It also discusses Lewis symbols, the octet rule, energetics of bonding, electron configurations of ions, and covalent bonding.
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8.1 Chemical Bonds, Lewis Symbols, Octet Rule • A chemical bond is a strong attractive force that exists between atoms • Three basic types of bonds: • Ionic • Electrostatic attraction between ions • Ex: Mg2+ + O2- MgO • Covalent • Sharing of electrons between atoms • Ex: Br2, S8, C12H22O11 • Metallic • Metal atoms bonded to several neighboring atoms
Lewis Symbols • The electrons involved in chemical bonding are the valence electrons, which usually reside in the outermost occupied shell of an atom. • A Lewis symbol is a simple way of showing these electrons. • Ex: Sulfur [Ne]3s23p4 • See Table 8.1 Octet Rule • Atoms gain, lose or share electrons until they are surrounded by eight valence electrons. • An octet consists of full s and p subshells in an atom. • An octet is a very stable configuration of electrons.
8.2 Ionic Bonding • The chemical reaction between Na(s) and Cl2(g) to produce NaCl(s) is very exothermic: [8.1] • Formation indicates an electron has been lost by sodium and gained by chlorine: • Ionization energy indicates how easily an electron can be lost by one atom, and electron affinity measures the ease with which another atom gains an electron.
Energetics of Ionic Bonding • As we saw in the last chapter, it takes 495 kJ/mol to remove electrons from sodium. • We get 349 kJ/mol back by giving electrons to chlorine. • If the transfer of an electron from Na(g) to Cl(g) were the only factor in forming an ionic bond, the process would rarely be exothermic. 496 -349 = 146 kJ/mol • The positive energy change indicates the ions are not interacting with each other.
Yet, we see the reaction as very exothermic. Both light and heat are given off by the reaction.
The principal reason that ionic compounds are stable is the electrostatic attraction between ions of opposite charge. • This attraction draws ions together, releasing energy , and causing the ions to form a stable, solid array, or lattice. • A measure of the stabilization that occurs is given by the lattice energy.
Q1Q2 d Eel = • Lattice energy is the energy required to completely separate a mole of a solid ionic compound into its gaseous ions. • For NaCl, ΔHlattice= +788 kJ/mol • This process is highly endothermic. The reverse process – the formation of NaCl – is highly exothermic: ΔH = -788 kJ/mol • The energy released by the attraction of ions of unlike charge more than makes up for the endothermic nature of ionization energies, making the formation of ionic compounds an exothermic process. • The energy associated with electrostatic interactions is governed by Coulomb’s law:
Q1Q2 d Eel = • The energy associated with electrostatic interactions is governed by Coulomb’s law: • Lattice energy, then, increases with the charge on the ions. • It also increases with decreasing ion size.
By accounting for all three energies (ionization energy, electron affinity, and lattice energy), we can get a good idea of the energetics involved in such a process.
Electron Configurations of Ions of the Representative Elements • The energetics of ionic bond formation explains why so many representative elements tend to have noble-gas electron configurations. • Na, for example, loses its one valence electron to form Na+, which has the noble gas configuration of Ne. • We never find Na2+ ions because the second electron removed would have to come from the inner shell of the Na atom, which would require a large amount if energy. • The increase in lattice energy is not enough to compensate for the energy needed to remove an inner shell electron.
Similarly, the addition of electrons to nonmetals is exothermic or slightly • endothermic as long as the electrons are being added to the valence shell. • The Cl atom easily adds an electron to form Cl-, which has the electron configuration of Ar. • To form the Cl-2 ion, the second electron would have to be added to the next higher shell of the Cl atom, which is energetically unfavorable.
Transition Metal Ions • The lattice energies of ionic compounds are generally large enough to compensate for the loss of only up to 3 electrons from atoms. • Most transition metals have more than three electrons beyond a noble gas core. • Silver, for example, has a [Kr]4d105s1 electron configuration. In forming Ag+, the 5s electron is lost leaving a full 4d subshell. • In forming ions, transition metals lose the valence shell s electrons first, then as many d electrons as necessary to reach the charge on the ion. • Ex: Fe2+ and Fe3+
8.3 Covalent Bonding • In these bonds atoms share electrons. • There are several electrostatic interactions in these bonds: • Attractions between electrons and nuclei • Repulsions between electrons • Repulsions between nuclei
Lewis Structures Lewis structures are representations of molecules showing all electrons, bonding and nonbonding.
Multiple Bonds • Single bonds share one pair of electrons. • Octets can also be obtained by sharing two pairs of electrons to form a double bond, or three pairs of electrons to form a triple bond. • As a general rule, the distance between bonded atoms decreases as the number of shared electron pairs increases
8.4 Bond Polarity and Electronegativity • The concept of bond polarity helps describe the sharing of electrons between atoms. • A nonpolar bond is one in which the electrons are shared equally between two atoms, as in Cl2 and N2. • In a polar covalent bond, one of the atoms exerts a greater attraction for the bonding electrons than the other. • If the difference in relative ability to attract electrons is large enough, an ionic bond is formed.
Electronegativity • The ability of an atom in a molecule to attract electrons to itself. • The greater an atom’s electronegativity, the greater its ability to attract electrons to itself. • Electronegativity is related to ionization energy and electron affinity. An atom with very negative electron affinity and a high ionization energy will both attract electrons from other atoms and resist having its electrons attracted away. • On the periodic table, electronegativity increases as you go from left to right across a row and from the bottom to the top of a column.
Electronegativity and Bond Polarity • When electrons are shared equally between two atoms, the bond is nonpolar, as in diatomic molecules. • Although atoms often form compounds by sharing electrons, the electrons are not always shared equally. • In HF, for example, fluorine has a higher electronegativity and pulls harder on the electrons it shares with hydrogen than hydrogen does. • Therefore, the fluorine end of the molecule has more electron density than the hydrogen end. • We can represent the charge distribution as • The δ+ and δ- (read “delta plus”, delta minus) symbolize the partial charges.
Dipole Moments • When two atoms share electrons unequally, a bond dipole results. • The dipole moment, , produced by two equal but opposite charges separated by a distance, r, is calculated: = Qr • It is measured in debyes (D).
Bond Length and Dipole Moments • The greater the difference in electronegativity, the more polar is the bond. • The bond length increases as the electronegativity differences get smaller.
Bond Types and Nomenclature • Ionic compounds (metal + nonmetal) are given names based on their component ions, to include the charge on the cation if that is variable (as in transition metals). • Molecular compounds use prefixes to indicate the number of atoms present. • The dividing line between the two ways to name compounds is not always clear. Ex: TiO2 is called titanium(IV) oxide but more commonly titanium dioxide. • Many compounds of metals with high oxidation states (usually above 3+) have properties more similar to covalent compound.
