Unit 2 Molecular and Ionic Compound Structure and Properties

1. Unit 2Molecular and Ionic Compound Structure and Properties Topic 2.1 Types of Chemical Bonds Text Reference Ch 8, sections 8.2, 8.3, 8.4; Ch 12 sections 12.3, 12.4

2. Understanding: Atoms or ions bond due to interactions between them, forming molecules. Learning Objective: Explain the relationship between the type of bonding and the properties of the elements participating in the bond. Essential Knowledge: Electronegativity values for the representative elements increase going from left to right across a period and decrease going down a group. These trends can be understood qualitatively through the electronic structure of the atoms, the shell model, and Coulomb’s law. Valence electrons shared between atoms of similar electronegativity constitute a nonpolar covalent bond. For example, bonds between carbon and hydrogen are effectively nonpolar even though carbon is slightly more electronegative than hydrogen.

3. Valence electrons shared between atoms of unequal electronegativity constitute a polar covalent bond. a. The atom with the higher electronegativity will develop a partial negative charge relative to the other atom in the bond. b. In single bonds, greater differences in electronegativity lead to greater bond dipoles. c. All polar bonds have some ionic character, and the difference between ionic and covalent bonding is not distinct but rather a continuum.

4. Chemical Bonds • A chemical bond is a strong attractive force between atoms or ions. • There are three basic types of bonds • Ionic: Electrostatic attraction between ions • Covalent: Sharing of electrons between atoms • Metallic: Free electrons hold metal atoms together

5. Octet Rule • When forming compounds, atoms tend to gain, lose, or share electrons until they are surrounded by eight valence electrons • An octet is a very stable electron configuration consisting of full s and p subshells in an atom. • The octet rule applies mostly to the representative elements that have s and p valence electrons. • Transition metals, having s and d valence electrons, do not follow the octet rule in achieving stability. • The duet rule applies to hydrogen which is stable when it shares two electrons for a full valence shell.

6. 8.2 Ionic Bonding • Between metals and nonmetals (except group 8A) • A metal readily gives up an electron (has a LOW ionization energy). • A nonmetal readily gains an electron (has a HIGH electron affinity). • Arrow(s) indicate the transfer of the electron(s).

7. Properties of Ionic Substances • Ions maintained in rigid, well-defined, 3-D structures • Usually crystalline • Ionic crystals can be cleaved along smooth, flat surfaces • Brittle with high melting points

8. Energetics of Ionic Bond Formation • As we saw in Topic 1.7, it takes energy to create a cation. For example, it takes 496 kJ/mol to remove electrons from sodium. (Ionization energy) • Energy is released by creating an anion. For example, we get 349 kJ/mol back by giving electrons to chlorine. (Electron affinity) • 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 = 147 kJ/mol • The positive energy change indicates the ions are not interacting with each other.

9. Yet, we see the reaction of forming ionic compounds from elements as very exothermic. Both light and heat are given off by the reaction.

10. 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. • Lattice energy is the energy required to completely separate a mole of a solid ionic compound into its gaseous ions.

11. 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 from elements an exothermic process.

12. The energy associated with electrostatic interactions is governed by Coulomb’s law: • Q1 and Q2 are the charges on the particles in Coulombs, with their signs; d is distance between their centers in meters; and  is a constant, 8.99 109 J-m/C2. • Lattice energy, then, increases with the charge on the ions. • It also increases with decreasing ion size.

13. Table 8.1 Lattice Energies for Some Ionic Compounds

14. Energetics of Ionic Bonding—Born–Haber Cycle The Born–Habercycle is an approach to analyzing reaction energies. It was named after and developed by the two German scientists Max Born and Fritz Haber. The cycle is concerned with the formation of an ionic compound from the reaction of a metal (often a Group I or Group II element) with a non-metal. • 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.

15. Electron Configuration of Ions of the s- and p-Block Elements. • Main group metals lose electrons, resulting in the electron configuration of the previous noble gas. • Nonmetals gain electrons, resulting in the electron configuration of the nearest noble gas.

16. Transition Metal Ions • Transition metals do Not follow the Octet rule. • Transition metals lose the Valence electrons First, Then lose the d- electrons necessary for the given ion charge.

17. 8.3 Covalent Bonding • In covalent bonds, atoms share electrons. • There are several electrostatic interactions in these bonds: • Attractions between electrons and nuclei • Repulsions between electrons • Repulsions between nuclei • For a bond to form, the attractions must be greater than the repulsions.

18. Lewis Structures • The Lewis symbol for an element consists of its chemical symbol plus a dot for each valence electron. • Sharing electrons to make covalent bonds can be demonstrated using Lewis structures. • We start by trying to give each atom the same number of electrons as the nearest noble gas by sharing electrons. • The simplest examples are for hydrogen, H2, and chlorine, C l2, shown below.

19. Number of Bonds for Nonmetals • The group number is the number of valence electrons. • To get an octet like the nearest noble gas, in the simplest covalent molecules for nonmetals, the number of bonds needed will be the same as the electrons needed to complete the octet. • Bonds between atoms that share one pair of electrons are called single bonds.

20. Electrons on Lewis Structures • Lone pairs or nonbonding pairs: electrons located on only one atom in a Lewis structure • Bonding pairs: shared electrons in a Lewis structure; they can be represented by two dots orone line, not both.

21. Multiple Bonds • Bonds between atoms that share two pairs of electrons in order to achieve an octet of electrons for each atom are called double bonds. • Bonds between atoms that share three pairs of electrons in order to achieve an octet of electrons for each atom are called triple bonds.

22. 8.4 Bond Polarity and Electronegativity • The electrons in a covalent bond are not always shared equally. • Bond polarity is a measure of how equally or unequally the electrons in a covalent bond are shared. • In a nonpolar covalent bond, the electrons are shared equally. • In a polar covalent bond, one of the atoms attracts electrons to itself with a greater force.

23. Electronegativity • Electronegativityis the ability of an atom in a molecule to attract electrons to itself. • On the periodic table, electronegativity generally increases as you go: • from left to right across a period. • from the bottom to the top of a group.

24. Electronegativity and Bond Polarity • Differences in electronegativity between two atoms can be used to gauge the polarity of the bond • In elemental fluorine, F2, the atoms pull electrons equally. The bond is a nonpolar covalent bond. Electronegativity difference for F2: • In HF, fluorine 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, making it a polar covalent bond. Electronegativity difference for HF:

25. Electrons tend to spend more time around the more electronegative atom, resulting in greater electron density. The result is a partial negative charge (nota complete transfer of charge). It is represented by δ- (“delta minus”). • The other atom has less electron density and a partial positive charge, or δ+ (“delta plus”). • See Fig 8.9 in text. • If the electronegativity difference between atoms is more than 2.0, the bond is considered to be ionic, i.e., the electrons are transferred between atoms, not shared. Electronegativity difference for LiF:

26. The greater the difference in electronegativity, the more polar is the bond. Table 8.2 Bond Lengths, Electronegativity Differences, and Dipole Moments of the Hydrogen Halides

27. Dipole Moments • Polarity helps to determine the properties of molecules. • Polar molecules align themselves with respect to each other: negative end of one molecule is attracted to the positive end of another (for example, as in water). • Polar molecules are likewise attracted to ions (for example, when NaCl dissolves in water). • How do we quantify the polarity of a molecule?

28. When two equal, but opposite, charges are separated by a distance, a dipole forms. • The quantitative measure of the magnitude of a dipole is called its dipole moment. • A dipole moment, μ,produced by two equal but opposite charges (Q+ and Q-) separated by a distance, r, is calculated: • It is measured in debyes (D), a unit that equals 3.34 x 10-30 coulomb-meters(C-m)

29. Dipole moment calculations:

30. Comparing Ionic and Covalent Bonding • Simplest approach: Metal + nonmetal is ionic; nonmetal + nonmetal is covalent. However, there is a continuous spectrum between the extremes of ionic and covalent bonding. • When covalent bonding is dominant, we expect compounds to exist as molecules: low melting and boiling points, nonelectrolyte behavior when dissolved in water. • When ionic bonding is dominant, we expect compounds to be brittle solids, with high melting points, extended lattice structures, and strong electrolyte behavior when dissolved in water • Properties of compounds are often best for determining type of bond. For example, lower melting points indicate covalent bonding.

31. However, there are many exceptions. Using electronegativity differences does not always take into account changes in bonding that accompany changes in the oxidation state of the metal. • Example: SnCl4 is a molecular compound, not ionic as its components would seem to make it (metal with a nonmetal). • Example: manganese (II) oxide, MnO, compared to manganese (VII) oxide, Mn2O7 • Metals in high positive oxidation states (+4 or higher) show significant covalency in bonds with nonmetals to form molecules such as Mn2O7 or polyatomic ions, such as MnO4- and CrO42-.

32. 12.3/12.4 Metallic Solids/Metallic Bonding • Consist entirely of metal atoms packed closely together. • Atoms are not held together by covalent bonds as there are too few valence electrons to form bonds. • Rather, valence electrons are delocalized (that is, spread) throughout the entire solid to form metallic bonds between atoms. • Metals can be visualized as an array of positively charged ions immersed in a “sea” of delocalized electrons.

33. Properties of metals • The electrons are confined to the metal by electrostatic attractions, yet are mobile as no individual electron is confined to any particular metal ion. This allows metals to conduct electricity (a flow of electrons) and heat (thermal conductivity). • Metals are lustrous, malleable and ductile. These properties indicate that metal atoms are capable of slipping past one another.

34. An alloy is a material that contains more than one element and has the characteristic properties of a metal. • It is an important means employed to change the properties of certain metals. • Substitutional alloys form when two metallic components have similar atomic radii and chemical-bonding characteristics. • Interstitial alloys form when the solute atoms have a much smaller radius than the solvent atoms. Alloys

35. In a heterogeneous alloy, the components are not dispersed equally. • Intermetallic compounds are compounds rather than mixtures. The atoms in an intermetallic compound are ordered rather than randomly distributed. These features make them attractive for high-temperature applications, though they can be very brittle.