Molecular and Ionic Compound Structure and Properties

1. Molecular and Ionic Compound Structure and Properties Unit 2

2. 2.1 Types of Chemical Bonds • Atoms or ions bond due to interactions between them, forming molecules or formula units • Ionic bonding occurs when oppositely charged ions are close enough to bond • They form a crystal lattice structure whose strength is relative to Coulomb’s law

3. Nonpolar Covalent Bond • Valence electrons shared between atoms of similar electronegativity constitute a nonpolar covalent bond • For example, bonds between C and H are effectively nonpolar even though C is slightly more electronegative than H • Nonpolar bonds are formed in diatomic molecules

4. Polar Covalent Bond • Valence electrons shared between atoms of unequal electronegativity constitute a polar covalent bond • The atom with the higher electronegativity will develop a partial negative charge • In a single bond, greater differences in electronegativity lead to greater bond dipoles

5. Bond Comparison • All polar bonds have some ionic character • The difference between ionic and covalent bonding is not distinct, but rather a continuum • Generally bonds between a metal and nonmetal are ionic and bonds between nonmetals are covalent • Examination of the properties of a compound is the best was to characterize the type of bonding

6. Metallic Bonding • The valence electrons in a metallic solid are considered to be delocalized • The valence electrons are not associated with any individual atom

7. 2.2 Intramolecular Force and Potential Energy • A graph of potential energy versus the distance between atoms describes the interactions between atoms • The following graph shows the separation between atoms at which the potential energy is lowest (when the bond will form) • It also shows the bond energy (the energy required to separate the atoms

8. Graph of PE vs. Distance

9. Bond Formation Example • When two hydrogen atoms approach each other, two “bad” energy things happen • Electron/electron repulsion and proton/proton repulsion • One “good” energy thing happens • Proton/electron attraction • When the attractive forces offset the repulsive forces, the energy of the two atoms decreases and a bond is formed

10. Before Bond Formation

11. Strength of Covalent Bonds • The bond length is influenced by both the size of the atom's core and the bond order (single, double, triple) • Bonds with a higher order are shorter and have larger bond energies (are stronger)

12. Ionic Interaction Strength • Coulomb’s law can be used to understand the strength of interactions between cations and anions • The interaction strength is proportional to the charge on each ion, larger charges lead to stronger interactions

13. Ionic Interaction Strength • The interaction strength increases as the distance between the centers of the ions decreases • Smaller ions lead to stronger interactions • This is when the ionic size trend becomes helpful in estimating or comparing bond strength

14. Relative Strength of the Interactions

15. 2.3 Structure of Ionic Solids • The cations and anions in an ionic crystal are arranged in a systematic, periodic 3-D array • This crystal lattice maximizes the attractive forces among the cations and anions while minimizing the repulsive forces • The energy of the crystal lattice is proportional to the strength of interactions between cations and anions

16. Typical Ionic Crystal Lattice

17. Comparative Strength

18. Practice • Explain why the lattice energy of LiF is so much stronger than the energy of LiI • Answer: this reasoning is totally due to atomic size. Iodine is a much larger atom than fluorine, so fluorine will be closer to Li in the lattice, making a stronger interaction • This should be answered using the formula for Coulomb’s Law

19. Another Practice • Explain why the lattice energy of NaCl is so much weaker than the energy of MgCl2 • Answer: this reasoning is due to two factors, atomic size and charge size. Mg is a smaller atom than sodium, so this enables atoms to be closer together, strengthening the interaction. Also, Mg has a 2+ charge and requires two Cl- charges to balance it, this gives a stronger interaction. (use the formula for Coulomb’s law)

20. 2.4 Structure of Metals and Alloys • Metallic bonding can be represented as an array of positive metal ions surrounded by delocalized valence electrons • This is referred to as an “electron sea” • This array allows for alloys to be easily formed • Alloy: a metallic mixture made by combining two or more metallic elements, especially to give greater strength or resistance to corrosion.

21. Electron Sea Model

22. Metal Alloys • Alloys typically retain a sea of mobile electrons and so remain conducting • Often the surface of a metal or alloy is changed through a chemical reaction (for example: the formation of a chemically inert oxide layer in stainless steel, through reaction with oxygen in the air • Two main types: substitutional and interstitial

23. Substitutional Alloy • Substitutional alloys form between atoms of comparable radius, where one atom substitutes for the other in the lattice • Example: brass, in which some Cu atoms are substituted usually with Zn • The density typically lies between those of the component metals and the alloy remains malleable and ductile

24. Interstitial Alloy • Interstitial alloys are formed when some of the interstices (holes) in the metal structure are occupied by small atoms • The presence of the interstitial atoms changes the properties of the host metal • Frequently the alloy is harder, stronger, and less ductile

25. Examples of the Two Types of Alloys

26. 2.5 Lewis Structures • The Lewis structure of a molecule show how the valence electrons are arranged among the atoms in the molecule • Named after G. N. Lewis • From experimentation, chemists learned that the most important requirement for the formation of a stable compound is that the atoms achieve noble gas electron configurations

27. Steps to Draw a Lewis Structure • Draw the Lewis structure for each atom • Add the total number of valence electrons • Add the total number of e-’s in lone pairs • Subtract the number of electrons in lone pairs from the total number of valence e-’s • Divide this number by 2, and this is usually the number of bonds you have • Draw sticks for the bonds that represent 2 e-’s, draw in lone pairs, check for octets

28. Helpful Hints • Hydrogen only needs to form a duet; it is stable with 2 valence electrons • Carbon forms 4 bonds (some can be multiple), oxygen forms 2 bonds (can be a double bond), nitrogen forms 3 bonds (can be multiple) • The atom with the lowest electronegativity is usually the central atom • If carbon is present, it is the central atom

29. Practice • Draw Lewis Structures for the following: • H2O • O2 • CH2O • PCl3 • CHI3 • C2H4 • CO2

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33. Ions and Lewis Structures • You must consider the absence of an electron or the extra electrons • Subtract an absent electron from the total and add the extra electrons to the total (it helps to draw them differently) • Do not change the number of electrons on the central atom • Place the entire structure in brackets and put the charge in the upper right corner outside the brackets

34. Practice • Draw Lewis Structures for the following: • CN- • NO+

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36. Exceptions to the Octet Rule • Fewer than 8: hydrogen has at most 2 electrons to form one bond, beryllium will have only 4 valence electrons, boron will have only 6 valence electrons (BF3) • Expanded valence: this can only happen if the central element has d-orbitals, the element is surrounded by more than four valence pairs in certain compounds • Odd-electron compounds: a few stable compounds contain an odd number of valence electrons, for example, NO, NO2, ClO2

37. Comments About the Octet Rule • The 2nd row elements C, N, O, and F obey the octet rule • 2nd row elements B and Be often have fewer than 8 e-’s. They are very reactive • 2nd row elements never exceed the octet rule • 3rd row and heavier elements often satisfy the octet rule, but can exceed the octet rule by using their empty valence d orbitals

38. Octet Rule Violations • Write the Lewis structure for the following: • PCl5 • I3- • ClF3 • RnCl2 • BeCl2 • ICl4-

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41. 2.6 Resonance and Formal Charge • Sometimes more than one valid Lewis structure is possible for a given molecule • Consider NO3- • The structure implies there are two types of bonds, but experiments clearly show that there is only one type of bond

42. Resonance • The structure for NO3- is valid, but does not represent the true bonding, so the model must be modified • There is no reason for choosing a particular O atom to have the double bond

43. Resonance • The nitrate ion doesn’t exist as any of the three structures, but as an average of all three • Resonance: a condition occurring when more than one valid Lewis structure can be written for a molecule

44. Example • Draw the Lewis structure for the following: • NO2- • CO32- • O3

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46. Formal Charge • Formal charge: the difference between the number of valence electrons on the free atoms and the number of valence electrons assigned to the atom in the molecule • Easy way to calculate: • # valence e- − # nonbonding e- − ½ #bonding e- • Formal charge is used to determine the correct Lewis structure when more than one is possible, but experimentation is needed to confirm

47. Example • There are two possible Lewis structures for the sulfate ion • Calculate the formal charge for all atoms on each structure realizing that 0 is ideal

48. Answer to Previous Slide • For the first structure: • S = 6 – 0 – 4 = 2 • For each O = 6 – 6 – 1 = -1 • For the second structure: • S = 6 – 0 – 6 = 0 • For O with double bond: 6 – 4 – 2 = 0 • For O with single bond: 6 – 6 – 1 = -1

49. Justification for Choice • Atoms in molecules try to achieve formal charges as close to zero as possible • We can expect resonance structures for the sulfate ion Usually you will see just these two

50. Examples • Give the possible Lewis structures after considering formal charge. • XeO3 (an explosive compound, strong oxidizer, unstable but slow acting) • SO2 (a reducing agent and is used for bleaching and as a fumigant and food preservative, large quantities of sulfur dioxide are used in the contact process for the manufacture of sulfuric acid)