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Understanding Chemical Bonds in Ionic and Molecular Compounds

Explore the formation of salt and water, key compounds in chemistry. Learn about ionic versus molecular compounds, Lewis dot structures, octet rule, and bond formation. Discover how atoms achieve stability and predict charges and formulas using the octet rule.

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Understanding Chemical Bonds in Ionic and Molecular Compounds

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  1. Chapter 9 The Chemical Bond

  2. Setting the Stage – Salt and Water • Earth is a complex world of chemicals • We will focus on two basic types of compounds: • Salt (sodium chloride) is a typical ionic compound • Water is a typical molecular compound • These substances are formed from other substances with very different properties. • We will explore these differences and reactions. Malone and Dolter - Basic Concepts of Chemistry 9e

  3. Setting a Goal – Part AChemical Bonds and the Nature of Ionic Compounds • You will learn how the octet rule is used to determine the charge on ions and the formulas of ionic compounds. Malone and Dolter - Basic Concepts of Chemistry 9e

  4. Objectives for Section 9-1 • Explain the significance of the octet rule. • Write the Lewis dot structure of the atoms of any representative element. Malone and Dolter - Basic Concepts of Chemistry 9e

  5. Professor’s Little Jokes • What do two dipoles say to each other? “Have you got a moment?” • Why did the white bear not get along with the other bears? He was polar. • My name is Bond – Ionic Bond Malone and Dolter - Basic Concepts of Chemistry 9e

  6. 9-1 Bond Formation and the Representative Elements • The s block and p block elements (often termed the representative elements) will often form bonds such that there are eight outer electrons surrounding each atom (the octet rule). Malone and Dolter - Basic Concepts of Chemistry 9e

  7. Bond Formation and the Representative Elements….Cont’d • Obtaining this configuration is the driving force for bond formation for many compounds formed by the representative elements. • The exceptions are: H and Li (which tend to follow a “duet” rule - filling the ns subshell); some compounds of group IIA and IIIA elements (e.g. BF3); odd-electron molecules (‘radicals’, e.g. NO); and hypervalent compounds of period 3+ elements (e.g. SF6). Malone and Dolter - Basic Concepts of Chemistry 9e

  8. Bond Formation and Noble Gases • Noble gases rarely form compounds. • They have filled s and p outer subshells. • This is a total of eight electrons, referred to as an octet. • Eight electrons in the outer s and p orbitals is a particularly stable configuration. • The energy required to remove an electron from these full subshells is particularly high. Malone and Dolter - Basic Concepts of Chemistry 9e

  9. Chemical Bonds • A chemical bond is the force that holds two or more atoms together. • Chemical bonds involve the valence electrons. • The valence electrons are the electrons in the outer s and p subshells. • A bond results if a more stable electron configuration results. Malone and Dolter - Basic Concepts of Chemistry 9e

  10. How Atoms Achieve an Octet • Metals can lose one to three electrons to form a cation with the electron configuration of the previous noble gas. • Nonmetals can gain one to three electrons to form an anion with the electron configuration of the next noble gas. • Atoms can share electrons. Malone and Dolter - Basic Concepts of Chemistry 9e

  11. Lewis Dot Symbols for Elements • G.N. Lewis – originator of dot structures and the octet rule Malone and Dolter - Basic Concepts of Chemistry 9e

  12. Lewis Dot Symbols for Elements • Since only the valence electrons are involved in bonding we can concentrate on those. • Lewis dot symbols are used to represent the valence electrons of an atom. • See Table 9-1. Malone and Dolter - Basic Concepts of Chemistry 9e

  13. Objective for Section 9-2 • Using the Lewis dot structure and the octet rule, predict the charges on the ions of representative elements and the formulas of binary ionic compounds. Malone and Dolter - Basic Concepts of Chemistry 9e

  14. 9-2 Formation of Ions and Ionic Compounds • Metals can lose electrons to form ions • Na ([Ne]3s1)  Na+ ([Ne]) + e- • If a metal loses all of its outer electrons, it acquires the octet of the previous noble gas • Nonmetals can gain electrons to form ions • Cl ([Ne]3s23p5+ e- Cl- ([Ne]3s23p6) • Lewis dot structures of the atoms can be very helpful here. Malone and Dolter - Basic Concepts of Chemistry 9e

  15. Forming Ionic Compounds • Reaction of Na with Cl • Na donates an electron to Cl • Na+ has the previous noble gas structure (Ne) • Cl- has the next noble gas structure (Ar) Malone and Dolter - Basic Concepts of Chemistry 9e

  16. Binary Ionic Compounds • In NaCl, each Na+ is surrounded by six Cl-, and each Cl- is surrounded by six Na+. • Ionic lattice is a three-dimensional array of ions, held together by electrostatic attractions. • These electrostatic attractions are called ionic bonds. Malone and Dolter - Basic Concepts of Chemistry 9e

  17. Setting a Goal – Part BChemical Bonds and the Nature of Molecular Compounds • You will learn how to apply the octet rule to draw structures that form the basis of our understanding of bonding in most molecular compounds. Malone and Dolter - Basic Concepts of Chemistry 9e

  18. Objective for Section 9-3 • Describe the covalent bond and how the octet rule determines the number of such bonds for simple compounds. Malone and Dolter - Basic Concepts of Chemistry 9e

  19. 9-3 The Covalent Bond • Covalent bonds result from electron sharing between two atoms. • We use Lewis dot structures to show the order and arrangement of the atoms in a molecule and all of the valence electrons. Malone and Dolter - Basic Concepts of Chemistry 9e

  20. Lewis Structure of F2 Malone and Dolter - Basic Concepts of Chemistry 9e

  21. Types of Covalent Bonds • Two nonmetals can share one, two or three electron pairs. • The bonds resulting from this sharing are referred to as single, double, or triple bonds respectively. • Multiple bonds are frequently observed in compounds of 2nd period elements. Malone and Dolter - Basic Concepts of Chemistry 9e

  22. Types of Covalent Bonds: Examples Malone and Dolter - Basic Concepts of Chemistry 9e

  23. Objective for Section 9-4 • Draw Lewis structures of a number of molecular compounds and polyatomic ions. Malone and Dolter - Basic Concepts of Chemistry 9e

  24. 9-4 Writing Lewis Structures • The octet rule and Lewis dot structures allow us to justify the formulas that we know. • We can also predict the formulas of new compounds. Malone and Dolter - Basic Concepts of Chemistry 9e

  25. Lewis Dot Structures • Count the number of valence electrons for the atoms in the molecule. • For ions, add one electron for each negative charge or subtract one electron for each positive charge. • Place the most electropositive atom in the center (the inner atom). Array the more electronegative atoms around this atom as outer atoms. Malone and Dolter - Basic Concepts of Chemistry 9e

  26. Lewis Dot Structures…Cont’d • Connect the outer atoms to the inner atoms with single bonds. • Subtract two electrons from the total number of valence electrons for each bond. • Array the remaining electrons around the outer atoms in pairs to complete their octet. Malone and Dolter - Basic Concepts of Chemistry 9e

  27. Lewis Dot Structures…Cont’d 7.Check the octets of all atoms. 8.Use lone pairs on the outer atoms to form multiple bonds to the inner atom if needed to complete the inner atom octet. Malone and Dolter - Basic Concepts of Chemistry 9e

  28. Exceptions to the Octet Rule • Some molecules do not follow the octet rule: one type has an odd number of electrons, such as NO and NO2. • These two molecules have an odd number of electrons (11 for NO; 17 for NO2). They are known as (free) radicals • Another exception involves central atom (period 3 and beyond) valence shell extension in high-valence compounds, like SF6 Malone and Dolter - Basic Concepts of Chemistry 9e

  29. Exceptions to the Octet Rule • A second type has a correct Lewis structure, but other evidence is inconsistent with the Lewis structure, such as O2. • Experiments show that O2 is also a free radical with two unpaired electrons. • More complex theories are needed to explain such substances. Malone and Dolter - Basic Concepts of Chemistry 9e

  30. Example: Write a Lewis structure for each of sulfur tetrafuoride and the triiodide ion These are both examples where the central atom disobeys the octet rule (expanded valence shell) Malone and Dolter - Basic Concepts of Chemistry 9e

  31. Objectives for Section 9-5 • Discuss the significance of resonance. • Write resonance structures of appropriate molecules or ions. Malone and Dolter - Basic Concepts of Chemistry 9e

  32. 9-5 Resonance Structures • In compounds with multiple bonds, sometimes you can draw structures which vary only by placement of the double bonds. • The structures are called resonance structures, and are an approximation of the true structure of the molecule. • Actually, the molecule is a superposition of all of the resonance structures. Malone and Dolter - Basic Concepts of Chemistry 9e

  33. Example of Resonance: Nitrate Ion • Nitrate has three resonance structures. • Each is identical except for the placement of the double bond and associated lone pairs. • Experimentally, all N-O bonds are identical. Malone and Dolter - Basic Concepts of Chemistry 9e

  34. Objective for Section 9-6 • Determine the validity of a Lewis structure based on formal charge considerations. Malone and Dolter - Basic Concepts of Chemistry 9e

  35. 9-6 Formal Charge • Formal charge is the charge that each atom in a molecule would have if the electrons in the bonds were divided equally between the two atoms. • This approach essentially treats all bonds as nonpolar (complete electron sharing). Malone and Dolter - Basic Concepts of Chemistry 9e

  36. Calculation of Formal Charge • Formal charge is calculated by subtracting the number of non-bonded electrons on the atom in question and half the shared electrons from the group number. Malone and Dolter - Basic Concepts of Chemistry 9e

  37. Other Applications of Formal Charge • Formal charge can help decide which is the best ordering of bonds. • Formal charge can also help decide whether a specific Lewis structure is legitimate (i.e. just because we can write a Lewis structure that follows the octet rule does not necessarily mean that the structure actually represents the bonding in the molecule). • The best Lewis structure is generally the one with all zero formal charges (where this is possible). Malone and Dolter - Basic Concepts of Chemistry 9e

  38. Choosing Best Lewis Structure -1 Nitrosyl chloride Malone and Dolter - Basic Concepts of Chemistry 9e

  39. Choosing Best Lewis Structure -2 Nitric oxide Malone and Dolter - Basic Concepts of Chemistry 9e

  40. Choosing Best Lewis Structure -3 • NOTE: some molecules can only be represented by Lewis structures with non-zero formal charges. A classic example is ozone O3. Malone and Dolter - Basic Concepts of Chemistry 9e

  41. Flowchart for Treating Molecules Malone and Dolter - Basic Concepts of Chemistry 9e

  42. Setting a Goal – Part CThe Distribution of Charge in Chemical Bonds • You will learn how the correct Lewis structure of a compound allows chemists a thorough understanding of its properties. Malone and Dolter - Basic Concepts of Chemistry 9e

  43. Objective for Section 9-7 • Classify a bond as being polar, nonpolar, or ionic. Malone and Dolter - Basic Concepts of Chemistry 9e

  44. 9-7 Electronegativity and Polarity of Bonds • Electronegativity - the ability of an atom to attract electrons in a bond to itself. • Difference in electronegativities of atoms that are bonded together results in a partial transfer of electron charge to the more electronegative atom. Malone and Dolter - Basic Concepts of Chemistry 9e

  45. Electronegativity and Polarity of Bonds • The bond is therefore a polar covalent bond. • The polar bond has a negative end and a positive end (a so-called dipole; which we indicate with a  with the appropriate sign added). Malone and Dolter - Basic Concepts of Chemistry 9e

  46. Cs 0.79 Na 0.93 H 2.20 C 2.55 N 3.04 O 3.44 Cl 3.16 F 3.99 Electronegativity • L. Pauling – joint originator (with R. Mulliken) of electronegativity scale Malone and Dolter - Basic Concepts of Chemistry 9e

  47. Polarity of Bonds • Bonds that involve atoms of differing electronegativities have a concentration of negative charge at the more electronegative atom, and a deficiency of charge at the less electronegative atom. • This unequal distribution of negative charge creates a dipole, where one end of the bond is slightly negative and the other is slightly positive. Malone and Dolter - Basic Concepts of Chemistry 9e

  48. Bond and Molecular Dipoles- see later for more details Malone and Dolter - Basic Concepts of Chemistry 9e

  49. Predicting the Polarity of Bonds • When two atoms compete for a pair of electrons in a bond, three possibilities exist: • 1. Both atoms share the electrons equally, forming a nonpolar bond (the electronegativity difference between the the atoms is less than 0.4). • Examples: X-X (where X is same element: C-C, H-H etc), C-H, Si-P, P-S, S-Se. Malone and Dolter - Basic Concepts of Chemistry 9e

  50. Bond Polarity….Contd. • 2. The two atoms share electrons unequally, forming a polar bond (the electronegativity difference between the two atoms is less than 1.8, but greater than 0.4). • Examples: Hd+-Hald-, Cld+-Fd-, >Cd+=Od-, Cd--Sid+, Od--Hd+ • 3. The electron pair is not shared, since one atom acquires the electrons (the electronegativity difference between the two atoms is greater than 1.8). • Examples: Cs+Cl-, Na+F-, Al3+(F-)3 Malone and Dolter - Basic Concepts of Chemistry 9e

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