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The Ties That Bind

The Ties That Bind. Chemical Bonding and Interactions. Chemical Bonding and Interactions. Stable Electron Configurations Electron-Dot (Lewis) Structures Drawing, Rules for Drawing The Octet Rule Some Exceptions to the Rule Ionic Bonding Naming ionic compounds Drawing Covalent Bonding

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The Ties That Bind

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  1. The Ties That Bind Chemical Bonding and Interactions

  2. Chemical Bonding and Interactions • Stable Electron Configurations • Electron-Dot (Lewis) Structures • Drawing, Rules for Drawing • The Octet Rule • Some Exceptions to the Rule • Ionic Bonding • Naming ionic compounds • Drawing • Covalent Bonding • Naming covalent compounds • Drawing • Electronegativity and Polar Covalent Compounds • Molecular Shapes and the VSEPR Theory • Intermolecular Forces of Attraction • H-bonds, Dipole-Dipole, Ion-Dipole, London Dispersion Forces

  3. It’s all about stability… • Fact: Noble gases are inert (undergo few, if any, chemical reactions) • Theory: The inertness of these gases is a result of their electronic structure • If elements could alter their electron structures like those of the noble gases, they would be less reactive.

  4. Electron Configurations of Cations and Anions Of Representative Elements Na [Ne]3s1 Na+ [Ne] Atoms lose electrons so that cation has a noble-gas outer electron configuration. Ca [Ar]4s2 Ca2+ [Ar] Al [Ne]3s23p1 Al3+ [Ne] H 1s1 H- 1s2 or [He] Atoms gain electrons so that anion has a noble-gas outer electron configuration. F 1s22s22p5 F- 1s22s22p6 or [Ne] O 1s22s22p4 O2- 1s22s22p6 or [Ne] N 1s22s22p3 N3- 1s22s22p6 or [Ne]

  5. -1 -2 -3 +1 +2 +3 Cations and Anions Of Representative Elements

  6. What neutral atom is isoelectronic with H- ? Na+: [Ne] Al3+: [Ne] F-: 1s22s22p6 or [Ne] O2-: 1s22s22p6 or [Ne] N3-: 1s22s22p6 or [Ne] Na+, Al3+, F-, O2-, and N3- are all isoelectronic with Ne H-: 1s2 same electron configuration as He

  7. Group e- configuration # of valence e- ns1 1 1A 2A ns2 2 3A ns2np1 3 4A ns2np2 4 5A ns2np3 5 6A ns2np4 6 7A ns2np5 7 Valence electrons are the outer shell electrons of an atom. The valence electrons are the electrons that particpate in chemical bonding. 9.1

  8. (Electron Dot Structures) *We generally place the electrons one four sides of a square around the element symbol. *Octet rule: we know that s2p6 is a noble gas configuration. We assume that an atom is stable when surrounded by 8 electrons (4 electron pairs).

  9. A violent bond: Sodium and Chlorine • Sodium is a soft, reactive metal. • Chlorine is a greenish-yellow gas which is toxic (WWI poison) • Sodium chloride is used for our food.

  10. Ionic Bonding Consider the reaction between sodium and chlorine: Na(s) + ½Cl2(g)  NaCl(s)

  11. Chemistry In Action: Sodium Chloride Mining Salt Solar Evaporation for Salt

  12. - - - - + Li+ Li Li Li+ + e- e- + Li+ Li+ + F F F F F F The Ionic Bond [He] [Ne] 1s22s1 1s22s22p5 1s2 1s22s22p6 IONIC BONDING -bonding due to electrostatic attraction arising from an exchange of electrons. 9.2

  13. ionic compounds consist of a combination of cations and an anions • the formula is always the same as the empirical formula • the sum of the charges on the cation(s) and anion(s) in each formula unit must equal zero The ionic compound NaCl 2.6

  14. Ionic Bonding *NaCl forms a very regular structure in which each Na+ ion is surrounded by 6 Cl- ions. Similarly, each Cl- ion is surrounded by six Na+ ions. There is a regular arrangement of Na+ and Cl- in 3D. Note that the ions are packed as closely as possible. Note that it is not easy to find a molecular formula to describe the ionic lattice.

  15. E = k Q+Q- r lattice energy cmpd MgF2 2957 Q= +2,-1 MgO 3938 Q= +2,-2 LiF 1036 LiCl 853 SIDEBAR: Electrostatic (Lattice) Energy Lattice energy (E) is the energy required to completely separate one mole of a solid ionic compound into gaseous ions. Q+ is the charge on the cation Q- is the charge on the anion r is the distance between the ions Lattice energy (E) increases as Q increases and/or as r decreases. r F < r Cl 9.3

  16. 9.3

  17. 11 protons 11 electrons 11 protons 10 electrons Na+ Na 17 protons 18 electrons 17 protons 17 electrons Cl- Cl An ion is an atom, or group of atoms, that has a net positive or negative charge. cation – ion with a positive charge If a neutral atom loses one or more electrons it becomes a cation. anion – ion with a negative charge If a neutral atom gains one or more electrons it becomes an anion. 2.5

  18. A monatomic ion contains only one atom Na+, Cl-, Ca2+, O2-, Al3+, N3- A polyatomic ion contains more than one atom OH-, CN-, NH4+, NO3- 2.5

  19. Test Yourself: • MgO • MgBr2 • The combination of the ion of magnesium and the ion of nitrogen

  20. 1 x +2 = +2 1 x +2 = +2 2 x +3 = +6 1 x -2 = -2 3 x -2 = -6 2 x -1 = -2 Formula of Ionic Compounds Al2O3 Al3+ O2- CaBr2 Ca2+ Br- Na2CO3 Na+ CO32- 2.6

  21. 2.6

  22. 2.7

  23. Chemical Nomenclature • Ionic Compounds • often a metal + nonmetal • anion (nonmetal), add “ide” to element name barium chloride BaCl2 potassium oxide K2O magnesium hydroxide Mg(OH)2 potassium nitrate KNO3 2.7

  24. Transition metal ionic compounds • indicate charge on metal with Roman numerals iron(II) chloride FeCl2 2 Cl- -2 so Fe is +2 FeCl3 3 Cl- -3 so Fe is +3 iron(III) chloride Cr2S3 3 S-2 -6 so Cr is +3 (6/2) chromium(III) sulfide 2.7

  25. Test Yourself • Write the formula and draw the ionic structure • Magnesium (II) bromide • Iron (III) Oxide • Sodium Azide • Ammonium Chloride • Magnesium Nitrate • Ammonium Phosphate • Write the name • FeO • Fe2O3 • LiBr • KCrO4 • K2Cr2O7

  26. 2.7

  27. Covalent Bonding

  28. Why should two atoms share electrons? + 8e- 7e- 7e- 8e- F F F F F F F F lonepairs lonepairs single covalent bond single covalent bond lonepairs lonepairs A covalent bond is a chemical bond in which two or more electrons are shared by two atoms. Lewis structure of F2 9.4

  29. single covalent bonds H H H H or H H O 2e- 2e- O 8e- O C O C O O double bonds 8e- 8e- 8e- double bonds O N N triple bond N N triple bond 8e- 8e- Lewis structure of water + + Double bond – two atoms share two pairs of electrons or Triple bond – two atoms share three pairs of electrons or 9.4

  30. Lengths of Covalent Bonds Bond Lengths Triple bond < Double Bond < Single Bond 9.4

  31. 9.4

  32. F H F H Polar covalent bond or polar bond is a covalent bond with greater electron density around one of the two atoms electron rich region electron poor region e- poor e- rich d+ d- 9.5

  33. X (g) + e- X-(g) Electronegativity is the ability of an atom to attract toward itself the electrons in a chemical bond. Electron Affinity - measurable, Cl is highest Electronegativity - relative, F is highest 9.5

  34. 9.5

  35. Bond Polarity and Electronegativity Electronegativity

  36. Drawing Lewis Structures 1. Add the valence electrons. 2. Identify the central atom (usually the one with the highest molecular mass, lowest electronegativity, or closest to the center of the periodic table). 3. Place the central atom in the center of the molecule and add all other atoms around it. 4.Place one bond (two electrons) between each pair of atoms. 5.Complete the octet for the central atom. 6.Complete the octets for all other atoms. Use double bonds if necessary. 7. Place remaining electrons on the central atom.

  37. Write the Lewis structure of nitrogen trifluoride (NF3). F N F F Step 1 – N is less electronegative than F, put N in center Step 2 – Count valence electrons N - 5 (2s22p3) and F - 7 (2s22p5) 5 + (3 x 7) = 26 valence electrons Step 3 – Draw single bonds between N and F atoms and complete octets on N and F atoms. Step 4 - Check, are # of e- in structure equal to number of valence e- ? 3 single bonds (3x2) + 10 lone pairs (10x2) = 26 valence electrons 9.6

  38. Write the Lewis structure of the carbonate ion (CO32-). O C O 2 single bonds (2x2) = 4 1 double bond = 4 8 lone pairs (8x2) = 16 O Total = 24 Step 1 – C is less electronegative than O, put C in center • Step 2 – Count valence electrons C - 4 (2s22p2) and O - 6 (2s22p4) • -2 charge – 2e- 4 + (3 x 6) + 2 = 24 valence electrons Step 3 – Draw single bonds between C and O atoms and complete octet on C and O atoms. Step 4 - Check, are # of e- in structure equal to number of valence e- ? 3 single bonds (3x2) + 10 lone pairs (10x2) = 26 valence electrons Step 5 - Too many electrons, form double bond and re-check # of e- 9.6

  39. Trial • Write the Lewis stucture for • H2O • CO2 • NH3 • CH4 • N2 • BH3

  40. Be – 2e- 2H – 2x1e- H Be H 4e- B – 3e- 3 single bonds (3x2) = 6 3F – 3x7e- F B F 24e- Total = 24 9 lone pairs (9x2) = 18 F Exceptions to the Octet Rule The Incomplete Octet BeH2 BF3 9.9

  41. N – 5e- S – 6e- N O 6F – 42e- O – 6e- 48e- 11e- F 6 single bonds (6x2) = 12 F F Total = 48 S 18 lone pairs (18x2) = 36 F F F Exceptions to the Octet Rule Odd-Electron Molecules NO The Expanded Octet (central atom with principal quantum number n > 2) SF6 9.9

  42. Strengths of Covalent Bonds Bond Enthalpies and Bond Length

  43. Increasing difference in electronegativity Covalent Polar Covalent Ionic partial transfer of e- share e- transfer e- Classification of bonds by difference in electronegativity Difference Bond Type 0 Covalent  2 Ionic 0 < and <2 Polar Covalent 9.5

  44. Classify the following bonds as ionic, polar covalent, or covalent: The bond in CsCl; the bond in H2S; and the NN bond in H2NNH2. Cs – 0.7 Cl – 3.0 3.0 – 0.7 = 2.3 Ionic H – 2.1 S – 2.5 2.5 – 2.1 = 0.4 Polar Covalent N – 3.0 N – 3.0 3.0 – 3.0 = 0 Covalent 9.5

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