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Bonding: General Concepts

Bonding: General Concepts. Chapter 8. Overview. Types of chemical bonds, Electronegativity, Bond polarity and Dipole Moments. The Ions: electron configurations, size, formula, lattice energy calculations. Covalent bonds: model, bond energies, chemical reactions.

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Bonding: General Concepts

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  1. Bonding:General Concepts Chapter 8

  2. Overview • Types of chemical bonds, Electronegativity, Bond polarity and Dipole Moments. • The Ions: electron configurations, size, formula, lattice energy calculations. • Covalent bonds: model, bond energies, chemical reactions. • Lewis structure, exceptions to the octet rule, resonance. • Molecular structure models from Valence Shell Electron Pair Model “VSEPR” for single and multiple bonds.

  3. Bonds • Forces that hold group of atoms together and make them function as a unit. NaCl – attraction is electrostatic since Na+ and Cl- are the stable forms of these elements. This is an example of “Ionic Bonding”

  4. Bond Energy • It is the energy required to break a bond. • It gives us information about the strengthof a bonding interaction, as well as radius.

  5. Bond Length • The distance where the system energy is a minimum.

  6. Figure 8.1: (a) The interaction of two hydrogen atoms.(b) Energy profile as a function of the distance between the nuclei of the hydrogen atoms.

  7. Change in electron density as two hydrogen atoms approach each other.

  8. Covalent Bond • No electron transfer • Electrons are shared between two atoms, positioned between the two nuclei • Example: H2, O2, H2O, CO2 etc. Among non-metals only

  9. Ionic Bonds • Formed from electrostatic attractions of closely packed, oppositely charged ions. • Formed when an atom that easily loses electronsreacts with one that has a high electron affinity.

  10. - - - - + Li+ Li Li Li+ + e- e- + Li+ Li+ + F F F F F F The Ionic Bond [He] [Ne] 1s22s1 1s22s22p5 1s2 1s22s22p6

  11. Cation is always smaller than atom from which it is formed. Anion is always larger than atom from which it is formed. Ionic Compound

  12. Ionic BondsandCoulomb’s Law • Q1 and Q2 = numerical ion charges • r = distance between ion centers (in nm) • Attractive forces are (-), repulsive are (+).

  13. Covalent or Ionic? • Covalent and ionic are simply extreme cases. • Most molecules share electrons but “Unequally” due to the difference in electronegativity and electron affinity. • This will give rise to a dipole moment and the molecule becomes polar.

  14. (H-H) + (X-X) (H-X)experimental# (H-X)expected = 2 Electronegativity • The ability of an atom in a molecule to attract shared electrons towards itself. • Linus Pauling simple model Δ= (H  X)actual (H  X)expected If  = 0 => No Polarity

  15. 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 to 0.1 Covalent  2 Ionic 0 < and <2 Polar Covalent

  16. 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

  17. F H F H Polar covalent bondorpolar bond is a covalent bond with greater electron density around one of the two atoms. The molecule is called “Dipolar”. electron rich region electron poor region e- poor e- rich d+ d- Dipole Moment

  18. The effect of an electric field on hydrogen fluoride molecules. • When no electric field is present, the molecules are randomly oriented. (b) When the field is turned on, the molecules tend to line up with their negative ends toward the positive pole and their positive ends toward the negative pole.

  19. Polyatomic Molecules • May exhibit dipole moment depending on their structure i.e. arrangement in space

  20. Comparison of Ionic and Covalent Compounds

  21. Figure 8.4: (a) The charge distribution in the water molecule.(b) The water molecule in an electric field. V-Shape

  22. Figure 8.5: (a) The structure and charge distribution of the ammonia molecule. The polarity of the N—H bonds occurs because nitrogen has a greater electronegativity than hydrogen. (b) The dipole moment of the ammonia molecule oriented in an electric field. Look for a “NET DIPOLE”

  23. Look for a “NET DIPOLE” N H H H Trigonal Pyramidal Structure

  24. The carbon dioxide molecule CO2: The opposed bond polarities cancel out, and the carbon dioxide has no dipole moment: Non-polar molecule Note:The C-O in CO2bond is polar, but the net dipole is Zero. Other examples:SO3, CCl4, etc.

  25. O O S H H H O O O C H H C H Which of the following molecules have a dipole moment? H2O, CO2, SO2, and CH4 dipole moment polar molecule dipole moment polar molecule no dipole moment nonpolar molecule no dipole moment nonpolar molecule

  26. Dipoles (polar molecules) and Microwaves

  27. Predicting Formulas of Ionic Compounds • Two nonmetalsreact: They share electrons to achieve NGEC(Noble Gas Electron Configurations) and form covalent bonds. • A nonmetal and a representative group metalreact (ionic compound): The valence orbitals of the metal are emptied (cation) to achieve NGEC. The valence electron configuration of the nonmetal (anion) achieves NGEC.

  28. Ions • Ionic compounds are always electrically neutral e.g. they have the same amount of +ve and –ve charges. • Common ions have noble gas configurations.

  29. 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]

  30. ns2np6 ns1 ns2np1 ns2np2 ns2np3 ns2np4 ns2np5 ns2 d10 d1 d5 4f 5f Ground State Electron Configurations of the Elements

  31. -1 -2 -3 +1 +2 +3 Cations and Anions Of Representative Elements

  32. Notes on Ions • Hydrogen may form H+ or H- • Tin forms Sn2+ and Sn4+. • Transition metals exhibit a more complicated behavior.

  33. Example of Ionic Compounds • MgO magnesium oxide is formed of Mg2+ and O2-. • CaO formed from Ca2+ and O2-. • Al2O3 is formed of 2Al3+ and 3O2-.

  34. Figure 8.7: Sizes of ions related to positions of the elements on the periodic table.

  35. PERIODIC TRENDS OF ELEMENTS

  36. Isoelectronic IonsContains the same number of electrons • 8O 1s22s22p4 O2-1s22s22p6 • 9F 1s22s22p5 F- 1s22s22p6 • 11Na 1s22s22p63s1 Na+ 1s22s22p6 • 12Mg 1s22s22p63s2 Mg2+1s22s22p6 • 13Al 1s22s22p63s23p1 Al3+1s22s22p6 Which one you expect to have the smallest radius? And why?

  37. Radii of Isoelectronic Ions • O2> F > Na+ > Mg2+ > Al3+ • largest smallest • 13 protons Vs 10 electrons

  38. Example • Choose the largest ion in each of the following groups: • Li+, Na+, K+, Rb+, Cs+ • Ba2+ , Cs+ , I- , Te2-

  39. To Reiterate • Cation is smaller than parent molecule. • Anion is larger than parent molecule. • Size increase down in a group (+ve or –ve). • The larger the mass number the smaller the size for isoelectronic cations and anions.

  40. Electron Configurations of Cations of Transition Metals When a cation is formed from an atom of a transition metal, electrons are always removed first from the ns orbital and then from the (n – 1) d orbitals. Fe: [Ar]4s23d6 Mn: [Ar]4s23d5 Fe2+: [Ar]4s03d6or [Ar]3d6 Mn2+: [Ar]4s03d5or [Ar]3d5 Fe3+: [Ar]4s03d5or [Ar]3d5

  41. Why Compounds Exist? The driving force behind the naturally occurring compounds (such as NaCl, H2O, etc.) is to yield a stable lower energy form. A stable form is a an arrangement of atoms, held together by bonding that prevent decomposition. This bonding energy when ions condense from gas phase into ionic solid is called Lattice Energy.

  42. Lattice Energy • The change in energy when separated gaseous ionsare packed together to form an ionic solid. • M+(g) + X(g)  MX(s) • Lattice energy is negative (exothermic) from the system point of view.

  43. Formation of an Ionic Solid • 1. Sublimation of the solid metal • M(s)  M(g) [endothermic] • 2. Ionization of the metal atoms • M(g)  M+(g) + e[endothermic] • 3. Dissociation of the nonmetal • ½ X2(g)  X(g) [endothermic] • 4. Formation (Electron Affinity) of X ions in the gas phase: • X(g) + e X(g) [exothermic] • 5. Formation of the solid (LATTICE) MX • M+(g) + X(g)  MX(s)[quite exothermic] • Lattice Energy

  44. The energy changes involved in the formation of solid lithium fluoride from its elements. From Gas to Solid Lattice Energy

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