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Chapter 16

Chapter 16. Covalent Bonding. Section 16.1 The Nature of Covalent Bonding. Objectives ? Use electron dot structures to show the formation of single, double, and triple covalent bonds

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Chapter 16

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  1. Chapter 16 Covalent Bonding

  2. Section 16.1 The Nature of Covalent Bonding • Objectives? • Use electron dot structures to show the formation of single, double, and triple covalent bonds • Describe and give examples of coordinate covalent bonding, resonance structures, and exceptions to the octet rule.

  3. Single Covalent Bonds • A bond in which two atoms share a pair of electrons is called a single covalent bond.

  4. Structural Formulas • When writing a formula for a covalent bond, the pair of electrons is represented as a dash. This is called a structural formula. These dashes are never used to show ionic bonds!

  5. Unshared Pairs • Pairs of valence electrons that are not shared between atoms are called unsharedpairs, also known as lone pairs or nonbonding pairs.

  6. Draw the electron dot (Lewis) structures for each molecule below. • Cl2 • HBr • PCl3

  7. Double and Triple Bonds • Double covalent bonds are bonds that involve two shared pairs of electrons. • Triple covalent bonds are bonds that involve three shared pairs of electrons.

  8. Draw the electron dot (Lewis) structure for the following: • N2 • HCN • BF41-

  9. Coordinate Covalent Bonds • A covalent bond in which one atom contributes both bonding electrons is called a coordinate covalent bond. • In structural formulas, you can show coordinate covalent bonds as arrows that point from the atom donating the pair of electrons to the atom receiving them.

  10. Draw the electron dot (Lewis) structure for the following: • CO • SO2 • NH4+

  11. Bond Dissociation Energies • The total energy required to break the bond between two covalently bonded atoms is known as the bond dissociation energy. H – H + 435 kJ → H. + H.

  12. Resonance • Resonance structures are structures that occur when it is possible to write two or more valid electron dot formulas that have the same number of electron pairs for a molecule or ion.

  13. Exceptions to the Octet Rule • If the central atom is located in row 3 or beyond, it is possible for it to hold more than 8 electrons during bonding. This is due to the fact that the d-sublevel is introduced in the 3rd energy level (row)

  14. Draw the electron dot (Lewis) structure for the following: • SF6 • XeF4

  15. It is possible for an unpaired electron to be present in Lewis structures. This represents instability in a molecule. • The concept of pairing is important in understanding the bonding and properties of molecules.

  16. Diamagnetic Vs. Paramagnetic • One can consider electrons as small, spinning, electric charges. These moving electric charges create magnetic fields, much as the current in an electric motor creates a magnetic field. • Since paired electrons spin in opposite direction, the magnetic effects essentially cancel. • Substances in which all of the electrons are paired are diamagnetic. • In contrast, paramagnetic substances are those that contain one or more unpaired electrons.

  17. Predict whether the following species are diamagnetic or paramagnetic. • BF3 • O2- • NO2 • F2

  18. Section 16.1 The Nature of Covalent Bonding • Did We Meet Our Objectives? • Use electron dot structures to show the formation of single, double, and triple covalent bonds • Describe and give examples of coordinate covalent bonding, resonance structures, and exceptions to the octet rule.

  19. Section 16.2 Bonding Theories • Objectives: • Describe the molecular orbital theory of covalent bonding, including orbital hybridization • Use VSEPR theory to predict the shapes of simple covalent bonded molecules

  20. Molecular Orbitals • There is a quantum mechanical model of bonding that describes the electrons in molecules by means of orbitals that exist only for grouping of atoms. • When two atoms combine, this quantum mechanical model assumes that their atomic orbitals overlap to produce molecular orbitals, or orbitals that apply to the entire molecule.

  21. Molecular orbitals are obtained by combining the atomic orbitals on the atoms in the molecule. Consider H2for example. • One of the molecular orbitals in this molecule is constructed by adding the mathematical functions for the two 1s atomic orbitals that come together to form this molecule. (σ1s) • Another orbital is formed by subtracting one of these functions from the other, as shown in the figure below. (σ*1s) Nucleus

  22. One of these orbitals is called a bonding molecular orbital because electrons in this orbital spend most of their time in the region directly between the two nuclei. It is called a sigma(σ) molecular orbital because it looks like an s orbital when viewed along the H-H bond. • Electrons placed in the other orbital spend most of their time away from the region between the two nuclei. This orbital is therefore an antibonding, or sigma star(σ*), molecular orbital. Nucleus

  23. Each orbital accommodates two electrons, and the two electrons in H-H fills the σ1smolecular orbital (MO). Obviously, as a result of the formation of H2 molecule, the energy of the system is lowered and become more stable. • Bond order will tell whether a single (BO = 1), double (BO = 2), or triple (BO = 3) bond will work.

  24. Example: He2 σ*1s ↑↓ • Bond Order = ½ (σ – σ*) = ½ (2e- – 2e-) = 0 • This is why He2 doesn’t exist! ↑↓ ↑↓ 1s 1s ↑↓ σ1s

  25. σ*2s ↑↓ Example: O22s22p4 ↑↓ ↑↓ 2s 2s • Bond Order ½ (σ– σ*) ½ (8e-– 4e-) 2 • This is why O2 has a double bond! ↑↓ σ2s σ*2p ↑ ↑ ↑↓ ↑ ↑ ↑↓ ↑ ↑ 2p 2p ↑↓ ↑↓ ↑↓ σ2p

  26. σ*2s ↑↓ Example: N22s22p3 ↑↓ ↑↓ 2s 2s • Bond Order ½ (σ– σ*) ½ (8e-– 2e-) 3 • This is why N2 has a triple bond! ↑↓ σ2s σ*2p ↑ ↑ ↑ ↑ ↑ ↑ 2p 2p ↑↓ ↑↓ ↑↓ σ2p

  27. σ*2s ↑↓ Example: F22s22p5 ↑↓ ↑↓ 2s 2s • Bond Order ½ (σ– σ*) ½ (8e-– 6e-) 1 • This is why F2 has a single bond! ↑↓ σ2s σ*2p ↑↓ ↑ ↑ ↑↓ ↑↓ ↑ ↑↓ ↑↓ ↑ 2p 2p ↑↓ ↑↓ ↑↓ σ2p

  28. Molecular orbital theory is a logical extension of the quantum mechanical description of the electron structure of atoms. Covalent bonding is described in terms of sigma and pi bonds. • Single bond– sigma bond • Double bond– sigma, pi bonds • Triple bond– sigma, pi, pi bonds Ethene

  29. The first two electrons being shared form a sigma bond. These orbitals overlap symmetrically. • The next set of electrons being shared form a pi bond. These orbitals overlap side-by-side, which is why they have lobes. Pi bonds are weaker than sigma bonds.

  30. Ex: What is the total number of sigma bonds and pi bonds in each molecule? Draw the three dimensional shape to represent these bonds. • 2 sigma, 2 pi bonds • 6 sigma, 2 pi bonds

  31. VSEPR • Valence Shell Electron Pair Repulsion (VSEPR) theory states that because electron pairs repel, molecular shape adjusts so the valence-electron pairs are as far apart as possible.

  32. Ex: Use VSEPR theory to predict the shapes of the following species. • CO2 • SiCl4

  33. Ex: Use VSEPR theory to predict the shapes of the following species. • XeF4 • IO3-

  34. Hybrid Orbitals • The VSEPR theory works well when accounting for molecular shapes, but does not help much in describing the types of bonds formed. • In hybridization, several atomic orbitals mix to form the same total number of equivalent hybrid orbitals. sp3 hybrid orbitals

  35. Ex: What types of hybrid orbitals are involved in the bonding of carbon atoms in the following molecules? • CH4 sp3 • HC≡CH sp • H2C=CH2 sp2 • N≡C–C≡N sp

  36. Section 16.2 Bonding Theories • Did We Meet Our Objectives? • Describe the molecular orbital theory of covalent bonding, including orbital hybridization • Use VSEPR theory to predict the shapes of simple covalent bonded molecules

  37. Section 16.3 Polar Bonds and Molecules • Objectives: • Use electronegativity values to classify a bond as nonpolar covalent, polar covalent, or ionic • Name and describe the weak attractive forces that hold groups of molecules together.

  38. Ionic Bonds

  39. Covalent Bond(aka NonPolar Covalent)

  40. Polar covalent

  41. Metallic Bond

  42. Bond Polarity • Covalent bonds involve electron sharing between atoms. However, covalent bonds differ in terms of how the bonded atoms share the electrons. • The bonding pairs of electrons in covalent bonds are pulled, as in a tug-of-war. • When the atoms are pulled equally, the bond is a nonpolar covalent bond. • When the pulling is unequal, the bond is a polar covalent bond, or simply polar bond.

  43. Ex: The bonds between the following pairs of elements are covalent. Arrange them according to polarity, naming the most polar bond first. • H―Cl, H―C, H―F, H―O, H―H, S―Cl • H―Cl = |2.1 – 3.0| = 0.9 • H―C = |2.1 – 2.5| = 0.4 • H―F = |2.1 – 4.0| = 1.9 • H―O = |2.1 – 3.5| = 1.4 • H―H = |2.1 – 2.1| = 0.0 • S―Cl = |2.5 – 3.0| = 0.5 • H―F, H―O, H―Cl, S―Cl, H―C, H―H

  44. Polar Molecules • In a polar molecule, one end of the molecule is slightly negative and the other end is slightly positive. • The electron-cloud picture of hydrofluoric acid shows that the fluoride atom attracts the electron cloud more than the hydrogen atom does. This is because fluorine has a higher electronegativity.

  45. A molecule that has two poles, or electrically charged regions, is called a dipolar molecule, or simply dipole. • Hydrochloric acid is a dipole:

  46. The effect of polar bonds on the polarity of an entire molecule depends on the shape of the molecule and the orientation of the polar bonds. • A carbon dioxide molecule has two polar bonds and is linear. Since the carbon and oxygens lie along the same axis, the polarities cancel (opposite directions). This is why CO2 is nonpolar!

  47. The water molecule also has two polar bonds. However, the water molecule is bent rather than linear. Therefore, the bond polarities do not cancel and this is why water is polar!

  48. Ex: Based on the information about molecular shapes in Section 16.2, which of these molecules would you expect to be polar? • SO2 • Polar • H2S • Polar • CO2 • Nonpolar • BF3 • Nonpolar

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