1 / 47

Covalent Bonds and Molecular Structure: The Basics

Learn about covalent bonds, bond strength, electron-dot structures, resonance, molecular shapes, and more in this comprehensive guide.

rsherrill
Download Presentation

Covalent Bonds and Molecular Structure: The Basics

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  1. Chapter 7 Covalent Bonds and Molecular Structure

  2. The Covalent Bond • Covalent bond – formed by the sharing of electrons between two nonmetal atoms • Forces involved in the bond • Electrostatic attraction between proton and electron • Electrostatic repulsion between electron and electron • When will a bond form?

  3. Strengths of Bonds • Bond dissociation energy – the amount of energy required to break a bond • Energy increases as the length of the bond gets shorter

  4. Internuclear distance (bond length) Covalent radius Internuclear distance (bond length) Covalent radius Internuclear distance (bond length) Covalent radius Internuclear distance (bond length) Covalent radius Bond length and covalent radius.

  5. Strengths of Bonds • Single bonds > double bonds > triple bonds

  6. Problem • Arrange the following bonds in order of increasing bond strength.  • A.  C-I < C-Br < C-Cl < C-F • B.  C-F < C-Cl < C-Br < C-I • C.  C-Br < C-I < C-Cl < C-F • D.  C-I < C-Br < C-F< C-Cl • E.  none of these orders is correct

  7. Problem • Select the strongest bond in the following group.  • A.  C-S • B.  C-O • C.  C=C • D.  C≡N • E.  C-F

  8. Electron Dot Structures • Electron dot symbols • Aid in understanding the formation of bonds between atomic nuclei • Elemental symbol represents the type of element and all core electrons; the valence electrons are represented by dots around the symbol

  9. Electron Dot Structures • A metal in an ionic loses its electrons to achieve an octet or pseudo-octet (transition elements) in its outermost shell • A nonmetal in an ionic compound gains electrons to achieve an octet in its outermost shell • Period 1 and 2 elements of a covalent compound share enough electrons to achieve an octet

  10. Electron-Dot Structures • Ionic: • Covalent:

  11. Electron Dot Structures • Using electron dot symbols build • H2, H2O, CH4, O2, N2, HCN, CO2 • Least electronegative atom is often central (except H)

  12. Naming Binary Molecular Compounds • Electronegativity – indicates how well an elements nuclei attract the electrons in a covalent bond

  13. The Periodic Table and Electronegativity

  14. Electron Dot Structures • Using electron dot symbols build • H2, H2O, CH4, O2, N2, HCN, CO2 • Least electronegative atom is often central (except H)

  15. Electron Dot Structures • Single bond: A covalent bond formed by sharing one electron pair. • Double bond: A covalent bond formed by sharing two electron pairs. • Triple bond: A covalent bond formed by sharing three electron pairs. • Single bonds are longer (weaker) than double bonds • Double bonds are longer (weaker) than triple bonds

  16. Electron-dot Structures • Step 1: Count the total valence electrons. • Step 2: Identify the central atom - Often least electronegative • Step 3: Place all other atoms around the central atom • Step 4: Draw a single bond between each external atom and the central atom subtracting 2 electrons for each bond drawn from the total valence electrons.

  17. Electron-dot Structures • Step 5: Distribute remaining valence electrons around the external atoms giving the external atoms an octet • Step 6: If valence electrons still remain, place them on the central atom in pairs • Step 7: Verify that each atom has an octet • Hydrogen needs only 2 electrons • Boron needs only 6 electrons

  18. Electron-Dot Structures • Step 8: If the central atom does not have an octet, form a multiple bond by bringing a pair of electrons in from the external atom • Step 9: Calculate formal charge and minimize the formal charge if acceptable • Period 3 elements and greater can have expanded octets if one is necessary to minimize formal charge • Formal charge = # valence electrons for the atom – 1 for every dot on the atom – 1 for every line around the atom

  19. Problems • BF3 • PF3 • C2H6 • I3+ • NH4+ • SO42- • KClO3

  20. Electron-dot Structures and Resonance • How is the double bond formed in O3? • The correct answer is that both are correct,but neither is correct by itself.

  21. Electron-Dot Structures and Resonance • When multiple structures can be drawn, the actual structure is an average of all possibilities. • The average is called a resonance hybrid. A straight double-headed arrow indicates resonance.

  22. Problem • S3 • PO43- • CO32- • NO2

  23. Molecular Shapes: The VSEPR Theory • The approximate shape of molecules is given by Valence-Shell Electron-Pair Repulsion (VSEPR).

  24. Molecular Shapes: The VSEPR Theory Molecular formula Step 1 Count all e- groups around central atom (A) – Single, Double and Triple bonds are all counted as 1 e- group Lewis structure Step 2 Electron-group arrangement Note lone pairs and double bonds Step 3 Count bonding and nonbonding e- groups separately. Bond angles Step 4 Molecular shape (AXmEn)

  25. Problem • Determine the shape of the molecules for which Lewis Structures have been developed.

  26. Valence Bond Theory • If, in order for a bond to form, a pair of electrons must be shared, then how does C form molecules with 4 bonds? • Valence Bond Theory – hybrid orbitals

  27. Valence Bond Theory Basic Principle A covalent bond forms when the orbtials of two atoms overlap and are occupied by a pair of electrons that have the highest probability of being located between the nuclei. Themes A set of overlapping orbitals has a maximum of two electrons that must have opposite spins. The greater the orbital overlap, the stronger (more stable) the bond. The valence atomic orbitals in a molecule are different from those in isolated atoms.

  28. Key Points Types of Hybrid Orbitals sp sp2 sp3 sp3d sp3d2 Valence Bond Theory The number of hybrid orbitals obtained equals the number of atomic orbitals mixed. The type of hybrid orbitals obtained varies with the types of atomic orbitals mixed.

  29. Step 1 Step 2 Step 3 Valence Bond Theory The conceptual steps from molecular formula to the hybrid orbitals used in bonding. Molecular shape and e- group arrangement Molecular formula Lewis structure Hybrid orbitals

  30. Problems • Carbon uses ______ hybrid orbitals in ClCN.  • A.  sp • B.  sp2 • C.  sp3 • D.  sp3d • E.  sp3d2

  31. both C are sp3 hybridized s-sp3 overlaps to s bonds sp3-sp3 overlap to form a s bond relatively even distribution of electron density over all s bonds The s bonds in ethane.

  32. overlap in one position - s p overlap -  electron density The s and p bonds in ethylene (C2H4)

  33. overlap in one position - s p overlap -  The s and p bonds in acetylene (C2H2)

  34. Polar Covalent Bonds: Electronegativity • Electronegativity – represents the ability of an atom to attract a shared pair of electrons • Higher the EN – the more the electrons in a bond will be pulled toward the atom • Most electronegative atom is F • EN ↓ down a group • EN↑ across a period from left to right w/ few exceptions

  35. Polar Covalent Bonds: Electronegativity

  36. Problem • Which of the following elements is the most electronegative?  • A.  S • B.  Ru • C.  Si • D.  Te • E.  Cs

  37. Problem • Arrange calcium, rubidium, sulfur, and arsenic in order of decreasing electronegativity.  • A.  S > As > Rb > Ca • B.  S > As > Ca > Rb • C.  As > S > Rb > Ca • D.  As > S > Ca > Rb • E.  None of these orders is correct.

  38. Polar Covalent Bonds: Electronegativity • % Ionic Character: As a general rule for two atoms in a bond, we can calculate an electronegativity difference (∆EN ): ∆EN = EN(Y) – EN(X) for X–Y bond. • If ∆EN < 0.5 the bond is covalent. • If ∆EN 0.5 - < 2.0 the bond is polar covalent. • If ∆EN > 2.0 the bond is ionic.

  39. Problem • Select the most polar bond amongst the following.  • A.  C-O • B.  Si-F • C.  Cl-F • D.  C-F • E.  C-I

  40. Molecular Orbital Theory • The molecular orbital (MO) model provides a better explanation of chemical and physical properties than the valence bond (VB) model. • Atomic Orbital: Probability of finding the electron within a given region of space in an atom. • Molecular Orbital: Probability of finding the electron within a given region of space in a molecule.

  41. Molecular Orbital Theory • Additive combinationof orbitals (s) is lower in energy than two isolated 1s orbitals and is called a bonding molecular orbital.

  42. Molecular Orbital Theory • Subtractive combinationof orbitals (s*) is higher in energy than two isolated 1s orbitals and is called anantibonding molecular orbital.

  43. Molecular Orbital Theory • Molecular Orbital Diagram for H2:

  44. Molecular Orbital Theory • Molecular Orbital Diagrams for H2– and He2:

  45. Molecular Orbital Theory • Additive and subtractive combination of p orbitals leads to the formation of both sigma and pi orbitals.

  46. Molecular Orbital Theory • Second-Row MO Energy Level Diagrams:

  47. Molecular Orbital Theory • Bond Order is the number of electron pairs shared between atoms. • Bond Order is obtained by subtracting the number of antibonding electrons from the number of bonding electrons and dividing by 2.

More Related