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Bonding and Molecular Structure: Fundamental Concepts

Bonding and Molecular Structure: Fundamental Concepts. Chapter 9. Valence Electrons. The electrons involved in bonding are called valence electrons . Valence electrons are found in the incomplete, outermost orbital shell of an atom.

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Bonding and Molecular Structure: Fundamental Concepts

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  1. Bonding and Molecular Structure: Fundamental Concepts Chapter 9 Chapter 9

  2. Valence Electrons • The electrons involved in bonding are called valence electrons. • Valence electrons are found in the incomplete, outermost orbital shell of an atom. • We can represent the electrons as dots around the symbol for the element. • These pictorial representations are called Lewis Structures or Lewis Dot Structures. Chapter 9

  3. Lewis Symbols and the Octet Rule Chapter 9

  4. Chemical Bond Formation • There are three types of chemical bonds • Ionic Bond - electrostatic attraction between ions of opposite charge (NaCl). • Covalent Bond - sharing of electrons between two atoms (Cl2). • Metallic Bond - sharing of electrons between several atoms (Ag). Chapter 9

  5. Ionic Bonding Consider the reaction between sodium and chlorine: Na(s) + ½Cl2(g)  NaCl(s) Chapter 9

  6. Ionic Bonding • Na(s) + ½Cl2(g)  NaCl(s) DH°f = -410.9 kJ • This reaction is very exothermic • Sodium loses an electron to become Na+ • Chlorine gains an electron to become Cl- • Na+ has an Ne electron configuration and Cl- has an Ar configuration Chapter 9

  7. Ionic Bonding • Energetics of Ionic Bond Formation • Lattice Energy (Hlattice) – The energy required to completely separate one mole of a solid ionic compound into its gaseous ions. • Lattice energy depends on • the charge on the ions • the size of the ions • Coulomb’s equation: • Q1, Q2 = charge on ions • k = 8.99 x 109 J-m/c2 • d = distance between ions Chapter 9

  8. Covalent Bonding When similar atoms bond, they share pairs of electrons to each obtain an octet. Example a pair of electrons connect the two nuclei. Chapter 9

  9. Covalent Bonding • Multiple Bonds • It is possible for more than one pair of electrons to be shared between two atoms (multiple bonds) • One shared pair of electrons - single bond (H2) • Two shared pairs of electrons - double bond (O2) • Three shared pairs of electrons - triple bond (N2). • Generally, bond distances decrease as we move from single through double to triple bonds. Chapter 9

  10. Lewis Symbols and the Octet Rule Octet rule – Atoms tend to gain, lose or share electrons until they are surrounded by eight valence electrons. Chapter 9

  11. Drawing Lewis Structures • Draw a skeleton structure of the molecule or ion showing the arrangement of the atoms and the connect each atom to another with a single bond. • Determine the total number of valence elections in the molecule or ion. • Deduct 2 electrons for each single bond used in step 1. • Distribute the rest of the electrons so that each atom (except H) has 8 electrons. • If you are “short” electrons, form multiple bonds • If you have “extra” electrons, one of the heavy atoms may be able to hold more that eight electrons. Chapter 9

  12. Drawing Lewis Structures • PCl3 Chapter 9

  13. Drawing Lewis Structures • PCl3 Chapter 9

  14. Drawing Lewis Structures Isoelectronic Species Molecules or ions having the same number of valence electrons and the same Lewis structure. Chapter 9

  15. Drawing Lewis Structures • Resonance Structures • Some molecules are not well described by Lewis Structures. • Example: Ozone Chapter 9

  16. Drawing Lewis Structures Resonance Structures - Experimentally, ozone has two identical bonds whereas the Lewis Structure requires one single and one double bond. Chapter 9

  17. Drawing Lewis Structures • Resonance Structures • Resonance structures are attempts to represent a real structure that is a mix between several extreme possibilities. • Each Lewis structure is call a Resonance Form • Resonance Form – Two or more Lewis structures having the same arrangements of atoms but a different arrangement of electrons Chapter 9

  18. Drawing Lewis Structures • Resonance Structures • In ozone the resonance forms have one double and one single bond. • The actual structure of O3 is a combination (or average) of the individual forms called a resonance hybrid. Chapter 9

  19. Exceptions to the Octet Rule • There are three classes of exceptions to the octet rule: • Molecules with an odd number of electrons • Molecules in which one atom has less than an octet • Molecules in which one atom has more than an octet Chapter 9

  20. Exceptions to the Octet Rule • Odd Number of Electrons • there are few molecules which fit this category • Examples. ClO2, NO, and NO2 Chapter 9

  21. Exceptions to the Octet Rule • Less than an Octet • This refers to the central molecule • Typical for compounds of Groups 1A, 2A, and 3A. • Examples: LiH, BeH2, BF3 Chapter 9

  22. Exceptions to the Octet Rule • More than an Octet • This starts for atoms in the 3rd period onwards. • This is due to vacant d orbitals which can hold the “extra” electrons. • Another factor is the size of the central atom, as they get bigger, it gets easier to place additional atoms around the central atom. Chapter 9

  23. Molecular Shapes Lewis structures give atomic connectivity (which atoms are connected to which). Chapter 9

  24. Molecular Shapes Molecular Shapes are determined by: Bond Distance – Distance between the nuclei of two bonded atoms along a straight line. Bond Angle – The angle between any two bonds containing a common atom. Chapter 9

  25. Molecular Shapes Chapter 9

  26. Molecular Shapes Valence Shell Electron Pair Repulsion Theory (VSEPR) - VSEPR theory is based on the idea that electrostatic repulsion of the electrons are reduced to a minimum when the various regions of high electron density assume positions as far apart as possible. Chapter 9

  27. Molecular Shapes • Predicting Molecular Geometries • draw the Lewis structure • count the total number of bonding regions and lone pairs around the central atom • arrange the bonding regions and lone pairs in one of the standard geometries to minimize e--e- repulsion • multiple bonds count as one bonding region Chapter 9

  28. Molecular Shapes Predicting Molecular Geometries Common Configuration for saturated molecules. Chapter 9

  29. Molecular Shapes • The “region of electron density” refers to: • Lone pairs • Covalent bonds (single, double, triple) • Remember, you can’t “see” lone-pairs but they do take-up space. Chapter 9

  30. Molecular Shapes Predicting Molecular Geometries Chapter 9

  31. Molecular Shapes Predicting Molecular Geometries Chapter 9

  32. Molecular Shapes Molecules with Expanded Valence Shells Chapter 9

  33. Molecular Shapes To minimize e--e- repulsion, lone pairs are always placed in equatorial positions. Chapter 9

  34. Molecular Shapes Molecules with Expanded Valence Shells Chapter 9

  35. Charge Distribution • Formal Charge • Used to predict the correct Lewis Structure. • Half of the electrons in a bond are assigned to each atom in a bond. • Both electrons of an unshared pair of electrons are assigned to the atoms to which the unshared pair belong. • The formal charge of an atom is equal to the valence electrons minus the number of electrons assigned to each atom. • Formal Charge = (group number) – (assigned electrons) • The sum of the formal charges will equal the charge on the molecule or polyatomic ion. Chapter 9

  36. Charge Distribution • Using Formal Charge • A Lewis structure in which all formal charges in a molecule are equal to zero is preferable to one in which some formal charges are not zero. • If a Lewis structure has non-zero formal charges, the one with the fewest nonzero formal charges is preferred. • A Lewis structure with one large formal charge is preferable to one with several small formal charges. • A Lewis structure with adjacent formal charges should have opposite signs. • When choosing between several Lewis structures, the structure with negative formal charges on the more electronegative atom is preferable. Chapter 9

  37. Bond Polarity and Electronegativity • Electrons in a covalent bond may not be shared evenly. • Electronegativity – The ability of an atom in a molecule to attract electrons to itself. • - The periodic trend for electronegativity is up and to the right across the periodic table. Chapter 9

  38. Bond Polarity and Electronegativity Electronegativity Chapter 9

  39. Bond Polarity and Electronegativity • Electronegativity and Bond Polarity • A chemical bond between elements with large differences in eletronegativity will shift the electrons to the atom with the higher electronegativity. • The positive end (or pole) in a polar bond is represented + and the negative pole -. • This is called a polar covalent bond. • If the electronegativity difference is small, the bond is nonpolar; if it is large, it is a polar bond. Chapter 9

  40. Polarity of Molecules • To determine if a molecule is polar, you need to know two things: • - polarity of the bonds in a molecule • - how the bonds are arranged • A molecule is considered polar if its center of negative and positive charge do not coincide. • Polar molecules have a dipole (a vector quantity) • If these dipoles act equally and in opposition to each other, the dipoles cancel-out and the molecule is considered nonpolar. Chapter 9

  41. Polarity of Molecules Dipole Moments of Polyatomic Molecules Example: CO2, each C-O dipole is canceled because the molecule is linear. H2O, the H-O dipoles do not cancel because the molecule is bent. Chapter 9

  42. Polarity of Molecules • Dipole Moments of Polyatomic Molecules • Two simple rules to help determine molecular polarity (most of the time) • If there are lone pairs on the central atom – the molecule is polar. • If there is more than one type of bond on the central atom – the molecule is polar. Chapter 9

  43. Polarity of Molecules Dipole Moments of Polyatomic Molecules Chapter 9

  44. Strengths of Covalent Bonds • Bond Enthalpy (Energy) - The energy required to break a covalent bond of a gaseous substance. • Cl2(g)  2Cl(g) DH = DCl-Cl • When more than one bond is broken the bond enthalpy is a fraction of H for the atomization reaction : • CH4(g)  C(g) + 4H(g) H = 1660 kJ • DC-H = ¼H = ¼(1660 kJ) = 415 kJ. • Bond enthalpies can either be positive or negative. Chapter 9

  45. Strengths of Covalent Bonds Chapter 9

  46. Strengths of Covalent Bonds • Bond Enthalpies and the Enthalpies of Reaction • Bond enthalpies can be used to calculate Hrxn. • Hrxn = D(bonds broken) -D(bonds formed). Chapter 9

  47. Bond Enthalpies and the Enthalpies of Reaction Hrxn = [1mol(614kJ/mol)+1mol(146kJ/mol)]-[2mol(358kJ/mol)+1mol(348kJ/mol)] = -304 kJ Chapter 9

  48. Bond Enthalpies and the Enthalpies of Reaction Hrxn = [2(3mol(200kJ/mol))]-[1mol(941kJ/mol)+3mol(242kJ/mol)] = -467 kJ Chapter 9

  49. Homework 2, 6, 18, 22, 26, 32, 46, 57 Chapter 9

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