Chemical Bonding


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Chemical Bonding

  1. 1. Chapter 6 CHEMICAL BONDING
  2. 2. <ul><li>Mutual electrical attraction between the nuclei and valence electrons of different atoms that finds that atom together </li></ul><ul><li>Why are most atoms chemically bonded together? </li></ul><ul><ul><li>Atoms are less stable existing by themselves than when they are combined </li></ul></ul><ul><ul><li>By bonding with each other, atoms decrease in potential energy, which means they create more stable arrangements of matter </li></ul></ul>CHEMICAL BOND
  4. 4. <ul><li>Ionic bonding </li></ul><ul><ul><li>Results from the electrical attraction between cations and anions </li></ul></ul><ul><ul><li>Atoms completely give up electrons to other atoms </li></ul></ul><ul><li>Covalent Boning </li></ul><ul><ul><li>Results from the sharing of electron pairs between two atoms </li></ul></ul><ul><ul><li>The shared electrons are “owned” equally by the two bonded atoms </li></ul></ul>TYPES OF BONDS
  7. 7. <ul><li>Non-polar </li></ul><ul><ul><li>The bonding electrons are shared equally by the bonded atoms </li></ul></ul><ul><ul><li>Results in a balanced distribution of electrical charge </li></ul></ul><ul><ul><li>Example: hydrogen-hydrogen bond </li></ul></ul><ul><li>Polar </li></ul><ul><ul><li>There is an uneven distribution of charge </li></ul></ul><ul><ul><li>The bonded atoms have an unequal attraction for the shared electrons. </li></ul></ul>TYPES OF COVALENT BONDS
  9. 9. <ul><li>Molecule </li></ul><ul><ul><li>Neutral group of atoms that are held together by covalent bonds </li></ul></ul><ul><li>Molecular Compound </li></ul><ul><ul><li>Chemical compound whose simplest units are molecules </li></ul></ul><ul><li>Chemical Formula </li></ul><ul><ul><li>Indicates the relative numbers of atoms of each kind in a chemical compound by using atomic symbols and numerical subscripts </li></ul></ul><ul><li>Molecular Formula </li></ul><ul><ul><li>Shows the types and numbers of atoms combined in a single molecule of a molecular compound </li></ul></ul><ul><li>Diatomic Molecule </li></ul><ul><ul><li>Molecule that contains only two atoms </li></ul></ul>MOLECULAR COMPOUNDS
  10. 10. <ul><li>Bond length </li></ul><ul><ul><li>The average distance between two bonded atoms </li></ul></ul><ul><li>Bond Energy </li></ul><ul><ul><li>The energy required to break a chemical bond and form neutral isolated atoms </li></ul></ul><ul><ul><li>Measure in kilojoules per mole (kJ/mol) </li></ul></ul><ul><ul><li>Example: </li></ul></ul><ul><ul><ul><li>436 kJ/mol of energy is needed to break hydrogen-hydrogen bonds in one mole of hydrogen molecules </li></ul></ul></ul>CHARACTERISTICS OF COVALENT BONDS
  11. 11. <ul><li>Chemical compounds tend to form so that each atom, by gaining, losing, or sharing electrons, has an octet of electrons in its highest occupied energy level </li></ul><ul><li>Exceptions to the rule: </li></ul><ul><ul><li>Hydrogen </li></ul></ul><ul><ul><ul><li>Forms bonds in which it is only surrounded by two electrons </li></ul></ul></ul><ul><ul><li>Boron </li></ul></ul><ul><ul><ul><li>Forms bonds in which it is surrounded by six electrons </li></ul></ul></ul>OCTET RULE
  12. 12. <ul><li>Electron configuration notation in which only the valence electrons of an atom of a particular element are shown </li></ul><ul><li>Valence electrons are indicated by dots placed around the element’s symbol </li></ul>ELECTRON DOT NOTATION
  13. 13. <ul><li>Formulas in which atomic symbols represent nuclei and inner-shell electrons </li></ul><ul><li>Dot pairs or dashes between two atomic symbols represent electron pairs in covalent bonds </li></ul><ul><li>Dots that are adjacent to only one atomic symbol represent unshared electrons </li></ul><ul><li>Unshared pair (lone pair) – a pair of electrons </li></ul><ul><li>that is not involved in bonding and that </li></ul><ul><li>belongs exclusively to one atom. </li></ul>LEWIS STRUCTURES
  14. 14. <ul><li>Indicates the kind, number, arrangement, and bonds but not the unshared pairs of atoms in a molecule </li></ul>STRUCTURAL FORMULAS
  15. 15. <ul><li>Single </li></ul><ul><ul><li>Covalent bond in which one pair of electrons is shared between two atoms </li></ul></ul><ul><li>Double </li></ul><ul><ul><li>Covalent bond in which two pairs of electrons are shared between two atoms </li></ul></ul><ul><li>Triple </li></ul><ul><ul><li>Covalent bond in which three pairs of electrons are shared between two atoms </li></ul></ul><ul><li>Multiple bonds </li></ul><ul><ul><li>Either double or triple bonds </li></ul></ul>BONDS
  16. 16. Multiple Covalent Bonds <ul><li>Double Bond – a covalent bond produced by the sharing of two pairs of electrons between two atoms. </li></ul>Triple Bond – a covalent bond produced by the sharing of three pairs of electrons between two atoms. N N O O N N O O
  17. 17. IONIC BONDING AND IONIC COMPOUNDS <ul><li>Ionic Compound – a compound composed of positive and negative ions that are combined so that the number of positive and negative charges are equal. </li></ul><ul><li>Formula Unit – the simplest collection of atoms from which an ionic compound’s formula can be established. </li></ul><ul><li>Ex. NaCl </li></ul>
  18. 18. FORMATION OF IONIC COMPOUNDS Na + + Cl - NaCl Ca 2+ + F - CaF 2 Ca 2+ + N 3- Ca 3 N 2
  19. 19. CHARACTERISTICS OF IONIC BONDING <ul><li>Crystal lattice – an orderly arrangement of ions. </li></ul><ul><li>Lattice energy – the energy released when one mole of an ionic crystalline compound is formed for gaseous ions. </li></ul><ul><li>Ionic bonds are stronger than molecular bonds. </li></ul>
  20. 20. POLYATOMIC IONS <ul><li>Polyatomic ions – a charged group of covalently bonded atoms. </li></ul>
  21. 21. METALLIC BONDING <ul><li>Within a metal, the vacant orbitals in the atoms’ outer energy level overlap, allowing outer electrons of atoms to roam freely throughout the entire metal. </li></ul><ul><li>Delocalized electrons – electrons that do not belong to one atom, but can freely move about the metal’s network of empty atomic orbitals. </li></ul><ul><li>Metallic Bonding – the chemical bonding that results from the attraction between metal atoms and the surrounding sea of electrons. </li></ul>
  22. 22. <ul><li>The shiny appearance of metals are due to the absorption of a high range of light frequencies, resulting in exciting/de-exciting electrons. </li></ul><ul><li>Malleability – the ability of a substance to be hammered or beaten into thin sheets. </li></ul><ul><li>Ductility – the ability of a substance to be drawn, pulled, or extruded through small opening to produce a wire. </li></ul><ul><li>Heat of Vaporization – the amount of heat required to vaporize a metal, which is the measure of the strength of the bonds that hold metal together. </li></ul>
  23. 23. MOLECULAR GEOMETRY <ul><li>Molecular Geometry – the three-dimensional arrangement of molecule’s atoms in space. </li></ul><ul><li>Molecular Polarity – the uneven distribution of molecular charge. </li></ul>
  24. 24. VSEPR THEORY <ul><li>Valence Shell Electron Pair Repulsion (VSEPR Theory) – the repulsion between the sets of valence-level electrons surrounding an atom cause these sets to be oriented as far apart as possible. </li></ul>
  25. 25. <ul><li>Example: CH 4 (bonding pairs only, no lone pairs) </li></ul>Key: Must consider both bonding and lone pairs in minimizing electron repulsion. Lewis Structure VSEPR Structure
  26. 26. • Example: NH 3 (both bonding and lone pairs). Lewis Structure VSEPR Structure Molecular Shape
  27. 27. VSEPR APPLICATIONS The previous examples illustrate the strategy for applying VSEPR to predict molecular structure: 1. Construct the Lewis Dot Structure 2. Arrange the bonding and lone electron pairs in space such that repulsions are minimized .
  28. 28. Case: Linear Structure ( AX 2 ): angle between bonds is 180° Example: BeF 2 180°
  29. 29. Case: Trigonal Planar Structure ( AX 3 ): The angle between bonds is 120° Example: BF 3 120°
  30. 30. Case: Pyramidal ( AX 3 E ): Bond angles are <120° structure is nonplanar due to repulsion of lone-pair. Example: NH 3 107° VSEPR Structure Molecular shape Lewis
  31. 31. Case: Tetrahedral ( AX 4 ): the angle between bonds is ~109.5° Example: CH 4 109.5°
  32. 32. Note: for ‘Tetrahedral’, the actual angle may vary slightly from 109.5°, due to size differences between bonding and lone pair electron densities bonding pair: more elongated, less repulsive lone pair: puffier, more repulsive
  33. 33. Example of distorted tetrahedron: water ( AX 2 E 2 ): the angle is reduced to 104.5° by repulsion of the lone pairs “ bent” VSEPR structure Molecular shape
  34. 34. Case: Trigonal Bipyramidal ( AX 5 ): non-equivalent bond positions: three in-plane (equatorial, 120°), and two at 90° to plane (axial) Example, PCl 5 90° 120°
  35. 35. Octahedral ( AX 6 ): all angles are 90° Example SF 6 90° Lewis VSEPR
  36. 36. INTERMOLECULAR FORCES <ul><li>Intermolecular Forces – the force of attraction between molecules. </li></ul><ul><li>Dipole – equal and opposite charges that are separated by a short distance. </li></ul><ul><li>Dipole-Dipole Forces – forces of attraction between polar molecules. </li></ul>H - Cl
  37. 37. <ul><li>Hydrogen Bonding – the intermolecular force in which a hydrogen atom that is bonded to a highly electronegative atom is attracted to an unshared pair of electrons of an electromagnetic atom in the nearby molecule. </li></ul><ul><li>London Dispersion Forces – the intermolecular attractions resulting from the constant motion of electrons and the creation of instantaneous dipoles. </li></ul><ul><ul><li>Act between all atoms and molecules </li></ul></ul><ul><ul><li>The only intermolecular forces acting among noble-gas atoms, nonpolar molecules, and slightly polar molecules </li></ul></ul><ul><li>Only intermolecular among noble gasses and non-polar molecules. </li></ul>