Chapter 6

Chemical Bonding

Objectives:

Define chemical bond.

Explain why most atoms form chemical bonds.

Describe ionic and covalent bonding.

Explain why most chemical bonding is neither purely ionic nor purely covalent.

Classify bonding type according to electronegativity differences.

  1. Section (6-1) Introduction to chemical Bonding
    1. What happens when you mix atoms?
      1. Compounds are formed
        1. Made up of chemical bonds
        2. Bonds- is a mutual electrical attraction between the nuclei and valence electrons of different atoms that binds the atoms together.
    2. Types of Chemical Bonding
      1. Ionic bonding-the electrical attraction between large numbers of cations and anions
      2. Covalent Bonding- results from the sharing of electron pairs between two atoms
      3. Is it ionic or covalent?
        1. It is usually never totally ionic or covalent
        2. They are some where in between
        3. Depends on its electronegativity (the ability to attract)
        4. To determine if ionic or covalent the electronegativity difference must be calculated.
        5. Refer to figure 6-2 pg 162
          1. electronegativity differences 1.7 or less is covalent
          2. Between two atoms of the same element it is totally covalent
          3. This is called a nonpolar bond an equal share between atoms
          4. Polar bonds have an uneven sharing of electrons.
          5. Many times called polar-covalent bonds.
          6. Polar bonds use symbols to indicate electronegativity
          7. d + or d -

     

     

     

    Objectives

    Define molecule and molecular formula.

    Explain the relationships between potential energy, distance between approaching atoms, bond length, and bond energy.

    State the octet rule.

    List the six basic steps used in writing Lewis structures.

    Explain how to determine Lewis Structure for molecules containing single bonds, multiple bonds, or both.

    Explain why scientist’s use resonance structures to represent some molecules.

     

  2. Section 6-2 Covalent Bonding and Molecular Compounds
    1. Definitions of some terms
      1. Molecule- is a neutral group of atoms that are held together by covalent bonds.
      2. Molecular compound-a compound whose simplest units are molecules.
      3. Chemical formula-indicates the relative numbers of atoms of each kind in a chemical compound by using atomic symbols and numerical subscripts.
      4. Molecular formula- shows the types and number of atoms combined in a single molecule of a molecular compound.
      5. Diatomic molecule- is a molecule containing only two atoms.
    2. Formation of a Covalent Bond.
      1. Nature favors bonding because it produces a lower potential energy.
      2. The protons pull on the electrons of another atom.
      3. The protons of one-atom pushes away from the protons of the other atom and the same occurs between electrons.
      4. There is eventually an even pull and push.
    3. Characteristics of the Covalent Bond
      1. Bond length- the distance between two bonded atoms at their minimum potential energy.
      2. Bond energy- the energy required to break a chemical bond and form neutral isolated atoms
      3. Orbitals overlap and create a region that atoms can share electrons.

       

       

       

    4. The octet Rule
      1. Definition-compounds tend to form so that each atom, by gaining or losing, or sharing electrons, has an octet of electrons in its highest occupied energy level.
      2. Atoms want to look like their nearest noble gas.
      3. Exceptions to the octet rule
        1. Boron can have six.
        2. Others can hold more than eight when they have a high electronegativity.
    5. Electron-dot notation
      1. Definition- is an electron-configuration notation in which only the valence electrons of atoms of a particular element are shown.
        1. The inner-shell electrons are represented by the symbol
        2. The outer shell electrons are represented by dots.
        3. Practice below.

       

       

       

       

    6. Lewis Structures
      1. Makes use of electron-dot notation
      2. Definition- formulas in which atomic symbols represent nuclei and inner-shell electrons, dot-pairs in covalent bonds, and dots adjacent to any one atomic symbol represent unshared electrons
        1. Unshared pair-or Lone Pair—a pain of electrons that is not involved in bonding
        2. Shared pair- electrons involved in bonding.
      3. Structural formulas- indicates the kind, number, arrangement, and bonds but no the unshared pairs of the atoms molecules.
      4. A single bond- is one pair of shared electrons
    7. Multiple covalent bonds
      1. Double bond-
        1. Sharing of two pairs
        2. Shown by parallel lines
      2. Triple bonds
        1. Sharing of three pairs
        2. Shown by three parallel lines

       

       

    8. Table 6-2

      9.      
           
           
           
           
           
           

    9. Resonance Structures- refers to bonding in molecules or ions that cannot be correctly represented by a single Lewis structure.

     

     

     

     

     

     

     

     

     

     

     

     

     

     

     

     

    Objectives

    Compare and contrast a chemical formula for a molecular compound with one for an ionic compound.

    Discuss the arrangements of ion in crystals

    Define lattice energy and explain its significance.

    List and compare the distinctive properties of ionic and molecular compounds.

    Write Lewis structure for a polyatomic ion given the identity of the atoms combined and other appropriate information.

  3. Section 6-3 Ionic Bonding and Ionic compounds
    1. Definitions
      1. Ionic compound-is composed of positive and negative ions that are combined so that the numbers of positive and negative charges are equal.
      2. Formula unit – is the simplest collection of atoms from which an ionic compound’s formula can be established.
    2. Formation of Ionic Compounds
      1. Occurs due to a transfer between electrons

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      2. Characteristics of Ionic Bonding
        1. Potential energy must be low
        2. Create an ionic crystal known as a crystal lattice.
    3. A comparison of ionic and molecular compounds
      4.

      Properties

      Covalent

      Ionic

      Bond strength

      		    

      Melting points

      		    

      Physical state

      		    

      Dissolved

      		    

    4. Polyatomic Ions
        1. Definition- a charged group of covalently bonded atoms.
        2. The charge comes from too many or too few electrons.
        3. Ex. pg 180
        4. Draw Lewis structures

     

     

     

     

     

     

     

     

    Objectives

    Describe the electron-sea model of metallic bonding and explain why metals are good electrical conductors

    Explain why metal surfaces are shiny.

    Explain why metals are malleable and ductile but ionic-crystalline compounds are not.

  4. Section 6-4 Metallic Bonding
    1. Definition- the attraction between metal atoms and the surrounding sea of electrons
      1. Metallic Properties-the freedom of electrons to move creates
        1. Good conductors of electricity
        2. Good conductors of heat
        3. Movement of electrons to an excited state – produces shiny surface.
      2. Metallic Bond Strength
        1. Varies with the nuclear charge of the metal
        2. Varies with the number of electrons

     

     

     

     

     

    Objectives

    Explain VSEPR theory.

    Predict the shapes of molecules of polyatomic ions using VESPR.

    Explain how the shapes of molecules are accounted for by hybridization theory.

    Describe dipole-dipole forces, hydrogen bonding, induced dipoles, and London dispersion forces.

    Explain what determines molecular polarity.

     

  5. Section 6-5 Molecular Geometry
    1. Molecular polarity- the uneven distribution of molecular charges is caused by
      1. Polarity of each bond
      2. And geometry of the molecule
    2. VESPR theory
      1. Definition- State that repulsion between the sets of valence-level electrons surrounding an atom causes these sets to be oriented as far apart as possible.
      2. Four main shapes
        1. Linear
        2. Bent
        3. Trigonal-planer
        4. Tetrahedron
      3. Other shapes table 6-5
      4. Examples
    3. Hybridization- is the mixing of two or more atomic orbitals of similar energies on the same atom to produce a new orbitals of equal energies.
    4. Intermolecular forces-
      1. Definition-the force of attraction between molecules.
        1. Are not between atoms but molecules
        2. Are weaker than bonds
      2. Molecular Polarity and Dipole-Dipole forces
        1. Dipole- is created by equal but opposite charges that are separated by a short distance.
          1. Represented by an arrow
          2. Between atoms
        2. Dipole-Dipole forces- the forces of attraction between polar molecules.
          1. Represented by an arrow
          2. Between the molecule
          3. Must take into account the geometry.

         

         

         

         

      3. 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 electron of an electronegative atom in a nearby molecule.
        1. Accounts for water being the universal solvent
        2. And many other properties
      4. London dispersion Forces- the intermolecular attractions resulting from the constant motion of electrons and the creations of instantaneous dipoles.