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Chapter 8 Chemical Bonds - Angelo State University · PDF fileChapter 8: Chemical Bonds (+ VSEPR) 5 Covalent Bonds Share and Share Alike 6 Covalent Bonds and Molecules • Covalent

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  • Chapter 8: Chemical Bonds (+ VSEPR)

    Mr. Kevin A. Boudreaux Angelo State University

    CHEM 1411 General Chemistry Chemistry: The Science in Context (Gilbert, 4th ed, 2015)

    www.angelo.edu/faculty/kboudrea

    Chapter Objectives:

    Understand the principal types of chemical bonds.

    Understand the properties of ionic and molecular compounds.

    Draw Lewis dot structures for molecular compounds, including resonance structures.

    (Chapter 9.1-9.3) Predict the shape of a molecule using VSEPR theory, and determine whether the molecule is polar.

    Chapter 8 Chemical Bonds

    (+VSEPR from Chapter 9)

    2

    Introduction

    Most of the substances that we encounter in daily life are not elemental substances, but compounds (and frequently, complex mixtures of compounds).

    Why do elements form compounds in the first place? Bonding lowers the potential energy between the charged particles that compose atoms.

    There are three major ways of modeling bonds between atoms, with varying degrees of complexity:

    Lewis theory (Lewis dot structures)

    Valence Bond theory

    Molecular Orbital theory

  • Chapter 8: Chemical Bonds (+ VSEPR)

    3

    Types of Chemical Bonds

    There are three types of bonds between elements:

    Ionic bonds result from a transfer of electrons from one species (usually a metal) to another (usually a nonmetal or polyatomic ion).

    Covalent bonds result from a sharing of electrons by two or more atoms (usually nonmetals).

    Metallic bonds result from a pooling of valence electrons by two or more metals into a delocalized electron sea.

    4

    Lewis Dot Structures for Elements

    In a Lewis dot structure [G. N. Lewis] for an element, the valence electrons are written as dots surrounding the symbol for the element.

    Place one dot on each side first; the remaining dots are paired with one of the first set of dots.

    A maximum of two dots are placed on each side.

    The unpaired dots indicate where covalent bonds can form, or where electrons can be gained or lost.

    Atoms tend to form bonds to satisfy the octet rule, having eight electrons in their valence shells.

    Li Be B C N O F Ne

  • Chapter 8: Chemical Bonds (+ VSEPR)

    5

    Covalent Bonds

    Share and Share Alike

    6

    Covalent Bonds and Molecules

    Covalent bonds form when two or more nonmetals share their electrons. The electrons are at their lowest potential energy when they are between the two nuclei that are being joined.

    Each atom in the bond holds on to the shared electrons, and the atoms are thus physically tied together.

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    The Formation of Diatomic Hydrogen

    As two isolated H atoms move closer together, the two positively-charged nuclei repel each other, and the two negatively-charged electrons repel each other, but each nucleus attracts both electrons.

    At some point, the attractions between the nuclei and the electrons are balanced against the repulsions between the nuclei and between the electrons.

    The shared electrons bind the two nuclei into an H2 molecule.

    The shared electrons act like they belong to both atoms in the bond.

    Covalent Bonding and Potential Energy

    When two isolated H atoms approach each other, the potential energy is lowered energy is released when an H2 molecule forms.

    Pushing the nuclei any closer causes the potential energy to rise.

    8 Figure 8.1

  • Chapter 8: Chemical Bonds (+ VSEPR)

    Covalent Bonding and Potential Energy

    The optimum distance between nuclei where the attractive forces are maximized and the repulsive forces are minimized is called the bond length. (For H2, the bond length is 74 pm.)

    In H2, the highest probability of finding electrons is in the space between the H nuclei.

    The increased attractive forces in this area help to lower the potential energy of H2 relative to isolated H atoms.

    9

    Lewis Structures and the Octet Rule

    The sharing of electrons by nonmetals to form molecules is modeled by the use of Lewis structures, in which sticks () represent pairs of shared electrons, and dots (:) represent unshared electrons (lone pairs).

    Atoms share electrons in such a way as to satisfy the octet rule, which gives each atom a total of eight valence electrons.

    Hydrogen is an exception to the octet rule, since its 1s orbital can only hold two electrons.

    Once the Lewis structure has been drawn, the 3D shape of the molecule and the polarity of the molecule can be predicted using the VSEPR model.

    10

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    Single Covalent Bonds

    The shared pairs of electrons are bonding pairs.

    The unshared pairs of electrons are lone pairs or nonbonding pairs.

    Na+ Cl- ionic bond

    HH covalent bond

    + H HorH H H H

    + H ForH F H F

    + F ForF FF F

    + H OorO OH H+ H H H

    12

    Double and Triple Covalent Bonds

    Some atoms can satisfy the octet rule by sharing two pairs of electrons to form a double bond:

    Double bonds are shorter and stronger than single

    bonds.

    Some atoms can share three pairs of electrons to form a triple bond:

    Triple bonds are even shorter and stronger than

    double bonds.

    O+ O OO OO

    + NN N N

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    Lewis Structures

    Drawing the Line

    14

    How To Write Lewis Structures

    1. Determine the total number of valence electrons in the molecule or ion.

    Add one electron for each unit of negative charge.

    Subtract one electron for each positive charge.

    2. Write the correct skeletal structure.

    For molecules of the formula ABn, place the least electronegative in the center, and the remaining atoms on the outside. Draw single bonds between the outer atoms and the central atom.

    H is NEVER a central atom, since it can only form one bond.

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    How To Write Lewis Structures

    3. Distribute the remaining valence electrons as lone pairs on the outer atoms first, making sure to satisfy the octet rule.

    Once all of the outer atoms have 8 electrons (or 2 for H), place any remaining electrons on the central atom.

    4. Compare the number of valence electrons in the Lewis structure to the number in Step 1 to make sure you havent miscounted electrons.

    5. Complete the octet on the central atom.

    This is done by sharing lone pairs from the outer atoms with the central atom.

    The formal charge can be used as a guideline for placing double bonds.

    16

    Formal Charges

    Formal charge is the difference between the number of valence electrons on the free atom and the number of valence electrons assigned to the atom in the molecule.

    The sum of the formal charges must equal the charge on the species.

    Smaller formal charges are better (more stable) than larger ones.

    Like charges on adjacent atoms are not desirable.

    When a formal charge cannot be avoided, negative formal charges should reside on more electronegative atoms.

    Formal charge = valence e- ( bonding e-) (lone pair e-)

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    Examples: Lewis Structures

    1. Write Lewis structures for the following molecules.

    a. CH4

    b. NH3

    c. NH4+

    d. CCl4

    e. H2O2

    18

    Examples: Lewis Structures

    1. Write Lewis structures for the following molecules.

    f. CCl2F2

    g. H2S

    h. C2H6

    i. C2H4

    j. N2O (atoms connected in the order NNO)

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    Examples: Lewis Structures

    1. Write Lewis structures for the following molecules.

    k. CO2

    l. CaCl2

    m COCl2

    n. acetamide, C2H5NO

    CC

    N

    O

    H

    H

    H

    HH

    20

    Examples: Lewis Structures

    1. Write Lewis structures for the following molecules.

    o. O3

    Neither of the two O3 structures is correct by itself.

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    Resonance Structures

    When there is more than one valid Lewis structure for a molecule, the actual electronic structure is an average of the different possibilities, called a resonance hybrid.

    Resonance forms differ only in the placement of the valence electrons, not the positions of the atoms!

    Ozone does not have a real double bond and a real single bond; the actual bond lengths in ozone are identical, and are between those of an OO bond and an O=O bond a one-and-a-half bond.

    When One Lewis Structure Aint Enough

    22

    Resonance Structures This molecule is shown more correctly with two

    Lewis structures, called resonance structures, with a two-headed arrow (T) between them:

    DO NOT USE THE STRAIGHT, DOUBLE-HEADED ARROW (T) FOR ANYTHING ELSE!

    Resonance structures are not real bonding depictions: O3 does not change back and forth between the two structures; the actual molecule is a hybrid of the two forms depicted.

    O

    OO O

    OO

    Bond order = 1 1/2

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    Delocalized Electrons and Charges

    In these resonance structures, one of the electron pairs (and hence the negative charge) is spread out or delocalized over the whole molecule.

    In contrast, the lone pairs on the oxygen in water are localized i.e., theyre stuck in one place.