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Bonding: General Concepts

Learn about valence electrons, types of chemical bonds (covalent & ionic), bond lengths, polar vs. nonpolar bonds, and bond polarity. Explore dipole moments and charge distributions in molecules.

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Bonding: General Concepts

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  1. Bonding: General Concepts

  2. Group e- configuration # of valence e- ns1 1 1A 2A ns2 2 3A ns2np1 3 4A ns2np2 4 5A ns2np3 5 6A ns2np4 6 7A ns2np5 7 Valence electrons are the outer shell electrons of an atom. The valence electrons are the electrons that participate in chemical bonding.

  3. Types of Chemical Bonds • What is a chemical bond? • Define bonds as the forces that hold groups of atoms together and make them function as a unit. • Why do atoms bond? • A bond will form if the potential energy of the bonded atoms is lower than that of the separated atoms. • Two principal classes of bonding forces. • Covalent bonding: occurs in molecules and involves the sharing of electrons. • Ionic bonding: involves the transfer of electrons between atoms; this produces ions.

  4. In Figure (a) two hydrogen atoms are being brought together to form a bond. • Two unfavorable potential energy interactions: proton-proton repulsion and electron-electron repulsion, and one favorable interaction: proton-electron interaction.

  5. When will the H2 molecule be favored over the separated hydrogen atoms? • Remember nature favors lower potential energy so the hydrogen atoms will position themselves to achieve the lowest possible energy.

  6. The distance where the energy is minimal is called the bond length. • Bond length: the distance between the nuclei of the two atoms connected by a bond or the distance where the total energy of a diatomic molecule is minimal.

  7. Types of Chemical Bonds • In ionic bonding the participating atoms are so different that one or more electrons are transferred to form oppositely charged ions, when then attract each other. • In covalent bonding (also called nonpolar covalent bonding) two identical atoms share electrons equally. • There are intermediate cases where the atoms are not so different that electrons are completely transferred but are different enough that unequal sharing results, forming what is called a polar covalent bond. • An example of this type of bond occurs in the hydrogen fluoride (HF) molecule.

  8. When molecules of HF are placed in an electric field, the molecules tend to orient so that the fluoride end is closed to the positive pole and the hydrogen end closest to the negative pole (Figure b). • This suggests the HF molecule has the following charge distribution. • H - F • δ+ δ- • Where δ is used to indicate a fractional charge.

  9. This partial negative and positive charges on the atoms is called bond polarity. • Explanation: electrons in the bond are not shared equally. • The different affinities of atoms for the electrons in a bond are described by a property called electronegativity. • Electronegativity: the ability of an atom in a molecule to attract shared electrons to itself.

  10. Electronegativity generally increases from left to right across a period and decreases down a group for the representative elements. • Range is from 4.0 for fluorine (most electronegative element) to 0.7 for cesium.

  11. For identical atoms (no electronegativity difference), the electrons in a bond are shared equally, and no polarity develops (nonpolar covalent). • For atoms with very different electronegativity values, electron transfer occurs (ionic bond). • Intermediate cases give polar covalent bonds with unequal electron sharing.

  12. Bond Polarity and Dipole Moments • A molecule such as HF which has a center of positive charge and a center of negative charge is said to be dipolar, or to have dipole moment. • The dipolar character of a molecule is often represented by an arrow pointing to the negative center charge with the tail of the arrow indicating the positive center of charge. H - F δ+ δ-

  13. Electrostatic potential diagrams can be used to represent charge distribution. • Colors of visible light are used to represent the variation in charge distribution. • Red indicates the most electron-rich region of the molecule and blue indicates the most electron-poor region.

  14. Any diatomic (two-atom) molecule that has a polar bond will show a molecular dipole moment. • In other words, if there is a difference in electronegativity between the two atoms the molecule is polar. • H - Cl • But what about polyatomic (more than two atoms) molecules?

  15. For example, the oxygen atom in water has a greater electronegativity value than hydrogen and the charge distribution is shown in Figure (a) below. • In Figure (b), when placed in an electric field the water molecule behaves as if it has two centers of charge. • Thus water has a dipole moment.

  16. This type of behavior is observed for the NH3 (ammonia) molecule also.

  17. Some molecules have polar bonds but do not have a dipole moment. • An example is the CO2 molecule, which is a linear molecule that has the charge distribution shown below. • The opposing bond polarities cancel out and the CO2 molecule does not have a dipole moment. • In other words, it is a nonpolar molecule.

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