Ch 11: Introduction to Chemical Bonding

Ch 11: Introduction to Chemical Bonding

Few atoms exist as independent particles. Most exist as combinations held together by chemical bonds. The way in which they bond affects the properties of the substance.

Atoms form bonds by obtaining stable electron configurations (i.e., a full valance shell of electrons), either by sharing electrons or by transferring them.

11.1 Types of Bonds

1. A chemical bond is a force that holds two or more atoms together and makes them function as a unit.

a. Some bonds are easily broken others are very hard to break.

b. Bond Energy – energy required to break a bond.

2. Ionic Bonds

a. Result from the electrical attraction between a metal ion (cation) and a nonmetal ion (anion).

b. An ionic compound is formed when an atom that loses electrons easily (a metal) reacts with an atom that has an affinity for electrons (a nonmetal). One or more electrons are transferred from the metal to the nonmetal.

c. There is generally a large electronegativity difference between the atoms (see section 11.2).

3. Covalent bonds

a. Result from the sharing of electron pairs between two nonmetal atoms.

b. One or both of the atoms can contribute the electrons to be shared.

c. Atoms share electrons because they have similar electronegativities (see figure 11-1 and section 11.2).

d. Shared electron pairs are considered to be localized between two atoms because this is where they spend most of their time.

e. There are two kinds of covalent bonds:

(1) Nonpolar Covalent (usually simply called “covalent”)

(a) Non-polar bondingresults when twoidentical non-metals equally shareelectrons between them. This type of bond is also formed between carbon and hydrogen. The atoms have the same or almost identical electronegativities. (See fig 11.1)

(b) Examples include H2, Cl2, N2, etc.

(2) Polar covalent bonds

(a) Result from an unequal sharing of electrons. The atom with the higher electronegativity attracts the electrons more strongly than the atom with less electronegativity, resulting in an unequal sharing of electrons.

(b) Shared pairs of electrons are shifted from the center between the two participating atoms, making one end of the molecule slightly positive and the other end slightly negative. The bond is polarized.

(c) Ex. Hydrogen fluoride (see fig 11-2)

(d) Delta (δ) indicates a partial or fractional charge.

(e) It is important to remember that polar molecules (unless ions) are overall neutral. The partial charges caused by the unequal sharing of electrons are not the same as ionic charges which are caused by the actual transfer of electrons between atoms.

4. Hydrogen Bond.

a. A weak attraction between molecules that occurs when a hydrogen atom is bonded to a small, highly electronegative atom, such as N, O, or F.

b. The shared electrons are pulled closer to the more electronegative atom, producing a partial positive charge on the hydrogen and a partial negative charge on the more electronegative atom.

c. The hydrogen end of one molecule has an attraction to the more electronegative side of an adjacent molecule.

5. Metallic bonds

a. Bonding between metal atoms in pure metals or alloys (combinations of two or more different metals).

b. Metal atoms are difficult to separate but can slide past each other fairly easily. Bonding in metals is strong but nondirectional, meaning the bonds occur in any direction.

c. Because metal atoms are relatively large they can lose their outer electrons easily. Large numbers of metal atoms share their valence electrons but in a manner different from covalent bonding.

d. The metal atoms in a sample pool their valence electrons into an evenly distributed “sea” of electrons that “flows” between and around the metal nuclei and core electrons. The electrons are delocalized and move freely throughout the piece of metal.

e. The distinctive properties of metals (malleable, ductile, conductive, high melting points, low ionization energy, etc.) are caused by the larger size of metal atoms, their delocalized electrons, and the way in which metal atoms slide past each other but do not easily separate.

f. In contrast, solid nonmetals are brittle/crumbly because their atoms are smaller with electrons that are not easily lost, nor do they slide past each other.

11.2 Electronegativity

1. Electronegativity is the tendency of an atom to attract shared electrons to itself. (See fig. 11.3)

2. Group 1 and 2 elements lose electrons easily. They have low electronegativities.

3. Group 7 elements, however, can attract electrons easily. They have high electronegativities.

4. Electronegativity generally increases going across the periodic table and decreases as you go down the groups.

5. Using electronegativity values:

a. If the electronegativity difference, DEN, between atoms is zero, the bond is covalent (nonpolar).

b. If the DEN is between 0.4 and 1.8, the bond is considered polar covalent. The more electronegative atoms hold the shared electrons more strongly than the less electronegative. This gives one end of the molecule a partial positive charge and one end a partial negative charge.

c. Bonds between metals and nonmetals generally have a high DEN. Electrons are transferred, not shared, and the bond is ionic.

d. See Fig 11.4 and below:

d. Most bonds have some covalent and some ionic character. An ionic bond simply has more ionic than covalent character, and a covalent bond has more covalent than ionic character. There is no definite dividing line or cut-off between bond types. They exist along a spectrum of ionic vs covalent character present in the bond. See Table 11.1 and below:

11.3 Bond Polarity and Dipole Moments

1. Molecules with polar bonds can have a dipole moment, defined as a property of a molecule whereby the charge distribution can be represented by a center of positive and a center of negative charge

2. The molecule behaves as though it has two centers of charge, one positive and one negative.

3. See example for HF and H2O (Fig 11-5) in text. The dipole moment is generally depicted as an arrow pointing away from the center of positive charge.

4. Water molecules are polar and have a dipole moment.

a. Since a water molecule has a bent shape its individual polar bonds result in a net dipole moment. Notice that the center of positive charge lies between the hydrogen atoms. The center of negative charge is by the oxygen. Since oxygen is more electronegative than the hydrogen the bonding electrons are not shared equally.

b. Because water is polar it is attracted to other water molecules by forming hydrogen bonds. This allows water to remain in the liquid state on the earth’s surface.

c. In contrast, carbon dioxide molecules, which are three times more massive than water molecules, remain in the gaseous state because they are nonpolar and do not attract each other.

d. If water was not polar, it would only exist as a gas at the normal range of temperatures on the Earth’s surface. There would be no liquid water – no oceans, lakes, etc.!! No life!

e. Water is able to dissolve most ionic compounds. They are attracted to positive ions by their negative ends, and to negative ions by their positive ends. (See fig 11.6 and below)

f. Notice how anions are surrounded by the positive ends of water molecules, and cations are surrounded by the negative ends of water molecules.

g. Ions are kept separated as long as water is present. When the water is removed the ions can reform the solid crystal.

5. Molecules can have polar bonds yet be nonpolar.

a. CO2 is an example of a molecule that has polar bonds: O=C=O, but because the two C=O dipole moments cancel each other out the molecule has no net polarity.

b. Carbon tetrachloride, CCl4, is another example of a molecule with polar bonds: C―Cl, but the four resulting dipole moments cancel each other out resulting in a nonpolar molecule with no net dipole moment.

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