Polar Bonds and Molecules Prentice-Hall Chapter 8.4 Dr. Yager

Objectives Describe how electronegativity values determine the distribution of charge in a polar molecule. Describe what happens to polar molecules when they are placed between oppositely charged metal plates. Evaluate the strength of intermolecular attractions compared with the strength of ionic and covalent bonds. Identify the reasons why network solids have high melting points.

What does a Polar Bond look like?

Snow covers approximately 23 percent of Earth’s surface Snow covers approximately 23 percent of Earth’s surface. Each individual snowflake is formed from as many as 100 snow crystals. The polar bonds in water molecules influence the distinctive geometry of snowflakes.

The bonding pairs of electrons in covalent bonds are pulled by the nuclei.

Bond Polarity nonpolar covalant bond Bond polarity is the distribution of charge between two atoms. Electronegativity values determine the charge distribution in a polar bond. Covalent bonds differ in how they share the electrons: nonpolar covalant bond polar covalant bond

Nonpolar Covalent Bond When the atoms in a bond pull equally (as occurs when identical atoms are bonded), the bonding electrons are shared equally, and the bond is a nonpolar covalent bond.

Polar Covalent Bond A polar covalent bond is a covalent bond between atoms in which the electrons are shared unequally. The more electronegative atom attracts electrons more strongly and gains a slight negative charge. The less electronegative atom has a slightly positive charge.

Polar Covalent Bond The chlorine atom attracts the electron cloud more than the hydrogen atom does.

HF is similar to HCl

Water

Water

Water

What are the following bond types? H and Br K and Cl C and O Cl and F Li and O Br and Br Moderately Polar Ionic Moderately to Very Polar Nonpolar

Polar Molecules The presence of one or more polar bonds in a molecule often makes the entire molecule polar. In a polar molecule, one end is slightly positive and the other end is slightly negative. A molecule with two poles (charged regions) is called a dipole.

Polar Molecules The number and orientation of polar bonds determine the overall polarity of a molecule. Consider CO2, a linear molecule with two sets of polar covalent bonds: :O = C = O: .. .. Because the bond polarities are in opposite directions, they cancel. CO2 is a non-polar molecule. Water (H2O) is a bent molecule with two polar covalent bonds. The bond polarities do not cancel, and water is a polar molecule.

Polar Molecules The shape of a molecule and the orientation of its polar bonds determines molecule polarity! When polar molecules are placed between oppositely charged electrical plates, they orient themselves with respect to the negative and positive plates.

A hydrogen chloride molecule is a dipole. Negative and positive plates caused by a battery.

Does a CO2 molecule orient in an electric field? Why/ why not? Although CO2 is made up of two polar bonds, the overall molecule is linear at 180o.

Attraction Between Molecules Intermolecular attractions are weaker than either ionic or covalent bonds. These attractions are responsible for determining whether a molecular compound is a gas, a liquid, or a solid at a given temperature. For example, water is liquid at room temperature.

van der Waals Forces The two weakest attractions between molecules are collectively called van der Waals forces, named after the Dutch chemist Johannes van der Waals (1837–1923).

Van der Waals Forces Dipole interactions: Opposite charges attract Dispersion forces: Weakest of all molecular interactions, caused by the motion of electrons. Generally increases with the number of electrons.

Hydrogen Bonds Hydrogen bonds are attractive forces in which a hydrogen atom covalently bonded to a very electronegative atom is also weakly bonded to an unshared electron pair of another electronegative atom.

Hydrogen Bonds, H2O

The relatively strong attractive forces between water molecules cause the water to form small drops on a waxy surface.

Kevlar

DNA Base Pairs

Intermolecular Attractions and Molecular Properties Network solids (or network crystals) are solids in which all of the atoms are covalently bonded to each other. Network solids consist of molecules that do not melt until the temperature reaches 1000°C or higher, or they decompose without melting at all.

Melting a network solid would require breaking covalent bonds throughout the solid. Diamond is an example of a network solid. Diamond does not melt. It vaporizes to a gas at 3500°C or above.

Silicon carbide is a network solid Silicon carbide is a network solid. It has a melting point of about 2700°C.

In a molecule, the atom with the largest electronegativity value repels electrons more strongly and acquires a slightly negative charge. repels electrons more strongly and acquires a slightly positive charge. attracts electrons more strongly and acquires a slightly positive charge. attracts electrons more strongly and acquires a slightly negative charge.

In a molecule, the atom with the largest electronegativity value repels electrons more strongly and acquires a slightly negative charge. repels electrons more strongly and acquires a slightly positive charge. attracts electrons more strongly and acquires a slightly positive charge. attracts electrons more strongly and acquires a slightly negative charge.

When polar molecules are placed between oppositely charged plates, the negative molecules stick to the positive plates. molecules stick to the negative plates. ends of the molecules turn toward the positive plates. ends of the molecules turn toward the negative plates.

When polar molecules are placed between oppositely charged plates, the negative molecules stick to the positive plates. molecules stick to the negative plates. ends of the molecules turn toward the positive plates. ends of the molecules turn toward the negative plates.

3. Which of the following bond types is the weakest? ionic bond van der Waals force covalent bond hydrogen bond

3. Which of the following bond types is the weakest? ionic bond van der Waals force covalent bond hydrogen bond