Lecture 3 Polar and non-polar covalent bonds Dr. A.K.M. Shafiqul Islam

Both lithium and sodium can lose a single electron to form a cation. Lithium has a single electron in a 2s orbital. If the single electron is lost, then a lithium cation is formed that has a filled outer shell. Sodium has a single electron in a 3s orbital. Sodium loses an electron to yield a sodium cation that has a filled outer shell.

Fluorine and chlorine contain seven electrons in their outer shell. The addition of another electron produces an octet.

Ionic bonds of sodium chloride Sodium gives up an electron to chlorine. The positively charged sodium ions and the negatively charged chloride ions are independent species held together by the attraction of opposite charges, called electrostatic attraction. This results in the formation of an ionic bond.

Structures of sodium chloride (a)Crystalline sodium chloride. (b)The electron-rich chloride ions are red and the electron- poor sodium ions are blue. Each chloride ion is surrounded by six sodium ions, and each sodium ion is surrounded by six chloride ions.

Covalent bonds in fluorine Two fluorine atoms can each attain a filled second shell by sharing their unpaired valence electrons, resulting in a covalent bond.

Covalent bonds of hydrogen Two hydrogen atoms can form a covalent bond by sharing electrons. Each hydrogen acquires a stable, filled first shell

Covalent bond of hydrogen chloride Hydrogen and chlorine can form a covalent bond by sharing electrons. Hydrogen fills its only shell and chlorine achieves an outer shell of eight electrons

Formation of a proton and a hydride A hydrogen atom loses its valence electron to form a proton, a positively charged hydrogen ion. A hydrogen atom gains an electron to form a hydride, a negatively charged hydrogen ion.

The Electronegativities of Selected Elements

Covalent bonds of oxygen, nitrogen, and carbon Because oxygen has six valence electrons, it needs to form two covalent bonds to achieve an outer shell of eight electrons. Nitrogen, with five valence electrons, must form three covalent bonds. Carbon, with four valence electrons, must form four covalent bonds.

Use the symbols  - on the most electronegative atom and the symbol  + on the most electropositive atom Oxygen is the most electronegative and thus is given a  - Carbon is the most electropositive and is given a  +

Polar covalent bonds of hydrogen chloride, water, and ammonia A polar covalent bond is the result of the differences in electronegativities of the two atoms involved in the bond. A polar covalent bond has a slight positive charge on the most electropositive atom and a slight negative charge on the most electronegative atom.

Polar bond of hydrogen chloride The direction of bond polarity is indicated with an arrow. The arrow is drawn along the bond, toward the most electronegative atom.

Bonding types The three major types of bonding found in most compounds are ionic, polar covalent, and nonpolar covalent bonds. An ionic bond is the transfer of electrons. A covalent bond is the sharing of electrons. A polar covalent bond is between atoms of different electronegativity. A nonpolar covalent bond is between atoms of similar electronegativity. Potassium fluoride and sodium chloride are examples of ionic bonds. An O-H bond and an N-H bond are examples of polar covalent bonds. A C-H bond and a C-C bond are examples of nonpolar covalent bonds.

Electrostatic potential maps for lithium hydride, hydrogen, and hydrogen fluoride For lithium hydride, the hydrogen atom has the greater electron density than the lithium atom. In hydrogen fluoride, the hydrogen has less electron density than a hydrogen in a hydrogen molecule.

Explanation of colors on electrostatic potential maps The colors indicate the distribution of charge in the molecule. Red is for electron-rich areas and blue is for electron-deficient areas.

Lewis structures of water, hydronium ion, hydroxide ion, and hydrogen peroxide Lewis structures show the valence electrons represented as dots. Lone-pair electrons or nonbonding electrons are valence electrons not used in bonding. A formal charge is the difference between the number of valence electrons minus the number of lone-pair electrons minus half the number of bonding electrons.

Electrostatic potential maps for the hydronium ion, water, and the hydroxide anion In the hydronium ion, the oxygen has a slight negative charge. In the hydronium ion, the oxygen has a negligible charge. The oxygen has the greater charge in the hydroxide ion.

Lewis dot structures of nitrogen compounds. Nitrogen has five valence electrons. In ammonia, nitrogen has five valence electrons. In the ammonium ion, nitrogen has lost one valence electron. Nitrogen has gained one valence electron in the amide anion. In hydrazine, one valence electron is shared between the two nitrogens.

Lewis structures of carbon compounds Carbon has four valence electrons. Carbon has lost one valence electron in the methyl cation. Carbon has gained one valence electron in the methyl anion. Carbon has four valence electrons in the methyl radical. In ethane, one valence electron is shared between the carbon atoms.

Lewis structures of hydrogen, bromide, bromine, and chlorine Hydrogen has one valence electron. Hydrogen ion has lost one electron. Hydride ion has gained one electron. Hydrogen radical has the valence electron. Bromide ion has gained one electron. Bromide radical has the seven valence electrons. Each bromine atom in the bromine molecule has seven electrons. Each chlorine atom in the chlorine molecule has seven electrons.

Number of covalent bonds with hydrogen, halides, oxygen, nitrogen, and carbon Generally, hydrogen forms one covalent bond, oxygen forms two covalent bonds, nitrogen forms three covalent bonds, halides form one covalent bond, and carbon forms four covalent bonds.

Lewis dot structures of methyl bromide, dimethyl ether, formic acid, methylamine, and nitrogen Hydrogen and bromide each forms one covalent bond. Oxygen forms two covalent bonds. Nitrogen forms three covalent bonds. Carbon forms four covalent bonds. A single bond contains two electrons, a double bond contains four electrons, and a triple bond contains six electrons.

Lewis structures of methyl bromide, dimethyl ether, formic acid, methylamine, and nitrogen A pair of shared electrons can be shown as a line between two atoms. In formic acid, two electron pairs are illustrated with a double bond. In nitrogen, three electron pairs are shown with a triple bond.

Drawing Lewis structures For HNO 2, 18 valence electrons are present. Place the valence electrons around the atoms that supply them. Check to see if all atoms have an octet. In the first structure, nitrogen does not have an octet. A pair of electrons from the oxygen is used to make a double bond between the nitrogen and the oxygen to give nitrogen an octet.

Kekulé structures of methyl bromide, dimethyl ether, formic acid, methylamine, and nitrogen In Kekulé structures, the bonding electrons are drawn as lines and the lone-pair electrons are left out entirely.

Condensed structures of methyl bromide, dimethyl ether, formic acid, methylamine, and nitrogen Condensed structures are simplified structures that omit some of the covalent bonds and listing atoms bonded to a particular carbon or other atom.

Kekulé and condensed structures of 2-bromo-5-chlorohexane In the Kekulé structure, all of the covalent bonds are shown. In the condensed structure, the atoms other than hydrogen are shown hanging from the carbon.

Kekulé and condensed structures of hexane In the Kekulé structure, all of the covalent bonds are shown. In the condensed structure, the repeating CH 2 groups can be shown in parentheses.

Kekulé and condensed structures of 4-methyl-2-hexanol In the Kekulé structure, all of the covalent bonds are shown. In the condensed structure, groups bonded to a carbon can be shown in parentheses to the right of the carbon or hanging from the carbon.

Kekulé and condensed structures of 4,4-dimethyl-1-pentanol In the Kekulé structure, all of the covalent bonds are shown. In the condensed structure, groups bonded to a carbon can be shown in parentheses to the right of the carbon or hanging from the carbon. Groups bonded to the far-right carbon are not put in parentheses.