Covalent Bonding In nature, only the noble gas elements exist as uncombined atoms. They are monoatomic - consist of single atoms. All other elements need to lose or gain electrons To form ionic compounds Some elements share electrons To form molecular compounds

Covalent Bonding A neutral group of atoms held together by covalent bonds is called a molecule Some molecules are made up of the same element Those molecules are called diatomic elements 7 naturally occurring diatomics are : H, N, O, F, Cl, Br, I

Covalent Bonding Molecules can also be made of atoms of different elements. A compound composed of molecules is called a molecular compound. A molecular formula gives the “recipe” for a molecular compound.

Covalent Bonding A molecular formula reflects the actual number of atoms in each molecule. The subscripts are not necessarily the lowest whole-number ratios. For example, the formula for peroxide is H2O2 Each molecule of peroxide contains 2 hydrogen atoms and 2 oxygen atoms

Covalent Bonding In covalent bonds, electron sharing usually occurs so that atoms attain the electron configurations of noble gases. The octet rule states that chemical compounds form so each atom (through gaining, losing, or sharing electrons) will have 8 valence electrons Exception: atoms trying to be like helium

Covalent Bonding We use Lewis Dot Diagrams to show covalent bonding However, we do not need to put the dots in the same order as before We need to put them in singles before we can pair them up

Covalent Bonding In the F2 molecule, each fluorine atom contributes one electron to complete the octet. Notice that the two fluorine atoms share only one pair of valence electrons. That is a single covalent bond When we show the bonding, we use Lewis structures Structural formulas are a neater way to show bonding

Covalent Bonding A pair of valence electrons that is not shared between atoms is called an unshared pair In F2, each fluorine atom has three unshared pairs of electrons.

Covalent Bonding A double covalent bond is a bond that involves two shared pairs of electrons.

Covalent Bonding Similarly, a bond formed by sharing three pairs of electrons is a triple covalent bond.

Covalent Bonding Even when electrons are being shared, the sharing is not equal The bonding pairs of electrons in covalent bonds are pulled between the nuclei of the atoms sharing the electrons.

Covalent Bonding When the atoms in the bond pull equally, the bonding electrons are shared equally, and each bond formed is a nonpolar covalent bond A nonpolar covalent bond always results when an element bonds with itself

Covalent Bonding A polar covalent bond, is a covalent bond between atoms in which the electrons are shared unequally. The more electronegative atom attracts more strongly and gains a slightly negative charge. The less electronegative atom has a slightly positive charge. δ+ δ– H—O

Covalent Bonding The electronegativity difference between two atoms tells you what kind of bond is likely to form. Electronegativity Differences and Bond Types Electronegativity difference range Most probable type of bond Example 0.0–0.30 Nonpolar covalent H—H (0.0) 0.31–2.00 Polar covalent δ+ δ– H—F (1.9) >2.00 Ionic Na+Cl– (2.1)

Covalent Bonding The polar nature of the bond may also be represented by an arrow pointing to the more electronegative atom. H—O

Covalent Bonding Electron dot structures fail to reflect the three-dimensional shapes of molecules. The electron dot structure and structural formula of methane (CH4) show the molecule as if it were flat and merely two-dimensional. Methane (structural formula) (Lewis Structure)

Covalent Bonding To determine the 3D shape of the molecule, we use VSEPR (valence shell electron pair repulsion) theory The theory states that repulsion between sets of valence level electrons surrounding an atom causes these sets to be oriented as far apart as possible Or simply, unshared pairs of electrons want to be as far apart as possible

Covalent Bonding The hydrogens in a methane molecule are at the four corners of a geometric solid called a regular tetrahedron. In this arrangement, all of the H–C–H angles are 109.5°, the tetrahedral angle.

Covalent Bonding The molecule ammonia (NH3) is trigonal pyramidal shape. However, one of the valence-electron pairs is an unshared pair and it repels the bonding pairs, pushing them together. The measured H—N—H bond angle is only 107°, rather than the tetrahedral angle of 109.5°. Unshared electron pair 107°

Covalent Bonding The water molecule is planar (flat) but bent. With two unshared pairs repelling the bonding pairs, the H—O—H bond angle is compressed to about 105°. 105°

Covalent Bonding CO2 is a linear molecule The carbon in a carbon dioxide molecule has no unshared electron pairs. The double bonds joining the oxygens to the carbon are farthest apart when the O=C=O bond angle is 180° Carbon dioxide (CO2) No unshared electron pairs on carbon 180°

Covalent Bonding In addition to bond polarity, there is also molecular polarity Some molecules may have polar bonds, but the molecule is overall nonpolar Linear Trigonal planar Tetrahedral

Bonding Many factors contribute to the properties of compounds Bond type Shape Intermolecular Forces

Ionic Bonding An ionic compound is a compound composed of cations and anions. Although they are composed of ions, ionic compounds are electrically neutral. Anions and cations have opposite charges and attract one another by means of electrostatic forces. The electrostatic forces that hold ions together in ionic compounds are called ionic bonds.

Covalent Bonding Molecules can be attracted to each other by a variety of different forces. Intermolecular attractions are weaker than either ionic or covalent bonds. Among other things, these attractions are responsible for determining whether a molecular compound is a gas, a liquid, or a solid at a given temperature.

Covalent Bonding Dipole interactions occur when polar molecules are attracted to one another. The electrical attraction occurs between the oppositely charged regions of polar molecules. Dipole interactions are similar to, but much weaker than, ionic bonds.

Covalent Bonding Hydrogen bonds are attractive forces in which a hydrogen covalently bonded to a very electronegative atom is also weakly bonded to an unshared electron pair of another electronegative atom. The other atom may be in the same molecule or in a nearby molecule Hydrogen bonds are the strongest of the intermolecular forces.

Covalent Bonding Dispersion forces are caused by the motion of electrons. They occur even between nonpolar molecules. When the moving electrons happen to be momentarily more on the side of a molecule closest to a neighboring molecule, their electric force influences the neighboring molecule’s electrons to be momentarily more on the opposite side. The strength of dispersion forces generally increases as the number of electrons in a molecule increases.

Covalent Bonding Fluorine and chlorine have relatively few electrons and are gases at ordinary room temperature and pressure because of their especially weak dispersion forces. Bromine molecules therefore attract each other sufficiently to make bromine a liquid under ordinary room temperature and pressure. Iodine, with a still larger number of electrons, is a solid at ordinary room temperature and pressure.

Characteristics of Ionic and Molecular Compounds Bonding This table summarizes some of the characteristic differences between ionic and covalent (molecular) substances. Characteristics of Ionic and Molecular Compounds Characteristic Ionic Compound Molecular Compound Representative unit Formula unit Molecule Bond formation Transfer of one or more electrons between atoms Sharing of electron parts between atoms Type of elements Metallic and nonmetallic Nonmetallic Physical state Solid Solid, liquid, or gas Melting point High (usually above 300°C) High (usually below 300°C) Solubility in water Usually high High to low Electrical conductivity of aqueous solution Good conductor Poor to nonconducting

Bonding Array of sodium ions and chloride ions Collection of water molecules Formula unit of sodium chloride Molecule of water Chemical Formula H2O NaCl Chemical Formula