Covalent Bonding Covalent Bond: a bond where atoms share electrons instead of gaining or losing electrons - A covalent bond that shares one pair of electrons is a single covalent bond. When two pairs of electrons are shared is a double covalent bond… Structural Formulas: how a formula is written to show a covalent bond (a dash is used to represent the two shared electrons) Example: H-Cl

- Covalent compounds are called molecules (subscripts are not necessarily in the lowest whole number ratio like ionic compounds) - The octet rule still applies (so each element still wants to have eight valance electrons) except for hydrogen. Unshared Pairs (Lone Pairs): pairs of electrons that are not involved in a covalent bond

We learned to write electron dot notation so that electrons that share an orbital will be written on the same side of the elemental symbol. - When electrons are given a little bit of energy electrons can move to higher energy sublevels for bonding purposes - This is why carbon makes four bonds

Double and Triple Covalent Bonds Double Bonds: when two pairs of electrons are shared between two atoms Triple Bonds: when three pairs of electrons are shared between two atoms - Atoms will make double or triple bonds to meet the requirements of the octet rule

Diatomic Molecules Because of the need to have eight valance electrons or obey the octet rule, some elements will never exist in nature as lone atoms. Diatomic Molecules – molecules consisting of two atoms of the same type. - H2, O2, F2, N2, I2, Cl2, Br2

Coordinate Covalent Bonds Typically in a bond, each atom contributes one electron to each bond. Coordinate Covalent Bond: a covalent bond where both electrons in the bond are contributed from one atom. - Atoms in polyatomic ions are coordinate covalently bonded.

Bond Dissociation Energy: the energy it takes to break a bond. - The more energy it takes to break a bond the stronger that bond is.

Resonance Structures Sometimes it is possible to draw the dot structures for one molecule in more than one way. - The molecule does not flip back and forth between the two structures; it is an average between the two. - The more resonance structures a molecule has the more stable that element is.

Exceptions to the Octet Rule When electrons share an orbital the have opposite spins. If a molecule has only paired electrons the spins will cancel out and it will be diamagnetic Paramagnetic: when a molecule has unshared electrons the spin of the electrons will not cancel each other out. - These molecules show a strong attraction in a magnetic field

Other Exceptions to the Octet Rule Other atoms may expand there octet and take on 10 or 12 electrons.

Properties of Covalent Molecules - nonmetalic - Could be a solid, liquid, or gas - Low melting points - Low solubility - Poor electrical conductors

Bonding Theories VSEPR Theory: Valence Shell Electron Pair Repulsion Theory - Electron pairs repel - Electron pairs are as far away from other pairs of electrons as possible in the geometric structure of the molecule

Shapes of Molecules Because each pair of electrons wants to be as far from another pair as possible, we can predict the shape of the molecule. Possible shapes: Linear: Pyramidal: Tetrahedral: Bent: Trigonal Planar: Trigonal Bypyramidal:

Hybridization: when electrons move between two different orbitals to make “hybrid” orbitals

Polar Bonds Nonpolar Covalent Bonds: the atoms pull the electrons toward them equally Example: the diatomic molecules Polar Covalent Bonds: two different atoms pull (share) the electrons unequally - Polarity all depends on the electronegativity that each atom has in the molecule or their attraction for electrons.

Look at page 405 in your book. - This is a table of electronegativities - To determine if the bond between two atoms is polar subtract the electronegativities of the two atoms involved in the bond.

Polar Molecules: molecules will become polar if one end of the molecule is more electronegative then the other end - All of the electrons will be pulled to the atom that is most electronegative - This will make the molecule negatively charged on one end and positively charged at the other end.

Rules of Thumb for Polar Molecules - If a bond in a two atom molecule is polar then the molecule is polar. - If the molecule is more than two atoms, then you must also look at the shape to determine if the molecule is polar. - Symmetric shapes, where the polar bonds cancel each other out lead to nonpolar molecules.

Attractions Between Molecules Molecules have attractions towards other molecules. - These attractions are not as strong as bonds Van der Waals Forces: 1. Dispersion Forces: the attractions that occur with nonpolar molecules - This happens because the nucleus of each atom is attracted to the electrons of other molecules

2. Dipole Attractions: the attractions between polar molecules. - The slightly positive end of one molecule is attracted to the slightly negative end of another molecule.

Hydrogen Bonds (a type of dipole attraction): Attractive forces that occur between the hydrogen of one molecule and the highly electonegative element of another molecule when hydrogen is covalently bonded to a highly electronegative element.

This is why water has such an unexpected high boiling point and why it gets larger when it freezes. - To boil water you have to add more energy to break all of the hydrogen bonds. - When water freezes the molecules of water are all organized with hydrogen bonds. This makes water expand when freezing. Hydrogen bonding is the strongest of the intermolecular forces.

Although most molecular compounds are unstable there is an exception: Network Solids (Network Crystals): solids in which all the atoms are covalently bonded to eachother. - These are very hard substances such as diamonds