Chapter 9 Covalent Bonds Read pgs

Covalent Bonds Covalent bonds form between atoms that share electrons. Covalent bonds form between two or more nonmetals. Shared electrons can be used or counted for both atoms in a bond. A pair of atoms can share more than one pair of valence electrons.

Multiple Bonds If a pair of atoms share 4 electrons, we call it a double bond. Share 6 electrons, a triple bond. The more electrons shared, shorter and stronger bond Strength of a bond depends on – Size of atoms Length of bonds – Number of shared electrons Shared bonds between the nucleii of atoms are called sigma bonds. Double and Triple bonds form outside of the central axis and are called Pi bonds.

Naming Covalent Compounds The rules for naming covalent compounds are very similar to ionic. Name the 1 st element Name the 2 nd element and change its ending to –ide. For covalent compounds we add a prefix to show how many of each element that is in the compound.

Prefixes – Mono – 1Di or Bi – 2 – Tri – 3Tetra – 4 – Penta – 5Hexa – 6 – Hepta – 7Octa – 8 The only exception to the rule is on the first element. If you only have one of the first element, you don’t use the mono- prefix.

Examples Name the following covalent compounds COCO 2 Carbon MonoxideCarbon Dioxide NCl 3 CCl 4 Nitrogen TrichlorideCarbon Tetrachloride SF 6 N 2 S 5 Sulfur Hexafluoride Dinitrogen Pentasulfide

Naming Acids We have already talked about acids. – Corrosive liquids – pH’s less than 7. – Produce H + ions in water. All acids start with hydrogen. You can have two types of acids: Binary, H and one other element, or Ternary, H with a polyatomic ion.

Naming Binary Acids – Use the prefix Hydro- – Name 2 nd element and change ending to –ic acid. Examples HCl – Hydrochloric Acid HF – Hydrofluoric Acid H 3 P – Hydrophosphoric Acid Phosphorus picks up an –or Sulfur picks up an - ur

Naming Ternary Acids – Identify the polyatomic ion. – If the ion ends in –ate, change ending to –ic Acid. – If the ion ends in –ite, change ending to –ous Acid. Examples HNO 3 – nitrate – Nitric Acid HNO 2 – nitrite – Nitrous Acid H 2 SO 4 – sulfate – Sulfuric Acid

Lewis Structures To show bonding electrons and unshared pairs of electrons in molecules we can use one of two systems: Dot Method or Math Method. The Lewis structure of a molecule will show bonds, the structure of the molecule and eventually the shape of the molecule. Choose one method or the other, but don’t try to do both.

Dot Method Find the number of valence electrons of each element in the molecule. Determine the central atom of the molecule. – The atom with the most single valence electrons or the element you have the least of. Match up single valence electrons to form bonds. – Coordinate covalent bonds form with a pair of electrons. Make sure every element has 8 valence electrons, (except hydrogen – it only needs two ).

Math Method Determine the total number of valence electrons you have in the molecule. Determine how many valence electrons you need for each element to complete its octet. Take the difference between the what you have and what you need and divide by two to get the number of bonds that form. Choose the central atom and distribute the bonds and the remaining electrons.

If an atom needs two electrons and another atom has a pair of electrons it can form a coordinate covalent bond where one atom donates both electrons in the bond. Resonance structures are when there is more than one correct Lewis structure for a molecule. Shapes of molecules play a big part in determining the chemical and physical properties of the molecule.

VSEPR Theory To determine the shape of a molecule we use the VSEPR Theory. VSEPR stands for Valence Shell Electron Pair Repulsion Bonding electrons have less repulsion than nonbonding electrons. The nonbonding electrons push and move the other bonds around to reduce repulsive forces.

There can be exceptions to the octet rule. Molecules can for with atoms having less than 8 valence electrons or more than 8 valence electrons. Because electrons are constantly moving around the atom it can create momentary positive or negative areas around the molecule. These weak forces of attraction are called dispersion or Van der Waal’s forces.

Polar Bonds and Molecules Stronger attractions are created when the positive/negative poles in a molecule are more constant. These dipoles are created when there is unequal sharing of electrons in a covalent bond. To determine if a bond is polar look at the difference in electronegativities.

If the difference in E.N. is less than.4, than the bond is nonpolar. The electrons are shared equally, no force of attraction. If the difference in E.N. is between.4 and 2.0, then the bond is polar. The molecule is dipolar, one end is always negative. If the difference is greater than 2.0, then the bond is ionic, electrons are transferred. Diff < 0.4nonpolar – equal sharing 0.4 ≤ Diff. ≤ 2.0polar – unequal sharing Diff. > 2.0ionic – transfers electrons