Water Chapter 6 2 Covalent Bonding and Molecular

Water Chapter 6. 2 Covalent Bonding and Molecular Compounds Fructose Carbon Dioxide Ammonia

Why Do Atoms Bond? ¡ ¡ To get eight valence electrons To become more stable In ionic bonds, metals lose electrons and non-metals gain electrons. What happens when both elements need electrons?

Molecules and Molecular Compounds that are NOT held together by an electrical attraction, but instead by a sharing of electrons. ¡ Atoms held together by sharing electrons and filling the outer energy levels are joined by a covalent bond. ¡ NONMETALS ONLY!! - No metals ¡

Molecules and Molecular Compounds A molecule is a neutral group of atoms joined together covalent bonds. A compound composed of molecules is called a molecular compound. ¡ The chemical formula for a molecule is called the molecular formula. ¡ A chemical formula tells you how many atoms of each element one molecule of a compound contains. ¡

Learning Check Indicate whether a bond between the following would be 1) Ionic 2) covalent Ionic ____ Covalent ____ A. sodium and oxygen B. nitrogen and oxygen C. phosphorus and chlorine D. calcium and sulfur E. chlorine and bromine

Monatomic (One Atom) Noble gases are monatomic. ¡ They exist as single atoms and do not combine with any other elements. ¡ Ex: He, Ne, Ar, Kr, Xe, Rn ¡

7 Diatomic Molecules Some elements will covalently bond to themselves to form a molecule composed of TWO atoms. ¡ Some elements occur as “diatomic” molecules in nature because they are more stable than individual atoms ¡ The 7 diatomic elements are all gases: H 2, O 2, N 2, Cl 2, Br 2, I 2, F 2 ¡

Strength of Covalent Bonds ¡ ¡ ¡ Distance between two bonding nuclei at the position of maximum attracting is bond length Bond length is determined by the size of the atoms and how many electron pairs are shared Bond energy is the energy required to break a chemical bond and form neutral isolated atoms.

Octet Rule in Covalent Bonds Remember that all compounds want to attain the electron configuration of noble gases. ¡ Hydrogen only needs 2, the rest need 8. ¡ Regarding covalent bonds, electrons are shared between the atoms so that they attain the electron configuration of noble gases. ¡

Exceptions to the Octet Rule A small group of molecules has an odd number of valence electrons and cannot form an octet around each atom -Ex: NO 2 2) O N O 1) Fewer than eight electrons: BORON is stable with 6! H-B-H - Ex: BH 3 H

Exceptions to the Octet Rule Some central atoms have more than eight valence electrons 3) - Referred to as an “expanded” octet Explained by d-orbitals PCl 5 (10 e-) SF 6 (12 e-)

Drawing Valence Electrons ¡“Electron-dot notation”: Electrons are represented as dots located around the symbol of the element. You must put one electron on each side before you double up. X Examples: Nitrogen = N Carbon = H Hydrogen = C

In-Class Examples ¡ Chlorine ¡ Neon ¡ Magnesium ¡ Sulfur ¡ Silicon

Homework ¡ Page 209 #10 -12, 15, 19

6. 2 Notes Continued

To draw Lewis structures for covalent bonds, use the NASB method: l l l N (Needed): Find the number of electrons needed to form full octets for all elements. For most nonmetals, they need 8. Hydrogen needs only 2. A (Available): Find the number of electrons available by adding up all of the valence electrons for all elements involved. S (Shared): Subtract the two numbers. S= N-A B (Bond): A bond is formed with 2 electrons, so divide by 2 to find how many bonds to draw between the elements. Draw the molecule. Put first atom in the center. H’s are always outside. Draw in the bonds, then fill in the rest of the electrons. Check to ensure all atoms have a full octet.

Draw the Lewis-dot-structure for the following molecules 1. HF 2. CCl 2 H 2

Draw the Lewis-dot-structure for the following molecules 1. H 2 O 2. CO 2

Types of Bonds one _____ pair • Each bond involves the sharing of _____ of electrons. 2 e-’s Single Bonds= __ 4 e-’s Double Bonds= __ 6 e-’s Triple Bonds=__

Resonance Structures ¡ ¡ Occurs when more than one valid Lewis Structure can be written for a molecule or ion Differ only in the position of electron pairs, never the atoms positions Actual molecule behaves as if it has one structure Example: O 3

Homework ¡ 6. 2 page 209 #16 -18, 21, 23