bonding

Electronegativity- the ability of an atom to attract a shared pair of electrons to itself. The difference in electro negativities of 2 elements can determine the type of bond that holds them together.

Differences in electronegativity 0 - .4 = covalent bond(nonpolar) .41 – 1.66 = polar covalent bond 1.67 or greater = ionic bond

Chemical bonds The force of attraction between the electrons of one atom and the nucleus of another atom. 3 types Ionic Covalent metallic

Bond Types Sigma bonds () – single covalent bond Pi bonds () – occur when multiple bonds are formed Single bond – sigma Double bond – 1 sigma & 1 pi Triple bond – 1 sigma & 2 pi

Pi Bonding in the Nitrate Ion Click in this box to enter notes. Go to Slide Show View (press F5) to play the video or animation. (To exit, press Esc.) This media requires PowerPoint® 2000 (or newer) and the Macromedia Flash Player (7 or higher). [To delete this message, click inside the box, click the border of the box, and then press delete.] Copyright © Houghton Mifflin Company. All rights reserved.

Pi Bonding and Antibonding Orbitals Click in this box to enter notes. Go to Slide Show View (press F5) to play the video or animation. (To exit, press Esc.) This media requires PowerPoint® 2000 (or newer) and the Macromedia Flash Player (7 or higher). [To delete this message, click inside the box, click the border of the box, and then press delete.] Copyright © Houghton Mifflin Company. All rights reserved.

Covalent Bonding Covalent bond –formed from the sharing of electrons Usually between 2 non metallic elements

Covalent bonds When 2 nonmetals combine. Both want the electrons available. The electrons in this type of bond are shared so that both elements can feel “full” These types of bonds produce molecules.

polarity The ability of an atom in a covalent bond to pull the shared pair of electrons more strongly to itself.

Molecules Molecule – two or more atoms covalently bound together Diatomic molecule – two of the same atom bound together H, N, O, F, Cl, Br, I H2, N2, O2, F2, Cl2, Br2, I2

Octet & Duet Rules Octet Rule – atoms want to have 8 valence electrons Duet Rule – H is the exception. It wants to be like He & is stable with only 2 valence electrons

Octet rule exceptions 3. B prefers 6 e’.

Ionic bonds When a metal and a nonmetal combine. E’ from the metal are transferred to the nonmetal. The goal is for both elements to achieve the noble gas configuration.

1. Predict if the following bonds are ionic or covalent Al-Si Li-S Ca-Cl Ba-O Ca-P F-S C-H B-Na Br-Rb

Comparison of a Molecular Compound and an Ionic Compound Click in this box to enter notes. Go to Slide Show View (press F5) to play the video or animation. (To exit, press Esc.) This media requires PowerPoint® 2000 (or newer) and the Macromedia Flash Player (7 or higher). [To delete this message, click inside the box, click the border of the box, and then press delete.] Copyright © Houghton Mifflin Company. All rights reserved.

Metallic bonds The electrons in this type of bond to not center around one atom but appear to float around each nucleus as a delocalized cloud. This causes metallic bonds to exhibit unique properties. 1. malleability 2.ductility

3. good conductor of heat and electricity. 4. hardness is determined by bond strength. 5. luster

Bond length & Strength As the number of bonds increases, the bond length decreases The shorter the bond, the stronger the bond Single> double>triple

Lewis diagrams & the octet rule Gilbert Newton Lewis (surrounded by pairs of electrons)

Lewis Structure Lewis Structures – shows how the valence electrons are arranged among the atoms of a molecule There are rules for Lewis Structures that are based on the formation of a stable compound Atoms want to achieve a noble gas configuration

Steps for drawing Lewis Structures Sketch a simple structure with a central atom and all attached atoms Add up all of the valence electrons for each individual atom If you are drawing a Lewis structure for a negative ion add that many electrons to create the charge If you are drawing a Lewis structure for a positive ion subtract that many electrons to create the charge

Steps for drawing Lewis Structures Subtract 2 electrons for each bond drawn Complete the octet on the central atom & subtract those electrons Complete the octet on the surrounding atoms & subtract those electrons Get your final number If 0  you are done! If +  add that many electrons to the central atom If -  need to form multiple bonds to take away that many electrons

Examples CCl4 Sketch a simple structure with a central atom and all attached atoms Cl │ Cl – C – Cl

Examples Add up all of the valence electrons for each individual atom 4 + 4(7) = 32 Subtract 2 electrons for each bond drawn 32-8 = 24 Complete the octet on the central atom & subtract those electrons Done

Examples Complete the octet on the surrounding atoms & subtract those electrons 24 – 24 = 0 Final number = 0…DONE! Final structure is… __ │Cl │ __ │ __ │Cl – C – Cl │ │

Draw the Lewis structures for the following molecules : HCl SCl2 PH3 SiF4

Polyatomic ion Lewis structures 1. determine the total number of valence electrons involved. 2. add the charge of the polyatomic to the total number. 3. arrange the elements as they would bond together. 4. place a pair of electrons everywhere a bond occurs.

Use the remaining electrons to fill the octet rule for the elements involved in the bond. * remember, electrons always travel in pairs.

When there are not enough e’ to satisfy everyone’s octet, double and triple bonds are formed!

Formation of C=C Double Bond in Ethylene Click in this box to enter notes. Go to Slide Show View (press F5) to play the video or animation. (To exit, press Esc.) This media requires PowerPoint® 2000 (or newer) and the Macromedia Flash Player (7 or higher). [To delete this message, click inside the box, click the border of the box, and then press delete.] Copyright © Houghton Mifflin Company. All rights reserved.

2. Draw the Lewis dot diagrams for the elements Z=7 thru Z=20 3. Draw the Lewis dot diagrams for each of the following polyatomic ions NH4+ OH- C2O42- O2-2

Molecular formulas- Used to tell exactly how many of each atom is present in the compound. Empirical Formulas- The simplest formulas possible indicating the ratio of atoms found in the molecule. Think of it as reducing a fraction to the common denominator.

Write the empirical formula for the following: C6H6 C2H2 C6H12O6 C4H10 P4O10 SO3 N2O4 NO2 Ag2C4H4O6 K2S4