Chemical bond: Two atoms combine to form a new substance.

1. Chemical Bonding Unit Chemical bond: Two atoms combine to form a new substance. The bond is created by an electromagnetic force produced by an exchange or sharing of electrons. (See video)

2. Bond Strength is measured by the amount of energy needed to break the connection between the atoms. - Bond Strength varies depending on the atoms involved. - Shorter bonds tend to be stronger. Bond Length This is measured by the distance between the nuclei of the two atoms. Electronegativity The ability of an atom to attract electrons. (Trend) To determine whether a bond is Covalent or Ionic, you can compare their EN values.. - A large difference in EN between atoms = Ionic Bond - Similar EN values in atoms = Covalent Bond ** Table is on page 303 in the text book…

3. He -- B 2.01 C 2.50 N 3.07 O 3.50 F 4.10 Ne -- Al 1.47 Si 1.74 P 2.06 S 2.44 Cl 2.83 Ar -- Zn 1.66 Ga 1.82 Ge 2.02 As 2.20 Se 2.48 Br 2.74 Kr -- Cd 1.46 In 1.49 Sn 1.72 Sb 1.82 Te 2.01 I 2.21 Xe -- Hg 1.44 Tl 1.44 Pb 1.55 Bi 1.67 Po 1.76 At 1.96 Rn -- Electronegativity Chart

4. Stable Octet Rule Atoms tend to gain or lose valence electrons in order to have a filled valence shell. (8 in the s,p sublevels) Ionic Bonds: A metal atom transfers electron(s) to a nonmetal atom. The resulting Cation and Anion bond due to their attraction.

5. Ionic Compounds: Tend to have high Melting Points, conduct electricity, typically soluble in water, and crystallize as sharply defined particles. Covalent Bonds: Two non-metal atoms share electrons in order to fulfill a stable octet. (Valence Shells Overlap) ** Hydrogen requires only 2 electrons (not 8) Covalent Compounds tend to have low Melting Points, are Brittle, and do NOT conduct electricity. Polyatomic Ions are groups of atoms covalently bonded that have an overall charge. Ex: PO4-3

7. LET’S CREATE SOME BONDS WITH LEWIS STRUCTURES!! • Basic rules for ionic and covalent bonding: • Lewis structures • I =Must empty metal’s valence and fill non • C =Single, double, triple bonds (even in shared) • Must have Central Atom

8. Two types of Covalent Bonds: 1. Nonpolar Covalent Bonds occur when the atoms share electrons equally. (*Diatomics*) 2. Polar Covalent Bonds occur when electrons are not being shared equally between the atoms due to a slight difference in Electronegativity. Metallic Bonds(Sea of Electrons) Metals in the solid state form crystals that allow their atoms’ outer level orbitals to overlap… Not a chemical bond! - The electrons move freely throughout the crystal arrangement. This movement between orbitals holds the metallic atoms together. Alloys are homogeneous mixtures between metals that strengthen weaker metals. Ex: ??

9. Determination of Bond Types Subtract the EN values of the two atoms in the bond. If the absolute value of the difference is greater than 1.67, it will form an ionic bond. Less than 1.68 = Covalent. - Between 1.67 and 0.31 is Polar Covalent whereas 0.3 or less is Nonpolar Covalent. Practice - Merrill Chem Textbook: Pg.304 #’s 1,2. Pg.306 #s 3 (a-f) ,4(all). Answers: 304: 1) In, Sb, Se, F 2) Fr, Zn, Ga, Ge, P 306: 3a) 0.27 = Covalent 3b) 2.53 = Ionic 3c) 0.30 = C 3d) 1.47 = C 3e) 1.02 = C 3f) 1 = C 4a) 0.01 = C 4b) 0.49 = C 4c) 2.49 = I 4d) 0.82 = C 4e) 1.66 = C 4f) 3.06 = I

10. e- N N e- e- e- Molecular Structures Covalent Molecule Shapes Important Terms: Bond Axis – The line drawn between the nuclei of the two atoms. Bond Angle – The angle formed between TWO bond axes around a central atom. Orbital Overlap – The shapes of the e- Orbitals of atoms in a bond are altered from their original form, due to both nuclei attracting the electrons.

11. VSEPR Theory Valence Shell Electron Pair Repulsion Vespur* theory helps us to determine the shape of molecules that have a central atom based on the repulsion of e- pairs. • Whether e- are shared in a bond or not, they will repel each other. • Therefore the VSEPR theory states that e- pairs around a central atom will spread out as far as possible to minimize repulsions. 1. Draw the lewis structures for each atom in the molecule. 2. Count the total # of Valence e- present. 3. Predict how the bonds will form to make all atoms stable. 4. Count the number of bonds, and the number of unbonded pairs around the central atom. 5. Determine the shape and draw the structure. **Make sure to label the shape type, all bond angles, and partial charges.

12. Exceptions • Some elements with an empty d-sublevel can expand their octet. This means that they can be stable with MORE than 8 valence e- . These elements will most often follow the octet rule, but can expand depending on the molecule. • Group 15 Elements: Can form up to 5 bonds. (8 or 10 e-) • Groups 16-18: Can form up to 6 bonds. (8,10, or 12 e-) • Unbonded Pairs are held closer to the central atom, which results in pushing the Bonded Pairs of e- closer to each other slightly. This changes the bond angles. • Multiple Bonds (Double or Triple Bonds) act as just one area of repulsion. • Sometimes they can be ions, and you will need to add or remove electrons to make a pair.

13. Dipoles A dipole is a molecule with two areas of charge due to a polar covalent bond. (partial + and partial - ) Ex: Water Resonance Resonance is a phenomenon which occurs within molecules with multiple bonds between the same atoms. The extra bond will constantly switch its location between the atoms. Ex: Fan effect!