Naming Molecules Ch. 9, Section 2: pg. 248

Naming Binary Molecular Compounds 1.The first element is always named first, using the entire element name. 2.The second element is named using the root of the element and adding “-ide”. (Just like ionic cpds.) 3. Use prefixes on the names to indicate the number of atoms of each type that are present in the compound.

Common Prefixes in Covalent Compounds # of atomsPrefix# of atomsPrefix 1mono-6hexa- 2di-7hepta- 3tri-8octa- 4tetra-9nona- 5penta-10deca-

What is the name of P 2 O 5? Step 1: phosphorous Step 2: oxide Step 3: diphosphorous pentoxide EASY!!!!!!

What are the names of the following binary molecular compounds? 1.CCl 4 2.As 2 O 3 3.CO 4.SO 2 5.NF 3 Answers: 1. carbon tetrachloride 2. diarsenic trioxide 3. carbon monoxide 4. sulfur dioxide 5. nitrogen trifluoride

Naming Acids There are two common types of acids: *binary acids *oxyacids

Binary acids contain HYDROGEN and one other element. Steps to naming binary acids: 1.Use the prefix hydro- to name the hydrogen part of the compound. 2.Use a form of the root of the second element plus the suffix –ic. 3.End with the word acid. Example: HCl hydrochloric acid

Name these binary acids. 1.HI 2.HF 3.H 2 S 4.H 2 Se 5.HBr Answers 1.hydroiodic acid 2.hydrofluoric acid 3.hydrosulfuric acid 4.hydroselenic acid 5.hydrobromic acid

Oxyacids are acids that contain hydrogen and an oxyanion. **Remember: An oxyanion is a polyatomic ion that contains oxygen. EX: PO 4 3- SO 4 2-

Steps to naming oxyacids: 1.Identify the oxyanion present in the acid. (You can use your ion chart.) 2.Use a form of the anion and **if the oxyanion ends in “-ate”, change to “-ic”. **if the oxyanion ends in “-ite”, change to “-ous”. 3.Add the word “acid” at the end. (We do NOT use “hydro” with oxyacids.)

“-ate” “-ic” “-ite” “-ous” Example: HClO 3 Step 1: ClO 3 is the chlorate anion. Step 2: Change chlorate to chloric. Step 3: Add the word “acid”. ANSWER: chloric acid

Name the following oxyacids 1.HClO 2 2.H 2 SO 4 3.H 2 SO 3 4.H 3 PO 4 5.HNO 3 Answers 1.chlorous acid 2.sulfuric acid 3.sulfurous acid 4.phosphoric acid 5.nitric acid

ASSIGNMENT **Page 250: #18 – 22 **Page 251: #23 – 29

Lewis Structures Lewis structure: when electron-dot diagrams are used to show how electrons are arranged in molecules EX: We use a line to show bonding pairs of electrons in a Lewis structure and a pair of dots to show electrons that are NOT being shared (called “lone pairs”).

Examples of Lewis Structures Methane Ammonia Ammonium ion Ammonia ion

Drawing Lewis Structures 1.Count ALL the valence electrons for the molecule. EX: CCl 4 C = 4 e -, Cl = 4x7 = 28 e - Total valence e - = 32 e - 2. Determine the central atom. *H is NEVER a central atom. *The halides are NEVER a central atom. *The element with the lowest electronegativity is the central atom. *The only single element would be the central atom EX: C is the only single element, so it would be the central atom in the Lewis structure.

3.Place two electrons in each bond by drawing a line to represent the bond. Cl l Cl – C – Cl l Cl 4.Complete the octet of the atoms attached to the central atom by adding electrons in pairs. (See board for example.)

5.Place any remaining electrons on the central atom in pairs. *****REMEMBER! The total electrons in the Lewis structure MUST equal the number of electrons in step #1.****** (See board for example.) 6.If the central atom does not have an octet after step #5, form double or triple bonds, as needed, between the central atom and one or more of the terminal atoms. (See board for example.)

Draw the Lewis Structures for the following molecules 1.PH 3 2.H 2 S 3.HCl 4.CCl 4 5.SiH 4

Multiple Bonds  Many molecules attain a noble-gas configuration by sharing more than one pair of electrons between two atoms, forming a multiple covalent bond.  C, N, O, and S most often form multiple bonds.

Lewis Structures Which of these is a **single bond? **triple bond? **double bond? O = C = O

Multiple Bonds: Draw the Structural Formulas for the following: 1.O 2 2.N 2 3.CO 2

Lewis Structures for Polyatomic Ions

1.Determine the number of valence electrons in the atoms present EX: PO 4 1 P = 5 electrons = 5 4 O = 6x4 electrons = electrons 2.Draw the Lewis structure for the ion. 3.Count the total electrons in the Lewis structure.

4.For negative ions: Subtract ions in the Lewis structure from the valence electrons found in step 1 and this is the charge. For positive ions: Subtract the valence electrons in step 1 from the total ions in the Lewis structure See next slide for examples……………………

1.ClO NH 4 +1

Resonance Structures  If a molecule or polyatomic ion has BOTH a double bond AND a single bond, it is possible to have more than one correct Lewis structure.  RESONANCE is a condition that occurs when more than one valid Lewis structure can be written for a molecule or ion.  Resonance structures differ only in the position of the electron pairs.

Examples of Resonance Structures

1.BCl 3 2.SO 2 3.SO 3 See board for answers.

Exceptions to the Octet Rule 1.Molecules with an odd # of valence electrons EX: NO 2 ClO 2 NO 2. Compounds with fewer than 8 valence electrons present (this is very rare) EX: BH 3

3. Compounds in which the central atom has more than 8 valence electrons -- called an EXPANDED OCTET ; occurs in energy levels of elements in period 3 and up **Extra lone pairs are added to the central atom OR more than 4 bonding atoms are present. EX: IF 4 -1

VSEPR Model : Valence Shell Electron Pair Repulsion model  Based on an arrangement that minimizes the repulsion of shared and unshared pairs of electrons around the central atom  READ pages 259 – 261.  Look at Table 9-3 as you read.  Complete #49 – 53 on pg. 262.

Electronegativity and Polarity  Electronegativity indicates the relative ability of an atom to attract electrons in a chemical bond.  It generally increases as the atomic number increases ACROSS A PERIOD and generally decreases as you go DOWN A GROUP.

Polar or Nonpolar???????????  Identical atoms, like N 2, have an electronegativity difference of zero, and the electrons in the bond are equally shared between the two atoms. This is a NONPOLAR COVALENT BOND, or a PURE COVALENT BOND.  A covalent bond between atoms of different elements does not have equal sharing of the electron pair, due to the difference in electronegativity.

 Unequal sharing results in a POLAR COVALENT BOND. The shared electrons are pulled toward one of the atoms and spend more time around that atom than the other atom. Partial charges occur at the ends of the bond. This bond is often referred to as a dipole (two poles).  To determine if the bond is polar or nonpolar, you must look at the shape of the molecule. Draw the molecular structure, using what you know from Table 9-3.  SYMMETRIC MOLECULES ARE USUALLY NONPOLAR AND ASYMMETRIC ARE POLAR AS LONG AS THE BOND TYPE IS POLAR.

Do #60 – 63 on page 266.

CLASSWORK: Pg. 273 Complete #94, 95, 96, 97, 98, 99: a,c,d, 100: a,d,e