Chapter : Chemical Bonding Cartoon courtesy of NearingZero.net

Chemical Bonds Lect1 Forces that hold groups of atoms together and make them function as a unit. Ionic bonds – transfer of electrons Metal + Nonmetal Ex) NaCl Li2O Covalent bonds – sharing of electrons. 2 nonmetals Ex) CH4 CO2 Metallic bonds- electrons are free to move throughout the material. Metals

Covalent Bonding Molecule- is the smallest unit quanitity of matter which can exist by itself and retain all the properties of the original substance. When a compound is formed by sharing electrons, the compound is called a ____Molecule__________________ Examples H2O & O2 Diatomic molecule- is a molecule containing 2 identical atoms. (H2 N2 O2 F2 Cl2 Br2 I2) H NO F

Covalent Bonding Diatomic molecule- is a molecule containing 2 identical atoms. (H2 N2 O2 F2 Cl2 Br2 I2) H NO F

Naming Covalent Compounds Two words, with prefixes Prefixes tell you how many. 1-mono, 6- hexa 2-di, 7- hepta (greek) septa (latin) 3-tri, 8-octa 4-tetra, 9-nona 5-penta, 10- deca

Naming Covalent Compounds First element whole name with the appropriate prefix, except mono Second element, -ide ending with appropriate prefix Practice

Naming Covalent Compounds CCl4 N2O4 XeF6 P2O9 H2O Carbon Dioxide Carbon Monoxide Carbon tetrachloride Dinitrogen tetroxide Xenon hexaflouride Diphosphorus Nonaoxide Dihydrogen Monoxide Water!

Covalent compounds The name tells you how to write the formula duh Sulfur dioxide SO2 diflourine monoxide F2O nitrogen trichloride NCl3 diphosphorus pentoxide P2O5

Empirical formula Shows the lowest, simplified ratio of elements in a compound Molecule Molecular formula Empirical formula 1 C2H4 CH2 2 C4H8 3 C3H8 ?

Chemical Formula- represents the relative # of atoms of each kind in a chemical compound by using atomic symbols and numerical subscripts. Example: H2O H=2 O=1 Molecular compound (Covalent Compounds) - simplest formula unit are molecules. Have low melting & boiling pts. Molecular formula- shows the types and numbers of atoms combined in a single molecule.

Bond Length- is the average distance between 2 bonded atoms. Bond Energy- is the energy required to break a bond. It gives us information about the strength of a bonding interaction.

I. Lewis Diagrams (p. 202 – 213) Lecture 2

The Octet Rule Chemical compounds tend to form so that each atom, by gaining, losing, or sharing electrons, has an (8) octet of electrons in its valence shell. 8 is Great! H He

Lewis Dot Shows how valence electrons are arranged among atoms in a molecule. Reflects central idea that stability of a compound relates to noble gas electron configuration. CH4 H2O I started it for you Try one yourself NH3

A. Octet Rule Remember… Most atoms form bonds in order to have 8 valence electrons.

B. Drawing Lewis Diagrams Find total # of valence e-. Arrange atoms - singular atom is usually in the middle. Form bonds between atoms (2 e-). Distribute remaining e- to give each atom an octet (recall exceptions). If there aren’t enough e- to go around, form double or triple bonds.

B. Drawing Lewis Diagrams CF4 1 C × 4e- = 4e- 4 F × 7e- = 28e- 32e- F F C F - 8e- 24e-

B. Drawing Lewis Diagrams CO2 1 C × 4e- = 4e- 2 O × 6e- = 12e- 16e- O C O - 4e- 12e-

B. Drawing Lewis Diagrams BeCl2 1 Be × 2e- = 2e- 2 Cl × 7e- = 14e- 16e- Cl Be Cl - 4e- 12e-

H O H N O F B F F F F F S F A. Octet Rule Very unstable!! Exceptions: Hydrogen  2 valence e- Groups 1,2,3 get 2,4,6 valence e- Expanded octet  more than 8 valence e- (e.g. S, P, Xe) Radicals  odd # of valence e-

C. Polyatomic Ions To find total # of valence e-: Add 1e- for each negative charge. Subtract 1e- for each positive charge. Place brackets around the ion and label the charge.

O O Cl O C. Polyatomic Ions ClO4- 1 Cl × 7e- = 7e- 4 O × 6e- = 24e-

H H N H C. Polyatomic Ions NH4+ 1 N × 5e- = 5e- 4 H × 1e- = 4e- 9e-

Electornegativity Lecture 3

Is it Covalent or ionic Nonpolar-Covalent bonds (H2) Electrons are equally shared Electronegativity difference of 0 to 0.5 Polar-Covalent bonds (HCl) Electrons are unequally shared Electronegativity difference between .5 and 2.1 3. Ionic Bonds 2.1- 3.3

Using Electronegativity differences CO 3.5-2.5= 1.0 polar covalent C=2.5 look on table pg 198 O= 3.5 2.1 thru 3.3= Ionic Bond

Electronegativity Nonpolar Covalent Polar covalent Ionic 0-------------0.5 --------------------------------------- >2.1 ---------3.3

practice NaCl Na=0.9 Cl=3.2 3.2-0.9=2.3 NaCl is ionic CO 3.4- 2.6= 0.8 Polar covalent CH4 2.6 -2.2= 0.4 Non polar covalent

D. Resonance Structures Molecules that can’t be correctly represented by a single Lewis diagram. Actual structure is an average of all the possibilities. Show possible structures separated by a double-headed arrow.

D. Resonance Structures O S O O O S O O O S O

D. Resonance-Occurs when more than one valid Lewis structure can be written for a particular molecule.

Structural Formula Shows shared pair of electrons by a dashed line.

Single bond- 1 pair of electrons Double bond- 2 pair of electrons Triple bond- 3 pair of electrons Try a couple: O2 N2

Ionic Bonding and Compounds Ionic Compound- is composed of positive and negative ions combined so that the positive & negative charges are equal. (Metal + nonmetal) Formula Unit- is the simplest collection of atoms from which a compounds formula can be established.

Writing Formulas Write the symbols of each element Put their charge in their upper right corner Crisscross the numbers down (Not the charges). Example: Write the formula for Magnesium Chloride Mg Cl Mg+2 Cl-1 MgCl2 MgCl2

Writing Formulas Practice +1 +2 +3 4 -3 -2 -1 Write the formula for: Aluminum Bromide Calcium Oxide Calcium Nitride Sodium Chloride

Lattice Energy- is the energy released when 1 mole of an ionic crystalline compound is formed from gaseous ions. Ionic Compounds have high melting points & boiling points, are hard and brittle, have crystalline structure.

Polyatomic Ions Many atoms with a charge. Example SO4-2

Metallic Bonding Metals- conduct heat, have low ionization energy & electronegativity, give up e- Metallic Bond- is a chemical bond resulting from the attraction between positive ions and surrounding mobile e-. Malleability and ductility

Molecular Geometry VSEPR Theory- “Valence- shell, electron-pair repulsion” states that repulsion between the sets of valence-level electrons surrounding an atom cause these sets to be oriented as far apart as possible.

Determining VSEPR H-O-H E2 B B A Determine the VSEPR for H2O Draw the Lewis Dot Draw the Structural Formula Label the central atom as A Label any atoms attached to the center atom as B Label any paired electrons on the central atom that are not used in the bond as E H-O-H E2 B B A VSEPR AB2E2 Shape Bent (look on chart)

VSEPR Chart VSEPR SHAPE AB or AB2 Linear AB2E Bent AB3 Trigonal-Planar Tetrahedral AB3E Trigonal-Pyramidal AB2E2 AB5 Trigonal-Bipyramidal AB6 Octahedral

Hybridization-The Blending of Orbitals. Dipole- is created by equal but opposite charges that are separated by a short distance. Dipole-Dipole Attractions-Attraction between oppositely charged regions of neighboring molecules. Hydrogen Bonding- Bonding between hydrogen and more electronegative neighboring atoms such as oxygen and nitrogen. Hydrogen bonding in Kevlar, a strong polymer used in bullet-proof vests. London Dispersion Forces- The temporary separations of charge that lead to the London force attractions are what attract one nonpolar molecule to its neighbors. London forces increase with the size of the molecules.

“Electronegativity chart”. Table. Aug. 9, 2006. http://www “Lewis Structures”. Drawings. Aug. 9, 2006. http://www.avon-chemistry.com/chem_bond_explain.html “Oscar”. Photo. Aug. 9, 2006. http://www.musicmerchant.com/22061.htm “Water Structural Formula”. Drawing. Aug. 10, 2006. http://www.accs.net/users/kriel/ch4notes/water_structural_formula.gif “Periodic Table of Elements”. Chart. Aug. 9, 2006. http://users.erols.com/kdennis/periodictable.jpg “Information”. Aug 11, 2006. http://www.sciencegeek.net/Chemistry/Powerpoint/Unit3/Unit3_files/frame.htm Holt, Rinehart and Winston. Modern Chemistry. Harcourt Brace & Company. 1999.