I. Introduction to Bonding (p. 161 – 163) Ch. 6 & 7 - Chemical Bonding I. Introduction to Bonding (p. 161 – 163)

A. Types of Bonds IONIC COVALENT Bond Formation e- are transferred from metal to nonmetal e- are shared between two nonmetals Type of Structure crystal lattice true molecules Physical State solid liquid or gas Melting Point high low Solubility in Water yes usually not Electrical Conductivity yes (solution or liquid) no Other Properties

A. Types of Bonds METALLIC e- are delocalized among metal atoms Bond Formation e- are delocalized among metal atoms Type of Structure “electron sea” Physical State solid Melting Point very high Solubility in Water no yes (any form) Electrical Conductivity Other Properties malleable, ductile, lustrous

A. Types of Bonds RETURN

A. Types of Bonds RETURN

Ionic Bonding - Crystal Lattice A. Types of Bonds Ionic Bonding - Crystal Lattice RETURN

Covalent Bonding - True Molecules A. Types of Bonds Covalent Bonding - True Molecules Diatomic Molecule RETURN

Metallic Bonding - “Electron Sea” A. Types of Bonds Metallic Bonding - “Electron Sea” RETURN

B. Vocabulary Chemical Bond electrical attraction between nuclei and valence e- of neighboring atoms that binds the atoms together bonds form in order to… decrease PE increase stability

NaCl CO2 B. Vocabulary CHEMICAL FORMULA IONIC COVALENT Formula Unit Molecular Formula NaCl CO2

NaCl NaNO3 B. Vocabulary COMPOUND more than 2 elements 2 elements Binary Compound Ternary Compound NaCl NaNO3

Na+ NO3- B. Vocabulary ION 1 atom 2 or more atoms Monatomic Ion Polyatomic Ion Na+ NO3-

C. Bond Polarity Most bonds are a blend of ionic and covalent characteristics.

C. Bond Polarity Nonpolar Covalent Bond e- are shared equally symmetrical e- density usually identical atoms

+ - C. Bond Polarity Polar Covalent Bond e- are shared unequally asymmetrical e- density results in partial charges (dipole) + -

C. Bond Polarity Nonpolar Polar Ionic View Bonding Animations.

C. Bond Polarity Electronegativity Attraction an atom has for a shared pair of electrons. higher e-neg atom  - lower e-neg atom +

C. Bond Polarity Electronegativity Trend (p. 151) Increases up and to the right.

C. Bond Polarity Difference in the elements’ e-negs determines bond type

II. Molecular Compounds (p. 164 – 172, 211 – 213) Ch. 6 & 7 - Chemical Bonding II. Molecular Compounds (p. 164 – 172, 211 – 213)

A. Energy of Bond Formation Potential Energy based on position of an object low PE = high stability

A. Energy of Bond Formation Potential Energy Diagram attraction vs. repulsion no interaction increased attraction

A. Energy of Bond Formation Potential Energy Diagram attraction vs. repulsion increased repulsion balanced attraction & repulsion

A. Energy of Bond Formation Bond Energy Energy required to break a bond Bond Energy Bond Length

A. Energy of Bond Formation Bond Energy Short bond = high bond energy

X O B. Lewis Structures 2s 2p Electron Dot Diagrams show valence e- as dots distribute dots like arrows in an orbital diagram 4 sides = 1 s-orbital, 3 p-orbitals EX: oxygen X 2s 2p O

Ne B. Lewis Structures Octet Rule Most atoms form bonds in order to obtain 8 valence e- Full energy level stability ~ Noble Gases Ne

B. Lewis Structures - + Nonpolar Covalent - no charges Polar Covalent - partial charges + -

C. Molecular Nomenclature Prefix System (binary compounds) 1. Less e-neg atom comes first. 2. Add prefixes to indicate # of atoms. Omit mono- prefix on first element. 3. Change the ending of the second element to -ide.

C. Molecular Nomenclature PREFIX mono- di- tri- tetra- penta- hexa- hepta- octa- nona- deca- NUMBER 1 2 3 4 5 6 7 8 9 10

C. Molecular Nomenclature CCl4 N2O SF6 carbon tetrachloride dinitrogen monoxide sulfur hexafluoride

C. Molecular Nomenclature arsenic trichloride dinitrogen pentoxide tetraphosphorus decoxide AsCl3 N2O5 P4O10

C. Molecular Nomenclature The Seven Diatomic Elements Br2 I2 N2 Cl2 H2 O2 F2 H N O F Cl Br I

III. Ionic Compounds (p. 176 – 180, 203 – 211) Ch. 6 & 7 - Chemical Bonding III. Ionic Compounds (p. 176 – 180, 203 – 211)

A. Energy of Bond Formation Lattice Energy Energy released when one mole of an ionic crystalline compound is formed from gaseous ions

B. Lewis Structures Covalent – show sharing of e- Ionic – show transfer of e-

B. Lewis Structures Covalent – show sharing of e- Ionic – show transfer of e-

C. Ionic Nomenclature Ionic Formulas Write each ion, cation first. Don’t show charges in the final formula. Overall charge must equal zero. If charges cancel, just write symbols. If not, use subscripts to balance charges. Use parentheses to show more than one polyatomic ion. Stock System - Roman numerals indicate the ion’s charge.

C. Ionic Nomenclature Ionic Names Write the names of both ions, cation first. Change ending of monatomic ions to -ide. Polyatomic ions have special names. Stock System - Use Roman numerals to show the ion’s charge if more than one is possible. Overall charge must equal zero.

C. Ionic Nomenclature Consider the following: Does it contain a polyatomic ion? -ide, 2 elements  no -ate, -ite, 3+ elements  yes Does it contain a Roman numeral? Check the table for metals not in Groups 1 or 2. No prefixes!

C. Ionic Nomenclature Common Ion Charges 1+ 2+ 3+ NA 3- 2- 1-

C. Ionic Nomenclature potassium chloride magnesium nitrate copper(II) chloride K+ Cl-  KCl Mg2+ NO3-  Mg(NO3)2 Cu2+ Cl-  CuCl2

C. Ionic Nomenclature NaBr Na2CO3 sodium bromide FeCl3 sodium carbonate iron(III) chloride