Bonding Unit III

I. Bond N A. What is a bond? Attraction of an electron by two nuclei B. What electrons are involved in bonding? Valence electrons Electrons in the outermost energy level Represented by an electron-dot diagram example N 2- 5 5 valence electrons N

C. Why do atoms bond? Atoms will form bonds to achieve 8 valence electrons Octet Rule “8 is GREAT” Stable configuration Noble gas configuration (Group 18) Compound has lower potential energy than the un-bonded atoms

D. Energy Terms Ionization Energy Electronegativity Amount of energy needed to remove an electron Compare Na Cl Ar Ionization Energy 496 kJ 1251 kJ 1521kJ Why? Electronegativity Attraction for electrons Compare Na Cl Ar Electronegativity 0.9 3.2 0

E. Types of Bonds Atomic Bonds Molecular Bonds Holds atoms together to make a molecule Molecular Bonds Holds molecules together to make a solid or liquid

II. Atomic Bonds Ionic Bond Transfer of electrons from metals to nonmetals Metals lose electrons Become positive Decreases in size Nonmetals gains electrons Become negative Increases in size Properties High melting points Hard Poor conductor as solid Good conductor in liquid or aqueous states Nonmetals Metals

Bond Diagram Write electron dot diagram for each element Transfer electrons from metal to nonmetal until all single electrons are paired up

Sharing of electrons between two nonmetals Types Nonpolar Covalent Covalent Bond Sharing of electrons between two nonmetals Types Nonpolar Covalent equal sharing of electrons Atoms have the same electronegativity values diatomic elements (H2 O2 F2 Br2 I2 N2 Cl2) Polar Covalent unequal sharing of electrons Atoms have the different electronegativity values higher electronegative element is the negative end of the molecule

Properties Low melting points Soft Poor conductor Bond Diagram Write electron dot diagram for each element Share electrons of any all single electrons Each “loop” is a bond

example of a nonpolar covalent bonds: Diatomic molecules H O F Br I N Cl

example of a polar covalent bond:

Network Solid Giant covalent bonds sharing of electrons occurs in the entire crystal examples C (diamond) C (graphite), SiO2 (silicon dioxide) SiC (silicon carbide) asbestos Properties Hardest solids Highest melting points Poor conductor

Metallic Bonding Metals do not have enough valence electrons to stabilize the structure Delocalize electrons Allow them to move anywhere in the structure Properties Hard High melting points Good conductor in the solid state

III. Nomenclature Naming Compounds 1. Binary Compounds Compounds that contain two elements Name the first element Name the second element using an “ide” ending Check the oxidation number of the first element If it has more than one positive value, use a Roman numeral to indicate the number that was used

CaCl2 Calcium chloride Ca is +2 only No Roman numeral needed

CCl4 Carbon chloride C is -4, +2 and +4 Since there are two positive choices, use a Roman numeral C Cl4 Carbon IV chloride +4 -4 =0 +4 -1

B. Ternary Compounds Compounds that contain more than two elements Polyatomic ion is present Found on Table E Identify the polyatomic used Name the first substance Name the second substance Check the oxidation number of the first element If it has more than one positive value, use a Roman numeral to indicate the number that was used

K2SO4 Potassium sulfate Potassium is +1 only No Roman numeral is needed

Cu(NO3) 2 Copper nitrate Copper can be +1 or +2 Copper II nitrate Needs a Roman numeral Cu (NO3)2 +2 -2 =0 +2 -1 Copper II nitrate

FeSO4 Iron sulfate Iron can be +2 or +3 Iron II sulfate Needs a Roman numeral Fe SO4 +2 -2 =0 +2 -2 Iron II sulfate

NH4Cl Ammonium Ammonium chloride Ammonium ion is +1 No choice means no Roman numeral needed

B. Writing Formulas Use ending to tell if compound is binary or ternary Binary usually ends in “ide” [except hydroxide and cyanide] Ternary usually ends in “ate” or “ite” Write the symbols Assign oxidation numbers Positive atom on left, negative on right Reduce if possible and “criss-cross” These numbers become the subscripts of the formula

Strontium phosphide “ide” means binary Strontium (Sr) Phosphide is phosphorus (P) Sr+2 P-3 Sr3P2

Iron II Oxide “ide” means binary Iron (Fe) Oxide is oxygen (O) reduce FeO

Manganese IV Carbonate “ate” means ternary Manganese (Mn) Carbonate is a polyatomic ion (CO3-2) Mn+4 CO3-2 Mn2(CO3 )4 reduce Mn (CO3 )2

IV Intermolecular Bonds Definition Holds completed molecules together to form a liquid or solid Strength is reflected by melting point or boiling point Types Ionic bonds, network solids and metallic bonds can be considered both atomic or intermolecular Covalent bonds will join by Hydrogen bonding Dipole bonding Van der Waals forces

Hydrogen Bonding Dipole Van der Waals forces Attraction between small highly electronegative elements H to F, O or N Dipole Attraction between asymmetrical molecules Polar molecule Greater polarity, stronger bond (electronegativity difference) Van der Waals forces Attraction between symmetrical molecules Nonpolar molecule Larger the molecule, stronger the bond “bigger is better”

Asymmetrical molecule Nonpolar molecule Asymmetrical molecule Polar molecule Dipole

H2O vs H2S vs H2Se vs H2Te Highest Boiling Point H2O vs H2S vs H2Se vs H2Te Highest Boiling Point? Strongest bond H2O Why? Hydrogen bonded Most polar

F2 vs Cl2 vs Br2 vs I2 Smallest F2 and Cl2 Br2 Largest I2 Gases Liquid Solid