Bonding Chapters 7-8

Octet Rule Atoms tend to lose or gain electrons to achieve a full valence shell (8) Exception: First Energy Level is full with 2 electrons

Electron Dot Structures Diagrams that show valence electrons, usually as dots AKA Lewis Electron Dot Diagrams Rules Start on any side First two get paired together Next three are separated Fill in as needed O

Examples H He F Ne Ar N Na Cl

Ions Atoms that have gained or lost electrons, and now have a charge Must show charge Na+ F- O-2

Practice Draw Lewis Electron Dot Structures for the following atoms and ions: Li, B, Mg, Al, P, S, Cl, Br, Kr H+, Li+, Mg2+, Al3+, O2-, F-, Cl-, S2-, P3-

Practice Al Li B Mg P S Kr Cl Br

Practice H+ Li+ Mg2+ Al3+ P-3 O-2 Cl- F- S-2

Compounds Two Main Types of Compounds Ionic Molecular (Covalent) Based on type of bonding involved

Bonding Bond Three Main Types Shared or exchanged electrons that hold two atoms together Three Main Types Covalent Ionic Metallic

Covalent Bonds Electrons are shared between two atoms to hold them together Each atom will try to achieve a full valence shell 2 nonmetals Two types of covalent bonds Non-Polar Covalent – Shared equally Polar Covalent – Shared unequally

Covalent Bonding H2 H H H Single Bond

Covalent Bonding H2O Bond O H H Bond H H

More Examples O2 O O O Double Bond

More Examples N2 N N N Triple Bond

More Examples Cl H HCl NH3 N H

More Examples CH4 CO2 C H C C O

Determining Bond Type Whether electrons are shared or exchanged is based on electronegativity difference between two bonding atoms Nonpolar Covalent Bond 2 same Nonmetals (no difference in electronegativity) Polar Covalent Bond 2 different Nonmetals (small difference in electronegativity)

Determining Bond Polarity The larger the difference in electronegativity, the more polar the bond.

Determining Bond Polarity Which is more polar? ΔEN H F Most Polar Biggest 1.8 H Cl 1.0 H Br 0.8 H I 0.5

Bonding Ionic Bond Electrons are transferred from one atom to another (one gives, one takes) Metal and nonmetal, NaCl Large electronegativity difference

Properties Ionic Compounds Most ionic compounds are hard, crystalline solids at room temperature High melting points Mostly soluble in water Can conduct an electric current when melted or dissolved in water(aq).

Properties Covalent Compounds Most molecular compounds tend to have relatively lower melting and boiling points than ionic compounds.

Ionic Compounds Formula Unit is the lowest whole-number ratio of ions in an ionic compound Ionic Compounds are repeating lattices of positive and negative ions

Ionic Compounds NaCl

Ionic Compounds Ionic compounds are electrically neutral, even though they are composed of charged ions Total positive charge equals total negative charge

Determining Formulas Must be electrically neutral Total positive charge must equal total negative charge Use oxidation numbers from Periodic Table Group 1  +1 Group 2  +2 Group 13  +3 Group 15  -3 Group 16  -2 Group 17  -1

Determining Formulas Determine number of each ion to balance out charge Use as subscript for element symbol Ex: CaCl2, Na3PO4, Mg(NO3)2 Write Positive Ion First Formula must be smallest whole-number ratio

Example Sodium and Chlorine Na+ Cl- Na1Cl1 NaCl

Example Calcium and Fluorine Ca+2 F- F- Ca1F2 CaF2

Examples Potassium and Oxygen K+ K+ O-2 K2O1 K2O

Polyatomic Ions Group of atoms that collectively have gained or lost electrons (Table E) Sodium and Nitrate Na+ (NO3)- Na1(NO3)1 NaNO3

More Examples Potassium and Sulfate Ammonium and Sulfur K2SO4 (NH4)2S

Short-cut (criss-cross method) Magnesium and Phosphate Mg+2 PO4-3 Mg3(PO4)2

Short-cut (criss-cross method) Magnesium and Carbonate Mg+2 CO3-2 Mg2(CO3)2 Must Simplify MgCO3

Na2SO4 Ca3(PO4)2 Do This Now Write out formulas for: Sodium Sulfate Calcium Phosphate Na2SO4 Ca3(PO4)2

[ ] Cl Na+ Cl- - Dot Structures Shows valence electrons Must show charge for Ions NaCl Cl [ ] - Na+ Cl-

Dot Structures MgO Mg+2 O-2

Dot Structures CaF2 Ca+2 F-

Polyatomics Compounds with polyatomic ions contain BOTH ionic and covalent bonds Example: NaNO3 - N O Na+

Network Solids All atoms in a network solid are covalently bonded together Network solids have very high melting and boiling points, since melting requires the breaking of many bonds throughout the compound. Some of the strongest materials known to man are network solids.

Network Solids Diamonds ( C ) Graphite ( C ) Silicon Dioxide (SiO2) Silicon Carbide (SiC)

Metallic Bonding Bonding within metallic samples is due to highly mobile valence electrons Free flowing valence electrons “Sea of Electrons”

Metallic Substances High melting points Conducts as a solid

Bond Energy When two atoms form a bond, energy is released Example: Cl + Cl  Cl2 + energy Energy needs to be added to break a bond Example: Cl2 + energy  Cl + Cl

O H N H O H Cl Structural Formulas Shared electrons are written as a line, unshared electrons are not written Each line represents 2 electrons O H N H O H Cl

Molecular Polarity Polar Molecule Nonpolar Molecule one end of a molecule is slightly negative(δ-) and the other end is slightly positive(δ+). Asymmetrical charge distribution Nonpolar Molecule Can not be separated into different ends Symmetrical charge distribution

O H Polar Molecule H2O Polar Covalent Bond Electrons shared Unequally δ- O H δ+

More Examples HCl NH3 H Cl δ- δ+ δ- N H N H δ+

Another Example CH4 Nonpolar Molecule δ+ C H δ- δ+ δ+ δ+

Polarity Ionic Compounds are Ionic Nonpolar Covalent Bonds always indicate Nonpolar Molecules Polar Covalent Bonds Determine Symmetry

“Like Dissolves Like” Polar and Ionic substances will dissolve in other Polar Substances Nonpolar substance will dissolve in other nonpolar substances

Intermolecular Forces Intermolecular Forces of Attraction attraction between two molecules or ions that hold them together (not a bond) Determines melting and boiling points of compounds Stronger intermolecular forces, higher melting and boiling points

Intermolecular Forces Van der Waals Dispersion Dipole-Dipole Molecule-Ion Hydrogen Bonding Weakest Strongest

Van der Waals Dispersion Electrons of one atom are attracted to the Protons of the next atom. Also called an induced dipole Attraction increases with increasing mass e e p p p p e e

H Cl H Cl Van der Waals Dipole-Dipole negatively charged end of polar molecule is attracted to positively charged end of another polar molecule δ- δ+ δ+ δ- H Cl H Cl

Molecule Ion Attraction between polar molecules and ions in solution O H O H Na+ Cl-

Hydrogen Bonding Hydrogen bonded to N, O, or F, is attracted to the N, O, or F of another molecule. Not actual bond, just attraction Hydrogen “Bond” H F

Boiling Point of H compounds

Boiling Point of H compounds