Chemical Bonding Notes

Neutral, “free” atoms are rarely found in nature Most elements exist as part of a compound Compounds are chemically bonded atoms

Chemical bond: strong attractive force that exists between atoms or ions in a compound (+ nuclei and – electrons)

So, why do atoms bond? To become more stable. Bonding involves only valence electrons.

Remember….. Noble gases are extremely STABLE because their outer shell is full of electrons With the exception of He, all noble gases have 8 valence electrons

Octet Rule: Atoms tend to gain, lose, or share electrons until they are surrounded by 8 valence electrons

Types of Chemical Bonds Ionic: formed when electrons are transferred from one atom to another Usually formed between metals and non-metals.

Creates positive and negative ions which are electrostatically attracted

electrons are shared between atoms Usually formed between non-metals 2. Covalent: formed when pairs of electrons are shared between atoms Usually formed between non-metals Polar covalent: electrons are shared unequally Non-polar covalent: electrons are shared equally

3.Metallic: – results from attraction between positive ions and highly mobile, delocalized electrons Delocalized electrons – valence electrons not held by any specific atom, but free to move from one atom to another See “electron sea” model

Ionic & Covalent Properties Ionic vs. Covalent___ “formula unit” “molecule” high melting & low melting & boiling points boiling points electrolyte non-electrolyte hard & brittle pliable Soluble in H2O polar: H2O soluble non-polar: insoluble in H2O

Metallic Properties 1. Luster: free electrons cause most metals to reflect light 2. High density: particles tightly packed in lattice 3. High melting and boiling points: strong attractive forces between positive ions and delocalized electrons

4. Good heat conductors: electrons transmit the energy vibrations 5. Good electrical conductors: mobile electrons flow in at one end and the same number flow out the other 6. Malleable and ductile: Can be bent into different shapes, and drawn into a long, thin wire. Distortion does not disrupt the bond

Ionic Bonds We use Electron Transfer Diagrams to represent Ionic Bonds Ionization: The formation of an ion 1) Cation: positive ion formed by the loss of one or more valence electrons Metals tend to form + ions 2) Anion: negative ion formed by the gain of one or more valence electrons Non-metals tend to form - ions

Oxidation Numbers The number of electrons that are transferred from an atom to form an ion. This becomes the positive or negative charge of the ion. Example: Na  loses 1 valence electron  becomes Na+1 It’s oxidation number is 1.

Electron Transfer Diagrams (used to represent ionic bonds) To write: 1) use equation format 2) left of arrow: Show electron dot diagrams 3) right of arrow: show ions formed, and coefficients to give proper ion ratio

Examples Na + Cl  Mg + I  Li + P  Al + S  Sr + O 

Lewis Structures (used to represent covalent bonds) Formulas in which: 1)Element symbols represent the nuclei and core electrons. 2) Dot pairs (a dash) between two symbols represent covalent bonds (shared pairs) 3) Dot pairs around the outside of a symbol represent “lone pairs” (unshared pairs)

Lewis Structures (used to represent covalent bonds) Single bond = 1 dash Double bond = 2 dashes Triple bond = 3 dashes Central atom is the least electronegative atom.

Drawing Lewis Structures Count valence electrons Choose center atom Connect atoms to central atom with single bond (dash) Add unshared pairs around all surrounding atoms (except H). Check for stability (each atom should have 8 valence electrons). If not, rearrange unshared electrons creating multiple bonds where necessary.

Lewis Structure Examples: HCl NH3 CH3F H2S CO2

VSEPR Theory VSEPR: Valence Shell Electron Pair Repulsion Theory states that the electrostatic repulsion between the valence-level electron pairs surrounding at atom causes these pairs to be oriented as far apart as possible.

VSEPR Theory VSEPR allows for the prediction of geometric shapes of molecules. “Lone Pairs” of electrons affect the shape but are not included in the shape.

Geometric Molecule Shapes Linear Bent Trigonal pyramidal Trigonal planar Tetrahedral

Linear

Bent

Trigonal Pyramidal

Trigonal Planar

Tetrahedral

THE END!