CHEMICAL BONDING What forces hold atoms and molecules together?

What is a bond? Bond: The force holding 2 or more atoms together making them function as a unit. Examples: Ionic, covalent, metal Bond Energy: The energy required to break a given chemical bond A measure of the general strength of the bond

Three General Types of Bonds Covalent bond: the sharing of valence electron pairs between atoms. Usually found between nonmetals. Ionic bond: the transfer of valence electrons from a metal to a nonmetal. Metallic bond: attractive force holding pure metals together.

Why do atoms bond? Octet Rule: Atoms tend to gain, lose, or share electrons until they are surrounded by 8 valence electrons (4 electron pairs). All noble gases except He have an s 2 p 6 configuration. Hydrogen follows the duet rule.

Ionic Bond Formation Neutral atoms come near each other. Electron(s) are transferred from the Metal atom to the Non-metal atom. They stick together because of electrostatic forces, like magnets. Non-MetalMetal

Why does a metal & nonmetal form an ionic bond? The difference in electronegativity determine bond type. Electronegativity: the tendency of an atom in a molecule to attract shared electrons to itself. High Electronegativity – high attraction of electrons Low Electronegativity – low attraction of electrons Increases as you move across a period Decreases as you move down a group Fluorine has highest value (4.0)

Ionic Bond cont’ In ionic compounds, there needs to be a large difference in electronegativity values between atoms. Greater than 1.7 Metals = low electronegativity Nonmetals = high electronegativity Example: NaCl Na = 0.93 Cl = – 0.93 = 2.23

Crystal Ionic Structure

Examples of Forming an Ionic Compound Potassium Chloride Barium Fluoride

Metallic Bonding Metals: low electronegativity = Don’t attract each other’s electrons Metals consist of closely packed cations floating in a “sea of electrons”. All of the atoms are able to share the electrons; the electrons are not bound to individual atoms.

“Sea of Electrons” Electrons are free to move through the solid. = Metals conduct electricity.

Alloys Alloys are mixtures of 2 or more elements, at least 1 is a metal Made by melting a mixture of the ingredients, then cooling Brass: Cu and Zn Bronze: Cu and Sn Steel: Fe and C Stainless Steel: Fe, C, Cr & V 14k Gold: 58% Au, 42% Ni, Cu & Ag

Covalent Bonding Atoms now share pairs of valence electron. Shared electron pair = bonding pair Can share 1, 2 or 3 pairs of electrons to form single, double or triple bonds. Occurs between 2 non-metals. Example: H 2 O

Covalent Bonding Two Types of Covalent Bonds: 1. Non-Polar Covalent Bond: A bond formed between 2 atoms in which electrons are shared equally between the 2 atoms. Electronegativity difference = less than 0.5 Example: Cl 2 Electronegativity difference = = 0

Example of Non-Polar Bonding

Covalent Bonding 2. Polar Covalent Bonding A bond formed between 2 atoms where electrons are not equally shared. Electronegativity difference = 0.5 – 1.7 The atom with the greater electronegativity value pulls the shared electrons closer to it’s nucleus. Example: Bond between hydrogen and oxygen Hydrogen = 2.1 Oxygen = 3.5 Difference = 1.4

Example of a Polar Bond

Electron Density Models PolarNon Polar Ionic

3 Different Bonds

Intramolecular vs. Intermolecular Bonding Intramolecular Forces: Attractive forces that occur between atoms in a molecule; chemical bonds Examples: ionic, covalent or metallic bonds Intermolecular Forces: Attractive forces that occur between molecules Examples: Dipole-Dipole attraction, hydrogen bonding, London Dispersion forces

Dipole-Dipole Attraction The attraction force between the positively charged end of one polar molecule with the negatively charged end of another polar molecule.

Hydrogen Bonding Special name for unusually strong dipole-dipole attraction that occur among molecules in which hydrogen is bonded to a highly electronegative atom (such as N, O or F)

Importance of Hydrogen Bonding in Water Responsible for water’s unique properties: High surface tension High melting and boiling points High specific heat Low density in solid form vs. liquid form Crystal structure of ice Insolublility in oil