Bonding

Electronegativity – attraction to another atom’s electrons Chemical Bonds A force of attraction between 2 atoms The “glue” that holds chemical compounds together **Involves VALENCE electrons** Electronegativity – attraction to another atom’s electrons Table S Outer electrons!

3 Types of Bonds Ionic Bonds – metal + nonmetal atoms “transfer of electrons” Electronegativity Difference (E.N.D.) greater than 1.7 Metallic Bonds – metal + metal atoms “sea of electrons” Covalent Bonds – nonmetal + nonmetal atoms “sharing electrons” less than 1.7 Catch! Tug-of-War

Energy of Bonds Exothermic – Energy Released Forming Bonds Atoms are more stable together, and so, RELEASE ENERGY to be at a lower, more stable energy state Endothermic – Energy Absorbed Breaking Bonds Overcome the attraction between the atoms to break a bond, so we ADD ENERGY The stronger the bond, the closer together the atoms are, the more energy is needed to break the bond

X Lewis Dot Structures Use dots to represent electrons Diagram of the valence electrons in an atom or ion Use dots to represent electrons X 3 7 Each consecutive dot will go opposite the one before it, before getting paired up. 1 6 5 2 Octet Rule – 8 Electron Rule Atoms will gain, lose, or share electrons to have 8 valence electrons (except H, He, Li, Be, B) 8 4

Electronegativity Difference E.N.D. > 1.7 One atom is far more electronegative than the other, results in a transfer of electrons 0.4 < E.N.D. < 1.7 One atom is slightly more electronegative than the other, results in an UNEQUAL sharing of electrons E.N.D. < 0.4 Two atoms of equal electronegativity, results in EQUAL sharing of electrons Ionic Bond Polar Covalent Bond Nonpolar Covalent Bond

Ionic Bonding E.N.D. > 1.7 Transfer of electrons Metal + Nonmetal One atom far more electronegative than the other Transfer of electrons Metals give up electrons to nonmetals Metal + Nonmetal Metal (cation) Nonmetal (anion) Lewis Structures Write the neutral atom electron dot structures Draw arrows to show the movement of electrons from metal(s) to nonmetal(s) Draw ion electron dot structures bonded together Na + Cl → [Na]+[ Cl ]-

Ionic Solids – Properties Solid substances formed from ionic bonds Crystalline structure High melting point Dissolve in water (aqueous solution) Don’t conduct heat or electricity as SOLID Conduct electricity as aqueous solutions or liquids

Metallic Bonding 2 or more metal atoms in a sea of electrons Strong bonds form between 2 metals in liquid or solid phase Evidence: high melting & boiling point Properties Conductive of heat & electricity LIQUID SOLID Malleable Ductile High Melting/Boiling Points

Covalent Bonding E.N.D. < 1.7 Sharing of electrons Nonmetal atoms of similar electronegativity will share electrons to complete octets Sharing of electrons Polar Covalent – unequal sharing Nonpolar Covalent – equal sharing

Molecules Smallest unit of a compound made of covalent bonds Exist in all 3 phases of matter (depending on forces of attraction between molecules) Properties Soft Low melting points Do not conduct heat/electricity

sand (silicon dioxide) Molecules Network Solids – bond in a strong network of atoms Have slightly different properties due to the network Harder High melting points Do not conduct heat/electricity (except graphite) Coordinate Covalent Solids – Occur in polyatomic ions One atom donates BOTH electrons to a bond Compounds with polyatomic ions contain BOTH IONIC & COVALENT BONDS!! diamond graphite aesbestos sand (silicon dioxide)

Lewis Structures Add up valence electrons The lowest electronegative element goes in the center (usually there is only 1 atom) Arrange all the other atoms around the central atom Draw a single bond from each outside (terminal) element to the central atom Draw dots to complete octets for outside atoms Add any extra electrons to complete the octet on the central atom If we don’t have enough electrons to complete this octet, make a multiple bond (double or triple) – C, N, O, P, S Count electrons – check for all valence electrons used

Molecule Geometry VSEPR (Valence Shell Electron Pair Repulsion) Electrons want to be as far apart as possible Electron-pair geometry refers to the number of electron pairs surrounding the central atom Molecule geometry refers to Lewis Structures and depends upon both bonding & nonbonding electrons NOTE – double and triple bonds “count” as ONE pair

Molecular Geometry (VSEPR) # Bonds (terminal e- pairs) Examples Shape 1, 2 H—F , O=C=O Linear 3 Trigonal Planar 4 Tetrahedral 5 Trigonal Bipyramidal 6 Octahedral

Molecular Geometry 2 1 Bent 3 Trigonal Pyramidal Linear T-Shaped 4 # Bonding Pairs # Nonbonding Pairs Shape 2 1 Bent 3 Trigonal Pyramidal Linear T-Shaped 4 See Saw 5 Square Pyramidal Square Planar

Molecule Polarity 2 factors affect molecule polarity Bond Polarity If a molecule contains NO polar bonds, it will be nonpolar If a molecule contains polar bonds it could be either polar or nonpolar Molecule Geometry (Shape) If a molecule is symmetrical there is an equal distribution of electrons and so is NONPOLAR If a molecule is asymmetrical there is an unequal distribution of electrons and so is POLAR

Intermolecular Forces Forces of attraction between molecules – based on polarity (a.k.a. Van der Waals Forces) More strongly polar moleculs will have greater Intermolecular Forces (IMFs) Affect melting and boiling points, surface tension, viscosity

Intermolecular Forces Dipole-Dipole Interactions The slightly positive end of one molecule is attracted to the slightly negative end of another Hydrogen Bonding A specific type of dipole interaction where the slightly positive H is attracted to the slightly negative N, O, or F of another

Hydrogen Bonding in Water WATER IS POLAR!!!!! Slightly + H is attracted to Slightly – O