Chapter 6 Chemical Bonding

Sect. 6-1: Introduction to Chemical Bonding Chemical bond – electrical attraction between nuclei and valence electrons of different atoms that binds them together Being bonded lowers potential energy and creates more stable arrangement of matter

Types of Chemical Bonding Ionic bonding – electrical attraction between cations and anions  Electrons are transferred from one atom to another forming ions Covalent bonding - sharing of electrons between two atoms

Difference in Electronegativities of the 2 atoms determines what type of bond it will have (chart on pg. 151) If difference is greater than 1.7, it is considered ionic, less than 1.7, covalent Atoms from same element always covalently bond

Nonpolar-covalent bond – electrons are shared equally between atoms (difference of about 0-0.3) Polar-covalent bond – unequal sharing of electrons (difference of )  One atom is partially negative (δ-)and the other partially positive (δ+)

Sect. 6-2: Covalent Bonding and Molecular Compounds Molecule – group of atoms held together by a covalent bond Chemical formula – indicates relative number and type of atoms in a compound Molecular formula – chemical formula for a molecular compound Diatomic molecule – only 2 atoms of same element

For a covalent bond to form, repulsive forces between electrons or protons must balance out the attractive force between protons and electrons Bond length – average distance between 2 bonded atoms Bond energy – amount of energy required to break a chemical bond (equal to amount released when bond was formed)  Pg. 168

Octet rule – atoms gain, lose, or share electrons in order to have 8 valence electrons  Exception – boron likes to only have 6  Exception – some elements are surrounded by more than 8 when paired with highly electronegative elements such as F, O, & Cl (called expanded valence)

Electron-dot Notation – shows valence electrons as dots around the element symbol, which represents nucleus and inside electrons Lewis-structures – uses pairs of dots for unshared (lone pair) electrons and dashes for shared electron pairs Structural formula – similar to Lewis- structure, but does not show lone pair electrons

Steps to drawing Lewis-structures 1.Determine # and type of atoms in compound 2.Write electron dot notation for each type of atom in compound 3.Determine total number of valence electrons in compound

4.Arrange atoms to form skeleton structure. Carbon, or least electronegative element goes in center. 5.Connect atoms by shared electron pairs. 6.Add unshared pairs of electrons until each atom has 8 electrons (except hydrogen which only wants 2) 7.Count # electrons used and be sure it matches # available, if not adjust to form multiple bonds.

Multiple Covalent Bonds Double bond – sharing of 2 pairs of electrons Triple bond – sharing of 3 pairs of electrons The more electron pairs shared, the shorter the bond length and higher the bond energy

Resonance – bonding that allows for more than one correct Lewis structure for a compound  Represented by double-headed arrow between the 2 possible structures

Sect. 6-3: Ionic Bonding and Ionic Compounds Ionic compound – composed of cations & anions, paired so that charges balance Formula unit – simplest ratio for which an ionic compound’s formula can be written Use dot structure to show how electrons transfer to form ionic compounds

Ionic compounds are arranged in a 3-D structure called a crystal lattice Lattice energy – energy released when one mole on an ionic compound is formed  Negative values show that energy is released

Comparison of Covalent/Ionic Covalent – low melting/boiling because weaker bonds; soft solids, liquid, or gas; do not conduct electricity Ionic – high melting/boiling because stronger bonds; very brittle & hard; conduct electricity in liquid state or dissolved in water, but not in solid state

Polyatomic ion – a group of covalently bonded atoms with a charge  Will bond with ions of opposite charge to form ionic compounds  Lewis structures are drawn with brackets around it and the charge written outside the brackets

Sect. 6-4: Metallic bonding Metallic bonding – chemical bond that results from the attraction between a sea of electrons and metal atoms Properties  Excellent conductor of electricity & heat  Lustrous - shiny  Malleable – hammered in to thin sheets  Ductile - drawn into wires Heat of vaporization – used to measure strength of bonds in metals

Sect. 6-5: Molecular Geometry VSEPR Theory (Valence Shell Electron Pair Repulsion Theory) – repulsion between valence electron pairs causes them to spread as far apart as possible  Central atom with 2 atoms around it and no lone electron pairs will spread apart with 180˚ bond angles and a straight configuration  See pg. 186 for full list

Unshared electron pairs act like an atom, but bond angles are slightly different Double and triple bonds act like single bonds

Hybridization – mixing of 2 or more orbitals of similar energies on the same atom to produce new orbitals of equal energy  A 2s and three 2p orbitals can combine to form four sp 3 orbitals that have an energy higher than the 2s, but less than the 2p

Intermolecular forces – force of attraction between molecules  Dipole-dipole  Hydrogen bonding  London dispersion forces

Dipole – created by equal, but opposite charges separated by a small distance  Direction is from positive to negative, represented by an arrow pointing toward negative and tail crossed at positive end Dipole-dipole – forces of attraction between polar molecules Induced dipole

Hydrogen bonding – very strong dipole between hydrogen bonded to a highly electronegative atom (N, O, F) attracted to an unshared pair of electrons of an electronegative nearby atom

London Dispersion Forces – weak intermolecular forces caused by instantaneous dipoles  The only type of intermolecular forces acting between noble gases and non-polar molecules  Forces increase with increasing atomic mass