CHAPTER 8 AP CHEMISTRY

COVALENT BONDING Polar bonds Nonpolar bonds Uneven attraction for the shared electrons Nonpolar bonds Bonding electrons are shared equally by bonded atoms, balanced distribution of charge

CONTINUE Molecule Group of two or more atoms (same or different type of elements) held together by covalent bonds (sharing electrons) and can exist in nature Diatomic - two atoms of the same element covalently bonded Chemical formulas have symbols, numerical subscripts, and states the type of atoms and the number of each atom Potential energy is at a minimum when attractive forces balance repulsive forces A bond is formed if the potential energy is lower as a compound than as single atoms

OCTET RULE Chemical compounds tend to form so that each atom, by gaining, losing, or sharing electrons, has an octet (eight) of electrons in its valence shell Exceptions to the octet rule Has less than a full octet around the central atom because the atom is too small (boron and beryllium) Has more than a full octet (10 or 12) around the central atom, central atom is large with empty d and f sublevels Central atom is the least electronegative atom Central atom is a non-metal found in periods 3, 4, and 5 and are surrounded by halogens

LEWIS STRUCTURE Lone pair Lewis structure Unshared pair of electrons, not involved in bonding and belongs to one atom Oxygen has two lone pair electrons Lewis structure FORMULAS IN WHICH ATOMIC SYMBOLS REPRESENT NUCLEI AND INNER-SHELL ELECTRONS, DOTS OR DASHES BETWEEN TWO SYMBOLS REPRESENT BONDS AND DOTS REPRESENT UNSHARED ELECTRONS

STRUCTURAL FORMULA Arrangement of atoms and where the bonds are formed Single bond: sharing two electrons Double bond: sharing four electrons, higher bond energy than a single bond so bond length is shorter Triple bond: sharing six electrons, higher bond energy than a double bond, so bond length is shorter

BONDS Bond energy Ionic bonds coulomb’s law energy required to break a bond Ionic bonds very stable, high melting point metal and nonmetal usually form these bonds coulomb’s law energy of interactions E = 2.31 x 10-19 J.nm (Q1Q2) r Q1 and Q2 numerical charge

CONTINUE A bond will form when the energy in a compound is lower than as separate atoms this will minimize the sum of positive (repulsive) energy and negative (attractive) energy terms and is called the BOND LENGTH

ELECTRONEGATIVITY Ability of an atom to attract shared electrons to itself page 334 polarity dipole moment +-------> б+ б- б- б- Б+

PREDICTING IONIC COMPOUNDS Ca + O <=> Ca2+ + O 2- calcium is transferring 2 electron‘s to the oxygen positive ions are negative ions are Isoelectronic ions have the same electron configuration as the noble gas it is close to.

LATTICE ENERGY change in energy when gaseous ions form ionic solids page 343 lattice energy = k(Q1Q2), k = proportionality ( r ) constant covalent bonds with ionic characteristics dipole moment gives a covalent bond a percent of ionic character (measured dipole moment X-Y) X 100 (calculated dipole moment X+ Y-) If the compound when melted can conduct an electrical current it is called a SALT

COVALENT BONDING forces which hold nonmetal atoms together because this bond consists of sharing electrons it is very stable

LEWIS STRUCTURES eight electrons in the valence shell. When there are eight electrons in the outer most energy level the atom is stable number of valence electrons = 1 to 8 In the lewis structure a covalent bond is a line. single bond = 2 electrons shared double bond = 4 electrons shared triple bond = 6 electrons shared

WRITING LEWIS STUCTURES draw a skeleton structure by joining the atoms by single bonds central atom written first in the formula the following are usually terminal atoms H, Br, F, I, Cl, O make sure all bonds are accountable determine the number of valence electrons that have not been accounted for

RESONANCE FORMS the concept of resonance is involved whenever a single Lewis structure does not explain all bonds Formal charge the charge on an atom in a molecule or ion that is calculated by assuming that all lone pairs of electrons shared by an atom, belong to the atom Cf = Ev - (Eu + 1/2Eb) Cf = formal charge, Ev = # of valence electrons Eu= # of unshared electrons, Eb= # of bonded electrons

EXCEPTIONS TO THE OCTET RULE molecules that contain an odd number of valence electrons species that contain an odd number of electrons are paramagnetic (attracted by a charge) less than a full octet of electrons expanded octet surrounded by more than four pairs of valence electrons

BOND ENERGY enthalpy change, ΔH, to break a bond in a mole of gaseous substance endothermic if ΔH > 0 exothermic if ΔH < 0 ΔH = (bond energies broken) -(bond energies formed)

CONTINUE bond strength and length bond polarity as the number of bonds increase the length decreases making then stronger bond polarity nonpolar bond atoms share electrons equally polar bond atoms share electrons unequally ionic bond atoms do not share electrons