8.1 - Molecular Compounds

bond - electrons are shared/transferred in the valence shell

ionic bond - an electron from 1 valence shell is transferred to another

covalent bond - share electrons in valence shell

molecule - neutral (p = e) group of atoms that are joined by covalent bonds

diatomic molecule - same atoms of an element that are covalently bonded

molecular compound

• diff atoms that are covalently bonded

• low melting/boiling points compared to ionic compounds which have high melting/boiling points

molecular formula - tells you…

• number of atoms in a molecule

• number of molecules

• types of atoms

H2O

• 1 water molecule

• 2 hydrogen

• 1 oxygen

• 2:1

C6H12O6

• 1 glucose molecule

• 6 carbon

• 12 hydrogen

• 6 oxygen

• 1:2:1

chg in composition or ratio → chg in substance

a straight line represents one bond/one pair of electrons

TYPES OF FORMULAS

electron sharing usually occurs so atoms can attain the electron configuration of noble gases

groups 4 - 7a are more likely to form covalent bonds

single covalent bond - 1 shared pair of electrons

double - 2

triple - 3

8.2

coordinate covalent bond - 1 atom contributes both bonding electrons

(represented by →)

polyatomic ion - tightly bound group of atoms that have a charge

bond dissociation energy - energy needed to break one mole of bonds (2 molecules covalently bonded)

resonance structure - occurs when it’s possible to write 2+ valid electron dot structure with the same number of electron pairs for a molecule/ion (way to envision double bonding)

(represented by double headed arrows)

main exceptions to octet rule

• ClO2

• NO

8.3 - Bonding Theories

molecular orbital - atomic orbital overlap during chemical bonds

bonding orbital - a molecular orbital in covalent bonds that can have 2 electrons occupy it

just as atomic orbitals belong to particular atoms, molecular orbitals belong to molecules as wholes

sigma bond - 2 atomic orbitals that combine to form a molecular orbital that is symmetrical around an axis connecting the 2 atomic nuclei

pi bond

• bonding electrons not as likely to be found in the sausage shaped regions above and below the bond axis

• not symmetrical

• don’t overlap as much as sigma

• weaker than sigma

VSEPR Theory

• valence shell electron pair repulsion theory

• explains 3D shape

• repulsion bc electron pairs cause molecular shapes to adjust so that the valence electron pairs satay as far apart as possible

tetrahedral angle - 109.5°

hybridization - several atomic orbitals mix and form same total number of hybrid orbitals

orbital hybridization provides info about both molecular bonding and shape

single bond stronger than double

double stronger than triple

ane - single bond between C

ene - double bond between C

yne - triple bond between C

8.4 - Polar Bonds and Molecules

nonpolar covalent bond

• bonding electrons shared equally

• EX: H2, O2, N2, CO2

polar covalent bond - electrons shared unequally

polar molecule - 1 end is slightly negative, other is slightly positive

EX: H2O

high EN - negative

low EN - positive

electronegativity - attractive force of an atom for electrons

if (EN diff. rang) then it is (bond):

0-0.4; nonpolar covalent

0.4-1; moderately polar covalent

1-2; very polar covalent

2+; ionic

dipolar molecule (dipole) - molecule that has 2 poles

intermolecular attractions are weaker than either ionic or covalent bonds

effect of polar bonds on polarity of an entire molecule → depends on shape of molecule and orientation of polar bonds

van der Waals forces

• weakest attraction of molecule to molecule

• weaker than ionic and covalent bonds

• consists of dipole interactions and dispersion forces

dipole interactions

• polar molecules are attracted to one another

• weaker than ionic bonds

dispersion forces

• weakest of all molecular interactions

• caused by motion of electrons

• occurs even between non polar molecules

high strength → high number of electrons in molecules

halogen diatomics

• Cl2

• F2

• Br2

• I2

• attracts mainly bc of dispersion forces

H Bonds

• attractive forces that covalently bond to a very electronegative atom

• also weakly bonded to unshared electron pair

• high EN atom

• 5% of strength of avg covalent bond

• strongest intermolecular forces

• determines properties of H2O and biological molecules such as proteins

Intermolecular Attractions and Molecular Properties

physical properties of a compound depend on type of bonding

network solid - all atoms are covalently bonded to each other