Bonding: General Concepts Chapter 8 Bonding: General Concepts

Chemical Bonds Bond Energy - energy required to break a bond Types: ionic and covalent

Ionic Bonding Electrostatic attractions between tightly packed, oppositely charged ions Between metals (cations) and nonmetals (anions)…can be polyatomic High melting/boiling points Ex: NaCl

Coulomb’s Law E = (2.31 X 10-19 J * nm)(Q1Q2/r) Calculates the energy between two ions E -> Energy in Joules R -> distance between ion centers in nm Q1 and Q2 are ions’ charges When E is negative, the pair of ions has a lower energy than the ions separated (ions are attracted) When E is positive, the repulsive energy is greater (occurs with two like-charged ions) In nature, a system wants the lowest possible energy

Covalent Bonding Molecules where electrons are SHARED by nuclei Equal sharing of electrons occurs between the two atoms Nonequal sharing = polar covalent

Identical Atoms Diatomic atoms… Systems of energy favor LOWER energy If a system can lower energy by forming bonds, it will happen

Bond Lengths Distance at which the system has minimum energy

Polar Covalent Bond Covalent bond where one atom pulls more on an electron than the other Ex: H-F (fluorine slightly negative) + and - are used to represent slight charges Water is another example Due to electronegativity

Electronegativity (electron affinity)…ability to attract shared electrons Higher electronegative atom will have the slight negative charge Higher difference in electronegativities, more likely to be an ionic bond

Bond Polarity Dipolar/dipole moment occurs when a molecule has a slight positive and slight negative charge. Arrow points towards the slight negative charge with a plus sign on the other end of the arrow Electrostatic potential diagram also used (red = electron-rich, blue = electron-poor)

Polar Bond without Dipole Moment Polar bonds where slight charges cancel each other out No lone electrons

Ion Electron Configuration and Sizes Atoms of stable compounds (ionic and covalent), have noble gas electron configurations Covalent: electrons are shared so that the valence electron configurations of both nonmetals attain noble gas e- configurations Ionic: Nonmetal achieves e- configuration of the next noble gas atom and the metal’s valance electrons are emptied so both ions achieve noble gas e- configurations

States of Ionic Compounds Usually “ionic compound” means a solid state where ions are close to one another/interacting, minimizing the - - and + + repulsions and maximizing the + - attractions

Trend in Ions’ Sizes Positive ion = loss of e- (cation is smaller than original atom) Negative ion = gain of e- (anion is larger than neutral atom) Ion size increases going down a group Horizontally is more complicated (metals vs. nonmetals)…..

Isoelectronic Ions Ions that have the same electron configuration (b/c have the same number of electrons) O2-, F-, Na+, Mg2+, Al3+ Same number of electrons, but different numbers of protons From O2- to Al3+ attraction to nucleus (increase in protons) increases so smaller ions

Bonds Single Bonds: one pair of electrons shared Double bond: two pairs of electrons shared Triple bond: three pairs of electrons shared More bonds = shorter bond length More bonds = stronger bonds Bond energies - amount of energy required to break bond

Breaking Vs. Forming Bonds To break bonds, energy must be added (required, endothermic, energy is positive) To form bonds, energy must be removed (given off, exothermic, energy is negative) Change in enthalpy: DH = sum of the energies required to break old bonds (positive signs) plus sum of energies released in formation of new bonds (negative signs)

Bond Energies Energy stored in bonds can be used to determine energy accompanying various chemical reactions. Bond dissociation energies = energy required to break bonds (positive kJ/mol, endothermic)… reverse for forming bonds. Table 8.4 in book DH = all bonds broken – all bonds formed Must use a balanced equation!

CH4(g) + 4F2(g)  CF4(g) + 4HF(g) EXAMPLE Use data 8.4 in your book to determine DH for the following reaction. CH4(g) + 4F2(g)  CF4(g) + 4HF(g) Answer: -1932 kJ

Localized Electron Bonding Model…Highlights Defined: A molecule is composed of atoms that are bound together by sharing pairs of electrons using the atomic orbitals of the bonded atoms Localized to one of the atoms (LONE PAIR) Localized to the space between atoms (BONDING PAIR)

LE Model Parts 1. Description of the valence electron arrangement in the molecule using Lewis structures 2. Prediction of the geometry of the molecule using the valence shell electron-pair repulsion (VSEPR) model 3. Description of the type of atomic orbitals used by the atoms to share electrons or hold lone pairs

LE Model Part 1: Lewis Structures All atoms want to have noble gas electron configurations Only valence electrons are included Determine total valence electrons Determine layout Determine bonds Place remaining valence electrons

Lewis Dot Structure Examples Ammonia: NH3 Water: H2O Acetylene: C2H2 Carbon Tetrachloride: CCl4 Dihydrogen Selenide: H2Se

Exceptions to Octet Rule If valence electron total is odd, the octet rule doesn’t work Some atoms do not require all 8 valence electrons (or have more than 8) These molecules can exist in stable form Boron: forms compounds where boron has less than 8 electrons (ex: BF3) More than 8 only happens with elements in period 3 and beyond ex: SF6 See pg. 371 purple box for rules if needed

Exception Examples Draw the Lewis dot structure for the following: ClF3 XeO3 RnCl2 BeCl2 ICl4-

Resonance Molecules with more than one possible electron dot structure Do not switch back and forth Molecules exist as a mixture (hybrid) of the resonance forms Use double headed arrow to signify Example:

Resonance Example Draw the Lewis dot (Localized Electron Model) structure(s) of CO32-

# valence e- assigned = (# lone pair e-) + (½ # shared e-) Formal Charges Can’t use oxidation numbers because electrons are not shared evenly between atoms (electronegativities) Atoms can be assigned formal charges using the following: Formal Charge = (# valence e- on atom) – (# valence e- assigned to the atom in the molecule) # valence e- assigned = (# lone pair e-) + (½ # shared e-) Atoms want to have formal charges close to zero Negative formal charges should be on the most electronegative atom ***ESTIMATES of charges (not exact charges)

Example Assign formal charges to each atom in CO2 Which is more likely? Answer: two double bonds b/c all formal charges are zero Draw all resonance structures and show the most stable one for SCN- Answer: two double bonds b/c N more electronegative than S

LE Model Part 2: VSEPR Model Molecular structure shows 3-D arrangement of molecules Based on minimizing electron-pair repulsions (bonding and nonbonding electrons will be placed as far apart as possible)

No Unshared Pairs Linear (180°) BeCl2 Trigonal Planar (120°) BCl3

No Unshared Pairs Tetrahedral (109.5°) CClF3 Square Planar (90°)

With Unshared Pairs Bent (104.5°) H2O Pyramidal (107°) NH3

More to Consider Tigonal bipyramidal See-saw T-shaped Octahedral  Deal with atoms that do not obey the octet rule

Examples H2S CO2 PCl3 OCS H2CO CH4 N2H4