Ch. 8 Bonding: General Concepts
chemical bond - force that holds groups of atoms together group function as a unit
bond NRG – NRG required to break bond indicates strength of bond
bond length – distance between two atoms bonded together indicates most stable(least amount of NRG)
state between two atoms
ionic bond – electromagnetic force that holds two oppositely charged ions together formed between cations and anions cations formed when metals lose e-‘s
▪ Na Na+ + e- (oxidation) anions formed when nonmetals gain e-‘s
▪ Cl + e- Cl- (reduction) opposite charges attract(electromagnetic force)
▪ Na+ + Cl- Na+Cl-
why do substances form ionic bonds??? lowest possible NRG for the
system(substances involved with bonding) see figure 8.9 on page 355
1. change Li(s) to Li(g) endothermic
Li(s) + 161 kJ/mol Li(g)
2. Li(g) oxidizes endothermic
Li(g) + 520 kJ/mol Li+ + e-
3. fluorine molecules separate and form fluorine atoms endothermic
½ F2(g) + 77 kJ/mol F(g)
4. fluorine reduces exothermicF(g) + e- F-(g) + 328 kJ/mol
5. ionic bond formed extremely exothermicLi+(g) + F-(g) Li+F- + 1047 kJ/mol
total endothermic processes = 758 kJ/moltotal exothermic processes = 1375 kJ/mol
NET NRG = 617 kJ/mol overall process is exothermic so an ionic bond
forms
metals lose e- (oxidation) nonmetals gain e- (reduction)
ionic compound formula total oxidation = total reduction
Mg2+ F- MgF2 utilize criss-cross method
Mg2+ F- MgF2
dissolving ionic crystals
covalent bond – force of attraction between 2 atoms when e-‘s are shared each nucleus attracts the other atoms e-‘s balance between attraction and repulsion
bonds result from system trying to attain lowest possible NRG state two driving forces in nature
1) lowest NRG2) highest entropy(disorder/chaos)
ionic and covalent bonds generally form to attain the lowest NRG state for atoms involved
single covalent bond – 1 pair e- shared H-H, F-F, Cl-Cl, Br-Br, I-I
double covalent bond – 2 pair e- shared O=O
triple covalent bond – 3 pair e- shared N≡N
molecular orbital most probable location of e-‘s when covalently
bonded
sigma () bond centers along the internuclear axis. single covalent bond
pi () bond occupies space above and below internuclear axis.
2nd or 3 covalent bond
electronegativity attraction an atom has for another atom’s e- ‘s arbitrary scale – based on F varies periodically (click here)
▪ generally increases across and decreases down
electronegativity difference can generally predict type of bond if diff. > 1.7 = ionic bond if 1.7 > diff. > 0.3 = polar covalent bond if diff. < 0.3 = nonpolar covalent bond
nonpolar covalent bond pure covalent bond sharing of e- ‘s is equal no poles/charges created
• Lewis structures localized e- model for diagramming bonds
and molecular shapes
1) total valence e- of all atoms in molecule• HCl = 8 valence e-
2) write symbols for each element• least number of = interior atom• hydrogen is always exterior atom• H Cl
3) add a pair of e- between atoms bonding together
• H : Cl valence e- remaining = 6
4) add remaining e- in pairs to exterior atoms to form octets
• if e- remain add them to interior atoms to form octets
• not H, H forms duets
• if necessary, move e- pairs to create octets
• H : Cl
5) all shared pair of e- become dashes
• H - Cl
::
:
:
::
6) follow VSEPR for shape• VSEPR – valence shell electron pair
repulsion theory• model used to predict geometry/shape of a
molecule based on the repulsion of e- pairs• e- pairs repel each other to maximum distance
• in 2-dimension = 90o
• in 3-dimension = 109.5o
Molecular Geometry• 3-dimensional shapes determine the
physical and chemical properties of molecules• example – sucrose- its molecular shape fits the
nerve receptors of the tongue for sweetness• sugar substitutes(Splenda, Nutrasweet, …) have similar
shapes as sucrose
1) linear• all atoms lie in a straight line• HCl, CO2
2) bent• atoms not in straight line• e- pairs point to 4 corners of tetrahedron
3) trigonal pyramidal• 3 atoms bonded to central atom and a pair
of nonbonding e-
4) trigonal planar• 3 groupings of e- around central atom• all atoms lie in same plane
5) tetrahedral• 4 groupings of e- around central atom
resonance ability to draw more than one acceptable
shape for a molecule originally believed molecule resonated
between different shapes▪ benzene, nitrate ion
actual structure is an average of all resonant images
exceptions to octet rule C, N, O, F always follow octet rule
Be and B often have less than 8
3rd period and heavier usually follow▪ some may exceed by putting e- in unoccupied d-
orbitals▪ when writing Lewis structures follow octet rule
▪ if e- remain add them to elements with d orbitals
• polar molecules• molecule with oppositely, partially charged
atoms on opposite sides• aka – dipoles, dipole moments
• molecule that has an asymmetrical distribution of charge
• partial charges not evenly distributed around central atom
• polar molecules must have:• polar bonds(partial charges)• unevenly distributed partial charges
• HCl• H = 2.1, Cl = 3.0• (electroneg diff = 0.9) = bond is polar• H = δ+, Cl = δ-• linear shape• polar molecule
• N2 • N = 3.0• (electroneg diff = 0) = bond is nonpolar• nonpolar molecule
• H2O• H = 2.1, O = 3.5• (electroneg diff = 1.4) = bonds are polar• H = δ+, O = δ-• bent shape• polar molecule
• NH3
• N = 3.0, H = 2.1• (electroneg diff = 0.9) = bonds are polar• H = δ+, N = δ-• trigonal pyramidal shape• polar molecule
• CCl4• C = 2.5, Cl = 3.0• (electroneg diff = 0.5) = bonds are polar• C = δ+, Cl = δ-• tetrahedral shape• nonpolar molecule
hybridization formation of hybrid orbitals from atomic orbitals of
similar NRG sp3 hybridization
sp2 hybridization
sp hybridization
naming binary molecules1) determine # of 1st element
▪ use prefix if more than one▪ 1=mono- 2=di-▪ 3=tri- 4=tetra-▪ 5=penta- 6=hexa-▪ 7=hepta- 8=octa-▪ 9=nona- 10=deca-
2) name element
3) determine # of 2nd element▪ use prefix except if bonded to H
4) use root of element name
5) end with -ide
polar covalent bond bond in which e- ‘s are shared unequally
▪ electronegativity of one atom is higher than other