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INTRO TO CHEMICAL BONDING / Ionic Bonding
IV. INTRODUCTION TO CHEMICAL BONDING
There are two principal types of chemical bonding: Ionic and Covalent. We will approach each separately, but later we will see that the type of bonding is not always black and white, but some shade of gray.
A. Ionic Bonding
1. Lewis symbols: For each valence electron the atom has place a dot (or mark) next to the symbol, keeping the electrons unpaired if possible.
LiBeBCNOF
Notice that when done correctly, the Lewis symbols reflect what we learned about the number of unpaired electrons in atomic orbital theory.
2. Formation of ionic bonds
a) Ionic bond: the mutual attraction between the oppositely charged ions.
- The ions form when an electron(s) is transferred form an element with a low IE to an element with a high EN.
- The elements with low IE’s are metals, and those with high EN’s are nonmetals
- Therefore, a metal and a nonmetal make an ionic compound.
- All ionic compounds are solid due to the strong attraction between the oppositely charged ions.
Example(1) Na with Cl
Example(2) Mg with S
Example(3) K with S
Example(4) Al with F
b) “Octet” “rule”: After chemically combining atoms end up with eight electrons in their outer most occupied principal energy level.
3. Writing formulas from charges
After you understand why and how ionic compounds form from elements, you can use this shortcut method to determine formulas for ionic compounds.
This method is based on two facts:
a) All compounds are electrically neutral.
b)The charges on many ions are predictable.
Group I Metals: Form 1+ charged ions upon making an ionic compound.
Remember elements do NOT have charges! The charges develop
as a result of forming a compound.
Group II:
Group III Metals:
Group VII:
Group VI:
Group V Nonmetals:
Example(5): What is the formula of the compound that forms between Ca and Br?
Example(6): What is the formula of the compound that forms between Na and N?
Example(7): What is the formula of the compound that forms between Al and S
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INTRO TO CHEMICAL BONDING / Covalent Bonding
B. Covalent Bonding
1. Formation of covalent bonds
a) When nonmetals bond to each other they share electrons, since one nonmetal is incapable of pulling an electron away from another nonmetal.
Example(1): F with F F1+ + F1-
b) Single covalent bond: the sharing of one pair of electrons between two atoms.
Example(2): F with F
Example(3): H with H
Example(4): H with F
2. Atomic orbital overlap
To understand how nonmetals form bonds, we must look at their atomic orbitals.
a) A covalent bond is formed between two atoms by overlapping their atomic orbitals. The attraction of the 2 nuclei for the region of increased electron cloud density holds the atoms together.
Example(5): H with H
b) The energy of the atoms depends on the distance between them.
c) If an orbital is completely filled it cannot overlap with another orbital that also contains electrons. This follows from the basic principal that no more than 2 electrons can occupy any single orbital.
Example(6): He with He
d) If an orbital contains only one electron (i.e. an unpaired electron) it is available for bonding.
Example(7): F with F
Example(8): H with F
Example(9): H2S
3. Dot formulas
Example(10): HF
Example(11): H2S
Example(12): SiCl4(For almost all simple molecules, such as this, the single atom is in the center and the multiple atoms are bonded to it.)
4. Using dot formulas to determine the formula of simple covalent compounds
Example(13): What is the formula of the compound that forms between S and Cl.
Example(14): What is the formula of the compound that forms between P and F.
5. Multiple bonding between two atoms
a) Double covalent bond: the sharing of two pair of electrons between two atoms
Example(15): O2
Example(16): CO2
b) Triple Covalent Bond: the sharing of three pair of electrons between two atoms
Example(17): N2
6. More dot formulas
a) Hydrocarbons (compounds containing only carbon and hydrogen)
Example(18): CH4
Example(19): C2H6
Example(20): C3H8
Example(21): C2H4
Example(22): C2H2
b) Oxygen attached to a central atom
When doing the dot formula of a compound or polyatomic ion containing oxygen attached to a central atom, put the electrons around the central atom in pairs.
Example(23): CO2
Example(24): SiO2
7. Polyatomic ions
a) Recognizing polyatomic ions
b) Dot formulas of anions
Example(25): SO42-
Example(26): PO43-
Example(27): ClO31-
c) Dot formulas of cations
Example(28): NH41+
Example(29): BrCl21+
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INTRO TO CHEMICAL BONDING / Covalent Bond Properties
C. Covalent Bond Properties
1. Bond Length (BL):The distance between two atoms, measured from nucleus to nucleus. BL is determined by two factors:
a) Atom size: The larger the atoms bonded together, the longer the bond.
Example(1): Group VII diatomic elements
F21.43 (143 pm)
Cl21.99 (199 pm)
Br22.28 (228 pm)
I22.66 (266 pm)
1 (angstrom) = 1 x 10-8 cm = 1 x 10-10 m
Example(2):
HF
HCl
HBr
HI
b) Number of bonds: The more bonds between two atoms the shorter the BL.
Example(3):
C2H6C —C1.54 (154 pm)
C2H4C = C1.34 (134 pm)
C2H2CC1.20 (120 pm)
2. Bond Energy (BE): The amount of energy required to completely dissociate two bonded atoms.
AB A + B H = BE
a) Breaking a bond requires energy - endothermic.
Making a bond releases energy - exothermic.
(Recall when we first bonded two H atoms together. We said that the energy was lower when bonded then when separate.)
The same two factors that determine BL also determine BE.
b) Atom size: the larger the atoms the weaker the bond (lower BE).
Example(4):
Cl2 60.7 kcal/mol (254 kJ/mol)
Br2 52.3 kcal/mol (218 kJ/mol)
I2 42.5 kcal/mol (179 kJ/mol)
(F2 is an exception, it can be explained, but we won’t worry about it.)
c) Number of bonds: the more bonds between two atoms the stronger the bond (higher BE).
Example(5):
C —C83 kcal/mol (348 kJ/mol)
C = C145 kcal/mol (615 kJ/mol)
CC194 kcal/mol (812 kJ/mol)
d) Effect of bond strength on chemical reactivity
Example(6): Nitrogen and chlorine both have an EN of 3.0, but at room temperature nitrogen is essentially inert and chlorine is very reactive.
Example(7): Oxygen is only slightly reactive at room temperature, but at elevated temperatures it is very reactive.
3. Vibrational frequency
The bond between atoms is not rigid, like masses attached by a spring, the atoms vibrate. When the bonded is stretched, the atoms are pulled together; when the bond is compressed the atoms are pushed apart.
a) The stronger the bond, the greater the vibration frequency.
Example(8): Which C to O bond has the higher vibrational frequency, H2C=O or H3COH ?
b) The more massive the atoms, the lower the vibrational frequency.
Example(9): Which Cl2 molecule has the higher vibrational frequency, 35Cl35Cl or , 37Cl37Cl ?
4. Resonance
a) A molecule exhibits resonance if it has more than one valid dot formula.
Example(8): SO2
b) The actual structure is the average structure of all of the dot formulas
c) Bond order (BO): The number of bonds between two atoms.
Example(9): What is the nitrogen to nitrogen bond order in N2?
Example(10): What is the carbon to oxygen bond order in CO2?
Example(11): What is the sulfur to oxygen bond order in SO2?
For a molecule with resonance: BO =
Example(12): What is the sulfur to oxygen bond order in SO3?
Example(13): What is the carbon to oxygen bond order in CO32-?
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INTRO TO CHEMICAL BONDING / Polar Covalent Bonding
D. Polar Covalent Bonding
1. Polar bonding
a) Role of electronegativity difference in determining bond type
Example(1): Na with F
Example(2): F with F
Example(3): H with F
b) Polar covalent Bonding: the unequal sharing of electrons between atoms. It occurs anytime nonmetals with a difference in EN are bonded.
Example(4): Cl with F
Example(5): Cl with S
Example(6): H with S (Assume that H has the lowest EN of the nonmetals.)
c) Degrees of polarity
Example(7):
HF
HCl
HBr
HI
Example(8):
NaCl
MgS
AlP
2. Polar Molecule: has a partial positive end and a partial negative end (a dipole).
FormulaStructure Polar Polar
Bonding?Molecule?
F2
HF
CO2
H2O
BF3
NH3