Lewis Structures and Covalent Bonding

• Covalent bonding in Honors Chemistry, taught by Mrs. Calder

Review of Matter

• Noble Gases exist as atoms (monoatomic elements)
• Metals and non-metals form ionic compounds/salts by transferring electrons
• Elements share electrons instead of giving them up, which is a covalent bond

About Covalent Compounds

• The smallest form is a molecule, not a formula unit
• They are sometimes called "Molecular" compounds

Molecular Information

• Diatomic molecules have only 2 atoms of a single element (HOFBrINCI)
• Molecular compounds consist of atoms when different elements combine in a specific ratio, like H2O and NH3
• Ionic compounds were lattices
• Melting and boiling points are low, compared to ionic compounds
• Consist of two or more non-metals combined
• Molecular compounds can exist as solids, liquids, or gases
• A molecular formula shows how many atoms of each element a molecule contains
• A molecular formula is the exact number of atoms, unlike ionic compounds
• Example of Ethane C2H6 illustrates this
• Structural formulas show the arrangement of atoms
• A ball and stick model is the best way to display atoms because it shows a 3-dimensional arrangement

Covalent Bonds

• Covalent bonds fulfill the octet rule by sharing electrons
• A single covalent bond is when only one pair of electrons is shared
• A pair of shared electrons is represented as a dash (-) in structural formulas
• Unshared electrons remain as dots
• Carbon joined with 4 hydrogens forms methane CH4

Octet Rule

• Octet Rule is when a compound tends to form so that each atom has an octet (8) of electrons in its highest energy level by gaining, losing, or sharing electrons
• Duet Rule applies to H and He

Drawing Lewis Structures

• Find the number of valence electrons in each atom and add them up
• Draw the atoms next to each other in the way they will bond (if only one atom of an element, make it the center)
• Add one bonding pair between each set of connected atoms (a dash = a pair)
• Add the remaining electrons as dots until all atoms have 8

Drawing Examples

• Methyl chloride (CH3Cl): C: 4 x 1 = 4, H: 1 x 3 = 3, Cl: 7 x 1 = 7, Total = 14 electrons, with carbon in the center
• Ammonia (NH3): N: 5 x 1 = 5, H: 1 x 3 = 3, Total = 8, with nitrogen in the center

More Steps for Drawing

• If you run out of electrons to add, and some atoms are short (they don't all have octets), take any electrons not involved bonding and create double or triple bonds
• Only carbon, oxygen, and nitrogen, (and a few others) can do this
• Example: Nitrogen gas (N2): N: 5 x 2 = 10, Total = 10 electrons
• Example: Formaldehyde (CH2O): C: 4 x 1 = 4, H: 1 x 2 = 2, O: 1 x 6 = 6, total = 12, with carbon central

Exceptions to Octet Rule

• Electron Deficient: less than 8
• Boron has 3 electrons in its outer energy level
• Beryllium has 2 electrons in its outer energy level
• Exceed Octet: more than 8
• Anything in the 3rd period or heavier
• Add extras to the middle atom
• Examples: S, P, I
• Examples: SF6, SF4, CIF3, PCl5
• For polyatomic ions, put the charge outside brackets
• Example: Ammonium Ion NH4+
• Boron Tetrafluoride Ion (BF4-) is another example as well as Sulfate (SO4¯) and Carbonate (CO32-)

Basic Rules to Remember

• C, N, O, F ALWAYS obey the octet rule
• Boron and Beryllium tend to be electron deficient
• 2nd row elements cannot exceed an octet
• 3rd row and up can exceed octet using d orbital but try to satisfy octet first

Coordinate Covalent Bond

• When one atom donates both electrons in a covalent bond.
• Carbon monoxide (CO) is a good example:
• Carbon and oxygen each give another single electron to share
• Shown with an arrow

Resonance Structures

• Drawn when there are 2 or more valid Lewis structures
• Allows envisioning of the bonding
• In reality, no back and forth charges occur
• The real bond is a hybrid of the 2 structures
• Example: NO2
• Ozone exhibits resonance
• Neither structure is correct, it is actually a hybrid of the two, so it is important to draw all varieties possible, and join them with a double-headed arrow
• The carbonate ion and acetate ion also demonstrate resonance

Bond Dissociation Energy

• Energy needed to break a covalent bond
• The stronger the covalent bond, the higher the energy
• Explains why some covalent compounds are so stable, as the energy to break them is too high

VSEPR Theory

• VSEPR (Valence shell electron pair repulsion) states that repulsion between electron pairs causes molecular shapes to adjust so that the valence-electron pairs stay as far apart as possible
• For VSEPR, an evaluation of the central atom is done while determining the 3D shape based on the # of bonds and lone pairs (unshared e- pairs) around that atom
• Linear: 2 Bonds, 0 unshared electrons (around central atom).
• The bond angle between the atoms attached to the center is 180° so the bonds are as far away from each other as possible
• 2 pair electrons and 0 unshared (both in bonds)
• BeH2 is an example
• Double bonds are ignored in double/triple bonds
• Carbon Dioxide is an example
• Bent: H2O, 2 bonds, 2 lone pairs of electrons around the central atom cause the 2 lone electron pairs to exert extra repulsion on the two bonding hydrogen atoms which creates a slight compression to a ~104°bond angle
• 4 pairs of electrons total – 2 bonds, 2 lone pairs
• 3 pairs with 1 unshared is called bent, but larger angle
• Trigonal Planar: 120° example molecules are BH3 with 3 bonds around the central atom and no lone pairs The hydrogen atoms are as far apart as possible at 120°
• The molecule is all in a plane and is two dimensional, which is flat
• 3 bonds no lone pairs Trigonal Pyramidal: Molecules like NH3 have 3 bonds & 1 lone pair around the central atom, pushing down more causing other 3 to be closer together, creating a slight compression to ~107° bond angle
• 4 pairs total, 3 bonds, 1 lone pair Tetrahedral: 109.5° Example: CH4
• We have 4 bonds with no unshared pair with The hydrogen atoms are as far apart as possible at ~109° degree bond angle
• 4 bonds, no lone pairs.
• Polarity is an uneven sharing of electrons in a covalent compound based on electronegativity difference
• It results in partial charges indicated by the symbol: ∂
• Polarity can determine the shape of the molecule as well as how it behaves

Electronegativity and Bond Types

• 0-0.4 difference is Non Polar Covalent, i.e. H2
• 0.5-1.9 difference is POLAR, i.e. HCl / H2O
• 2.0 + difference is lonic i.e. NaCl
• Technically, 1.7-2.0 is tricky because if a metal is involved, it's ionic, but if the element pairing involves 2 nonmetals, it's polar covalent

Bond Polarity

• Written using symbols Hδ+Clδ-
• The positive and minus signs (with the lower case delta: δ+ and δ- ) denote partial charges, the alternative way shows an arrow pointing towards the direction polarity: +→
• The arrow will always point to the mores electronegative

Using Electronegativity to Determine Polarity

• Difference between H-H is 0.0, which is pure covalent
• The difference between C - H is 0.4, which is (weakly) polar covalent
• The difference between O - H is 1.4, which is polar covalent
• The difference between H-F is 1.9, which is polar covalent
• The difference between S - O is 1.0, which is polar covalent
• The difference between C - O is 1.0, which is polar covalent
• The difference between Al - C is 1.0, which is polar covalent
• The difference between Na - Cl is 2.1, which is ionicThe difference between Mg - O is 2.3, which is ionic
• The difference between Mg - C is 1.3, which is polar covalent
• A bond can be polar, however, that DOES NOT necessarily mean the whole molecule is polar
• CO2 is linear, so the C-O polar bonds cancel and the MOLECULE is non-polar
• H2O is bent with lone pairs, making the MOLECULE polar

Naming/Formulas for Covalent Compounds

• Naming binary covalent compounds requires prefixes
• The prefix tells how many atoms of each element are in the compound
• You do NOT have to put Mono on the first element, but you do on the second
• 1: Mono, 2: Di, 3: Tri, 4: Tetra, 5: Penta, 6: Hexa, 7: Hepta, 8: Octa, 9: Nona, 10: Deca
• CO = Carbon Monoxide
• CO2 = Carbon Dioxide
• P2O3 = Diphosphorus Trioxide
• BF3 = Boron Trifluoride
• PF5 = Phosphorus Pentafluoride

Network solids

• Network solids are special covalent compounds where all atoms are covalently bonded to each other
• Breaking the solid requires breaking covalent bonds, not just intermolecular forces, for example, diamonds and silicon carbide

Test your knowledge of Lewis structures, octet rule exceptions, and covalent bonding. Questions cover valence electrons, coordinate covalent bonds, and differences between covalent and ionic compounds. Includes identifying elements likely to form electron-deficient molecules and molecules exceeding the octet rule.