Unit 2: Bonding & Lewis Structures Flashcards
What is a Chemical bond ?
the attraction between positive and negative ions, or two nonmetal atoms.
Ionic Bonds
are due to the electrostatic attractions between oppositely charged ions.
- solids with mp > 400C;
- soluble in polar solvent;
- conductive,
- high ΔEn
Covalent Bonds
molecules are formed by the sharing of electrons between atoms
- covalent – gases, liquids, solids rt, mp usually <300 C;
- many insoluble in polar solvent,
- nonconductive,
- low ΔEn diff
Lewis Formula
Chemical symbol of element surrounded by dots representing valence electrons
Lewis symbol: inner electrons/nucleus as symbol; valence as dots
Primary contribution = valence electrons
G.N. Lewis – express valence electrons
Octet Rule
Atoms often gain, lose or share electrons to achieve the same number of electrons as the noble gas closest to them in the periodic table.
Noble gases have very stable electron arrangements, because all noble gases except He have eight valence electrons, many atoms undergoing reactions end up with eight valence electrons. This observation led to octet rule.
Octet Rule: Atoms tend to gain, lose or share electrons until they are surrounded by 8 valence electrons
Ionic Bonding
Ionic substances generally result from the interaction of metals on the left side of the periodic table with nonmetals on the right side
For Example:
One element oxidized, another reduced – electron transfers
Ex: NaCl
2Na (s) + Cl2 (g) à 2 NaCl (s)
Ionic Bonding (continue example)
Ionic Bonding
Ionic Bonding
Electron configuration for Zn2+ ion:
Electron configuration for Co2+ and Co3+ ion:
Covalent Bonding
A chemical bond formed by sharing a pair of electrons is a covalent bond.
Represented by a line between atoms in a Lewis structure: X-X
Ionic compound – strong intermolecular, weak intramolecular
Covalent – weak intermolecular, strong intramolecular
Covalent bond formation follows energy curve w/ minimum distance at bond distance
Covalent Bonding
Covalent Bonds & Lewis Structures
if a polyatomic ion is formed, the charge is included as in ionic compounds
NH4+
Writing Lewis Formulas
- Obey octet rule
- Differentiate bonding and nonbonding electrons
Determine bonding electrons:
S = N - A
S = shared
N = needed
A = available
- For F2 as an example
- A = 2 F = 2x7 available;
- N = 2 F = 2x 8 needed
- S = 16 - 14 = 2 shared
- n2 electrons = 1 bond
Guide to Dot Structures
What is a Oxidation State (number)
a means devised for keeping track of electrons being gained or lost by a substance.
Track electronic charges
Single atom ions = actual charge
Polyatomics spread over ion; assign # to each atom
General rules but no absolutes
Oxidation Number Rules
(1-4)
- The oxidation number of an element in its free (uncombined) state is zero
- The oxidation number of a monatomic (one-atom) ion is the same as the charge on the ion.
- The sum of all oxidation numbers in a neutral compound is zero.
- The sum of all oxidation numbers in a polyatomic (many-atom) ion is equal to the charge on the ion.
Oxidation Number Rules
(5-7)
- The oxidation number of fluorine is always –1.
- The oxidation number of oxygen in a compound is usually –2, except in peroxides (- 1), superoxides (- ½), or combined with fluorine (+1 or +2)
- The oxidation state of hydrogen in a compound is usually +1, except as a binary compound with a metal (-1)
Oxidation Number Rule 8
- Periodic position contributes to oxidation state:
a. Group IA = 1+
b. Group IIA = 2+
c. Group IIIA = 3+ (some exceptions)
d. Group VA = 3- w/ metal, H, NH4+; apply 3 & 4 to element to the right of group V
e. Group VIA = 2- w/ metal, H, NH4+; apply 3 & 4 to element to O or halogen in lower n
f. Group VIIA = 1- in binary compounds; apply 3 & 4 when combined w/ O or halogen in lower n
If En are equal, negative to element closest to F
Determine ox state of N in N2O4, NH3, HNO3, NO3-
Resonance
Resonance – situations where a single Lewis structure does not present the actual configuration for a molecule, and it is capable of representation by multiple valid Lewis structures.
Resonance structure – different electron arrangement
Resonance hybrid – real molecule
Ex: CO32-
Formal Charges
Rules for assigning formal charge for main group elements:
Formal Charge – the charge an atom would have if each bonding electron pair in the molecule were shared equally between its two atoms.
Rules for assigning formal charge for main group elements
- FC on an atom is = group number – [number bonds + number unshared electrons].
- If # bonds = group number, F.C. = 0
- a. in molecule, sum charge = 0
b. in polyatomic, sum charge = charge on ion
Keep in mind for Formal Charges
Keep in mind for formal charges:
a. most likely formula charges minimize charge – closest to 0
b. put negative charges mostly on more electronegative atom
c. If adjacent atoms have same sign (+ or -) usually incorrect
EX: NOCl
Octet Exceptions
Some instances octet rule fails: must modify S = N - A formula for these:
A. covalent compounds of Be – use 4 for N
B. group 3A elements – use N = 6
C. Some compounds w/ unpaired electrons NO and NO2
D. Expanded Valence – > 8 electrons; usually S and P
Octet Exceptions
Expanded Valence
Write Lewis symbols for Na, F, Si, Ar, S
Write Lewis structures for NaF, SiF4, BeCl2, BCl3, PF5,
Write resonance structures for NO3-
