Chemical Bonding

Chemical Bonds Compound are formed from chemically bound atoms or ions Bonding only involves the valence electrons

Chemical Bonds Defn – force holding two atoms together How are they formed? (2 ways) a) atoms gain/lose valence electrons b) share valence electrons Why does bonding occur? Stability – achieve octet rule

Electron Dot Structure Shows valence electrons around atomic symbol hydrogen nitrogen chlorine (group 5) (group 7) (group 1) H N Cl

Types of Chemical Bonds 3 Types – covalent bond – ionic bond – metallic bond

Covalent Bond Defn – two atoms share one pair of electrons ABA B Electrons shared

Covalent Bonds Where are these bonds found? - molecules (molecular compounds) - polyatomic ions

Ionic Bond Defn – force holding cations and anions together ABA+A+ B-B- Ionic bond

Ionic Bond Where are these bonds found? Ionic Compounds

Metallic Bonding Defn – attraction of metallic cations Occurs only in metals

Covalent Bonding What’s going on? Molecule – formed when 2 or more atoms bond covalently Sharing of electrons

Two Types of Covalent Bonds i) nonpolar covalent – equal sharing of e - ii) polar covalent – UNequal sharing of e -

Nonpolar vs. Polar NONPOLARPOLAR

Nonpolar vs. Polar

Single Bond Defn – one pair (2) of e - shared Lewis Structures – represents how atoms in molecules are arranged –atoms MUST obey octet rule (except hydrogen)

Lewis Structures bonded electrons – occur between bonded atoms A B AB single bond or

Lewis Structures Unshared or Lone Pairs – electron pairs NOT involved in bonding AB A B lone pairs

Lewis Structures Examples H 2 O H H (8 valence e - or 4 pairs) O O H H O H H

Lewis Structures Examples NHF 2 (20 v.e. or 10 pairs) N F F H F N F H N F F H

Multiple Covalent Bonds Double Bond – two pairs (4) e - shared A B AB O O2O2 O (12 v.e. = 6 pairs) O O O O O O

Multiple Covalent Bond Triple Bond – three pairs (6) e - shared A B AB N2N2 (10 v.e. = 5 pairs) N N N N N N N N

Comparing single, double, and triple bonds Bond Strength Bond Length Triple > Double > Single Single > Double > Triple The shorter the bond, the stronger it is

Polyatomic Ions Defn – CHARGED group of atoms covalently bonded - ex: SO 4 2-, NH 4 1+, NO 3 1-

Polyatomic Ions SO 4 2- (32 v.e. = 16 pairs) O O O O S 2- O O O O S 2-

Polyatomic Ions NH 4 1+ (8 v.e. = 4 pairs) H H HHN 1+ H H HH N 1+

VSEPR Valence Shell Electron Pair Repulsion Defn – determines the shape of molecule Valence electron pairs repel each other; try to stay far away as possible

# atoms bonded to central atom # lone pairs on central atom shape 40tetrahedral

Tetrahedral

# atoms bonded to central atom shape tetrahedral trigonal planar # lone pairs on central atom

Trigonal Planar

shape tetrahedral trigonal planar # atoms bonded to central atom # lone pairs on central atom trigonal pyramidal

Trigonal Pyramidal

shape tetrahedral trigonal planar linear # atoms bonded to central atom # lone pairs on central atom trigonal pyramidal

Linear 2 atoms only

shape tetrahedral trigonal planar trigonal pyramidal linear bent # atoms bonded to central atom # lone pairs on central atom

Bent

Predict these shapes N HH H NH 3 trigonal pyramidal H2OH2OCO 2 O HH bent C OO linear

Ionic Bonding What’s going on? If I gave you a compound, how can you tell if it is ionic or not? combo of metal + nonmetal giving/taking of valence electrons

Formation of Ionic Bonds NaCl NaCl + Na 1+ + Cl 1- 2s 2 2p 6 3s 1 3s 2 3p 5 2s 2 2p 6 3s 2 3p 6 8 v.e.

Formation of Ionic Bonds CaBr 2 CaBr + Br Ca 2+ + Br 1- Br 1-

Using electronegativity to determine bond type Recall electronegativity: how much an atom wants electrons Each atom is assigned a number between to determine electronegativity strength

We know 3 types of bonds: - nonpolar covalent - polar covalent - ionic To determine bond type, subtract electronegativity values and see scale Using electronegativity to determine bond type

Scale polar covalent nonpolar covalent ionic

H-H H-O 2.2 – 2.2 = – 2.2 = 1.24 nonpolar polar H-C 2.55 – 2.2 = 0.35 polar K-F 3.98 – 0.82 = 3.16 ionic

Metallic Bonding Defn – bond formed from attraction between positive nuclei and delocalized electrons –holds metals together Delocalized Electrons – electrons detached from parent atom –“lost electron away from home”

Electron Sea Model Defn – electrons move freely within other molecular orbitals

Properties of Metals Electron sea model gives metals certain physical properties 1)Shiny – due to photoelectric effect 2)Conduct electricity and heat – electrons move easily from one place to another 3)Malleable (pound into sheets) 4)Ductile (put into wires)

Why malleable and ductile? atoms can also move from one place to another and still remain in contact with and bonded to the other atoms and electrons around them shifted atoms Shape #1 Shape #2

Dipole Moment defn – imbalance of electron density in a covalent bond –Due to electronegativity of atoms  - ( partial negative) = signifies more EN atom  + (partial positive) = signifies less EN atom = shows direction of dipole moment

Examples H = 2.2 C = 2.6 N = 3.0 Cl = 3.2 O = 3.4 F = 4.0 HO ++ -- ClC -- ++ NH -- ++ CF ++ --

Intermolecular Forces Defn – attractive forces between 2 molecules; weak compared to bonds

Intermolecular Forces London Dispersion Forces – very weak, exist in all molecules due to temporary shifts in the electron cloud

Electrons evenly distributed Temporary dipole London force

Intermolecular Forces Dipole-Dipole – attraction between oppositely charged regions of polar molecules ++ -- ++ -- ++ --

Intermolecular Forces Hydrogen Bonding – special type of dipole-dipole occurring btwn molecules that have a H bonded to either a N,O, or F of another molecule; strongest force - Water is prime example O H H O ++ ++ --

O H H O O H H O O H H O O H H O O H H O hydrogen bond ++ ++ -- ++ ++ --

Strength Ranking Hydrogen > dipole-dipole > London