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? Atoms gain, lose, or 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+
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 Using electronegativity to determine bond type
H and Cl3.0 – 2.1= 0.9 polar covalent C and S2.5 – 2.5= 0 nonpolar covalent Na and F4.0 – 0.9= 3.1 ionic Using electronegativity to determine bond type
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
Intermolecular Forces Dipole-Dipole – attraction between oppositely charged polar molecules ++ -- ++ -- ++ --
Intermolecular Forces London Dispersion Forces – very weak, very brief dipole moment created in nonpolar molecules
Electrons evenly distributed Temporary dipole London force
Intermolecular Forces Hydrogen Bonding – strong bond between H and N,O, or F of another molecule - 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
VSEPR Valence Shell Electron Pair Repulsion Defn – determines the shape of molecule Electron pairs try to stay far away as possible
# atoms bonded to central atom # lone pairs shape 40tetrahedral
Tetrahedral
# atoms bonded to central atom # lone pairs shape tetrahedral trigonal pyramidal
Trigonal Pyramidal
# atoms bonded to central atom # lone pairs shape tetrahedral trigonal pyramidal bent
Bent
# atoms bonded to central atom # lone pairs shape tetrahedral trigonal pyramidal bent trigonal planar
Trigonal Planar
# atoms bonded to central atom # lone pairs shape tetrahedral trigonal pyramidal bent trigonal planar linear
Linear