Transcript pps
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
Lewis Structures
•
Electron Dot Diagrams
– show valence e as dots – distribute dots like arrows in an orbital diagram – 4 sides = 1 s-orbital, 3 p-orbitals – EX: oxygen
X 2s 2p O
Lewis Structures
•
Octet Rule
– Most atoms form bonds in order to obtain 8 valence e – Full energy level stability ~ Noble Gases
Ne
Electron Dot Structure
• Shows valence electrons around atomic symbol hydrogen nitrogen chlorine (group 1) (group 5) (group 7)
H
• •• • • • •• •• ••
Types of Chemical Bonds
• 3 Types – ionic bond –covalent bond – metallic bond
Ionic Bonding
Ionic Bonding Vocabulary
Ionic compounds are referred to as Formula Units.
Compounds are composed to two or more elements. Binary Compound – 2 elements - NaCl Ternary Compound – 3 or more elements – NaHCO 3 Ion – A charged atom Monatomic Ion – 1 atom Na 1+ Polyatomic Ion - 2 or more atoms NO 3 1-
Ionic Bond
• Defn – force holding cations and anions together
A
• •
B A +
• •
B Ionic bond
Cations – positively charged ions Anions – negatively charged ions
Ionic Bond
• Where are these bonds found?
In
Ionic Compounds
Ionic Bonding Properties
Bond Formation
Type of Structure Physical State Melting Point Solubility in Water Electrical Conductivity
Electrons are transferred from metal to nonmetal
Crystal Lattice Solid High Yes Yes – in solutions or liquid
Ionic Bonding
• What’s going on?
giving/taking of valence electrons
• If I gave you a compound, how can you tell if it is ionic or not?
combo of metal + nonmetal
Formation of Ionic Bonds
• NaCl
Na • + 2s 2 2p 6 3s 1 • • • Cl • • •• 3s 2 3p 5 Na 1+ 2s 2 2p 6 + •• • • Cl • • •• 1 3s 2 3p 6 8 v.e.
Formation of Ionic Bonds
• CaBr 2
• Ca • + • • • • Br • • •• • • Br • • •• Ca 2+ + •• • • Br • • •• 1 •• • • Br • • •• 1-
Ionic Formulas
• Don’t show charges in the final formula.
• Overall charge must equal zero .
– If charges cancel, just write symbols.
– If not, use subscripts to balance charges.
• Use parentheses to show more than one polyatomic ion .
• Stock System the ion’s charge.
Roman numerals indicate
Ionic Formula Names
• Write the names of both ions, cation first. • Change ending of monatomic anions to ide .
• Polyatomic ions have special names. • Stock System - Use Roman numerals to show the ion’s charge if more than one is possible. Overall charge must equal zero.
1+ 2+
Common Ion Charges (Oxidation States)
3+ 4+/4 3- 2- 1 0
Ionic Formulas
potassium chloride K + Cl
magnesium nitrate
Mg 2+ NO
copper(II) chloride Cu 2+ Cl
KCl Mg(NO 3 ) 2 CuCl 2
Ionic Formula Names
NaBr sodium bromide
Na 2 CO 3 sodium carbonate
FeCl 3 iron(III) chloride
Metallic Bonding
Metallic Bonding
• Defn – attraction of metallic cations
Occurs only in metals
Metallic Bonding Properties
Bond Formation Type of Structure Physical State Melting Point Solubility in Water Electrical Conductivity Delocalized among metal atoms Electron Sea Solid Very High No Yes
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 Shape #1 shifted atoms Shape #2
Covalent Bonding
Covalent Bond
• Defn – two atoms share one pair of electrons
A
• •
B A
• •
B Electrons shared
Covalent Bonds
• Where are these bonds found?
- molecules (molecular compounds) - polyatomic ions
Covalent Bonding
• What’s going on?
Sharing of electrons • Molecule – formed when 2 or more atoms bond covalently
Covalent Bonding Properties
Bond Formation Type of Structure Physical State Melting Point Solubility in Water Electrical Conductivity Electrons are shared between two nonmetal atoms True molecules Liquid or gas, brittle solids Low Usually not No
Two Types of Covalent Bonds
i) nonpolar covalent – equal sharing of e ii) polar covalent – UN equal sharing of e -
Nonpolar vs. Polar NONPOLAR POLAR
Nonpolar vs. Polar
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 or single bond A B
Lewis Structures
• Unshared or Lone Pairs – electron pairs NOT involved in bonding ••
A
• •
B
•• ••
A B
••
lone pairs
Lewis Structures Examples
• H 2 O (8 valence e or 4 pairs)
H
•
H
• • ••
O
•
H
• • •• •
H
•
H
••
O
••
H
Lewis Structures Examples
• • NHF 2 (20 v.e. or 10 pairs) ••
N
•
H
• • • ••
F
•• • •
F
• • • •• • • •• • • • •• •
H
• ••
F
• • ••
N
••
F
• •
H
Multiple Covalent Bonds
• • • Double Bond – two pairs (4) e shared
A
• • ••
B A B O 2 (12 v.e. = 6 pairs)
•• • ••
O
• • • ••
O
• ••
O
• •• ••
O
•• ••
O
••
O
••
Multiple Covalent Bond
• Triple Bond – three pairs (6) e shared
A
• • ••••
B A B N 2 (10 v.e. = 5 pairs)
• ••
N
• • ••
N
• • ••
N
• ••
N
• ••
N
••
N
• •
N
• • •• •• • •
N
Comparing single, double, and triple bonds
• Bond Strength
Triple > Double > Single
• Bond Length
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
•
2-
• • • •• • • • ••
O S
• • ••
O
•
2-
Polyatomic Ions NH 4 1+ (8 v.e. = 4 pairs) H
• • •• •
H
•
H 1+ H H N H H 1+
Using electronegativity to determine bond type
•
Recall electronegativity: how much an atom wants electrons
•
Each atom is assigned a number between 0-4.0 to determine electronegativity strength
Bond Polarity
• Most bonds are a blend of ionic and covalent characteristics.
• Difference in electronegativity determines bond type.
Lewis Structures
•
Nonpolar Covalent
- no charges •
Polar Covalent
- partial charges
+
-
+
Bond Polarity
•
Electronegativity
– Attraction an atom has for a shared pair of electrons.
– higher e neg atom – lower e neg atom +
Bond Polarity
•
Electronegativity Trend
• Increases up and to the right.
Using electronegativity to determine bond type
•
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 0 Scale nonpolar covalent polar covalent 0.3
1.7
ionic 4.0
Using electronegativity to determine bond type H and Cl C and S Na and F 3.0 – 2.1 = 0.9
polar covalent 2.5 – 2.5 = 0 nonpolar covalent 4.0 – 0.9 = 3.1
ionic
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 H = 2.2
C = 2.6
N = 3.0
Cl = 3.2
O = 3.4
F = 4.0
N
O
+ H
Cl
+ C
+ C
F
Naming Covalent Compounds
•
Prefix System
(binary compounds) 1.
Less e neg comes first. atom 2. Add prefixes to indicate # of atoms. Omit mono prefix on first element.
3. Change the ending of the second element to -ide .
C. Molecular Nomenclature
PREFIX
mono di tri tetra penta hexa hepta octa nona deca-
NUMBER
1 2 3 4 5 6 7 8 9 10
Covalent Compound Names
CCl 4 carbon tetrachloride
N 2 O
dinitrogen monoxide SF 6 sulfur hexafluoride
Covalent Formulas
arsenic trichloride AsCl 3
dinitrogen pentoxide N 2 O 5
tetraphosphorus decoxide P 4 O 10
C. Molecular Nomenclature
H
•
The Seven Diatomic Elements Br 2 I 2 N 2 Cl 2 H 2 O 2 F 2 N O F Cl Br I
Intermolecular Forces
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 London force Temporary dipole
Intermolecular Forces
• Hydrogen Bonding – strong bond between H and N,O, or F of another molecule -
Water is prime example
-
+ H H
+
-
+ H H H
+
H
+ H H H
+ hydrogen bond H H H
Strength Ranking Hydrogen > dipole-dipole > London
VSEPR
•
V
alence
S
hell
E
lectron
P
air
R
epulsion • Defn – determines the shape of molecule •
Electron pairs try to stay far away as possible
# atoms bonded to central atom 4 # lone pairs 0 shape tetrahedral
Tetrahedral
# atoms bonded to central atom 4 3 # lone pairs 0 1 shape tetrahedral trigonal pyramidal
Trigonal Pyramidal
# atoms bonded to central atom 4 3 2 # lone pairs 0 1 2 shape tetrahedral trigonal pyramidal bent
Bent
# atoms bonded to central atom 4 3 2 3 # lone pairs 0 1 2 0 shape tetrahedral trigonal pyramidal bent trigonal planar
Trigonal Planar
# atoms bonded to central atom 4 3 2 3 2 # lone pairs 0 1 2 0 0 shape tetrahedral trigonal pyramidal bent trigonal planar linear
Linear