Chemical Bonding Chapter 6. Types of Chemical Bonds Chemical Bond: mutual electrical attraction b/...

Chemical Bonding

Chapter 6

Types of Chemical Bonds

Chemical Bond: mutual electrical attraction b/ the nuclei and valence e- of different atoms

Atoms make bonds b/c they become more stable

Types of Chemical Bonds

Ionic Bonding: Results from the electrical attraction

between cations and anions Atoms completely give up e-s to other

atoms

Types of Chemical Bonds

Covalent Bonding: Results from sharing of e- pairs

between two atoms Shared e- are “owned” equally by two

atoms

Determining the Type of Bond Bonds fall somewhere b/ purely Ionic

and purely Covalent Depends on electronegativity – how

much is the atom pulling the e-s Calculate the difference in

electronegativities

Determining the Type of Bond

Types: Ionic: electronegativity dif: 1.7 – 3.2 Polar-Covalent: dif: 0.4 – 1.6 Nonpolar-Covalent: dif: 0 – 0.3

Determining the Type of Bond

Ionic: One atom is so electronegative it strips

the other atom of electrons making cation and anion

NaCl

Determining the Type of Bond

Polar-Covalent: Bond where atoms have unequal

attraction for shared e-s More electronegative atom has

stronger attraction for e-s HCl

Determining the Type of Bond

Nonpolar-Covalent: Bond where atoms have equal sharing

of e-s, balanced distribution of charge Bonds b/ two atoms of the same

element, F2

C-H

Determining the Type of Bond

Determine the type of bond b/ these elements and sulfur, which is more electronegative element?

H, Cs, Cl

Covalent Bonding

Covalent Bonding

Molecule: neutral group of atoms that are held together by covalent bonds.

Molecular = “covalent”

Covalent Bonding

Covalent Bonding

Chemical formula: Ionic or Covalent NaCl, MgCl2, H2O

Molecular formula: only for molecules (covalent)

Covalent Bonding

Forming Covalent Bonds

Share e-s to get noble gas configuration

Octet Rule

Def: atoms gain, lose or share e-s to have octet (8) of electrons in outer energy level

H – exception, only needs 2 e-s Ex: F2

Exceptions to Octet

B – happy with 6 e-s in outer level Other elements can have more than

8 valence e-s – PF5, SF6 – d orbitals invovled in bonding

Electron-Dot Notation

Use element symbol and dots to indicate valence e-s.

Period 2:

1A 2A 3A 4A 5A 6A 7A8A

Li Be B C N O FNe

Lewis Structures

Def: Use electron-dot notation for molecules

H2

F2 – shared pair with dash (-) lone pair – unshared pair

Lewis Structures

Single bond: covalent bond where one pair of e- being shared b/ two atoms

H-H and HCl :

Lewis Structures

Draw Lewis Structure:1. Start with e-dot diagram of each2. Atom with most bonding sites in

middle3. Circle unpaired e-s to make bonds4. Replace circles with dashes

NH3

H2S

Multiple Covalent Bonds

C, N, and O can share more than 1 pair of e-s

Double bond: two pairs, 4 e-s, being shared

C2H4

Triple Bond: three pairs, 6 e-s, being shared

N2

Bond Lengths Bond length: average distance b/

two bonded atoms Forming bonds – atoms release energy Same amount of energy needed to

break bond bond energy (kJ/mol) Lengths of multiple bonds?

More bonds – shorter – more energy to break

p.187

Resonance Structures

Molecules can’t be shown with one Lewis structure

Ex: O3

Lewis Structure

CH2O

CH3Br

C2HCl

SO3

Recap Ch. 6

Bonds: Ionic and Covalent Ionic, Polar-Covalent, and Nonpolar-

Covalent Drawing Lewis Structures

C2H2

Type of bonds?

Ionic Bonding

Ionic Compounds

Def: (+) and (-) ions that combine so charges balance out

Crystalline solids Formula Unit: simplest unit of

ionic compound where charges are balanced NaCl: Na+ Cl-

Video (68)

Forming Ionic Compounds

NaCl – use electron dot diagrams Compound with Ca and F:

Characteristics of Ionic Bonds

Ions in crystal lattice are more stable – lower potential energy

Lattice energy: energy released when 1 mole of gaseous ions form a lattice

More negative energy = more energy released = lower potential energy = more stable = stronger bonds

Ionic vs. Covalent

Ionic stronger bonds b/ formula units than b/

molecules in covalent compounds HIGHER melting and boiling points hard but brittle conduct electricity in molten or

dissolved state

Ionic vs. Covalent

Covalent Weak bonds b/ molecules Most compounds are gases at room

temp. LOW boiling and melting points

Polyatomic Ions

Def: charged group of covalently bonded atoms Result from excess or lack of electrons in

bonding

N HH

HH

+

N OO

O

Ammonium ion Nitrate ion

S OO

O

Sulfate ion

2

O

Metallic Bonding

Excellent conductors in solid state – due to highly mobile valence electrons

Filled outermost sublevel is s Vacant p and d orbitals overlap -

valence e-s are delocalized, do not belong to any one atom but move freely

Metallic Properties

High electrical and thermal conductivity

Malleability – hammered into sheets Ductility – drawn into wires

Molecular Geometry

3-D arrangement of molecules

VSEPR Theory

Valence-shell, electron-pair repulsion

Def: repulsion b/ valence e- pairs around atom causes them to be as far apart as possible

Shapes

NO lone pairs on CENTRAL atom

Symmetrical Linear Trigonal-Planar Tetrahedral Trigonal-

bipyramidal Octahedral

WITH lone pairs on CENTRAL atom

Non-symmetrical Trigonal-pyramidal Bent (angular)

Shapes – NO lone pairs on central atom

1. Linear (AB2):

• A – central atom B-bonded atoms- 3 atom molecules CO2

- 2 atom molecules, O2, HCl, etc.

- bond angles: 180o

Shapes – NO lone pairs on central atom

2. Trigonal Planar (AB3):

- BCl3- bond angles: 120o

Shapes – NO lone pairs on central atom

3. Tetrahedral (AB4):

- CCl4- bond angles: 109.5o

Shapes – NO lone pairs on central atom

4. Trigonal-bipyramidal (AB5):

- PCl5- bond angles: 120o and 90o

Shapes – NO lone pairs on central atom

5. Octahedral (AB6):

- SF6

- bond angles: 90o

Shapes – WITH lone pairs on central atom

6. Trigonal-Pyramidal (AB3E):

• A – central atom B – bonded atoms E – lone pair- NH3

- triangular sides - bond angles: 107o

Shapes – WITH lone pairs on central atom

7. Bent or Angular (AB2E2):

- H2O

- bond angles: 105o

Molecular Polarity

Polarity of each bond Molecular polarity

Molecular Polarity

1. Has ALL bonds NONPOLAR nonpolar molecule

2. Has bonds nonpolar AND polar polar molecule

3. Has ALL bonds POLAR depends on shape

Symmetrical shape (linear - octahedral) NONPOLAR

Non-symmetrical shape (bent & trigonal pyramidal) POLAR

Molecular Polarity Examples

CCl4 PH3

CBr3H

Intermolecular Forces

Intermolecular Forces

“between molecule” forces Generally weaker than bonds b/

atoms Boiling point – good to measure

intermolecular forces

Dipole-Dipole Forces

Dipole- equal but opposite charges separated by a short distance

Video 124H Cl

Dipole-Dipole Forces

Induced Dipole: polar molecule makes a dipole on a nonpolar molecule

Ex: O2 dissolved in H2O Weaker than regular dipole forces

Hydrogen Bonding

Type of dipole-dipole force Def: H-atom bonded to highly e-neg

atom is attracted to lone pair of the e-neg atom in nearby molecule

Ex: HF, H2O, NH3

Hydrogen Bonding

London Dispersion Forces

Def: constant motion of e-s and creation instantaneous dipoles

Video 133