Chemical Bonding

Chapter 6

General Chemistry

Valence Electrons• Valence electrons _____________________________________________________________________________

– The s and p electrons in the outer energy level– Fluorine [He] 2s2 2p5 = 7 valence e-

• The electrons responsible for the chemical properties of atoms are those in the outer energy level

• ________________those in the energy levels below the outer energy level

2s2 2p5

Lewis Dot (Electron Dot) Diagrams

• Lewis Dot (electron dot) diagrams ___________________________________

X

Lewis Dot Diagrams of Selected Elements

Element Valence e- Lewis dot diagram

Sodium 1 Na •

Magnesium

Phosphorus

Chlorine

Octet Rule

• The octet rule: _______________________

_____________________________________

• Want to achieve the e- configuration of a noble gas

• Why named “octet”?

• Exceptions?

Chemical Bonding

• When atoms bond, the valence electrons are redistributed to make the atom more stable

• _______________results from the electrical attraction between large numbers of cations and anions

• Covalent bonding: results from the _______ of electrons between two atoms

Ionic Bonding

Remember Ions ?

• Ions: charged atoms

• Cations: positively charged atoms – _________________________________ to

create a noble gas configuration (cations)

• Anions: negatively charged atoms– _________________________________to

create a noble gas configuration (anions)

Ionic Bonds• Formed between ________________ atoms• Anions and cations are held together by

opposite charges• The bond is formed through the transfer of

electrons • Ionic compounds are called ___________• Simplest ratio is called the ____________

– Example: Na+ will bond with Cl- to make sodium chloride, NaCl

Electronegativity

• Electronegativity: reflects an atom’s ability to attract electrons in a chemical bond

• Metals generally have ___ electronegativity

• Nonmetals generally have high electronegativity

How Determine if Ionic?

• Ionic bonds form between 2 atoms with difference in electronegativity of 2.0 or greater

Properties of Ionic Compounds

• Conduct electricity in aqueous form– are __________________

• High melting and boiling points

• Usually solids at room temperature

• Have ________________

• Example: sodium chloride (table salt)

Lattice Energy

• The strength of an ionic bond compared to another ionic bond is determined by the lattice energy

• Lattice energy _______________________

_____________________________________

• Examples: – NaCl -787.5 kJ/mol (weaker bond)– MgO -3760 kJ/mol (stronger bond)

Crystalline structure

Ionic Bonding Lewis Dot Diagrams

Na Cl

Ionic Bonding Lewis Dot Diagrams

Na+ Cl-

Ionic Bonding Lewis Dot Diagrams

Ca2+ P 3-

Ca2+P

3-

Ca2+

Ionic Bonding Lewis Dot Diagrams

= Ca3P2Formula Unit

Metallic Bonding

Metallic Bonds

• Metallic bonding is the bonding that results from the attraction between ______________________________________________________________– Bond between two metal atoms

+ + + ++ + + +

+ + + +

Sea of Electrons• Metals hold on to their valence electrons very weakly.

• Think of them as positive ions (cations) floating in a ______________

• Electrons are free to move through the solid.

• Metals conduct electricity.

Covalent Bonding

Covalent Bonds

• Two nonmetals_____ electrons to achieve full octet of electrons

• By sharing, both atoms get to count the electrons toward a noble gas configuration.

• Form molecules - _____________________

___________________

Examples of Molecules

How determine if covalent?

• Covalent bonds form between 2 atoms with difference in electronegativity of ________________

Properties of Covalent Compounds

• Do not conduct electricity in aqueous solution– Are __________________

• Relatively low melting and boiling points

• Can be gasses, liquids or solids @ room temp– Examples: sugar, wax, carbon dioxide

Comparison of MP, BP in Ionic and Covalent Compounds

Bond Energy

• The strength of an covalent bond compared to another covalent bond is determined by the bond energy

• Bond Energy: _______________________

________________________________________________________________________– Stronger covalent bonds have a higher bond

energy

Bond Energy and Bond Length

Bond Length

• Bond Length: ________________________

_____________________________________

• The longer the bond, the ___________ the bond energy (the ___________the bond)

• The shorter the bond, the ___________ the bond energy (the ___________ the bond)

Types of Covalent Bonds

• Single covalent

• Double covalent

• Triple covalent

• Share __ e- (one pair)

• Share __e- (two pairs)

• Share __e- (three pairs)

Covalent bonding

• Fluorine has seven valence electrons

F

Covalent bonding• Fluorine has seven valence

electrons• A second F atom also has seven• By sharing electrons…

F F

Covalent bonding• Fluorine has seven valence electrons• A second atom also has seven• By sharing electrons…• …both end with full orbitals

F F8 Valence electrons

Bonding and Nonbonding Electrons

• _______________________ are involved in a chemical bond

• ___________________________________are not involved in bonding and belong exclusively to one atom

Nonbonding electronsBonding

electrons

Diatomic Elements

• __________ pure elements that exist as pairs in nature

• Are covalently bonded– H2 N2 O2 F2 Cl2 Br2 I2

• Ways to remember:– BrINClHOF – H, N, O, Halogens

Polarity

Bond Polarity

• Atoms of elements do not always share electrons equally

• ___________________unequal sharing of electrons (dif electroneg 0.5 – 1.9)

• ___________________equal sharing of electrons (dif electroneg 0.0-0.4)

Bond Polarity• When two different atoms bond

covalently, there is an unequal sharing– the more electronegative atom will have a

_____________ attraction and will acquire a slightly ________charge

– called a polar covalent bond or simply polar bond.

Bond Polarity• Refer to Periodic Table values of

Electronegativity• Consider HCl

H = electronegativity of 2.1Cl = electronegativity of 3.0– the bond is _________– the chlorine acquires a slight negative

charge, and the hydrogen a slight positive charge

Bond Polarity• Only partial charges, much less than a

true 1+ or 1- as in ionic bond

• Written as:

H Cl• the positive and minus signs (with the

lower case delta ) denote partial charges.

Bond Polarity• Can also be shown:

• the arrow points to the more electronegative atom.

H Cl

Calculate Polarity of Bond

Difference in Electronegativity

Type of Bond

0.0-0.4

0.5-1.9

2.0 and greater

Geometry

VSEPR Theory

• Valence Shell Electron Pair Repulsion Theory

• ______________________________________________________________________

• Allows chemists to predict shapes of simple molecules

Predict shape and polarity

• Shape affects polarity of molecule

• Even though atoms may have dif electroneg > 0.5, the shape may cancel out the effects

• Example: CO2

Intermolecular Forces

Intermolecular Forces

• Polar molecules, such as water (H2O) attract other polar molecules.

• The forces of attraction between molecules are known as _______________________– Stronger IM Forces result in ________ MP, BP

(solids and liquids)– Weaker IM Force result in __________MP, BP

(liquids, gases)

Types of Intermolecular Forces

• Dipole-dipole forces

• Hydrogen bonding

• London Dispersion Forces

Dipole-Dipole

• Dipole-dipole forces– Attractions between ________________– Example: BrF

Hydrogen Bonding

• Hydrogen bonding– Is a special type of dipole-dipole attraction– Not really a “bond” but a stronger attraction– ______________________________________

______________________________________

– Example: H2O NH3

London Dispersion Forces

• London Dispersion Forces– Generally only significant IM force in nonpolar

molecules– Attraction between large massed atoms (that

have lots of electrons)– ______________________________________

______________________________________