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Chapter 8 Covalent Bonding. General Chemistry I T.ARA. Chemical Bonding. Now that we know something about electron configurations, we can take a closer look at the ways atoms form bonds. There are two main types of chemical bonds: Ionic Bonds Covalent Bonds. A. Ionic Bonding.

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Chapter 8 covalent bonding

Chapter 8Covalent Bonding

General Chemistry I


Chemical bonding
Chemical Bonding

  • Now that we know something about electron configurations, we can take a closer look at the ways atoms form bonds.

  • There are two main types of chemical bonds:

    • Ionic Bonds

    • Covalent Bonds

A ionic bonding
A. Ionic Bonding

  • Ionic bonding involves the transfer of valence electrons from one atom (usually a metal) to another atom (a nonmetal) such that each atom gains a noble gas configuration.

  • The ionic bond is the electrostatic attraction between the cation (+) and the anion (-) that result from the electron transfer.

  • The two bonded atoms do not share electrons.

A ionic bonding sodium chloride
A. Ionic Bonding: Sodium Chloride

  • A sodium atom will readily lose a valence electron:

    Na [Ne]3s1 Na+ [Ne] + e-

  • A chlorine atom will readily acceptan electron:

    Cl [Ne]3s23p5 + e- Cl- [Ar]

  • The sodium cation & the chloride anion are held together by an electrostatic (coulombic) attraction– opposite charges attract.

    Na+ + Cl- NaCl (ionic bond)

A ionic bonding sodium chloride1
A. Ionic Bonding: Sodium Chloride

Using Lewis symbols:

1 polyatomic ions
1. Polyatomic Ions

  • Polyatomic Ion:a group of covalently bonded atoms with an overall positive or negative charge

  • The atoms within

    a polyatomic ion are

    covalently bonded

    together, but

    polyatomic ions form

    ionic bonds with other


eg. NO3- is a polyatomic anion.

B covalent bonding
B. Covalent Bonding

  • In a covalent bondthe electrons are the “glue” that holds the atoms together.

  • A covalent bond is formed when two atoms share one or more pairs of electrons.

  • Each atom will form enough covalent bonds to achieve a noble gas configuration.

G.N. Lewis

B covalent bonding hydrogen
B. Covalent Bonding: Hydrogen

H• + H•  H : H

  • The two electrons are shared evenly between the two hydrogen atoms.

  • It is as if each atom has two electrons – the noble gas configuration of [He].

    H : H

B covalent bonding hydrogen1
B. Covalent Bonding: Hydrogen

  • The actual electron probability cloud(2) for the two electrons in the H-H bond looks like this.

1 lewis structures
1. Lewis Structures

  • A Lewis structure is a way of drawing a molecule that shows all valence electrons as dots or lines that represent covalent bonds.

    eg. The Lewis structure for H2 can be drawn in two ways:

    H : H or H–H

    a) A single line represents two covalently shared electrons – also known as a single bond.

1 lewis structures1
1. Lewis Structures

  • The Lewis Structure for F2:

  • Two electrons are shared between the two F atoms (one single covalent bond).

  • Each F atom also has three unshared electron pairs. These non-bonding electron pairs are called lone pairs.

2 the octet rule
2. The Octet Rule

  • Note that by sharing electrons, it is as if each F atom has eight electrons - the noble gas configuration of [Ne].

  • The Octet Rule:Main group elements with more than two valence electrons gain, lose, or share electrons to achieve a noble gas configuration characterized by eight valence electrons.

Draw a Lewis structure for water (H2O) that obeys the octet rule.

Draw a Lewis structure for ammonia (NH3) that obeys the octet rule.

Draw a Lewis structure for methane (CH4) that obeys the octet rule.

2 the octet rule1
2. The Octet Rule

  • As the previous examples illustrated, the number of covalent bonds an atom must form to achieve an octet is equal to eight minus it group number.

A multiple covalent bonds
a) Multiple Covalent Bonds

  • It is not always possible for atoms to gain a full octet by sharing single electron pairs with other atoms.

  • In other words, it is not always possible to construct a valid Lewis structure using only single bonds.

    eg. N2

In this structure, each nitrogen atom would have only six valence electrons – two short of an octet.

A multiple covalent bonds1
a) Multiple Covalent Bonds

  • Two atoms can share more than one electron pair to gain a full octet.

    • Double Bond: when 2 electron pairs (4 electrons) are shared between 2 atoms

    • Triple Bond:when 3 electron pairs (6 electrons) are shared between two atoms

  • Double & triple bonds are referred to as multiple covalent bonds.

A multiple covalent bonds2
a) Multiple Covalent Bonds

Ethane (C2H6), Ethylene (C2H4) & Acetylene (C2H2)

A multiple covalent bonds3
a) Multiple Covalent Bonds

4 total bonds

3 total bonds

2 total bonds

1 bond

When necessary, atoms will form any combination of single, double and triple covalent bonds to gain a full octet.

Draw a valid Lewis structure for the cyanide ion (CN-).

Lewis Structures of Ions: Add one additional valence electron for every negative charge & subtract one valence electron for every positive charge.

B formal charge
b) Formal Charge

  • As you just saw, even when the octet rule is obeyed, some compounds have an overall charge.

  • This means that the compound contains one or more charged atoms.

    Formal charge: the charge a bonding atom would have if its bonding electrons were shared equally

    Formal Charge = Atomic Group #

    – # lone pair electrons

    – ½ (# bonding electrons)

Draw all possible Lewis structures for N2O (O-N-N) & assign formal charges.

Draw all possible Lewis structures for N2O (O-N-N) & assign formal charges.

3 resonance structures
3. Resonance Structures

  • When more than one Lewis structure can be drawn for a molecule, the structures are called resonance structures.

    • Each resonance structure contributes to the overall structure of the molecule.

    • Individual resonance structures DO NOT ACTUALLY EXIST– we use resonance structures conceptually to help us understand molecular structures!!!

  • The actual structure is a weighted average of the resonance structures – called a resonance hybrid.

3. Resonance Structures

eg. Resonance Structures of Ozone

  • Each bond is in between a single & a double bond.

  • Each terminal O has a partial negative charge.

3 resonance structures1
3. Resonance Structures

  • Resonance structures are connected by double-headed arrows ( ).

  • Each resonance structure must have the same overall charge as the molecule.

  • Lower energy resonance structures contribute more to the overall structure of the molecule.

  • Resonance structures are lowest in energy when:

    • All atoms with full octets

    • The minimum # of formal charges

    • Negative charges on electronegative atoms (more on this in a minute)

3 resonance structures2
3. Resonance Structures

Which N2O resonance structure(s) are lowest in energy?




B & C both have only two formal charges – lower in energy than A.

High in Energy:

Too many formal charges!!

C is lowest in energy because the negative charge is on O (more electronegative than N) – more on this in a minute!

4 bond lengths bond strengths
4. Bond Lengths & Bond Strengths

  • How does the type of bond (bond order) affect the properties of the bond?

  • Bonds get shorter as the number of electrons shared between the two atoms increases. (Bonds get shorter as the bond order increases.)

    Bond Length:

    Single Bond (1) > Double Bond (2) > Triple Bond (3)

4 bond lengths bond strengths1
4. Bond Lengths & Bond Strengths

  • The bond order also affects the strength of the bond (the bond enthalpy).

  • As the bond order increases, the bond gets stronger (harder to break).

C ionic vs covalent bond polarity
C. Ionic vs. Covalent: Bond Polarity

  • Now we know the difference between ionic and covalent bonds.

  • Given a specific compound, how do you know which type of bond to expect?

1 the continuum
1. The Continuum

  • In reality, all bonds have some covalent character & some ionic character.

  • In other words, all bonds fit somewhere on a continuum between covalent & ionic.

  • The real questions are:

    • How evenly are the electrons being shared between the two atoms?

    • How can you predict how evenly the electrons will be shared?

1 the continuum1
1. The Continuum

Covalent Polar Ionic


- electrons shared - electrons shared - electrons not

equally unequally shared at all

- No separation - some separation - complete separation

of charge of charge of charge

- held together by - held together by - held together by

shared electrons shared electrons electrostatic


2 electronegativity
2. Electronegativity

  • In 1932, Linus Pauling proposed the idea of electronegativity to explain the ways atoms share electrons in bonds.

Electronegativity: the ability of an atom in a covalent bond to attract shared electrons to itself; the “electron-pulling” power of an atom

2 electronegativity1
2. Electronegativity

The higher the EN, the more tightly an atom holds its electrons.

A bond polarity
a) Bond Polarity

  • The polarity of a bond describes how evenly the electrons are shared between the two atoms.

  • The more polar a bond, the less evenly the electrons are shared.

  • A polar bond is indicated by using + and - to represent the partial chargeson the atoms.

+ = less electronegative atom

- = more electronegative atom

A bond polarity1
a) Bond Polarity

  • The greater the difference in electronegativity between the two atoms, the more polar the bond between them will be.

A bond polarity2
a) Bond Polarity

  • “Approximate” Guidelines:

  • If EN ≤ 0.5, the electrons are shared fairly evenly & the bond is nonpolar covalent.

  • If EN  1.9, the electrons are localized on the more electronegative atom & the bond is ionic.

  • If 1.9 > EN > 0.5, the electrons are shared unequally & the bond is polar covalent.

The Continuum

H2 LiF HCl MgO CsI N2


Polar Covalent










D exceptions to the octet rule
D. Exceptions to the Octet Rule

  • Most structures containing main group elements follow the octet rule, but there are exceptions in certain compounds:

    • Some atoms have fewer than eight electrons (less than an octet).

    • Some atoms have more than eight electrons (more than an octet).

1 less than an octet
1. Less than an Octet

  • Elements in group 3A have three valence electrons.

  • Using those three electrons to form three bonds gives the central atom only six valence electrons (less than an octet).

  • eg. BH3

1 less than an octet1
1. Less than an Octet

  • Because boron has less than an octet, BH3 is very reactive – it will react to form an octet.

  • The nitrogen in ammonia uses its lone pair electrons to form a bond with boron – both B and N now have an octet.

2 more than an octet expanded valence
2. More than an Octet – Expanded Valence

  • Elements in the third period & lower can form stable compounds in which the central atom has more than 8 electrons – an expanded valence.

  • Unlike elements in the first two periods (like N & O), third row elements (like P & S) can use their empty 3d orbitals to accommodate extra electrons.