PRINCIPLES OF CHEMISTRY I

CHEM 1211

CHAPTER 9

DR. AUGUSTINE OFORI AGYEMANAssistant professor of chemistryDepartment of natural sciences

Clayton state university

CHAPTER 9

CHEMICAL BONDING

- The attractive force that holds atoms together

- The result of interactions between electrons in the combining atoms

Two types of chemical bonds - Covalent and Ionic (electrovalent) bonds

CHEMICAL BOND

Covalent Bond

- Formed through the sharing of one or more pairs of electrons between two atoms

- Always involve two nonmetals

- Electron sharing

CHEMICAL BOND

Ionic Bond

- Formed by attraction between two oppositely charged ions

- Formed as a result of the transfer of electron(s) from atom(s) to another atom(s)

- Often formed between metal and nonmetal ions through electrostatic attraction

- Electron transfer

CHEMICAL BOND

CHEMICAL BOND

Two concepts

- Valence Electrons

- Octet Rule

VALENCE ELECTRONS

- Not all electrons in a given atom participate in bonding

- Only valence electrons are available for bonding (electrons in the outer most shell)

- For representative and noble-gas elements these electrons are always found in the s or p subshells

VALENCE ELECTRONS

- Using electron configuration to determine the number of valence electrons

C: 1s22s22p2

O: 1s22s22p4

Na: 1s22s22p63s1

- Using electron-dot structure (Lewis symbol) to designatethe number of valence electrons

(place first 4 dots separately on four sides and pair up as needed)

∙C∙ :O∙ Na∙.

.

..

.

VALENCE ELECTRONS

Three important facts about valence electrons

- Representative elements in the same group of the periodic table have the same number of valence electrons

- The number of valence electrons for representative elements is the same as the group number (with A) in the periodic table

- The maximum number of valence electrons for any given element is eight

OCTET RULE

- Electrons arranged with 8 valence electrons are more stable than all others

- The valence electron configuration of the noble gases are considered the most stable

(all have 8 valence electrons; helium has 2)

- All noble gases have the outermost s and p subshells completely filled

OCTET RULE

- The noble gases are the most unreactive of all elements

- Atoms of many elements tend to acquire the 8 valence electron configuration through chemical reactions

- Atoms of elements tend to gain, lose, or share electrons to produce a noble-gas electron configuration

- This results in the formation of compounds

- This tendency is known as the OCTET RULE

IONIC BOND

- Electron transfer

- Metals donate electrons to form positive ions

- Nonmetals accept electrons to form negative ions

- The electrons lost by the metal are the same ones gained by the nonmetal

- The positive and negative ions attract one another to form ionic compounds

- Ions combine in ratios to obtain charge neutrality (net charge = 0)

- The symbol for positive ions is always written first

IONIC BOND

Lewis Structures- Lewis structures involve compounds

- Lewis symbols involve individual elements

Na∙ + ∙Cl: [Na]+ [:Cl:]- NaCl

CaCl2

..

.. ..

..

∙Ca∙ +..

..∙Cl:

∙Cl:..

..[Ca]2+

[:Cl:]-

[:Cl:]-

..

..

..

..

IONIC BOND

Energetics

Removing an electron from Na(g) to form Na+(g)Na(g) → Na+(g) + e- E = +496 kJ/mol

Adding an electron to Cl(g) to form Cl-(g) Cl(g) + e- → Cl-(g) E = -349 kJ/mol

- Attraction between the unlike charges draws ions togethercausing energy to be released

Heat of formation of ionic substances is quite exothermicNa(s) + 1/2Cl2(g) → NaCl(s) Hf

o = -410.9 kJ

IONIC BOND

Energetics

- Ionic compounds do not contain discrete molecules but ordered arrays of positive and negative ions

(result of energy released)

NaCl- Formula unit that indicates combining ratio

- A given sodium ion has six immediate chloride ion neighbors

- A given chloride ion has six immediate sodium ion neighbors

IONIC BOND

Lattice Energy

- The energy required to completely separate one mole of a solid ionic compound into its gaseous ions

- Increases with increasing charge on the ions and decreasing distance between the radii of the ions (from electrostatic

potential energy, Eel)

NaCl(s) → Na+(g) + Cl-(g) Hlattice = +788 kJ/mol

IONIC BOND

Lattice Energy

- Highly endothermic indicating ions are strongly attracted to one another

- Reason why ionic compounds are hard, brittle, and have high melting points

Melting point of NaCl is 801 oC

IONIC BOND

- Generally, transition metals do not form ions that have the noble-gas configuration

- Transition metals first lose valence-shell s electrons and then as many d electrons as required to form ions

- Transition metals can form different cations

Fe: Fe2+ and Fe3+

Sn: Sn2+ and Sn4+

Pb: Pb2+ and Pb4+

TRANSITION METAL IONS

COVALENT BONDING

- Involve electron sharing

- Usually occurs between two nonmetals

- The basic structural unit in covalent bonding is a molecule

- Forms molecular compounds

H H∙ ∙ :H H

Two hydrogen atoms H + H

Hydrogen molecule H H

1s electrons Shared electron pair

COVALENT BONDING

- Two neclei attract the same shared electrons to form a covalent bond

- Orbitals containing the valence electrons overlap to create a common orbital

- The electrons move throughout the common orbital

- The electrons are shared by both nuclei

COVALENT BONDING

- The valence electrons help each atom achieve a noble-gas configuration

H∙ ∙H H : H H H

:F∙ :F : F: :F F:

H : F: F:H

H2

∙F:

∙F:

..

.. ..

..F2

HF H∙..

..

.. ..

.. ..

..

..

..

....

..

..

..

bonding electrons nonbonding

electrons

LEWIS STRUCTURES

Bonding Electrons - The pairs of valence electrons involved in the

covalent bond formation

Nonbonding Electrons (Lone Pairs of Electrons)- The pairs of valence electrons not involved in

electron sharing

LEWIS STRUCTURES

H2O

H ∙

H ∙

O : O

H

H : OR O

H

H

- Oxygen (O) has six valence electrons - Gains two more through electron sharing with H

- Achieves a noble-gas configuration

..

..:

..

.. .. :

LEWIS STRUCTURES

NH3

H ∙

H ∙

N N

H

H : OR N

H

HH ∙

H H

: : :.

..

..

..

LEWIS STRUCTURES

- Nitrogen (N) has five valence electrons - Gains three more through electron sharing with H

- Achieves a noble-gas configuration

CH4

∙ C ∙ C

H

H : OR C

H

H

H H

H ∙

: H H.

.

..

..H ∙

H ∙

H ∙

LEWIS STRUCTURES

- Carbon (C) has four valence electrons - Gains four more through electron sharing with H

- Achieves a noble-gas configuration

SINGLE COVALENT BOND

- Two atoms share one pair of valence electrons

- Represented by one line

- Bond order is one

Bond Order- Number of electron pairs that are shared between two atoms

Bond Length- The minimum energy distance between the nuclei

of two bonded atoms in a molecule

DOUBLE COVALENT BOND

- Two atoms share two pairs of valence electrons

- Represented by two lines

- Approximately twice as strong as a single covalent bond between the same two atoms

- Bond order is two

DOUBLE COVALENT BOND

CO2

- C has four valence electrons and needs four more

- Each O atom has six valence electrons and needs two more

:O::C::O: or O C O

- Possible for elements that need two electrons to complete their octet

.. ..

TRIPLE COVALENT BOND

- Two atoms share three pairs of valence electrons

- Represented by three lines

- Approximately thrice as strong as a single covalent bond between the same two atoms

- Bond order is three

- Bond length decreases with increasing bond order

TRIPLE COVALENT BOND

N2

- Nitrogen has five valence electrons and needs three more tocomplete its octet

- Each nitrogen must share three of its electrons with the other

:N:::N: or :N N:

- Possible for elements that need three or more electrons to complete their octet

COORDINATE COVALENT BOND

- Both electrons come from only one of the two bonding atoms

- Oxygen often forms coordinate covalent bonds

: +X Y :X Y

filled orbital vacant orbital shared electron pair

H : O : Cl :

coordinate covalent bond

Chlorous acid (HClO2)Hypochlorous acid (HOCl)

.. H : O : Cl : O : ..

..

.. .. ..

.. ..

.. ..

ELECTRONEGATIVITY

- The ability of an atom to attract to itself the electrons in a chemical bond

- Electronegativity depends on atom size

nuclear charge number of inner shell electrons

- Increases from left to right across periods on the periodic table

- Increases from bottom to top within groups on the periodic table

- Flourine is the most electronegative of all the elements

- Nonmetals are more electronegative than metals

- Indicative of the fact that nonmetals gain electrons and metals lose electrons

ELECTRONEGATIVITY

LEWIS STRUCTURES

- Calculate the total number of valence electrons in the molecule (use group numbers in the periodic table)

HClO2

H (group IA) has 1 valence electronCl (group VIIA) has 7 valence electronsO (group VIA) has 6 valence electrons

Total electron count = 1 + 7 + 2(6) = 20

- Determine the central atom

The central atom - mostly appears only once (SO3, SO2, CH4)

- is usually any additional element other than H and O (HNO3, H2SO4)

- is C in almost all carbon-containing compounds- is neither H nor F (can make only one covalent bond)

- for O and H containing compounds O is bonded to the central atom and H to O

HClO2 (Cl is the central atom)

LEWIS STRUCTURES

- Write the atoms in the order in which they are bonded together

- Place a pair of electrons between each pair of atoms

H : O : Cl : O

HClO2

LEWIS STRUCTURES

- Add nonbonding electron pairs to all atoms except the central atom- Each atom should have eight electrons

- H needs only 2 electrons

HClO2

H : O : Cl : O :

16 out of the 20 electrons have been used up

..

.. ..

..

LEWIS STRUCTURES

HClO2

H : O : Cl : O :

20 out of the 20 electrons have been used up

..

.. ..

..

LEWIS STRUCTURES

- Place any remaining electrons on the central atom of the structure

.. ..

HClO2

H : O : Cl : O : ..

.. ..

..

LEWIS STRUCTURES

.. ..

- This step is not needed in this case since Cl has completed its octet

- If the central atom has less than eight move nonbonding electron pairs to form double or triple bonds

HClO2

H : O : Cl : O : ..

.. ..

..

LEWIS STRUCTURES

.. ..

- Count the total number of electrons in the Lewis structure(must equal the initial number)

20 electrons equal to the intial 20

HCN

H (group IA) has 1 valence electronC (group IVA) has 4 valence electronsN (group VA) has 5 valence electrons

Total electron count = 1 + 4 + 5 = 10

LEWIS STRUCTURES

- Calculate the total number of valence electrons in the molecule (use group numbers in the periodic table)

HCN (C is the cental atom)

LEWIS STRUCTURES - Determine the central atom

The central atom

- mostly appears only once (SO3, SO2, CH4)- is usually any additional element other than H and O

(HNO3, H2SO4)- is C in almost all carbon-containing compounds

- is neither H nor F (can make only one covalent bond)- for O and H containing compounds

O is bonded to the central atom and H to O

HCN

H : C : N

LEWIS STRUCTURES

- Write the atoms in the order in which they are bonded together

- Place a pair of electrons between each pair of atoms

HCN

H : C : N :

LEWIS STRUCTURES

.. ..

10 out of the 10 electrons have been used up

- Add nonbonding electron pairs to all atoms except the central atom- Each atom should have eght electrons

- H needs only 2 electrons

HCN

H : C : N :

LEWIS STRUCTURES

.. ..

10 out of the 10 electrons have been used up- Nothing left to be placed on the central atom

- Place any remaining electrons on the central atom of the structure

HCN

H : C : N :

LEWIS STRUCTURES

.. ..

H : C ::: N :

- If the central atom has less than eight move nonbonding electron pairs to form double or triple bonds

HCN

H : C : N :

LEWIS STRUCTURES

.. ..

H : C ::: N :

- Count the total number of electrons in the Lewis structure(must equal the initial number)

10 electrons equal to the initial 10

POLYATOMIC IONS

The total number of electrons for negative charges

- increase the number of electrons by the magnitude of the charge

SO42-

S (group VIA) has 6 valence electronsO (group VIA) has 6 valence electrons

Charge of -2

Total number of electrons = 6 + 4(6) + 2 = 32

NH4+

N (group VA) has 5 valence electronsH (group IA) has 1 valence electron

Charge of +1

Total number of electrons = 5 + 4(1) - 1 = 8

The total number of electrons for positive charges

- decrease the number of electrons by the magnitude of the charge

POLYATOMIC IONS

Ionic compound containing polyatomic ion

- The cation and anion are treated separately

Na2SO4

[Na]+

[Na]+S

:O:

:O:

O::O

2- ..

.. ..

.. ..

..

POLYATOMIC IONS

BOND POLARITY

Nonpolar Covalent Bond

- Two atoms involved in electron sharing have equal or similar electronegativity

- Typically less than 0.4

- Equal sharing of electrons

F2, H2, O2

BOND POLARITY

Polar Covalent Bond

- There exists unequal sharing of electrons

- One atom is more electronegative than the other

- One atom attracts electrons more strongly than the other

- Electronegativity difference is between 0.4 and 1.5

HCl, CO

BOND POLARITY

- Increasing bond polarity renders a bond more ionic

- Ionic bonds have electronegativity difference greater than 2.0

- Most bonds are a mixture of pure ionic and pure covalent

- No natural boundary between ionic and covalent bonding

For electronegativity difference between 1.5 and 2.0 - ionic bond if metal and a nonmetal are involved

- polar covalent bond if two nonmetals are involved

- Polar covalent bonds create partial positive and negative charges on the atoms involved

- Delta (δ) is used to designate these partial charges δ+ for less electronegative atom δ- for more electronegative atom

H Cl:

BOND POLARITY

..

..δ+ δ-

H Cl:

- An arrow with a cross can also be used

- The arrowhead is near the more electronegative end of the bond

BOND POLARITY

..

..

DIPOLE MOMENTS

- A dipole establishes whenever two electrical charges of equal magnitude but opposite sign are separated by a distance

- The quantitative measure of the magnitude of the dipole is known as the dipole moment

µ = Qr

µ = dipole moment Q = electrical charge (two equal and opposite charges Q+ and Q-)

r = distance between the centers of Q+ and Q- Units: debyes (D)

1 D = 3.34 x 10-30 coulomb-meters (C-m)

FORMAL CHARGE

- Used to predict stability and connectivity

To Calculate the Formal Charge- All nonbonding (unshared electrons) are assigned to the atom

on which they are found- Half of the number of bonding electrons are assigned to each

atom in the bond

Formal Charge = Number of electronsassigned to the atom

Number of valenceelectrons in the

isolated atom-

- Sum of formal charges equals the overall charge- Sum of formal charges in neutral atoms equals zero

FORMAL CHARGE

Formal Charge = Number of electronsassigned to the atom

Number of valenceelectrons in the

isolated atom-

[:C N:]-

Six electrons in the triple bondC: 2 nonbonding electrons + 3 bonding electrons = 5

Number of valence electrons = 4N: 2 nonbonding electrons + 3 bonding electrons = 5

Number of valence electrons = 5

Formal Charge of C = 4 - 5 = -1Formal Charge of N = 5 - 5 = 0

[:C N:]-

RESONANCE STRUCTURESOzone (O3)

SO3

O

O O OO

O

:

:: :

::

:::

: :

:

S S S

O O

OO

OO OO

O

: :

: ::: : :

::

::

::: ::::

: ::

::

EXCEPTIONS TO THE OCTET RULE

Odd Number of Electrons

(NO, ClO2, NO2)

N O N O

: ::.

: :

: .and

- Called radicals and are very reactive

For exampleThe immune system uses NO to fight bacteria

Less Than an Octet of Valence Electrons(Electron Defficient)

- Usually in compounds of boron, beryllium, and aluminum

- BF3 (only six valence electrons around boron)- BeH2

- BeF2

- BH3

- AlH3

EXCEPTIONS TO THE OCTET RULE

More Than an Octet of Valence Electrons(Expanded)

- Occurs in elements of period 3 and beyond

- No d orbitals in periods 1 and 2 to hold extra electrons

- PCl5 (10 valence electrons around phosphorus)

- SF6 (12 valence electrons around sulfur)- XeF4

EXCEPTIONS TO THE OCTET RULE

STRENGTH OF COVALENT BONDS

- Determined by the energy required to break the bonds

- Bond enthalpy is the enthalpy change for breaking the bond in one mole of a gaseous substance

- D(Cl — Cl) denotes bond enthalpy in Cl2

- Bond enthalpies are always positive (energy is consumed)

- To decompose CH4 into C and 4H, H = 1660 kJ There are 4 equivalent C — H bonds

Average C — H bond enthalpy = D(C — H) = (1660/4) kJ/mol = 415 kJ/mol

BOND ENTHALPIES

- Bond breaking is an endothermic process

- Bond formation is an exothermic process

formedbondsofenthalpiesbondbrokenbondsofenthalpiesbondΔH rxn

- Bond enthalpy increases with increasing number of bonds

- Bond length decreases with increasing number of bonds