Chemical Bonding
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Chemical Bonding. Introduction to Bonding. Vocabulary. Chemical Bond attractive force between atoms or ions that binds them together as a unit bonds form in order to… decrease potential energy (PE) increase stability. Vocabulary. CHEMICAL FORMULA. IONIC. COVALENT. Formula Unit.

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Chemical Bonding

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Chemical Bonding

Introduction toBonding


Vocabulary

  • Chemical Bond

    • attractive force between atoms or ions that binds them together as a unit

    • bonds form in order to…

      • decrease potential energy (PE)

      • increase stability


Vocabulary

CHEMICAL FORMULA

IONIC

COVALENT

Formula

Unit

Molecular

Formula

NaCl

CO2


Vocabulary

COMPOUND

more than 2

elements

2 elements

Binary

Compound

Ternary

Compound

NaCl

NaNO3


Vocabulary

ION

2 or more atoms

1 atom

Monatomic

Ion

Polyatomic

Ion

Na+

NO3-


Types of Bonds

COVALENT

IONIC

e- are transferred from metal to nonmetal

e- are shared between two nonmetals

Bond Formation

Type of Structure

true molecules

crystal lattice

Physical

State

liquid or gas

solid

Melting

Point

low

high

Solubility in

Water

yes

usually not

yes (solution or liquid)

Electrical Conductivity

no

Other

Properties

odorous


Types of Bonds

METALLIC

e- are delocalized among metal atoms

Bond Formation

Type of Structure

“electron sea”

Physical

State

solid

Melting

Point

very high

Solubility in

Water

no

yes (any form)

Electrical Conductivity

malleable, ductile, lustrous

Other

Properties


Types of Bonds

Ionic Bonding


Types of Bonds

Covalent Bonding


Types of Bonds

Metallic Bonding - “Electron Sea”

RETURN


Types of Bonds

Ionic Bonding - Crystal Lattice

RETURN


Types of Bonds

Covalent Bonding - True Molecules

Diatomic Molecule

RETURN


Bond Polarity

  • Most bonds are a blend of ionic and covalent characteristics.

  • Difference in electronegativity determines bond type.


Bond Polarity

  • Electronegativity

    • Attraction an atom has for electrons.

    • higher e-neg atom  -

    • lower e-neg atom +


Bond Polarity

  • Electronegativity Trend (p. 151)

    • Increases up and to the right.


Bond Polarity

  • Nonpolar Covalent Bond

    • e- are shared equally

    • symmetrical e- density

    • usually identical atoms


-

+

Bond Polarity

  • Polar Covalent Bond

    • e- are shared unequally

    • asymmetrical e- density

    • results in partial charges (dipole)


Bond Polarity

Because chlorine is more electronegative than hydrogen, the electron density is greater around the chlorine atom.


Bond Polarity

  • Nonpolar

  • Polar

  • Ionic


Bond Polarity

  • Examples:

  • Cl2

  • HCl

  • NaCl

3.0-3.0=0.0

Nonpolar

3.0-2.1=0.9

Polar

3.0-0.9=2.1

Ionic


Bond Polarity

In a polar covalent bond, the bonding electrons will spend a greater amount of time around the atom that has the stronger affinity for electrons. A good example of a polar covalent bond is the hydrogen-oxygen bond in the water molecule.

Close


Chemical Bonding

II. Molecular Compounds


Energy of Bond Formation

  • Potential Energy

    • based on position of an object

    • low PE = high stability


no interaction

increased attraction

Energy of Bond Formation

  • Potential Energy Diagram

attraction vs. repulsion


increased repulsion

balanced attraction & repulsion

Energy of Bond Formation

  • Potential Energy Diagram

attraction vs. repulsion


Bond Energy

Bond Length

Energy of Bond Formation

  • Bond Energy

    • Energy required to break a bond


C. Molecular Nomenclature

  • Prefix System (binary compounds)

    1.Less e-neg atom comes first.

    2.Add prefixes to indicate # of atoms. Omit mono- prefix on first element.

    3.Change the ending of the second element to -ide.


PREFIX

mono-

di-

tri-

tetra-

penta-

hexa-

hepta-

octa-

nona-

deca-

NUMBER

1

2

3

4

5

6

7

8

9

10

Molecular Nomenclature


Molecular Nomenclature

  • CCl4

  • N2O

  • SF6

  • carbon tetrachloride

  • dinitrogen monoxide

  • sulfur hexafluoride


Molecular Nomenclature

  • arsenic trichloride

  • dinitrogen pentoxide

  • tetraphosphorus decoxide

  • AsCl3

  • N2O5

  • P4O10


Molecular Nomenclature

  • The Seven Diatomic Elements

    Br2 I2 N2 Cl2 H2 O2 F2

H

N

O

F

Cl

Br

I


Chemical Bonding

III. Ionic Compounds


Energy of Bond Formation

  • Lattice Energy

    • Energy released when one mole of an ionic crystalline compound is formed from gaseous ions


Lewis Structures

  • Covalent – show sharing of e-

  • Ionic – show transfer of e-


Lewis Structures

  • Covalent – show sharing of e-

  • Ionic – show transfer of e-


Ionic Nomenclature

Ionic Formulas

  • Write each ion, cation first. 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 - Roman numerals indicate the ion’s charge.


Ionic Nomenclature

Ionic Names

  • Write the names of both ions, cation first.

  • Change ending of monatomic ions 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.


Ionic Nomenclature

  • Consider the following:

    • Does it contain a polyatomic ion?

      • -ide, 2 elements  no

      • -ate, -ite, 3+ elements  yes

    • Does it contain a Roman numeral?

      • Check the table for metals not in Groups 1 or 2.

    • No prefixes!


Ionic Nomenclature

Common Ion Charges

1+

0

2+

3+

NA

3-

2-

1-


Ionic Nomenclature

  • potassium chloride

  • magnesium nitrate

  • copper(II) chloride

KCl

  • K+ Cl-

  • Mg2+ NO3-

Mg(NO3)2

CuCl2

  • Cu2+ Cl-


Ionic Nomenclature

  • NaBr

  • Na2CO3

  • FeCl3

  • sodium bromide

  • sodium carbonate

  • iron(III) chloride


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