Unit 4: Chemical Bonding and Molecular Structure Chapter 6 Notes ...


Unit 4: Chemical Bonding and Molecular Structure Chapter 6 Notes ...

Unit 4: Chemical Bonding and Molecular Structure

Chapter 6 Notes: Chemical Bonding

6.1: Introduction to Chemical Bonding

I. Chemical Bonds

Chemical Bond: attractive force between atoms or ions that binds them

together as a unit

** Concept: Why do most atoms form chemical bonds Independent particles are at

relatively high potential energy. (Most atoms are less stable existing by themselves.)

Nature favors arrangements in which potential energy is minimized-­‐

1. decreasing potential energy

2. creating more stable arrangements of matter.

II. Types of Chemical Bonds:

1. Ionic bonding: Chemical bonding that results from the electrical attraction

between positively charged cations and negatively charged anions. (In purely

ionic bonding, electrons are transferred from one ion to another.)

2. Covalent bonding: Atoms joined by covalent bonding share electrons. Covalent

bonding results from the sharing of electron pairs between two atoms. (In

purely covalent bonding, the shared electrons are “owned” equally by the two

bonded atoms.)

3. Metallic bonding: the chemical bonding that results from the attraction between

metal atoms and the surrounding sea of electrons

III. Ionic or Covalent

Using Electronegativity: a measure of the ability of an atom in a chemical compound

to attract electrons. The degree to which bonding between atoms of two elements is

ionic or covalent can be estimated by calculating the difference in the elements’

electronegativities. (END-­‐ electronegativity difference)

Ionic: Bonding between atoms with an electronegativity difference of greater than 1.7

Example: Fluorine (4.0) and Cesium (0.7)

4.0 – 0.7 = 3.3 will be an ionic bond.

Covalent: Bonding between atoms with an electronegativity difference of 1.7 or less

will be a covalent bond.

IV. Two types of Covalent bonds:

Covalent bonds will be either polar-­‐covalent or non-­‐polar covalent.

A. Non-­‐polar covalent bonds: a covalent bond in which the bonding electrons are

shared equally by the bonded atoms, resulting in a balanced distribution of electrical


Example: Bonding between atoms of the same element is completely covalent.

(The electronegativity difference will equal zero.) Generally, electronegativity

differences between 0 and 0.3 are considered non-­‐polar covalent bonds.

B. Polar-­‐covalent bonds: a covalent bond in which the bonded atoms have an

unequal attraction for the shared electrons. Generally, electronegativity differences

between 0.3 and 1.7 are classified as polar.

Example: The electronegativity difference between Chlorine and hydrogen is

3.0 – 2.1 = 0.9 indicating a polar-­‐covalent bond.

The electrons in this bond are closer to the more-­‐electronegative chlorine atom than to

the hydrogen atom. Therefore, the chlorine end of the bond has a partial negative

charge, indicated by the δ -­‐ . The hydrogen end of the bond then has an equal partial

positive charge, δ + .

δ + δ -

H Cl

** Concept: Why is most chemical bonding neither purely ionic nor purely covalent

Bonding between different elements is rarely purely ionic or purely covalent. It

usually falls somewhere between these two extremes, depending on how strongly the

atoms of each element attract electrons.

6.2: Covalent Bonding and Molecular Compounds

I. Important Definitions:

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

B. Molecular compound: a chemical compound whose simplest units are molecules

C. Chemical formula: indicates the relative numbers of atoms of each kind of a

chemical compound by using atomic symbols and numerical subscripts

D. Molecular formula: shows the tpes and numbers of atoms combined in a single

molecule of a molecular compound

II. Nature of the Covalent Bond

A. A bond between two nonmetals. Nonmetal to nonmetal

B. electrons are shared (the unpaired valence electrons)

C. Very strong bond

1. strong intramolecular forces (force of attraction within molecule)

2. can not be broken up by dissolving in water or melting, so never conduct

electricity (exception later in the year)

3. bond length and bond energy vary depending on the types of atoms bonded

D. Intermolecular forces (force of attraction between molecules)

1. much weaker than forces between ionic formula units

2. compared to ionic compounds:

a. lower melting points and boiling points

b. do not conduct electricity

III. The Octet Rule

A. Chemical compounds tend to form so that each atom, by gaining, losing, or sharing

electrons, has an octet of electrons in its highest occupied energy level.

B. Covalent compounds tend to form so that each atom, by sharing electrons,

completes an octet of electrons in its highest occupied energy level.

C. Exceptions to the octet rule:

1. Hydrogen: forms bonds in which it’s surrounded by two electrons

2. Boron: (3 valence electrons) forms bonds in which it’s surrounded by six


IV. Electron-­‐Dot Notation

A. Definition: an electron-­‐configuration notation in which only the valence electrons

of an atom of a particular element are shown, indicated by dots placed around the

element’s symbol. The inner-­‐shell electrons are not shown.

V. Lewis Structures:

A. Electron dot notation can also be used to represent molecules.

B. Unshared Pairs of electrons (Lone Pairs)

1. A pair of electrons that is not involved in bonding and that belongs

exclusively to one atom

2. represented by two dots

C. Shared pairs of electrons

1. Electrons pairs involved in covalent bonds (shared)

2. represented by dash

D. Drawing Lewis Structures (example: trichloromethane, CHCl3)

1. Determine the type and number of atoms in the molecule

1 x C, 1 x H, 3 x Cl

2. Determine the total number of valence electrons to be accounted for

C 1 x 4 e -­‐ = 4e -­‐

H 1 x 1 e -­‐ = 1e -­‐

Cl 3 x 7 e -­‐ = 21 e -­‐


26 e -­‐

3. Arrange atoms to form a skeleton structure for the molecule. If carbon is

present, it is the central atom. Otherwise the least elecronegative element

atom is central (except for hydrogen, which is never central). Then connect

the atoms by electron-­‐ pair bonds.

4. Add unshared pairs of electrons so that each nonmetal is surrounded by

eight electrons (except hydrogen-­‐ just two)—octet rule.

5. Count the electrons in the structure to be sure that the number of valence

electrons used equals the number available.

VI. Multiple Covalent Bonds

A. Double Bonds

1. A covalent bond produced by the shring of two pairs of electrons between

two atoms

2. Higher bond energy and shorter bond length than single bonds

B. Triple Bonds

1. A covalent bond produced by the sharing of three pairs of electrons between

two atoms

2. Higher bond energy and shorter bond length than single bonds

VII. Resonance Structures

A. Resonance Structure: bonding in molecules or ions that cannot be correctly

represented by a single Lewis Structure

B. To indicate resonance, a double-­‐headed arrow is placed between a molecule’s

resonance structures.

6.3 Ionic Bonding and Ionic Compounds

I. Important Definitions

A. Ionic compound: composed of positive and negative ions that are combined so that

the numbers of positive and negative charges are equal

B. Formula unit: the simplest whole number ratio of atoms from which an ionic

compound’s formula can be established

C. Crystal lattice structure: an orderly arrangement of ions that minimizes potential


D. Lattice energy: the energy released when one mole of an ionic crystalline

compound is formed from gaseous ions

E. Polyatomic ions: a charged group of covalently bonded atoms

II. Formation of Ionic Compounds

A. A bond between a positive ion (cation) and a negative ion (anion)

B. Metal ions bonded to nonmetal ions (metal to nonmetal)

C. Electron Configuration Changes

1. Electrons are trasferred from the highest energy level of one atom to the

highest energy level of a second atom, creating stable ions.

2. Opositely charged ions come together in a ratio that produces a net charge

of zero.

a. Na = 3s 1 Cl = 3s 2 3p 5

b. Na +1 = 2s 2 2p 6 Cl -­‐1 = 3s 2 3p 6

c. NaCl

III. Polyatomic ions

A. charged group of covalently bonded atoms

B. Mostly anions, one main exception: NH4 +1

C. Lewis structures:

1. Net charge equals charge of the ion

2. Written in brackets to show that the group as a whole has a charge.

IV. Comparing Ionic and Covalent Compounds

Ionic Compounds Covalent Compound

Structure Crystal lattice Molecule

Melting Point High Usually low

Boiling Point High Lower

Electrical Conductivity Yes, when dissolved, or

molten (not solid)

No, usually(exception noted


Solubility in water Yes Polar = yes, Non-­‐polar = no

6.4 Metallic Bonding

I. The Metallic Bond Model

A. Metallic Bonding

1. The chemical bonding that results from the attraction between metal atoms

and the surrounding sea of electrons

B. Electron Delocalization in Metals

1. Vacant p and d orbitals in metal's outer energy levels overlap, and allow

outer electrons to move freely throughout the metal

2. Valence electrons do not belong to any one atom

II. Metallic Properties

A. Metals are good conductors of heat and light

B. Metals are shiny

1. Narrow range of energy differences between orbitals allows electrons to be

easily excited, and emit light upon returning to a lower energy level

C. Metals are Malleable

1. Can be hammered into thin sheets

D. Metals are ductile

1. Ability to be drawn into wire

a. Metallic bonding is the same in all directions, so metals tend not to be


E. Metals atoms organized in compact, orderly crystalline patterns

F. Different metallic elements (and carbon) can be mixed to form alloys

1. Sterling silver

a. Ag = 92.5%, Cu = 7.5%

2. Brass

a. Cu = 60%, Zn = 40%

6.5 Molecular Geometry

A. VSEPR theory

1. valence-­‐shell, electron-­‐pair repulsion

2. states that repulsion between the sets of valence-­‐level electrons surrounding

an atom causes these sets to be oriented as far apart as possible

3. shared pairs of electrons will be as far apart from each other as possible

B. Unshared Electron Pairs

1. Lone pairs (unshared electron pairs) occupies space around the atom just as

bonding pairs do.

2. Lone pairs have a relatively greater effect on geometry than shared pairs

(Have greater repusive forces & tend to compress the angles between

bonding pairs)

C. Double and triple bonds are treated in the same way as single bonds

D. Polyatomic ions are treated similarly to molecules.

E. Molecular Polarity

1. The uneven distribution of molecular charge

2. A dipole is created by equal but opposite charges that are separated by a

short distance.

3. The direction of the dipole is from the positive pole to its negative pole

4. A dipole is represented by an arrow with a head pointing toward the negatie

pole and a crossed tail situated at the positive pole.

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