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Chapter 6: Chemical Bonding

C. Goodman, Doral Academy Charter High School, 2011-2014

Based on a PowerPoint by Mrs. S. Temple

1. What is a chemical bond?

2. Why do elements bond?

3. How can you classify bonds using the electronegativity of the elements in the compound?

Essential Questions: Section 6.1

Chemical bond

Ionic bond

Covalent bond

Non-polar covalent bond

Polar-covalent bond

Electronegativity

Vocabulary for section 6.1

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Chemical Bond A mutual electrical attraction between the nuclei and valence electrons of different atoms that binds the atoms together.

• Bonding occurs

so atoms

become more

stable

• Stability = full

valence

(8 electrons)

Why do elements bond?

Covalent Results from the

sharing of electron pairs between two atoms

Ionic Transfer electrons

to complete valence

This creates ions

Ions attract each other

Types of Bonding

8e-

8e-

6e-

6e-

6e-

6e-

6e-

6e-

Ionic Bonding

Remember an ion is

an atom or molecule

that has gained or

lost a(n) electron(s)

and now has a

negative or positive

charge

Formation of a Covalent

Bond

http://www.youtube.com/watch?v=lJxq37Xr18k

Bond length decreases

as the strength of the

bond (bond energy)

increases

Small bond length =

large bond energy

Large bond length =

small bond energy

Bond Length vs

Bond Energy

{ Ionic bonds:

metals + nonmetals

Covalent bonds: nonmetals + nonmetals

Metallic bonds metals + metals

The basics

P. 166 for this graphic

Electronegativity –

tendency of an element to

hold tightly to electrons –

p. 153 P. Table

w/electronegativity values

To find the type of bond

simply subtract the

electronegativity values of

the atoms bonded

See below…

Electronegativity and Bonding

Polar-covalent

Compound EN of Element 1

EN of element 2

delta EN Bond type

H-O

Li-F

Ca-Cl

C-O

Practice time

Nonpolar-covalent Bonded atoms have equal attraction for

shared electrons

Balanced distribution of electrical charge

EN = 0-0.4

Polar-covalent Bonded atoms have unequal attraction for

shared electrons

Unbalanced distribution of electrical charge

EN = 0.4–1.7

Types of Covalent Bonds

1. What is a molecule?

2. What is the octet rule?

3. How can Lewis structures be used to make models of molecules?

Essential Questions: Section 6.2

Compound

Molecule

Chemical formula

Molecular formula

Octet rule

Dual rule

Lewis structure

Resonance structure

Vocabulary for section 6.2

A NEUTRAL group of

atoms that are held together

by covalent bonds

Diatomic molecules:

molecules containing

only 2 atoms , both

the same element

Molecule

Chemical Formula

Shows relative numbers

of elements in a chemical

compound by using atomic

symbols and numerical

subscripts

Molecular Formula

Chemical formula for

molecular compounds

Chemical compounds tend to form so

that each atom, by gaining, losing, or

sharing electrons, has an octet of

electrons in its valence

Octet = 8 = ns2 + np6

The Octet Rule

Hydrogen/Helium 2 electrons (dual rule)

Boron tends to form with 6 electrons

Sometimes bonding occurs with d orbitals

Exceptions to the Octet

Boron trifluoride

BF3

When elements bond,

their orbitals Overlap

Diagrams that show only valence electrons (s and p only) as dots placed around the element's symbol

Inner shell electrons are not shown

Electron-Dot Notation

Let’s make some electron dot diagrams for individual elements

Element Lewis Structure

Hydrogen H

Carbon C

Boron B

Oxygen O

Nitrogen N

Fluorine F

Neon Ne

Practice

Time!

Electron dot diagrams for molecules

Diagrams representing covalent bonds

Atomic symbol = nuclei & inner-shell electrons

Dashes = covalent bond (each - = 2 e-)

Pair of dots on a side = lone pair electrons; they do

not bond

Single dots = unpaired electrons; they can bond

Lewis Structures

Now, how do you make Lewis Structures for actual molecules?

Steps for Drawing Lewis Structures Step 1: Draw

… electron dot structures for each element Step 2: Arrange

… elements – central and terminal Step 3: Connect

… the dots – Complete octets by bonding lone electrons

Step 4: Check your math (dual and octet rules)

1. How is the chemical formula for an ionic compound similar to the chemical formula for a molecular compound?

2. What are the characteristics of ionic compounds?

3. How are ionic compounds arranged, as solids?

4. How can Lewis structures be used to make models of polyatomic ions?

Essential Questions: Section 6.3

Ionic compound

Formula unit

Lattice energy

Polyatomic ions

Vocabulary for section 6.3

Ionic bonding

An atom or

molecule that has

gained or lost a(n)

electron(s) and now

has a negative or

positive charge

Review: Ions

• Composed of cations and anions

• Positive and negative charges are balanced

• Electrically neutral

• Metal cation + nonmetal anion

Ionic Compounds

Electrons transferred

Electromagnetic attraction between cations and anions holds ions together

Ionic Bonds

Chemical formulas are for molecules only

(covalently bonded)

Ionic compounds have formula units, instead

The lowest whole-number ratio of ions in an ionic

compound

i.e. Sodium chloride = NaCl = 1:1

Na+ Cl-

Magnesium chloride = MgCl2 = 1:2

Mg2+ Cl-

Formula Units

1. Crystalline solids at room temperature

2. High melting points

(strong forces of attraction between oppositely

charged ions)

3. Conduct electricity when melted or dissolved

(because ions are free to move in solution)

Electrical conductivity of ionic solutions

Properties of Ionic

Compounds

Unit cell smallest basic portion of crystal lattice that, when stacked together at various directions, can reproduce the entire crystal structure.

Ionic Crystals

Ionic crystals are brittle

When force is applied the ions shift. “Like charges” line up, then repel … end result, the crystal breaks

For ionic compounds, crystals are stable.

Minimum potential energy

Energy released (lattice energy) when

crystallizing: stabilizing energy of the

compound

Energy in Ionic Compounds

Ionic crystals: Really cool video

Crystallization of ionic solutions

A charged group of COVALENTLY bonded atoms

Polyatomic Ions

Polyatomic ions

Polyatomic Ions, continued

Some names of polyatomic anions relate to their oxygen content.

An -ate ending is used to name an ion with more oxygen atoms.

Examples: sulfate (SO42–), nitrate (NO3

–), chlorate (ClO3

–)

An -ite ending is used to name an ion with fewer oxygen atoms.

Examples: sulfite (SO32–), nitrite (NO2

–), chlorite (ClO2

–)

Memorize these! (in book, p. 226)

1. What is the “electron sea” model of metallic bonding?

2. How does the bond structure of metals explain their characteristics of ductility, malleability, and conductivity?

Essential Questions: Section 6.4

Metallic bond

Malleability

Ductility

Vocabulary for section 6.4

Metals contain closely packed cations with a “sea

of electrons”

Valence electrons are mobile

Drift freely from one part of the metal to another

Metallic bonds: attraction of valence electrons

and the positively charged metal ions

Metallic Bonds and

Properties

Good electrical conductors =

Electrons can flow freely throughout the

metal

Cations can slide past each other easily =

Ductile

Malleable

Properties of Metals

Explained

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1. What is VSEPR theory?

2. How can one predict the shapes of molecules or polyatomic ions using VSEPR theory?

Essential Questions: Section 6.5

Molecular geometry

Molecular polarity

Bond angle

VSEPR theory

Hybridization

Vocabulary for section 6.5

Molecular geometry and

molecular polarity affect

properties of molecules

Molecular geometry: the three-dimensional

arrangement of a molecule’s atoms.

Molecular polarity: the uneven distribution of

electrons in the molecular shape.

Molecular polarity influences forces between

molecules, e.g. water is polar, has

intermolecular bonds that affect its

properties (ice crystals, capillary action).

“valence-shell electron-pair repulsion.” Electron pair = lone pair OR covalent

bond Repulsion between sets of valence-

level electrons causes sets to be oriented as far apart as possible.

VSEPR Theory

BeF F

central atom = A Terminal atoms = B

Types of molecules, based on

molecular formula

Hybridization also affects behavior of molecules

Hybrid orbitals: orbitals of equal energy produced by the combination of two or more orbitals on the same atom.

Hybridization

Diatomic molecules Linear

AB

1 bond only, so no bond

angle

Hybridization: sp

Ex: HCl

AB2 molecules Linear 180o bond angle Hybridization: sp Ex: BeF2, CO2

If there are no lone pairs in the

Lewis structure, molecules are

planar

AB3 molecules equilateral triangle – trigonal planar 120° bond angles. Hybridization: sp2

Ex: BF3

• Boron central

AB4 molecules Tetrahedral 109.5o bond angles Hybridization: sp3

Ex: CH4

They repel other electron pairs more strongly than bonding pairs do.

Which groups have lone pairs?

Lone pairs act like bonded

atoms, but…

When these elements have bonded fully, they have essentially FOUR bonded pairs.

The lone pairs also take up space, so which affects the geometry of these molecules, as follows:

Group 15 and Group 16!

Lone pair at top – makes molecule polar Trigonal pyramidal shape 107o bond angle Hybridization: sp3

Example: NH3

Group 15

Two lone pairs at top – makes molecule polar Bent shape 105o bond angle Hybridization: sp3

Group 16