chem bond 3

8/7/2019 chem bond 3

1/11

Chemistry Lecture Chapter 6- Chemical

BondingSection 6.1: Introduction to Chemical Bonding- 1-

2 daysBell work: Define chemical bond, ionic bond, covalent bond,

nonpolar covalent bond, and polar covalent bond.

Types of Chemical Bonding-Ionic or Covalent?

Atoms seldom exist as independent particles in nature. Why?

Because when they bond with each other, atoms are more

stable and have a lower potential energy.

When atoms bond their valence electrons are redistribute inways that make the atoms more stable.

When they do they form chemical bonds. Chemical bonds =

a mutual electrical attraction between the nuclei and valence

electrons of different atoms that binds the atoms together.

The first type of bond that we will look at is the ionic bond.

Ionic bond = chemical bonding that results from theelectrical attraction between large numbers of cations and

anions. See figure 6-1.

Covalent bonding results from the sharing of electron pairs

between two atoms. This is the second type of chemical

bond. Atoms with covalent bonds form independent

molecules.

Remember that electronegativity is the measure of an

atoms ability to attract electrons. The degree to which

bonding between atoms of two elements is ionic or covalent

can be estimated by calculating the difference between the

elements electronegativities.

1


2/11

Bonding between atoms with an electronegativity

difference of 1.7 or less has a covalent bond. Bonds

between two atoms of the same element are also covalent

because they have no electronegativity.

There are two kinds of covalent bonds. The first is a

nonpolar covalent bond = a covalent bond in which the

bonding electrons are share equally by the bonded atoms,

resulting in a balanced distribution of electrical charge. Their

electronegativities are between 0-.3.

A polar covalent bond is a covalent bond in which the

bonded atoms have an unequal attraction for the sharedelectrons. That means that electrons tend to spend more time

closer to one element in the compound than the other. The

electronegativity is going to be between .3-1.7.

Recognize that electronegativity difference is only a general

guide for determining bonding type. There are some

exceptions. See figure 6-2 for the general rule.

Section Review, pg. 163, #1-4.

Study Guide 6.1

2


3/11

Section 6.2- Covalent Bonding and Molecular

Compounds- 2 days.Bell work- Define molecule and molecular formula.

2

nd

day: Define bond length, bond energy

Introduction

1. Molecule = a neutral group of atoms that are held

together by covalent bonds.

2. Molecular compound = a chemical compound whose

simplest units are molecules.

3. Chemical formula = the relative numbers of atoms of

each kind of a chemical compound by using atomic

symbols and numerical subscripts.

4. Diatomic molecule = a molecule containing only two

atoms of the same element.

I. Formation of a Covalent Bond

Visualize a kick ball rolling toward a ditch. The ball rolls

down the near slope of the ditch, passes the bottom of the

ditch, and rolls part of the way up the opposite slope of the

ditch. It then rolls back down, passes the bottom, and rollspart of the way up the near side of the ditch. This pattern

continues, each time rolling to a lesser height, until it

eventually comes to rest at the bottom of the ditch. At this

point, the balls potential energy is at a minimum. The

formation of a covalent bond between two atoms is like

the ball in the ditch. The atoms alternately attract and

repel each other until they reach a distance (or bond

length) at which their potential energy is minimized.

See example of hydrogen on pg. 165.

3


4/11

II. Characteristics of the Covalent Bond Atoms that form covalent bonds are nonmetals.

The bottom of the ditch represents the balance between

attraction and repulsion in a stable covalent bond.

Bond length = the distance between two bonded atoms

at their minimum potential energy, that is, the average

distance between two bonded atoms.

Energy is released when the bond is formed.

Bond energy = the energy required to break a chemical

bond and form neutral isolated atoms. The energy

released to form the bond is the same amount of energy

needed to break the bond.

Look at table 6-1 and note that bond length decreases asthe strength of the bond increases.

III. The Octet Rule

Noble gases exist in nature at independent atoms. The

reason is the stable valence electron configurations.

Other main group atoms can fill their outermost s and p

orbitals by sharing electrons through covalent bonds.

The octet rule = chemical compounds tend to form so thateach atom, by gaining, losing, or sharing electrons, has an

octet of electrons in its highest occupied energy level.

Exceptions to the Octet Rule- Hydrogen with two

electrons, boron with its three and three others (BF3).

Others expand past 8 valence electrons usually by

involving the d orbitals as well as the s and p orbitals.

IV. Electron-Dot Notation Covalent bond formation usually involves only he

electrons in an atoms outermost energy levels, or the

atoms valence electrons.

To keep track of these valence electrons it is helpful to use

electron-dot notation.

4


5/11

Electron-dot notation = an electron-configuration notation

in which only the valence electrons of an atom of a

particular element are shown, indicated by the dots placed

around the elements symbol.

Do sample problem 6-2 on pg. 170.

V. Lewis Structures Lewis structures = formulas in which atomic symbols

represent nuclei and inner-shell electrons, dot-pairs or

dashes between two atomic symbols represent electron

pairs in covalent bonds, and dots adjacent to only one

atomic symbol represent unshared electrons.

The pair of dots between the two chemical symbolsrepresents the shared pair of the covalent bond. An

unshared pair or lone pair is a pair of electrons that is not

involved in bonding and that belongs exclusively to one

atom.

Structural formulas indicated the kind number,

arrangement, and bonds but not the unshared pairs of the

atoms in a molecule. (F-F, H-Cl are structural formulas)

See sample problem 6-3 on pg. 171.

VI. Multiple Covalent Bonds

Double bond = a covalent bond produced by the sharing of

two pairs of electrons between two atoms.

Carbon, Nitrogen, and Oxygen can share more than

one electron pair.

Illustrate this by two side-by-side pairs of dots or

by two parallel dashes. Triple bond = a covalent bond produced by the sharing of

three pairs of electrons between two atoms.

Nitrogen has 5 valence electrons so it is seeking

three bonds.

Illustrate this by 3 side-by-side dots, 3 parallel lines

5


6/11

Double and triple bonds are referred to as multiple bonds

Sample problem 6-4 on page 174. There are additional

sample problems on pg. 199A.

VII. Resonance Structures

Resonance = bonding in molecules or ions that cannot be

correctly represented by a single Lewis structure.

Ozone resonates or shifts between two Lewis structures. It

has one single and one double oxygen-oxygen bond and

ozone alternates between the two structures.

VIII. Covalent-Network Bonding

Sometimes covalent bonding occurs between largenumbers of molecules.

This is called a covalent network.

This is similar to the network structures of ionic bonds.

This will be discussed in chapter 11.

Interactive Tutor, Module 4- Covalent Bonding- How and why

covalent bonds form

Section Review, PG. 175, #1-4

Study Guide 6-2

Quiz over first two sections

6


7/11

Section 6.3- Ionic Bonding and Ionic Compounds-

2 daysBell work: define ionic compound, formula unit, and lattice

energy2nd day: Lab B6

Introduction:

1. Ionic compound = a compound composed of positive and

negative ions that are combined so that the numbers of

positive and negative charges are equal.

2. Formula unit = the simplest collection of atoms from which

an ionic compounds formula can be established.

I. Formation of Ionic Compounds- Characteristics of Ionic

Bonding

Most ionic compounds exist as crystalline solids.

Draw blocks to demonstrate the difference between ionic

and covalent bonds. A molecule with covalent bonds

should be flush on all sides. An ionic compound looks like

a block with pieces jutting out to represent cations and

anions in a huge aggregate of ions with minimized potentialenergy. Remember that the elements are looking for ways

to minimize potential energy.

Example: Na becomes Na+ and Cl becomes Cl- as they

approach each other. Now each has a noble gas

configuration.

Then the attractive forces at work within an ionic crystal

come into play. There are 2 attractive forces. The first is

the one between oppositely charged ions and the second isbetween the nuclei and electrons of adjacent ions.

But there is also the repulsive force between like-charged

ions and between electrons of adjacent ions.

7


8/11

The distance between ions in a crystal creates a balance

between all these forces. Draw a lattice structure like figure

6-14.

To compare bond strengths in ionic compounds, chemists

compare the amounts of energy released when separatedions in a gas come together to form a crystalline solid.

Lattice energy = the energy released when one mole of an

ionic crystalline compound is formed from gaseous ions.

The negative energy values in table 6-3 indicate that energy

is released when bonds are formed.

Teaching Tip : Lattice energy is negative because it

indicates a release of energy in the formation of a bond.

Bond energy is the amount of energy needed to break abond so it is a positive number. Both are measured in

kJ/mol. See table 6-2 on pg. 173 and table 6-3 on pg. 179.

II. A Comparison of Ionic and Molecular Compounds

Ionic bonds are stronger than covalent bonds. This

difference in the strength of attraction between these bonds

gives rise to different properties in the two types of

compounds. Ionic compounds have a higher melting point and boiling

point than covalent molecules. This relates to the strength

of the bond and the difficulty in breaking the bond.

Ionic compounds are hard but brittle. Why? If a layer in an

ionic compound is moved the repulsive forces make the

layers part completely, meaning that ionic compounds are

brittle.

Ionic compounds are not conductors when solid but dobecome conductors when melted or dissolved. Why?

Because ions separate from each other and are free to move

about and conduct electricity due to the charged particles

that are no longer held in neutral bonds.

8


9/11

III. Polyatomic Ions

Polyatomic ion = a charged group of covalently bonded

atoms.

These ions have characteristics of both covalent and ionic

bonds. Certain atoms form covalent bonds with each otherbut have an overall charge like an ion.

The ammonium ion has nitrogen that normally has 7

electrons (5 valence electrons) and 4 hydrogens that have 4

electrons. One of the electrons is pulled away (NH4) so it

has an overall charge of +1. Nitrate (NO3) is -1, sulfate

(SO4) is -2, and phosphate (PO4) is -3.

These bond with other ions to form ionic compounds.

Section Review, pg. 180, #1-4

Study Guide 6-3.

ChemFile Lab B6: Chemical Bonds. Compare properties of ionic

and covalent compounds.

9


10/11

Chemistry Lecture 6.4- Metallic Bonding- 1 dayBell work: define metallic bonding, malleability, and ductility.

The Metallic-Bond Model

Chemical bonding is different in metals than it is in

ionic, or covalent compounds. That is why metals have

unique characteristics.

Metals are great electrical conductors in the solid state.

Why? Because of the highly mobile valence electrons

Lesson Starter: Have samples of NaCl and mossy zinc.

Hip both with a hammer. The salt is brittle and shatters.

Zinc will flatten out because it is malleable.

Metals have several empty orbitals, especially in

the d-block elements.

The vacant orbitals tend to overlap between

metal atoms.

This causes the electrons to become delocalized,

which means that they dont belong to anyparticular atom but move freely about the

metals network of empty atomic orbitals. This

is also called a sea of electrons around the metal

atoms.

Metallic bonding = the chemical bonding that

results from the attraction between metal atoms

and the surrounding sea of electrons.

2. Metallic Properties

This freedom of movement among electrons

account forhigh conductivity of metals.

The large number of electron orbitals means that

metals absorb a wide range of light frequencies.

10


11/11

When light is absorbed it excites the electrons to

higher energy levels.

When they fall back to lower levels, energy is

emitted in the form of light.

This de-excitation is why metals have shiny

surfaces.

Metals can also be easily formed into new

shapes. The are malleable and ductile.

Malleability = the ability of a substance to be

hammered or beaten into thin sheets.

Ductility = the ability of a substance to be

drawn, pulled, or extruded through a small

opening to produce a wire. The electrons can move easily and so metals can

be easily shaped.

3. Metallic Bond Strength

The amount of heat required to vaporize the metal

is a measure of the bond strength of the metal.

Heat of vaporization = the heat at which bonded

atoms are converted into individual metal atomsin the gaseous state.

Bond strength varies with the nuclear charge and

number of electrons in the particular metal. Bond

strength increases across a period and decreases

down a group.

Section Review

Study Guide 6.4Chapter 6 test

11

chem bond 3

Documents