Ch 6 Chemical Bonding

What you should learn in this section (objectives):

Define chemical bond

Explain why most atoms form chemical bonds

Describe ionic and covalent bonding

Explain why most chemical bonding is neither purely ionic or purely covalent

Classify bonding type according to electronegativity differences.

Introduction to Chemical Bonding

There are very few atoms that exist as individual particles in nature. Most atoms are bonded to

other atoms to form compounds. A chemical bond is a mutual electrical attraction between the nuclei

and valence electrons of different atoms that binds the atoms together. One reason atoms bond is to

decrease their amount of potential energy. When atoms exist by themselves they have relatively high

potential energy. Nature favors arrangements that have minimum potential energy. When atoms bond

it decreases the amount of potential energy and creates a more stable arrangement of matter. It takes

less energy to bond atoms together than to break the bonds between atoms. Bond energy is the

amount of energy it takes to break a chemical bond.

Types of Chemical Bonds

Ionic bonding- chemical bonding that results from the electrical attraction between cations and

anions. In purely ionic bonding atoms completely give up electrons to other atoms. Ionic bonding

generally involves metals and nonmetals

Covalent bonding- the sharing of electron pairs between two atoms. In a pure covalent bond

the electrons are shared by the two bonded atoms. Covalent bonding generally involves two nonmetals.

A nonpolar-covalent bond is when the bonding electrons are shared equally by the bonded atoms,

resulting in a balanced distribution of electrical charges. When the distribution of charge is uneven we

call this polar. Polar covalent bonds occur when the bonded atoms have an unequal attraction for the

shared electrons.

Ionic or Covalent?

Most chemical bonds are somewhere in between purely ionic and purely covalent. We use the

difference in electronegativity values to determine the type of bond that is formed. Remember from Ch

5 that electronegativity is an atom’s ability to attract electrons to its self. We use the following image

to determine the bond type.

Problem A

Use electronegativity differences to classify bonding between sulfur and the following elements:

hydrogen, cesium, and chlorine. In each pair, which atom will be more negative?

Ionic- electronegativity difference is

greater than 1.67

Polar covalent- electronegativity

difference is less than 1.67

Nonpolar covalent-electronegativity

difference is less than 0.4

Covalent Bonding and Molecular Compounds

What you should learn in this section (objectives):

Define molecule and molecular formula

Explain the relationships among potential energy, distance between approaching atoms, bond

length and bond energy.

State the octet rule

List the six basic steps used in writing Lewis structures

Explain how to determine Lewis structures for molecules containing single bonds, multiple

bonds, or both.

Explain why scientists use resonance structures to represent some molecules.

A molecule is a neutral group of atoms that are held together by covalent bonds. They can exist as two

or more of the same elements bonded together or two or more different elements bonded together. A

chemical compound whose simplest units are molecules is called a molecular compound. A chemical

formula shows the relative numbers of atoms of each kind in a chemical compound by using atomic

symbols and numerical subscripts. Remember sub means below. A molecular formula shows the types

and numbers of atoms combined in a single molecule of a molecular compound.

Formation of a Covalent Bond

Nature favors chemical bonding because most atoms have lower potential energy when they

are bonded to other atoms than they have as they are independent particles.

When 2 hydrogen atoms approach each other 2 “bad” things happen: electron/electron

repulsion and proton/proton repulsion. One “good” thing that happens: proton/electron attraction.

When the attractive forces offset the repulsive forces, the energy of the tow atoms decreases and a

bond is formed. Remember, nature is always striving for a lower energy state.

too FAR

too CLOSE

just RIGHT

Bond length is the distance between the two nuclei where the energy is minimal between the two

nuclei. In other words, it is the average distance between two bonded atoms.

When bonds form individual atoms release energy as they change from isolated individual atoms to

molecules. Bond energy is the amount of energy that is required to break the bond. The units for bond

energy are usually kj/mol (kilojoule per mole).

Octet Rule

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 (outer most ring of the atom).

Exceptions to the Octet Rule

Fewer than 8: H at most only 2 electrons (one bond),BeH2, only 4 valence electrons around Be (only 2

bonds), Boron compounds only 6 valence electrons (three bonds)

Expanded valence (more than 8): can only happen if the central element had d-orbitals which means it is

from the 3rd period or greater and can thus be surrounded by more than four valence pairs in certain

compounds. The number of bonds depends on the balance between the ability of the nucleus to attract

electrons and the repulsion between the pairs. Some of the more elements are fluorine, oxygen,

chlorine and noble gases.

Electron-Dot Notation

• When two atoms form a covalent bond, their shared

electrons form overlapping orbitals.

• This achieves a noble-gas configuration.

• The bonding of two hydrogen atoms allows each atom to

have the stable electron configuration of helium, 1s2.

• To keep track of valence electrons, it is

helpful to use electron-dot notation.

• Electron-dot notation is 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.

Lewis Structures

An unshared pair, also called a lone pair, is a pair of electrons that is not involved in bonding

and that belongs exclusively to one atom.

Lewis Structures are 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.

A structural formula indicates the kind, number, arrangement, and bonds but not the unshared pairs of

the atoms in a molecule. Example F-F and

H-Cl.

Single bonds (sigma bonds) are a covalent bond in which one pair of electrons is shared between two

atoms. They are represented by two dots (electrons) or one dash. These are the longest bonds, but also

the weakest

Double bonds (pi bonds) are covalent bonds in which two pairs of electrons are shared between two

atoms. They are represented by four dots (electrons) or 2 dashes = Example C=C

Triple bonds (pi bonds) are covalent bonds in which three pairs of electrons are shared between two

atoms. They are represented by six dots (electrons) or 3 dashes. Example These are shortest,

but also the strongest.

Carbon, nitrogen, oxygen, phosphorous, and sulfur are the most common elements that form multiple

bonds.

Drawing Lewis Structures

1. H is always a terminal atom. ALWAYS connected to only one other atom.

2. Lowest electronegativity is the central atom in a molecule.

3. Find the total number of valence electrons by adding up group numbers of the elements. For

ions add electrons for negative charges and subtract electrons for positive charges.

4. Place one pair of electrons (sigma bond) between each pair of bonded atoms.

H:H

5. Subtract from the total number of bonds you just used.

6. Place lone pairs about each terminal atom (except H) to satisfy the octet rule. Left over pairs are

assigned to the central atom.

7. If the central atom is not yet surrounded by four electron pairs, convert on or more terminal

atom lone pairs to a double or triple bond ( pi bonds). Only C, N, O, P, and S can form multiple

bonds (pi bonds)

Resonance Structures

Resonance refers to bonding in molecules or ions that cannot be correctly represented by a single Lewis

structure. Ozone (O3) exists as an average of these two images, so it must be shown both ways.

Ionic Bonding and Ionic Compounds

What you should learn in this section (objectives):

Compare and contrast a chemical formula for a molecular compound with one for an ionic

compound.

Discuss the arrangements of ions in crystals

Define lattice energy and explain its significance

List and compare the distinctive properties of ionic and molecular compounds

Write the Lewis structure for a polyatomic ion given the identity of the atoms combined and

other appropriate information.

Formation of an Ionic Compound Ionic compounds are composed of positive (cation) and negative (anion) ions that are combined so that the numbers of positive and negative charges are equal Ionic Bond - Completely transfer electrons.

Positive charge – cation – lost electrons to the anion. Negative charge – anion – gained electrons from the cation.

Positive charge must equal and, therefore, cancel the negative charge. Example: Sodium Chloride – sodium wants to lose one electron to become stable and chlorine wants to gain one electron to become stable. Formula unit – a chemical formula of the smallest sample of an ionic compound. Ionic Character

Ionic compounds have the greatest ionic character with full on charged ions. The further

the ions are apart in electronegativity, the more the ionic character.

Molecular compounds have very low electronegativity. The closer the ions are in

electronegativity, the less the ionic character.

Characteristics of Ionic Bonding

Ionic compounds are crystalline solids at room temperature. They are arranged in

repeating three-dimensional pattern called a crystal lattice. These structures are very

orderly and stable. Example: In solid NaCl, each Na is surrounded by six Cl and each Cl is

surrounded by six Na.

These crystalline solids also have very high melting points. It is extremely hard to break

the attraction between the cations and anions because of their stability. Example: NaCl

melts at 800 ˚Celsius. The energy released when one mole of an ionic crystalline

compound is formed from gaseous ions is called lattice energy.

Ionic compounds conduct electric currents when molten (liquid) or dissolved in water

(aqueous). The cations and anions are then able to migrate freely.

Ionic compounds are electrically neutral salts (solids). Many of these compounds appear

as minerals in the Earth’s crust.

Comparing Ionic and Molecular Compounds

The force that holds ions together in an ionic compound is a very strong electrostatic

attraction.

In contrast, the forces of attraction between molecules of a covalent compound are

much weaker.

This difference in the strength of attraction between the basic units of molecular and

ionic compounds

gives rise to different properties between the two

types of compounds.

Molecular compounds have relatively weak forces between individual molecules.

They melt at low temperatures.

The strong attraction between ions in an ionic compound gives ionic compounds some

characteristic properties, listed below.

o very high melting points

o hard but brittle

o not electrical conductors in the solid state, because the ions cannot

move

Polyatomic Ions

Polyatomic Ions – a group of atoms that acts as a unit with a single charge

Begin memorizing polyatomic ions…get the list from the website and make flashcards.

Know the formula, the charge, and the correct spelling of the name of the polyatomic ions listed

on the website. We will have a quiz over these.

Metallic Bonding

What you should learn in this section (objectives):

Describe the electron-sea model of metallic bonding, and explain why metals are good

electrical conductors

Explain why metal surfaces are shiny

Explain why metals are malleable and ductile but ionic crystalline compounds are not.

The Metallic-Bond Model

Metallic bonding is the chemical bonding that is a result from the attraction between metal

atoms and the surrounding sea of electrons. A sea of electrons refers to the free moving valence

electrons in an atom. These electrons are delocalized which means that they can freely move to any

other atom

Properties of Metals

Good conductors of electricity – electrons enter one end of the metal bar and leave the other.

Ductile – can be stretched into wires.

Malleable – can be pounded into shapes. Metals ions slide passed one another in a sea of

drifting

Molecular Geometry

What you should learn in this section (objectives):

Explain VSEPR theory

Predict the shapes of molecules or polyatomic ions using VSEPR theory

Explain how the shapes of molecules are accounted for by hybridization theory

Describe dipole-dipole forces, hydrogen bonding, induced dipoles, and London dispersion forces

and their effects on properties such as boiling and melting points

VSEPR Theory - Valence Shell Electron Pair Repulsion

The VSEPR theory states that repulsion between the sets of valence-level electrons surrounding an atom

causes these sets to be oriented as far apart as possible. The electron dot structures are not flat 2D

structures, but are 3D in real life.

Molecules adjust their shapes so that the valence electron pairs are as far apart as possible.

**See the chart on pg. 200 of your textbook

Linear – angles are 180 degrees – definitely will be linear if only have two atoms in the molecule. No

lone pairs and two covalent bonds or three lone pairs and one covalent bond around central atom.

Example: CO2

Bent – again, unshared pair(s) strongly repels the covalent bonding pairs. Two lone pairs and two shared

pairs around central atom. All angles are 105 degrees. Example: H2O

Trigonal-Planar – three shared pairs (covalent bonds ) separate as much as possible, but are unaffected

by a lone pair (no lone pairs) of electrons like the pyramidal structure. Example: BF3

Trigonal-Pyramidal – one unshared pair strongly repels the three shared pairs (covalent bonding),

pushing them closer together. All angles are 107 degrees. Example: NH3

Tetrahedral – four faced – four shared pairs and no lone pairs, all angles are 109.5 degrees. Example:

CH4

Trigonal Bipyramidal – five shared pairs separate as much as possible, but are unaffected by a lone pair

of electrons (no lone pairs). Example: PCl5

Octahedral – six shared pairs separate as much as possible, but are unaffected by a lone pair of

electrons (no lone pairs). Example: SF6

Hybridization

Hybridization - two atoms combine, their atomic orbitals overlap to produce molecular orbitals. One

electron from each atomic orbital combines to create a shared pair in a molecular orbital.

sp hybridization – has electrons in 2 orbitals

sp2 hybridization – has electrons in 3 orbitals

sp3 hybridization – has electrons in 4 orbitals

Intermolecular Forces (IMF)

The forces of attraction between molecules are known as intermolecular forces. These forces vary in

strength and are generally weaker than bonds

The strongest IMF exists between polar molecules. A dipole is created by equal but opposite charges

that are separated by a short distance. The direction of a dipole is from the dipole’s positive pole to its

negative pole.

Dipole Interactions - when polar molecules are attracted to one another: opposite charged regions of polar

molecules are attracted.

Hydrogen Bonds – a particularly strong dipole interaction specifically involving hydrogen at the partially positive

pole. Hydrogen is covalently bonded to a very electronegative atom AND to an unshared pair of another atom.

The negative region in one polar molecule attracts the positive region in

adjacent molecules. So the molecules all attract each other from opposite

sides.

Such forces of attraction between polar molecules are known as dipole-dipole

forces.

Dipole-dipole forces act at short range, only between

nearby molecules.

Hydrogen is able to bond with the unshared pair of electrons from another molecule because its’ valence

electrons are not shielded from the nucleus by another layer of electrons (hydrogen’s valence electrons are

directly up against the nucleus). Example: H2O

The more electronegative the element that hydrogen is bonded to the stronger the intermolecular attractions.

Dispersion Forces - weakest of all molecular interactions – caused by the motion of electrons.

Vibrating electrons may end up moving randomly closer to one atom or another creating a momentary dipole.

The more electrons there are the greater the interaction between nonpolar molecules.