Chapter 6

Chemical Bonding

I: Introduction to Chemical Bonding

chemical bond - definition

high potential energy to a lower potential energy

A: Types of Chemical Bonding

valence electrons

cations

anions

ionic bonding - definition

giving up one or more electrons

opposite attraction of the ions is what forms the bond

covalent bonding - definition

figure 6-1 page 162

metallic bonding

valence electrons

kernel

1. Ionic or Covalent?

usually a blend of the two

electronegativity differences

figure 6-2 page 162

nonpolar covalent bond - definition

polar covalent bond - definition

Polar - definition

figure 6-3 page 163

electron density

d- and d+ e.g. HCl

Homework: 6.1

II: Chemical Bonding and Molecular Compounds

molecule - definition

two or more of the same type of atom or two or more different atoms - joined by covalent bond(s)

figure 6-4 page 164

molecular compound - definition

chemical formula - definition

molecular formula - definition

diatomic molecule - definition

A: Formation of a Covalent Bond

hydrogen-hydrogen bond

electron configuration

orbital notation of valence electrons

how can each element best get an octet?

nuclei of one atom attracted to electron cloud of other atom and visa versa

eventually the nuclei repel each other and so an optimal distance between the nuclei is reached

figure 6-5 page 165

figure 6-6 page 165

B: Characteristics of the Covalent Bond

the electron cloud and, therefore, the orbitals overlap

figure 6-7 page 167

bond length - definition

bond energy - definition

in general, the shorter the bond length, the higher the bond energy

table 6-1 page 168

sharing electrons to get an octet or a noble gas configuration e.g. hydrogen molecule

figure 6-8 page 168

Homework: 6.2

C: The Octet Rule

stable electron configuration -- filled s and p sublevel of the outermost energy level -- octet

noble gases - minimum of potential energy

Octet rule - statement

fluorine molecule F2

figure 6-9a page 169

hydrogen chloride molecule HCl

figure 6-9b page 169

1. Exceptions to the Octet Rule

applies to most main group elements that form covalent bonds

exceptions: B - three pairs of electrons instead of four e.g. BF3

Some elements can be surrounded by more than four pairs of electrons when bonding with highly electronegative elements fluorine, oxygen, and chlorine -- these are said to have expanded valence involving d as well as s and p orbitals

2. Electron Dot Notation

electron dot notation - definition

shows paired and unpaired valence electrons

figure 6-10 page 170

electron dot notation for hydrogen, nitrogen

3. Lewis Structures

electron dot formula vs electron dot notation

unshared pair or lone pair - definition

Lewis structure replaces the bonding pair of electrons with a dash

Structural formula - definition

e.g. H-H or F-F or H-Cl

Single covalent bond (single bond) - definition

Homework: 6.3

4. Multiple Covalent Bonds

double covalent bond - definition

two pairs of side by side dots or two parallel dashes

C2H4 is the compound ethene - double bond betweent the two carbons

triple covalent bond - definition

N2

figure 6-11 page 173

C2H2 is the compound ethyne - triple bond betwen the two carbons

multiple bonds - definition

bond energy of single vs double vs triple bond

length of bonds of single vs double vs triple

carbon, oxygen and nitrogen can form multiple bonds with the same element

Homework: 6.4

5. Resonance Structures

Some molecules and ions cannot be represented by a single Lewis structure e.g. O3

diagram page 175

O = O - O <----> O - O = O

the two bonds are identical but we don't have a good way to represent it using a single Lewis structure - thus resonance

the two structures are resonance structures or resonance hybrids

Resonance - definition

double headed arrow - only place this is allowed - between the two resonance structures

Homework: 6.5

6. Covalent Network Bonding

molecules with covalent bonding usually consist of molecules

molecules THEMSELVES ARE held together by forces holding the molecules together - NOT COVALENT BONDS

more on this in a later chapter

Homework: Chapter 6, 6.6

III: Ionic Bonding and Ionic Compounds

ionic compound - definition

most ionic compounds are crystalline solids - figure 6-12 page 176

Empirical formula - definition

the ratio of ions in an empirical formula depends on the charges of the ions combined e.g. calcium fluoride v sodium fluoride; CAF2 V NAF

A: Formation of Ionic Compounds

electron dot equation for reaction of sodium and chlorine to yield sodium chloride page 177

electron dot equation for the reaction between calcium and fluorine to yield calcium fluoride page 177

1. Characteristics of Ionic Bonding

ions minimize their potential energy by combining in an orderly arrangement known as a crystal lattice

figure 6-13 page 177

attractive forces present in ionic crystal include: a) those between oppositely charged ions; b) those between the nuclei and electrons of adjacent ions

repulsive forces in an ionic crystal include: a) those between like-charged ions; b) those between electrons of adjacent ions

distance between ions and their arrangement in a crystal represent a balance among all these forces

figure 6-14 page 177

the arrangements of ions and the strengths of attraction between the ions vary with a) the sizes and charges of the ions and b) the numbers of ions of different charges

figure 6-15 page 178

lattice energy - definition

it is used to compare bond strengths in ionic compounds

table 6-3 page 179

negative values indicate energy is released (exothermic)

B: A Comparison of Ionic and Molecular Compounds

The forces between molecules are much weaker than the forces of ionic bonding -- thus different properties

Melting point, boiling point and hardness depend on how strongly its basic units are attracted to each other

Many molecular compounds melt at low temperatures while many ionic compounds have higher melting and boiling points

Ionic cpds do not vaporize as readily at room temperature as molecular cpds do

Ionic cpds are hard but brittle

figure 6-17 page 179

In the molten state (melted) or when dissolved in water, ionic cpds conduct electricity but not in the solid state

Those ionic cpds that are not soluble (dissolve) in water are cpds in which the water molecules cannot overcome the attraction between the ions of the cpd

C: Polyatomic Ions

Polyatomic ions - definition

As with any ion, these result from the shortage or excess of electrons

Difference is that they contain more than one type of element

polyatomic ions - page 180

Homework: 6.7

IV: Metallic Bonding

The bonding in metals reflects their properties

Mobile valence electrons

A: Metallic Bond Model

In metals, usually, the s sublevel is filled and the three orbitals of the p sublevel are empty

Also, some have some vacant d orbitals

The vacant orbitals and the atoms' outer energy levels overlap. This overlapping of orbitals allows the outer electrons of the atoms to roam freely throughout the entire metal

delocalized

sea of electrons

Metallic bonding - definition

1. Metallic Properties

delocalization of electrons explain the high electrical and thermal conductivity

because they contain many orbitals separated by extremely small energy differences, metals can absorb a wide range of light frequencies - luster

Malleability - definition

Ductility - definition

Metallic bonding is the same in all directions throughout the solid. One plane of atoms in a metal can slide past another without encountering any resistance or breaking any bonds, unlike ionic crystals

2 Metallic Bond Strength

Varies with the nuclear charge of the metal atoms and the number of electrons in the metal's electron sea

Both are reflected in the metal's heat of vaporization

Heat of vaporization - definition

The amount of heat is a measure of the strength of the bonds that hold the metal together

table 6-4 page 182

V: Molecular Geometry

properties depend on bonding and molecular geometry

polarity of each bond and the geometry of the molecule determines the molecular polarity

molecular geometry - definition

molecular polarity - definition

molecular polarity influences the forces that act between molecules in liquids and solids

the chemical formula does not tell us directly about molecular polarity

two theories are prevalent a) molecular bond angles; b) describe the orbitals that contain the valence electrons

A: VSEPR Theory

figure 6-20 page 183

when there are only two atoms (diatomic) the geometry must be linear

more complicated molecules - consider all electron pairs surrounding the bonded atoms -- VSEPR theory

VSEPR - definition

we will consider molecules with no unpaired (unshared) electron pairs then those with unpaired (unshared) electron pairs and see how they differ

first example: BeF2

methodology: a) inert gas electron configuration; b) orbital notation for valence electrons; c) dot structure for each element; d) dot formula for molecule; e) consider effect of electron pairs on geometry

BeF2 dot formula

VSEPR states that electron pairs orient themselves to be as far away from each other as possible

figure 6-21 page 184

AB2

second example: BF3

methodology: same as above

AB3

third example: CH4

methodology: same as above

carbon uses its four valence electrons to bond with four other atoms; the four atoms occupy the corners of a tetrahedron, bond angles are 109.5 degrees; a tetrahedron consists of four identical triangles put together

other molecules on table 6-5 page 186 - know the general formulas, the geometry, the bond angles and examples

keep in mind that if the "B" atoms are not all the same, this will distort the molecule's geometry

1. VSEPR and Unshared Electron Pairs

central atom has unshared electron pairs e.g. water and ammonia

How VSEPR theory handles this: electron dot formula for ammonia, NH3; one pair occupies space just as the bonding pairs do

ammonia is described as an AB3E, where A is the central atom, B represents atoms bonded to the central atom, E represents lone pair(s)

bond angle is 107 degrees, less than the 109.5 degrees of a tetrahedron -- lone pairs repel electrons more than bonding pairs do; geometry is trigonal pyramidal

water molecule: the oxygen has two lone pairs; it is an AB2E2; oxygen is at the center of a modified tetrahedron, with two hydrogen's occupying two of the three corners of the base, one lone pair at the third corner of the base and one lone pair at the top. Geometry is described as bent.

Figure 6-22 page 185

Links to websites:

http://www.mhhe.com/physsci/chemistry/animations/chang_7e_esp/bom3s2_7.swf

http://www.hnhsoakland.org/faculty/kconover/Sites/PreviousPages/shapes.html

http://intro.chem.okstate.edu/1314F00/Lecture/Chapter10/VSEPR.html

Homework: 6.9

B: Hybridization

hybridization - definition

the compound methane, CH4

or carbon: electron configuration and orbital notation; change that results from hybridization that involves the 2s and 2p orbitals to form a new orbital called the sp3 hybridized orbital

figure 6-23 page 188

new electron configuration to show hybridization

hybrid orbital - definition

explain the geometry of molecules formed by Group 15 and 16 elements

Homework: 6.10

C: Intermolecular Forces

boiling point is a good measure

intermolecular forces - definition

weaker than chemical bonds

comparing boiling points of metals and ionic compounds with boiling points of molecular substances table 6-7 page 190

1. Molecular Polarity and Dipole-Dipole Forces

strongest intermolecular forces exist between polar molecules

polar molecules act as tiny dipoles

dipole - definition

The direction of the dipole is from the positive to the negative end of the molecule

Indicated by an arrow pointing toward the negative end of the molecule. The tail of the arrow is crossed and is at the positive end of the molecule

dipole-dipole forces - definition

short range forces

figure 6-25 page 191

For molecules containing more than two atoms, molecular polarity depends on both the polarity and the orientation of each bond; e.g. water

figure 6-26 page 191

e.g. ammonia

In some molecules, individual bond dipoles cancel one another, causing the resulting molecular polarity to be zero e.g. carbon tetrachloride and carbon dioxide figure 6-26 page 191

A polar molecule can induce a dipole in a nonpolar molecule by temporarily attracting its electrons; The short range intermolecular force is somewhat weaker than the dipole-dipole force

The force of an induced dipole accounts for the solubility of nonpolar oxygen in water

Figure 6-27 page 192

2. Hydrogen Bonding

strong type of dipole-dipole force

Occurs between the hydrogen of one molecule and the F, O, or N of an adjacent molecule

examples are HF, water and ammonia

hydrogen bonding - definition

represented by dotted lines connecting the hydrogen of one molecule to the highly electronegative atom (fluorine, oxygen or nitrogen) of an adjacent molecule

Figure 6-28 page 192

effect can be seen by comparing boiling points of phosphine and ammonia; hydrogen sulfide and water on table 6-7 page 190

3. London Dispersion Forces

noble gases and nonpolar molecules still experience a weak intermolecular attraction

because of the random motion of the electrons, the distribution of the electrons may become uneven; creates a temporary dipole in the molecule; this temporary dipole can induce a dipole in an adjacent molecule and the two molecules are attracted for an instant then the effect disappears

figure 6-29 page 193

London dispersion forces - definition

These forces operate between all atoms and molecules

they are the only intermolecular forces acting among noble gas atoms, nonpolar molecules, and slightly polar molecules

Notice the low bp's of the noble gases, etc. on table 6-7 page 190

Because these forces depend on the motion of the electrons, their strength increases with the number of electrons in the interacting atoms or molecules i.e. the forces increase with increasing atomic or molar mass

Note the BP's of helium and argon, hydrogen and oxygen, chlorine and bromine

Homework: 6.11

end of notes

London dispersion forces are the intermolecular attractions resulting from the constant motion of electrons and the creation of instantaneous dipoles

Hydrogen bonding is the intermolecular force in which a hydrogen atom that is bonded to a highly electronegative atom is attracted to an unshared pair of electrons of an electronegative atom in a nearby molecule

Dipole-dipole forces are the forces of attraction between polar molecules

A dipole is created by equal but opposite charges that are separated by a short distance

Intermolecular forces are the forces of attraction between molecules

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

Hybridization is a mixing of two or more atomic orbitals of similar energies on the same atom to produce new orbitals of equal energies

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.

VSEPR stands for valence shell electron pair repulsion

Molecular polarity is the uneven distribution of molecular charge

Molecular geometry is the three dimensional arrangement of a molecule's atoms in space

Heat of vaporization is the heat necessary to convert a metal from the solid state to individual metal atoms in the gaseous state

Ductility is the ability of a substance to be drawn, pulled, or extruded through a small opening to produce a wire

Malleability is the ability of a substance to be hammered or beaten into thin sheets.

Metallic bonding is the bonding that results from the attraction between metal atoms and the surrounding sea of electrons.

A polyatomic ion is a charged group of covalently bonded atoms

Lattice energy is the energy released when one mole of an ionic crystalline compound is formed from gaseous ions

An empirical formula (formula unit) indicates what elements are present and the simple whole number ratio of those elements

Ionic compound is composed of positive and negative ions that are combined so that the numbers of positive and negative charges are equal

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

A triple covalent bond is a covalent bond produced by the sharing of three pairs of electrons between two atoms

Multiple bonds are double or triple covalent bonds

A double covalent bond is a covalent bond produced by the sharing of two pairs of electrons between two atoms

A single covalent bond is a covalent bond produced by the sharing of one pair of electrons between two atoms

Structural formula indicates the kind, number, arrangement, and bonds but not the unshared pairs of the atoms in a molecule

An unshared pair of electrons is a pair of electrons that is not involved in bonding and that belongs exclusively to one atom

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

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.

Bond energy is the energy required to break a chemical bond and form neutral isolated atoms

The distance between two bonded atoms at their minimum potential energy is the bond length

A diatomic molecule is a molecule containing only two atoms.

The diatomic molecules are fluorine, chlorine,bromine, iodine, hydrogen, oxygen, and nitrogen.

A molecular formula shows the types and numbers of atoms combined in a single molecule of a molecular compound

A chemical formula indicates the relative numbers of atoms of each kind in a chemical compound by using atomic symbols and numerical subscripts

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

A molecule is a neutral group of atoms that are held together by covalent bonds

Polar means an uneven distribution of charge

A polar covalent bond is a covalent bond in which the bonded atoms have an unequal attraction for the shared electrons

A nonpolar covalent bond is a covalent bond in which the bonding electrons are shared equally by the bonded atoms, resulting in a balanced distribution of electrical charge

Covalent bonding results from the sharing of electron pairs between two atoms

Ionic bonding is a chemical bond that results from the electrical attraction between large numbers of cations and anions

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