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Topic 6: Chemical Bonding & Molecular Geometry

Chemical Bonding (Chapter 6 in Modern Chemistry)

Atoms seldom exist as independent particles in nature. Most substances consist of combinations of atoms that are held together by chemical bonds. A chemical bond is a mutual electrical attraction between the nuclei and valence electrons of different atoms that binds the atoms together. Why do atoms make bonds? It turns out that most atoms are less stable existing by themselves than when they are combined. By bonding with each other, atoms decrease in potential energy, thereby creating more stable arrangements of matter.

When atoms bond, their valence electrons are redistributed in way that make the atoms more stable. The way in which the electrons are redistributed determines the type of bonding. In Topic 4 you learned the main-group metals tend to lose electrons to form positive ions, or cations. Nonmetals tend to gain electrons to form negative ions, or anions. Chemical bonding that results from the electrical attraction between cations and anions is called ionic bonding. If a bond is purely ionic, an atom will completely give up electron(s) to another atom.

In contrast, atoms joined by covalent bonding share electrons. Covalent bonding results from the sharing of electron pairs between two atoms. In a purely covalent bond, the shared electrons are  “owned”  equally  by  the  two  bonded  atoms.

an example of an ionic bond an example of a covalent bond

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Ionic or Covalent?

Bonding between atoms 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.    Recall  that  electronegativity  is  a  measure  of  an  atom’s  ability to attract electrons. To determine whether a bond is ionic or covalent you have to calculate the difference in electronegativities of the two atoms involved.

An electronegativity difference greater than 1.67 is referred to as an ionic bond.

Electronegativity differences of 1.67 or less have an ionic bond character of 50% or less. These compounds are typically classified as covalent. Bonding between two atoms of the same element is completely covalent. This is called a nonpolar-covalent bond. This is a covalent bond in which the bonding electrons are shared equally by the bonded atoms, resulting in a balanced distribution of electrical charge. Bonds having only 0% to 5% ionic character, or an electronegativity difference equal to or less than 0.3, are considered nonpolar-covalent.

Bonds having an ionic character between 5% and 50%, or with corresponding electronegativity differences of 0.3 to 1.67, are classified as polar-covalent. A polar-covalent bond is a covalent bond in which the bonded atoms have an unequal attraction for the shared electrons.

Here are some example of determining bond character based on the electronegativities from the periodic table on the next page.

Bonding Pair Electronegativity

difference Bond type Li and F 3.98 - 0.98 = 3.00 ionic

Cu and S 2.58 – 1.90 = 0.68 polar-covalent

I and Br 2.96 -2.66 = 0.30 nonpolar-covalent

In summary, subtract the electronegativities, if:

Greater than 1.67 ionic Above 0.3 to 1.67 polar-covalent 0.3 or less nonpolar-covalent�

3.3

Ionic

100%

1.67

Polar -covalent

50%

0.3 Nonpolar-covalent 5%

0

0%

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Electronegativities of the elements

Periodic table of electronegativity using the Pauling scale

Please note, if you do not have the electronegativities, ionic compounds are generally made up of elements that are far apart on the periodic table. For example: K & Cl make an ionic bond. Sr & Br make an ionic bond. S & O make a covalent bond. (Watch out for hydrogen; even though it is in group 1, it has a high electronegativity.)

Task 6a

1. Using the electronegativity chart, determine the type of bond between the atoms in the following compounds.

a. AgCl b. K2O c. Br2 d. HCl

→  Atomic radius decreases  →  Ionization energy increases  →  Electronegativity increases  → Group

(vertical) 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18

Period (horizontal)

1

H 2.20

He

2

Li 0.98

Be 1.57

B 2.04

C 2.55

N 3.04

O 3.44

F 3.98

Ne

3

Na 0.93

Mg 1.31

Al 1.61

Si 1.90

P 2.19

S 2.58

Cl 3.16

Ar

4

K 0.82

Ca 1.00

Sc 1.36

Ti 1.54

V 1.63

Cr 1.66

Mn 1.55

Fe 1.83

Co 1.88

Ni 1.91

Cu 1.90

Zn 1.65

Ga 1.81

Ge 2.01

As 2.18

Se 2.55

Br 2.96

Kr 3.00

5

Rb 0.82

Sr 0.95

Y 1.22

Zr 1.33

Nb 1.6

Mo 2.16

Tc 1.9

Ru 2.2

Rh 2.28

Pd 2.20

Ag 1.93

Cd 1.69

In 1.78

Sn 1.96

Sb 2.05

Te 2.1

I 2.66

Xe 2.60

6

Cs 0.79

Ba 0.89

*

Hf 1.3

Ta 1.5

W 2.36

Re 1.9

Os 2.2

Ir 2.20

Pt 2.28

Au 2.54

Hg 2.00

Tl 1.62

Pb 2.33

Bi 2.02

Po 2.0

At 2.2

Rn 2.2

7

Fr 0.7

Ra 0.9

**

Rf

Db

Sg

Bh

Hs

Mt

Ds

Rg

Cn

Uut

Uuq

Uup

Uuh

Uus

Uuo

Lanthanoids

*

La 1.1

Ce 1.12

Pr 1.13

Nd 1.14

Pm 1.13

Sm 1.17

Eu 1.2

Gd 1.2

Tb 1.1

Dy 1.22

Ho 1.23

Er 1.24

Tm 1.25

Yb 1.1

Lu 1.27

Actinoids

**

Ac 1.1

Th 1.3

Pa 1.5

U 1.38

Np 1.36

Pu 1.28

Am 1.13

Cm 1.28

Bk 1.3

Cf 1.3

Es 1.3

Fm 1.3

Md 1.3

No 1.3

Lr 1.3

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2. Only using the periodic table, determine if the bonds below are covalent or ionic. Do not use the electronegativity chart.

a. XeCl6 b. CsF c. MgCl2 d. NO2

Ionic Bonding and Ionic Compounds

An ionic compound is composed of positive and negative ions that are combined so that the numbers of positive and negative charges are equal. The chemical formula of an ionic compound shows the ratio of ions present in a sample of any size. A formula unit is the simplest  collection  of  atoms  from  which  an  ionic  compound’s  formula  can  be  established. For example, one formula unit of sodium chloride, NaCl, is one sodium cation plus one chloride anion. (In the naming  of  a  monatomic  anion,  the  ending  of  the  element’s  name  is  replaced  with    –ide.) The ratio of ions in a formula unit depends on the charges of the ions combined. They must achieve electrical neutrality. In other words, the sum of the charges must equal zero.

For example: calcium fluoride

calcium has a 2+ charge, Ca2+ fluorine has a 1- charge, F-

In order for their charges to equal zero, there has to be two fluoride ions for every calcium ion.

CaF2

The Formation of Ionic Compounds Consider that a sodium atom and a chlorine atom are approaching each other. The two atoms are neutral and have one and seven valence electrons, respectively.

Sodium loses its electron to chlorine forming Na+ and Cl-.

Sodium atom + Chlorine atom Æ Sodium ion + Chloride ion

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Covalent Bonding

Most substances are composed of molecules. A molecule is a neutral group of atoms that are held together by covalent bonds. A molecule may consist of two of more atoms of the same element, as in oxygen, or of two or more different atoms, as in water or sugar.

Oxygen molecule, Water molecule, Sucrose molecule,

O2 H2O C12H22O11

A chemical compound whose simplest units are molecules is called a molecular compound. The composition of a compound is given by its chemical formula. A chemical formula indicates the relative numbers of atoms of each kind in a chemical compound by using atomic symbols and numerical subscripts. The chemical formula of a molecular compound is referred to as a molecular formula. A molecular formula shows the types and numbers of atoms combined in a single molecule of a molecular compound. The molecular formula for water, for example, is H2O, which reflects the fact that a single water molecule consists of one oxygen atom joined by separate covalent bonds to two hydrogen atoms. A molecule of oxygen, O2, is an example of a diatomic molecule. A diatomic molecule is a molecule containing only two atoms.

Formation of a Covalent Bond

Remember that nature favors chemical bonding because most atoms have lower potential energy when they are bonded to other atoms than they have when they are not bonded.

Using the example of two hydrogen atoms, if they are separated by enough distance, they will not influence each other. However, if the two hydrogen atoms approach each other, there will come to a position in which the nucleus of one hydrogen atoms attracts the electron from the other hydrogen atom. The attraction corresponds to a decrease in the total potential energy of the atoms. At the same time the electrons of the two hydrogen atoms are repelling each other. The protons in the nuclei of the two hydrogen atoms are also repelling each other. The repulsion results in an increase in potential energy. The relative strength of attraction and repulsion between the charged particles is dependent on the distance separating the two hydrogen atoms.

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Here the arrows indicate the attractive and repulsive forces between the electrons and nuclei of two hydrogen atoms. Attraction (red) between particles corresponds to a decrease in potential energy of the atoms, while repulsion (blue) corresponds to an increase.

The attractive force dominates and continues to pull the two hydrogen atoms closer together until they get to a distance at which the repulsion between like charges equals the attraction between opposite charges. At this position the two hydrogen atoms (now a hydrogen molecule, H2) have their minimum potential energy and are close enough to share electrons. They are now covalently bonded.

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Hydrogen Hydrogen atoms molecule

All individual hydrogen atoms contain a single, unpaired electron in a 1s atomic orbital. When two hydrogen atoms form a molecule, they share electrons in a covalent bond. The sharing of the electrons allows each atom to have the stable electron configuration of helium, 1s2. The tendency is for atoms to achieve a noble-gas configuration by sharing electrons. The Octet Rule

Unlike other atoms, the noble-gas atoms exist independently in nature. They possess a minimum of energy existing on their own because of the special stability of their electron configuration. This stability results from the fact that, with the exception of helium and its two electrons in a completely filled outer shell, the noble-gas  atoms’  outer  s and p orbitals are completely filled by a total of eight electrons.

Other main-group atoms can effectively fill their outermost s and p orbitals with electrons by sharing electrons through covalent bonding. Such bond formations follow the octet rule. Chemical compounds tend to form so that each atom, by gaining, losing, or sharing electrons, has an octet of electrons init highest occupied energy level. Here is an example of the covalent bonding of hydrogen and fluorine to make hydrofluoric acid, HF. H H F F

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Exceptions to the Octet Rule

Most main-group elements tend to form covalent bonds according to the octet rule, However, there are exceptions. Hydrogen forms bonds in which there are only two electrons. Boron, B, and aluminum, Al, has just three valence electrons so it makes bonds surrounded by 6 electrons. Other atoms can be surrounded by more than eight electrons. Examples of elements that make bonds with expanded valences are phosphorus, P, in PF5 and sulfur, S, in SF6.

Task 6b

1. Use orbital notation to illustrate the bonding in the chlorine molecule, Cl2.

2. Describe the general location of the electrons in a covalent bond.

Metallic Bonding

Chemical bonding is different in metals than it is in ionic or molecular compounds. The highest energy levels of most metal atoms are occupied by very few electrons. The properties of metals are due to the highly mobile valence electrons of the atoms that make up a metal. The highest energy levels of most metal atoms are occupied by very few electrons. They have many vacant orbitals.     The   vacant   orbitals   in   the   atoms’   outer   energy   levels   overlap.     This   overlapping   of  orbitals allows the outer electrons of the atoms to roam freely throughout the entire metal. The electrons are delocalized, which means that they do not belong to any one atom but move freely about   the   metal’s   network   of   empty   atomic   orbitals.     These   mobile   electrons   form   a   sea of electrons around the metal atoms, which are packed together in a crystal lattice. The chemical bonding that results from the attraction between metal atoms and the surrounding sea of electrons is called metallic bonding.

Here are two diagrams to help you understand how the electrons surround the cation in a metallic bond.

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Van der Waals force

Van der Waals forces are a group of weak intermolecular forces that vary in strength but are generally weaker than bonds (ionic, covalent, and metallic). We will discuss hydrogen bonding, dipole-dipole interactions, and London dispersion forces (LDF).

Hydrogen Bonding

Some hydrogen-containing compounds, such as hydrogen fluoride, water, and ammonia, have unusually high boiling points. This is explained by the presence of a particularly strong type of dipole-dipole force. In compounds that contain hydrogen and a very electronegative element (N, O, F) an intermolecular attraction occurs between the molecules. Note that this is not a bond between atoms within the molecule, but an attraction among the molecules. Below, the dotted lines represent the hydrogen bonds between water molecules. This allows water to have relatively high melting and boiling points for hydrogen compounds that bond with group 16 nonmetals. This also explains why ice expands and floats.

Dipole-Dipole Attractions

A dipole is a molecule or a part of a molecule that contains both positively and negatively charged regions. For example, the molecule, HCl is made with hydrogen and the very electronegative atom chlorine. The electrons in this molecule tend to gather around the chlorine. This makes the hydrogen end more positively charged (G�� read as partially positive) and the chlorine end more negative (G-, read as partially negative). Each end is a dipole. H Cl

G+ G-

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That means that when two hydrogen chloride molecules get close together and they are oriented correctly, the different dipoles will be attracted to each other.

London Dispersion Forces

There is some type of intermolecular force among all atoms and molecules, even noble gases. These are the weakest of the van der Waals forces and are called the London dispersion forces. London dispersion forces (LDF) are intermolecular attractions resulting from the constant motion of electrons and the creation of instantaneous dipoles. The more electrons there are the greater the London dispersion forces. So it goes to reason that the bigger the atomic mass or molecular mass the greater the London dispersion forces. That is why boiling points generally increase down a group on the periodic table.

Summary of Bonding and their relationship to properties

Ionic (Giant lattice)

+ve ion and -ve ion formed via transfer of electrons held together in giant lattice with

strong electrostatic interaction

Metal + Non-Metal ex: KCl, MgF2, Na2SO4, etc.

Strong Bonds

high m.p./b.p. poor conductors of

electricity when solid (ions not free to move), good

when liquid or in solution (dissolved)

will dissolve in polar solvents

Covalent (Individual molecules)

Small groups of atoms covalently bonded together by sharing electrons

Non-Metal + Non-Metal ex: CO2, PCl5, etc.

Strong bonds within molecules, but weak between molecules

low m.p/b.p., often liquids or gases at RT;

LDF- (induced dipoles) Dipole - perm. dipoles) H-Bonds - (H - N/O/F)

poor conductors

Metallic (Mixtures of Metals)

Close packed array of atoms (ions) with "sea" of free moving electrons

Metals ex: Na, Al, Au, Stainless Steel (Fe/C/Cr),

Bronze (Cu/Sn), Brass (Cu/Zn)

Strong (but flexible) bonds

high m.p/b.p. good conductors of electricity ("sea' of electrons) and heat (close packed), malleable (can be shaped), ductile (drawn into

thin wires), luster (shiny)

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Task 6c

1. Compound B has lower melting and boiling points than compound A. At the same temperature, compound B vaporizes faster than compound A. If one of these compounds is ionic and the other is molecular, which would you expect to be molecular? Ionic? Explain your reasoning.

2. Analyzing Data. The melting points for the compounds Li2S, Rb2S, and K2S are 900oC, 530oC, and 840oC, respectively. List these three compounds in order of increasing lattice energy.

3. Explain why most metals are malleable and ductile but ionic crystals are not.

4. Explain why metals are good electrical conductors.

5. What is the difference between a formula unit and a molecule.

Lewis Structures

In Topic 3, you learned about electron dot diagrams for elements. These can be used to represent molecules also. When representing molecules, they are called Lewis structures. 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. If only the shared pairs (bonds) are written using dashes, then this will be a structural formula. A structural formula indicates the kind, number, arrangement, and bonds but not the unshared pairs of the atoms in a molecule.

It is important to learn how to draw the Lewis structures of molecules and to predict the molecular geometry of the molecule. To predict the geometry, you have to consider shared (bonded) electrons and lone (nonbonded) electron pairs surrounding the central atom. To do this, we will use the VSEPR theory.

VSEPR stands   for   “valence   shell   electron   pair   repulsion.     That  means   that   the   electron   pairs  around the atoms (usually the central atom) repel each other to affect the shape of the molecule. VSEPR theory states that the repulsion between the sets of valence-level electrons surrounding an atom causes these sets to be oriented as far apart as possible. Unshared electrons repel the most. Remember these molecules are 3-D even though we draw them in 2-D.

Drawing Lewis structures

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Here are the steps for writing a Lewis structure for compounds for which you do not know the formula.

1. Calculate the total number of valence electrons taking into account any charges. (add electrons for negative charges and subtract electrons for positive charges)

2. Decide which element is the central atom. Usually this is obvious; if in doubt it will be the least electronegative atom (except for H: it only wants 2 electrons). Put that element in the center and add the other elements around the central atom using lines to represent bonding pairs of electrons.

3. Arrange the remaining electrons to complete the octet of the terminal atoms by placing pairs of dots around the atoms. Place any remaining electrons (dots) on the central atom, if necessary expanding the octet.

4. If the central atom lacks an octet, form multiple bonds (double or triple bonds) by converting non-bonding electrons on terminal atoms into bonding pairs. (Sometimes atoms will remain electron deficient).

For example:

CCl4 Valence electrons = 4 + 7(4) = 32

H2O Valence electrons = 1(2) + 6 = 8

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Now practice with your teacher the following:

NH3

HF

PF5

NH4+

PCl6-

CO32-

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Task 6d

1. Draw a Lewis structure for each of the following molecules. You will come back to this section for later tasks, so be sure to be neat and orderly.

a. SCl2

b. PI3

c. Cl2O

d. NH2Cl

e. SiCl3Br

f. ONCl

g. SO42-

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h. ClO2-

i. BeCl2

You can also use VSEPR theory to predict the shapes of a molecule. Here is a link that might help you understand the shapes.

http://library.thinkquest.org/C006669/data/Chem/bonding/shapes.html

You will need to learn the chart on the next page.

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Task 6e

1. Using the Lewis structures that you drew in 6d, determine the molecular geometry for each.

Hybridization

VSEPR theory helps determine the shapes of a molecule but it does not reveal the relationship between  a  molecule’s  geometry  and  the  orbitals  occupied  by  its  bonding  electrons.    To  explain  how the orbitals of an atom become rearranged when the atom forms covalent bonds, a different model is used. This model is called hybridization, which is the mixing of two or more atomic orbitals of similar energies on the same atom to produce new hybrid atomic orbitals of equal energies.

Methane, CH4, provides a good example of how hybridization is used to explain the geometry of molecular  orbitals.    The  orbital  notation  for  carbon’s  valence  electrons  is  2s2 2p2. We know form experiments that a methane molecule has tetrahedral geometry. How does carbon form four equivalent, tetrahedrally arranged covalent bonds? The one s orbital and the three p orbital hybridize to form four new, identical orbitals called sp3 orbitals.

2p sp3

2s

Carbon’s  orbitals Carbon’s orbitals after Before hybridization sp3 hybridization

Hybrid orbitals are orbitals of equal energy produced by the combination of two or more orbitals on the same atom. The number of hybrid orbitals produced equals the number of orbitals that have combined. A molecule with 3 bonding areas will be sp2, while a molecule with 4 bonding areas will be sp3. Trigonal bipyramidal molecules will be dsp3 and octahedral molecules will be d2sp3.

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Polarity & Dipole forces The strongest intermolecular forces exist between polar molecules. Polar molecules act as tiny dipoles because of their uneven charge distribution. A dipole is created by equal but opposite charges that are separated by a short distance. The direction of  the  dipole  is  from  the  dipole’s  positive pole to its negative pole. A dipole is represented by an arrow with a head pointing toward the negative pole and a crossed tail situated at the positive pole. The negative pole will be the atom that is the most electronegative. For example, hydrochloric acid, HCl:

H__Cl

This is a polar bond and a polar molecule because the charges are unevenly distributed. If the charge distribution is evenly distributed, the molecule will be nonpolar. For example CH4:

H

H __ C __ H

H

Notice that all the dipole point toward the C. Here the charge is evenly distributed. Therefore the molecule is nonpolar even though the individual bonds are polar.

Task 6f

1. Go back to Task 6d and label the hybridization and the polarity of each molecule.