Concepts of Chemical Bonding - Spada UNS

Concepts of Chemical

Bonding

Chemical Bonds

Three types:

– Ionic

Electrostatic attraction between ions

Covalent

Sharing of electrons

Metallic

Metal atoms bonded to several other atoms

Ionic Bonding

When a metal and a non-metal get together

is a chemical bond formed by the electrostatic attraction between positive and negative ions.

An ionic bond forms when one or more electrons are transferred from the valence shell of one atom to the valence shell of another atom.

Na ([Ne]3s¹) + Cl ([Ne]3s²3p⁵) →

Na⁺ ([Ne]) + Cl^- ([Ne]3s²3p⁶) The atom that transferred the electron(s) becomes a cation.

The atom that gained the electron(s) becomes an anion.

A Lewis electron-dot symbol is a notation in

which the electrons in the valence shell of an atom or ion are represented by dots placed around the chemical symbol of the element.

Note: Dots are placed one to a side, until all four sides are occupied.

Table 9.1 illustrates the Lewis electron-dot symbols for second- and third-period atoms.

? Represent the transfer of electrons in forming calcium oxide, CaO, from

atoms.

+ Ca²⁺

Ca O +

[

O

]

^2-

Let’s look next at the energy involved in forming ionic compounds.

The energy to remove an electron is the ionization energy.

The energy to add an electron is the electron affinity.

The combination of ionization energy and electron affinity is still endothermic; the process requires energy.

However, when the two ions bond, more than enough energy is released, making the overall process exothermic.

The lattice energy is the change in energy that occurs when an ionic solid is separated into gas- phase ions.

It is very difficult to measure lattice energy directly.

It can be found, however, by using the energy changes for steps that give the same result.

Energetics of Ionic Bonding

it takes 495 kJ/mol to remove 1 electron from sodium.

495x2 = 990 kJ/2 Na

2Na(s) + Cl₂(g) ---> 2NaCl(s)

Energetics of Ionic Bonding

We get 349 kJ/mol Cl back by

giving 1 electron to each to 1 mole of Cl₂.

-349x2 = -700 kJ/mol Cl₂

990 kJ/2Na – 700 kJ/Mol Cl₂ = 290kJ

Energetics of Ionic Bonding

• But these numbers don’t explain why the reaction of

sodium metal and chlorine gas to

form sodium chloride is so exothermic!

990 kJ/2Na – 700 kJ/Mol Cl₂ = 290kJ

Energetics of Ionic Bonding

• There must be a third piece to the puzzle….

• The electrostatic attraction

• Between Na⁺ and Cl^- .

• The ionic Bond!

Ionic bonding

Ionic bonding involves 3 steps (3 energies) 1) loss of an electron(s) by one element,

2) gain of electron(s) by a second element, 3) attraction between positive and negative

Na Ionization energy e

^–

+ Na

⁺

Cl + e

^–

Electron affinity Cl

^–

Lattice energy

•Note that although we

represent this as a three step process it actual

Lattice Energy

• This third piece of the puzzle is the lattice energy:

The energy required to completely separate a mole of a solid ionic compound into its gaseous ions.

• The energy associated with electrostatic interactions is governed by Coulomb ’ s law:

E

_el

=  Q

¹

Q

²

d

Lattice Energy

• Lattice energy, then, increases with the charge on the ions.

• It also increases with decreasing size of ions.

For example, to find the lattice energy for NaCl, we can use the following steps.

The process of finding the lattice energy indirectly from other thermochemical reactions is called the Born–Haber cycle.

Energetics of Ionic Bonding

By accounting for all three energies

(ionization energy, electron affinity, and lattice energy), we can get a good idea of the energetics

involved in such a process.

Na(s) + 1/2Cl₂(g) ---> NaCl(s)

Energetics of Ionic Bonding

• These phenomena also help explain the

“octet rule.”

• Elements tend to lose or gain electrons once they attain a noble gas configuration because energy would be expended that cannot be overcome by lattice energies.

Exercises

Lattice Energy

Ionic substances are typically high-melting solids.

There are two factors that affect the strength of the ionic bond. They are given by Coulomb’s law:

The higher the ionic charge, the stronger the force;

the smaller the ion, the stronger the force.

2 2 1

r

Q

F = kQ

Based on this relationship, we can predict the relative melting points of NaCl and MgO.

The charge on the ions of MgO is double the

charge on the ions of NaCl. Because the charge is double, the force will be four times stronger.

The size of Na⁺ is larger than that of Mg²⁺; the size of Cl^- is larger than that of O^2-. Because the

distance between Mg²⁺ and O^2- is smaller than the distance between Na⁺ and Cl^-, the force between Mg²⁺ and O^2- will be greater.

Based on the higher charge and the smaller

distance for MgO, its melting point of MgO should be significantly higher than the melting point of

NaCl.

The actual melting point of NaCl is 801°C; that for MgO is 2800°C.

When we examine the electron configuration of main-group ions, we find that each element gains or loses electrons to attain a noble-gas

configuration.

? Give the electron configuration and the Lewis symbol for the chloride ion, Cl-.

Cl

]

^-

[

For chlorine, Cl, Z = 17, so the Cl^- ion has 18 electrons. The electron configuration for Cl^- is

1s² 2s² 2p⁶ 3s² 3p⁶ The Lewis symbol for Cl^- is

Group IIIA to VA metals often exhibit two different ionic charges: one that is equal to the group

number and one that is 2 less than the group number.

The higher charge is due to the loss of both the s subshell electrons and the p subshells electron(s).

The lower charge is due to the loss of only the p subshell electron(s).

For example, in Group IVA, tin and lead each form both +4 and +2 ions. In Group VA, bismuth forms +5 and +3 ions.

Polyatomic ions are atoms held together by

covalent bonds as a group and that, as a group, have gained or lost one or more electron.

Transition metals form several ions.

The atoms generally lose the ns electrons before losing the (n – 1)d electrons.

As a result, one of the ions transition metals generally form is the +2 ion.

? Give the electron configurations of Mn and Mn2+.

Manganese, Z = 25, has 25 electrons;. Its electron configuration is

1s² 2s² 2p⁶ 3s² 3p⁶3d⁵ 4s²

Mn²⁺ has 23 electrons. When ionized, Mn loses

the 4s electrons first; the electron configuration for Mn²⁺ is

1s² 2s² 2p⁶ 3s²3p⁶ 3d⁵

Ionic radius is a measure of the size of the spherical region around the nucleus of an ion within which the electrons are most likely to be found. While ionic radius, like atomic radius, can be somewhat arbitrary, it can be measured in ionic compounds.

A cation is always smaller than its neutral atom.

An anion is always larger than its neutral

The term isoelectronic refers to different species having the same number and configuration of

electrons.

For example, Ne, Na⁺, and F^- are isoelectronic.

Ionic radius for an isoelectronic series decreases with increasing atomic number.

? Using the periodic table only, arrange the following ions in order of increasing ionic radius: Br-, Se2-, Sr2+.

35

Br

34

Se

38

Sr

These ions are isoelectronic,

so their size decreases with increasing atomic number:

Sr²⁺ < Br^- < Se^2-

Covalent Bonding

• In these bonds atoms share electrons.

• The electrons that can be shared are the Valence electrons.

What happens when nonmetals get together

Covalent compounds

• Covalent compounds are formed when non-metal atoms react together.

• As these atoms come near their outer electrons are attracted to the nucleus of both atoms and become shared by the atoms.

• The shared electrons count towards the shells of both atoms and therefore help fill up incomplete electron shells.

Covalent bonds

• Covalent compounds are held together by this sharing of electrons.

• A pair of electrons shared in this way is known as a covalent bond.

• It is sometimes represented in full bonding diagrams (see figure 1). Often these bonds are just shown as a pair of electrons (xx) or even just a line (see figure 2).

F

^X_X

F F - F

Figure 1 Figure 2

Small covalent structures

• Sometimes just a few atoms join together in this way.

• This produces small covalent molecules – often known as simple molecular structures.

a simple molecular structure

covalent bonds

Giant covalent structures

• Sometimes millions of atoms are joined together by covalent bonds.

• This produces a rigid 3-D network called a giant lattice.

a giant lattice

Covalent bonding in chlorine

Chlorine (2.8.7) needs 1 more electron to attain a full electron shell.

Cl

(2,8,7)

Cl

(^2,8,7)

Cl Cl

Cl-Cl

Both fluorine and chlorine needs 1 more electron to attain a full electron shell.

Cl

(2,8,7)

F

(^2,7)

Copy this diagram and add the electron

arrangements that could exist in fluorine chloride (FCl).

F Cl

Covalent bonding in hydrogen chloride

Both hydrogen (1) and chlorine (2.8.7) needs 1 more electron to attain a full outer shell.

H

(2)

Cl

(2,8,8)

Cl

H-Cl

(2,8,7)

H

(1)

Covalent bonding in water

Hydrogen (1) needs 1 more electron but oxygen (2.6) needs 2 more. Therefore, we need 2 hydrogens.

O H

H

O H

H

O H

H

• Hydrogen (1) needs 1 more electron.

• How many does nitrogen (2.5) need?

• How many hydrogens per 1 nitrogen?

• Draw bonding diagrams for ammonia.

N H

H

H N H

H H

3 3

• Hydrogen (1) needs 1 more electron.

• How many does carbon (2.4) need?

• How many hydrogens per 1 carbon?

• Draw bonding diagrams for methane.

4 4

C H

H H

H

C H

H H

H

H H

O O

H H

O O

• Copy the atoms below.

• Complete the diagram showing how each atom can achieve full shells.

Covalent bonding - multiple bonds

• Mostly electrons are shared as pairs.

• There are some compounds where they are shared in fours or even sixes.

• This gives rise to single, double and triple covalent bonds.

• Again, each pair of electrons is often represented by a single line when doing simple diagrams of molecules.

Cl-Cl

Single bond

O=O

Double bond

N=N

Triple bond

Covalent bonding in oxygen

Oxygen (2.8.6) needs 2 more electrons to attain a full electron shell.

O O

O=O

O O

4 electrons

Nitrogen (2.8.5) needs 3 more electrons to attain a full electron shell and forms a triple bond.

Draw a bonding diagram of nitrogen.

6 electrons

N N

N N

N=N

1. Hydrogen fluoride (HF) 2. Hydrogen sulphide (H₂S)

3. Ethane (C₂H₆ and the carbons are joined by a single covalent bond)

4. Carbon dioxide (CO₂ and the carbon oxygen bonds are double bonds)

H F

H H

H

C C H O C O

Draw ‘dot and cross’ type bonding diagrams for each of the following:

Giant covalent structures

1. Carbon atoms form giant structures.

2. What is interesting is that there is more than one possible arrangement for the atoms.

3. Although this does not affect the chemical properties it can make a huge difference to the physical

properties such as hardness, slipperiness, melting point and density.

Different arrangements of the same element are called allotropes.

C

Giant covalent structures: diamond

• One form of carbon is diamond.

• Each diamond

consists of millions of carbon atoms

bonded into a single giant structure.

• It is very hard.

Diamond

strong covalent

bonds

carbon atoms

Giant covalent structures: graphite

• A more common form of carbon is graphite.

• Millions of carbon atoms are bonded into a giant structure but within this

structure the layers are only weakly

joined.

Graphite _covalent^strong

bonds

carbon atoms

weak

attraction

Giant covalent structures: carbon footballs!

• During the last 20 years new forms of carbon have been

discovered some of which have

“closed cage” arrangements of the atoms.

• These are large but are not really giant molecules.

One of them contains 60 carbon atoms and bears remarkable

similarities to a football!

Giant covalent structures: sand

• Sand is an impure form of silicon dioxide.

• Although it is a

compound, it has a

giant covalent structure with certain similarities to diamond.

silicon atoms oxygen atoms

Covalent Bonding

• There are several

electrostatic interactions in these bonds:

– Attractions between electrons and nuclei

– Repulsions between electrons – Repulsions between nuclei

• Covalent bond, sharing electrons,

• But electron sharing not always equal.

• Fluorine pulls harder on the shared electrons than hydrogen does.

• Therefore, the fluorine end has more electron density than the hydrogen end.

• But how do you know who pulls hardest?

Electronegativity:

Developed 1^st by Linus Pauling like this:

H-F -→ H + F >> H-H -→ H + H or F-F → F + F

In other words, H-F bond much stronger than H-H or F-F bond.

Why?

Because there is an ionic component to attraction in H- F

F more – and H more + so the ionic component makes bond stronger.

Electronegativity:

• The ability of atoms in a molecule to attract electrons to itself.

• On the periodic table, electronegativity increases as you go…

E . N

_A

- E . N

_B

= eV

^{-1/ 2}

D

_AB

- (D

_AA

+ D

_BB

) /2

The dissociation energy of the A-B bond versus the A-A And B-B bond gives Pauling electronegativity

A measure of how much an atom attracts electrons when It is in a molecule.

Developed 1^st by Linus Pauling like this:

Refinements have occurred since, but this is pretty close.

Electronegativity:

• On the periodic chart, electronegativity increases as you go…

Polar Covalent Bonds

• When two atoms share

electrons unequally, a bond dipole results.

• The dipole moment,



,

produced by two equal but opposite charges separated by a distance, r, is calculated:



= Qr

• It is measured in debyes (D).

Polar Covalent Bonds

The greater the difference in

electronegativity, the more polar is the bond.

Lewis symbols

• A convenient way to keep track of the

valence electrons in an atom or molecule

• Lewis dot symbol

Each dot is one valence electron

• Lewis structures for 16 elements

• It is rare to use Lewis pictures for other

elements (transition metals, etc.)

One dot = 1 electron

Lewis Structures

Diagrams for bonding in molecules

Lewis structures are representations of molecules showing all valence electrons, bonding and nonbonding.

Lines correspond to 2 electrons in bond

Covalent Bonding

• A bond where electrons from each atom are shared

• Each covalent bond has 2 electrons that are shared.

• Only the Valence electrons are involved in these covalent bonds.

• Why? Why don’t electrons in n-1 levels get shared?

1. Find the sum of

valence electrons of all atoms in the

polyatomic ion or molecule.

– If it is an anion, add one electron for each

negative charge.

– If it is a cation, subtract one electron for each positive charge.

PCl 3

5 + 3(7) = 26

Lewis Structures

A way to keep track of those valence

electrons

Lewis Structures

A way to keep track of those valence electrons

2. The central atom is the least

electronegative element that isn’t hydrogen (why?).

Connect the outer atoms to it by single bonds.

Keep track of the electrons:

Writing Lewis Structures

3. Put eight electrons around the outer atoms (“fill their octet”)

Keep track of the electrons:

26 − 6 = 20 −

Writing Lewis Structures

4. Fill the octet of the central atom.

Keep track of the electrons:

Writing Lewis Structures

5. If you run out of

electrons before the central atom has an octet…

…form multiple bonds until it does.

Writing Lewis Structures

• Then assign formal charges.

– For each atom, count the electrons in lone pairs and half the electrons it shares with other atoms.

– Subtract that from the number of valence electrons for that atom: The difference is its formal charge.

Writing Lewis Structures

• The best Lewis structure…

– …is the one with the fewest charges.

– …puts a negative charge on the most electronegative atom.

-2 0 +1 -1 0 0 0 0 -1

Exceptions to the Octet Rule

• There are three types of ions or

molecules that do not follow the octet rule:

– Ions or molecules with an odd number of electrons.

– Ions or molecules with less than an octet.

– Ions or molecules with more than eight valence electrons (an expanded octet).

Odd Number of Electrons

Though relatively rare and usually quite unstable and reactive, there are ions

and molecules with an odd number of

electrons.

Odd Number of Electrons

• Example: NO N O ..

.. ..

.

What ’ s nitric oxide good for?

N O ..

.. .

- +

..

Fewer Than Eight Electrons

Draw the Lewis structure for BF

₃

:

Fewer Than Eight Electrons

• Consider BF₃:

– Giving boron a filled octet places a negative charge on the boron and a positive charge on fluorine.

– This would not be an accurate picture of the

Fewer Than Eight Electrons

Therefore, structures that put a double bond between boron and fluorine are much less important than the one that leaves boron with only 6 valence electrons.

Double bonds to halogens don’t happen.

+

- - -

+ +

Fewer Than Eight Electrons

The lesson is: If filling the octet of the central atom results in a negative charge on the

central atom and a positive charge on the

more electronegative outer atom, don’t fill the octet of the central atom.

+

+ +

- - -

More Than Eight Electrons

Draw the Lewis structure for PCl

₅

More Than Eight Electrons

• The only way PCl₅ can exist is if phosphorus has 10 electrons

around it.

• atoms on the 3rd row or below can go over an octet of electrons

– Presumably d orbitals in these atoms participate in bonding.

More Than Eight Electrons

• Draw the Lewis structure for phosphate

• PO₄^-3

More Than Eight Electrons

Even though we can draw a Lewis structure for the phosphate ion that has only 8 electrons around the central phosphorus, a common Lewis structure puts a double bond between the phosphorus and one of the oxygens.

More Than Eight Electrons

• This eliminates the charge on the phosphorus and the charge on one of the oxygens.

• The lesson is: When the central atom is on the 3rd row or below and expanding its octet

eliminates some formal charges, you can do so.

• But, you don’t have to.

More Practice

• Draw lewis structures for:

• SO₄^-2, CO₃^-2, CHCl₃, CN₃H₆⁺ (H’s are

attached to the N’s). SO₂, PO₃^3-, NO₂^-1, BrO₃^-, ClO₄^-,

Covalent Bond Strength

• The strength of a bond is measured by

determining how much energy is required to break the bond.

• This is the bond enthalpy.

• The bond enthalpy for a Cl—Cl bond,

H = 242 kJ/mol

Average Bond Enthalpies

• Average bond enthalpies are positive, because bond

Average Bond Enthalpies

NOTE: These are average bond enthalpies, not absolute bond

enthalpies; the C—H bonds in methane, CH₄, will be a bit different than the C—H bond in

chloroform, CHCl .

Enthalpies of Reaction

• Can use bond enthalpies to estimate H for a reaction

H_rxn = (bond enthalpies of bonds broken) −

(bond enthalpies of bonds formed)

This is a fundamental idea in chemical reactions. The

Enthalpies of Reaction

CH₄^(g) + Cl₂^(g) ⎯⎯→

CH₃Cl^(g) + HCl^(g) In this example, one

C—H bond and one

Cl—Cl bond are broken;

one C—Cl and one H—Cl bond are formed.

Enthalpies of Reaction

So,

H_rxn = [D(C—H) + D(Cl—Cl) − [D(C—Cl) + D(H—Cl)

= [(413 kJ) + (242 kJ)] − [(328 kJ) + (431 kJ)]

= (655 kJ) − (759 kJ)

= −104 kJ

CH₄(g) + Cl₂(g) ⎯⎯→ CH₃Cl(g) + HCl(g)

Bond Enthalpy and Bond Length

• We can also measure an average bond length for different bond types.

• As the number of bonds between two atoms increases, the bond length decreases.

Covalent bonding and electron structures

• The driving force for covalent bonding is

again the attainment of outer electron shells that are completely full.

• This is achieved by sharing electrons where the shared electrons count towards the outer shells of both atoms.

• Sometimes this is achieved with equal numbers of each type of atom.

Sometimes it is not!

Cl Cl C

H

H H

H

N

H

H

H Cl ^H

METALLIC BONDING

Metallic bonding

• Metal atoms form a giant lattice similar to ionic compounds.

• The outermost electrons on each metal are free to move throughout the structure and form a

“sea of electrons”.

• Having released electrons into this “sea” the metal atoms are left with a + charge.

Metallic bonding is the attraction of + metal ions for

the “sea of electrons.”

BONDING AND PHYSICAL

PROPERTIES

Bonding and physical properties

These are things such as:

• Density

• Conductivity

• Malleability/ brittleness

• Melting point

The type of structure that substances have has a huge effect upon physical properties.

The next few slides illustrate just a few of the general patterns.

• Ionic compounds are very brittle.

• Opposite charges attract, so

neighbouring ions are pulled together.

• When something hits the substance a layer of ions will be pushed so that they are next to ions with the same charge.

Attraction becomes:

+ - - +

+

- -

+ +

- + +

-

- -

-

+ + + - - +

+

- -

+ +

+ - - +

+

- -

+ +

Blow

+ - - +

+

- -

+ +

repulsion!

Bonding and physical properties

• Metals are not brittle.

• The metal atoms are the same and exist in simple structures.

• If something hits the substance, it simply moves to the next layer along.

Blow

Bonding and physical properties

• Covalent substances do not conduct electricity.

• This is because in covalent substances the outer

electrons are fixed (localised) between specific atoms.

• Metals conduct electricity.

• In metals the electrons can, given a potential, move anywhere throughout the structure.

H H

H H

H

C C H

electrons fixed in covalent

bonds

electrons free to move

Bonding and physical properties

• Ionic substances do not conduct electricity as solids.

• When molten or dissolved they will conduct (and also undergo electrolysis).

• This is because the electricity is carried through the solution by the ions which are free to move when the ionic compound is molten or in solution.

+ - - +

+

- -

+ +

- + +

-

- -

-

+ + - +

- +

+

- -

+ +

Solid – not free to move

-

+ +

- ⁺-

Molten – mobile

Bonding and physical properties

• Generally substances with giant structures have high melting points and boiling points.

• Small molecules have melting points and boiling points that increase as the size of the molecule increases.

+ - - +

+

- -

+ +

- + +

-

- -

-

+ + - +

- +

+

- -

+ +

In giant structures all the atoms are tightly bonded together.

Usually they are high melting-point solids.

Small molecules tend to be gas, liquid solids with low melting points.

weak forces between

Bonding and physical properties

• Generally substances with giant structures do not dissolve easily (although many ionic compounds dissolve in water for a special reason).

• Again this is because in giant structures separating the particles involves breaking chemical bonds.

Small molecules usually dissolve in a range of solvents. We just

separate one molecule from another.

weak forces

+ - - +

+

- - - + +

- - ⁺

+ - - +

+

-

Giant structures generally don’t dissolve easily.

strong bonds between the

atoms/ions

Bonding and physical properties

• The density of substances depends upon how closely the atoms are packed together.

• Giant structures, metals especially, tend to be dense because all atoms/ions are pulled tightly together.

• Small molecules often have lower densities.

Small molecules tend to have low densities because of space wasted between the molecules.

weak forces between

+ - - +

+

- -

+ +

- + +

-

- -

-

+ + - +

- +

+

- -

+ +

Giant structures generally have high densities.

atoms / ions held closely together

Bonding and physical properties

Copy the Table and fill in the blank columns.

Activity

Which of the following will have covalent bonding?

A. Sodium chloride B. Iron

C. Bronze

D. Nitrogen dioxide

Which of the following will have metallic bonding?

A. Copper chloride B. Graphite

C. Bronze

D. Phosphorus chloride

Which is a true statement about covalent bonds?

A. Usually formed between metals and non- metals

B. Involve transfer of electrons between atoms.

C. Form full electron shells by sharing of electrons.

D. Always involve 2 electrons per atom.

Which of the following exists as a giant molecular structure?

A. Water

B. Carbon dioxide C. Sodium chloride D. Diamond

What will be the formula of the compound formed by hydrogen and sulphur?

A.HS B.H

₂

S C.HS

₂

D.H

₂

S

₂

32

S

16 1

H

1

1 2.8.6

Which of these will conduct as both solid and liquid?

A. metal B. ionic

C. small molecules D. giant molecules

Which of these will conduct when liquid but not when solid?

A. small molecules B. giant molecules C. metal

D. ionic

Which of these will dissolve in solvents like petrol?

A. small molecules B. giant molecules C. metal

D. ionic

Which of these will not conduct at all and is hard to melt?

A. small molecules B. giant molecules C. metal

D. ionic