Bonding

What is a bond?

• A bond is an attraction between two atoms that holds them together.

Chemical Bonds and Energy

• The forces holding atoms together involve energy.

Bond Formation

• The process of forming chemical bonds releases energy (exothermic).

A Spring as an Analogy for Bond Formation

• Unbonded atoms that are apart are like a stretched spring. They have potential energy. As they are attracted and come closer to each other, they lose or release potential energy until they are bonded (like a spring at rest position).

Bond Breaking

• It takes energy to break a bond holding atoms together. Bond breaking is an endothermic process.

+ energy

The Covalent Bond

• When atoms share two or more electrons between them, this is called a covalent bond.

The Formation of a Covalent Bond between Hydrogen Atoms

• The negative electrons of a hydrogen atom attract to its own nucleus and also to the nucleus of a neighbor hydrogen atom, bringing the atoms closer.

Covalent Bonding between Hydrogen Atoms

• As hydrogen atoms approach each other, they overlap their orbitals, forming a molecular orbital.

• In the molecular orbital there are two electrons which are shared between the atoms, giving each atom a full level, n=1.

Attracting and Repelling Forces between Hydrogen Atoms

• Attracting forces are found between electrons and their own nucleus as well as the other atom’s nucleus.

• Repelling forces are found between the two nuclei as well as the electrons with other electrons.

Potential Energy is Lost With Bond Formation

At lowest potential energy, atoms have a given bond distance between them. If they

are forced closer, their identical positive nuclear charges force them apart.

Ways of Representing a Covalent Bond

• A covalent bond can be shown as a probability distribution, as a Lewis Dot Diagram or as a line drawn between two atoms (structural formula).

• A single line (covalent bond) always represents two shared electrons (an electron pair) between the atoms.

Valence Electrons

The outermost electrons in an atom are called valence electrons.

Lewis Dot Diagrams

A Lewis Dot Diagram is an element’s symbol with a number of dots around it equal to the number of its valence electrons.

Noble Gas Dot Diagrams

The Noble Gases•The chemical family called the Noble Gases (Inert Gases) have full shells or shells with 8 outer electrons.•All noble gases have 8 valence electrons except helium which has 2 valence electrons.

What is a Covalent Bond?

•Covalent bonds are the bonds formed when atoms share electrons between them. By sharing each other’s electron(s) they get octets or 2 electrons like helium

Sharing Electrons To Get Octets or Two (Like He)

•Covalent bonds are the bonds formed when atoms share electrons between them. By sharing each other’s electron(s) they get octets or 2 electrons (for hydrogen) like helium.

Representing Covalent Bonds

•Covalent bonds can be represented with a Lewis Diagram as two dots between atoms, or as a single line drawn between two atoms. A single line represents an electron pair (two electrons).

Sharing Electrons Creates Molecules

•When atoms share electrons, they form a unit or group called a molecule.

The Kinds of Atoms Forming Covalent Bonds

•Covalent bonds form when a nonmetal combines with another nonmetal. Ex: H with H to form H2

Covalent Compounds : Nonmetal + Nonmetal

Lewis Dot Diagrams Show Covalent Bonds

Often Lewis Dot Diagrams are used to show covalent bonds.

Covalent Compounds and Lewis Diagrams

Carbon’s 4 valence electrons are arranged around the symbol. Each hydrogen shares one valence electron, having 2 like He and C gains 4 by sharing to get an octet.

Sharing More Than One Pair Of Electrons

Two oxygen atoms can each get an outer octet by sharing four electrons (2 pairs). This results in a double bond.

Sharing More Than One Pair Of Electrons

Two nitrogen atoms can each get an outer octet by sharing six electrons (3 pairs). This results in a triple bond.

Examples of Covalent Compounds

Steps in Completing Lewis Dot Diagrams

1. Decide which atoms are bonded.For example, lets draw the Lewis structure for a SO3 molecule.

2. Count all valence electrons.There are a total of 24 valence electrons in SO3 (6 from the sulfur, and 6 each from the 3 oxygens).

3. Place two electrons in each bond.

O

S O

O

Steps in Completing Lewis Dot Diagrams

4. Complete the octets of the atoms attached to the central atom by adding electrons in pairs.

5. Place any remaining electrons on the central atom in pairs.(already has 24, no remaining electrons in this example)

6. If the central atom does not have an octet, form double bonds. If necessary form triple bonds.

O

S O

O

Electronegativity : An Atom’s “Electron-Pull” Power

•Electronegativity is the degree to which an atom attracts electrons.•In the periodic chart, atoms have increasing electronegativity moving to the right and from bottom to top.

Uneven Sharing Due To Different Electronegativities

When a fluorine atom (e. neg.=4) shares electrons with a chlorine atom (e. neg.=3.0), the fluorine pulls the electrons more to itself, making the molecule slightly negative on the fluorine end and slightly positive on the chlorine end.

Uneven Sharing Causes A Polar Covalent Bond

The positive and negative ends of the molecule are called poles. The bond is called a polar covalent bond, one which has uneven sharing of electrons.

Valence for Polar Covalent Compounds I

In polar covalent compounds (uneven sharing of electrons), the element with the greater electronegativity is written second and the element with the lesser electronegativity is written first.

•Example: HF

Valence for Polar Covalent Compounds II

Example: •The most electronegative atom becomes negative (has the shared electrons most of the time) while the least electronegative atom becomes positive (having the shared pair less of the time).

Valence for Polar Covalent Compounds III

The most electronegative atom is assigned a negative valence while the least electronegative atom is assigned a positive valence.

Examples:

1. ClF Cl+F -

2. HF H+F -

3. OF2 O+2F2-1

Valence for Polar Covalent Compounds IV

The valence number assigned to a polar covalent element is the number of electrons that element is sharing while the valence sign indicates if it is stronger (negative) or weaker (positive) in attracting electrons.

Example: compound NF3 (called nitrogen trifluoride)

Valences Assigned: N+3F3-1

Note that this pattern is similar to the ionic compounds like Ca+2Cl2

-1 where the + valence element is written first.

Ionic Bond : Extreme Uneven Electron Sharing

When atoms have an extreme uneven sharing of electrons because their difference in electronegativity is so large, the bond is referred to as an ionic bond and the pair of electrons is virtually transferred from one atom to the other atom.

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Ex: Na + Cl NaCl

Determining Bond Character

The difference in electronegativity of the two elements forming a bond will determine if the bond is covalent, polar covalent or ionic.

Table of Electronegativities (Heath p333)

Sample ProblemWhat is the bond type in CH4 ?

Step 1 : Find the electronegativities of C (2.5) and H (2.1).

Step 2 : Find the difference in these electronegativities (.4)

Step 3 : Assign the bond type based on EN diff. table

Step 3 : This bond would be classified as a nonpolar covalent bond

Sample Problem 2What is the bond type in Li3N ?

Step 1 : Find the electronegativities of Li (1.0) and N (3.0).

Step 2 : Find the difference in these electronegativities (2)

Step 3 : The bond type would be ionic based on EN diff. table

Ionic Compounds and Ionic Bonds• Ionic compounds

are produced when their atoms form ionic bonds. An ionic bond is an extreme polar covalent bond in which the electron is so unevenly shared that it is effectively transferred from one atom to another. If atoms have an electronegativity difference of 1.7 or larger, they will form ionic bonds.

Ionic Solids and Bond Strength• Ionic bonds are very

strong bonds so compounds with these bonds have high melting and boiling points.

• Ionic bonds produce solids in which oppositely-charged ions alternate in a crystal lattice.

• Covalent bonds (polar and nonpolar) produce neutral molecules and molecular substances.

Comparison of Melting Points of Ionic Solids• In the unit on periodicity, it was noted that within a chemical family, the

nucleus of larger atoms (lower in the vertical column) has less ability to hold its outer valence electrons due to their greater distance from the nucleus and the screening effect of inner electrons. Thus larger atoms in a chemical family bond with weaker bonds than smaller atoms in the same family. Thus a compound like LiF has stronger bonds than a compound like NaCl since the nuclei of Li and F have stronger attractions than the nuclei of Na and Cl. This causes the melting point of LiF (845 C) to be higher than the melting point of NaCl (801 C) or KBr (734 C). Likewise, mp BeO: (2570 C) and MgS: (2000)

Elements forming Ionic or Covalent Compounds

•Metals and nonmetals form ionic bonds (substances).•Nonmetals with nonmetals form polar covalent bonds.

Comparison of Bond Lengths• The distance between bonded atoms varies for single

double and triple bonds. The interatomic distance becomes less as the bond goes from single to double to triple bonds.

Resonance (Heath p 347) of SO3

• When drawing Lewis structures for chemicals, often more than one structure can be composed. The question must be asked, “Which one is correct?”.

Resonance

• Since the distance between double bonds is less, bond lengths can be measured experimentally to identify any differences.

• Experimental measurements show that all bonds are identical and intermediate in length between single and double bonds.

Resonance

• Resonance is a concept that tries to explain the intermediate bond length and different resonance structures.

Resonance• The actual molecule is thought to rapidly oscillate (resonate) between

the differing Lewis structures to effectively produce an intermediate or average of the structures. In this case, the extra pair of electrons (extra bond) is rapidly moved between all the oxygen atoms to produce an average, intermediate 1 1/3 bond between the C and O atoms.

Resonance Between The Benzene Structures

• Benzene, C6H6 , has two resonance structures.

• Experimental measurement of bond distances places the actual bonds intermediate between single and double bonds.

Exceptions to the Octet Rule• Elements like Be and B tend

not to get octets of their valence electrons while elements in period three (P,S, Cl) may at times get valence electrons more than an octet.

Bonding Orbitals• When atoms approach each other, the energy they release can

promote paired electrons into unoccupied orbitals. Example is 4Be, showing only valence electrons.

• The orbitals then hybridize, a transformation in which the orbital shapes are blended into new hybrid orbital shapes. Ex: BeF2

Bonding Orbitals• The new hybrid bonding orbitals are named according to the number

and type of atomic orbitals they originated from (sp orbitals in the previous example). Note that there are the same number of orbitals before and after hybridization. The new hybrid orbitals differ only in shape but not in number.

• The example below show B bonding to F. The 5B atom’s ground state is 1s22s22p1 , its promoted configuration is 1s22s12px

12py1 , its

hybridized bonding orbitals are 3 identical orbitals called sp2 sp2 sp2 .

Bonding Orbitals

• Carbon’s four bonding orbitals are called sp3 orbitals (made up from one s hybridized with three p orbitals).

The Spatial Arrangement of Bonding Orbitals• The VSEPR (valence shell electron pair repulsion) theory explains how

bonding orbitals will arrange themselves in space (in three dimensions).

• The VSEPR theory states that bonding orbitals will arrange themselves spatially so as to obtain the greatest distance away from each other (largest bond angle) since they are all repelling each other (like negative charges of electrons in the orbitals).

Period Three Bonding• Phosphorus bonds with hydrogen either as PH3 or as PH5. In the

molecule of PH, P promotes an s electron into an unfilled d orbital on the same principal energy level.

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When Polar Bonds Make Polar Molecules• If the polarity of a dipole bond is balanced by one or more

other dipole bonds, there will be no net polarity on the molecule. A molecule’s 3-D geometry determines if the molecule is polar or not when it has polar bonds. The small Greek letter delta, δ, is used to denote poles in bonds and on molecules.

When Polar Bonds Make Polar Molecules• If the polarity of a dipole bond is balanced by one or more

other dipole bonds, there will be no net polarity on the molecule. A molecule’s 3-D geometry determines if the molecule is polar or not when it has polar bonds.

When Polar Bonds Make Polar Molecules• If the polarity of a dipole bond is balanced by one or more

other dipole bonds, there will be no net polarity on the molecule. A molecule’s 3-D geometry determines if the molecule is polar or not when it has polar bonds.

When Polar Bonds Make Polar Molecules• If the polarity of a dipole bond is balanced by one or more

other dipole bonds, there will be no net polarity on the molecule. A molecule’s 3-D geometry determines if the molecule is polar or not when it has polar bonds.

Dipole Molecular Behaviour• Dipoles align themselves in electric fields• Dipoles align themselves with respect to each other (form

attractions between molecules which raises boiling and melting points).

• Dipoles align themselves with respect to ions. Dipole liquids will dissolve ionic solids.

Solvation (Dissolving) of an Ionic Solid by a Polar Solvent

• Polar solvent molecules orient their opposite poles to charged ions. As polar solvent molecules bond to an ion, they weaken that ion’s charges for its neighbor ions, freeing it from the ionic crystal lattice. In solution, each ion has a cohort of solvent molecules surrounding it.

Hydrogen Bonds

• A slightly positive hydrogen on one molecule will bond to an electronegative oxygen, nitrogen or fluorine one another molecule (or another part of a very large molecule.)

Hydrogen Bonds• A slightly positive hydrogen on one molecule will bond to

an electronegative oxygen, nitrogen or fluorine one another molecule (or another part of a very large molecule.)

Hydrogen Bonds• Hydrogen bonds are weak bonds but many of them

together can form strong attractions. Wood and paper are composed of cellulose fibers which are hydrogen-bonded to each other.

Making Paper• A pulp fiber source (trees, recycled paper etc.) is ground

up and made into a slurry in water. The slurry is deposited onto screens and the water is pressed out. The damp pulp is then dried. Commercially, the pulp is pulled on rollers as it dries which gives commercial papers a fiber bias that makes them tear more easily along one axis than the other.

Making Paper• As paper dries, water is removed between cellulose fibers and

they form hydrogen bonds. Artist papers and papers for money are deposited on screens and just pressed (not pulled). These papers have no bias and are therefore stronger.

Forces Between Nonpolar Molecules

• Nonpolar molecules do not attract each other and as a result, substances with these molecules are gases with low boiling and melting points. Fluorine (F2) has a mp of -219.6 celsius and a bp of -188.1 celsius. Methane (CH4) has a mp of -183 and a bp of -162 celsius.

Forces Between Nonpolar Molecules

• At very low temperatures (very little molecular motion), nonpolar gases do form liquids and solids. This could only happen if some forces of attraction exist at very low temperatures.

Momentary Dipoles and London Forces• At very low temperatures, momentary uneven electron

distributions in one molecule/atom (in a split second) causes a momentary dipole. This momentary dipole induces neighboring molecules to develop similar dipoles.

Momentary Dipoles and London Forces• These dipoles oscillate back and forth numerous times per

second creating opposite charges that weakly attract neighboring molecules so they form liquids and solids. These weak forces due to momentary, oscillating dipoles are called London Forces. These forces only develop when nonpolar molecules are relatively still and/or relatively close (at very low temperatures).

van der Waals Forces

• Dipole-dipole forces (steady poles on molecules) and London Forces (momentary oscillating dipoles) are collectively known as van der Waals forces.

London Forces in Carbon Chains• The longer a carbon chain gets, the more temporary

dipoles that can set up along the chain. With more momentary dipoles in longer chains, the London Forces increase which leads to higher melting and boiling points. Propane, a three carbon chain, has a mp of -190 and a bp of -45 celsius. Hexane, a six carbon chain has a mp of -95 and a bp of 69 celsius.

Network Covalent Bonds• In some substances, covalent bonds link atoms

throughout the substance. This is called a network solid. In network solids, electrons are shared between atoms so that no ions are formed, nor are the electrons free to move except narrowly between adjacent atoms within the covalent bonds.

Network Covalent Bonds• Quartz and glass (silicon dioxide) and carborundum (SiC –

grit in sandpaper) are network solids. Diamond, the hardest solid on earth, is made up just carbon. Network solids have the strongest bonds, followed by ionic solids.

Allotropes: Different Forms of the Same Element

• A number of elements have more than one form. These forms result from different connections of the same kind of atom. Phosphorus, oxygen, sulfur and carbon all have different allotropes. Allotropes of phosphorus are shown.

Allotropes of Carbon: Diamond and Graphite• Two allotropes of carbon are diamond and graphite. In

graphite, flat layers are weakly linked by pi bonds (delocalized electrons), held together by London forces. The weak bonds and delocalized electrons allow flat layers to slide past each other (making graphite powder a good lubricant) and also allow electric current (electrons) to flow through it.

Metallic Bonds• Metallic solids are composed of positive metal ions

surrounded by a “sea” of valence electrons donated by all the positive ions. These electrons freely wander throughout the metal. These communally shared electrons hold the metal ions together.

Metallic Solid Properties• The “sea” of valence electrons in metals helps explain

their physical properties. These communal electrons easily move when a voltage difference is applied across the metal (conduct electricity). These electrons likewise freely convey heat, extra kinetic vibration. Metals exhibit malleability (can be re-formed) since their ions are embedded in a fluid “electron sea”.

Comparison of Various Bond Strengths• Network solid (covalent bonds) > ionic solid (ionic bonds)

> metallic solid (valence electron sea) > hydrogen bond > dipole-dipole (fixed dipoles) > London forces (momentary dipoles at low temperatures)

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