Unit III: Bonding

Textbook Chapters 5,6 & 11

materials held together by the simultaneous attraction of electrons to two nuclei

What electrons?

How attracted? 2 ways!

Bonding

Chemical Energy -a form of potential energy (PE) PE- energy stored in

molecules Substances possess

energy (chemical) due to their composition and structure

Energy Changes in Bonding

When a chemical bond is formed, energy is released.

When a chemical bond is broken, energy is absorbed.

When 2 atoms are held together by a chemical bond, generally at a lower energy condition than when separated.

Bonding and Stability Generally, systems at

low energy levels are more stable than at high energy levels.

So bonding will more often occur if change leads to lower energy condition (MORE STABLE).

Bond is formed -more energy released means more stable, stronger bond

-less energy given off means less stable weaker bond

Is Oxygen complete? Forms bond, releases energy

exothermic reaction Later, breaks bond and

decomposes endothermic rxn (reverse of exothermic)

Exact same amount of energy required for reaction either released or absorbed (3351 kiloJoules)

2Al2O3(s) 4Al(s) + 3O2(g) +3351kJ

4Al(s) + 3O2(g) 2Al2O3(s) -3351kJ

*from Ref Table I

Electronegativity Measure of the ability

of an atom to attract the electrons that form a bond between it and another atom.

Highest electronegativity is 4.0

Reference Table S

Can use electronegativities to predict type of bonds formed.

Bonds Between Atoms Electrons (valence)

involved in bond formation can be transferred, shared equally or shared unequally between 2 atoms.

Why Bond? Become more like

“inert” or noble gases by completing the valence shell

Most stable

Ionic Bonds aka Electrovalent/Electrostatic Transfer of one or more electrons from metals to nonmetals form ions ions attract (+ and -) {electrostatic force} ionic bond formed

Ionic Bonds continued Form between

elements with electronegativity difference 1.7 with few exceptions.

*Remember metals with nonmetals.

Ionic bonds may form between monatomic or polyatomic ions.

Monatomic ions one atom with charge Na+, F-, Al+3

Polyatomic ions compound of 2 or

more covalently bonded atoms with a charge.

OH-, NH4+, S2O3

-2

Reference Table E.

AgNO3(aq) + NaCl(aq) --> AgCl(aq) + NaNO3(aq)

What is polyatomic ion in above reaction?

What happens to compounds above in water?

Lattice (Binding) Energyof an Ionic Solid

Measure of the energy required to completely separate a mole of a solid ionic compound into its separate ions.

The higher the lattice energy, the stronger the ionic bond.

Taken from mikeblaber.org on 7/27/11.

Ionic Solids Structural unit: made up of ions. Ionic bonding High melting points-->strong forces Geometric structure, ions held in crystal lattice by

electrostatic attraction Do not conduct electricity. When melted or dissolved in water, crystal lattice

is destroyed and ions move freely allowing for electrical conductivity.

Brittle Ex: NaCl, KClO3,MgO, KBr, Li2SO4

Covalent Bonds Simultaneous attraction of 2 nuclei for the same electrons resulting in the sharing of those electrons.

Difference in electronegativities is less than 1.7(some exceptions).

Bonding Continuum

Building Molecules with Lewis Dot Structures

Nonpolar Covalent Bonds Electrons shared

equally between atoms of the same element.

Ex: Diatomics “identical twins”

H2,N2,O2,F2,Cl2,Br2,I2

Difference in electronegativities is ZERO.

Diatomics can have single, double or triple covalent bonds.

Structures are symmetrical/nonpolar.Building Molecules with Lewis Dot Structures

Sigma (s) and Pi (p) Bonds•Single bonds are sigma bonds, electron density is concentrated along the line that represents the bond joining the two atoms.(overlapping s orbitals)•Double bonds contain one sigma and one pi bond. A pi bond occurs when the electron density is concentrated above and below the line that represents the bond joining the two atoms. (overlapping p orbitals)•Triple bonds are one sigma and two pi bonds.

Polar Covalent Bonds Electrons are shared unequally between atoms of

different elements.

Polar Covalent bonds can help create both polar(asymmetrical) and non-polar(symmetrical) structures.

Remember that the words polar and nonpolar can be used to describe both bonds and overall symmetry of molecules!

Polar Molecules due to either bonding or symmetry

Types of Molecular Shape that influence Overall Symmetry

Molecule Bonding Diagram Shape Symmetry of Molecule

Water polar   bentasymmetrical

(polar)

Carbon dioxide polar   linear

symmetrical (non-polar)

Methane polar   tetrahedralsymmetrical (non-polar)

3D image

3D image

Resonance Structures A hybrid of the possible drawings because

no one Lewis structure can represent the situation.

Taken from sv.wikipedia.org on 7/27/11.

Taken from en.wikipedia.org on 7/27/11.

The VSPER Model-molecule will assume the shape that most minimizes electron pair repulsions

Total number of single bonds, double bonds, and lone pairs on the central atom

Structural pair geometry Shape

2 Linear

3 Trigonal Planar

4 Tetrahedral

5 Trigonal Bipyramidal

6 Octahedral

planar

VSPER is Valence Shell Electron Pair Repulsion.

Coordinate Covalent Bonds When 2 shared electrons forming covalent bond are both

donated by one of the atoms “Free loader” Once formed, same as ordinary covalent bond Often involved with forming polyatomic ions

Another Coordinate Covalent Bonding Example

Bonding by Jarod Gagnon

ElectronegativityIt is the ability

of an atom to takesome electrons to make

it’s outer shell stableBut how is it able?

What is electronegatvity for?Sometimes the difference is 1.7 or more

That means one atom is so much strongerit takes the other’s electrons who has them no longer.

If the number is lower than thatthe electrons just share and sit where they sat

Then a covalent bond has begun,but that isn’t the end of the fun.

Sometimes a freeloader comes to the tableand shares some electrons so it can be stable.

The previous compound had no electrons to lack,which is good for the freeloader gives nothing back.

Covalent Bonds form molecules. Molecule-discrete particle formed by covalently

bonded atoms where atoms share electrons so that final electron configuration of each atom is similar to an inert gas

Examples: O2, HCl, H2O, CH4, NH3, C6H12O6, CO2

Molecules form molecular substances. Molecular substances may be gases, liquids or

solids depending on attraction that exists between the molecules.

Just as we did for Ionic Bonding, let’s compare properties & characteristics of Molecular substances as solids!

Covalent or Molecular Solids Structural unit: made

of molecules (covalent bonds)

Low melting points Relatively weak forces Soft Poor heat conductors Do not conduct

electricity (good electrical insulators)

Ex: I2, H2O, CO2, cellulose C5H10O5

Predicting Bonds and Structure?

Building Molecules with Lewis Dot Structures

Network Bonding Certain solids consist of covalently bonded atoms

linked in a network that extends throughout sample with an absence of discrete particles

“One Big Giant” Molecule

Just as we did for Ionic & Covalent bonding, let’s compare properties & characteristics of Network solids!

Network Solids Structural unit: made

up of atoms, “One Giant Molecule”

Very high melting pts. Very strong covalent

bonds (Network bonds) Do not conduct in any

phase Very hard Ex: Diamond (C),

Quartz/Sand (SiO2), Silicon Carbide (SiC)

Silicon Carbide grinding wheels

Metallic Bonding Occurs between atoms that have a small number of valence electrons (metals) leaving them with many vacant valence orbitals and low ionization energies

“Electron Sea Model”

Electron Sea Model Held together by positive kernel and negative

valence electrons. Do you remember kernel? Valence e- free to move from atom to atom. How

can this explain a metal’s conductivity, ductility and malleability?

Just as we did for Ionic, Covalent & Network bonding, let’s compare properties & characteristics of Metallic solids!

Metallic Solids Structural units: made of

positive kernels and valence electrons.

Intermediate melting pts. Relatively intermediate

forces (metallic bonds) Conduct electricity and

heat in all phases. Malleability, Ductility

and Luster. Ex: Cu, Na, Fe, K, Au

Can you complete a summary chart that compares characteristics of all four types of solids?

4 Types of Solids: Ionic, Molecular, Network & Metallic solids

At least 6 properties/characteristics:

type of bonding structural unit melting point conductivity examples other

So far we have talked about bonds between atoms inside molecules or particles! Intramolecular Forces:

found within molecule or particle called chemical bonds.

Think intramurals or intravenous!

6 examples: Ionic Polar Covalent Non-polar covalent Coordinate covalent Network Metallic

Molecular Attractions aka Intermolecular Forces

Bonds (Intramolecular forces) build particles. Now how do these particles come together to build

something we can actually see?

forces beween moleculesThink international or interscholastic

Dipoles Dipoles are polar

molecules “2 poles.” Have asymmetrical

distribution of electrical charge within molecule.

Dipole-dipole attraction is force of attraction between polar molecules.

Hydrogen Bonding Special case of dipole-

dipole that occurs when hydrogen is bonded to a small, highly electronegative atom (N, O, F)

Slightly stronger than other dipole-dipole attraction.

Hydrogen bonding is FON!!! Or I’ve had NOF of Hydrogen bonding!!!

Why doesn’t H2O fit with pattern of others? All are polar, same family, all have same bonding. Smaller mass, should be lower boiling pt.? H2O has hydrogen bonding.

London Dispersion Forcesaka Van Der Waals

Weak attractive forces exist between non-polar molecules (no dipoles, no H bonding)

Caused by momentary dipole, a chance distribution of electrons.

Make it possible for small, non-polar molecules to exist as solids or liquids under low temp and high pressure ex. H2, He, O2, N2

London Dispersion Forces act over short distances. Can these forces be

increased? London Dispersion

forces increase with increase in size of molecules and # of electrons or a decrease in distances between molecules.

Pentane

Changing London Dispersion Forces influences phase of matter. Look at halogens for example!

Molecule-Ion Attraction Ionic compounds

(salts) are generally soluble in polar solvents.

Remember (aq)! Why soluble?

Molecule-Ion Attraction Explained Polar solvents are asymmetrical (+ & - ends). + & - end of liquid are attracted to + & - ends of

ionic salt. This pulls apart ions, breaking crystal lattice structure (salt is dissolved and broken!)

If water is solvent, creates hydrated ions: water molecule surrounded by ions.

Now that ions are broken free of lattice, what can the solution now do?

Chemical Symbols How much lithium is

represented? One Li atom or One

mole of Li atoms (6.02 X 1023)

Chemical Formulas Both a qualitative and

a quantitative expression of the composition of an element or compound.

How much sulfuric acid is represented?

One H2SO4 molecule containing 2 H atoms, 1 S atom & 4 O atoms OR One mole of H2SO4 molecules containing 2 moles of H atoms, 1 mole of S atoms and 4 moles of O atoms.

2 Types of Chemical Formulas Molecular Formula:

indicates total number of atoms of each element needed to form a molecule.

Empirical Formula: simplest ratio in which atoms combine to form a compound

Empirical formulas do not always exist in nature.

For the next section in this unit, you will need to do the following: Write chemical

formulas. Name chemical

compounds. Balance chemical

equations.

Writing Formulas for Ionic Compounds If you use the criss-cross method, remember to

simplify to lowest values. (Use parentheses if multiple polyatomic ions are needed).

1. Write the symbol for the metal ion.2. Write the symbol for the nonmetal or

polyatomic ion.3. Check the oxidation numbers of each ion. If they

add up to zero, this is the formula. A Roman numeral after the name of the metal ion

denotes its oxidation number.

Writing Formulas for Ionic Compounds continued

4. Use the proper subscripts after the symbol for each ion so that when multiplied times the oxidation #, the total algebraic sum is zero. (Use parentheses if multiple polyatomic ions are needed).

Examples: calcium carbonate, ammonium sulfite, Nickel (III) sulfide, Copper (II) chloride

Naming Ionic Compounds Text p.176

1. Write the name of the cation (metal or ammonium ion).

2. Write the name of the anion (nonmetal or polyatomic ion). Nonmetals ending to ide.

3. If the metal can have more than one charge (oxidation number) place a Roman Numeral after its name to denote the charge.

4. Examples: FeBr3, K2Cr2O7, Mn(C2H3O2)3

Naming Covalent Compounds Text pp.206-207

For covalent compounds composed of two elements, name 1st element and then 2nd element’s name ending is changed to –ide.

Use prefixes to indicate number of atoms of each element.

Examples: CO, CO2, CCl4, P2S5,

1 = mono- 2 = di- 3 = tri- 4 = tetra- 5 = penta-

6 = hexa- 7 = hepta- 8 = octa- 9 = nona- 10 = deca-

Balancing Chemical EquationsText pp.267-274

Equations must be balanced to support the Law of Conservation of Mass (matter or mass can not be created nor destroyed, only rearranged).

To balance, you need to make # of atoms of each element the same on both reactants’ and products’ side.

Balancing Chemical Equations Continued

Can’t change the formula, only the coefficients. Coefficient- a small whole number that appears as

a factor in front of a formula in a chemical rxn. Hint: When balancing, start with uncommon

elements first. If polyatomic ions appear on both reactant and product side, balance as a group. Often it is helpful to save H and O until the end.

When finished, your answer should be simplified into smallest whole #’s.

Let’s try some examples.

Balancing Equations Examples1. Fe3O4 + H2 Fe + H2O

2. Al + Pb(NO3)2 Al(NO3)3 + Pb

3. Fe2O3 + CO Fe + CO2

4. Ca(OH)2 + (NH4)2SO4 CaSO4 + NH3 + H2O

5. H3PO4 + CaCl2 Ca3(PO4)2 + HCl