Unfollowing the Octet Rule of 8

Incomplete octet

An atom with less than eight electrons in its valence shell. An atom with less than eight total bonding and nonbonding electrons in a Lewis structure, for example, B in BH3 has an incomplete octet.

There are certain atoms of certain elements that can exist in stable compounds forming bonds with less than eight valence electrons.  When this occurs, the atom of the element within the molecule is said to contain an incomplete octet.  The common examples of such elements are hydrogen (stable with only 2 valence electrons), beryllium (stable with only 4 valence electrons) and boron and aluminum (stable with only 6 valence electrons).   For hydrogen 2 valence electrons give it a noble gas structure (like He) so this is much like the octet rule for everything else below period 1.   But covalent compounds in groups 2 and 3 can form stable compounds in which the valence electrons are not in the noble gas structure.   However, for these compounds you will find that they do form compounds that minimize formal charge.

Example:

 Good examples of the first type of exception are provided by BeCl₂ and BCl₃. Beryllium dichloride, BeCl₂, is a covalent rather than an ionic substance. Solid BeCl₂ has a relatively Complex structure at room temperature, but when it is heated to 750°C, a vapor which consists of separate BeCl₂ molecules is obtained. Since Cl atoms do not readily form multiple bonds, we expect the Be atom to be joined to each Cl atom by a single bond. The structure is

Instead of an octet the valence shell of Be contains only two electron pairs. Similar arguments can apply to boron trichloride, BCl₃, which is a stable gas at room temperature. We are forced to write its structure as

Expanded octet

A hypervalent molecule (the phenomenon is sometimes colloquially known as expanded octet) is a molecule that contains one or more main group elements formally bearing more than eight electrons in their valence shells. Phosphorus pentachloride (PCl₅), sulfur hexafluoride (SF₆), chlorine trifluoride (ClF₃), and the triiodide (I₃⁻) ion are examples of hypervalent molecules.

Examples :

      SF₄

1. This compound is covalent.

2. Determine the total number of valence electrons available:

One sulfur has 6 valence electrons
Four fluorine, each with 7 valence electron, totals 28
This means there are 34 valence electrons, making 17 pairs, available.

3. Organize the atoms so there is a central atom (usually the least electronegative) surrounded by ligand (outer) atoms. Hydrogen is never the central atom.

I put the fluorines like I did because I knew I needed an open space for the unbonded pair on the sulfur. If you put the 4 fluorines around the S like in CH₄, that's OK. It gets tiresome making all this little graphic files, but hey, I knew that when I took on this project. OK, enough complaining, back to work.

4. Determine a provisional electron distribution by arranging the electron pairs (E.P.) until all available pairs have been distributed:

 

5. The formal charge (F) on the central atom is zero. The right-hand structure in step 4 is the correct answer.

Odd molecule

Term invented by Gilbert N. Lewis in 1916 for a molecule containing an odd number of electrons. Taking the p-shell elements, such molecules are rare; they are usually colored and paramagnetic, that is, attracted by a magnet. Odd molecules are 'radicals.' A fine example is nitric oxide, q.v.; nitrogen dioxide is another; chlorine dioxide is also an example, being a reddish-yellow gas. They are all fairly reactive. When including d-shell elements, i.e., the transition metals, the concept mostly doesn't apply, and this 'odd' state is not so unusual.

Example:

1.  Nitrogen dioxide (NO₂)

This molecule has a total of 17 electrons to place - five from the nitrogen and 12 from the oxygens. I will go immediately to the final structure:

Geometry Molecules

Number of Electron Groups	Electron-Group Geometry	Number of Lone Pairs	VSEPR Notation	Molecular Geometry	Ideal Bond Angles	Examples
2	linear	1	AX2		180°	BeH2
3	trigonal-planar	0	AX3		120°	CO32-
		1	AX2E		120°	O3
4	tetrahedral	0	AX4	Tetrahedral	109.5°	S042-
		1	AX3E		109.5°	H3O+
		2	AX2E2		109.5°	H2O
5	trigonal-bipyramidal	0	AX5		90°, 120°	PF5
		1	AX4Eb		90°, 120°	TeCl4
		2	AX3E2		90°	ClF3
		3	AX2E3		180°	I3-
6	octahedral	0	AX6	octahedral	90°	PF6-
		1	AX5E		90°	SbCl52-
		2	AX4E2		90°	ICl4-

Steps Used to Find the Shape of the Molecule

To sum up there are four simple steps to apply the VSEPR theory.

1. Draw the Lewis Structure.

2. Count the number of electron groups and identify them as bond pairs of electron groups or lone pairs of electrons. Remember electron groups include not only bonds, but also lone pairs!

3. Name the electron-group geometry. (State whether it is linear, trigonal-planar, tetrahedral, trigonal-bipyramidal, or octahedral.)

4. Looking at the positions of other atomic nuclei around the central determine the molecular geometry. (See how many lone pairs there are.)

(Use the chart as a guide)

Reference:

http://chemwiki.ucdavis.edu/Wikitexts/UCD_Chem_2A/ChemWiki_Module_Topics/Chemical_Bonding/Geometry_of_Molecules

orbital hybridization

In chemistry hybridization  is the concept of mixing atomic orbitals to form new hybrid orbitals suitable for the qualitative description of atomic bonding properties. Hybridised orbitals are very useful in the explanation of the shape of molecular orbitals for molecules. It is an integral part of valence bond theory. The hybrids are named based on the atomic orbitals that are involved in the hybridization, and the geometries of the hybrids are also reflective of those of the atomic-orbital contributors

 Chemist Linus Pauling first developed the hybridization theory in order to explain the structure of molecules. This concept was developed for such simple chemical systems, but the approach was later applied more widely, and today it is considered an effective heuristic for rationalizing the structures of organic compounds.

EXAMPLES :

Methane, the three 2p orbitals of the carbon atom are combined with its 2s orbital to form four new orbitals called "sp³" hybrid orbitals. The name is simply a tally of all the orbitals that were blended together to form these new hybrid orbitals. Four hybrid orbitals were required since there are four atoms attached to the central carbon atom. These new orbitals will have an energy slightly above the 2s orbital and below the 2p orbitals as shown in the following illustration. Notice that no change occurred with the 1s orbital.

Ammonia, the three 2p orbitals of the nitrogen atom are combined with the 2s orbital to form four sp³ hybrid orbitals. The non-bonded electron pair will occupy a hybrid orbital. Again we need a hybrid orbital for each atom and pair of non-bonding electrons. Ammonia has three hydrogen atoms and one non-bonded pair of electrons when we draw the electron-dot formula. In order to determine the hybridization of an atom, you must first draw the electron-dot formula.

homogenous equilibrium

------The equilibrium between different chemical species present in the same or different phases is called chemical equilibrium. There are two types of chemical equilibrium.

The equilibrium reactions in which all the reactants and the products are in the same phase are called homogeneous equilibrium reactions.

 

C₂H₅OH(l) + CH₃COOH(l)   CH₃COOC₂H₅(l) + H₂O(l)

N₂(g) + 3H₂(g)  2NH₃(g)

2SO₂(g) + O₂(g)   2SO₃(g)

The equilibrium system, ethanol-acetic acid-ethyl acetate-water, is well known. The equation for this equilibrium is: C₂H₅OH + CH₃COOH <-----> CH₃COOC₂H₅ + H₂O

Equilibrium is established very slowly at room temperature when these four substances are mixed. The rate of reactions (both forward and reverse) can be accelerated by using a catalyst so that the equilibrium can be more rapidly established. 3 M HCl is used in this experiment as a catalyst, and the equilibrium at room temperature should be established in a few days. This means that in order to successfully complete this experiment, a series of combinations of reactants and/or products must be prepared a few days before the actual experiment. The equilibrium constant is then calculated by the equation:

Using aA<----->bB

Kc=[B]^a/[A]^b           

We derive this     Kp=P^b_B/P^a_B to get Kp=Kc(RT)^N

K_c = [pure ethyl acetate] [water] / [ethanol] [glacial acetic acid]

The four quantities in the above equation may be replaced by the number of moles of the four substances involved since the volume is constant in each mixture.

Using aA<----->bB

Kc=[B]^a/[A]^b           

We derive this formula     Kp=P^b_B/P^a_B     to get    Kp=Kc(RT)^N

Gas Equilibrium Constants, Kc And Kp

Rosiel Mariano

K_c and K_p are the equilibrium constants of gaseous mixtures.

—     -The Equilibrium Constant, Kc, relates to a chemical reaction at equilibrium. It can be calculated if the equilibrium concentration of each reactant and product in a reaction at equilibrium is known.

  with each concentration raised to the power of the corresponding stoichiometric coefficient.

1. In equilibrium equations, even though the arrows point both ways () we usually associate the right as reactants and the left as products.
2. The products are on the TOP of the fraction (the numerator).
3. The reactants are on the BOTTOM of the fraction (the denominator).
4. The concentrations of the products and reactants are always raised to the power of their coefficient in the balanced chemical equation.
5. If any of the reactants or products are solids or liquids, their concentrations are equal to one because they are pure substances.

2.) H₂ (g) + I₂ (g) 2HI (g)

REFERENCE
http://chemwiki.ucdavis.edu/Physical_Chemistry/Chemical_Equilibrium/The_Equilibrium_Constant/Calculating_An_Equilibrium_Concentration_From_An_Equilibrium_Constant/Writing_Equilibrium_Constant_Expressions_Involving_Gases/Gas_Equilibrium_Constants,_Kc_And_Kp#Problems

lewis acid and bases

 

The Lewis definitions of acids and bases are even more inclusive than the Bronsted definitions. The Lewis definitions are that an acid is an electron-pair acceptor and a base is an electron-pair donor.

 

The term Lewis acid refers to a definition of acid published by Gilbert N. Lewis in 1923, specifically: An acid substance is one which can employ an electron lone pair from another molecule in completing the stable group of one of its own atoms. The term Lewis acid refers to a definition of acid published by Gilbert N. Lewis in 1923, specifically: An acid substance is one which can employ an electron lone pair from another molecule in completing the stable group of one of its own atoms.

The modern-day definition of Lewis acid, as given by IUPAC is a molecular entity (and the corresponding chemical species) that is an electron-pair acceptor and therefore able to react with a Lewis base to form a Lewis adduct, by sharing the electron pair furnished by the Lewis base.^[2] This definition is both more general and more specific—the electron pair need not be a lone pair (it could be the pair of electrons in a π bond, for example), but the reaction should give an adduct (and not just be a displacement reaction).

A Lewis base, then, is any species that donates a pair of electrons to a Lewis acid to form a Lewis adduct. 

 

The Lewis acid-base theory can also be used to explain why nonmetal oxides such as CO₂ dissolve in water to form acids, such as carbonic acid H₂CO₃.

CO₂(g) + H₂O(l)  H₂CO₃(aq)

In the course of this reaction, the water molecule acts as an electron-pair donor, or Lewis base. The electron-pair acceptor is the carbon atom in CO₂. When the carbon atom picks up a pair of electrons from the water molecule, it no longer needs to form double bonds with both of the other oxygen atoms as shown in the figure below

One of the oxygen atoms in the intermediate formed when water is added to CO₂ carries a positive charge; another carries a negative charge. After an H⁺ ion has been transferred from one of these oxygen atoms to the other, all of the oxygen atoms in the compound are electrically neutral. The net result of the reaction between CO₂ and water is therefore carbonic acid, H₂CO₃.

bronsted-lowry acid-base theory

In this section we will consider the Brønsted-Lowry concept. This concept focuses on what an acid or base does.In chemistry, the Brønsted–Lowry theory is an acid-base theory, proposed independently by Johannes Nicolaus Brønsted and Thomas Martin Lowry in 1923. The Bronsted-Lowery concept of acids and bases is that acid-base reactions can be seen as proton-transfer reactions. This results in acids and bases being able to be defined in terms of this proton (H+) transfer.  In this system,Brønsted acids and Brønsted bases are defined, by which an acid is a molecule or ion that is able to lose, or "donate," a hydrogen cation (proton, H+), and a base is a species with the ability to gain, or "accept," a hydrogen cation (proton).

An Acid is a Proton Donor and a Base is a Proton Acceptor

So what does this mean? For a reaction to be in equilibrium a transfer of electrons needs to occur. The acid will give an electron away and the base will receive the electron. Acids and Bases that work together in this fashion are called a conjugate pair made up of conjugate acids and conjugate bases.

So what does this look like? (A stands for an Acidic compound and Z stands for a Basic compound)

HA + Z ↔ A- + HZ+

A Donates H to form HZ+.

Z Accepts H from A which forms HZ+

A- becomes conjugate base of HA and in the reverse reaction it accepts a H from HZ to recreate HA in order to remain in equilibrium

HZ+ becomes a conjugate acid of Z and in the reverse reaction it donates a H to A- recreating Z in order to remain in equilibrium

Properties of acids and bases

It follows that, if a compound is to behave as an acid, donating a proton, there must be a base to accept the proton. So the Brønsted–Lowry concept can be defined by the reaction:

acid + base   conjugate base + conjugate acid.

The conjugate base is the ion or molecule remaining after the acid has lost a proton, and the conjugate acid is the species created when the base accepts the proton. The reaction can proceed in either forward or backward direction; in each case, the acid donates a proton to the base.

Example:

reaction of hydrochloric acid with ammonia

which reduces to:

reduced reaction of hydrochloric acid with ammonia

because there is a Cl-(aq), on each side. We now have the net ionic equation after we cancel out the "spectator ions"(Cl-).

What happens in this reaction in aqueous solution is that a proton is transferred from H3O+ to NH3. This results in H3O+ losing a (H+), resulting in H2O. The NH3 gains the transferred proton, resulting in NH4+. We call H3O+ the proton donor, or acid. We call NH3 the proton acceptor, or base

The Bronsted-Lowery concept defines something as either an acid or base depending on its function in the acid-base (proton transfer) reaction. Some things can act as either an acid or a base. These are called amphiprotic species, they can either lose or gain a proton, depending on the other reactant. An example of an amphiprotic species would be:

example of an amphiprotic species
In the presence of OH-, it acts as an acid. In the presence of HF it acts as a base. Water is also amphiprotic, as are most anions with ionizable hydrogens and certain solvents. Water as an amphiprotic species is very important to the acid-base reactions.

In the Bronsted-Lowery concept we have found that:

A base is a species that accepts protons, while an acid is a species that donates protons.

Acids and bases can be ions as well as molecular substances.

Some species can act as either acids or bases, depending on what the other reactant is.

strengths Of acids and bases

What Makes a Strong Acid or Strong Base?

Strong electrolytes are completely dissociated into ions in water. The acid or base molecule does not exist in aqueous solution, only ions. Weak electrolytes are incompletely dissociated.

Strong Acids

Strong acids are strong electrolytes that, for practical purposes, are assumed to ionize completely in water. Strong acids completely dissociate in water, forming H⁺ and an anion. There are six strong acids. The others are considered to be weak acids. Most of them are inorganic acids. You should commit the strong acids to memory:

·         HCl - hydrochloric acid

·         HNO₃ - nitric acid

·         H₂SO₄ - sulfuric acid

NOTE: this is a diprotic; we show only the first stage of ionization here.

·         HBr - hydrobromic acid

·         HI - hydroiodic acid

·         HClO₄ - perchloric acid

Weak Acids

A weak acid only partially dissociates in water to give H⁺ and the anion. Examples of weak acids include hydrofluoric acid, HF, and acetic acid, CH₃COOH. Weak acids include:

·         Molecules that contain an ionizable proton. A molecule wih a formula starting with H usually is an acid.

·         Organic acids containing one or more carboxyl group, -COOH. The H is ionizable.

·         Anions with an ionizable proton. (e.g., HSO₄^- → H⁺ + SO₄^2-)

·         Cations                                           

·         transition metal cations

·         heavy metal cations with high charge

·         NH₄⁺ dissociates into NH₃ + H⁺

Strong Bases

Strong bases are all strong electrolytes that ionizes completely in water. It dissociate 100% into the cation and OH^- (hydroxide ion). The hydroxides of the Group I and Group II metals usually are considered to be strong bases.

·         LiOH - lithium hydroxide                                        CsOH - cesium hydroxide

·         NaOH - sodium hydroxide                                      *Sr(OH)₂ - strontium hydroxide

·         KOH - potassium hydroxide                                               Ba(OH)₂ - barium hydroxide  

·         RbOH - rubidium hydroxide                                               *Ca(OH)₂ - calcium hydroxide

 

* These bases completely dissociate in solutions of 0.01 M or less. The other bases make solutions of 1.0 M and are 100% dissociated at that concentration. There are other strong bases than those listed, but they are not often encountered.

Weak Bases

Weak bases, like weak acids, are weak electrolytes. Examples of weak bases include ammonia, NH₃  (note: ammonia does not ionizelike an acid because it does not split up to form ions the way, say, HCL does), and diethylamine, (CH₃CH₂)₂NH.

·         Most weak bases are anions of weak acids.

·         Weak bases do not furnish OH^- ions by dissociation. Instead, they react with water to generate OH^- ions.

 

 

TABLE1 lists some important conjugate acid-base pairs, in order of their relative strengths. Conjugate acid-base have the following properties:

 

1.      If an acid is strong, its conjugate acid-base has no measurable strength. Thus, the  ion, which is the conjugate base of the strong acid HCl, is an extremely weak base.

2.         is the strongest acid that can exist in aqueous solution. Acids stronger than  react with water to produce  and their conjugate bases. Thus, HCl, which is a stronger acid than , reacts with water completely to form  and :

HCl(aq) +            (aq) + 

 

Acids weaker than  react with water to a much smaller extent, producing  and their conjugate bases. For example, the following equilibrium lies primarily to the left:

HF(aq)  +             (aq)  + 

    

3.      The  ion is the strongest base that can exist in aqueous solution. Bases stronger than  react  with water to produce and their conjugate acids. For example, the oxide ion () is a stronger base than  , so it reacts with water completely as follows:

   (aq) +           

For this reason the oxide ion does not exist in aqueous solutions.

EXAMPLE:

Calculate the pH of (a) a  M HCl olution and (b) a 0.020 M  solution.

STRATEGY:

 Keep in mind that HCl is a strong acid and  is a strong base. Thus, these species are completely ionized and no HCl and  will be left in solution.