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CHEMICAL BONDING

Chemical Bonding

We have covered the basic ideas of atomic structure, but it is worth realizing that of all the elements only the noble gases are found naturally in a form such that their atoms occur as individuals, widely separated from all other atoms. Under the conditions that prevail on the surface of the earth, almost all atoms are linked by chemical bonds to other atoms. Oxygen, for example, is the most common element on earth. It is found in combination with metals in rocks, with hydrogen in water, with carbon and hydrogen in living organisms, or as the diatomic molecule O2 in the atmosphere, but individual oxygen atoms are quite rare. Most other elements behave in a similar way. Thus, if we want to understand the chemistry of everyday matter, we need to understand the nature of the chemical bonds which hold atoms together.

Theories of chemical bonding invariably involve electrons. When one atom approaches another, the valence electrons, found in the outermost regions of the atoms, interact long before the nuclei can come close together. Electrons are the least massive components of an atom, and so they can relocate to produce electrostatic forces which hold atoms together. According to Coulomb’s law, such electrostatic or coulombic forces are quite large when charges are separated by distances of a few hundred picometers—the size of an atom. Coulombic forces, then, are quite capable of explaining the strengths of the bonds by which atoms are held together.

An important piece of evidence relating electrons and chemical bonding was noted by G. N. Lewis shortly after the discovery that the atomic number indicated how many electrons were present in each kind of atom. Most chemical formulas correspond to an even number of electrons summed over all constituent atoms. Thus H2O has 2 electrons from the 2 H’s and 8 from O for a total of 10, NCl3 has 7 + (3 × 17) = 58 electrons, and so on. This is a bit surprising when you consider that half the elements have odd atomic numbers so that their atoms have an odd number of electrons. Lewis suggested that when atoms are bonded together, the electrons occur in pairs, thus accounting for the predominance of even numbers of electrons in chemical formulas. These pairs often leave bonded atoms with eight electrons in the outer most shell, known as the octet rule.

There are two important ways in which the valence shells of different atoms can interact to produce electron pairs and chemical bonds. When two atoms have quite different degrees of attraction for their outermost electrons, one or more electrons may transfer their allegiance from one atom to another, pairing with electrons already present on the second atom. The atom to which electrons are transferred will acquire excess negative charge, becoming a negative ion, while the atom which loses electrons will become a positive ion. These oppositely charged ions will be held together by the coulombic forces of attraction between them, forming an ionic bond. Since electrons are shifted so that one neutral atom becomes positively charged, and the other becomes negatively charged, substances formed from ionic bonding are often in pairs and are then referred to as binary ionic compounds. Binary ionic compounds are not too common, but the existence of polyatomic ions greatly extends the number of ionic substances. Because oppositely charged ions are held in crystal lattices by strong coulombic forces, physical properties of ionic compounds include hardness, brittleness, and having high melting and boiling points. The majority of them dissolve in water, and in solution each ion exhibits its own chemical properties. Ionic compounds obey the octet rule, which explains why ions are generally in noble-gas electron configurations.

On the other hand, when two atoms have the same degree of attraction for their valence electrons, it is possible for them to share pairs of electrons in the region between their nuclei. Such shared pairs of electrons attract both nuclei, holding them together with a covalent bond. Sharing one or more pairs of electrons between two atoms attracts the nuclei together and usually results in an octet around each atom. Covalent bonding often produces individual molecules, like CO2 or CH3CH2OH, which have no net electrical charge and little attraction for each other. Thus covalent substances often have low melting and boiling points and are liquids or gases at room temperature. Occasionally, as in the case of SiO2, an extended network of covalent bonds is required to satisfy the octet rule. Such giant molecules result in solid compounds with high melting points.
A number of atomic properties, such as ionization energy, electron affinity, van der Waals radii, covalent atomic radii, and ionic radii, are important in determining whether certain elements will form covalent or ionic compounds and what properties those compounds will have. In the next sections we will consider the formation of covalent and ionic bonds and the properties of some substances containing each type of bond. The figure below previews how each of atomic properties important to understanding bonding varies according to an atom’s position in the periodic table. The figure also includes one atomic property, electronegativity, which will be covered when we explore further aspects of covalent bonding in more depth.

Periodic variations of atomic properties.
 
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Posted by on August 2, 2012 in physical chemistry

 

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PHYSICAL CHEMISTRY

Not long ago I heard a student say she wasn’t going to declare the biochemistry major because she was “afraid” of pchem. Afraid of pchem? What? Get it together.

I know some students are scared because of the word “physics.” Lets face it; many of us are not excellent physicists, but the truth is that how well you did in physics is not a predictor of how well you will do in pchem. Pchem is not physics. It’s general chemistry on steroids,, we aren’t even required to perform calculus! I say perform because, though you will be expected to know how to integrate by parts, you will never have to do it on an exam. That’s pretty amazing.

Doing well in pchem does not require unconceivable levels of intelligence. It’s all about repetition.

First, attend every lecture, because reading the book chapter is not enough. The book will never tell you what the professor really cares about, because the book cares about every topic from “the chemical potential of the solute in terms of the mole fraction” to………….

Second, you must make it a point to write down every bit of seemingly trivial information that the professor throws, quickly and surreptitiously, your way. Pages one to three of the midterm are inevitably constructed around true-false, multiple-choice, theoretical armchair physicist questions. Words can hardly express the kind of anxiety questions like “which best describes the second law of thermodynamics” can cause the unsuspecting soul.

It is such an elementary question, yet during my studying I’d never bothered to remember what law was the first, the third, or the zillionth. It seemed completely unimportant and incidental, and I had decided very early on that I’d focus on knowing how to apply the laws rather than knowing the order in which they came. Prodigious mistake, so be sure to memorize any and all definitions as well as key concepts.

Third, do not believe any professor who tells you not to memorize the things he says. Also, do not believe a professor who tells you that if you just focus on the underlying ideas, you will have the ultimate key to understanding physical chemistry. LIES. Blatant lies. Sometimes understanding doesn’t come until years later!

Which brings up to tip number four. If, like me, you’re good with numbers and if, like me, you wonder where theory fits in, then treat calculations and concepts separately. I know it sounds infinitely foolish, but I’ve done it very successfully. Good old rote memorization and perfect lecture attendance go a long way.

If you’d like to work on mastering the numbers, this next short paragraph is for you. While homework problems are largely inadequate for conceptual thinking, doing them two to three times over will ensure you are ready to tackle the fun part of the midterm: calculations. I am really good with numbers and I can choose, rearrange and apply formulas (which are provided to you on the midterm, by the way) without a problem. This kind of proficiency comes easily if you do the homework over, and over, and over again. Eventually you start figuring out what formulas can do for you, and then you become free to solve problems any way you choose. This kind of skill will be tremendously helpful because many problems can be solved at least two different ways. That’s literally all there is to it.

Physical chemistry is every bit as interesting as it sounds, it is not fundamentally difficult, and you can do great. Identify your weakness ahead of time, and be prepared to fix it. Nobody need be afraid of physical chem

 
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Posted by on July 29, 2012 in physical chemistry

 

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Rate law for the reaction

One of the first steps in studying the kinetics of a chemical reaction is to determine the rate law for the reaction. One method for making this determination is to experimentally measure how the concentration of a reactant or product varies with time and then make characteristic kinetics plots. Another strategy for determining the rate law is to use the method of initial rates.

The Method of Initial Rates involves measuring the rate of reaction, r, at very short times before any significant changes in concentration occur. Suppose one is studying a reaction with the following stoichiometry:

A + 2 B   –>   3 C

While the form of the differential rate law might be very complicated, many reactions have a rate law of the following form:

r = k [A]a [B]b

The initial concentrations of A and B are known; therefore, if the initial reaction rate is measured, the only unknowns in the rate law are the rate constant, k, and the exponents a and b. One typically measures the initial rate for several different sets of concentrations and then compares the initial rates.

Consider the following set of data:

Trial Rate
(mole L-1 sec-1)
Initial Concentration
of A (mole L-1)
Initial Concentration
of B (mole L-1)
1 2.73 0.100 0.100
2 6.14 0.150 0.100
3 2.71 0.100 0.200

If simple multiples are chosen for the concentrations and only one concentration is varied at a time, one can determine a and b by inspection. In this case the values of a and b may not be obvious. One can employ the following algebraic technique for determining the exponents.

First, write the ratio of the rate laws for two trials.

r1

r2
= k [A]1a [B]1b

k [A]2a [B]2b

Next, substitute the numerical values into the equation.

2.73 mole L-1 sec-1

6.14 mole L-1 sec-1
= k (0.100 mole L-1)a (0.100 mole L-1)b

k (0.150 mole L-1)a (0.100 mole L-1)b

Notice that the units for each quantity and the rate constant can be removed, and in this case the exponent b is removed when the concentrations of B divide. The equation simplifies to

2.73

6.14
= 0.100a

0.150a

0.4446 = 0.6667a

To convert a from an exponent into a coefficient, take the logarithm of both sides of the equation.

ln[0.4446] = ln[0.6667a]

-0.8106 = -0.4054a

The value of a may now be readily determined.

a = -0.8106

-0.4054
= 1.9995

In most cases, the exponents are integers (or less commonly fractions such as 1/2). In this case the reaction is second order in A (a = 2). A similar strategy can be employed to determine the value of b. (Actually, it should be obvious from inspection of trials 1 and 3 that the reaction is zero-order in B.) Once the exponents are known, the rate constant can be calculated. Because the data generally suffers from experimental error, it is best to calculate the rate constant for each trial and use the average value.

 

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Chemical bonding

the hydrogen bond desiraju 2011

hydrogen bond can depicted as an X-Y…H-Z interaction involving 0.5 to 40 kcal/moles of energy with the strongest hydrogen bond stronger than weakest covalent bond and the weakest H-bond in the range of a van der Waals bond. In an Angewandte essay Gautam R. Desiraju elaborates on the nature of the hydrogen bond and a proper definition with unusual clarity. (DOI). For those of you without an Angewandte subscription or with without patience to read all 8 pages, here is a brief capture. To start with just a few quotes:
On definitions: The history of chemistry is strewn with names that were once hotly contested and disputed by various protagonists and why do chemists fight over names so much, good question.
On trivial namesa term is acceptable if the largest numbers of chemists are in a maximum degree of agreement about what it means. A trivial name is the victory of such a consensus
On chemical bonds: the word bond has an almost religious connotation for chemists (…) every chemist has his or her own idea as to what constitutes a bond
On the origin of life : did life on earth become water-based because the hydrogen bond was the only interaction of choice to facilitate life?
So what is the problem with these hydrogen bonds. Strong hydrogen bonds or hydrogen bridges found in O-H–O-R configurations pose no problem. Problems arise with weaker bonds, not always possible or difficult to detect it experimentally More disagreement: in the original Pauling definition and later IUPAC definition Y and Z are both electronegative as in R-O-H…O-R (electropositive hydrogen partners with two electronegative partners) while in an scenario with C-H…O=C-R the carbon atom is electronegative and the oxygen atom electronegative. So not a hydrogen bond then?
An elite team of IUPAC chemists of which Desiraju was part of have now decided on the future of the hydrogen bond. In a new definition the hydrogen bond is an attractive interaction between a hydrogen atom from a molecule or a molecular fragment X-H in which X is more electronegative than H, and an atom or a group of atoms in the same or a different molecule, in which there is evidence of bond formation. So far so good. The definition goes on with the evidence for hydrogen-bond formation may be experimental or theoretical, or ideally, a combination of both. Here Desiraju stresses that theory and experiment have equal status which may come as a surprise to those who believe in the ancient scientific principle that theories are validated by experiment. The new definition also stipulates that the actual bridge is formed by Y-H but if Desiraju has had his way the bridge would be formed by all 4 atoms in X-Y-H-Z.
The X to Y distance no longer needs to be smaller than the sum of the van der Waals radii of X and Y: the precision of X-ray analysis does not allow it.

DR VIJAYA SHASTRY

RJ COLLEGE, GHATKOPAR

MUMBAI,INDIA

mob 09819794779

 
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Posted by on July 20, 2012 in physical chemistry

 

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Stoichiometry

Stoichiometry

Amounts of reactantsLet’s start with how to say this word. Five syllables: STOY-KEE-AHM-EH-TREE. It’s a big word that describes a simple idea. Stoichiometry is the part of chemistry that studies amounts of substances that are involved in reactions. You might be looking at the amounts of substances before the reaction. You might be looking at the amount of material that is produced by the reaction. Stoichiometry is all about the numbers.

All reactions are dependent on how much stuff you have. Stoichiometry helps you figure out how much of a compound you will need, or maybe how much you started with. We want to take the time to explain that reactions depend on the compounds involved and how much of each compound is needed.

Reactions are limitedWhat do you measure? It could be anything. When you’re doing problems in stoichiometry, you might look at…
– Mass of Reactants (chemicals before the reaction)
– Mass of Products (chemicals after the reaction)
– Chemical Equations
– Molecular Weights of Reactants and Products
– Formulas of Various Compounds

Now, an example. Let’s start with something simple like sodium chloride (NaCl). You start with two ions and wind up with an ionic/electrovalent compound. When you look at the equation, you see that it takes one sodium ion (Na+) to combine with one chlorine ion (Cl) to make the salt. When you usestoichiometry, you can determine amounts of substances needed to fulfill the requirements of the reaction. Stoichiometry will tell you that, if you have ten million atoms of sodium and only one atom of chlorine, you can only make one molecule of sodium chloride. Nothing you can do will change that. It’s like this:

10,000,000 Na + 1 Cl –> NaCl + 9,999,999 Na

Hydrogen and Oxygen moleculesLet’s bump it up a level. When you mix hydrogengas (H2) and oxygen gas (O2), nothing much happens. When you add a spark to the mixture, all of the molecules combine and eventually form water (H2O). You would write it like this:

2H2 + O2 –> 2H2O

What does stoichiometry look at here? First, look at the equation. Four hydrogen atoms and two oxygen atoms are on each side of the equation. It’s an important idea to see that you need twice as many hydrogen atoms as you do oxygen atoms. The number of atoms in the equation will help you figure out how much of each substance you will need to make the reaction happen. If you make this an extreme example and fill a sealed container with one million hydrogen molecules and only one oxygen molecule, the spark won’t make an explosion. There is no monster reaction to be created when there is only one oxygen molecule around. You will make two water molecules and be done.

DR VIJAYA SHASTRY

 

 

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DR VIJAYA SHASTRY
ASST PROF, PHYSICAL CHEMISTRY
RJ COLLEGE
GHATKOPAR
MUMBAI, INDIA

 

 

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