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Ionic Bonds


Ionic bonds form when two atoms have a large difference in electronegativity. (Electronegativity is the quantitative representation of an atom’s ability to attract an electron to itself). Although scientists do not have an exact value to signal an ionic bond, the amount is generally accepted as 1.7 and over to qualify a bond as ionic. Ionic bonds often occur between metals and salts; chloride is often the bonding salt. Compounds displaying ionic bonds form ionic crystals in which ions of positive and negative charges hover near each other, but there is not always a direct 1-1 correlation between positive and negative ions. Ionic bonds can typically be broken through hydrogenation, or the addition of water to a compound.

Covalent Bonds


Covalent bonds form when two atoms have a very small (nearly insignificant) difference in electronegativity. The value of difference in electronegativity between two atoms in a covalent bond is less than 1.7. Covalent bonds often form between similar atoms, nonmetal to nonmetal or metal to metal. Covalent bonding signals a complete sharing of electrons. There is usually a direct correlation between positive and negative ions, meaning that because they share electrons, the atoms balance. Covalent bonds are usually strong because of this direct bonding.

Polar Covalent Bonds


Polar covalent bonds fall between ionic and covalent bonds. They result when two elements bond with a moderate difference in electronegativity moderately to greatly, but they do not surpass 1.7 in electronegativity difference. Although polar covalent bonds are classified as covalent, they do have significant ionic properties. They also induce dipole-dipole interactions, where one atom becomes slightly negative and the other atom becomes slightly positive. However, the slight change in charge is not large enough to classify it entirely as an ion; they are simply considered slightly positive or slightly negative. Polar covalent bonds often indicate polar molecules, which are likely to bond with other polar molecules but are unlikely to bond with non-polar molecules.

Hydrogen Bonds


Hydrogen bonds only form between hydrogen and oxygen (O), nitrogen (N) or fluorine (F). Hydrogen bonds are very specific and lead to certain molecules having special properties due to these types of bonds. Hydrogen bonding sometimes results in the element that is not hydrogen (oxygen, for example) having a lone pair of electrons on the atom, making it polar. Lone pairs of electrons are non-bonding electrons that sit in twos (pairs) on the central atom of the compound. Water, for example, exhibits hydrogen bonding and polarity as a result of the bonding. This is shown in the diagram below.

Because of this polarity, the oxygen end of the molecule would repel negative atoms like itself, while attracting positive atoms, like hydrogen. Hydrogen, which becomes slightly positive, would repel positive atoms (like other hydrogen atoms) and attract negative atoms (such as oxygen atoms). This positive and negative attraction system helps water molecules stick together, which is what makes the boiling point of water high (as it takes more energy to break these bonds between water molecules).

In addition to the four types of chemical bonds, there are also three categories bonds fit into: single, double, and triple. Single bonds involve one pair of shared electrons between two atoms. Double bonds involve two pairs of shared electrons between two atoms, and triple bonds involve three pairs of shared electrons between two atoms. These bonds take on different natures due to the differing amounts of electrons needed and able to be given up.

Now, let’s look at determining what types of bonds we see in different compounds. We’ve already looked at the bonds in H2O, which we determined to be hydrogen bonds. However, now let’s look at a few other types of bonds as examples.

Compound: HNO3 (also known as Nitric acid)

There are two different determinations we can make as to what these bonds look like; first we can decide whether the bonds are covalent, polar covalent, ionic, or hydrogen. Then, we can determine if the bonds are single, double, or triple.

In order to decide whether the bonds are covalent, polar covalent, ionic or hydrogen, we need to look at the types of elements seen and the electronegativity values. We look at the elements and see hydrogen, nitrogen, and oxygen—no metals. This rules out ionic bonding as a type of bond seen in the compound. Then, we would look at electronegativity values for nitrogen and oxygen. Oftentimes, this information can be found on a periodic table, in a book index, or an educational online resource. The electronegativity value for oxygen is 3.5 and the electronegativity value for nitrogen is 3.0. The way to determine the bond type is by taking the difference between the two numbers (subtraction). 3.5 – 3.0 = 0.5, so we can determine that the bond between nitrogen and oxygen is a covalent bond. We can also determine, from past knowledge, that the bond between oxygen and hydrogen is a hydrogen bond as it was in water.

Now, we need to count the electrons and draw the diagram for HNO3. For more help counting electrons, please see the page on Electron Configuration. For more help drawing the Lewis structures, please see the page on Lewis Structures. This process combines both of these in order to determine the structure and shape of a molecule of the compound.

First, we determine that N follows the octet rule, so it needs eight surrounding electrons. This is important to keep in mind as we move forward. Next we count up how many valence electrons the compound has as a whole. H gives us 1, N gives us 5, and each O gives us 6. We can discern this from looking at the tops of the columns in the periodic table (see above). We then add these numbers together (3 x 6 = 18, + 1 = 19, + 5 = 24), and we get 24 electrons that we need to distribute throughout the molecule. First, we need to draw the molecule to see how many initial bonds we’ll be putting in. Our preliminary structure looks like this:

Now, we can count how many electrons we have used by counting 2 electrons for each bond placed. We see that we have placed 4 bonds, so we have used 8 electrons. 24 – 8 = 16 electrons that we need to distribute. In order to correctly place the rest of the electrons, we need to determine how many electrons each atom needs to be stable.

The central atom, N, has three bonds attached (equivalent of 6 electrons) so it needs 2 more electrons to be stable. The O to the right has one bond (two electrons) so it needs 6 more to be stable. The O above the N has one bond (two electrons) so it also needs 6 electrons to be stable. The O to the left of the N is bonded both to N and to H, so it has two bonds (4 electrons); therefore, it needs 4 more electrons to be stable. We add up the total amount of electrons needed, 2 + 6 + 6 + 4 = 18, and see that we need 18 electrons to stabilize the compound. We know this is not possible, since we only have 16 available electrons. When this happens, we need to insert a double bond in order to resolve the problem of lack of electrons. This is because, although we count each bond as 2 electrons, the elements joined together in the bond are actually sharing the electrons. Therefore, when we count out the bonds, we are counting some electrons twice because they are shared. This is normal and expected, and resolves not having enough valence electrons. Now, we need to decide where to put the double bond in this compound. We know that the double bond cannot go between O and H, because H does not have enough room to accept another electron. Therefore, we know we must place the bond between N and O. You might be thinking, how do I decide where to put the bond? In this particular example, we can place the bond either between the top O and N, or the right O and N. This is because HNO3 displays resonance.

Here are the ways you can place the double bond:



or

We are going to keep the bond between N and the right O in our example. After we add in the bond, we subtract two more electrons from our available electrons (16) and are left with 14 electrons to distribute. Now we need to make sure we have the correct number of electrons. After placing in the double bond, N is now stable because it has 4 bonds (8 electrons) surrounding it. It does not need any additional electrons. The top O (above N) needs 6 electrons, the right O now only needs 4 electrons (because it has a double bond now, which is 4 electrons), and the left O still needs 4 electrons to become stable. We add these numbers together, 6 + 4 + 4 = 14, and we see that 14 is the number of electrons we have, so we can go ahead and distribute them, like this:



Now, our compound is stable with appropriately distributed valence electrons. We can see that there are three single bonds (H—O, N—O, and N—O) and one double bond (N==O). [5]



Periodic Law and Periodic System

The early innovation of Periodic law by a Russian chemist, Dmitry I. Mendeleev in the mid-19th century, has been of great value in the growth of chemistry. In chemistry, Periodic law is the arrangement of theelements in order of increasing atomic number, that is, the total number of protons in the atomic nucleus and their physical and chemical properties recur in a pattern with the increasing atomic number. In the periodic table, the horizontal rows, known as the “periods” exemplify these relations. In fact, until the second decade of the 20th century, it was not documented that the order ofelements in the periodic system is that of their atomic numbers. 


What is the History of Periodic Law?


Classification became necessary due to the increase in the number of elements discovered. In 1817, J.W. Dobereiner, a German chemist , explained that the atomic weight of strontium rests in the middle between that of calcium and barium, and few years later he demonstrated that other such “triads” such as chlorine, bromine, and iodine  and lithium, sodium, and potassium are also present. Other scientists, during1827 and 1858 developed analogous relationships which extended more than the triads of elements, fluorine being added to the halogens and magnesium to the alkaline-earth metals, while oxygen, sulfur, selenium, and tellurium were classed as one family and nitrogen, phosphorus, arsenic, antimony, and bismuth as one more element family.  De Chancourtois, a French scientist, in 1862 anticipated a classification of the elements based on the new values of atomic weights. It was finally the  Russian chemist, Mendeleev, who proposed the periodic law, which indicated that the elements arranged in the increasing order of atomic weights show a periodic change of properties. In 1869, the first periodictable was tabulated by Mendeleev. It had 17 columns, with two nearly complete periods of elements, frompotassium to bromine and rubidium to iodine, lead by two fractional periods of seven elements each (lithium to fluorine and sodium to chlorine), and three incomplete periods. 

What is the  importance of the Periodic law?

- Innovation of New Elements: Through Mendeleyev’s efforts in 1871, the grand significance of the periodic law was made apparent in predicting that the properties of the 17 elements could be interrelated with those of other elements by relocating the 17 to new positions from those shown by their atomic weights. The subsistence of many of the properties of uninvented elements of that time like eka-boron, eka-aluminum, and eka-silicon, now identified with the elements scandium, gallium, and germanium, was also predicted by Mendeleev. Likewise, following the discovery of helium and argon, the periodic law allowed the discovery of the subsistence of neon, krypton, xenon, and radon. Besides, the absence of element 72 was anticipated, from its position in the periodic system, to be alike to zirconium in its properties rather than to the rare earths, was shown by Niels Bohr, a Danish physicist. In 1922, other scientists also examined zirconium ores following Bohr’s prediction.

- Implication of atomic numbers: A number of of the elements in the Mendeleyev periodic tables were necessary to arrange elements according to their atomic weight . For example, in the pair’s of argon and potassium, cobalt and nickel, and tellurium and iodine, the first element had the previous place in the periodic system but with a bigger atomic weight. When the structure of the atom was studied, the explanation to this complexity was solved. Research done by Ernest Rutherford on the scattering of alpha particles by the nuclei of heavy atoms, in1910, helped the prediction of the nuclear electrical charge. Approximately, the ratio of the nuclear charge to that of the electron was noted to be one-half the atomic weight. Another scientist, in 1911 recommended that this quantity, the atomic number, might be recognized with the ordinal number of the element in the periodic table. In 1913, this proposal was vividly established by the English physicist, H.G.J. Moseley’s measurements of the wavelengths of the characteristic X-ray spectral lines of many elements, which showed that the wavelengths relied in a usual way on the atomic numbers indistinguishable with the ordinal numbers of the periodic table elements.


 

How did the  periodic law evolve ?


Starting in 1913, thorough knowledge of the the elements and their properties had improved. Followed by the exclusion principle by the Austrian theoretical physicist, Wolfgang Pauli in 1925, the invention of the spin of the electron by George E. Uhlenbeck and Samuel Goudsmit in 1925, and the development of quantum mechanics by Werner Heisenberg were the significant advanced inventions which developed the periodic law. The development of the electronic theory of valence and molecular structure, beginning with the postulate of the shared electron pair by  the chemist, Gilbert N. Lewis in 1916, also played a vital role in elucidating the periodic law. [6]

The Periodic Table

The Periodic Table

"If all the elements are arranged in the order of their atomic weights, a periodic repetition of properties is obtained. This is expressed by the law of periodicity." 


Dmitri Mendeleev, Principles of Chemistry, Vol. 2, 1902, P. F. Collier, p17. 

The periodic table we use today is based on the one devised and published by Dmitri Mendeleev in 1869.

Mendeleev found he could arrange the 65 elements then known in a grid or table so that each element had:

1. A higher atomic weight than the one on its left. For example, magnesium (atomic weight 24.3) is placed to the right of sodium (atomic weight 23.0):



23.0
Na

24.3
Mg

 

2. Similar chemical properties to other elements in the same column - in other words similar chemical reactions. Magnesium, for example, is placed in the alkali earths' column:



9.01
Be

24.3
Mg

40.1
Ca

87.6
Sr

 

Mendeleev realized that the table in front of him lay at the very heart of chemistry. And more than that, Mendeleev saw that his table was incomplete - there were spaces where elements should be, but no-one had discovered them.

Just as Adams and Le Verrier could be said to have discovered the planet Neptune on paper, Mendeleev could be said to have discovered germanium on paper. He called this new element eka-silicon, after observing a gap in the periodic table between silicon and tin:


28.1
Si

??
??

119
Sn

Similarly, Mendeleev discovered gallium (eka-aluminum) and scandium (eka-boron) on paper, because he predicted their existence and their properties before their actual discoveries.

Although Mendeleev had made a crucial breakthrough, he made little further progress. With the benefit of hindsight, we know that Mendeleev's periodic table was underpinned by false reasoning. Mendeleev believed, incorrectly, that chemical properties were determined by atomic weight. Of course, this was perfectly reasonable when we consider scientific knowledge in 1869.

In 1869 the electron itself had not been discovered - that happened 27 years later, in 1896.

In fact, it took 44 years for the correct explanation of the regular patterns in Mendeleev's periodic table to be found. [7]




The law of conservation of mass. 

The law of conservation of mass or principle of mass conservation states that for any system closed to all transfers of matter andenergy, the mass of the system must remain constant over time, as system mass cannot change quantity if it is not added or removed. Hence, the quantity of mass is "conserved" over time. The law implies that mass can neither be created nor destroyed, although it may be rearranged in space, or the entities associated with it may be changed in form, as for example when light or physical work is transformed into particles that contribute the same mass to the system as the light or work had contributed. The law implies (requires) that during anychemical reactionnuclear reaction, or radioactive decay in an isolated system, the total mass of the reactants or starting materials must be equal to the mass of the products.

The concept of mass conservation is widely used in many fields such as chemistrymechanics, and fluid dynamics. Historically, mass conservation was discovered in chemical reactions by Antoine Lavoisier in the late 18th century, and was of crucial importance in the progress from alchemy to the modern natural science of chemistry.

The closely related concept of matter conservation was found to hold good in chemistry to such high approximation that it failed only for the high energies treated by the later refinements of relativity theory, but otherwise remains useful and sufficiently accurate for most chemical calculations, even in modern practice.

In special relativity, needed for accuracy when large energy transfers between systems is involved, the difference between thermodynamically closed and isolated systems becomes important, since conservation of mass is strictly and perfectly upheld only for so-called isolated systems, i.e. those completely isolated from all exchanges with the environment. In this circumstance, the mass–energy equivalence theorem states that mass conservation is equivalent to total energy conservation, which is the first law of thermodynamics. By contrast, for a thermodynamically closed system (i.e., one which is closed to exchanges of matter, but open to exchanges of non-material energy, such as heat and work, with the surroundings) mass is (usually) only approximately conserved. The input or output of non-material energy must change the mass of the system in relativity theory, although the change is usually small, since relatively large amounts of such energy (by comparison with ordinary experience) carry only a small amount of mass (again by ordinary standards of measurement).

In special relativity, mass is not converted to energy, since mass and energy cannot be destroyed, and energy in all of its forms always retains its equivalent amount of mass throughout any transformation to a different type of energy within a system (or translocation into or out of a system). Certain types of matter (a different concept) may be created or destroyed, but in all of these processes, the energy and mass associated with such matter remains unchanged in quantity (although type of energy associated with the matter may change form).

In general relativity, mass (and energy) conservation in expanding volumes of space is a complex concept, subject to different definitions, and neither mass nor energy is as strictly and simply conserved as is the case in special relativity and in Minkowski space. For a discussion, see mass in general relativity.

An important idea in ancient Greek philosophy was that "Nothing comes from nothing", so that what exists now has always existed: no new matter can come into existence where there was none before. An explicit statement of this, along with the further principle that nothing can pass away into nothing, is found in Empedocles (approx. 490–430 BC): "For it is impossible for anything to come to be from what is not, and it cannot be brought about or heard of that what is should be utterly destroyed."

A further principle of conservation was stated by Epicurus (341–270 BC) who, describing the nature of the Universe, wrote that "the totality of things was always such as it is now, and always will be".[5]

Jain philosophy, a non-creationist philosophy based on the teachings of Mahavira (6th century BC),[6] states that the universe and its constituents such as matter cannot be destroyed or created. The Jain text Tattvarthasutra (2nd century AD) states that a substance is permanent, but its modes are characterised by creation and destruction. A principle of the conservation of matter was also stated by Nasīr al-Dīn al-Tūsī (1201–1274). He wrote that "A body of matter cannot disappear completely. It only changes its form, condition, composition, color and other properties and turns into a different complex or elementary matter".

Mass conservation in chemistry[edit]


The principle of conservation of mass was first outlined by Mikhail Lomonosov (1711–1765) in 1748. He proved it by experiments—though this is sometimes challenged.[9]Antoine Lavoisier (1743–1794) had expressed these ideas in 1774. Others whose ideas pre-dated the work of Lavoisier include Joseph Black (1728–1799), Henry Cavendish(1731–1810), and Jean Rey (1583–1645).

The conservation of mass was obscure for millennia because of the buoyancy effect of the Earth's atmosphere on the weight of gases. For example, a piece of wood weighs less after burning; this seemed to suggest that some of its mass disappears, or is transformed or lost. This was not disproved until careful experiments were performed in which chemical reactions such as rusting were allowed to take place in sealed glass ampoules; it was found that the chemical reaction did not change the weight of the sealed container and its contents. The vacuum pump also enabled the weighing of gases using scales.

Once understood, the conservation of mass was of great importance in progressing from alchemy to modern chemistry. Once early chemists realized that chemical substances never disappeared but were only transformed into other substances with the same weight, these scientists could for the first time embark on quantitative studies of the transformations of substances. The idea of mass conservation plus a surmise that certain "elemental substances" also could not be transformed into others by chemical reactions, in turn led to an understanding of chemical elements, as well as the idea that all chemical processes and transformations (such as burning and metabolic reactions) are reactions between invariant amounts or weights of these chemical elements.

Following the pioneering work of Lavoisier the prolonged and exhaustive experiments of Jean Stas supported the strict accuracy of this law in chemical reactions, even though they were carried out with other intentions. His research indicated that in certain reactions the loss or gain could not have been more than from 2 to 4 parts in 100,000. The difference in the accuracy aimed at and attained by Lavoisier on the one hand, and by Morley and Stas on the other, is enormous.



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