Ionic Bonding
The mass of a single chlorine atom =35g/6.023×10^23 = 5.81×10–23g The mass of a single hydrogen atom= the mass of a carbon-12 isotope of 1/12 = 1.66×10^-24g Therefore,the average mass of a single chlorine atom= 5.81×10^-23g/1.66×10^-24g= 35.5 304 views.
- Chlorine is a chemical element with atomic number 17 which means there are 17 protons in its nucleus. Total number of protons in the nucleus is called the atomic number of the atom and is given the symbol Z. The total electrical charge of the nucleus is therefore +Ze, where e (elementary charge) equals to 1,602 x 10-19 coulombs.
- Chlorine is a commonly used household cleaner and disinfectant. Chlorine is a potent irritant to the eyes, the upper respiratory tract, and lungs. Chronic (long-term) exposure to chlorine gas in workers has resulted in respiratory effects, including eye and throat irritation and airflow obstruction. No information is available on the carcinogenic effects of chlorine in humans from inhalation.
Soduim atom'>We have learned that atoms tend to react in ways that create a full valence shell, but what does this mean? Sodium (Na), for example, has one electron in its valence shell. This is an unstable state because that valence shell needs eight electrons to be full. In order to fill its valence shell, sodium has two options:
- Find a way to add seven electrons to its valence shell, or
- Give up that one electron so the next lower energy shell (already full) could become its new valence shell.
Which do you think is more easily accomplished?
Right, giving up one electron! Atoms like sodium, with only one or two electrons in a valence shell that needs eight electrons, are most likely to give up their valence electrons to achieve a stable state. All these atoms need is another atom that can attract their electrons!
An atom such as chlorine (Cl), that contains seven electrons in its valence shell, needs one more electron to have a full valence shell. If we use the same logic as we did for sodium, we should conclude that chlorine would rather gain one more electron than lose all seven of its valence electrons to achieve stability. Under the right conditions, atoms like chlorine will steal an electron from nearby atoms like sodium. This ability to pull electrons away from other atoms is termed electronegativity. Atoms with a valence shell that is almost full are more likely to be electronegative because they have greater reason to pull electrons towards them. Electronegative atoms are not negatively charged, but they are more likely to become negatively charged.
When an electron moves from one atom to another, both atoms become ions. An ion is any atom that has gained electrons to have a negative charge (anion) or lost electrons to have a positive charge (cation). An easy way to remember that a cation has a positive, or +, charge is to think of the letter t in the word cation as a + sign.
Once an atom becomes an ion, it has an electrical charge. Ions of opposite charge are attracted to one another, forming a chemical bond, an association formed by attraction between two atoms. This type of chemical bond is called an ionic bond because the bond formed between two ions of opposite charge. The sodium cation (Na+) and the chlorine anion (Cl-) are attracted to one another to form sodium chloride, or table salt.
Although ionic bonds are very strong, they can be relatively easily broken if another attractive ion (or polar molecule) comes around. An ionic bond is formed when two ions of opposite charge come together by attraction, NOT when an electron is transferred.
Think of ionic bond formation as the second step in a two step process:
- Two atoms each become ions. The atoms could have become ions in previous reactions with other atoms or the atoms may have reacted with each other, transferring the electron(s) from one to the other.
- The two ions of opposite charge “see” one another and become attracted enough to form the bond.
Ionic Bonding
This video visually illustrates how atoms form ionic bonds.
Create an Ionic Bond
In this activity, you'll create cations and anions and watch the ionic bond form.
Covalent Bonding
Carbon atom'>In ionic bonding, we looked at atoms with either one or two electrons in their valence shell and atoms that only needed one or two electrons to fill their valence shell. What happens when an atom, carbon (C) for example, has four valence electrons? Carbon would need to either lose four electrons or gain four electrons in order to have a full valence shell. Both of these situations would cause carbon to have a very strong charge, which would probably make it as unstable as having an incomplete valence shell! For atoms like carbon, there is another option: sharing.
When two atoms each need additional electrons to fill their valence shells, but neither is electronegative enough to steal electrons from the other, they can form another kind of chemical bond called a covalent bond. In covalent bonds, two atoms move close enough to share some electrons. The electrons from each atom shift to spend time moving around both atomic nuclei.
In the most common form of covalent bond, a single covalent bond, two electrons are shared, one from each atom’s valence shell. Double covalent bonds where four electrons are shared, and triple covalent bonds where six electrons are shared, are also commonly found in nature.
How do we know if and when covalent bonds will form? Atoms will form as many covalent bonds as it takes to fill their valence shell. This means that carbon, our previous example, will need to form four covalent bonds in order to fill its outermost shell. In each of the four bonds, carbon will contribute one electron and the other atom will contribute one electron, supplying carbon with eight electrons effectively orbiting its nucleus. Atoms like oxygen (O) will form two covalent bonds because they already have six valence electrons and only need two more electrons obtained by sharing. Put another way, oxygen shares two of its six valence electrons in covalent bonds while keeping four valence electrons for itself (4 unshared oxygen electrons + 2 shared oxygen electrons + 2 shared electrons from the other atoms = 8 total electrons).
In shell models, the shared electrons are shown within the overlapping region of the valence shells to represent the fact they are shared, but the electrons are actually moving around both nuclei and could be found anywhere around either nucleus at a given time.
For simplicity, we often draw covalent bonds as straight lines between atoms to represent the structural formula. Each line represents a single covalent bond (two shared electrons), so double lines represent a double covalent bond (four shared electrons).
Covalent bonds are usually found in atoms that have at least two, and usually less than seven, electrons in their outermost energy shell, but this is NOT a hard and fast rule. Hydrogen, for example, is a unique atom that bears closer examination. Because hydrogen only has one electron surrounding its nucleus, its valence shell is the first energy shell, which only needs two electrons to be full. Hydrogen tends to form covalent bonds because a single covalent bond will fill its shell. However, the one-proton nucleus is very weak and has trouble keeping the shared electrons around the hydrogen atom for very long.
Covalent Bonding
This video visually illustrates how atoms form covalent bonds.
Other Kinds of Bonding
When a covalent bond forms between atoms of similar electronegativity, the shared electrons tend to spend equal time around each nucleus. What happens if a bond is formed between atoms like oxygen, which is highly electronegative, and hydrogen, which is not? The oxygen atom tends to pull the shared electrons over to its side of the bond more often than the hydrogen atom does, resulting in polarity, or a partial separation of charge between atoms. The electrons don’t actually leave the less electronegative atom, but they do spend less time on that side. This results in two poles, one slightly positive and one slightly negative. This is called a polar bond, while covalent bonds where electrons are shared equally are called nonpolar.
Polar covalent bonds are the source of an additional kind of association called a hydrogen bond. In a hydrogen bond, the partially positive end of a polar covalent molecule is attracted to the partially negative end of another polar covalent molecule. For example, water is composed of two hydrogen atoms covalently bonded to a single oxygen atom. In a container full of water molecules, the hydrogen atoms of each water molecule are attracted to the oxygen atoms of other water molecules, forming hydrogen bonds between all of the molecules in the container of water. Hydrogen bonds are weak compared to covalent bonds, but they are strong enough to affect the behavior of the atoms involved. This leads to many important chemical properties in water and other molecules.
Chlorine is a greenish yellow gas at room temperature and atmospheric pressure. It is two and a half times heavier than air. It becomes a liquid at −34 °C (−29 °F). It has a choking smell, and inhalation causes suffocation, constriction of the chest, tightness in the throat, and—after severe exposure—edema (filling with fluid) of the lungs. As little as one part per thousand in air causes death within a few minutes, but less than one part per million may be tolerated. Chlorine was the first gas used in chemical warfare in World War I. The gas is easily liquefied by cooling or by pressures of a few atmospheres at ordinary temperature.
Chlorine has a high electronegativity and a high electron affinity, the latter being even slightly higher than that of fluorine. The affinity of chlorine for hydrogen is so great that the reaction proceeds with explosive violence in light, as in the following equation (where hν is light):
In the presence of charcoal, the combination of chlorine and hydrogen takes place rapidly (but without explosion) in the dark. A jet of hydrogen will burn in chlorine with a silvery flame. Its high affinity for hydrogen allows chlorine to react with many compounds containing hydrogen. Chlorine reacts with hydrocarbons, for example, substituting chlorine atoms for the hydrogen atoms successively. If the hydrocarbon is unsaturated, however, chlorine atoms readily add to the double or triple bond.
Chlorine Atomic Mass
Chlorine molecules are composed of two atoms (Cl2). Chlorine combines with almost all the elements, except for the lighter noble gases, to give chlorides; those of most metals are ionic crystals, whereas those of the semimetals and nonmetals are predominantly molecular.
The products of reaction with chlorine usually are chlorides with high oxidation numbers, such as iron trichloride (FeCl3), tin tetrachloride (SnCl4), or antimony pentachloride (SbCl5), but it should be noted that the chloride of highest oxidation number of a particular element is frequently in a lower oxidation state than the fluoride with the highest oxidation number. Thus, vanadium forms a pentafluoride, whereas the pentachloride is unknown, and sulfur gives a hexafluoride but no hexachloride. With sulfur, even the tetrachloride is unstable.
Aside from the −1 oxidation states of some chlorides, chlorine exhibits +1, +3, +5, and +7 oxidation states, respectively, in the following ions: hypochlorite (ClO−), chlorite (ClO−2), chlorate (ClO−3), and perchlorate (ClO−4). Five oxides—chlorine monoxide (Cl2O), chlorine dioxide (ClO2), chlorine perchlorate (Cl2O4), dichlorine hexoxide (Cl2O6), and dichlorine heptoxide (Cl2O7)—all highly reactive and unstable, have been indirectly synthesized. Chlorine can undergo addition or substitution reactions with organic compounds.
Chlorine displaces the heavier, less electronegative halogens, bromine and iodine, from compounds. The displacement of bromides, for example, occurs according to the following equation:
Chlorine Atomic
Furthermore, it converts several oxides into chlorides. An example is the conversion of iron trioxide to the corresponding chloride:
Chlorine is moderately soluble in water, yielding chlorine water, and from this solution a solid hydrate of ideal composition, Cl2∙7.66H2O, is obtained. This hydrate is characterized by a structure that is more open than that of ice; the unit cell contains 46 molecules of water and 6 cavities suitable for the chlorine molecules. When the hydrate stands, disproportionation takes place; that is, one chlorine atom in the molecule is oxidized, and the other is reduced. At the same time, the solution becomes acidic, as shown in the following equation:
in which the oxidation numbers are written above the atomic symbols. Chlorine water loses its efficiency as an oxidizing agent on standing, because hypochlorous acid gradually decomposes. The reaction of chlorine with alkaline solutions yields salts of oxyacids.
Chlorine Atom Info
The first ionization energy of chlorine is high. Although ions in positive oxidation states are not very stable, high oxidation numbers are stabilized by coordination, mainly with oxygen and fluorine. In such compounds bonding is predominantly covalent, and chlorine is capable of exhibiting the oxidation numbers +1, +3, +4, +5, +6, and +7.
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