For example, indium tribromide InBr 3 and lanthanide tribromide LnBr 3 are all high-melting-point solids that are quite soluble in water. As the oxidation state of the metal increases, the covalent character of the corresponding metal halides also increases due to polarization of the M—X bond.
All halogens react vigorously with hydrogen to give the hydrogen halides HX. As a result, liquid HF is a polar solvent that is similar in some ways to water and liquid ammonia; after a reaction, the products can be recovered simply by evaporating the HF solvent. Hydrogen fluoride must be handled with extreme caution, however, because contact of HF with skin causes extraordinarily painful burns that are slow to heal.
Glass etched with hydrogen flouride. Except for fluorine, all the halogens react with water in a disproportionation reaction, where X is Cl, Br, or I:.
The acid strengths of the oxoacids of the halogens increase with increasing oxidation state, whereas their stability and acid strength decrease down the group. Although all the oxoacids are strong oxidants, some, such as HClO 4 , react rather slowly at low temperatures.
Consequently, mixtures of the halogen oxoacids or oxoanions with organic compounds are potentially explosive if they are heated or even agitated mechanically to initiate the reaction. Because of the danger of explosions, oxoacids and oxoanions of the halogens should never be allowed to come into contact with organic compounds.
In all cases, the heavier halogen, which has the lower electronegativity, is the central atom. The maximum oxidation state and the number of terminal halogens increase smoothly as the ionization energy of the central halogen decreases and the electronegativity of the terminal halogen increases. The interhalogen compounds are also potent oxidants and strong fluorinating agents; contact with organic materials or water can result in an explosion.
Classify the type of reaction. Using periodic trends in atomic properties, thermodynamics, and kinetics, explain why the observed reaction products form. The halogens are highly reactive. All halogens have relatively high ionization energies, and the acid strength and oxidizing power of their oxoacids decreases down the group.
Their chemistry is exclusively that of nonmetals. Consistent with periodic trends, ionization energies decrease down the group. Fluorine, the most reactive element in the periodic table, has a low F—F bond dissociation energy due to repulsions between lone pairs of electrons on adjacent atoms. Fluorine forms ionic compounds with electropositive elements and covalent compounds with less electropositive elements and metals in high oxidation states. All the halogens react with hydrogen to produce hydrogen halides.
Except for F 2 , all react with water to form oxoacids, including the perhalic acids, which contain the halogens in their highest oxidation state. Halogens also form interhalogen compounds; the heavier halogen, with the lower electronegativity, is the central atom. Learning Objectives To understand the periodic trends and reactivity of the group 17 elements: the halogens. Reactions and Compounds of the Halogens Fluorine is the most reactive element in the periodic table, forming compounds with every other element except helium, neon, and argon.
There are three reasons for the high reactivity of fluorine: Because fluorine is so electronegative, it is able to remove or at least share the valence electrons of virtually any other element. These five toxic, non-metallic elements make up Group 17 of the periodic table and consist of: fluorine F , chlorine Cl , bromine Br , iodine I , and astatine At. Although astatine is radioactive and only has short-lived isotopes, it behaves similar to iodine and is often included in the halogen group.
Because the halogen elements have seven valence electrons, they only require one additional electron to form a full octet. This characteristic makes them more reactive than other non-metal groups.
The bonds in these diatomic molecules are non-polar covalent single bonds. However, halogens readily combine with most elements and are never seen uncombined in nature.
As a general rule, fluorine is the most reactive halogen and astatine is the least reactive. All halogens form Group 1 salts with similar properties. In these compounds, halogens are present as halide anions with charge of -1 e.
Replacing the -ine ending with an -ide ending indicates the presence of halide anions; for example, Cl - is named "chloride. Therefore, most of the chemical reactions that involve halogens are oxidation-reduction reactions in aqueous solution. The halogens often form single bonds, when in the -1 oxidation state, with carbon or nitrogen in organic compounds.
When a halogen atom is substituted for a covalently-bonded hydrogen atom in an organic compound, the prefix halo- can be used in a general sense, or the prefixes fluoro- , chloro- , bromo- , or iodo- can be used for specific halogen substitutions. Halogen elements can cross-link to form diatomic molecules with polar covalent single bonds. Chlorine Cl 2 was the first halogen to be discovered in , followed by iodine I 2 , bromine Br 2 , fluorine F 2 , and astatine At, discovered last in The name "halogen" is derived from the Greek roots hal- "salt" and -gen "to form".
Together these words combine to mean "salt former", referencing the fact that halogens form salts when they react with metals. Halite is the mineral name for rock salt, a natural mineral consisting essentially of sodium chloride NaCl. Lastly, the halogens are also relevant in daily life, whether it be the fluoride that goes in toothpaste, the chlorine that disinfects drinking water, or the iodine that facilitates the production of thyroid hormones in one's body.
Fluorine - Fluorine has an atomic number of 9 and is denoted by the symbol F. Elemental fluorine was first discovered in by isolating it from hydrofluoric acid. Fluorine exists as a diatomic molecule in its free state F 2 and is the most abundant halogen found in the Earth's crust.
Fluorine is the most electronegative element in the periodic table. It appears as a pale yellow gas at room temperature. Fluorine also has a relatively small atomic radius. Its oxidation state is always -1 except in its elemental, diatomic state in which its oxidation state is zero.
Fluorine is extremely reactive and reacts directly with all elements except helium He , neon Ne and argon Ar. In H 2 O solution, hydrofluoric acid HF is a weak acid. In addition, fluorine produces very powerful oxidants. For example, fluorine can react with the noble gas xenon and form the strong oxidizing agent Xenon Difluoride XeF 2. There are many uses for fluorine, which will be discussed in Part VI of this article. Chlorine - Chlorine has the atomic number 17 and the chemical symbol Cl.
Chlorine was discovered in by extracting it from hydrochloric acid. In its elemental state, it forms the diatomic molecule Cl 2.
At room temperature it appears as a light green gas. Since the bond that forms between the two chlorine atoms is weak, the Cl 2 molecule is very reactive. Chlorine reacts with metals to produce salts called chlorides. Chloride ions are the most abundant ions that dissolve in the ocean. Chlorine also has two isotopes: 35 Cl and 37 Cl. Sodium chloride is the most prevalent compound of the chlorides. Bromine - Bromine has an atomic number of 35 with a symbol of Br. It was first discovered in In its elemental form, it is the diatomic molecule Br 2.
At room temperature, bromine is a reddish- brown liquid. Bromine is more reactive than iodine, but not as reactive as chlorine. Also, bromine has two isotopes: 79 Br and 81 Br. Bromine consists of bromide salts, which have been found in the sea. The world production of bromide has increased significantly over the years, due to its access and longer existence.
Like all of the other halogens, bromine is an oxidizing agent, and is very toxic. Iodine - Iodine has the atomic number 53 and symbol I. Iodine exists as a diatomic molecule, I 2 , in its elemental state. At room temperature, it appears as a violet solid. Iodine has one stable isotope: I.
It was first discovered in through the use of seaweed and sulfuric acid. Currently, iodide ions can be isolated in seawater. Although iodine is not very soluble in water, the solubility may increase if particular iodides are mixed in the solution. Iodine has many important roles in life, including thyroid hormone production.
This will be discussed in Part VI of the text. Astatine - Astatine is a radioactive element with an atomic number of 85 and symbol At. It is the only halogen that is not a diatomic molecule and it appears as a black, metallic solid at room temperature. Astatine is a very rare element, so there is not that much known about this element. In addition, astatine has a very short radioactive half-life , no longer than a couple of hours.
It was discovered in by synthesis. If your number of, number of protons, and this is for an atom or molecule. A molecule's just a bunch of, a bunch of atoms bonded together. If the number of protons does not equal the number of electrons.
And you can have positive ions if the protons are more than the number of electrons, protons, or positive electrons or negative. And you can have negative ions if the number of electrons are greater than the number of protons.
For example, for example, if you just had Hydrogen in it's neutral state has one proton and one electron, but if you were to take one of those electrons away then Hydrogen would have a positive charge and essentially it would just be, in its most common isotope it would just be a proton by itself.
And so, when we talk about a positive ion like this where our protons are more than our electrons, the number of protons are more than the number of electrons, we call these cations, cations. Cation, once again, just another word positive ion. Likewise, we can have negative ions. So, say for example, Fluorine. So, Fluorine gains an electron, it's going to have a negative charge.
It's gonna have a negative charge of negative one, and a negative ion we call an anion. And the way that I remember this is a kind of means the opposite or the negation of something. So, this is a negative ion. We've negating, you can somehow think we are negating the ion. So, with that out of the way, let's think about how hard it will be ionize different elements in the periodic table. In particular, how hard it is to turn them into cations.
And to think about that, we'll introduce an idea called ionization energy. Ionization energy And this is defined, this is defined as the energy required, energy required to remove an electron, to remove an electron.
So, it could've even been called cationization energy because you really see energy required to remove an electron and make the overall atom more positive. So, let's think about the trends. And we already have a little bit of background on the different groups of the periodic table. So, for example, if we were to focus on, especially we could look at group one, and we've already talked about how Hydrogen's a bit of a special case in group one but if we look at everything below Hydrogen.
If we look at the Alkali, if we look at the Alkali metals here we've already talked about the fact that these are very willing to lose an electron.
Because if they lose an electron they get to the electron configuration of the noble gas before it. So, if Lithium loses an electron then it has an outer shell electron configuration of Helium. It has two outer electrons and that's kind of, we typically talk about the Octet Rule but if we're talking about characters like Lithium or Helium they're happy with two 'cause you can only put two electrons in that first shell.
But all the rest of 'em, Sodium, Potassium, etc. Lithium, if you remove an electron, it would get to Helium and it would have two electrons in its outer shell. So, you can imagine that the ionization energy right over here, the energy required to remove electrons from your Alkali Metals is very low. So, let me just write down this is So, when I say low, I'm talking about low ionization energy. Now, what happens as we move to the right of the periodic table?
0コメント