Phyiscal Properties of Element in the Halogen Family
Group 17: Physical Backdrop of the Halogens
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Some chemical and physical properties of the halogens are summarized in Table \(\PageIndex{1}\). It can be seen that at that place is a regular increment in many of the properties of the halogens proceeding down group 17 from fluorine to iodine. This includes their melting points, boiling points, intensity of their color, the radius of the corresponding halide ion, and the density of the element. On the other manus, in that location is a regular decrease in the first ionization energy as nosotros go downwardly this group. Equally a outcome, there is a regular increase in the ability to form high oxidation states and a subtract in the oxidizing strength of the halogens from fluorine to iodine.
| Property | F | Cl | Br | I |
|---|---|---|---|---|
| Atomic number, Z | ix | 17 | 35 | 53 |
| Basis country electronic configuration | [He]2s2 2pfive | [Ne]3s2 3p5 | [Ar]3dx 4s2 4p5 | [Kr]4d10 5s2 5p5 |
| color | pale yellow gas | yellow-green gas | red-brown liquid | blue-blackness solid |
| Density of liquids at diverse temperatures, /kg m-3 | 1.51 (85 °K) | 1.66 (203 °K) | 3.19 (273 °K) | 3.96 (393 °K) |
| Melting point, /K | 53.53 | 171.6 | 265.eight | 386.85 |
| Boiling signal, /Chiliad | 85.01 | 239.18 | 331.93 | 457.five |
| Enthalpy of atomization, ΔaH° (298K) / kJ mol-one | 79.08 | 121.viii | 111.7 | 106.vii |
| Standard enthalpy of fusion of Ten2, ΔfusH°(mp) / kJ mol-ane | 0.51 | half-dozen.4 | 10.57 | fifteen.52 |
| Standard enthalpy of vaporization of X2, ΔvapH°(bp) / kJ mol-1 | six.62 | 20.41 | 29.96 | 41.57 |
| Starting time ionization energy, IEi / kJ mol-one | 1681 | 1251.1 | 1139.ix | 1008.iv |
| ΔEAHane°(298K) / kJ mol-one | -333 | -348 | -324 | -295 |
| ΔhydH°(Ten-,k) / kJ mol-1 | -504 | -361 | -330 | -285 |
| ΔhydDue south°(Ten-,yard) / JK-one mol-i | -150 | -90 | -seventy | -50 |
| ΔhydK°(X-,m) / kJ mol-1 | -459 | -334 | -309 | -270 |
| Standard redox potential, E°(Xtwo /2X-) /V | ii.87 | 1.36 | 1.09 | 0.54 |
| Covalent radius, rcov = ½ X-Ten bond length /pm | 72 | 100 | 114.ii | 133.3 |
| Ionic radius, rion for X- /pm | 133 | 181 | 196 | 220 |
| van der Waals radius, rfive /pm | 135 | 180 | 195 | 215 |
| X-X(one thousand)bond energy /kJ mol-one | 159 | 243 | 193 | 151 |
| H-X(k)bond free energy /kJ mol-ane | 562 | 431 | 366 | 299 |
| C-X(k)bond energy /kJ mol-1 | 484 | 338 | 276 | 238 |
| Pauling electronegativity, χP | 3.98 | 3.sixteen | ii.96 | 2.66 |
Color
The origin of the color of the halogens stems from the excitation betwixt the highest occupied π* Molecular Orbital and the lowest unoccupied σ* Molecular Orbital. The energy gap between the HOMO and LUMO decreases according to F2 > Cl2 > Br2 > Itwo. The amount of energy required for excitation depends upon the size of the atom. Fluorine is the smallest element in the group and the strength of attraction between the nucleus and the outer electrons is very large. As a issue, it requires a large excitation free energy and absorbs violet light (loftier free energy) and and so appears pale yellow. On the other hand, iodine needs significantly less excitation energy and absorbs yellow low-cal of low free energy. Thus information technology appears night violet. Using similar arguments, it is possible to explain the greenish yellowish color of chlorine and the reddish brown color of bromine.
Effigy \(\PageIndex{ane}\): Molecular Orbital diagram for fluorine.
The halogens show a diversity of colors when dissolved in different solvents. Solutions of iodine can exist bright violet in CCl4, pinkish or cherry-red brown in aromatic hydrocarbons and deep brown in alcohols for instance. This tin be explained by weak donor-acceptor interaction and complex formation. The presence of charge-transfer bands farther supports this since they are thought to be derived from interaction with the HOMO σu* orbital.
The X-ray construction of some of these take been obtained and often the intense color can exist used for label and determination such as the brilliant blueish colour of iodine in the presence of starch. In the example of the solid formed between dibromine and benzene, the structure is shown beneath and a new charge transfer band occurs at 292 nm. The Br-Br bond length is substantially unchanged from that of dibromine (228 pm).
Figure \(\PageIndex{ii}\):
In a study of the reaction of dibromine with substituted phosphines in diethyl ether, all but 1 showed a tetrahedral arrangement where one bromine was linked to the phosphorus.[three]
\[R_3P + Br_2 (Et_2O, N_2/r.t.) rightarrow R_3PBr_2 \label{i}\]
The X-ray study of the triethylphosphine was interpreted as [EtiiiPBr]Br where the Br-Br separation was 330 pm. This is considerably longer than the 228 pm establish above and was taken to mean that the compound was ionic In the case of the tri(perfluorophenyl)phosphine notwithstanding the structure showed both bromines linked to give a trigonal bipyramid arrangement with D3 symmetry. Why (Chalf-dozenFfive)threePBrtwo was the merely R3PBr2 compound that adopted trigonal bipyramidal geometry was reasoned to be due to the very low basicity of the parent tertiary phosphine.
Melting and Boiling Points
Intermolecular forces are the attractive forces between molecules without which all substances would be gases. The various types of these interactions bridge large differences in energy and for the halogens and interhalogens are generally quite pocket-size. The dispersion forces involved in these cases are called London forces (afterward Fritz Wolfgang London, 1900-1954). They are derived from momentary oscillations of electron accuse in atoms and hence are nowadays betwixt all particles (atoms, ions and molecules).
The ease with which the electron cloud of an atom can exist distorted to go asymmetric is termed the molecule's polarizability. The greater the number of electrons an atom has, the farther the outer electrons volition be from the nucleus, and the greater the chance for them to shift positions inside the molecule. This means that larger nonpolar molecules tend to accept stronger London dispersion forces. This is evident when considering the diatomic elements in group 17, the Halogens. All of these diatomic elements are nonpolar, covalently bonded molecules. Descending the grouping, fluorine and chlorine are gases, bromine is a liquid, and iodine is a solid. For nonpolar molecules, the farther y'all become downwards the grouping, the stronger the London dispersion forces.
To picture show how this occurs, compare the situation 1) where the electrons are evenly distributed and so consider 2) an instantaneous dipole that would arise from an uneven distribution of electrons on i side of the nucleus. When two molecules are shut together, the instantaneous dipole of one molecule tin induce a dipole in the second molecule. This results in synchronized motion of the electrons and an attraction between them. 3) Multiply this consequence over numerous molecules and the overall result is that the allure keeps these molecules together, and for diiodine is sufficient to make this a solid.
Figure \(\PageIndex{3}\): On average the electron deject for molecules can exist considered to be spherical in shape. When two not-polar molecules approach, attractions or repulsions between the electrons and nuclei tin can lead to distortions in their electron clouds (i.due east. dipoles are induced). When more molecules interact these induced dipoles lead to intermolecular allure.
The changes seen in the variation of MP and BP for the dihalogens and binary interhalogens van be attributed to the increment in the London dispersion forces of attraction betwixt the molecules. In full general they increase with increasing atomic number.
Figure 4:
Redox Properties
The almost characteristic chemical characteristic of the halogens is their oxidizing strength. Fluorine has the strongest oxidizing ability, then that a simple chemical grooming is almost impossible and it must be prepared by electrolysis. Annotation that since fluorine reacts explosively with h2o oxidizing it to dioxygen, finding reaction weather for any reaction can be difficult. When fluorine is combined with other elements they generally exhibit loftier oxidation states. Chlorine is the next strongest oxidizing agent, just it tin can exist prepared past chemic oxidation. Most elements react direct with chlorine, bromine and iodine, with decreasing reactivity going down the Grouping, but often the reaction must be activated past estrus or UV low-cal. [2] The free energy changes in redox process are:
- Enthalpy of atomization,
- ΔEAH1,
- ΔhydH°(Ten-,g)
The redox potential, Due east°, 102/2X-, measures a free-energy alter, commonly dominated by the ΔH term. The values in the Table evidence that there is a subtract in oxidizing forcefulness proceeding down the group (2.87, ane.36, one.09, 0.54 V). This can exist explained by comparing the steps shown higher up.
- 1) atomization of the dihalide is the energy required to suspension the molecule into atoms
\[½ X_{2(g)} \rightarrow X_{(g)} \label{2}\]
annotation that only F2 and Cl2 are gases in their natural land and then the energies associated with atomization of Brtwo and I2 requires converting the liquid or solid to gas first.
- ii) ΔEAH1 is the energy liberated when the cantlet is converted into a negative ion and is related to the Electron Analogousness
\[X_{(grand)} + e^- \rightarrow 10^-_{(g)} \label{3}\]
Addition of an electron to the small F cantlet is accompanied by larger e-/e- repulsion than is institute for the larger Cl, Br or I atoms. This would suggest that the process for F should be less exothermic than for Cl and not fit the tendency that shows a general decrease going down the grouping.
- 3) ΔhydH°(X-,yard) is the free energy liberated on the hydration of the ion, the Hydration energy.
\[X^-_{(m)} + H_2O \rightarrow 10^-_{(aq)} \label{four}\]
The overall reaction is so:
\[½ X_{two(one thousand)} \rightarrow Ten^-_{(aq)} \characterization{5}\]
| Halogen | atomization energy (kJ mol-1) | ΔEAH1 (kJ mol-1) | hydration enthalpy (kJ mol-i) | overall (kJ mol-i) |
|---|---|---|---|---|
| F | +79.08 | -333 | -504 | -758 |
| Cl | +121.8 | -348 | -361 | -587 |
| Br | +111.7 | -324 | -330 | -542 |
| I | +106.seven | -295 | -285 | -473 |
This shows a very negative energy change for the fluoride compared to the others in the grouping. This comes near because of two main factors: the high hydration energy and the low atomization energy. For Ftwo 2) is less than for Cl2 just since the energy needed to break the F-F bond is also less and the hydration more, the full energy drop is much greater. In spite of their lower atomization energies, Br2 and I2 are weaker oxidizing agents than Cl2 and this is due to their smaller ΔEAHi and smaller ΔhydH°.
Information technology can be seen that the ΔEAH1 value for fluorine is in betwixt those for chlorine and bromine and so this value alone does non provide a good explanation for the observed variation.
Each of the halogens is able to oxidize whatever of the heavier halogens situated below it in the grouping. They can oxidize hydrogen and nonmetals such as:
\[X_2 + H_{two(g)} \rightarrow 2HX_{(g)} \label{six}\]
In h2o, the halogens asymmetric according to:
\[X_2 + H_2O_{(fifty)} \rightarrow HX_{(aq)} + HXO_{(aq)} \characterization{7}\]
where \(X=Cl, Br, I\). When base of operations is added so the reaction goes to completion forming hypohalites, or at higher temperatures, halates, for example heating dichlorine:
\[3Cl_{2(g)} + 6OH^-_{(aq)} \rightarrow ClO^-_{3(aq)} + 5Cl^-_{(aq)} + 3H_2O(l) \label{8}\]
Starting time Ionization Energies
The trend seen for the consummate removal of an electron from the gaseous halogen atoms is that fluorine has the highest IE1 and iodine the lowest. To overcome the attractive strength of the nucleus means that energy is required and so that the Ionisation Energies are all positive. The variation with size can exist explained since as the size increases information technology take less free energy to remove an electron. This inverse relationship is seen for all the groups, not only group 17. As the distance from the nucleus to the outermost electrons increases, the allure decreases so that those electrons are easier to remove. The high value of IEi for Fluorine is such that information technology does not exhibit whatsoever positive oxidation states, whereas Cl, Br and I can be equally loftier every bit seven.
Oxidation states
Fluorine is the near electronegative element in the periodic table and exists in all its compounds in either the -one or 0 oxidation state. Chlorine, bromine, and iodine however can be found in a range of oxidation states including: +i, +3, +5, and +7, equally shown below.
| Oxidation States | Examples |
|---|---|
| -1 | CaFii, HCl, NaBr, AgI |
| 0 | Fii, Cltwo, Br2, I2 |
| 1 | HClO, ClF |
| three | HClO2, ClFthree |
| 5 | HClO3, BrFv, [BrFsix]-, IFv |
| 7 | HClOiv, BrFhalf-dozen +, IF7, [IFeight]- |
In full general, odd numbered groups (similar group 17) form odd-numbered oxidation states and this can be explained since all stable molecules contain paired electrons (free radicals are obviously much more reactive). When covalent bonds are formed or broken two electrons are involved so the oxidation state changes by 2.
When difluorine reacts with diiodine initially iodine monofluoride is formed.
\[I_2 + F_2 \rightarrow 2IF\]
Calculation a 2nd difluorine uses ii more iodine valence electrons to form two more than bonds:
\[2IF + F_2 \rightarrow IF_3\]
more than on this in the next lecture
References
1. "Inorganic Chemistry" - C. Housecroft and A.G. Sharpe, Prentice Hall, 3rd Ed., December 2007, ISBN13: 978-0131755536, ISBN10: 0131755536, Chapter 17.
2. "Chemistry. The Molecular Nature of Thing and Modify" - M.S. Silberberg, McGraw Hill Higher Education, quaternary Ed., 2006, ISBN13: 978-0072558203, Capacity viii, 12 and 14.
3. Stephen M. Godfrey, Charles A. McAuliffe, Imran Mushtaq, Robin G. Pritchard and Joanne Thou. Sheffield, J. Chem. Soc., Dalton Trans., 1998, 3815-3818
Source: https://chem.libretexts.org/Bookshelves/Inorganic_Chemistry/Supplemental_Modules_and_Websites_%28Inorganic_Chemistry%29/Descriptive_Chemistry/Elements_Organized_by_Block/2_p-Block_Elements/Group_17:_The_Halogens/0Group_17:_Physical_Properties_of_the_Halogens
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