Category: Chemistry Notes PPT

The elements in group 17 are the halogens. The Halogen elements are fluorine, chlorine, bromine, iodine, and astatine. These elements are too reactive to occur freely in nature, but their compounds are cosmopolitan. Chlorides are the foremost abundant; although fluorides, bromides, and iodides are less common, they’re reasonably available. During this section, we’ll examine the occurrence, preparation, and properties of halogens. Next, we’ll test halogen compounds with the representative metals followed by an examination of the interhalogens. This section will conclude with some applications of halogens.

Halogen Group Periodic Table

Properties of the Halogens

Halogen properties of fluorine could be a pale-yellow gas, chlorine could be a greenish-yellow gas, bromine could be a deep reddish-brown liquid, and iodine could be a grayish-black crystalline solid. Liquid bromine contains a high-pressure level, and therefore the reddish vapour is quickly visible in. Iodine crystals have a standard pressure level. When gently heated, these crystals sublime and form a good-looking deep violet vapour.

Chlorine could be a pale yellow-green gas (left), gaseous bromine is deep orange (centre), and gaseous iodine is purple (right). Fluorine is so reactive that it’s too dangerous to handle.

Bromine is merely slightly soluble in water, but it’s miscible altogether proportions in less polar (or nonpolar) solvents like chloroform, carbon tet, and compound, forming solutions that fluctuate from yellow to reddish-brown, counting on the concentration.

Iodine is soluble in chloroform, carbon tet, compound, and lots of hydrocarbons, giving violet solutions of I₂ molecules. It’s quite soluble in aqueous solutions of iodides, with which it forms brown solutions. These brown solutions result because iodine molecules have empty valence d orbitals and may act as weak Lewis acids towards the iodide ion. 

Halogen Group Numbers 

Fluorine is the most electronegative element within the tabular array. As a result, it’s a number -1 altogether. Because chlorine, bromine, and iodine are less electronegative, it’s possible to arrange compounds within which these elements have oxidation numbers of +1, +3, +5, and +7, as shown within the table below.

Standard Halogen Group Elements Number

Halogens Characteristics

• All, these elements are incredibly reactive.

• Due to the present tendency towards high reactivity, the halogens cannot exist within the environment as pure elements. They are usually found occurring as compounds or as ions.

• Most halogen ions and atoms are often found together with other compounds present within the sea or drinking water. It can be because halogen elements tend to form salt after they are available in contact with the metals and mix with them to make compounds.

• As mentioned previously, halogens are the sole elemental group within the entire tabular array, which consists of elements that belong to any or all three classical states of matter (solid, liquid, and gas). It can be proven by the fact that when kept under temperature and normal pressure, astatine and iodine take the shape of solids, bromine appears as a liquid, and chlorine and fluorine occur as gases.

• All halogen elements form hydrogen halides, which are potent acids after they combine with hydrogen and form binary compounds.

• The halogen group reacts among itself; these elements form diatomic interhalogen compounds.

• Halogens get their high tendency to react with other matter thanks to high levels of electronegativity of their atoms, which could be a result of the high sufficient nuclear charge of all halogen atoms.

Key Concepts and Summary

The halogens form halides with less electronegative elements. Halides of the metals vary from ionic to covalent; halides of non-metals are covalent. Interhalogens are created by the mixture of two or more different halogens. Minerals are directly reacted by all the representative halogen elements or with solutions of the hydrohalic acids (HF, HCl, HBr, and HI) to provide representative metal halides. Basic anions, like hydroxides, oxides, or carbonates, are involved in the addition of aqueous hydrohalic acids.

What is Helium?

Helium is the lightest noble gas that has been detected and is the only element of the Helium periodic table and was discovered by the French astronomer Pierre Janssen, who also detected a bright yellow line in the solar chromosphere spectrum during an eclipse in 1868. The helium element can be found on the top right side of the periodic table, and the atomic number of helium is 2, where it comes first amongst the noble gases family.

It held a single atomic orbital and was named by Frankland and Lockyer. The term helium is derived from the Greek word “Helios,” which means Sun. Before it was discovered, scientists knew there was an enormous amount of helium in the Sun.

Since its outermost electron orbital is full with two electrons, helium falls under inert gas. Also, helium can be found in compressed air tanks, lasers, and as a coolant in nuclear reactors. Helium holds the lowest melting and boiling points among all the other elements. The Nuclear fusion of hydrogen in stars generates helium in a significant amount.

The Helium symbol is  He.

Abundance of Helium

Helium is the second lightest and most abundant element in the universe, which is observable (where hydrogen is the most abundant and lightest). It exists about 24% of the total elemental mass, which is more than 12 times the mass of all the combined heavier elements. In both the Sun and in Jupiter, its abundance is similar to the same.

This is because of the very high nuclear binding energy (per nucleon) of helium-4, concerning the next three elements after helium. Also, Helium-4 binding energy is accountable for why it is a product of both radioactive decay and nuclear fusion. The most amount of helium in the universe is helium-4, the huge majority of which was formed at the time of the Big Bang. Excess amounts of new helium are being created by the nuclear fusion of hydrogen in stars.

The abundance of helium-3 and helium-4 is equivalent to 0.0002% and 99.9998%, respectively. This difference in abundances can be seen in the Earth’s atmosphere, where the ratio of 4He atoms to 3He atoms is nearly 1000000:1.

Isotopes of Helium

Although there are eight isotopes of helium (He) (standard atomic mass: 4.002602(2) u), which are known, from those, only helium-3 (3He) and helium-4 (4He) are stable. All the radioisotopes are short-lived, and the longest-lived is being 6He, with a half-life of

806.7 milliseconds. The least stable is 5He, with a half-life of 7.6×10−22 seconds, even though it is possible that 2He has a shorter half-life ever.

There is one 3He atom for every million 4He atoms in the Earth’s atmosphere. However, helium is an unusual element in that its isotopic abundance varies highly based on its origin. The proportion of 3He is up to a hundred times higher in the interstellar medium. Different formation processes of two stable isotopes of the helium produce different isotope abundances.

Properties of Helium

The helium element is an odorless, colourless, insipid, and non-toxic, gas. Other than any has, it is less soluble in water. It is also the less reactive element and doesn’t form chemical compounds essentially. The viscosity and density of helium vapour are very low. The thermal conductivity and the caloric content of helium are exceptionally high. Helium can be liquefied, but its condensation temperature is the lowest when compared to all the known substances. The physical and chemical properties of Helium gas are given below briefly.

The physical properties of helium are nothing but the characteristics that are seen without changing the substance into another. The physical properties of helium are tabulated below.

Physical Properties of Helium

Chemical Properties of Helium

Chemical properties are simply the characteristics that define how the element reacts with other substances or changes from one to another substance. And, the chemical properties are seen only during a chemical reaction.

The Chemical Properties of Helium are Tabulated Below.

Health Effects of Helium

Humans have no sense that they can detect the presence of Helium. Although Helium is non-toxic and inert, it can act as a simple asphyxiant by displacing oxygen in the air to the below-required levels to support life. Excessive inhalation of Helium can cause dizziness, nausea, vomiting, loss of consciousness, and death.

Death may result from confusion, errors in judgment, loss of consciousness that prevents self-rescue. At low concentrations of oxygen, unconsciousness and death may occur in seconds without any warning. Personnel, including rescue workers, should not enter areas where the concentration of oxygen is lower to 19.5% unless offered with a self-contained breathing apparatus or an air-line respirator.

Elimination rules apply to the process of taking a compound and producing a more simplified and simpler form. It can be done to reduce the number of steps, yield, purity, etc.

Mechanism of the Elimination Reaction:

When elimination takes place, one more molecule is removed and replaced by another compound of a similar or the same structure. Thus, it is known as, “reduction process”. It is always the end result that matters! It is not just a reaction that results in elimination! If the reaction will be reduced from a compound to an atom, it is known as, “elimination.”

The reaction does not take place at the bond which is to be reduced (eliminated). It occurs at a bond that is weaker and less strong than the bond which is to be reduced.

The most reactive carbon in an elimination process is the C-C bond or C-C-C bond. So, the molecule will be in a highly excited state when it takes place. A molecule that contains a high concentration of carbon is the source of energy for the elimination reaction. So, the formation of the carbocation is an important factor in the initiation of the elimination.

The position of the C-C bond to be reduced (eliminated) depends on the structure of the starting material and the product. Generally, the most reactive C-C bond is the end of a ring. Generally, the most reactive C-C bond is the end of a ring.

A structure that contains the highest concentration of energy is an advantage and is more reactive in an elimination reaction. When two or more C-C bonds are involved in the formation of a new C-C bond during elimination, one bond will have a stronger and shorter life. While the bond with a weak bond will be broken and it is eliminated.

Examples of Elimination

1. A compound of the formula

CH2 = CH – CO – CH = CH = CH CH2Cl becomes CH3Cl

It is an example of an elimination reaction with two carbons involved (in the form of two bonds that are breaking and one bond which is being formed).

2. The same formula in the 1st edition becomes COCl2 

when the compound has one carbon that is broken (CH2) and one carbon that is being formed (CH = CHCl).

3. The same formula in the 2nd edition becomes CHCl2 

4. In 1st edition C = C – C – Cl becomes CCl4

It is an example of an elimination reaction with three carbons involved (two bonds that are breaking and one bond which is being formed).

Hofmann Elimination and the Underlying Mechanism

The Hofmann elimination reaction requires a leaving group to be generated in one step. From the point of view of the mechanism, the elimination reaction is a rearrangement reaction. We will now consider the different processes possible and examine how we might predict the most likely reaction to take place. This will give an insight into the nature of the reaction.

If we look at a typical peptide, it is likely that we will observe the formation of the following intermediate.

For the purposes of this article, it is easier to consider the Hofmann elimination of a carboxylic acid (such as an alpha-amino acid) using sodium hydride. The reaction can take place using either the sodium alkoxide or hydroxide as a base, thus, we have an example of the hydroxide route or a route involving a hydroxide species as shown in. 

In either case, we have a primary alkyl alkoxide. This can be considered a nucleophile and will react with the reactive center to form a reactive species of the formula (a) shown in This species, known as the “activated species” will react with the side chain of the carboxylic acid or any other reactive center to give the intermediate shown in (b). It is important to realize that the reactive center will not remain static and has a high probability of moving around to react with the activated species. In fact, the reactive center will not react at all and the reaction will be aborted or will proceed to the other side of the molecule.

When looking at a peptide or protein, there are two ways that we can activate the side chain. The first is to activate the carboxylic acid, but before it can react with the active species, we have to generate the leaving group. The second is to generate the active species from the carboxylic acid, and in doing so we generate the leaving group and the active species.

In the context of the example, sodium methoxide is an effective base to generate the sodium hydroxide, an effective nucleophile to react with the carboxylic acid to form the reactive species. An important point here is that both methods for activation are effective, but for our purposes of describing activation of side-chain carboxyl groups in peptides and proteins, we will refer to them as the “base” and “nucleophile” route, respectively.

Hofmann elimination reactions are the elimination reactions of quaternary ammonium salts producing tertiary amines and alkenes. It is also known as Hofmann exhaustive methylation and Hofmann degradation. The products of this reaction, tertiary amines and alkenes are known as Hofmann products. Silver hydroxide and heat is used in the reaction to get the product. This reaction was given by August Wilhelm Von Hofmann in 1851.

 

Hofmann Elimination Reaction – 

()

 

Hofmann Elimination of Cyclic Amines – 

 

()

 

Hofmann Rule

According to Hofmann rule the less stable alkene will be the main product in Hofmann Elimination reactions or other such kinds of elimination reactions. This rule is used in the prediction of regioselectivity of elimination reactions. 

 

So, when you have a bulky leaving group in the elimination reaction, the least substituted alkene will be the main product. 

 

Examples of reactions which follow Hofmann rule – 

Cope elimination also follows Hofmann rule. 

 

Reaction – 

 

()

 

Reaction – 

 

()

 

Hofmann Elimination Mechanism 

Hofmann Elimination Reaction Mechanism involves E2 elimination mechanism. It involves 3 steps. The Hofmann exhaustive methylation mechanism starts with formation of the ammonium iodide salt then ammonium iodide salt reacts with silver oxide and gives silver iodide as precipitate. Deprotonation of water also takes place by silver  oxide. It results in the formation of hydroxide ions. Now, this mixture is heated. This facilitates the elimination reaction and gives rise to alkene. 

Here we have explained all three steps in detail – 

Step 1. 

Ammonium iodide salt is formed in this step. Tertiary amine reacts with methyl iodide and forms ammonium iodide salt. 

Reaction – 

()

 

Step 2. 

In this step substitution of iodide anion with hydroxide formed by deprotonation of water molecules. 

Reaction – 

()

Step 3. 

In this step the reaction mixture is heated to start the elimination reaction and to get the required product alkenes. 

Reaction – 

()

 

In the above step the tertiary amine product is also formed. 

E2 transition state has been shown in the diagram below – 

()

 

Three steps involved in the Hofmann elimination mechanism have been explained above can be explained in the following way as well – 

()

 

Uses of Hofmann Elimination Reaction 

The primary use of the Hofmann Elimination reaction is to synthesize alkenes. 

• Physiological importance of Hofmann elimination reaction – the design of skeletal muscle relaxant is superior to tubocurarine because of its less cardiac side effects and self-destruction mechanism into blood by Hofmann Elimination reaction. 

• It has uses in the pharmaceutical field as well. 

• It is used in the synthesis of anthranilic acid. 

• It is used in the formation of primary reactants involved in the production of benzene and its derivatives. 

• It is used in the production of precursors of tryptophan. 

• It is used in the production of sweetening agents.

 

Hofmann Elimination reaction is an important name reaction for CBSE Class 12 Board examination perspective. Generally, direct questions are asked related to name reactions in class 12 CBSE board exams. Although, this reaction is an important name in competitive exams point of view as well. Generally, questions based on Hofmann reactions are asked in JEE and NEET exams as well. So, you should understand this reaction nicely and should practice its mechanism many times. Still if you have any doubts refer to free PDFs of study material and NCERT Solutions of class 12 Chemistry available on website.

Hybridization is a concept used in organic chemistry to explain chemical bonding in cases where the valence bond theory does not provide satisfactory clarification. This theory is especially useful to explain the covalent bonds in organic molecules.

 

Basically, hybridization is intermixing of atomic orbitals of different shapes and nearly the same energy to give the same number of hybrid orbitals of the same shape, equal energy and orientation such that there is minimum repulsion between these hybridized orbitals.

Types of Hybridization

There are different types of hybridization-based on the mixing of the orbitals.

• sp³ hybridization: When one s orbital and three p orbital from the same shell of atom mix together to form a new equivalent orbital then this is called sp³ hybridization.

• sp² hybridization: It is observed when one s orbital and two p orbitals undergo mixing of energy for equivalent orbitals.

• sp hybridization: When one s and one p orbital goes in the process of mixing of energy to form a new orbital such kind of hybridization is called sp hybridization. The molecules possessing sp hybridization used to have a linear shape with an angle of 180°.

The above are three basic hybridizations along with them there are other hybridizations based on the mixing of orbitals such as sp³d hybridization, sp³d² hybridization and sp³d² hybridization.

Explanation of Hybridization Through Examples

Example 1:  Consider an example of the simplest hydrocarbon molecular Methane. CH₄. According to experimental observations, the Methane molecule has 4 identical C-H bonds with equal length and equal bond energy. All the four hydrogen atoms are arranged in a manner such that the four hydrogen atoms form corners of a regular tetrahedron. 

                         

                                                 Image: Structural Formula of Methane

Based on the valence theory, a covalent bond is formed between two atoms in a molecule when there is an overlapping of half-filled atomic orbitals containing unpaired electrons. In the case of the methane molecule, we first write down the electronic configuration of each atom – C and H

      Image: Electronic configuration of carbon and hydrogen for hybridization

Each carbon atom has two unpaired electrons (in the 2pₓ and 2pᵧ orbitals). Based on the valence theory, only two hydrogen molecules could be paired to the two unpaired electrons of the carbon atom and there will be a formation of only 2 C-H bonds in the molecule. This will lead to an incomplete octet in the 2nd orbital of the carbon molecule (2pz orbital is unfilled) and so the molecule should be unstable. However, we see that actually the methane molecule is extremely stable in nature and has 4 C-H bonds and not two. Thus, the valence theory doesn’t explain the covalent bond of the methane molecule. 

The hybridization concept explains the formation of identical 4 C-H bonds and the tetrahedral shape of the molecule. 

According to this concept, when a carbon atom reacts with a hydrogen atom, the electrons in the carbon atom initially go into an excited state as shown here:

    Image: Electronic configuration of carbon in the ground state and in the excited state

Post excitation, hybridization can be imagined as the process where these 4 excited s and the p orbitals combine together to give a homogenous mixture and divide themselves into 4 identical orbitals having identical energy. These new orbitals have been termed hybridized orbitals. Since there one s orbital and 3 p orbitals have combined to form the hybrid orbital, the hybridized orbitals are called sp³ orbitals. The energy of these hybrid orbitals lie in between the energy levels of the s and the p orbitals as shown here:

                                       Image: Formation of the hybridized orbital sp³

Each sp³ hybrid orbitals has one unpaired electron. Since these 4sp³ orbitals are identical in terms of energy, there is a tendency amongst these electrons to repel each other. To minimize the repulsion between electrons, the sp³ hybridized orbitals arrange themselves around the carbon nucleus in a tetrahedral arrangement. The resulting carbon atom is termed as sp³ hybridized carbon atom. 

                          Image: Tetrahedral arrangement of sp³ hybridized orbital

Overlap of each of the 4sp³ orbitals of the hybridized carbon atom with the s orbital of the hydrogen atoms leads to the formation of a methane molecule. The methane molecule can be shown as:

                              Image: sp³-s overlapping to form C-H bonding

It can be seen from the above that there are 4 identical sp³-s overlaps forming 4 identical C-H bonds which are consistent with the observations. Moreover, since these sp³ orbitals are oriented in the form of tetrahedrons, the geometry of the methane molecule is tetrahedral.

Thus, adding the concept of hybridization to the valence theory helps to understand the bonding in the methane molecule.

Example 2: The above example of methane had sp³ hybridization formed because of hybridization of 1 s and 3 p orbitals of the carbon atom. There are other types of hybridization when there are hybrid orbitals between 2 p orbitals and 1 s orbital called sp² hybridization. In case, there are hybrid orbitals between 1 s and 1 p orbitals, it is called sp hybridization. 

Let us consider the case of sp² hybridization. The structure of ethylene can be explained using the concept of sp² hybridization. The structure of the ethylene molecule observed is as:

                         Image: Electronic configuration of sp² orbital

Experimentally, the four carbon-hydrogen bonds in the ethylene molecule are identical and the geometry at each carbon atom in the ethylene molecule is planar trigonal.

Since the carbon atom has only 2 unpaired electrons, the valence bond theory cannot explain the formation of 4 bonds by each of the carbon atoms. Hence, we have to consider the excited state of both the carbon atoms in order that each carbon atom forms 4 bonds. 

We first consider the two carbon atoms and the double bond between them. 

For each of the excited carbon atoms, the one 2s orbital and two 2p orbitals (of the three 2p orbitals) form hybridization resulting in 3 hybrid orbitals called sp2 sp² orbitals. (1 s and 2 p orbitals). These 3 sp² orbitals try to be as distant from each other as possible and hence form a planar trigonal structure. The third 2p orbital in each of the carbon atoms does not participate in hybridization and remains as 2p orbital. 

                        Image: Structural Formula of C₂H₄

Each of the three sp² hybrid orbitals and the non-hybrid 2p orbital has 1 unpaired electron. To minimize repulsion of this non-hybrid 2p orbital with the 3 sp² orbitals, the 2p orbital stands perpendicular to each of the sp² hybrid orbitals. Hence, post-hybridization, the sp2 hybridized carbon atom looks as:

                                      Image:  sp2 hybridized carbon atom 

Each carbon atom in the ethylene molecule is bonded to two hydrogen atoms. Thus, overlap two sp²-hybridized orbitals with the 1s orbitals of two hydrogen atoms Also, the covalent C-C bond forms by overlapping of sp² orbitals of the two carbon atoms as:

                        Image: C-C bond forms by the overlapping of sp² orbitals

The two 2p orbitals of the carbon atoms overlap laterally to form a weak bond called a pi bond. 

Thus, the ethylene molecule is said to have sp2- sp² s bonds (4 C-H bonds), one sp²-sp² bond (C-C bond) and one p-p pi bond (C-C bond).   

Thus, the sp² hybridization theory explains the double bond, the trigonal planar structure in ethylene molecules. 

Example 3: Similarly, for a triple bond formation, like that of an acetylene molecule, there is sp hybridization between 1 s and 1 p orbital of the carbon atom. 

 Image: Structural Formula of C₂H₂

Here, there are 2 C-H bonds and a triple C-C bond. 

In each of the excited carbon atoms, one 2s and one 2p orbital form hybrid molecules called sp hybrid orbitals and the non-hybrid two 2p orbitals do not participate in hybridization. Because there are electron molecules in each of the orbitals, they tend to repel each other and the 2sp orbitals form a linear arrangement. The non-hybrid 2p orbital position themselves as far away as possible from each sp-hybridized orbital when perpendicular to each sp-hybridized orbital. So the resulting sp hybrid carbon atom looks like this:

                                                  Image:  sp hybridization 

The s orbitals of the hydrogen atom overlap with one sp hybrid orbital of each of the carbon atoms forming the 2 C-H bonds. The C-C covalent bond is formed by overlapping the sp-sp orbitals of the two-hybrid carbon atoms. In order to complete octet, the two non-hybrid 2p orbitals of each of the carbon atoms overlap laterally forming 2 pi bonds as shown:

Thus, sp hybridization explains the triple bond in acetylene molecules and the linear structure as well.

Nature of  the Types of Hybridization 

Hence, from the above text, we understand that hybridization is mathematically a concept of mixing atomic orbitals to form new hybrid orbitals suitable for the pairing of electrons to form chemical bonds in valence bond theory. Also, an entirely new orbital formed is different from its components and hence being called a hybrid orbital.

Hydrogen Peroxide is a highly unstable chemical compound. Two molecules of Hydrogen combine with two molecules of oxygen to form Hydrogen Peroxide. Hence, its chemical formula is H₂O₂. Hydrogen Peroxide is a pale blue, clear liquid, slightly more viscous than water in its pure form. It is the simplest Peroxide (since it is a compound with an oxygen-oxygen single bond). Hydrogen Peroxide has basic uses as an oxidizer, bleaching agent and antiseptic. Concentrated Hydrogen Peroxide, also known as “high-test Peroxide”, is a reactive oxygen species and has been used as a propellant in rocket propulsions. As this compound is unstable, it slowly decomposes in the presence of light. Hydrogen Peroxide is generally stored with a stabilizer in a weakly acidic solution since it is unstable. It can be found in biological systems including the human body. Peroxidases are the enzymes that use or decompose Hydrogen Peroxide.

 

History of Discovery

In 1799, Alexander von Humboldt synthesized barium Peroxide, one of the first synthetic Peroxides, as a by-product of his attempts to decompose air. After nineteen years, Louis Jacques Thénard stated that this compound could be used for the preparation of a previously unknown compound. He described it as eauoxygénée (French: oxygenated water) – which came to be known as Hydrogen Peroxide. An advanced version of Thénard’s method used hydrochloric acid, followed by addition of sulfuric acid to precipitate the barium sulfate by-product. This method was followed from the end of the 19th century until the middle of the 20th century. In 1811, Thénard and Joseph Louis Gay-Lussac synthesized sodium Peroxide. Pure Hydrogen Peroxide was initially believed to be unstable since early experiments to separate it from the water, which is present during synthesis,- all failed. This instability was present owing to traces of impurities (transition-metal salts). These impurities catalyze the decomposition of the Hydrogen Peroxide. Richard Wolfenstein first obtained pure Hydrogen Peroxide in 1894. He produced it by vacuum distillation.

Structure of Hydrogen Peroxide

The structure of Hydrogen Peroxide is non-planar which means that it has a three-dimensional quality. The structure for this compound is also popularly known as open book structure. The following diagram will explain the statement.

 

The diagram shows that there are two planes in the structure and each plane has one O-H pair, the angle between both the planes is 90.2°, The length of O-O bond is 145.8 pm and the O-H bond length is 98.8 pm(which equals to 9.88 × 10-13 m). Two pairs of unbound electrons will be present in both oxygen atoms. This proves the valence shell electron repulsion theory. The Hydrogen atoms will always repel the un-bonded electrons of oxygen. Hence, the bent molecular shape is formed.

 

Various methods are involved in the preparation of Hydrogen Peroxide:

• When acidification of barium Peroxide takes place and the excess water is removed by the process of evaporation under reduced pressure, we obtain Hydrogen Peroxide. The following reaction vouches for the fact:

• P/ρ+gz+v²/2=k

• P/ρ+gz+v²/2=k

• BaO₂.8H₂O(s)+H₂SO₄(aq)→BaSO₄(s)+H₂O₂(aq)+8H₂O(l)

• When acidified sulfate solution is electrolyzed at high current density, peroxodisulphate can be obtained. Peroxodisulphate then needs to be hydrolyzed to get Hydrogen Peroxide.

• 2HSO−₄(aq) Electrolysis−→−−−−−−− HO₃SOOSO₃H(aq) Hydrolysis−→−−−−−−− 2HSO−₄(aq)+2H+(aq)+H₂O₂(aq)

 

Properties of Hydrogen Peroxide

The properties of Hydrogen Peroxide are as follows:

• Hydrogen Peroxide is almost colorless (very pale blue) in a pure state.

• Its boiling point has been extrapolated at a temperature as high as 150.2 C which is almost 50 C higher than the boiling point of water.

• The melting point of Hydrogen Peroxide is -0.43 C.

• It forms a homogeneous mixture in water in all proportions and forms hydrates.

• 34.0147 g/mol is the molar mass of Hydrogen Peroxide.

• It has a slightly sharp odor.

• Its density is 1.11g/cc in aqueous solution and 1.450 g/cc in its pure form.

• Hydrogen Peroxide is soluble in ether, alcohol but insoluble in petroleum ether.

 

Chemical properties of Hydrogen Peroxide are as follows:

Hydrogen Peroxide acts both as an oxidizing as well as the reducing agent in acidic and also in basic medium. The following reactions depict the picture:

• Oxidizing nature in an acidic medium

• PbS(s)+4H₂O₂(aq)→PbSO₄(s)+4H₂O(l)

• The reducing nature in an acidic medium

• HOCl+H₂O2→H₃O++Cl−+O₂

• Oxidizing nature in a basic medium

• Mn₂++H₂O₂→Mn₄++2OH−

• Reducing nature in a basic medium

• I₂+H₂O₂+OH−→2I−+2H₂O+O₂

 

Storage of Hydrogen Peroxide

When exposed to sunlight, Hydrogen Peroxide decomposes. This decomposition process is catalyzed by traces of alkali metals. Therefore, Hydrogen Peroxide can be stored in wax-lined glass or plastic containers and are kept in the dark. It must also be kept away from dust particles as dust can induce explosive decomposition of this compound.

 

Uses of Hydrogen Peroxide

Hydrogen Peroxide has a number of uses. Some of them are listed below:

• Bleaching: Hydrogen Peroxide is used as a bleaching agent in the textile and paper industry. Data suggests that almost about 60% of the world’s production of Hydrogen Peroxide is used for pulp- and paper-bleaching.

• Hydrogen Peroxide can be used as hair bleach in our daily life and also as a mild disinfectant.

• Detergents: One of the major industrial applications of this compound is the manufacture of sodium percarbonate and sodium perborate, which are used as mild bleaches in laundry detergents. Sodium percarbonate is an active ingredient in products such as OxiClean and Tide laundry detergent. It releases Hydrogen Peroxide and sodium carbonate when dissolved in water.

• Production of organic compounds: Hydrogen Peroxide is used in the production of several organic Peroxides with dibenzoyl Peroxide being a high volume example. It is used in polymerizations, as a flour bleaching agent and also as a treatment for acne.

• Hydrogen Peroxide has also been used in certain waste-water treatment processes to remove organic impurities.

• Hydrogen Peroxide is used to sterilize various surfaces including surgical tools and it may be deployed as a vapor (VHP) for room sterilization.

• Hydrogen Peroxide is an environmentally secure substitute to chlorine-based bleaches, as it degrades to form water and oxygen and it is normally accepted as safe as an antimicrobial agent by the U.S. Food and Drug Administration (FDA).

• Since time immemorial, Hydrogen Peroxide had been used for disinfecting wounds, partly because of its low cost and prompt availability compared to other antiseptics. However, in recent times, it is thought to inhibit healing and to induce scarring because it destroys newly formed skin cells.

• Dermal exposure to dilute solutions of Hydrogen Peroxide can cause bleaching or whitening of the skin

• Hydrogen Peroxide has wide application in certain cosmetic production as well.

• It is present in most whitening toothpaste. One can mix this compound with baking soda and salt to make a home-made toothpaste.

• This compound is widely used in the production of alternative medicine. The use of Hydrogen Peroxide can cure various conditions, including emphysema, influenza, AIDS, and cancer although there is no proper evidence of effectiveness and in some cases, it may even be life-threatening.

• Hydrogen Peroxide can be used as a propellant in the rocket industry. Rocket-belt Hydrogen Peroxide propulsion systems are used in a jet back.

• Other uses of Hydrogen Peroxide include:

• Glow sticks

• Horticulture: Some horticulturalists and users of hydroponics suggest the use of a weak Hydrogen Peroxide solution in watering solutions.

• Fish aeration: Hence, Hydrogen Peroxide is one of the most important compounds.

An ideal solution is a composition where the molecules of separate species are identifiable, however, as opposed to the molecules in an ideal gas, the particles in an ideal solution apply force on each other. When the forces applied across all molecules are the exact same, irrespective of the species, a solution is said to be ideal.

For example, consider two liquids X and Y, and combine them. The solution formed will undergo various intermolecular forces of attraction in it, such as:

• X – X intermolecular forces of attraction

• Y – Y intermolecular forces of attraction

• X – Y intermolecular forces of attraction

For an ideal solution to be formed, the intermolecular forces of attraction between A – A, B – B and A – B must be just about equal.

If we go by the simplest definition of an ideal solution, it is described as a solution that is formed by mixing two elements that are of the same molecular size, structure and that have uniform intermolecular forces. An ideal solution abides by Raoult’s law at almost every range of concentration and temperatures. 

Characteristics of an Ideal Solution

Most of the time an ideal solution has physical properties closely related to those of the pure components, some of them are as follows:

The thermodynamics of a solution is zero. If the thermodynamics of the solution is closer to zero, then it is more probable to exhibit ideal behavior. ΔmixH = 0

The mixing volume is zero too. ΔmixV = 0

To get an ideal solution, it can be helpful to mix a solvent and a solute with identical molecular size and structure. The student can even take two substances X and Y and then mix them together. It is observed that there are several intermolecular forces existing between them.

In spite of the fact that getting a  balanced and stable ideal solution is an infrequent situation, certain solutions tend to show ideal behavior at times.

The ideal solution is a homogeneous combination of compounds that has physical properties that are linearly related to the quality of the elements. The classic example of this situation is the rule of Raoult, which refers to certain heavily diluted solutions and to a small class of condensed solutions, including those under which the interactions between the solute and solvent molecules are the same as those between the molecules themselves of each substance.

Benzene and toluene solutions, which have very close molecular structures, are ideal: each combination of the two has a volume equivalent to the sum of the concentrations of the respective elements, and the combining phase takes place without heat absorption or evolution. The vapor pressures of the solutions are defined mathematically by the linear structure of the molecular composition.

What is Raoult’s Law and Derive it?

Raoult’s law implies that the partial vapor pressure of a solvent in a solution (or mixture) is identical to or equal to the vapor pressure of a pure solvent multiplied by the mole fraction of the solution. 

Raoult’s law equation can be written mathematically as;

P_solution = Χ_solvent P^o_solvent

Where,

P_solution = vapour pressure of the solution

Χ_solvent  = mole fraction of the solvent

P⁰_solvent = vapour pressure of the pure solvent

As an example to define Raoult’s Law, consider a solution of volatile liquids in a container A and B as listed below. Since A and B both are volatile, in the vapor phase, there would be both particles of A and B.

Both A and B vapor particles, hence, exert partial pressure contributes to the total pressure above the solution.

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Further, Raoult’s law states that at equilibrium,

PA = P°A xA ,PB = P°B xB

Where PA is the partial pressure of A

P°A is the vapour pressure of pure A at the corresponding temperature.

xA is the mole fraction which is in the liquid phase

Similarly,

PB , P°B xB

Hence,

PT = PA + PB (Dalton’s Law) = P°A x_A + P°B xB = P°A + xB (P°B – P°A)

What is the Importance of Raoult’s Law?

Imagine we have a locked container loaded with volatile liquid A. For some time, due to evaporation, A vapour particles may begin to form. As time passes, the A vapor particles will be in dynamic equilibrium with the liquid particles (on the surface). The pressure exerted by the vapor particles of A at some given temperature is called the vapor pressure of A at that temperature.

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Vapor pressure is exhibited on both solids and liquids and relies solely on the temperature and the type of liquid. 

Now assume that we’re attaching another liquid B (solute) to this container. This will result in the B particles occupying the space between the A particles on the surface of the solution.

For any liquid that is given, there are a fraction of molecules on the surface that will have sufficient energy able to escape to the vapour phase.

Because we now have a lower number of A particles on the surface, the amount of A vapor particles in the vapor process would be greater. It would lead to the lower vapour pressure of A. 

However, if we believe that B is still volatile, we would have fewer B particles in the vapor phase compared to pure B’s liquid.

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Now, this new partial pressure of each (A and B) is given by Raoult’s law and depends completely on the concentration of each component in the liquid phase.

PA α XA, PB α PB = XB = XA P°A = XA P°B

It is evident from Raoult ‘s law that, as the mole fraction of the component lessens, its partial pressure also lessens during the vapour phase. 

The below graphs demonstrate the pressure for mole fractions A and B.

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Combining the two graphs, we get the resultant one as below.

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We have also added the graph for the total vapour pressure of the solution in the above diagram, i.e., PA + PB.

As far as Raoult’s law application goes, it is also useful to calculate the molecular mass of an unknown solute.

Properties of an Ideal Solution

Most of the time, Raoult’s ideal solution has physical properties that are closely related to the properties of pure components. 

Some of its properties are,

The solution enthalpy is zero. If the enthalpy of the solution reaches near to zero, it is more likely to exhibit ideal behavior. 

ΔmixH=0

The volume mixing is therefore zero.

ΔmixV=0

Characteristics of Ideal Solution

The ideal solution can be obtained by combining a solute with a solvent composed of a common molecular structure with size. If we take two substances as X and Y, and mixed together, we can see there are quite few intermolecular forces between them.

For instance,

X and X experience the intermolecular forces of attraction.

Y and Y experience the intermolecular forces of attraction.

X and Y experience the intermolecular forces of attraction.