Category: Chemistry Notes PPT

The study of chemical links between atoms or molecules is known as chemical bonding. This chapter discusses why only particular atoms can combine to form a new product and why they must be arranged in a specific shape. VSEPR, valence bond theory and molecular orbital theory are some of the theories that can explain all of the occurrences in detail. Bonding isn’t simply an example; it’s nature’s way of bringing every atom or molecule to the most stable state possible.

Every structure in the universe is the consequence of the development of specific types of bonds. In reality, bonding is nothing more than the joining of two atoms.

 

Kossel – Lewis Approach to Chemical Bonding 

Chemical bond is an attraction force between atoms of a molecule. 

In 1916 Kossel and Lewis succeeded in explaining the chemical bonding in terms of electrons. 

Octet Rule – Atoms of different elements try to attain electronic configuration like noble gas atoms or to complete their octet by chemical bonding. In other words, atoms of all main group elements tend to bond in such a way that each atom has 8 electrons in its valence shell so that the atoms will attain electronic configuration like noble gasses. Thus, by chemical bonding atoms get stability like noble gasses. 

Lewis postulated that atoms achieve the stable octet when they are linked by chemical bonds. 

Lewis Symbols – In the formation of a molecule, only the valence electrons (electrons of the outermost shell of an atom) take part in chemical bonding. An American chemist G. N. Lewis introduced simple notations to represent valence electrons in an atom. These notations are known as Lewis symbols. For example, carbon has 4 electrons in its outermost shell. So, its Lewis symbol will be as follows –

Electrovalent Bond – The bond formed, as a result of electrostatic attraction between the positive and negative ion is known as electrovalent bond. It is also called an ionic bond. 

The number of electrons lost or gained during the formation of an electrovalent linkage is termed as the electro valency of the element.

Factors Which Affect the Formation of Ionic Bond – 

• Ionization Enthalpy – Ionization enthalpy of any element is the amount of energy required to remove an electron from the outermost shell of an isolated atom in gaseous phase so as to convert it into a gaseous positive ion. It is also known as ionization energy. Lesser the ionization enthalpy, easier will be the removal of an electron, that is formation of a positive ion and hence greater the chances of formation of an ionic bond. 

• Electron Gain Enthalpy – Electron affinity or Electron gain enthalpy of an element is the enthalpy change that takes place when an extra electron is added to an isolated atom in the gaseous phase to form a gaseous negative ion. 

Higher is the electron affinity, more is the energy released and more stable will be the negative ion produced. Consequently, the probability of formation of ionic bonds will increase. 

Characteristics of Ionic Compounds –

• Ionic compounds usually exist in the solid state. 

• In ionic compounds ions get arranged in regular patterns. So, they have a definite crystal structure. For example, NaCl has octahedral crystal structure. 

• Ionic compounds usually possess high melting and boiling points. This is because ions are held together by strong electrostatic forces in ionic compounds. 

• Ionic compounds are usually soluble in polar solvents such as water.

• Solutions of ionic compounds are good conductors of electricity. They are good conductors of electricity in their molten state as well. 

• Ionic compounds take part in various chemical reactions. 

Covalent Bond – The bond formed by the sharing of electrons between atoms, is called covalent bond. This term and the idea was introduced by Langmuir in 1919.

The dots in the above structure represent the valence electrons (Lewis symbols). Such structures are called Lewis dot structures. Two atoms may have single, double or triple covalent bonds. 

Formal Charge – Formal charge of an atom in a polyatomic molecule can be defined as the difference between the number of valence electrons of that atom in an isolated or free state and the number of electrons assigned to that atom in the Lewis structure. It helps to select the most stable structure, that is, the one with least energy out of the different possible Lewis structures.

Formal charge on an atom in a Lewis structure = total number of valence electrons in free atom – Total number non bonding electrons – 1/2 Total number of shared electrons.

Example – Formal charge on atoms in carbonate ions. 

Lewis Structure of CO₃-2 Ion –

Formal charge on C atom = 4 – 0 – ½ (8) = 0

Formal charge on double bonded O atom = 6 – 4 – ½(4) = 0

Formal charge on single bonded O atom = 6 – 6 –  ½(2) = -1

Bond Parameters 

Bond Length – Bond length is defined as the equilibrium distance between the nuclei of two bonded atoms in a molecule. 

Factors Affecting Bond Length 

HI > HBr > HCl > HF 

C  ≡ C < C = C < C – C 

• Type of Hybridisation – As an s – orbital is smaller in size, greater the s-character, shorter is the hybrid orbital and hence shorter is the bond length. For example, 

Bond lengths – sp3 C–H > sp2 C–H > sp C–H 

s-character – (25%) (33%) (50%)

Bond Angle – Bond angle is defined as the angle between the orbitals containing bonding electron pairs around the central atom in a molecule or complex ion. It can be determined experimentally by spectroscopic methods. It is expressed in degrees. 

Bond Enthalpy – Bond enthalpy is defined as the amount of energy required to break one mole of bonds of a particular type between two atoms in a gaseous state. 

Factors Affecting Bond Energy –

• Size of the Atoms – Greater the size of the atoms, greater is the bond length and less is the bond dissociation enthalpy, i.e., less is the bond strength. 

• Multiplicity of Bonds – For the bond between the same two atoms, greater is the multiplicity of the bond, greater is the bond dissociation enthalpy. This is firstly because atoms come closer and secondly, the number of bonds to be broken is more. For example, bond dissociation enthalpies of H₂, O₂ and N₂ are in the order – H–H < O = O < N  N

• Number of Lone Pairs of Electrons Present – Greater the number of lone pairs of electrons present on the bonded atoms greater is the repulsion between the atoms and hence less is the bond dissociation enthalpy. 

Bond Order – In the Lewis description of covalent bond, the bond order is given by the number of bonds between the two atoms in a molecule. 

Isoelectronic molecules and ions have identical bond orders. Bond order of F₂ and O₂^2- is 1 and they are isoelectronic. 

With increase in bond order, bond enthalpy increases and bond length decreases. Thus, we can write –

Bond order ∝ bond enthalpy ∝ 1 Bond Length

Dipole Moment – Dipole moment can be defined as the product of the magnitude of the charge and the distance between the centers of the positive and negative charge.it is denoted by the Greek letter ‘μ’ . Mathematically, it can be expressed as follows –

Dipole moment = charge x distance of separation 

μ = Q × d

It is expressed in Debye units (D). 1D = 3.33564  10-30 C m (C = coulomb and m = meter)

It is a vector quantity. It is depicted by a small arrow with a tail on the positive center and head pointing towards the negative center. For example, dipole moment of HCl is represented as follows –

Significance of Dipole Moment – The molecules having zero moment are non-polar molecules and those having μnet ≠ 0  are polar in nature. 

The value of the dipole moment can be used for determining the amount of ionic character in a bond. The formula for determining percentage of ionic character is given below –

Percentage of ionic character = [frac{text{Experimental value of dipole moment}}{text{Theoretical value of dipole moment}} times 100 ]

The Valence Shell Electron Pair Repulsion (VSEPR) Theory

Valence shell electron pair repulsion theory explains the shapes of molecules. It is based on the repulsive interactions between electron pairs in the valence shell of the atoms. This theory was given by Sidgwick and Powell in 1940 and was further improved by Nyholm and Gillespie in 1957. Main postulates of VSEPR theory are as follows –

• The shape of the molecule depends upon the number of valence shell electron pairs either bonded or non-bonded around the central atom. 

• Pairs of electrons in the valence shell repel one another since their electron clouds are negatively charged. 

• These pairs of electrons tend to occupy such positions in space that minimize repulsion and thus maximize distance between them. 

• The valence shell is taken as a sphere with the electron pairs localizing on the spherical surface at maximum distance from one another. 

• A multiple bond is treated as if it is a single electron pair and the two or three electron pairs of a multiple bond are treated as a single super pair. 

• Where two or more resonance structures can represent a molecule, the VSEPR model is applicable to any such structure. 

Decreasing Order of the Repulsive Interaction of Electron Pairs –

Lone pair – lone pair > lone pair – bond pair > Bond pair – Bond pair

Or 

lp – lp > lp – bp > bp – bp 

Valence Bond Theory 

Valence bond theory was given by Heitler and London in 1927 and further developed by Pauling and others. It is based on electronic configuration of elements, the overlap criteria of atomic orbitals, the hybridization of atomic orbitals and the principle of variation and superposition. In class XI Chemistry, valence bond theory is discussed at a basic level in qualitative and non – mathematical terms only. 

1. nucleus of one atom and its own electron. 

2. nucleus of one atom and electron of another atom.

1. Electrons of the two atoms 

2. Nuclei of the two atoms

If the magnitude of attraction force is more than repulsive force, then the formation of bonds takes place and potential energy decreases. Ultimately, net force of attraction balances the force of repulsion and molecules acquire minimum energy and become stable.  

Sigma Bond – The strongest covalent bond which is formed by the head on overlapping atomic orbitals is called sigma bond. It is denoted by  . We find the sigma bond in alkanes, alkenes, alkynes. It is formed by s-s overlapping, s-p overlapping and p-p overlapping. Formation of sigma bond is given below between the orbitals- 

Pi Bond – The covalent bond which is formed by lateral overlapping of the half-filled atomic orbitals of atoms is called pi bond. It is denoted by . We find pi bonds in alkenes and alkynes. Formation of pi bond is given below between the two orbitals – 

Hybridization 

The concept of hybridization was introduced by Pauling. Atomic orbitals combine to form a new set of equivalent orbitals known as hybrid orbitals. The phenomenon is known as hybridization. These hybrid orbitals take part in bond formation. 

The Main Features of Hybridization Are As Follows –

• Number of hybrid orbitals = Number of atomic orbitals that take part in hybridization 

• The hybridized orbitals are always equivalent in energy and shape. 

• The hybrid orbitals are more effective in forming stable bonds than the pure atomic orbitals. 

• These hybrid orbitals are directed in space in some preferred direction to have minimum repulsion between electron pairs. Thus, it forms a stable arrangement. 

• The type of hybridization indicates the geometry of the molecules. 

Conditions for Hybridization –

• The orbitals present in the valence shell of the atom are hybridized. 

• The orbitals underg
  oing hybridization should have almost equal energy. 

• Promotion of electrons is not necessary before hybridization. 

• It is not necessary that only half – filled orbitals can take part in hybridization. Even filled orbitals of the valence shell take part in hybridization. 

Types of Hybridization – 

Hybridization takes place by involvement of various orbitals such as s, p, and d orbitals. It has following types –

• sp Hybridization – one s and one p orbitals take part in hybridization and each sp hybrid orbital has 50% s – character and 50% p – character. Two sp- hybrid orbitals are formed. 

• sp2 Hybridization – one s and two p orbitals participate in this hybridization and form three hybrid orbitals. Each of these hybrid orbitals show 33% s character and 67% p character. 

• sp3 Hybridization – one s and three p orbitals participate in this hybridization and form four hybrid orbitals. Each of these hybrid orbitals show 25% s character and 75% p character. 

• dsp2 Hybridization – one – s, two p and one d orbitals participate in this hybridization and form four hybrid orbitals. 

• sp3d Hybridization – one s, three p and one d orbitals participate in this hybridization and form five hybrid orbitals. 

• sp3d2 Hybridization – one – s, three p and two d orbitals participate in this hybridization and form three hybrid orbitals. 

• d2sp3 Hybridization – one – s, three p and two d orbitals participate in this hybridization and form three hybrid orbitals. 

Molecular Orbital Theory 

Molecular orbital theory is another approach to explain chemical bonding in molecules. It was given by Mulliken and Hund in 1932. The molecular orbital theory considers the entire molecule as a unit with all the electrons moving under the influence of all the nuclei present in the molecule. Salient features of molecular orbital theory are as follows:

• Like an Atomic orbital which is around the nucleus of an atom there are molecular orbital which are around the nuclei of a molecule. 

• The molecular orbitals are entirely different from the atomic orbitals from which they are formed. Atomic orbitals fuse together and form molecular orbitals. Conditions for atomic orbitals to form molecular orbitals –

1. The combining atomic orbitals should be of comparable energy.

2. The combining atomic orbitals must overlap to a large extent. Greater the overlap, stable the molecule form. 

• The valence electrons of the constituent atoms are considered to be moving under the influence of nuclei of participating atoms in the molecular orbital.

• The molecular orbitals possess different energy levels like atomic orbitals in an isolated atom. 

• The shape of molecular orbitals is dependent upon the shapes of atomic orbitals from which they are formed.

• Molecular orbitals are arranged in order of increasing energy just like atomic orbitals. 

• The number of molecular orbitals formed is equal to the number of atomic orbitals combining in bond formation. 

• Like atomic orbitals, the filling of electrons in molecular orbitals is governed by the three principles – Aufbau principle, Hund’s rule and Pauli’s exclusion principle.

Hydrogen Bonding 

Hydrogen bonding can be defined as the attraction force which binds the hydrogen atom of one molecule with the electronegative atom of another molecule. It is also called the hydrogen bridge. It is a very weak bond. 

Conditions For Hydrogen Bonding – The molecule must contain a highly electronegative atom such as F, Cl, Br etc. and the size of the electronegative atom should be small. 

Types of Hydrogen Bonding – 

It Is of Two Types –

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Intermolecular Hydrogen Bonding – When hydrogen bonding takes place between different molecules of the same or different compounds, it is called intermolecular hydrogen bonding. Example – water alcohol. 

1. Intramolecular Hydrogen Bonding – The hydrogen bonding which takes place within a molecule itself. The bond is formed between the Hatom of one group with the more electronegative atom of the other group. Example – Hydrogen bonding in an o-nitrophenol molecule. 

This ends our coverage on the topic “Chemical bonding and chemical structure”. We hope you enjoyed learning and were able to grasp the concepts. You can get separate articles as well on various subtopics of this article such as ‘Bond angle’, ‘Molecular orbital theory’ etc. on website. We hope after reading this article you will be able to solve problems based on the topic. If you are looking for solutions to NCERT Textbook problems based on this topic, then log on to website or download Learning App. By doing so, you will be able to access free PDFs of NCERT Solutions as well as Revision notes, Mock Tests and much more.

The following are some of the ways that distinct atoms or “bonds” come together:

• Ionic Bond: An ionic bond is formed between the various molecules when all of the electrons in a molecule have been exchanged.

• Covalent Bond: A covalent bond is formed when electrons are shared.

• Hydrogen Bond: A hydrogen atom forms a link with some of the most electronegative atoms, such as oxygen, nitrogen, and fluorine.

• Van der Waal Forces Of Attraction: Atoms are attracted to one another and form bonds.

The Chemical Bond’s General Structure 

All macromolecules in our bodies, such as DNA, RNA, proteins, and others, are bound together by chemical bonding. This chemical bond holds all of the structures together, whether the relationship is stronger or weaker. The structure’s stability is dictated by the strength of the bonds. The melting point, boiling point, and other key qualities are determined by the chemical bond strength. This chapter is popular among students since it is simple and provides a good distribution of marks in board exams and other competitive exams such as JEE and NEET.

How can You prepare for the Chapter of Chemical Bonding?

Even though this chapter is part of physical Chemistry, it is entirely theoretical. It is not necessary to practise formulas or numbers to prepare this chapter. To answer the problems in this chapter, you must first have a thorough understanding of the atomic structure, element classification, and periodicity from the previous chapters. To learn about Chemical Bonding and Molecular Structure, you need to first read its theory from the NCERT book and then work through the examples and questions. Apart from that, you can read O.P Tandon or P. Bahadur’s book if you want to prepare for advanced level competitive exams like JEE and NEET. Meanwhile, you can also give mock tests to know your preparation level. 

There are many different numbers of chemical reactions in our day-to-day life that we notice. We often tend to see a flaky brown coloured layer that appears on the surface of several iron items, for example, the bodies of vehicles, gates, etc. During the festivities and celebrations of Diwali, we burn crackers which when burned, gives a bright light and sound. The bright light is a result of the burning of the components like magnesium that is used amongst one of the ingredients of the crackers and sparklers that we burn. In this article, we will study these chemical processes in everyday life that we see around us.

Types of Chemical Reaction in Daily Life

Some common types of chemical reactions are given below.

1.  Photosynthesis:

The process of photosynthesis refers to the process by which autotrophs tend to manufacture their food. It is one of the natural chemical reactions in everyday life. In the presence of chlorophyll and sunlight, the plants produce glucose in the form of energy from water and carbon dioxide from the environment.

6CO₂ + 6H₂O → C₆H₁₂O₆(aq) + 6O₂ (g)

2.  Rusting: 

The process of oxidation, which refers to the reaction in the presence of oxygen, gives a brown flaky layer that we often notice over the metal surfaces like the one on the iron. This layer is formed because of the oxidation of the topmost layer for forming the metal oxide. We know this and call it rust. Similarly, several layers get formed on other metals as well, for example, in silver in which a green coloured layer gets formed when it undergoes an oxidation reaction.

Fe + 3O₂ + xH₂O → Fe₃O₄ . xH₂O

3.  Cellular Respiration:

The process of respiration in humans in their lungs also involves a chemical reaction. In this reaction, the glucose molecules tend to undergo an oxidation reaction and produces water and carbon dioxide along with energy. This process happens at a cellular level and hence, is called cellular respiration.

C₆H₁₂O₆ (aq) + 6O₂ (g) → 6CO₂ (g) + 6H₂O(l) + energy

4.  Anaerobic Respiration: 

When a living organism tends to undergo the process of fermentation, the components present in it which contain starch or sugar get converted to alcohol, acids, and gases.

C₆H₁₂O₆ → 2C₂H₅OH + 2CO₂

5. Combustion: 

The process of combustion is amongst the chemical changes in everyday life. It tends to involve the oxidation reaction of a material when it is exposed to the presence of both heat and oxygen from the surroundings to produce smoke, ash, and several other gases.

C + O₂ → CO₂

6.  Acid–-base Reactions: 

The acid–-base reactions are also a part of the chemical reactions in our daily life. In these reactions, the weaker acids that are produced inside our mouth overnight because of the sugary foods and bacteria get neutralized in the presence of base that is present in the toothpaste that we use in the morning every day.

Acid + base → salt + water

In a similar manner, the oxide layer that is found on the metallic artefacts gets removed with the help of weaker acids like vinegar. This process helps to give us the layer of the metal clean and a fresh new layer of the metal gets exposed.

CH₃COOH + FeO → CH₃COOFe + H₂O

Conclusion

In everyday life, we witness a number of chemical reactions. These reactions are very essential for the existence of human beings as well as for all living beings. In this article, we get to know about some common everyday life chemical reactions with their chemical equations.

What Is Chromium? 

The Chromium element has an atomic number 24 and is represented in the periodic table as the Cr element. It is the first element in group six. The chromium element is a very hard metal. It has a silver-grey appearance. The chromium electronic configuration is given as [Ar] 3d⁵ 4s¹. This logic behind the chromium electronic configuration is explained by the half-filled d orbital that is responsible for offering stability. The Cr electronic configuration is trivalent and is a very vital nutrient for the human body. This nutrient is found in different traces of human insulin, sugar, lipid metabolism, fatty acids and more. 

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A Few Common Chromium Properties

The atomic number of chromium (Z): 24

Group: group 6

Period: period 4

Block: d-block

Element category: Transition metal

Properties Of Chromium

The Chromium melting point is near about 2180 K ​(1907 °C)

Chromium boiling point: 2944 K ​(2671 °C, ​4840 °F)

Density: 7.19 g/cm3

The heat of fusion: 21.0 kJ/mol

The heat of vaporization: 347 kJ/mol

Uses Of Chromium

• Chromium is used in the stainless steel industry to manufacture and produce several alloys. A plating of chromium gives a mirror finish to stainless steel. Cars and trucks have number plates made up of stainless steel. These stainless steel parts contain chromium.

• Chromite is said to be the principal ore of Chromium. It is found in Africa, India, Turkey and Afghanistan. 

• Rubies get the bright red colour from chromium. Similarly, when treated with glass and emerald green colour is produced. 

• In the leather industry, the Cr element is used for tanning the leather products. But over the course of time, this process has proved to be highly toxic and as a result, an alternative process is being looked into. 

• Chromium when polished, shines like a mirror and reflects the toxic outline of the metal. 

• Various automobile decorations contain chromium or alloys made out of chromium. 

Symptoms Of Chromium Toxicity

Accuse exposure to chromium, in any form, leads to several diseases. These diseases are triggered along with symptoms such as nausea, vomiting, vertigo, fever and muscle cramps. The intensity of these symptoms can escalate and lead to more chronic symptoms such as renal failure, liver damage, respiratory tract discomfort and nephritis. Lung cancer is a very common disease linked with exposure to chromium. In addition to the lung, even the kidney and liver become a common target for chromium attacking particles. One should be extra careful while consuming chromium-rich food items. 

Chemical Properties Of Chromium 

1. Chromium Reacts with the Halogens: 

Chromium reacts with fluorine directly i.e F2 at 400°C and 200-300 atmospheric pressure to give chromium(VI) fluoride, CrF6.

Cr(s) + 3F2(g) → CrF6(s) [yellow]

Under milder conditions, chromium(V) fluoride, CrF5, is formed.

2Cr(s) + 5F2(g) → 2CrF5(s) [red]

In further milder conditions, Cr reacts with the halogens to give trihalide Chromium fluoride (III). 

2Cr(s) + 3F2(g) → 2CrF3(s) [green]

2Cr(s) + 3Cl2(g) → 2CrCl3(s) [reddish-violet]

2. Chromium Reacts with Acids: 

The Cr element dissolves in a dilute solution of hydrochloric acid to form solutions which contain the Cr(II) ion together along with H2 gas. Similarly while reacting with sulphuric acid the molecules become attack resilient. Chromium metal does not react with nitric acid or HNO3. 

Cr(s) + 2HCl(aq) → Cr2+(aq) + 2Cl–(aq) + H2(g)

3. Chromium Reacts with Sulphuric Acid: 

Chromium reacts with sulphuric acid to give chromium sulphate and hydrogen. The sulphuric acid must be a diluted solution. 

Cr + H2SO4 → CrSO4 + H2  

Fun Facts Related To The Chromium Element 

1. Chromium was first discovered in the year 1797 by a French Chemist named Louis Nicolas Vauquelin. 

2. The name means colour as the chromium element comes in different colours. Chromium, the word, has been derived from a Greek word “Chroma”. 

3. The Cr element is not found in free metal form. 

4. Each year about 20,000 tons of chromium is produced every year. 

5. The chromium electronic configuration gives the 4 natural isotopes. 

6. Stones of ruby and emerald contain traces of chromium. 

7. The most basic use of chromium is that it is used in the production of stainless steel. Chromium has anti-rusting properties which makes it suitable for being used in the production of stainless steel. Mostly all types of steel have a minimum chromium content of 10-11%. 

8. Considered to be the hardest metal present in nature, chromium is even stronger than iron, tungsten and titanium. In spite of this, the Cr element is very brittle. It can break into pieces and crumble easily if not handled properly. 

Meaning of Modern Periodic Table

The Modern Periodic Table, which is also known as the Periodic Table, is a tabular arrangement in which chemical elements, electron configurations, and chemical properties are arranged according to their atomic numbers. The Modern Periodic Table consists of four blocks, namely the blocks “S”, “P”, “D” and “F.” On the left-hand side of the Current Periodic Table, metallic elements are found, and on the right-hand side are non-metallic elements.

For the classification of elements, a periodic table is still used today. A simplified version of the modern periodic table can be seen in the figure below. By increasing the atomic number instead of the atomic mass, the modern table arranges the elements. The atomic number is the number of protons in an atom and for each particle, this number is unique. The modern periodic table has rows that are known as periods. Throughout a time, from left to right, each element has one more proton than the element before it.

Below is a figure of the modern periodic table, where you can see a simple version of the table. 

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How are Elements in the Modern Periodic Table Classified?

The method by which elements are classified based on their characteristics is a periodic classification of elements i.e. we hold the elements that are similar in one group and the rest of the elements in the other group. In the periodic table, several empty spaces have been left to position the elements that will be found in the future without disrupting the trending periodicity of the elements. 

7 periods of the modern periodic table, also known as horizontal rows, and 18 groups, also known as vertical columns have been grouped into the modern periodic table. According to the build-up theory, the elements are arranged ascendingly according to their atomic numbers and the way electrons fill their atomic energy sublevels, where each element has one electron more than the element that accompanies it.

In the modern periodic table, the classification of elements has been done in a way where most of the elements are metals. On the left-hand side of the periodic table, these metals occur. Non-metals are some of the elements in the table and they are less than 20 in number and on the right side of the table they exist. Metalloids are defined as some of the elements that occur on the boundary of metals and non-metals. Their properties are identical to those of both metals and non-metals. Most of these elements are solids, out of them just 11 elements i.e, noble gases, oxygen, nitrogen, fluorine, chlorine, hydrogen exist as gases and two of them are liquids i.e., mercury and bromine.

What is the Need for the Classification of Elements?

It is difficult to individually study every element and to know its properties and uses. Therefore, based on their similarities in properties, they were classified. In the Periodic Table, the structured classification of the elements has helped chemists to study and understand the properties of the elements and their compounds more systematically and orderly.

Advantages of Classification of Elements in the Modern Periodic Table

1. If the position of an element is known, It becomes easier to remember the properties of an element.

2. Unlike Mendeleev’s periodic table, positions of isotopes are taken care of within one element itself. 

3. The metals, non-metals, transition metals, gases are separately placed in a particular location with a specific identity in the modern periodic table.

4. The classification of elements is colour on the atomic number, which is a more basic property.

5. Many scientists have contributed their priceless efforts to systematically arrange the elements in a table. This lead to the development of the Periodic Table that we use today.

The Constituents of the Modern Periodic Table 

The periodic table classification has been done with mainly four types of elements. The elements present in the modern periodic table are :

1. The noble gases: a group of rare gases that consist of helium, radon, neon, argon, krypton, and xenon 

2. The main transition element: Electrons that can engage in chemical bond formation in two shells instead of just one shell.

3. The representative elements:  the group of elements from the group l and ll and the last six groups on the periodic table 

4. The inner transition element: the chemical elements, normally shown below all other elements

The colligative properties can be defined as the properties of solutions which is wholly determined by the ratio of the number of solute particles and the number of solvent molecules in a particular solution, and are completely independent of the nature of the chemical species present. The number ratio can be calculated by using various units that determine the concentration of solutions. The general assumption is that in the case of an ideal solution the properties are independent of the nature of solute particles present in it and are somewhat approximate for dilute real solutions. In other words, colligative properties can be considered as a set of properties of the solution that can be reasonably approached if we follow the assumption that the solution is ideal.

 

Here we have considered only those properties which are formed from the dissolution of a non-volatile solute in a volatile liquid solvent. Essentially the solvent properties are changed due to the presence of the solute. The solute particles actually displace some solvent molecules in the liquid phase and thus result in the reduction of the concentration of solvent. Thus we may conclude that the colligative properties are independent of the nature of the solute. The word colligative has been derived from the Latin colligates which means bound together.

 

Colligative Properties of a Solution Include:

• Relative lowering of vapor pressure

• Elevation of boiling point

• Depression of freezing point

• Osmotic pressure

 

For a particular mass ratio of solute and solvent, all colligative properties should be inversely proportional to solute molar mass.

 

Relative molar masses can be determined by measuring colligative properties for a dilute solution of a non-ionized solute. Such solutions include urea or glucose in water or another solvent. This is applicable for both small molecules and for polymers. In an alternate manner, the percentage of dissociation taking place can be estimated by measuring ionized solutes.

 

Colligative properties are mostly applicable for dilute solutions as their behavior may often be approximated. This is because they are ideal solutions.

 

This unit will focus mainly on the relative lowering of vapor pressure.

 

Relative Lowering of Vapour Pressure

On dissolving the non-volatile solute in a pure solvent, the vapour pressure of a pure solvent gets gradually decreased.

 

Let us assume that p is the vapour pressure of the solvent and ps is the vapour pressure of the solution, then the lowering of vapour pressure can be written as (p – ps). This lowering of vapour pressure relative to the vapour pressure of the pure solvent is called the Relative lowering of Vapour pressure. Thus,

 

Relative Lowering Vapour Pressure = [frac{p – p_s}{p}]

 

After extensive experimentation, Raoult (1886) gave an empirical relation to establishing the connection between the relative lowering of vapor pressure and the concentration of the solute in a solution. This is now referred to as Raoult’s Law. As per this law, the relative lowering of the vapor pressure of a dilute solution is equal to the mole fraction of the solute present in dilute solution.

 

Mathematically Raoult’s Law can be Expressed in the Form:

where n = number of moles or molecules of solute

 

[frac{p – p_s}{p} = frac{n}{n + N} ]

 

 

Derivation of Raoult’s Law

The vapor pressure of the pure solvent is the result caused by the number of molecules evaporating from its surface. When a non-volatile solute is dissolved in solution, due to the presence of solute molecules in the surface a fraction of the surface gets blocked and here no evaporation can take place.

 

By Lowering of vapor pressure by a non-volatile solute the particles of the solute prohibit the escape of solvent molecules from the surface of the solution. This finally results in the lowering of the vapor pressure. The vapor pressure of the solution is, therefore, dependent on the number of molecules of the solvent found at any time in the surface which is again proportional to the mole fraction. That is,

 

[p_s  alpha  frac{n}{n + N}]

 

where N = moles of solvent and n = moles of solute.

 

Or we can write it in the form

 

[p_s = K frac{N}{n + N}]

 

k being a proportionality factor.

 

In case of pure solvent n = 0 and hence Mole fraction of solvent

 

[p_s = frac{N}{n + N} = frac{N}{0 + N} = 1]

 

Now from equation (1), the vapor pressure p = k Therefore the equation (1) assumes the form:

 

[p_s = p frac{N}{n + N}]

 

[frac{p_s}{p} = frac{N}{n + N}]

 

[1 – frac{p_s}{p} = 1- frac{N}{n + N}]

 

[frac{p – p_s}{p} = frac{n}{n + N}]

 

This is Raoult’s law.

 

Ideal Solutions and Deviations From Raoult’s Law

A solution which strictly follows Raoult’s law strictly is called an ideal solution. A solution which shows even slight deviations from Raoult’s law is called a non-ideal or Real solution. Let us consider that the molecules of the solvent and solute are represented by A and B respectively. Now let γAB be the representation of the attractive force acting between A and B, and γAA between A and A. If

 

[gamma_{AB} = gamma_{AA}]

 

The solution will have the same vapor pressure as predicted by Raoult’s law and it is an ideal solution. However, if

 

[gamma_{AB} > gamma_{AA}]

 

molecule A will escape comparatively less readily and the vapor pressure will then be less than that of the predicted one, which is calculated by obeying Raoult’s law (Such deviation is called Negative deviation). On the other hand, if

 

[gamma_{AB} < gamma_{AA}]

 

A molecule will escape from the solution surface more rapidly and then the vapor pressure of the solution will become higher than predicted by Raoult’s law (Such deviation is then called the Positive deviation). If we consider very dilute solutions of nonelectrolytes, the solvent and solute molecules are very much similar be it in terms of molecular size or be its molecular attractions. Thus, under such situations, the given solutions have the tendency to approach the ideal behavior and obey Raoult’s law more or less accurately.

 

Determination of Molecular Mass From Vapour Pressure Lowering

The molecular mass of a non-volatile solute can be calculated from the measurement of the lowering of vapor pressure (p – ps ) produced by dissolving a known weight of it in a known weight
of the solvent. It is considered that w grams of solute is dissolved in W grams of the solvent, and let m and M are molecular masses of the solute and solvent respectively, we have :

 

No of Moles of solute (n) = [frac{w}{m}]

 

No of Moles of solvent (N) = [frac{W}{M}]

 

Substituting these values in Raoult’s law Equation

 

[frac{p – p_s}{p} = frac{n}{n + N}]

 

[frac{p – p_s}{p} = frac{frac{w}{m}}{frac{w}{m} + frac{W}{M} }]

 

Considering an extremely diluted solution, the number of moles (molecules) of solute (w/m), is very very small, it can be then neglected in the denominator. The equation (1) can then be written as

 

[frac{p – p_s}{p} = frac{wM}{mW} ]

 

With the known experimental value of [frac{p – p_s}{p} ], and the molecular mass of the solvent (M), the molecular weight of solute (m) can be calculated from the above equations.

 

Measurement of Lowering of Vapour Pressure

Barometric Method:

The individual vapor pressure of a liquid was calculated by Raoult and then the same process was followed to calculate the vapor pressure of the solution as well. He poured the liquid or the solution into the Torricellian vacuum of a barometer tube and calculated the depression of the mercury level. This method was later found to be neither practicable nor accurate as the lowering of vapor pressure is almost negligible.

 

Manometric Method:

The vapor pressure of a liquid or solution can be fairly measured with the help of a manometer. Let us assume a bulb is charged with the liquid or solution. The air in the connecting tube of the instrument is then removed with a vacuum pump. With the stopcock being closed, the pressure inside is only due to the vapor evaporating from the solution or liquid. This method can be applied to aqueous solutions. The manometric liquid used can be either mercury or n-butyl phthalate which has low density and low volatility.

 

Ostwald and Walker’s Dynamic Method (Gas Saturation Method):

In this method, the relative lowering of vapor pressure can be calculated in an easy simple Procedure. The apparatus used by Ostwald and Walker consist of two sets of bulbs : (a) Set A contains the solution (b) Set B contains the solvent. The weight of each set is calculated separately. A slow stream of dry air is then removed by a suction pump through the two sets of bulbs. At the end of the operation, the weight of these sets is again measured. From the weight loss in each of the two sets, the lowering of vapor pressure is measured. But here the temperature of the air, the solution, and the solvent must be kept constant all throughout. As the air bubbles through set A reaches saturation up to the vapor pressure ps of the solution and then up to vapor pressure p of solvent in set B, the amount of solvent taken up in set A becomes proportional to ps and the amount taken up in set B becomes proportional to (p – ps ).

 

Ostwald-Walker method of measuring the relative lowering of vapor pressure

 

If w1 and w2 ne the loss of weight in set A and B respectively,

 

[w_1] [alpha] [p_s]

 

[w_2] [alpha] [p – p_s]

 

Adding (1) and (2), we have

 

[w_1 + w_2]   [alpha]    [p_s  +  p – p_s]

 

[alpha]    [p_s  +  p – p_s]

 

Dividing (2) by (3), we can write

 

[frac{p – p_2}{p} = frac{w_2}{w_1 + w_2}]

 

Knowing the loss of mass in set B (w2) and the net loss of mass in the two sets (w1 + w2), we can find the relative lowering of vapor pressure. If we use water as the solvent, a    set of calcium chloride tubes (or a set of bulbs containing conc. H2SO4) is linked to the end of the apparatus to capture the escaping water vapor. Therefore, the gain in mass of the CaCl2-tubes will be equal to (w1 + w2), the total loss of mass in sets A and B.

A colligative property is a property of an answer that is reliant upon the proportion between the absolute number of solute particles (in the answer for) the complete number of dissolvable particles. Colligative properties are not reliant upon the compound idea of the arrangement’s parts. Subsequently, colligative properties can be connected to a few amounts that express the convergence of an answer, like molarity, ordinariness, and molality. The four colligative properties that can be displayed by an answer are:

“Colligative” has been adjusted or taken from the Latin word “colligatus” which means “bound together”. With regards to characterizing an answer, colligative properties assist us with seeing how the properties of the arrangement are connected to the grouping of solute in the arrangement.

What are Colligative Properties?

Weaken arrangements containing non-unpredictable solutes display a few properties which rely just upon the quantity of solute particles present and not on the kind of solute present. These properties are called colligative properties. These properties are for the most part seen in weakened arrangements.

We can additionally consider colligative properties as those properties that are acquired by the disintegration of a non-unpredictable solute in an unstable dissolvable. By and large, the dissolvable properties are changed by the solute where its particles eliminate a portion of the dissolvable atoms in the fluid stage. This likewise brings about the decrease of the convergence of the dissolvable.

In the meantime, when we talk about the given solute-dissolvable mass proportion, colligative properties are supposed to be conversely corresponding to the solute molar mass.

Colligative Properties Examples

We can notice the colligative properties of arrangements by going through the accompanying models. Assuming we add a touch of salt to a glass brimming with water its frigid temperature is brought down extensively than the typical temperature. On the other hand, its bubbling temperature is additionally expanded and the arrangement will have a lower fume pressure. There are changes in its osmotic tension too.

Likewise, assuming we add liquor to water, the arrangement’s edge of freezing over goes down underneath the ordinary temperature that is noticed for either unadulterated water or liquor.

Various Types of Colligative Properties of Solution

There are various kinds of colligative properties of an answer. These incorporate, fume pressure bringing down, edge of boiling over height, edge of freezing over gloom and osmotic strain.

In an unadulterated dissolvable, the whole surface is involved by the atoms of the dissolvable. Assuming a non-unstable solute is added to the dissolvable, the surface presently has both solute and dissolvable particles; along these lines part of the surface covered by dissolvable atoms
gets diminished. Since the fume strain of the arrangement is exclusively because of dissolvable alone, at a similar temperature the fume tension of the arrangement is viewed as lower than that of the unadulterated dissolvable.

1. Bringing Down of Vapor Pressure

On the off chance that P0 is the fume strain of unadulterated dissolvable and Ps is the fume tension of the arrangement. The distinction Po – Ps is named as bringing down in fume pressure. The proportion (Po – Ps)/Po is known as the overall bringing down of fume pressure.

Raoult, in 1886, set up a connection between relatives bringing down in fume tension and mole division. The relationship is known as Raoult’s law. It expresses that the overall bringing down in fume strain of a weaken arrangement is equivalent to the mole part of the solute present in the arrangement.

On the off chance that n moles of solute is broken down in N moles of the dissolvable, then, at that point, as per Raoult’s law

[ frac{P_o – P_s}{Po}  = frac{n}{n + N}]

2. Height in Boiling Point

The limit of a fluid is the temperature at which the fume pressure is equivalent to climatic tension. We realize that on the expansion of a non-unpredictable fluid to an unadulterated dissolvable, the fume tension of an answer declines. Consequently to make fume pressure equivalent to barometrical tension we need to expand the temperature of the arrangement. The distinction in the limit of the arrangement and the edge of boiling over of the unadulterated dissolvable is named as height in edge of boiling over.

In the event that T0b is the limit of the unadulterated dissolvable and Tb is the edge of boiling over of the arrangement then rise in edge of boiling over is given as

∆Tb = T0b – Tb

Trial results show that there is a connection between rise in edge of boiling over and molality ‘m’ of the solute present in arrangement

∆Tb ∝ m

∆Tb = kb m

Where,

kb = molal rise steady

Subbing the worth of ‘m’ in the above connection we get

∆Tb = [frac{1000 times kb times m_2}{M_2 times m_1}]

Where,

m2 = mass of dissolvable in g

M1 = mass of dissolvable in kg

M2 = molar mass of solute

3. Sadness in Freezing Point

The edge of freezing over a substance is characterized as the temperature at which the fume tension of its fluid is equivalent to the fume of the comparing strength. As indicated by Raoult’s law when a non-unpredictable strong is added to the dissolvable its fume pressure diminishes and presently it would become equivalent to that of strong dissolvable at a lower temperature. The contrast between the edge of freezing over of the unadulterated dissolvable and its answer is called wretchedness in edge of freezing over.

In the event that T0f is the limit of the unadulterated dissolvable and Tf is the edge of boiling over of the arrangement then wretchedness in edge of freezing over is given as

∆Tf = T0f  –  Tf

Very much like height at the edge of boiling over, sadness at the edge of freezing over is likewise straightforwardly connected with the morality ‘m’.

∆Tf = [frac{1000 times kf times m_2}{M_2 times m_1}]

Where,

k f = molal gloom consistent

m2  = mass of dissolvable in g

M1 = mass of dissolvable in kg

M2 = molar mass of solute

4. Osmotic Pressure

At the point when a semipermeable layer is put between an answer and dissolvable, it is seen that dissolvable particles enter the arrangement through the semipermeable film and the volume of the arrangement increases. The semi-porous film permits just dissolvable atoms to go through it however forestalls the section of greater particles like solute. This peculiarity of the unconstrained progression of dissolvable atoms through a semipermeable film from an unadulterated dissolvable to an answer or from a weakened to a concentrated arrangement is called assimilation.

The progression of dissolvable particles through the semipermeable layer can be halted on the off chance that some additional strain is applied from the arrangement side. This strain that simply stops the progression of dissolvable is called osmotic tension of the arrangement.

Osmotic strain is a colligative property as it relies upon the quantity of solute present and not on the idea of the solute. Tentatively it was demonstrated that osmotic tension (⫪) is straightforwardly relative to molarity(C) and temperature(T).

Numerically, ℼ = CRT where R is the gas consistent.

[Rightarrow pi = (frac{n_2 RT}{V})]

Here, V is the volume of arrangement in liters and n2 is moles of solute

Assuming m2 is the heaviness of solute and M2 molar mass of solute, then, at that point, [n_2 = frac{m_2}{M_2}]

[pi = frac{W_2 RT}{M_2V} ]

Accordingly by knowing the upsides of ℼ,w2, T and V we can work out the molar mass of the solute.

Various Solutions

Isotonic arrangement: Two arrangements having a similar osmotic tension at a given temperature are known as an isotonic arrangement. At the point when such arrangements are isolated by a semi-porous film then there is no assimilation.

Hypotonic arrangement: A hypotonic arrangement has a lower osmotic strain than that of the encompassing i.e, the convergence of solute particles is not exactly that of the encompassing. On the off chance that the hypotonic arrangement is isolated by a semipermeable layer then, at that point, water moves out of the hypotonic arrangement.

Hypertonic arrangement: A hypertonic arrangement has a higher osmotic strain than that of the encompassing i.e, the centralization of solute particles is more than that of the encompassing. Assuming the hypertonic arrangement is isolated by semipermeable film then, at that point, water moves inside the hypertonic arrangement.

The molar mass of a chemical compound in chemistry is defined as the mass of a sample divided by the amount of substance in the sample that is measured in moles. 

The molecular weight is usually used as a synonym of molar mass especially for molecular compounds, although it is very different as is known as molecular mass.

What is Molar Mass?

Molar mass is a smallest unit of a compound with one twelfth of the mass of one carbon – 12 atoms. When we know the number of moles needed, the concept of molar mass can be used to calculate how many grams of substance are required. The molar mass otherwise known as molecular weight is the sum of the total mass in grams of all the atoms that make up a mole of a molecule. The unit required to measure is grams per mole.

Formula of Mole

1. Mass of 1 mole of atom = Ar in grams

2. Number of moles of atom = mass of the element in grams/ relative atomic mass, Ar

3. Mass of 1 mole of molecules = Mr in grams

4. Number of moles of molecules = mass of the substance in grams / relative molecular mass, Mr

5. Mass of substance the contains 1 mole of particle = molar mass

6. Percentage yield = actual mass of product obtained / theoretical mass of product obtainable.

Chemical Computation with Mole and Avogadro’s Number

A mole is the SI measure of the quantity of a ‘chemical entity’ such as electrons, protons, or atoms. 1 mole contains 6.022 × 1023  elementary entities of the substance. Avogadro’s number is fundamental for understanding both the makeup of molecules and their combinations and interactions. For instance, since one atom of oxygen combines two atoms of hydrogen to create one molecule of water (H2O), one mole of oxygen (6.022 × 1023  of O atoms) combines with two moles of hydrogen (2 × 6.022 × 1023  of H atoms) to make one mole of H2O. 

Another Avogadro’s number property is that the mass of one mole of a substance is equal to the molecular weight of the substance. For instance, the mean molecular weight of water is 18.015 atomic mass units. Therefore, one mole of water weighs 18.015 grams. It simplifies many chemical computation.

 

Steps to Find Molar Mass for Compounds

Compounds are those substances that are made up of one or more elements. Examples of some common compounds include glucose, salt, acetic acid, and sodium bicarbonate. 

Sodium chloride compound is made up of two elements i.e., sodium and chloride. We will use sodium chloride as one of the examples to calculate the molar mass for the compounds 

Step 1: Find the atomic mass of individual element in the periodic table

We have to first find the atomic mass for each element. The element sodium has the atomic mass of 22.98976g/mol. Chlorine has an atomic mass of 35.453 g/mol.

Step 2: Count atoms for each element

As there are no subscripts for compound sodium chloride, it means it has only one sodium and one chlorine atom for the compound. 

Step 3: find the molar mass

Now, we are able to find the molar mass as we know the number of atoms for each element. Here we calculate first the mass of the sodium atoms that is 22.98976 g/mol. we will repeat the same for mass of chlorine atoms that is 35.453 g/mol, now we have to add these two masses together to find the total mass of molecules of sodium chloride. The total sodium chloride molecules are 58.44276 g/mol which we can round up to 58.44 g/mol.

Na = 1 × (22.98976 g/mol) = 22.98976 g/mol

Cl = 1 × (35.453 g/mol) = 35.453 g/mol

Molar mass = 22.98976 + 35.453 g/mol

Molar mass = 58.44276 or 58.44 g/mol

Solved Examples of Molar Mass

1. When you have 1.25 grams of a molecule with a molecular weight of 134.1 g/mol, find out how many moles of that molecule you have?

Solution: 1.25 g × 1 mole / 134.1 g = 0.0093 grams.

2. Calculate the mass of 6.022 × 1023 molecules of NH4Cl?

Solution: Molar mass in grams of NH4Cl = 14 + 4 +35.5 = 53.5 g

No. of moles of NH4Cl = 6.022 × 1023  / 6.022 × 1023  = 1 mole.

Now, mass of NH4Cl = number of moles × molar mass

                                         = 1 × 53.5 g

      = 53.5 g

3. Calculate the number of methane molecules and the number of carbon and hydrogen atoms in 25 g of methane?

Solution: Molar mass of methane = 16

Number of methane molecules = 25/16 × 6.022 × 1023  

    = 9.411 × 1023  

Number of carbon molecules = 1 × 9.411 × 1023  

  = 9.411 × 1023  

Number of hydrogen molecules = 4 × 9.411 × 1023  

      = 3.74 × 1023