Category: Chemistry Notes PPT

We all know that matter exists in three different physical states. All matter exists in solid, liquid or gaseous states. But as students go up to higher classes, they will learn that there are other states where matter exists in more sophisticated forms. Still, most of the substances in our daily life exist in the three basic physical states.

The Solid State

A solid-state the greater is characterized by strong bonds of attraction between its elementary particles. They have a fixed density and shape. They are the most stable state of any substance. They are rigid and relatively strong.

The liquid State

The liquid state is characterized by its appearance. It is not rigid and has relatively low intermolecular attraction between its particles. Liquids do not have a specific shape and take the shape of any container that they are placed in. They have a fixed density, unlike gasses.

The Gaseous State

The gaseous state is the most volatile physical form of an object. The particles are in constant motion and possess a lot of kinetic energy. They have large interatomic spaces and are less stable. They do not have any definite shape or density. They have low intermolecular and interatomic interaction due to their kinetic energy.

Even when this is the case, no substance permanently remains in its basic physical state. This is because of external forces acting on it. Every substance in the world is influenced by both internal and external forces. So they undergo a constant physical transformation. These transformations are facilitated by different processes that can be either spontaneous or forced. Some of them are boiling, melting, evaporation, sublimation, fusion, condensation etc. 

All these processes involve some amount of energy. Energy is either consumed or released during physical transformation. Energy consuming processes are called endothermic processes while energy releasing processes are called exothermic processes. Whether energy is used or released is determined by the nature of intermolecular or interatomic forces acting upon different substances. Solids have the greatest interatomic interaction, so any process that starts with the solid-state consumes energy since these strong bonds need to be broken. The reverse processes that result in the solid-state will be exothermic. Similar principles apply for liquids and gasses and their energy utilization is determined by the extent of atom-atom interaction and the energy these particles possess.

Have you tried making Maggie on the top of the mountain? not yet? Now, cooking Maggie at the top of the mountain will  test your knowledge of chemistry as well as your cooking skills. So before you reach the top of the mountain and test your cooking skills, uncover the chemistry behind it and improve your experience. 

Water boils at 100 ° C on the surface of the sea, but water in highlands such as mountain peaks begins to boil at low temperatures. When you try to cook Maggie at the top of the mountain,  the  water boils at a low temperature and evaporates quickly. Because of this, Maggie remains half cooked and needs more water to cook it well. Therefore, it takes  time to cook Maggie in the mountains. 

In this article, to get a clear understanding of the concept, we will first elaborate on vaporization and boiling, and then explain the differences between them. 

What is boiling? 

Boiling means that the liquid evaporates rapidly when  heated to the boiling point. For example, when  water is heated to 100 ° C on the surface of the sea, the water begins to boil and turns into steam. Boiling water is used to kill and soften the microorganism that are present there. It is also used in a variety of cooking methods.

In the boiling process, the vapor pressure of the liquid can overcome atmospheric pressure. That’s why bubbles can form and rise upwards in boiling. 

 

What is Boiling Point? 

• The boiling point of a substance is the temperature at which the vapor pressure of the liquid corresponds to the pressure surrounding the liquid and the liquid turns into vapor. The boiling point of a substance depends on the ambient pressure. High-pressure liquids have a high boiling point, and low-pressure liquids have a low boiling point. For example, at sea level, it boils at 100 ° C, but in the highlands where atmospheric pressure is low, water boils at temperatures below 100. ° C. 

•  factors affecting boiling point 

•  The following factors affect the boiling point- 

•  Atmospheric pressure-The boiling point of a substance changes depending on the ambient pressure. The higher the atmospheric pressure, the more energy is needed to break the bonds between the particles, and the boiling point rises. Therefore, high pressure raises the boiling point and low pressure lowers the boiling point. 

•  Impurities-The boiling point of a compound is used as a reference for its pure form. For example, pure water boils at 100 ° C, while water containing various other substances and dirty water boils at a higher temperature.

•  Thus, impurities in a pure substance elevate the boiling point. 

What is Evaporation?  

Take a beaker and put water in it. Then place this mug on the flame and continue heating. After a while, you will notice that the water begins to boil and turns into steam. This phenomenon is called evaporation. Did you notice now that when a glass of water falls to the floor and no one wipes it off, it dries after a while? Wet clothing will dry out after a while. Do you know how to do it? This means that they have different amounts of kinetic energy at different temperatures, as we know that particles of matter are always moving and  never resting. Even in the case of liquids, a small portion of the particles on the surface with high kinetic energy can be converted into vapors away from the gravitational pull of other particles. This phenomenon of converting a liquid into vapor at any temperature below its boiling point is known as evaporation.

Factors Affecting Evaporation

Following factors affect the rate of evaporation –

• Surface Area – Evaporation is a surface phenomenon. As the surface area increases, so does the rate of evaporation. For example, distribute clothing and let it dry faster. The relationship between evaporation rate  and surface area can be written as 

evaporation rate ∝ surface area. 

• Temperature-Evaporation rate  increases at higher temperatures. As with the rise in temperature, more particles receive enough kinetic energy to change to a vapor state. For example, wet clothing dries quickly in the sun. The relationship between evaporation rate  and temperature can be written as:

Rate of Evaporation ∝ Temperature 

• Humidity – Humidity is the amount of water vapor present in the air. Air cannot hold more than a definite amount of water vapor at a given temperature. If the amount of water in the air is already high or maximum, the rate of evaporation decreases. Relation between rate of evaporation and surface area can be written as –

Rate of Evaporation  ∝ 1∕ Humidity

• Wind Speed – With the increase in wind speed, the particles of water vapor move away from the wind, decreasing the amount of water vapor in the surrounding area. Thus, an increase in wind speed increases the rate of evaporation as well. For example, clothes dry faster on a windy day. The Relation between rate of evaporation and surface area can be written as –

Rate of Evaporation ∝ Wind Speed 

Differences Between Evaporation and Boiling 

This ends our coverage on the Difference between evaporation and boiling point. We hope you enjoyed learning and were absolutely able to grasp the concepts. We hope after reading this article you will be able to answer questions related to this topic. If you are someone who is  looking for solutions to NCERT Textbook problems based on this topic, then log on to the 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. 

Rutherford Atomic model is also known as the Rutherford model, nuclear atom, or planetary model of the atom was established in the year 1911 which explained the structure of atoms and was developed by the New Zealand-born physicist Ernest Rutherford. The model derived that the atom is nothing but a small tiny dense mass that has a positively charged body present in the core which is presently known as the nucleus where the entire mass of the atom is concentrated and around it revolves the negatively charged light electrons at a certain distance much like the planets revolving around the sun. 

In the gold foil experiment, the nucleus was postulated as a dense and small mass which was responsible for the scattering of the alpha particles. It was observed in a series of experiments that were carried out by the undergraduate Ernest Marsden under the guidance of Rutherford and German physicist Hans Geiger in 1909. The Rutherford model as a supplement for the  “plum-pudding” atomic model of English physicist Sir J.J. Thomson worked on the fact claimed by the plum-pudding atomic model that the electrons are embedded into the positively charged mass that was claimed as the atom-like plums in a pudding. 

Rutherford’s model was also obsoleted by Bhor’s atomic model purely based on classical physics. Bohr’s atomic model has also been seen to be incorporating some of the early concepts of quantum theory.     

Rutherford’s Alpha Scattering Experiment Explanation

Rutherford conducted a light scattering experiment where he placed a gold foil and bombarded the gold sheet with the alpha particles. The trajectory of the alpha particles was then studied after they interacted with the gold foil. There was a radioactive source that emitted Alpha particles which are positively charged particles that were enclosed within a lead shield in a protective manner.

The radiation then passed in a narrow beam after it passed through a slit which was made in the lead screen. A very thin section of a gold foil is placed before the lead screen and the LED screen was covered with zinc sulphide so as to give it a fluorescent nature that served as a counter detection to the Alpha particles. 

As soon as the Alpha particles right the fluorescent screen it’s shattered into a burst of light which is known as scintillation. It was visible from the viewing microscope that was attached to the back of the screen. As the screen was movable, it allowed Rutherford to study whether or not Alpha particles get deflected by the gold foil.

Observations of Rutherford’s Alpha Scattering Experiment

The observation that was made by Rutherford let him conclude that:-

1. Most of the Alpha particles that bombarded the gold fell passed through without any deflection that shows that the nucleus is made up of a large empty space.

2. Few of the Alpha particles that bombarded against the gold foil experienced a very minor deflection that shows that there is a presence of a counter positive charge.

3. Still, some of the Alpha particles that bombarded against the gold foil deflected to a larger angle and some of them even bounced back showing that the positive charge is concentrated in a very small volume and its distribution is non-uniform.

4. All the above points show that the volume occupied by positively charged particles in an atom is very small as compared to the total volume of the atom.

Result of Rutherford’s Alpha Scattering Experiment 

On the basis of his experiment, observation and result, Rutherford put forward Rutherford’s atomic model, which had the following features:

• The entire mass and positive charge of an atom are concentrated in a very small region at the centre known as the nucleus. 

• The positive charge on the nucleus is due to protons. Since the number of protons is different for atoms of different elements, therefore, the magnitude of positive charge on the nucleus is different for atoms of different elements. 

• The nucleus is surrounded by negatively charged electrons. The number of electrons in an atom is equal to the number of protons (positively charged) in the nucleus. Thus, the atom as a whole is neutral. 

• The electrons are revolving around the nucleus at a very high speed. 

• Most of the space in an atom is empty.

Drawbacks of Rutherford’s Atomic Model

Rutherford’s Atomic Model had the Following Limitations:

• This atomic model failed to explain the stability of atoms.

• According to the model, electrons revolve around the positively charged nucleus. It’s not possible for the long run as we know atoms are stable while any particle in a circular orbit would undergo acceleration. During acceleration charged particles would radiate energy. Revolving electrons will lose energy and finally fall into the nucleus. 

• This model of the atom also failed to explain the existence of definite lines in the hydrogen spectrum.

This was all about Rutherford’s atomic model. If you are looking for NCERT Solutions of Science Class IX, then download the learning app or register yourself on .

Introduction to Electrochemistry

Electrochemistry is the study of electricity production from energy released during spontaneous chemical reactions and how the electrical energy usage brings about non-spontaneous chemical transformations. A large number of metals such as sodium and magnesium, compounds like sodium oxide, and gases like chlorine and fluorine are produced by electrochemical processes. Batteries and fuel cells convert chemical energy into electrical energy and are used in large scale and devices.  Let’s get introduced to some important terms related to an electrochemical cell, chapter 20 electrochemistry for Science subject of class 12. 

Oxidation and Reduction

During oxidation, a loss of electrons takes place whereas during reduction, a gain of electrons is happening. Both oxidation and reduction reactions take place simultaneously in a redox reaction. Direct redox reactions involve a redox reaction in the same vessel and chemical energy gets converted to heat energy. During indirect redox reactions, oxidation and reduction takes place in different vessels and chemical energy is converted into electrical energy and the device where this takes place is called an electrochemical cell.

Electrochemical Cell

An electrochemical cell is a device which is capable of either producing an electric current due to chemical action or of producing chemical action due to the passage of electricity.

Types of Electrochemical Cell

There are two types of electrochemical cells, named as follows:

1. Voltaic Cells: In voltaic cells, the chemical energy of a spontaneous redox reaction is converted into electrical energy. These are also called galvanic cells. The electrical energy produced by such batteries can be used for powering cell phones, radios, and other devices. 

2. Electrolytic Cells: In electrolytic cells, electrical energy is needed to carry out a non-spontaneous chemical reaction. When we charge a cell phone battery, it’s like running an electrolytic cell.

Voltaic Cell: It consists of a half cell where oxidation takes place and it is known as oxidation half-cell; the other half-cell where reduction takes place is called reduction half-cell. Two half cells must be connected to build a voltaic cell. Oxidation takes place at negatively charged anode and reduction takes place at positively charged cathode. There will be a transfer of electrons taking place from anode to cathode when electric current flows in the opposite direction. An electrochemical cell also consists of an electrode which is prepared by dipping the metal plate into the electrolytic solution of its soluble salt. A salt bridge is a U shaped tube consisting of an inert electrolyte in agar-agar and gelatine. It maintains electrical neutrality and also allows the flow of electric current by completing the electric circuit. 

How is an Electrochemical Cell Represented?

1. Anode is on the left side while the cathode is written on the right sife.

2. Anode represents the oxidation half-cell and can be written as metal/metal ion (concentration).

3. Cathode represents the reduction half-cell and is written as metal ion (concentration)/metal

4. There’s a salt bridge which is indicated by placing double vertical lines between the anode and the cathode. 

5. Electrode potential is defined as the potential difference developed between the electrode and its electrolyte. It is represented as P.D. and P.D. between the metal and the solution of its ions is the result of the separation of charges at the equilibrium state and it is the measure of tendency of an electrode in the half cell in losing or gaining electrons. 

Standard Electric or Electrode Potential: When the concentration of all the species involved in a half cell is unity, it is called standard electrode potential. 

Anode is a negative electrode in the voltaic cell and cathode is a positive electrode in a voltaic cell. 

Functioning of Daniel Cell or Voltaic Cell

The half reaction for the Daniel cell are written as follows:

Left electrode : Zn(s) → Zn₂+ (aq, 1 M) + 2 e–

Right electrode: Cu₂ + (aq, 1 M) + 2 e– → Cu(s)

The net ionic equation for the daniel cell is as follows:

Zn(s) + Cu₂ + (aq) → Zn₂ + (aq) + Cu(s)

Zinc metal reacts with Copper ions to form Zinc ions and copper metal.

Electrochemical Cell and Gibbs Energy of the Reaction

Electrical potential multiplied by total charge passed is equal to the electrical work done in one second. If we need to obtain maximum work from a galvanic cell then charge has to be passed reversibly. The reversible work done by a galvanic cell is equal to the amount decreased in its  Gibbs energy and therefore, if the EMF of the cell is E and amount of charge passed is nF, delta G is the Gibbs energy of the reaction. Then,

ΔrG = – nFE(cell)

To understand the electronic configuration in Group 15 elements, we have to first understand the basic fundamental principles by which the electrons in an atom are arranged. The theoretical method by which electrons of an element are arranged in their subshells and orbital shells is referred to as electronic configuration. The structure of an atom is such that it consists of a nucleus which is surrounded by electrons that move in an orbital path around the particle. So, when there is an interaction of an atom with another particle, then the electrons on the outer side are the first ones that make contact.

It is important to note that if the outermost shell of an atom does not have the complete set of electrons, then that atom is considered to be reactive. As the reactivity is based on the number of electrons in the outer shell, the chemical properties of an element are also influenced by the electrons in its outer shell. These properties are found to be similar for the elements which have the same number of electrons in their outer shell.

We also use the electronic configurations of an element to tell whether it is found in nature in a stable form or not. The stability of the atoms depends upon the energy levels and number of atoms in their orbital path corresponding to the number of electrons required to complete its octet. So, those atoms which have a complete octet in their outermost shell are considered to be stable. The noble gases are a good example to show which elements are considered stable. We can also predict the reactivity of an element by using the octet rule. To study the electronic configuration of group 15 elements, such as the electronic configuration of nitrogen, the key is to remember the three basic laws, which are Paulie’s Exclusion Principle, Hund’s Law, and Aufbau’s Principle. 

Pauli’s Exclusion Principle

Wolfgang Pauli came up with Pauli’s exclusion principle for electrons in 1925. Now to understand the concept clearly, you must know these terms:

Quantum number = n

Azimuthal quantum number = m

Principle quantum number= l

The methodology to fill electrons in an atom is from lower energy levels to higher energy levels. According to Pauli’s exclusion principle, the electrons of an atom should not possess the same n, m, and l simultaneously.  For instance, the n,m, and l for an electron in the same orbital path are the same. Then their magnetic quantum number would be the same as well, and hence they would possess opposite integer spins, i.e. ½ and -½. Pauli’s exclusion principle does not hold for bosons (particles with integer spin). Since several bosons are capable of holding the same quantum state. Also, this principle helps get a clear picture of orbital shells of an atom. 

Hund’s Rule

According to Hund’s rule, when filling orbitals with energy levels, an electron looks to fill subshells with the same energy levels before pairing them with other electrons. In other words, all the orbitals must be filled with single electrons first. By filling all the orbitals with individual particles having the same energy levels allows maximizing the total spin. This process is due to all single filled electrons having the same integer spin. 

 

Aufbau’s Principle

Aufbau’s principle states that while filling the orbitals with electrons in an atom, it should always be in a manner of increasing energy level. In other words, the orbitals with low energy levels are to be filled with electrons first and then the orbitals with higher energy levels. You can use this principle correctly in the first eighteen elements of the periodic table. And after that, the efficiency will start to decrease.

The Different Rules on the Electronic Configuration of Group 15 Elements

Group 15 in the periodic table consists of five elements. They are also known as nitrogen group elements. The total number of valence electrons in group 15 is five. Expanding on the electronic configuration rules, we can write the electronic configuration of group 15 elements. Let us write the electronic configuration of nitrogen, phosphorus, arsenic, antimony, and bismuth.

Nitrogen is one of the two non-metallic gases in the group 15 elements. Its atomic number is 7, and its symbol is N. The nitrogen atom has an s-orbital with two electrons and p-subshell with three electrons. This configuration is because to pair with other electrons, the p subshell needs to be half-filled. Its electronic configuration is as follows-

[left [ He right ]]2s²2p³

Phosphorus is another metallic gas in group 15 elements. It has an atomic number of 15, and its symbol is P. The electronic configuration of phosphorus is as follows-

[left [ Ne right ]]3s²3p³

Arsenic is the third element in the group 15 elements. Its atomic number is 33, and its symbol is As. Its electronic configuration is as follows-

[left [ Ar right ]]3d¹⁰4s²4p³

Antimony is the fourth element in the group 15 elements. Its atomic number is 51, and its symbol is Sb. Its electronic configuration is as follows-

[left [ Kr right ]]4d¹⁰5s²5p³

Bismuth is the last element in the group 15 elements. Its atomic number is 83, and its symbol is Bi. Its electronic configuration is as follows-

[left [ Xe right ]]4f¹⁴5d¹⁰6s²6p³

The measure of the ability of the elements, mainly metals, to donate electrons for the formation of the positive ions is called electropositivity. On the other hand, the elements which can easily accept the electrons for the formation of negative ions are known as electronegative elements. Non-metals are examples of electronegative elements. 

 

It should be noted that electropositivity is opposite to electronegativity, which is a measure of atomic metals having the propensity to receive electrons and form poorly charged anions. And hence, light-emitting substances have very low electronegativities and very high-energy electrons have very low electropositivity. Electronegative elements usually have no metals and have the propensity to lose electrons to form cations, and electropositive metals substances usually do not receive electrons to form anions. The highest electropositive elements normally form ionic salts with electronegative elements. For instance, sodium is a highly electropositive metal that easily provides an electron to obtain a stable electron suspension. And Chlorine is a highly potent element that absorbs electrons to achieve a stable octet.

 

Define Electropositivity

Electropositivity can be defined as the tendency of an atom to donate electrons to form positively charged cations. The property to form positively charged cations is most probably exhibited by the metallic elements in the periodic table, especially the alkali metals and the alkaline earth metals. 

 

Electropositivity is just the opposite of electronegativity. The highly electronegative elements have very low electropositivity whereas the highly electropositive elements have very low electronegativity.  The electronegative elements usually form ionic salts with the electronegative elements. Sodium, which is a highly electropositive element, gives up an electron to obtain a stable electronic configuration. On the other hand, chlorine is a highly electronegative element that accepts an electron to achieve a stable octet. 

 

Therefore, the electropositive element sodium and the electronegative element chlorine can form an ionic bond with each other to give sodium chloride. Sodium chloride is also known as common salt which is consumed every day. 

 

Periodic Trends in The Electropositivity of The Elements

The electropositivity of an element depends on various factors like the metallic character of an element, the ionization energy of an element, the distance between the nucleus and the valence shell, and also the effective nuclear charge acting on the valence shell. The periodic trends in electropositivity exhibited by the elements are always opposite to the periodic trends in the electronegativities of the elements. 

 

Electronegativity of the elements increases across a period whereas the electropositivity of the elements decreases across the periods, the electronegativity of the elements decreases down the group, and the electropositivity of the elements increases while traversing down a group. This is the reason why the elements at the top right of the periodic table are the least electropositive and the elements at the bottom left of the periodic table are always electropositive. 

 

Electropositivity is a metallic attribute; it is dependent on the metallic character of an element.  This is the only reason why all the alkali metals are regarded as the most electropositive elements in the periodic table. Caesium and francium are the highest electropositive elements in the entire periodic table. Whereas, fluorine, chlorine, and oxygen are the most electronegative elements in the periodic table which also means that they are the least electropositive elements in the periodic table. 

 

Electropositivity is primarily a metallic property; hence it is influenced by the element’s metallic character. This is why alkali metals are considered to be the most electropositive elements e.g. with caesium and francium being the most electropositive elements in the entire periodic table.

Fluorine, oxygen, and chlorine are the least electropositive elements in the periodic table since they are the most electronegative.

Top 5 electropositive elements

The most selective feature in the timeline is Cesium (Cs). A list of the top five energy selection items is given below:

1. Caesium

2. Rubidium

3. Potassium

4. Sodium

5. Lithium

What is Electronegativity

The metals of an atom to a molecule to attract electrons distributed by itself is known as electronegativity. It is flawless material because it is just a habit. It shows the complete effect of the atomic inclination on various objects to attract pairs that form electrons. Measure electronegativity on several scales. The most widely used scale is designed by Linus Pauling. According to this scale, fluorine is a highly electronegative substance with a value of 4.0 and caesium is a small non-electrical substance with a value of 0.7.

Electropositive Character

The tendency of an element to lose the electrons to form the positive ions is called the electropositive character. It is also called the metal character. The elements which have very low ionization energies have a higher tendency to lose electrons and therefore they are electropositive or metallic in their behavior. The alkali metals are always the most highly electropositive elements. 

 

Periodicity

1. The reactivity of the metal decreases from the left to the right in a period as the tendency of an element to lose electrons decreases. 

2.  The reactivity of the metals increases from the top to the bottom in a group because the tendency of an element to lose electrons increases in a period. 

3. In a periodic table, the electropositivity or the metallic characters increases from the top to the bottom of the group. 

 

Electropositive Elements

In a periodic table, the elements are usually divided into two groups. The first group of elements is called metals and the second group of the elements is known as the non-metals. The metals and the non-metals are also divided into two categories which include electropositive and electronegative. 

 

The electropositive is those elements or groups that give up electrons i.e. metals and acidic hydrogen. Electropositive elements are those elements whose electrode potential is more positive than that of a standard hydrogen electrode which is assigned an arbitrary value of zero. Examples of univalent alkali metals are Li+, Na+, K+, etc. An example of divalent alkaline earth metals is Be2+
, Mg2+, Ca2+. 

 

What are Electropositive Radicals?

Electropositive radicals are atoms, ions, or molecules that can lose an electron and carry a positive electrical charge. An electropositive radical is formed due to the electropositive nature of a chemical species, which means a particular chemical species has the tendency to lose electrons in order to form positive radicals. Moreover, some examples of electropositive radicals include calcium cation (Ca+2), sodium cation (Na+), etc.

What are Electronegative Radicals?

Electronegative radicals are atoms, ions, or molecules that can gain an electron and carry a negative electrical charge. An electronegative radical is formed due to the high electronegativity of a chemical species, meaning, a particular chemical species have the tendency to gain electrons and form negatively charged radicals.

A radical, in chemistry, is an atom, molecule, or ion that contains an unpaired valence electron. The key difference between electropositive and electronegative radicals is that electropositive radicals are radical compounds having the capacity to lose electrons and carry a positive charge whereas electronegative radicals are radical compounds having the capacity to gain electrons and carry a negative charge. Most of the time, radicals are highly reactive chemical species which makes them undergo dimerization and polymerization reactions.

As the study in the field of chemistry furthered over the centuries, scientists preferred rationalizing the otherwise anomalous behavioral patterns of atoms and molecules, thus introducing the Periodic Table put forward by Mendeleev and the Atomic Structure by Ernest Rutherford, later corrected by Bohr. The idea of the atomic structure much like a solar model is flawed in many ways. Still, the perk here is it can give the fundamental understanding of the Nucleus posing as the Sun at the center and the Electrons orbiting the Nucleus matching planets with the disparity that the Electrons populate the regions of space, which are called the Orbitals and the orbitals have varied orbital energy levels.

 

As the magnitude of the changes increase, the magnitude of force also increases, and the forces decrease when the separation of charges is more. Therefore, the force of attraction between an electron and its nucleus is directly proportional to the distance between them. The electron is bound more tightly to the nucleus if the electron is closer to the nucleus. Hence, the electrons in the different shells which are at different distances have different energies.

 

What is Principal Quantum Number (N)?

The foundation of orbitals chemistry starts with Bohr who established that electron orbitals represent an energy level in terms of their distance from the Nucleus. The Orbitals are named K, L, M, N… or 1, 2, 3, 4… in ascending order. These numbers are the Principal Quantum Numbers. A Principal Quantum number is denoted as ‘n’. For example, for the K-orbital n = 1, for L-orbital n = 2, for M-orbital n = 3.

 

What is the Azimuthal Quantum Number (L)? 

Arnold Sommerfeld delved deeper into the orbitals chemistry, he viewed every orbital energy level or shell is made up of many subshells. He imagined that other than the circular orbits that Bohr established, there are elliptical orbits as well. The Azimuthal or Subsidiary quantum number helps to determine the ellipticity of the subshells. It is generally denoted as ‘l’.

 

To denote the value of ‘l’ instead of 1, 2, 3… some spectroscopic symbols are used –

 

It should be noted that the subshells are energy levels as well, called Subsidiary orbital energy levels, so if we sort the subshells in ascending order in terms of their energy levels, it would be s < p < d < f.

 

Orbital Energies

The energy of the electron in a hydrogen atom depends only on the principal quantum number, n. 

 

The nucleus of a hydrogen atom has a charge of +1. If the electron is bound to a nucleus of arbitrary charge +Z, then the energy of the electron is

 

E=−RyZ2n2E=−RyZ2n2

 

Where, Ry is the Rydberg unit of energy where 

1 Ry = 2.179877125595425×10−18×10−18J 

         = 13.60572374378387 eV

 

This equation is used for a single electron orbiting a single nucleus of charge +Z. 

 

With the increase in quantum number, n increases (holding Z constant), and the energy increases i.e. it becomes less negative. In the limit that n goes to infinity then the energy goes to zero. With the increase in artitary charge, z, the energy decreases i.e. it becomes more negative.

 

A higher Z means a more positively charged nucleus, therefore it holds the electron tighter.

 

Calculating the Energy Level of an Orbital

In a single electron, Hydrogen-like atom, the orbital energy i.e. the energy of that one electron depends just on the principal quantum number (n). In orbitals chemistry when it comes to filling up the atom with electrons, the Aufbau principle tells the lower energy level orbitals always come first. The order followed here is 1s <2s < 2p < 3s <3p < 4s… 

 

To easily memorize this anomalous behavior I strongly suggest following this diagram-

• Since the electrons are negatively charged particles, they repel each other. The stability of an atom depends on the attraction between the electrons, positively charged Nucleus and repulsive force within the electrons. The particle can only be stable if the total attractive interaction is more than the whole repulsive interaction. 

• As we go down the periodic table, the atomic number increases and another factor comes into play here, i.e., shielding. Due to the presence of electrons in the inner shells, the total positive charge exerted by the Nucleus (Ze) is slightly hindered for the electrons in outer shells. The net positive charge felt by the electrons in the outer shells are termed as an effective nuclear charge (Zeffe ). 

• The closer the orbital is to the Nucleus more tightly bound it would be. So an s-orbital electron will be more tightly bound to the Nucleus of that atom than a p-orbital.

• The s-orbital particles will be of a lesser charge as it has a lower orbital energy, which means it would be a more negative charge than the electrons in the p-orbital, which will have smaller energy for its higher orbital energy compared to the d-orbital electrons.

• In some cases, two orbitals may have the same n+l value; in those instances, the orbital with a lower n (principal quantum number) count will have low energy.

• In some cases two orbitals may have the same n+l value; in those instances the orbital with a lower n (principal quantum number) will have a low energy. 

 

Which Factors Affect Orbital Energy?

The s orbital electron is more tightly bound to the nucleus in comparison to the p orbital electron, which is more tightly bound with respect to a d orbital electron for a given value of the principal quantum number.

 

The orbital energy decreases in the same subshell with the increase in the atomic number (Zeff).

 

As compared to p orbital electrons, s orbital electrons have lesser amounts of energy and are more negative. In this case, the p orbital electrons will have lesser energy than that of d orbital electrons.

 

The extent of shielding from the nucleus is different for the electrons in different orbitals. Hence, it leads to the splitting of energy levels that have the same principal quantum number. Hence, the orbital energy depends on the values of both the principal quantum number and azimuthal quantum number, which are symbolized as n and l respectively. Therefore, the lower the value of (n + 1) for an orbital, the lower is its energy.

 

The Energy of Orbital in Hydrogen Atom

The energy of an electron in a hydrogen atom is calculated solely by the principal quantum, m (n). Therefore, the energy of the orbitals in hydrogen atom increases as follows :

1s < 2s = 2p < 3s = 3p = 3d <4s = 4p = 4d = 4f <..

 

The shapes of 2s and 2p orbitals are different, an electron has the same energy when it is in the 2s orbital as when it is present in a 2p orbital. The orbitals which have the same energy are called degenerate orbitals, whereas the orbitals which have the same energy are called degenerate orbitals.

 

The 1s orbital in a hydrogen atom is the most stable condition and is called the ground state and an electron living in this orbital is most strongly sustained by the nucleus. An electron in the 2s, 2p or higher orbitals in a hydrogen atom is in an excited state.

 

Solved Examples 

1. Which of these orbitals has a lower orbital energy level 3d or 4s?

The n + l value of 3d orbital is (3 + 2) = 5, Similarly the (n + l) value of 4s is (4 + 0) = 5.

So the 4s orbital has a higher (n+l) value, thus has a higher orbital energy level. 

Ans: 3d

 

2. Which of these orbitals has a higher orbital energy level 3d or 4p?

The (n+l) value of 3d orbital is (3+2) = 5, and 4p orbital is (4+1)=5. Both have the same (n+l) value with 3d having a lower n-count; thus, it is weaker and has a lower orbital energy level. 

Ans: 3d 

 

Fun Facts

The letters s, p, d ,f represent the shape of the orbitals. The s-orbital is spherical, and the Nucleus is in its center. The p-orbital has a form of a pair of lobes on each side of the Nucleus, somewhat has a dumbbell kind of structure.