essay about atomic structure

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• Atomic Structure

Atomic Structure - Discovery of Subatomic Particles

The atomic structure refers to the structure of an atom comprising a nucleus (centre) in which the protons (positively charged) and neutrons (neutral) are present. The negatively charged particles called electrons revolve around the centre of the nucleus .

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The history of atomic structure and quantum mechanics dates back to the times of Democritus, the person who first proposed that matter is composed of atoms. The study of the structure of an atom gives a great insight into the entire class of chemical reactions, bonds and their physical properties. The first scientific theory of atomic structure was proposed by John Dalton in the 1800s.

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Structure of Atom – Important Topics

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What is atomic structure, atomic models, dalton’s atomic theory, thomson atomic model, rutherford atomic theory, subatomic particles, atomic structure of isotopes, bohr’s atomic theory, dual nature of matter.

The advances in atomic structure and quantum mechanics have led to the discovery of other fundamental particles. The discovery of subatomic particles has been the base for many other discoveries and inventions.

The atomic structure of an element refers to the constitution of its nucleus and the arrangement of the electrons around it. Primarily, the atomic structure of matter is made up of protons , electrons and neutrons.

The protons and neutrons make up the nucleus of the atom, which is surrounded by the electrons belonging to the atom. The atomic number of an element describes the total number of protons in its nucleus.

Neutral atoms have equal numbers of protons and electrons. However, atoms may gain or lose electrons in order to increase their stability, and the resulting charged entity is called an ion.

Atoms of different elements have different atomic structures because they contain different numbers of protons and electrons . This is the reason for the unique characteristics of different elements.

In the 18th and 19th centuries, many scientists attempted to explain the structure of the atom with the help of atomic models. Each of these models had its own merits and demerits and was pivotal to the development of the modern atomic model . The most notable contributions to the field were by the scientists such as John Dalton, J.J. Thomson, Ernest Rutherford and Niels Bohr. Their ideas on the structure of the atom are discussed in this subsection.

The English chemist John Dalton suggested that all matter is made up of atoms, which were indivisible and indestructible. He also stated that all the atoms of an element were exactly the same, but the atoms of different elements differ in size and mass.

Chemical reactions, according to Dalton’s atomic theory, involve a rearrangement of atoms to form products. According to the postulates proposed by Dalton, the atomic structure comprises atoms, the smallest particle responsible for the chemical reactions to occur.

The following are the postulates of his theory:

• Every matter is made up of atoms.
• Atoms are indivisible.
• Specific elements have only one type of atom in them.
• Each atom has its own constant mass that varies from element to element.
• Atoms undergo rearrangement during a chemical reaction.
• Atoms can neither be created nor destroyed but can be transformed from one form to another.

Dalton’s atomic theory successfully explained the Laws of chemical reactions , namely, the Law of conservation of mass, the Law of constant properties, the Law of multiple proportions and the Law of reciprocal proportions.

Demerits of Dalton’s Atomic Theory

• The theory was unable to explain the existence of isotopes.
• Nothing about the structure of the atom was appropriately explained.
• Later, scientists discovered particles inside the atom that proved the atoms are divisible.

The discovery of particles inside atoms led to a better understanding of chemical species; these particles inside the atoms are called subatomic particles. The discovery of various subatomic particles is as follows:

The English chemist Sir Joseph John Thomson put forth his model describing the atomic structure in the early 1900s.

He was later awarded the Nobel Prize for the discovery of “electrons” . His work is based on an experiment called the cathode ray experiment . The construction of working of the experiment is as follows:

Cathode Ray Experiment

It has a tube made of glass which has two openings, one for the vacuum pump and the other for the inlet through which a gas is pumped in.

The role of the vacuum pump is to maintain a “partial vacuum” inside the glass chamber. A high-voltage power supply is connected using electrodes, i.e., cathode and anode , which are fitted inside the glass tube.

Observations:

• When a high voltage power supply is switched on, there are rays emerging from the cathode towards the anode. This was confirmed by the ‘Fluorescent spots’ on the ZnS screen used. These rays were called “Cathode Rays”.
• When an external electric field is applied, the cathode rays get deflected towards the positive electrode, but in the absence of an electric field, they travel in a straight line.

• With all this evidence, Thompson concluded that cathode rays are made of negatively charged particles called “electrons”.
• On applying the electric and magnetic field upon the cathode rays (electrons), Thomson found the charge-to-mass ratio (e/m) of electrons. (e/m) for electron: 17588 × 10 11 e/bg.

From this ratio, the charge of the electron was found by Mullikin through an oil drop experiment . [Charge of e – = 1.6 × 10 -16 C and Mass of e – = 9.1093 × 10 -31 kg].

Conclusions:

Based on conclusions from his cathode ray experiment, Thomson described the atomic structure as a positively charged sphere into which negatively charged electrons were embedded.

It is commonly referred to as the “plum pudding model” because it can be visualised as a plum pudding dish where the pudding describes the positively charged atom and the plum pieces describe the electrons.

Thomson’s atomic structure described atoms as electrically neutral, i.e., the positive and the negative charges were of equal magnitude.

Limitations of Thomson’s Atomic Structure:  Thomson’s atomic model does not clearly explain the stability of an atom. Also, further discoveries of other subatomic particles couldn’t be placed inside his atomic model.

Rutherford, a student of J. J. Thomson, modified the atomic structure with the discovery of another subatomic particle called “Nucleus” . His atomic model is based on the Alpha ray scattering experiment.

Alpha Ray Scattering Experiment

Construction:.

• A very thin gold foil of 1000 atoms thick is taken.
• Alpha rays (doubly charged Helium He 2+ ) were made to bombard the gold foil.
• Zn S screen is placed behind the gold foil.
• Most of the rays just went through the gold foil, making scintillations (bright spots) in the ZnS screen.
• A few rays got reflected after hitting the gold foil.
• One in 1000 rays got reflected by an angle of 180° (retraced path) after hitting the gold foil.
• Since most rays passed through, Rutherford concluded that most of the space inside the atom is empty.
• A few rays got reflected because of the repulsion of its positive with some other positive charge inside the atom.
• 1/1000th of the rays got strongly deflected because of a very strong positive charge in the centre of the atom. He called this strong positive charge “nucleus”.
• He said most of the charge and mass of the atom resides in the nucleus.

Rutherford’s Structure of Atom

Based on the above observations and conclusions, Rutherford proposed his own atomic structure , which is as follows.

• The nucleus is at the centre of an atom, where most of the charge and mass is concentrated.
• The atomic structure is spherical.
• Electrons revolve around the nucleus in a circular orbit, similar to the way planets orbit the sun.

Limitations of the Rutherford Atomic Model

• If electrons have to revolve around the nucleus, they will spend energy and that too against the strong force of attraction from the nucleus, a lot of energy will be spent by the electrons, and eventually, they will lose all their energy and will fall into the nucleus so the stability of atom is not explained.
• If electrons continuously revolve around the ‘nucleus, the type of spectrum expected is a continuous spectrum. But in reality, what we see is a line spectrum.

Atomic Structure – Rutherford’s Model, J.J Thomson’s Model

• Protons are positively charged subatomic particles. The charge of a proton is 1e, which corresponds to approximately 1.602 × 10 -19
• The mass of a proton is approximately 1.672 × 10 -24
• Protons are over 1800 times heavier than electrons.
• The total number of protons in the atoms of an element is always equal to the atomic number of the element.
• The mass of a neutron is almost the same as that of a proton, i.e., 1.674×10 -24
• Neutrons are electrically neutral particles and carry no charge.
• Different isotopes of an element have the same number of protons but vary in the number of neutrons present in their respective nuclei.
• The charge of an electron is -1e, which approximates to -1.602 × 10 -19
• The mass of an electron is approximately 9.1 × 10 -31 .
• Due to the relatively negligible mass of electrons, they are ignored when calculating the mass of an atom.

Nucleons are the components of the nucleus of an atom. A nucleon can either be a proton or a neutron. Each element has a unique number of protons in it, which is described by its unique atomic number . However, several atomic structures of an element can exist, which differ in the total number of nucleons.

These variants of elements having a different nucleon number (also known as the mass number) are called isotopes of the element. Therefore, the isotopes of an element have the same number of protons but differ in the number of neutrons.

The atomic structure of an isotope is described with the help of the chemical symbol of the element, the atomic number of the element and the mass number of the isotope. For example, there exist three known naturally occurring isotopes of hydrogen , namely, protium, deuterium and tritium. The atomic structures of these hydrogen isotopes are illustrated below.

The isotopes of an element vary in stability. The half-lives of isotopes also differ. However, they generally have similar chemical behaviour owing to the fact that they hold the same electronic structures .

Atomic Structures of Some Elements

The structure of an atom of an element can be simply represented via the total number of protons, electrons and neutrons present in it. The atomic structures of a few elements are illustrated below.

The most abundant isotope of hydrogen on the planet Earth is protium. The atomic number and the mass number of this isotope are 1 and 1, respectively.

Structure of Hydrogen Atom: This implies that it contains one proton, one electron and no neutrons (Total number of neutrons = Mass number – Atomic number)

Carbon has two stable isotopes – 12C and 13C. Of these isotopes, 12C has an abundance of 98.9%. It contains 6 protons, 6 electrons and 6 neutrons.

Structure of Carbon Atom: The electrons are distributed into two shells, and the outermost shell (valence shell) has four electrons. The tetravalency of carbon enables it to form a variety of chemical bonds with various elements.

There exist three stable isotopes of oxygen – 18O, 17O and 16O. However, oxygen-16 is the most abundant isotope.

Structure of Oxygen Atom: Since the atomic number of this isotope is 8 and the mass number is 16, it consists of 8 protons and 8 neutrons. 6 out of the 8 electrons in an oxygen atom lie in the valence shell.

Neils Bohr put forth his model of the atom in the year 1915. This is the most widely used atomic model to describe the atomic structure of an element which is based on Planck’s theory of quantization .

Postulates:

• The electrons inside atoms are placed in discrete orbits called “stationery orbits”.
• The energy levels of these shells can be represented via quantum numbers.
• Electrons can jump to higher levels by absorbing energy and move to lower energy levels by losing or emitting their energy.
• As long as an electron stays in its own stationery, there will be no absorption or emission of energy.
• Electrons revolve around the nucleus in these stationary orbits only.
• The energy of the stationary orbits is quantised.

Limitations of Bohr’s Atomic Theory:

• Bohr’s atomic structure works only for single electron species such as H, He+, Li2+, Be3+, ….
• When the emission spectrum of hydrogen was observed under a more accurate spectrometer, each line spectrum was seen to be a combination of a number of smaller discrete lines.
• Both Stark and Zeeman’s effects couldn’t be explained using Bohr’s theory.

Heisenberg’s uncertainty principle: Heisenberg stated that no two conjugate physical quantities could be measured simultaneously with 100% accuracy. There will always be some error or uncertainty in the measurement.

Drawback: Position and momentum are two such conjugate quantities that were measured accurately by Bohr (theoretically).

Stark effect: Phenomenon of deflection of electrons in the presence of an electric field.

Zeeman effect: Phenomenon of deflection of electrons in the presence of a magnetic field.

The electrons, which were treated to be particles, and the evidence of the photoelectric effect show they also have a wave nature. This was proved by Thomas Young with the help of his double-slit experiment .

De-Broglie concluded that since nature is symmetrical, so should light or any other matter wave be.

Quantum Numbers

• Principal Quantum Number (n): It denotes the orbital number or shell number of an electron.
• Azimuthal Quantum Numbers ( l ): It denotes the orbital (sub-orbit) of the electron.
• Magnetic Quantum Number: It denotes the number of energy states in each orbit.
• Spin Quantum number(s): It denotes the direction of spin, S = -½ = Anticlockwise and ½ = Clockwise.

Electronic Configuration of an Atom

The electrons have to be filled in the s, p, d and f in accordance with the following rule.

1. Aufbau’s principle: The filling of electrons should take place in accordance with the ascending order of energy of orbitals.

• Lower energy orbital should be filled first, and higher energy levels.
• The energy of orbital α(p + l) value it two orbitals have the same (n + l ) value, E α n
• Ascending order of energy 1s, 2s, 2p, 3s, 3p, 4s, 3d, . . .

2. Pauli’s exclusion principle: No two electrons can have all four quantum numbers to be the same, or if two electrons have to be placed in an energy state, they should be placed with opposite spies.

3. Hund’s rule of maximum multiplicity: In the case of filling degenerate (same energy) orbitals, all the degenerate orbitals have to be singly filled first, and then, only pairing has to happen.

Atomic Structure Solved Problems and Solutions

Atomic Structure – Important Questions

Structure of Atom Class 11 – Full Chapter Revision

Structure of Atom – Top 12 Most Important JEE Main Questions

Frequently Asked Questions on Atomic Structure

What are subatomic particles.

Subatomic particles are the particles that constitute an atom. Generally, this term refers to protons, electrons and neutrons.

How do the atomic structures of isotopes vary?

They vary in terms of the total number of neutrons present in the nucleus of the atom, which is described by their nucleon numbers.

What are the shortcomings of Bohr’s atomic model?

According to this atomic model, the structure of an atom offers poor spectral predictions for larger atoms. It also failed to explain the Zeeman effect. It could only successfully explain the hydrogen spectrum.

How can the total number of neutrons in the nucleus of a given isotope be determined?

The mass number of an isotope is given by the sum of the total number of protons and neutrons in it. The atomic number describes the total number of protons in the nucleus. Therefore, the number of neutrons can be determined by subtracting the atomic number from the mass number.

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UNIT 5- Atomic Structure Essay

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Atomic structure.

By Wasima Uddin In this essay, I will explain the distribution of mass and charges within an atom, highlighting the distinct characteristics of protons, neutrons, and electrons concerning their mass and charge. The focus will extend to the arrangement of mass and charges within an atom. Additionally, I will expound on the quantities of protons and neutrons within an atomic nucleus, as well as electron numbers in atoms and ions. The variations among isotopes will be explored based on differing neutron numbers. Furthermore, the essay will delve into the numbers and relative energies of the s, p, and d orbitals for principal quantum numbers 1-3, outlining the shapes of these orbitals. The concept of the first ionisation energy of an element will be discussed, taking into account electronic configurations obtained from successive ionisation data.

Atoms have a basic structure of a nucleus and shells. All atoms contain subatomic particles; each element will have a different number of these. An atom's mass, shape and structure are all determined by its number of subatomic particles. It is also important to note that the position of an element on the periodic table as well as the properties in which the element will withhold are also decided by its number of subatomic particles. “Most of the atoms are empty. The rest consists of a positively charged nucleus of protons and neutrons surrounded by a cloud of negatively charged electrons” (Trefil, 2019) Contained in the nucleus are protons and neutrons both of which have a relative mass of 1. Protons have a positive charge and neutrons have no charge the nucleus in shells are electrons which have a relative mass of 1/ and a negative charge. Due to the neutrons and protons having a larger mass, it means that most of an atom's mass is within the nucleus. Positively-charged protons in the nucleus of an atom are exactly balanced by the negatively charged electrons, meaning an atom is always neutral. Electrons move within electron shells around the nucleus. Although electron shells provide a spatial structure for electrons, they are not physical entities; instead, they represent energy levels extending outward from the nucleus. The number of electrons varies among different elements, but atoms of the same element share an identical electron count. Each shell has distinct energy levels, with higher-energy electrons located in shells farther from the nucleus. Shells closer to the nucleus can accommodate fewer electrons than those situated farther away, resulting in an increasing electron-holding capacity as the distance from the nucleus expands. Examining the periodic table provides crucial information about an element's atom. Each element on the table is accompanied by two numbers. The larger one, the mass number, differs for each element and represents the sum of protons and neutrons. The smaller number, the atomic number, corresponds to the electrons in atoms. Atoms with incomplete outer shells tend to be highly reactive and unstable,

called 2s and 2p and shell three has three sub-shells called 3s, 3p and 3d. Electrons behave in a number of different ways in order for this to be understood in greater detail, the principles of quantum mechanics should be explored. There are several different principles of quantum mechanics; firstly we have the Heisenberg uncertainty principle which states “you cannot determine the position and momentum of an electron at the same time” it essentially explains how the speed and location of an atom cannot be precisely calculated at the same time (Caltech, n) Then Pauli’s exclusion principle which explains “no two electrons can have the same four quantum numbers” or otherwise explained as more than two electrons can not occupy the same subshell and two electrons and these two electrons in the same subshell have opposite spins (sciencedirect, n.) There is also the Aufbau principle which states “electrons enter the lowest available energy level” which is to be thought of that electrons want to take the easiest route and therefore will go to the lower energy level. Then finally there are also Hund’s rules of maximum multiplicity which explores that "When possessing equal energy levels, electrons strive to avoid pairing." Considering these principles provides insight into the guidelines governing electron orbital filling. Shells are thought of to be spherical-shaped orbits around the nucleus in the centre; this model is primarily used as a straightforward way to represent them on paper, however the way in which the electrons in different subshells orbit the nucleus differs for each quantum number of the shell. As atoms are in a three-dimensional space, they need to be visualised as such, this can be done using the x,y and z-axis. S, d and p subshells all have different three-dimensional shapes (see appendix 2) As previously discussed, shells and subshells can be thought of to be levels of energy. The shells which are closest to the nucleus fill up first, this is because they are lower-energy shells, energy level increases the further the shells are from the nucleus. The order in which subshells fill up can be illustrated (see appendix 3)

Ionisation energy can be defined as the amount of energy it is required to remove one mole of electrons from one mole of atoms in the gaseous state. IE or ionisation energy is measured in kilojoules per mole (KJ mol -1). How much energy it takes to remove the first electron is called the First ionisation energy and it will produce a 1+ ion then the amount of energy it takes to remove the second electron is known as the second ionisation energy and will produce a 2+ion from a 1+ ion. The third ionisation energy will produce a 3+ ion from a 2+ ion. By looking at the periodic table we can make some conclusions about the ionisation energies of different elements, this is all to do with their location there. Generally, the ionisation energy increases going left to right along the periods as well as going from the bottom to the tops of groups on the periodic table. This general trend is correct, however (see appendix 4) looking at the trends in period 3 elements for first ionisation energies we can see the troughs on the graph which are essentially anomalies to the trend. A drop can be seen between aluminium and magnesium this is because the outer electron in aluminium is in a p subshell however the outer electron in magnesium is in an s subshell which is a lower energy level than the P subshell meaning it takes less energy to remove the outer electron in magnesium. A trough can be seen between sulphur and phosphorus, in sulphur, the outer shell 3p has 4 electrons or essentially 2 pairs, paired electrons are easier to remove than single ones, phosphorus only has 3 single electrons in its outer shell, which explains why the ionisation energy is lower for sulphur than phosphorus. What Is already known about IE data can be used to work out an atom's electron configuration that shows the sub-shells at peaks and troughs of ionisation energy. Elements with low ionisation energy can be considered reactive, and elements with high ionisation energy will be less reactive. From the information already known about ionisation energies, the correct position of an element on the periodic table can be justified.

Appendix 1 3 Isotopes of Carbon Diagram (Isotopes, MrReid) Appendix 2 S, p and d Subshells Diagram (s and d orbitals, general chemistry.steps)

Appendix 3 Order of Subshell Filling Diagram (electronic configuration, byjus) Appendix 4 First ionisation trends of period 3 elements Diagram (physical properties of period 3, secondaryscience4all)

Caltech (n.). What Is the Uncertainty Principle and Why Is It Important? Caltech Science Exchange. Available at: scienceexchange.caltech/topics/quantum-science-explained/ uncertainty-principle. (Accessed: 3rd December 2023) sciencedirect. (n.). Pauli Exclusion Principle - an overview Available at: sciencedirect/topics/pharmacology-toxicology-and-phar maceutical-science/pauli-exclusion-principle. (Accessed: 3rd December 2023) Appendices Mrreid. (n.) Potassium iodide pills are radioactive. Available at: wordpress.mrreid/2011/03/23/potassium-iodide-pills-are-radioa ctive/ Generalchemistrysteps. (n.). S-p-d-f Atomic orbitals. Available at: general.chemistrysteps/s-p-d-f-atomic-orbitals/ Secondary science 4 All. (2014) Physical properties of Period 3 Elements. Available at: secondaryscience4all.wordpress/2014/08/09/physical-proper ties-of-period-3-elements/ byjus.(n.). Electronic Configuration. Available at: byjus/chemistry/electron-configuration/

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Module : Unit 5 Atomic Structure Essay

University : stonebridge college.