1. Dalton's atomic theory proposed that atoms are indivisible, hard spheres that differ in properties based on their unique structure.
2. Rutherford's gold foil experiment showed that atoms have a small, dense nucleus containing their mass, with electrons orbiting the nucleus.
3. Bohr used Planck's quantum theory to explain electron orbits in hydrogen atoms, stabilizing Rutherford's model and beginning modern atomic theory.
Cbse 8 chemical effects of electric currentSupratim Das
This document contains a chemistry assignment with multiple choice and short answer questions about electrical conductivity and chemical effects of electric current. Some key points covered include:
- Materials that can and cannot be easily charged by friction.
- The charges acquired when a glass rod is rubbed with silk cloth.
- Definitions of terms like electrolysis, electroplating, and electrodes.
- Why a charged body loses its charge when touched.
- Drawings and explanations of instruments that can detect a charged body.
- Differences in conductivity between pure water and other liquids.
- Processes involved in electrolysis and electroplating.
Transistors and integrated circuits are important components in electronics. A transistor is a semiconductor device with three electrical contacts that can be used as an amplifier, detector, or switch. An integrated circuit is a circuit composed of transistors, resistors, and capacitors constructed on a single semiconductor chip, where the components are interconnected to perform a given function. A bipolar junction transistor consists of a three-layer sandwich of doped semiconductor materials (either PNP or NPN) where a small base current controls a larger collector current flowing between the emitter and collector. Integrated circuits allow many transistors to be packed onto a single chip to construct more complex circuits.
Metallic bonding involves a lattice of positive metal ions surrounded by delocalized electrons that form a negative "electron cloud". This electron cloud binds the positively charged ions together and is responsible for metals' properties. Metals exhibit high electrical conductivity because the mobile electron cloud allows electrons to move freely throughout the structure. The strength of metallic bonding depends on electron density and ionic radius - greater electron density and smaller ionic radius lead to stronger bonding and higher melting points. Alloys are mixtures of metals designed to enhance properties like strength, corrosion resistance, magnetism, and ductility.
Metallic bonding occurs when metal atoms bond through delocalized electrons. Metals form a lattice structure with positive metal ions surrounded by a sea of delocalized electrons that are attracted to the positive ions. This bonding allows metals to have distinctive physical and chemical properties including malleability, conductivity, and high melting points. Metallic bonding is important for creating alloys and is found in many materials used in daily life.
Presentation on Metallic Bond and its nature, presented by Engineer S.M. Wahid Mahmud from Daffodil International University from the department of Electronics & Telecommunication, Faculty of Science & IT.
Metallic bonds and the properties of metalsKamal Metwalli
This document describes metallic bonding and the properties of metals and alloys. It explains that metallic bonding results from the delocalization of valence electrons among metal atoms, forming a "sea" of electrons. Metals are good conductors of heat and electricity due to the mobility of these electrons. The properties of a metal, such as its hardness and melting point, depend on the number of delocalized electrons. Alloys are mixtures of elements that are substitutional, with one metal replacing atoms of another, or interstitial, with smaller atoms filling holes in the metallic crystal structure.
The document discusses the different types of bonding mechanisms that hold atoms together in solids, including ionic bonding, covalent bonding, metallic bonding, van der Waals bonding, and hydrogen bonding. Ionic bonding involves the transfer of electrons between metals and nonmetals, covalent bonding involves the sharing of electrons between nonmetals, and metallic bonding involves the delocalization of electrons among metal atoms. Weaker van der Waals forces result from induced dipole interactions between molecules. Hydrogen bonding is a special type of dipole-dipole interaction that occurs when hydrogen is bonded to highly electronegative atoms like oxygen or fluorine.
Cbse 8 chemical effects of electric currentSupratim Das
This document contains a chemistry assignment with multiple choice and short answer questions about electrical conductivity and chemical effects of electric current. Some key points covered include:
- Materials that can and cannot be easily charged by friction.
- The charges acquired when a glass rod is rubbed with silk cloth.
- Definitions of terms like electrolysis, electroplating, and electrodes.
- Why a charged body loses its charge when touched.
- Drawings and explanations of instruments that can detect a charged body.
- Differences in conductivity between pure water and other liquids.
- Processes involved in electrolysis and electroplating.
Transistors and integrated circuits are important components in electronics. A transistor is a semiconductor device with three electrical contacts that can be used as an amplifier, detector, or switch. An integrated circuit is a circuit composed of transistors, resistors, and capacitors constructed on a single semiconductor chip, where the components are interconnected to perform a given function. A bipolar junction transistor consists of a three-layer sandwich of doped semiconductor materials (either PNP or NPN) where a small base current controls a larger collector current flowing between the emitter and collector. Integrated circuits allow many transistors to be packed onto a single chip to construct more complex circuits.
Metallic bonding involves a lattice of positive metal ions surrounded by delocalized electrons that form a negative "electron cloud". This electron cloud binds the positively charged ions together and is responsible for metals' properties. Metals exhibit high electrical conductivity because the mobile electron cloud allows electrons to move freely throughout the structure. The strength of metallic bonding depends on electron density and ionic radius - greater electron density and smaller ionic radius lead to stronger bonding and higher melting points. Alloys are mixtures of metals designed to enhance properties like strength, corrosion resistance, magnetism, and ductility.
Metallic bonding occurs when metal atoms bond through delocalized electrons. Metals form a lattice structure with positive metal ions surrounded by a sea of delocalized electrons that are attracted to the positive ions. This bonding allows metals to have distinctive physical and chemical properties including malleability, conductivity, and high melting points. Metallic bonding is important for creating alloys and is found in many materials used in daily life.
Presentation on Metallic Bond and its nature, presented by Engineer S.M. Wahid Mahmud from Daffodil International University from the department of Electronics & Telecommunication, Faculty of Science & IT.
Metallic bonds and the properties of metalsKamal Metwalli
This document describes metallic bonding and the properties of metals and alloys. It explains that metallic bonding results from the delocalization of valence electrons among metal atoms, forming a "sea" of electrons. Metals are good conductors of heat and electricity due to the mobility of these electrons. The properties of a metal, such as its hardness and melting point, depend on the number of delocalized electrons. Alloys are mixtures of elements that are substitutional, with one metal replacing atoms of another, or interstitial, with smaller atoms filling holes in the metallic crystal structure.
The document discusses the different types of bonding mechanisms that hold atoms together in solids, including ionic bonding, covalent bonding, metallic bonding, van der Waals bonding, and hydrogen bonding. Ionic bonding involves the transfer of electrons between metals and nonmetals, covalent bonding involves the sharing of electrons between nonmetals, and metallic bonding involves the delocalization of electrons among metal atoms. Weaker van der Waals forces result from induced dipole interactions between molecules. Hydrogen bonding is a special type of dipole-dipole interaction that occurs when hydrogen is bonded to highly electronegative atoms like oxygen or fluorine.
Definition of ionic compounds
Property of ionic compoundIn crystal formSolubility in waterHigh melting and boiling pointConductivity of electricity
Use of ionic compoundUse as dryersUse to make pans
The document provides an overview of chemical bonding principles including ionic bonding, covalent bonding, and molecular geometry. It discusses how atoms interact to achieve stable electronic configurations through ionic bonding by exchanging electrons or covalent bonding by sharing electrons. Lewis structures are introduced as a way to represent valence electrons in molecules and determine molecular geometry and polarity based on electron pair arrangements around central atoms. Key concepts covered include octet rule, electronegativity, bond polarity, and using Lewis structures to systematically determine molecular structure characteristics.
Metallic bonding results from the attraction between metal cations and delocalized electrons in the "sea" of electrons. This allows electrons to move freely throughout the metal and form metallic bonds between atoms. Metallic bonding gives metals properties like high melting points, conductivity of heat and electricity, and malleability.
This document provides an overview of ionic bonding. Ionic bonds form between elements when one atom loses electrons to become a positively charged ion and another atom gains those electrons to become a negatively charged ion. For example, in sodium chloride, sodium atoms lose electrons to form Na+ ions and chloride atoms gain electrons to form Cl- ions. The oppositely charged ions are then held together by electrostatic forces in a repeating crystal lattice structure. Ionic compounds have high melting points, are brittle, and do not conduct electricity in solid form but do conduct when molten or dissolved in water as the ions become mobile.
1. Materials science is the study of relationships between the structure and properties of materials. It relates how the atomic and molecular structure of a material influences its properties.
2. A material's properties determine how it responds to external forces and the environment. Key properties include mechanical, electrical, thermal, optical, and chemical properties. Mechanical properties describe response to forces like strength and toughness.
3. There are three main classes of materials: metals, ceramics, and polymers. Metals are strong, ductile, and conductive. Ceramics are brittle but heat resistant. Polymers are lightweight and insulating. Materials science helps understand materials and design new components.
Ionic, covalent, metallic, hydrogen, and Van Der Wall's bonding are the main types of bonding. Ionic bonding occurs between positive and negative ions, covalent bonding involves sharing electrons between atoms, and metallic bonding involves loosely bound electrons in metal atoms that form an electron cloud. Hydrogen bonding arises from electrostatic attraction between hydrogen atoms partially charged by covalent bonds. Van Der Wall's bonding involves weak, temporary bonds between molecules of the same substance.
Calcium ion and bromide ion would form Calcium bromide. Potassium ion and sulfide ion would form Potassium sulfide. Aluminum ion and selenide ion would form Aluminum selenide. Metallic bonds form between metal cations through a "sea of electrons" shared between all the metal atoms. Metals are ductile and good conductors of electricity due to the free movement of delocalized valence electrons between metal cations. Alloys can have superior properties to their component elements.
This document summarizes the different types of chemical bonding: ionic bonding occurs between oppositely charged ions and is electrostatic in nature; covalent bonding involves sharing electrons between atoms and can range from partially ionic to completely covalent; metallic bonding is characterized by positive ions floating in a "sea" of delocalized electrons allowing metals to conduct electricity; van der Waals bonding is the weakest type and only present between inert gases due to induced dipole interactions. The document also discusses the different arrangements of spheres to model ion packing in cubic close, body-centered cubic, and hexagonal close packing of metals.
Metallic bonds form macromolecular structures that are ductile and malleable, allowing metals to be shaped. Metallic bonds can also conduct electricity due to free electrons that carry charges. Examples of materials with macromolecular structures are diamond and graphite, with diamond having a tetrahedral structure where each carbon atom is bonded to four other carbons.
This document discusses three main types of chemical bonds: ionic bonds, covalent bonds, and metallic bonds.
1. Ionic bonds form between metals and nonmetals by the transfer of electrons from one atom to another, resulting in positively and negatively charged ions.
2. Covalent bonds form between nonmetals of similar electronegativity by the sharing of electron pairs between atoms.
3. Metallic bonds form between metallic elements and involve a sea of electrons that hold the metal atoms together strongly.
The document discusses the properties of metals and their crystalline structure. It begins by explaining that metals have a closely packed crystalline structure, usually face centered cubic, body centered cubic, or hexagonal close packed. This gives metals their high conductivity of heat and electricity as well as their malleability, ductility, and high melting and boiling points. Metals also have metallic luster and can emit electrons through thermionic or photoemissive processes due to their mobile electrons.
This document discusses three main types of chemical bonding: ionic bonds, covalent bonds, and metallic bonds. Ionic bonds form between metals and nonmetals through the transfer of electrons from one atom to another, producing charged ions. Covalent bonds form between nonmetals of similar electronegativity through the sharing of electron pairs. Metallic bonds form between metallic elements through a shared "electron cloud" that holds the atoms together strongly.
Metallic bonding is the strong attraction between closely packed positive metal ions and delocalized electrons. This bonding gives metals high melting and boiling points. Metals are good conductors of heat and electricity due to their delocalized electrons, which are free to move and transfer energy. Their dense, closely packed atomic structure results in high density as well as malleability and ductility, as the layers of atoms can easily slide past one another.
This document discusses ionic compounds and the formation of ionic bonds. Ions form when atoms gain or lose electrons to achieve stable electron configurations like noble gases. Ionic bonds occur between oppositely charged ions and result in crystalline solids with high melting points. The document explains how to name ionic compounds based on the cation and anion present and write chemical formulas from compound names. It also briefly discusses metallic bonding and the properties of metals and alloys.
Interatomic bonds determine the macroscopic properties of materials. Ionic bonds in ceramics result in high melting temperatures, elastic moduli, and low thermal expansion. Metallic bonds produce intermediate properties due to varying bond strengths. Covalent and secondary bonds in polymers lead to low melting points, moduli and high thermal expansion, with secondary bonds dominating behavior. Bonding energy correlates with melting temperature, while curvature relates to thermal expansion and initial elastic response.
Ionic bonds form between oppositely charged ions. They result from the transfer of electrons from one atom to another. Ionic bonding typically occurs between metals and nonmetals. Metals tend to lose electrons to fill their outer shell, becoming positively charged ions, while nonmetals gain electrons to fill their outer shell, becoming negatively charged ions. The electrostatic attraction between the opposite charges of the ions forms the ionic bond.
This document discusses conducting materials used in dye-sensitized solar cells. It begins by providing background on solar cells and photovoltaics, and describes dye-sensitized solar cells. It then focuses on different types of conducting materials used in these cells, including conductive polymers and electrolyte systems using ionic liquids. The document concludes by discussing the promising results of using ionic liquid electrolytes to optimize the performance of dye-sensitized solar cells.
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Metallic bonding occurs when metal atoms lose valence electrons to become positively charged ions embedded in a "sea" of delocalized electrons. This electron sea model explains several properties of metals, including their ability to conduct electricity and heat, as well as their malleability, ductility, lustrous appearance, and high melting and boiling points. The mobile electrons allow for heat and charge conduction, while the metallic lattice structure enables atoms to slide past one another under pressure.
IUPAC NOMENCLATURE_ORGANIC_for JEE(MAIN)-JEE(ADVANCED)-NEETSupratim Das
This document discusses IUPAC nomenclature rules for naming organic compounds. It begins by listing common names and IUPAC names for some simple organic molecules. It then describes the system for naming hydrocarbons based on identifying the parent chain, numbering carbons, and indicating substituents. Rules are provided for naming saturated and unsaturated compounds, cyclic compounds, branched compounds, and compounds containing common functional groups like alcohols, aldehydes, ketones, acids, and others. Substituted benzene compounds are also discussed. The goal is to systematically name compounds to identify parent structures and functional groups.
The four quantum numbers - principal (n), azimuthal (l), magnetic (m), and spin (s) - are used to completely characterize electrons in an atom. The principal quantum number indicates the electron's energy level. The azimuthal number denotes the subshell and shape of its orbital. The magnetic number determines the number of orbitals in a subshell. The spin number arises from the electron's intrinsic spin and can have values of ±1/2. Together, the four quantum numbers uniquely identify each electron in an atom.
Definition of ionic compounds
Property of ionic compoundIn crystal formSolubility in waterHigh melting and boiling pointConductivity of electricity
Use of ionic compoundUse as dryersUse to make pans
The document provides an overview of chemical bonding principles including ionic bonding, covalent bonding, and molecular geometry. It discusses how atoms interact to achieve stable electronic configurations through ionic bonding by exchanging electrons or covalent bonding by sharing electrons. Lewis structures are introduced as a way to represent valence electrons in molecules and determine molecular geometry and polarity based on electron pair arrangements around central atoms. Key concepts covered include octet rule, electronegativity, bond polarity, and using Lewis structures to systematically determine molecular structure characteristics.
Metallic bonding results from the attraction between metal cations and delocalized electrons in the "sea" of electrons. This allows electrons to move freely throughout the metal and form metallic bonds between atoms. Metallic bonding gives metals properties like high melting points, conductivity of heat and electricity, and malleability.
This document provides an overview of ionic bonding. Ionic bonds form between elements when one atom loses electrons to become a positively charged ion and another atom gains those electrons to become a negatively charged ion. For example, in sodium chloride, sodium atoms lose electrons to form Na+ ions and chloride atoms gain electrons to form Cl- ions. The oppositely charged ions are then held together by electrostatic forces in a repeating crystal lattice structure. Ionic compounds have high melting points, are brittle, and do not conduct electricity in solid form but do conduct when molten or dissolved in water as the ions become mobile.
1. Materials science is the study of relationships between the structure and properties of materials. It relates how the atomic and molecular structure of a material influences its properties.
2. A material's properties determine how it responds to external forces and the environment. Key properties include mechanical, electrical, thermal, optical, and chemical properties. Mechanical properties describe response to forces like strength and toughness.
3. There are three main classes of materials: metals, ceramics, and polymers. Metals are strong, ductile, and conductive. Ceramics are brittle but heat resistant. Polymers are lightweight and insulating. Materials science helps understand materials and design new components.
Ionic, covalent, metallic, hydrogen, and Van Der Wall's bonding are the main types of bonding. Ionic bonding occurs between positive and negative ions, covalent bonding involves sharing electrons between atoms, and metallic bonding involves loosely bound electrons in metal atoms that form an electron cloud. Hydrogen bonding arises from electrostatic attraction between hydrogen atoms partially charged by covalent bonds. Van Der Wall's bonding involves weak, temporary bonds between molecules of the same substance.
Calcium ion and bromide ion would form Calcium bromide. Potassium ion and sulfide ion would form Potassium sulfide. Aluminum ion and selenide ion would form Aluminum selenide. Metallic bonds form between metal cations through a "sea of electrons" shared between all the metal atoms. Metals are ductile and good conductors of electricity due to the free movement of delocalized valence electrons between metal cations. Alloys can have superior properties to their component elements.
This document summarizes the different types of chemical bonding: ionic bonding occurs between oppositely charged ions and is electrostatic in nature; covalent bonding involves sharing electrons between atoms and can range from partially ionic to completely covalent; metallic bonding is characterized by positive ions floating in a "sea" of delocalized electrons allowing metals to conduct electricity; van der Waals bonding is the weakest type and only present between inert gases due to induced dipole interactions. The document also discusses the different arrangements of spheres to model ion packing in cubic close, body-centered cubic, and hexagonal close packing of metals.
Metallic bonds form macromolecular structures that are ductile and malleable, allowing metals to be shaped. Metallic bonds can also conduct electricity due to free electrons that carry charges. Examples of materials with macromolecular structures are diamond and graphite, with diamond having a tetrahedral structure where each carbon atom is bonded to four other carbons.
This document discusses three main types of chemical bonds: ionic bonds, covalent bonds, and metallic bonds.
1. Ionic bonds form between metals and nonmetals by the transfer of electrons from one atom to another, resulting in positively and negatively charged ions.
2. Covalent bonds form between nonmetals of similar electronegativity by the sharing of electron pairs between atoms.
3. Metallic bonds form between metallic elements and involve a sea of electrons that hold the metal atoms together strongly.
The document discusses the properties of metals and their crystalline structure. It begins by explaining that metals have a closely packed crystalline structure, usually face centered cubic, body centered cubic, or hexagonal close packed. This gives metals their high conductivity of heat and electricity as well as their malleability, ductility, and high melting and boiling points. Metals also have metallic luster and can emit electrons through thermionic or photoemissive processes due to their mobile electrons.
This document discusses three main types of chemical bonding: ionic bonds, covalent bonds, and metallic bonds. Ionic bonds form between metals and nonmetals through the transfer of electrons from one atom to another, producing charged ions. Covalent bonds form between nonmetals of similar electronegativity through the sharing of electron pairs. Metallic bonds form between metallic elements through a shared "electron cloud" that holds the atoms together strongly.
Metallic bonding is the strong attraction between closely packed positive metal ions and delocalized electrons. This bonding gives metals high melting and boiling points. Metals are good conductors of heat and electricity due to their delocalized electrons, which are free to move and transfer energy. Their dense, closely packed atomic structure results in high density as well as malleability and ductility, as the layers of atoms can easily slide past one another.
This document discusses ionic compounds and the formation of ionic bonds. Ions form when atoms gain or lose electrons to achieve stable electron configurations like noble gases. Ionic bonds occur between oppositely charged ions and result in crystalline solids with high melting points. The document explains how to name ionic compounds based on the cation and anion present and write chemical formulas from compound names. It also briefly discusses metallic bonding and the properties of metals and alloys.
Interatomic bonds determine the macroscopic properties of materials. Ionic bonds in ceramics result in high melting temperatures, elastic moduli, and low thermal expansion. Metallic bonds produce intermediate properties due to varying bond strengths. Covalent and secondary bonds in polymers lead to low melting points, moduli and high thermal expansion, with secondary bonds dominating behavior. Bonding energy correlates with melting temperature, while curvature relates to thermal expansion and initial elastic response.
Ionic bonds form between oppositely charged ions. They result from the transfer of electrons from one atom to another. Ionic bonding typically occurs between metals and nonmetals. Metals tend to lose electrons to fill their outer shell, becoming positively charged ions, while nonmetals gain electrons to fill their outer shell, becoming negatively charged ions. The electrostatic attraction between the opposite charges of the ions forms the ionic bond.
This document discusses conducting materials used in dye-sensitized solar cells. It begins by providing background on solar cells and photovoltaics, and describes dye-sensitized solar cells. It then focuses on different types of conducting materials used in these cells, including conductive polymers and electrolyte systems using ionic liquids. The document concludes by discussing the promising results of using ionic liquid electrolytes to optimize the performance of dye-sensitized solar cells.
FellowBuddy.com is a platform which has been setup with a simple vision, keeping in mind the dynamic requirements of students.
Our Vision & Mission - Simplifying Students Life
Our Belief - “The great breakthrough in your life comes when you realize it, that you can learn anything you need to learn; to accomplish any goal that you have set for yourself. This means there are no limits on what you can be, have or do.”
Like Us - https://www.facebook.com/FellowBuddycom-446240585585480
Metallic bonding occurs when metal atoms lose valence electrons to become positively charged ions embedded in a "sea" of delocalized electrons. This electron sea model explains several properties of metals, including their ability to conduct electricity and heat, as well as their malleability, ductility, lustrous appearance, and high melting and boiling points. The mobile electrons allow for heat and charge conduction, while the metallic lattice structure enables atoms to slide past one another under pressure.
IUPAC NOMENCLATURE_ORGANIC_for JEE(MAIN)-JEE(ADVANCED)-NEETSupratim Das
This document discusses IUPAC nomenclature rules for naming organic compounds. It begins by listing common names and IUPAC names for some simple organic molecules. It then describes the system for naming hydrocarbons based on identifying the parent chain, numbering carbons, and indicating substituents. Rules are provided for naming saturated and unsaturated compounds, cyclic compounds, branched compounds, and compounds containing common functional groups like alcohols, aldehydes, ketones, acids, and others. Substituted benzene compounds are also discussed. The goal is to systematically name compounds to identify parent structures and functional groups.
The four quantum numbers - principal (n), azimuthal (l), magnetic (m), and spin (s) - are used to completely characterize electrons in an atom. The principal quantum number indicates the electron's energy level. The azimuthal number denotes the subshell and shape of its orbital. The magnetic number determines the number of orbitals in a subshell. The spin number arises from the electron's intrinsic spin and can have values of ±1/2. Together, the four quantum numbers uniquely identify each electron in an atom.
1) The document discusses the topic-structure of an atom and summarizes the key discoveries that led to modern atomic theory, including the discovery of the electron, proton, neutron, and development of atomic models.
2) It describes Michael Faraday's experiments in the 1830s that provided early insights into atomic structure and the discovery of the electron in the 1850s from cathode ray experiments.
3) The document also summarizes Bohr's 1913 model of the hydrogen atom which explained its spectral lines by postulating stable electron orbits, and the development of quantum mechanics and Schrodinger's equation to more fully describe atomic structure.
This Presentation was prepared to help the readers to get the basic ideas for learning about the concepts of Quantum Numbers in Elementary Partcles ...
Chemistry deals with the composition, structure, and properties of matter. Matter exists as solids, liquids, or gases depending on temperature and pressure. Pure substances consist of either one type of atom (elements) or combinations of atoms in a fixed ratio (compounds). Mixtures contain varying ratios of different substances. Dalton's atomic theory states that matter is made of atoms, atoms of different elements have different masses, and compounds form when atoms combine in fixed ratios.
Chemistry question paper for cbse 10 carbon & its compoundsSupratim Das
This document contains a chemistry question paper for CBSE with 26 multiple choice questions about carbon and its compounds. The questions cover topics like carbon dioxide bonding, structural isomers of pentane, properties of carbon enabling many compounds, functional groups of compounds, oxidation reactions, soap and detergent properties, and mechanisms of chemical reactions.
This document provides study material for class 12 on p-block elements. It includes 16 questions with answers on topics like why pentahalides are more covalent than trihalides, the conditions required to maximize ammonia yield in the Haber process, and the nature of the bonds in SO2. It also lists important oxides, oxyacids, and properties of elements in the oxygen family and provides information on interhalogens, hydrogen halides, and xenon compounds.
1) Rutherford conducted an experiment where he bombarded a thin gold foil with alpha particles. He observed that most alpha particles passed through without deflection, but some were deflected at large angles, indicating the positive charge in an atom is concentrated in a small nucleus.
2) Bohr modified Rutherford's model by proposing electrons orbit the nucleus in fixed, quantized energy levels. Electrons can jump between these levels, emitting or absorbing photons of specific frequencies.
3) Frank and Hertz conducted an experiment where they observed sharp drops in current through a mercury vapor cathode at multiples of 4.9V. This provided direct evidence that electrons in atoms can only occupy discrete energy levels, as predicted by Bohr
The document discusses the history of atomic structure models from Democritus' idea of atoms to Bohr's model. Some key points:
1. J.J. Thomson's experiments in 1897 led him to propose the "plum pudding" model where electrons were embedded in a uniform positive charge.
2. Rutherford's gold foil experiment in 1911 showed that the atom has a small, dense, positively charged nucleus at its center.
3. Bohr modified Rutherford's model in 1913 to propose that electrons orbit the nucleus in discrete energy levels, explaining atomic line spectra. When electrons fall from higher to lower orbits, photons are emitted.
1. Atoms are the basic building blocks of matter and consist of a small, dense nucleus surrounded by electrons.
2. Rutherford's gold foil experiment in 1911 showed that the atom has a small, dense nucleus containing positively charged protons and uncharged neutrons.
3. Niels Bohr proposed his model of the atom in 1913 where electrons orbit the nucleus in fixed shells at specific energy levels, explaining atomic spectra. However, it did not explain more complex atomic structures.
1. Atoms are the basic building blocks of matter and consist of a small, dense nucleus surrounded by electrons.
2. Rutherford's gold foil experiment in 1911 showed that the atom has a small, dense nucleus containing positively charged protons and uncharged neutrons.
3. Niels Bohr proposed his model of the atom in 1913 in which electrons orbit the nucleus in fixed shells at specific energy levels, explaining atomic spectra.
1. Atoms are the basic building blocks of matter and consist of a small, dense nucleus surrounded by electrons.
2. Rutherford's gold foil experiment in 1911 showed that the atom has a small, dense nucleus containing positively charged protons and uncharged neutrons.
3. Niels Bohr proposed his model of the atom in 1913 where electrons orbit the nucleus in fixed shells at specific energy levels, explaining atomic spectra. However, it did not explain more complex atomic structures.
Ernest Rutherford's alpha ray scattering experiment led him to propose the nuclear model of the atom. The key findings were:
1) Most alpha particles passed through the thin gold foil with little deflection, but a small percentage were deflected by large angles, including backwards.
2) This could only be explained if the positive charge of the atom was concentrated into a very small, dense nucleus.
3) Rutherford concluded atoms have a small, dense nucleus containing its positive charge and mass, with electrons orbiting the nucleus.
This nuclear model replaced the plum pudding model, but had its own limitations that were later addressed by Niels Bohr's model of electron orbits and quantization
This presentation is specially made for the students of grades 11 and 12 of High School. This is the presentation of chapter Atomic Structure with proper diagrams, figures, facts, mnemonics, and some repeated past questions. Here you will get a chance to know about Atomic theory, Daltons Law, particles and so on.
In these slides, I covered the following topics with PYQ's of CH-12 (Atom) of class 12th Physics:
-Alpha-particle scattering experiment
-Rutherford's model of the atom
-Bohr model,
-Energy levels,
-Hydrogen spectrum
The document discusses the development of atomic theory and models of atomic structure based on experiments. Key points include:
1. Early experiments with cathode ray tubes led Thomson to discover the electron and determine its small mass and negative charge.
2. Rutherford's gold foil experiment showed that the mass and positive charge of atoms are concentrated in a very small, dense nucleus.
3. Later experiments discovered the proton in the nucleus and the neutron, establishing the main subatomic particles that make up all elements.
4. Models evolved from Thomson's "plum pudding" model to Rutherford's nuclear model to better explain experimental results and the stability of atoms.
The document discusses atomic structure and models of the atom. It describes J.J. Thomson's discovery of the electron and the plum pudding model. It then summarizes Rutherford's gold foil experiment, which led to the nuclear model of the atom with a small, dense nucleus surrounded by electrons. The document also discusses atomic spectra and how they provided evidence for discrete energy levels of electrons within atoms.
In your previous class you have already studies about the structure of an atom but some of the exception you can learn here in this chapter how the structure of an atom is fully defined
This document provides an overview of general chemistry concepts related to atomic structure. It discusses several atomic models proposed over time, including Dalton's atomic theory, Thomson's "plum pudding" model, Rutherford's nuclear model, and Bohr's model. It also describes experiments like Thomson's cathode ray experiment and Rutherford's alpha scattering experiment that helped develop understanding of atomic structure. Key topics covered include the discovery of subatomic particles like protons, neutrons, and electrons, isotopes, ionization, and atomic spectra.
1) Experiments with cathode ray tubes led to the discovery of the electron as a negatively charged fundamental particle.
2) Further experiments showed that atoms are mostly empty space and contain a small, dense nucleus made up of protons and neutrons, around which electrons orbit.
3) The photoelectric effect showed that light behaves as a particle (photon) rather than just a wave, transferring its energy in discrete quantized amounts to electrons and ejecting them from metal surfaces.
The document discusses the atomic structure and models of the atom. It begins with an acknowledgement and table of contents. It then covers Dalton's atomic theory, subatomic particles like electrons and protons, cathode rays and the discovery of electrons. It discusses the charge to mass ratio of electrons, the discovery of protons and neutrons, and models of the atom including Thomson's model and Rutherford's nuclear model. It also addresses isotopes, limitations of models, wave nature of radiation, and the electromagnetic spectrum.
Here is a semi-log plot of the data with an exponential trendline:
The equation of the trendline is:
y = 12456e-0.4693x
Taking the natural log of both sides:
ln(y) = ln(12456) - 0.4693x
The slope is -0.4693
Using the equation:
t1/2 = 0.693/λ
λ = 0.4693
t1/2 = 0.693/0.4693 = 1.5
Therefore, the half-life of the isotope is 1.5 intervals, or 1.5 x 30 s = 45 seconds.
The document summarizes three major atomic models:
1. J.J. Thomson's model from 1898 viewed the atom as a uniform positive sphere with electrons embedded inside.
2. Rutherford's 1911 model resulted from experiments showing most alpha particles passed through a gold foil, but a small fraction bounced back, indicating a small, dense nucleus.
3. Niels Bohr's 1913 model incorporated Planck's quantum theory, proposing electrons orbit in discrete energy levels and do not radiate while in the same energy level.
1) Atoms are the basic building blocks of matter and consist of a nucleus surrounded by electrons. The nucleus contains protons and neutrons, while electrons orbit around the nucleus.
2) Rutherford's gold foil experiment provided evidence that atoms have a small, dense nucleus and that most of an atom's mass and positive charge is concentrated in the nucleus.
3) Bohr's model improved upon earlier models by proposing that electrons orbit in fixed shells and energy levels around the nucleus, explaining the stability of atoms and emission of photons during changes in electron energy levels.
- The document discusses the structure of the atom, beginning with John Dalton's model of indivisible atoms and progressing to modern atomic structure based on experiments like Rutherford's gold foil experiment.
- It describes subatomic particles like electrons, protons, and neutrons that make up atoms, and models of atomic structure proposed by scientists like Thomson, Rutherford, Bohr, and Moseley.
- Key aspects of atomic structure covered include the Bohr model of electron orbits, calculation of orbital radii and velocities, and the relationship between potential energy, kinetic energy and total energy of electrons in atoms.
The document discusses the history of the development of atomic structure models from Thomson's plum pudding model to Rutherford's nuclear model. Key events include J.J. Thomson's discovery of the electron, Millikan's oil drop experiment determining the charge of an electron, discovery of the proton through canal ray experiments, Rutherford's alpha particle scattering experiment revealing the dense nucleus at the center of the atom, and Rutherford proposing the nuclear model of the atom. The nuclear model represented a major breakthrough but did not fully explain electron stability.
1. The document discusses the discovery of subatomic particles like electrons, protons, and neutrons through various experiments in the late 19th and early 20th centuries.
2. Key discoveries include J.J. Thomson identifying electrons through cathode ray experiments, E. Goldstein discovering protons through canal ray experiments, and Chadwick discovering neutrons through bombardment of beryllium with alpha particles.
3. The document provides details on the charge, relative charge, and mass of these fundamental subatomic particles.
Similar to Structure Of Atom_STUDY MATERIALS_JEE-MAIN_AIPMT (20)
- The document discusses the p-block elements in the boron family, including their occurrence, characteristics, chemical properties, and compounds.
- Boron is a rare element found in borax and boric acid, while aluminum is the most abundant metal found in bauxite. Gallium, indium, and thallium are less common and found in sulfide ores.
- The elements have the outer electronic configuration ns2np1 and exist in oxidation states of +1 and +3, with the stability of +1 increasing down the group. They form covalent bonds and react to form oxides, nitrides, halides, and hydroxides.
Some basic concepts of chemistry (JEE - NEET)Supratim Das
This document provides an overview of key chemistry concepts for exams. It defines matter and its three states. Compounds have a constant composition while mixtures have a variable composition. It also defines valence and provides formulas for common cations and anions. Additional concepts covered include the mole, atomic and molecular weights, gas laws, empirical and molecular formulas, and several important reaction types. Concentration terms like density, specific gravity, and percentages are also defined.
Basic concepts of chemistry class 8, 9, 10Supratim Das
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Transition elements have several key characteristics:
(1) They exhibit variable oxidation states unlike s-block and p-block elements.
(2) Many of their compounds are colored due to electron transitions between d-orbital energy levels.
(3) They have a great tendency to form complex compounds due to their small, highly charged ions and vacant low energy d-orbitals.
This document summarizes several important gas laws and concepts related to the behavior of gases:
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3) Finally, it briefly touches on the van der Waals equation, critical constants, liquefaction of gases, and the ratio of molar heat capacities for different types of gases.
The document provides answers to 11 questions about acids, bases and salts. It defines strong and weak acids, and lists differences between acids and bases. It explains that acetic acid is in vinegar, ammonia is basic. Litmus solution comes from lichens and is used as a pH indicator. Distilled water is neutral as shown by litmus. Neutralization reactions form salts and water. Antacids neutralize excess stomach acid; calamine neutralizes ant bites. Factory waste is neutralized before disposal. The last questions discuss indicator color changes with acids and bases.
Optical isomerism, a form of stereoisomerism, arises when molecules have the same molecular formula but different spatial arrangements of atoms. Stereoisomers that are non-superimposable mirror images of each other are called enantiomers. Enantiomers rotate polarized light in opposite directions and are described as (+) and (-) forms. Molecules with four different substituents attached to a carbon atom, making it chiral, can exhibit optical isomerism through enantiomers. Skeletal formulas can identify chiral carbons if each has four distinct bonds indicated. Rings and multiple chiral carbons can lead to many possible stereoisomers.
Unit 10 studynotes-haloarenes-for Jee-NeetSupratim Das
Haloarenes are unreactive towards nucleophilic substitution due to resonance effects, differences in carbon-halogen hybridization, and instability of phenyl cations. They undergo substitution only under drastic conditions, but electron-withdrawing groups increase reactivity. Haloarenes also undergo halogenation, sulfonation, and Friedel-Crafts reactions, with halogens directing ortho/para substitution due to inductive and resonance effects. They react with metals in Wurtz-Fittig and Fittig reactions to form alkylarenes and diphenyls, respectively.
Solid State & Solution_Unit 1-2_Class-12_Board-JEE-NEETSupratim Das
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LAND USE LAND COVER AND NDVI OF MIRZAPUR DISTRICT, UPRAHUL
This Dissertation explores the particular circumstances of Mirzapur, a region located in the
core of India. Mirzapur, with its varied terrains and abundant biodiversity, offers an optimal
environment for investigating the changes in vegetation cover dynamics. Our study utilizes
advanced technologies such as GIS (Geographic Information Systems) and Remote sensing to
analyze the transformations that have taken place over the course of a decade.
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Remote Sensing and Geographic Information Systems
9
Changes in vegetation cover refer to variations in the distribution, composition, and overall
structure of plant communities across different temporal and spatial scales. These changes can
occur natural.
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A workshop hosted by the South African Journal of Science aimed at postgraduate students and early career researchers with little or no experience in writing and publishing journal articles.
Natural birth techniques - Mrs.Akanksha Trivedi Rama University
Structure Of Atom_STUDY MATERIALS_JEE-MAIN_AIPMT
1. Study Materials for BOARD-JEE-MAIN-AIPMT
Dalton’s theory of atom
According to this theory the Atom is considered to be hard, dense and smallest particle of matter,
which is indivisible, the atoms belonging to a particular element, is unique. The properties of
elements differ because of the uniqueness of the atoms belonging to particular elements. This
theory provides a satisfactory basis for the laws of chemical combination. The atom can neither
be created nor be destroyed i.e., it is indestructible.
It fails to explain why atoms of different kinds should differ in mass and
valency etc.The discovery of isotopes and isobars showed that atoms of same elements may have
different atomic masses (isotopes) and atoms of different kinds may have same atomic masses (isobars).
Atom is not the smallest indivisible particle but had a complex structure of its own and was made
up of still smaller particles like electrons, protons, neutrons etc. At present about 35 different
subatomic particles are known but the three particles namely electron, proton and neutron are
regarded as the fundamental particles.
The existence of electrons in atoms was first suggested, by J.J. Thomson, as a result of
experimental work on the conduction of electricity through gases at low pressures and at high
voltage, which produces cathode rays consisting of negatively charged particles, named as
electrons. The e/m ratio for cathode rays is fixed whose value is 8
1.76 10 C/ g× .
Properties & Characteristics of the three fundamental particles are:
Electron Proton Neutron
Symbol e or –1
e P n
Approximate relative mass 1/1836 1 1
Approximate relative charge –1 +1 No charge
Mass in kg –31
9.109 10× –27
1.673 10× –27
1.675 10×
Mass in amu –4
5.485 10× 1.007 1.008
Actual charge (coulomb) –19
1.602 10× –19
1.602 10× 0
Actual charge (e.s.u.) –10
4.8 10× –10
4.8 10× 0
The atomic mass unit (amu) is 1/12 of the mass of an individual atom of
12
6 C , i.e.,
1.660 10h – 27× kg. The neutron and proton have approximately equal masses of 1 amu and the
electron is about 1836 times lighter, its mass can sometimes be neglected as an approximation.
The electron and proton have equal, but opposite, electric charge while the neutron is not
charged.
Different moDels of atom
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Thomson's Model: Thomson assumed that an atom is a sphere of positive charges uniformly
distributed, with the electrons scattered as points throughout the sphere. This was known as
plum-pudding model at that time. However this idea was dropped due to the success of α-particle
scattering experiments studied by Rutherford and Mardson.
Rutherford’s Model: α -particle emitted by radioactive substance were shown to be
dipositive Helium ions (He )++
having a mass of 4 units and 2 units of positive charge.
Rutherford allowed a narrow beam of α -particles to fall on a very
thin gold foil of thickness of the order of –4
4 10× cm and
determined the subsequent path of these particles with the help of a
zinc sulphide fluorescent screen. The zinc sulphide screen gives off
a visible flash of light when struck by an α − particle, as ZnS has
the remarkable property of converting kinetic energy of particle
into visible light.
[For this experiment, Rutherford specifically used α − particles
because they are relatively heavy resulting in high momentum].
Experimental Observation
a) Majority of the α -particles pass straight through the gold strip with little or no deflection.
b) Some -particles are deflected from their path and diverge.
c) Very few -particles are deflected backwards through angles greater than 90°.
d) Some were even scattered in the opposite direction at an angle of 180° [Rutherford was
very much surprised by it and remarked that "It was as incredible as if you fired a 15 inch
shell at a piece of tissue paper and it came back and hit you"].
Conclusions
1. The fact that most of the α particles passed straight through the metal foil indicates the
most part of the atom is empty.
2. The fact that few α particles are deflected at large angles indicates the presence of a
heavy positively charged body i.e., for such large deflections to occur α -particles must
have come closer to or collided with a massive positively charged body, and he named it
nucleus.
3. The fact that one in 20,000 have deflected at 180° backwards indicates that volume
occupied by this heavy positively charged body is very small in comparison to total
volume of the atom.
Rutherford’s Atomic model: On the basis of the above observation, and having realized
that the rebounding α -particles had met something even more massive than themselves inside
the gold atom, Rutherford proposed an atomic model as follows.
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a) All the protons (+ve charge) and the neutrons (neutral charge) i.e. nearly the total mass of
an atom is present in a very small region at the centre of the atom. The atom's central core
is called nucleus.
b) The size of the nucleus is very small in comparison to the size of the atom. Diameter of
the nucleus is about –13
10 while the atom has a diameter of the order –8
10 of cm. So, the
size of atom is 5
10 times more than that of nucleus.
c) Most of the space outside the nucleus is empty.
d) The electrons, equal in number to the net nuclear positive charge, revolve around the
nucleus with high speed in various circular orbits.
e) The centrifugal force arising due to the high speed of an electron balances the columbic
force of attraction of the nucleus and the electron remains stable in its path. Thus
according to him atom consists of two parts (a) nucleus and (b) extra nuclear part.
Defects of Rutherford's atomic model
1. Position of electrons: The exact positions of the electrons from the nucleus are not
mentioned.
2. Stability of the atom: Neils Bohr pointed out that Rutherford's atom should be highly
unstable. According to the law of electro-dynamics, the electron should therefore,
continuously emit radiation and lose energy. As a result of this a moving electron will
come closer and closer to the nucleus and after passing through a spiral path, it should
ultimately fall into the nucleus.
absorption spectrum-atomic spectrum
If the atom gains energy the electron passes from a lower energy level to a higher energy level,
energy is absorbed that means a specific wave length is absorbed. Consequently, a dark line will
appear in the spectrum. This dark line constitutes the absorption spectrum.
For Hydrogen Atom: If an electric discharge is passed through hydrogen gas taken in a
discharge tube under low pressure, and the emitted radiation is analysed with the help of
spectrograph, it is found to consist of a series of sharp lines in the UV, visible and IR
regions. This series of lines is known as line or atomic spectrum of hydrogen. The lines in
the visible region can be directly seen on the photographic film.
Six Series of Spectrum:
Each line of the spectrum corresponds to a light of definite wavelength. The entire spectrum
consists of six series of lines each series, known after their discoverer as the Balmer, Paschen,
Lyman, Brackett, Pfund and Humphrey series. The wavelength of all these series can be
expressed by a single formula.
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2 2
1 2
1 1 1
R –
n n
= ν =
λ
ν = wave number
λ = wave length
R = Rydberg constant (109678 –1
cm )
1n and 2n have integral values as follows
Series 1n 2n Main spectral lines
Lyman 1 2, 3, 4, etc Ultra-vio
Balmer 2 3, 4, 5 etc Visible
Paschen 3 4, 5, 6 etc Infra-red
Brackett 4 5, 6, 7 etc Infra-red
Pfund 5 6, 7, 8, etc Infra-red
The pattern of lines in atomic spectrum is characteristic of hydrogen.
Emission Spectra:
a) Continuous spectra: When white light from any source such as sun or bulb is
analysed by passing through a prism, it splits up into seven different wide bands of colour
from violet to red (like rainbow). These colour also continuous that each of them merges
into the next. Hence the spectrum is called as continuous spectrum.
b) Line spectra: When an electric discharge is passed through a gas at low pressure light
is emitted. If this light is resolved by a spectroscope, it is found that some isolated
coloured lines are obtained on a photographic plate separated from each other by dark
spaces. This spectrum is called line spectrum. Each line in the spectrum corresponds to a
particular wavelength. Each element gives its own characteristic spectrum.
planck’s Quantum theory
When a black body is heated, it emits thermal radiations of different wavelengths or frequency.
To explain these radiations, Max Planck put forward a theory known as Planck’s quantum theory.
The main points of quantum theory are:
a) Substances radiate or absorb energy discontinuously in the form of small packets or
bundles of energy.
b) The smallest packet of energy is called quantum. In case of light the quantum is known as
photon.
c) The energy of a quantum is directly proportional to the frequency of the radiation. E ∝ ν
(or E = hν were v is the frequency of radiation and h is Planck’s constant having the value
–27
6.626 10× erg-sec or –34
6.626 10× J-sec.
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d) A body can radiate or absorb energy in whole number multiples of a quantum hv, 2hv,
3hv …. nhν where n is the positive integer. Nelis Bohr used this theory to explain the
structure of atom.
bohr’s atomic moDel
Bohr developed a model for hydrogen and hydrogen like atoms one-electron species (hydrogenic
species). He applied quantum theory in considering the energy of an electron bond to the nucleus.
Bohr’s Postulates:
An atom consists of a dense nucleus situated at the center with the electron revolving around it in
circular orbits without emitting any energy. The force of attraction between the nucleus and
an electron is equal to the centrifugal force of the moving electron.
Of the finite number of circular orbits possible around the nucleus, and electron can revolve only
in those orbits whose angular momentum (mvr) is an integral multiple of factor h/ 2π.
nh
mvr
2
=
π
where, m = mass of the electron, v = velocity of the electron
n = orbit number in which electron is present
r = radius of the orbit
As long as an electron is revolving in an orbit it neither loses nor gains energy. Hence these
orbits are called stationary states. Each stationary state is associated with a definite amount of
energy and it is also known as energy levels. The greater the distance of the energy level from the
nucleus, the more is the energy associated with it. The different energy levels are numbered as 1,
2, 3, 4, (from nucleus onwards) or K, L, M,N etc. Ordinarily an electron continues to move in a
particular stationary state without losing energy. Such a stable state of the atom is called as
ground state or normal state.
If energy is supplied to an electron, it may jump (excite) instantaneously from lower energy (say
1) to higher energy level (say 2, 3, 4, etc) by absorbing one quantum of energy. This new state of
electron is called as excited state. The quantum of energy absorbed is equal to the difference in
energies of the two concerned levels.
Since the excited state is less stable, atom will lose it’s energy and come back to the ground state.
Energy absorbed or released in an electron jump, ( E)∆ is given by
2 1E E – E hv∆ = =
Where 2E and 1E are the energies of the electron in the first and second energy levels, and v is
the frequency of radiation absorbed or emitted.
Important Note: If the energy supplied to hydrogen atom is less than 13.6 eV, it will accept or absorb
only those quanta which can take it to a certain higher energy level i.e., all those photons having
energy less than or more than a particular energy level will not be absorbed by hydrogen atom.
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But if energy supplied to hydrogen atom is more than 13.6 eV then all photons are absorbed and
excess energy appear as kinetic energy of emitted photo electron.
Bohr’s theory explains the followings:
a) The experimental value of radii and energies in hydrogen atom are in good agreement
with that calculated on the basis of Bohr’s theory.
b) Bohr’s concept of stationary state of electron explains the emission and absorption spectra
of hydrogen like atoms.
c) The experimental values of the spectral lines of the hydrogen spectrum are in close
agreement with the calculated by Bohr’s theory.
Limitations of Bohr’s Theory
a) It does not explain the spectra of atoms or ions having more than one electron.
b) Bohr’s atomic model failed to account for the effect of magnetic field (Zeeman effect) or
electric field (Stark effect) on the spectra of atoms or ions. It was observed that when the
source of a spectrum is placed in a strong magnetic or electric field, each spectral line
further splits into a number of lines. This observation could not be explained on the basis
of Bohr’s model.
c) de-Broglie suggested that electrons like light have dual character. It has particle and wave
character. Bohr treated the electron only as particle.
d) Another objection to Bohr’s theory came from Heisenberg’s Uncertainty Principle.
According to this principle “it is impossible to determine simultaneously the exact
position and momentum of a small moving particle like an electron”. The postulate of
Bohr, that electrons revolve in well defined orbits around the nucleus with well defined
velocities is thus not attainable.
Applications of Bohr’s theory:
a) Radius and Energy levels of hydrogen atom: Consider an electron of mass `m’
and charge `e’ revolving around a nucleus of charge Ze (where, Z = atomic number and e
is the charge of the proton) with a tangential velocity v.r is the radius of the orbit in which
electron is revolving.
By Coulomb’s Law, the electrostatic force of attraction between the moving electron and
nucleus is Coulombic force
2
2
KZe
r
=
0
1
K
4
=
πε
(where 0ε is permitivity of free space)
9 2 –2
K 9 10 Nm C= ×
In C.G.S. units, value of K = 1 dyne 2 –2
cm (esu)
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The centrifugal force acting on the electron is
2
mv
r
Since the electrostatic force balance the centrifugal force, for the stable electron orbit.
2 2
2
mv KZe
r r
= … (i)
(or)
2
2 KZe
v
mr
= … (ii)
According to Bohr’s postulate of angular momentum quantization, we have
nh
mvr
2
=
π
nh
v
2 mr
=
π
2 2
2
2 2 2
n h
v
4 m r
=
π
… (iii)
Equating (2) and (3)
2 2 2
2 2 2
KZe n h
mr 4 m r
=
π
Solving for r we get
2 2
2 2
n h
r
4 mKZe
=
π
where n = 1, 2, 3, … ∞
Hence only certain orbits whose radii are given by the above equation are available for
the electron. The greater the value of n, i.e., farther the level from the nucleus the greater
is the radius.
The radius of the smallest orbit (n = 1) for hydrogen atom (Z = 1) is 0r .
2 2 2 –34 2
–11
0 2 2 2 –31 –19 2 9
n h 1 (6.626 10 )
r 5.29 10 m
4 me K 4 (3.14) 9 10 (1.6 10 ) 9 10
× ×
= = = ×
π × × × × × × ×
0r 0.529Å=
Radius of nth orbit for an atom with atomic number Z is simply written as
2 2
n 0
n n
r r 0.529 Å
z z
= = ×
b) Energy level of Hydrogen atom: The total energy, E of the electron is the sum of
kinetic energy and potential energy. Kinetic energy of the electron = 21
2 mv
Potential energy
2
–KZe
r
=
Total energy
2
21
2
KZe
mv –
r
= … (4)
From equation (1) we know that
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2 2
2
mv KZe
r r
=
2
21
2
KZe
mv
2r
∴ =
Substituting this in equation (4)
Total energy (e)
2 2 2
KZe KZe KZe
– –
2r r 2r
= =
1
KE – PE, KE –TE
2
= =
Substituting for r, gives us
2 2 4 2
2 2
2 mZe e K
E
n h
π
= where n = 1, 2, 3, …
This expression shows that only certain energies are allowed to the electron. Since this
energy expression consists so many fundamental constant, we are giving you the
following simplified expressions.
2
–12
2
Z
E –21.8 10 erg
n
= × × per atom.
2
–19
2
Z
– 21.8 10
n
= × × J per atom = –
2
2
Z
13.6
n
× eV per atom
2
n 2
z
E –13.6
n
= eV per atom
(1eV = –23
3.83 10 Kcal)×
(1eV = –12
1.602 10× erg)
(1eV = –19
1.602 10× J)
2
2
Z
E –313.6
n
= × kcal/mole (1 cal = 4.18 J)
The energies are negative since the energy of the electron in the atom is less than the
energy of a free electron (i.e., the electron is at infinite distance from the nucleus) which
is taken as zero. The lowest energy level of the atom corresponds to n = 1, and as the
quantum number increases, E become less negative.
When n = ∞ , E = 0 which corresponds to an ionized atom i.e., the electron and nucleus
are infinitely separated.
–
H H e+
→ + (ionization).
c) Velocity of electron
We know that, mvr
nh nh
; v
2 2 mr
= =
π π
By substituting for r we are getting
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2
2 KZe
v
nh
π
=
Where excepting n and z all are constant, v =
8 Z
2.18 10 cm /sec.
n
×
Further application of Bohr’s work was made, to other one electron species (Hydrogenic
ion) such as He+
and 2
Li +
. In each case of this kind, Bohr’s prediction of the spectrum
was correct.
d) Explanation for hydrogen spectrum by Bohr’s theory: According to the Bohr’s
theory electron neither emits nor absorbs energy as long as it stays in a particular orbit.
However, when an atom is subjected to electric discharge or high temperature, and
electron in the atom may jump from the normal energy level, i.e., ground state to some
higher energy level i.e., exited state. Since the life time of the electron in excited state is
short, it returns to the ground state in one or more jumps.
During each jumps, energy is emitted in the form of a photon of light of definite
wavelength or frequency. The frequency of the photon of light thus emitted depends upon
the energy difference of the two energy levels concerned ( 1 2n , n ) and is given by
2 4 2
2 1 2 2 2
1 2
–2 mZ e K 1 1
hv E – E –
h n n
2
π
= =
2 2 4 2
3 2 2
1 2
2 mZ e K 1 1
v –
h n n
π
=
The frequencies of the spectral lines calculated with the help of above equation are found
to be in good agreement with the experimental values. Thus, Bohr’s theory elegantly
explains the line spectrum of hydrogen and hydrogenic species.
Bohr had calculated Rydberg constant from the above equation.
2 2 4 2
3 2 2
1 2
C 2 mZ e K 1 1
–
h n n
π
ν = =
λ
2 2 4 2
3 2 2
1 2
1 2 mZ e K 1 1
–
h c n n
π
= ν =
λ
where
2 4 2
–7 –1 –1
3
2 me K
1.097 10 m or109678 cm
h c
π
= ×
i.e. Rydberg constant (R)
2
2 2
1 2
1 1 1
RZ –
n n
∴ ν = =
λ
ν =wave number.
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Photoelectric effect
It was observed by Hertz and Lenard around 1880 that when a clean metallic surface is irradiated
by monochromatic light of proper frequency, electrons are emitted from it. This phenomenon of
ejection of the electrons from metal surface was called as Photoelectric Effect.
It was observed that if the frequency of incident radiation is below a certain minimum value
(threshold frequency), no emission takes place however high the intensity of light may be.
According to Einstein, when a quantum of light (photon) strikes a metal surface, it imparts its
energy to the electrons in the metal. In order for an electron to escape from the surface of the
metal, it must overcomes the attractive force of the positive ions in the metal. So a part of the
photon's energy is absorbed by the metal surface to release the electron, this is known as work
function of the surface and is denoted by φ . The remaining part of the energy of the photon goes
into the kinetic energy of the electron emitted. If E is the energy of the photon, KE is the kinetic
energy of the electron and φ be the work function of the metal then we have;
0 ih and E hφ = ν = ν
⇒ i 0 0KE E – KE h – h h( – )= φ ⇒ = ν ν = ν ν
Also, if m be the mass and v be the velocity of the electron ejected then
21
2 0KE mv h( – )= = ν ν .
Important Note: The electromagnetic Radiation (or wave) now emerges as an entity which shows dual
nature i.e., sometimes as Wave and sometimes as Particle (quantum aspect).
Quantum mechanical model of atom
The atomic model which is based on the particle and wave nature of the electron is known as
wave or quantum mechanical model of the atom. This was developed by Schrodinger in 1926.
This model describes the electron as a three dimensional wave in the electronic field of positively
charged nucleus. Schrodinger derived an equation which describes wave motion of positively
charged nucleus. Schrodinger derived an equation which describes wave motion of an electron.
The differential equation is
2 2 2 2
2 2 2 2
d d d 8 m
(E – V) 0
dx dy dz h
ψ ψ ψ π
+ + + ψ =
where x, y, z are certain coordinates of the electron, m = mass of the electron, E = total energy of
the electron. V = potential energy of the electron; h = Planck's constant and ψ (psi)= wave
function of the electron.
Wave Function (ψ):
The wave function may be regarded as the amplitude function expressed in terms of coordinates
x, y and z. The wave function may have positive or negative values depending upon the value of
coordinates. The main aim of Schrodinger equation is to give solution for probability approach.
When the equation is solved, it is observed that for some regions of space the value of ψ is
negative. But the probability must be always positive and cannot be negative, it is thus, proper to
use 2
ψ in favour of ψ .
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Probability Density (ψ2
):
ψ2
is a probability factor. It describes the probability of finding an electron within a small space.
The space in which there is maximum probability of finding an electron is termed as orbital. The
important point of the solution of the wave equation is that it provides a set of numbers called
quantum numbers which describe energies of the electron in atoms, information about the shapes
and orientations of the most probable distribution of electrons around nucleus.
Quantum numbers
To explain the Quantum Numbers, we must known the meaning of some terms clearly so as to
avoid any confusion.
Energy Level & Sub-Energy Level:
The non-radiating energy paths around the nucleus are called as Energy Levels of Shells. These
are specified by numbers having values 1, 2, 3, 4, ... or K, L, M, N, ... in order of increasing
energies. The energy of a particular energy level is fixed.
The phenomenon of splitting of spectral lines in electric and magnetic fields reveals that there
must be extra energy levels within a definite energy level. These were called as Sub-Energy
Levels or Sub-Shells. There are four types of sub-shells namely; s, p, d, f.
First energy level (K or ) has one sub-shell designated as 1s, the second energy level (L or 2) has
two sub-shell as 2s & 2p, the third energy level (M or 3) has three sub shell as 3s, 3p and 3d, and
the fourth energy level (N or 4) has four sub-shells as 4s, 4p, 4d and 4f. The energy of sub-shell
increases roughly in the order: s < p < d <f.
Concept of Orbital:
Each sub-energy level (sub-shell) is composed of one or more orbitals. These orbitals belonging
to a particular sub-shell have equal energies and are called as degenerate orbitals. s-sub-shell has
one orbital, p has three orbitals, d have five orbitals and f has seven orbitals.
To describe or to characterize the electrons around the nucleus in an atom, a set of four numbers
is used, called as Quantum Numbers. These are specified such that the states available to the
electrons should follows the laws of quantum mechanics or wave mechanics.
Principal Quantum Number: (n): This quantum number represents the main energy
levels (principal energy levels) designated as n = 1, 2, 3, ... or the corresponding shells are named
as K, L, M, N, ... respectively. It gives an idea of position and energy of an electron. The energy
level n = 1 corresponds to minimum energy and subsequently n = 2, 3, 4, ..., are arranged in order
of increasing energy.
Higher is the value of n, greater is its distance from the nucleus, greater is its size and also greater
is its energy.
It also gives the total electrons that may be accommodated in each shell, the capacity of each
shell is given by the formula 2
2n , where n : principal quantum number.
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Azimuthal Quantum Number: (l): This number determines the energy associated with
the angular momentum of the electron about the nucleus. It is also called as the angular
momentum quantum number. It accounts for the appearance of groups of closely packed spectral
lines in electric field.
It can assume all integral values from 0 to n–1. The possible values of l are :
0, 1, 2, 3, ..., n–1.
Each value of l describes a particular sub-shell in the main energy level and determines the shape
of the electron cloud.
When n = 1, l = 0, i.e., its energy level contains one sub-shell which is called as a s-sub-shell. So
for l = 0, the corresponding sub-shell is a s-sub-shell. Similarly when l = 1, 2, 3, the sub-shells
are called p, d, f sub-shells respectively.
As you known for n = 1, l = 0, there is only one sub-shell. It is represented by 1s. Now for n = 2,
l can take two values (the total number of values taken by l is equal to the value of n in a
particular energy level). The possible values of l are 0, 1. The two sub-shell representing the IInd
energy level are 2s, 2p. In the same manner, for n = 3, three sub-shells are designated as 3s, 3p,
3d corresponding to l = 0, 1, 2, and for n = 4, four sub-shells are designated as 4s, 4p, 4d, 4f
corresponding to l = 0, 1, 2, 3.
The orbital Angular momentum of electron =
h
( 1)
2
+
π
l l .
Note that its value does not depend upon value of n.
Magnetic Quantum Number (m): An electron with angular momentum can be thought as
an electric current circulating in a loop. A magnetic field due to this current is observed. This
induced magnetism is determined by the magnetic quantum number. Under the influence of
magnetic field, the electrons in a given sub-energy level prefer to orient themselves in certain
specific regions in space around the nucleus. The number of possible orientations for a sub-
energy level is determined by possible values of m corresponds to the number of orbitals in a
given sub-energy level).
m can have any integral values between –l to +l including 0, i.e., m = –l, 0 +, l, …, 0, 1, 2, 3, 4, .
. ., l–1 + l. We can say that a total of (2l + 1) values of m are there for a given value of l– 2, –1, 0,
1, 2, 3.
In s sub-shell there is only one orbital [l = 0, ⇒ m = (2l +1) = 1].
In p sub-shell there are three orbitals corresponding to three values of m : –1, 0 +1. [l = 1 ⇒ m =
(2l +1) = 3]. These three orbitals are represented as x y zp , p , p along X, Y, Z axes perpendicular
to each other.
In d sub-shell, there are five orbitals corresponding to –2, –1, 0 +1, +2, [l = 2 ⇒ m = (
2 2 1) 5]× + = . These five orbitals are represented as 2 2
2
xy zx x –y
d ,d ,d ,dz .
In f sub-shell there are seven orbitals corresponding to –3, –2, –1, 0, +1, +2, +3 [l = 3 ⇒ m =
(2 3 1) 7]× + = .
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Spin quantum Number (s): When an electron rotates around a nucleus it also spins about
its axis. If the spin is clockwise, its spin quantum number is +1/2 and is represented as ↑ . If the
spin is anti-clockwise, its value is –1/2 and is represented as ↓ . If the value of s is +1/2, then by
convention, we take that electron as the first electron in that orbital and if the value of s is –1/2, it
is taken as second electron.
Filling up orbital - auFbau principle
Aufbau is a German word meaning `building up'. This gives us a sequence in which various sub-
shells are filled up depending on the relative order of the energy of the subs-hells. The sub-shell
with minimum energy is filled up first and when this obtains maximum quota of electrons, then
the next sub-shell of higher energy starts filling.
1s
2s
3s
4s
5s
6s
7s
2p
3p
4p
5p
6p
3d
4d
5d
4f
1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p.
Sub-shell N L (n+l)
1s 1 0 1
2s 2 0 2
2p 2 1 3 Lowest value of n
3s 3 0 3
3p 3 1 4 Lowest value of n
4s 4 0 4
3d 3 2 5 Lowest value of n
4p 4 1 5
5s 5 0 5
4d 4 2 6 Lowest value of n
5p 5 1 6
6s 6 0 6
4f 4 3 7 Lowest value of n
5d 5 2 7
6p 6 1 7
7s 7 0 7
5f 5 3 8 Lowest value of n
6d 6 2 8
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7p 7 1 8
Exceptions to Aufbau Principle: In some cases it is seen that actual electronic arrangement is
slightly different from arrangement given by Aufbau principle. A simple reason behind this is that half-
filled and full-filled sub-shell have got extra stability.
2 2 6 2 6 4 2
Cr(24) 1s , 2s 2p , 3s 3p 3d , 4s→ (wrong)
23 2 6 2 6 5 1
1s ,2s 2p ,3s 3p 3d ,4s→ (right)
2 2 6 2 6 9 2
Cu(29) 1s ,2s 2p , 3s 3p 3d ,4s→ (wrong)
2 2 6 2 6 10 1
1s , 2s 2p , 3s 3p 3d ,4s→ (right)
Similarly the following elements have slightly different configurations than expected.
4 1
Nb [Kr] 4d 5s→
4 1
Mo [Kr] 4d 5s→
7 1
Ru [Kr] 4d 5s→
8 1
Rh [Kr] 4d 5s→
10 0
Pd [Kr] 4d 5s→
10 1
Ag [Kr] 4d 5s→
14 9 1
Pt [Xe] 4f 5d 6s→
14 10 1
Au [Xe] 4f 5d 6s→
Shapes of Atomic Orbitals
i) S-orbital: An electron in considered to be immersed out in the form of a cloud. The shape
of the cloud is the shape of the orbital. The cloud is not uniform but denser in the region
where the probability of finding the electron in maximum.
The orbital with the lowest energy is the 1s orbital. It is a sphere with its center of the
nucleus of the atom. The s-orbital is said to spherically symmetrical about the nucleus, so
that the electronic charge is not concentrated in any particular direction. 2s orbital is also
spherically symmetrical about the nucleus, but it is larger than (i.e., away from) the 1s
orbit.
nucleus
Z radial node
2s1s
x
y
ii) p-orbitals: There are three p-orbitals: x y zp ,p and p . they are dumb-bell shaped, the two
levels being separated by; a nodal plane, i.e., a plane where there is no likely hood of
finding the electron. The p-orbitals have a marked direction character, depending as
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whether x y zp ,p and p orbital is being considered. The p-orbitals consist of two lobes with
the atomic nucleus lying between them. The axis of each p-orbital is perpendicular to the
other two. The x y zp ,p and p orbitals are equivalent except for their directional property.
They have same energy; orbitals having the same energy are said to be degenerated.
x
y
z
z
y
px
x
pz
z
y
x
py
iii) d-orbitals: There are five d-orbitals. The shapes of four d-orbitals resemble four leaf
cloves. The fifth d-orbital loops different. the shapes of these orbitals are given below.
y
dxy
x
z
z
dyz
y
dzx
zy
x
2
z
d22 yx
d
−
x
Wave particle Dual character : De-broglie
In case of light some phenomenon like diffraction and interference can be explained on the basis
of its wave character. However, the certain other phenomenon such as black body radiation and
photoelectric effect can be explained only on the basis of its particles nature. Thus, light is said to
have a dual character. Such studies on light were made by Einstein in 1905.
Louis de-Broglie, in 1942 extended the ideal of photons to material particles such as electron and
he proposed that matter also has a dual character-as wave and as particle.
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Derivation of de-Broglie equation: The wavelength of the wave associated with any
material particle was calculated by analogy with photon. In case of photon, if it is assumed
to have wave character, its energy is given by
E = hν … (i) (according to the Planck’s quantum theory)
where ν is the frequency of the wave and `h’ is Planck’s constant
If the photon is supposed to have particle character, its energy is given by
2
E mc= … (ii) according to Einstein’s equation)
where `m’ is the mass of photon, `c’ is the velocity of light.
By equating (i) and (ii)
2
h mcν =
But c/ν = λ
2c
h mc=
λ
(or) h / mcλ =
The above equation is applicable to material particle if the mass and velocity of photon is
replaced by the mass and velocity of material particle. Thus for any material particle like
electron.
h
h / mv (or)
p
λ = λ =
where mv = p is the momentum of the particle.
Heisenberg’s Uncertainty PrinciPle
All moving objects that we see around us e.g., a car, a ball thrown in the air etc, move along
definite paths. Hence their position and velocity can be measured accurately at any instant of
time. Is it possible for subatomic particle also?
As a consequence of dual nature of matter. Heisenberg, in 1927 gave a principle about the
uncertainties in simultaneous measurement of position and momentum (mass × velocity) of
small particles. This principle states.
It is impossible to measure simultaneously the position and momentum of a small microscopic
moving particle with absolute accuracy or certainty i.e., if an attempt is made to measure any one
of these two quantities with higher accuracy, the other becomes less accurate.
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The product of the uncertainty in position ( x)∆ and the uncertainty in the momentum
( p m. v∆ = ∆ where m is the mass of the particle and v∆ is the uncertainty in velocity) is equal to
or greater than h / 4π where h is the Planck’s constant.
Thus, the mathematical expression for the Heisenberg’s uncertainty principle is simply written as
h
x. p h / 4 or E t
4
∆ ∆ ≥ π ∆ ×∆ ≥
π
Electronic Configuration:
Quantum numbers can now characterize the electrons in an atom. To describe the arrangements
and distribution of electrons for different elements, following rules an selective principles are
used. The distributions of electrons in an atom is known as the electronic configuration of that
element.
Aufbau Principle: An atom in its lowest state of energy is said to be in ground state. The
ground state is the most stable in an atom. According to Aufbau principle.
“Electrons are added progressively to the various orbitals in their order of increasing energy
starting with the orbital of lowest energy”.
The order of increasing energy may be summed up as follows
1s, 2s, 2p, 3s, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, …
As a working rule, a new electron enters an empty orbital for which the value of (n + l) is
minimum. If the value (n + l) is same for two or more orbitals, the new electron enters an orbital
having lower value of n.
Pauli’s Exclusion Principle: According to this principle
No two electrons in an atom can have the same set of all the quantum numbers or one can say
that no two electrons can have the same quantised states.
Consider an electronic arrangement in 1st
energy level (n = 1). For n = 1. l= 0, and m = 0. Now s
can have to values corresponding to each value of m i.e. s = +1.2, –1/2 (n, 1, possible designation
of an electron in a state with n = 1 is 1, 0, 0, +1/2 and 1, 0, 0, –1/2 (n, l, m, s) i.e., two quantised
states. This implies that an orbital can accommodate (for n = 1, m = 0, ⇒ one orbital) maximum
of two electrons having opposite spins.
The maximum number of electrons in the different subshells = 2 (2l +1).
s-sub-shell = 2, p-sub-shell = 6, d-sub-shell = 10 and f-sub-shell = 14.
Hund’s Rule of Multiplicity
According to this rule: “Electrons never pair until no available empty degenerate orbitals are left
to him.”
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This means an electron always occupies a vacant orbital in the same sub-shell (degenerate
orbital) and pairing starts only when all of the degenerate orbitals are filled up. This means that
the pairing starts with 2nd
electron in a sub-shell, 4th
electron in p-sub-shell, 6th
electron in d-sub-
shell and 8th
electron in f-sub-shell.
By doing this, the electrons stay as far away from each other as possible. This is highly
reasonable if we consider the electron-electron repulsion. Hence electrons obey Hund’s rule as it
results in lower energy state and hence more stability.
Extra Stability of Half and fully Filled Orbitals: A particularly stable system is
obtained when a set of equivalent orbitals (degenerate orbitals) is either fully filled or half filled,
i.e., each containing one or a pair of electrons. This effect is more dominant in d and f sub-shells.
This means three or six electrons in p-sub-shell, five or ten electrons in d-sub-shell, and seven or
fourteen electrons in f-sub-shell forms a stable arrangement. Note this effect when filling of
electrons takes place in d sub-shells (for atomic number Z = 24, 25, and 29, 30).
In the following table you should analyse how to employ the above rules to write electronic
configuration of various elements.
Electronic configuration of an element is represented by the notation x
nl .
x : number of electrons present in an orbital
l : denotes the sub-shell
n : principal quantum number.
node and nodal Plane
Node is defined as a region where the probability of finding an electron is zero.
The planes passing through the angular nodal points are called nodal planes.
Nodes
No. of radial or spherical nodes = n – l – 1
No. of angular nodes = l, Total no. of nodes = n–1
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