This document provides a review of key concepts in atomic theory and chemistry for an upcoming unit test. It begins with a history of atomic theory models from Dalton to Heisenberg. It then covers quantum mechanics, the four quantum numbers, Bohr's theory, Pauli's exclusion principle, and electron configuration diagrams. Additional sections explain concepts like the VSEPR theory, molecular polarity, intermolecular forces, Lewis structures, and valence bond theory. The document concludes with information about atomic spectra experiments and different types of crystal structures.
1) Atoms are the building blocks of matter and are composed of a nucleus containing protons and neutrons surrounded by electrons that orbit in shells.
2) There are different subatomic particles that make up an atom including protons, neutrons, and electrons. Protons and neutrons are in the nucleus while electrons orbit in shells around the nucleus.
3) Isotopes are atoms of the same element that have differing numbers of neutrons. For example, hydrogen has isotopes of deuterium and tritium that have extra neutrons compared to common hydrogen.
Applied Chemistry, atomic and molecular structure, part 1, by Shiraz mahbob PhDMaqsoodAhmadKhan5
applied chemistry lecture and slide,
Applied Chemistry, atomic and molecular structure, part 1, by Shiraz mahbob PhD, lecturer in chemistry in pakistan institute of engineering and applied sciences
1. Atoms are composed of protons, neutrons, and electrons. Protons and neutrons are located in the nucleus, while electrons orbit the nucleus.
2. Rutherford's gold foil experiment showed that atoms are mostly empty space, with a small, dense nucleus at the center containing positively charged protons and uncharged neutrons.
3. Bohr's model of the atom explained the stability of atoms by proposing that electrons orbit the nucleus in fixed, quantized energy levels.
This document discusses atomic structure and the discovery of subatomic particles. It describes J.J. Thomson's discovery of the electron in cathode rays in 1897. The discovery of the proton in anode rays by Goldstein in 1886 is also discussed. The discovery of the neutron by Chadwick in 1932 when bombarding beryllium with alpha particles is summarized. Rutherford's gold foil experiment in 1911 is briefly described, which provided evidence for the nuclear model of the atom with electrons orbiting a small, dense nucleus. Bohr's model improved upon Rutherford's by introducing electron energy levels and orbits. Key concepts like electron configuration, atomic number, mass number, and orbital shapes are defined.
BE UNIT-1 basic electronics unit one.pptxharisbs369
1. The document discusses the atomic structure of matter, which is made up of protons, electrons, and neutrons. Atoms contain protons and neutrons in their nucleus, surrounded by electrons.
2. Atoms of different elements have different atomic structures because they contain different numbers of protons and electrons. Neutral atoms have equal numbers of protons and electrons, but atoms can gain or lose electrons to become ions.
3. The document then discusses subatomic particles like protons, neutrons, and electrons in more detail, including their relative masses and charges. It also discusses isotopes and how they have the same number of protons but different numbers of neutrons.
Secondary Education
Chemistry
Chapter 1
Lesson 1
if you have any question don't hesitate to contact me
join the facebook group
http://www.facebook.com/#!/group.php?gid=17663120872&v=info
Best of luck
Mr.Ehab Mohamed
module 1 electronic structure of matter.pptxMaryroseBudhi1
Module 1: Electronic Structure of Matter
Objectives: Know atom and its sub - particles
determine the characteristics colors that metal salts emit
what is atom?
atom is the basic unit of chemical element
it composes three subatomic particle
proton with a positively electric charge
electron with a negatively electric charge
neutron no electric charge
What minerals produce the color in fireworks?
Mineral elements provide color in fireworks. Barium produces bright greens; strontium yields deep reds;' copper produces blues/ and sodium yields yellow. other colors can be made by mixing elements; strontium and sodium produce brilliant orange; titanium, zirconium, and magnesium alloys make silvery white; copper and strontium make lavender. gold sparks are produced by iron fillings and small pieces of charcoal. bright flashes and loud bangs come from aluminum powder.
This document discusses the structure of the atom. It begins by describing Bohr's model of the atom and its limitations. It then introduces shells and subshells, as well as quantum numbers and the shapes of atomic orbitals. Rules for filling electrons into orbitals, such as the Aufbau principle and Pauli exclusion principle, are also covered. The document discusses atomic spectra, photoelectric effect, and the dual wave-particle nature of light and matter. It provides an overview of concepts like de Broglie wavelength, Heisenberg uncertainty principle, and atomic electron configuration.
1) Atoms are the building blocks of matter and are composed of a nucleus containing protons and neutrons surrounded by electrons that orbit in shells.
2) There are different subatomic particles that make up an atom including protons, neutrons, and electrons. Protons and neutrons are in the nucleus while electrons orbit in shells around the nucleus.
3) Isotopes are atoms of the same element that have differing numbers of neutrons. For example, hydrogen has isotopes of deuterium and tritium that have extra neutrons compared to common hydrogen.
Applied Chemistry, atomic and molecular structure, part 1, by Shiraz mahbob PhDMaqsoodAhmadKhan5
applied chemistry lecture and slide,
Applied Chemistry, atomic and molecular structure, part 1, by Shiraz mahbob PhD, lecturer in chemistry in pakistan institute of engineering and applied sciences
1. Atoms are composed of protons, neutrons, and electrons. Protons and neutrons are located in the nucleus, while electrons orbit the nucleus.
2. Rutherford's gold foil experiment showed that atoms are mostly empty space, with a small, dense nucleus at the center containing positively charged protons and uncharged neutrons.
3. Bohr's model of the atom explained the stability of atoms by proposing that electrons orbit the nucleus in fixed, quantized energy levels.
This document discusses atomic structure and the discovery of subatomic particles. It describes J.J. Thomson's discovery of the electron in cathode rays in 1897. The discovery of the proton in anode rays by Goldstein in 1886 is also discussed. The discovery of the neutron by Chadwick in 1932 when bombarding beryllium with alpha particles is summarized. Rutherford's gold foil experiment in 1911 is briefly described, which provided evidence for the nuclear model of the atom with electrons orbiting a small, dense nucleus. Bohr's model improved upon Rutherford's by introducing electron energy levels and orbits. Key concepts like electron configuration, atomic number, mass number, and orbital shapes are defined.
BE UNIT-1 basic electronics unit one.pptxharisbs369
1. The document discusses the atomic structure of matter, which is made up of protons, electrons, and neutrons. Atoms contain protons and neutrons in their nucleus, surrounded by electrons.
2. Atoms of different elements have different atomic structures because they contain different numbers of protons and electrons. Neutral atoms have equal numbers of protons and electrons, but atoms can gain or lose electrons to become ions.
3. The document then discusses subatomic particles like protons, neutrons, and electrons in more detail, including their relative masses and charges. It also discusses isotopes and how they have the same number of protons but different numbers of neutrons.
Secondary Education
Chemistry
Chapter 1
Lesson 1
if you have any question don't hesitate to contact me
join the facebook group
http://www.facebook.com/#!/group.php?gid=17663120872&v=info
Best of luck
Mr.Ehab Mohamed
module 1 electronic structure of matter.pptxMaryroseBudhi1
Module 1: Electronic Structure of Matter
Objectives: Know atom and its sub - particles
determine the characteristics colors that metal salts emit
what is atom?
atom is the basic unit of chemical element
it composes three subatomic particle
proton with a positively electric charge
electron with a negatively electric charge
neutron no electric charge
What minerals produce the color in fireworks?
Mineral elements provide color in fireworks. Barium produces bright greens; strontium yields deep reds;' copper produces blues/ and sodium yields yellow. other colors can be made by mixing elements; strontium and sodium produce brilliant orange; titanium, zirconium, and magnesium alloys make silvery white; copper and strontium make lavender. gold sparks are produced by iron fillings and small pieces of charcoal. bright flashes and loud bangs come from aluminum powder.
This document discusses the structure of the atom. It begins by describing Bohr's model of the atom and its limitations. It then introduces shells and subshells, as well as quantum numbers and the shapes of atomic orbitals. Rules for filling electrons into orbitals, such as the Aufbau principle and Pauli exclusion principle, are also covered. The document discusses atomic spectra, photoelectric effect, and the dual wave-particle nature of light and matter. It provides an overview of concepts like de Broglie wavelength, Heisenberg uncertainty principle, and atomic electron configuration.
The document discusses the evolution of atomic models over time from Dalton's model to the current quantum mechanical model. It summarizes key developments including the plum pudding model, Bohr's model of electrons in orbits, and how the Schrodinger equation led to the quantum mechanical model where electrons occupy distinct energy levels and orbitals. The modern model describes electron probability clouds rather than set orbits and accounts for properties like electron spin and the Pauli exclusion principle.
The document discusses the historical development of atomic models from Dalton to Bohr and beyond. It introduces John Dalton's early model of atoms as indivisible particles with no internal structure in 1863. Later models incorporated the discoveries of the electron by J.J. Thomson in 1897 and the nuclear structure of atoms by Ernest Rutherford in 1911. Niels Bohr's 1913 model proposed that electrons orbit the nucleus in fixed, quantized energy levels. This laid the foundations for understanding atomic emission spectra and the quantum mechanical model that later replaced Bohr's model.
1) The document discusses the electronic structure of atoms, including the quantum mechanical model of the atom and how it explains experimental observations.
2) Key aspects covered include the wave-particle duality of electrons and light, the development of quantum numbers to describe electron orbitals and energies, and how the organization of electrons in atoms is reflected in the periodic table.
3) The document also notes some anomalies that arise when s and d orbitals are partially filled due to their similar energies.
The document summarizes Dalton's atomic theory and provides information about atomic structure and subatomic particles. It discusses Dalton's four main postulates, including that atoms are indivisible and atoms of different elements combine in whole number ratios. The document also outlines the discoveries of key subatomic particles like electrons, protons, and neutrons by scientists such as Thomson, Rutherford, and Chadwick. It describes Bohr's model of the atom and introduces concepts like orbitals, electron configuration, and quantum numbers.
This document provides an overview of atomic structure, bonding, and electron distribution. It begins by defining the basic subatomic particles that make up atoms. It then discusses several historical atomic models including Thomson's plum pudding model, Rutherford's nuclear model, and Bohr's early quantum model. The document introduces concepts like electron orbitals and quantum numbers. It also covers bonding theories such as ionic and covalent bonding as well as localized and delocalized bonding. Hybridization of atomic orbitals is discussed through examples like sp, sp2, and sp3 hybridization. The summary concludes with an introduction to molecular orbital theory.
The document provides information on the structure of atoms, including key experiments and models that helped reveal the internal structure of atoms. It discusses J.J. Thomson's cathode ray experiment that discovered electrons, Rutherford's alpha particle scattering experiment that showed atoms have a small, dense nucleus, and Bohr's model of electron orbits around the nucleus. It also covers topics like isotopes, mass number, atomic number, electron configuration, and valency.
ELECTRON-theory ppt industrials arts part2GalangRoxanne
Electron theory aims to explain the structure and properties of matter through its electronic structure. It states that all matter is comprised of molecules made up of atoms, which contain protons, neutrons, and electrons. Electrons play a key role in electricity as their movement constitutes electric current. Within atoms, electrons exist in specific energy levels or shells around the nucleus, and their behavior is described by quantum mechanics.
The document discusses atomic structure and bonding. It describes the structure of atoms including protons, neutrons, and electrons. It explains how atomic number determines the element and how isotopes have the same number of protons but different neutrons. Electron configuration and quantum numbers are also summarized. The three main types of bonds - ionic, covalent, and metallic - are introduced along with how they influence material properties.
The document discusses the development of atomic models from Dalton to Bohr and beyond. It describes Rutherford's discovery of the nucleus and Bohr's model of electrons in fixed orbits around the nucleus. Later, the quantum mechanical model was developed, restricting electrons to specific energy levels rather than exact orbits. This modern model determines the probability of finding electrons in different locations around the nucleus.
MODULE 1 ELETCRONIC STRUCTURE OF MATTER.pptxKathelynRuiz1
- The document discusses the history and development of atomic models from Dalton to Bohr. It describes key discoveries and experiments.
- Rutherford's gold foil experiment showed that the atom is mostly empty space, with a small, dense nucleus at the center containing positive charge. This led him to propose the nuclear model of the atom.
- Bohr incorporated Rutherford's nuclear model into his planetary model of the atom. He proposed that electrons orbit the nucleus in defined energy levels. This model had limitations when describing multi-electron atoms.
- The quantum mechanical model was developed based on discoveries that electrons can be described as waves using quantum physics principles like wave-particle duality and the uncertainty principle. This model uses atomic orbit
Atomic structure is composed of protons, neutrons, and electrons. Rutherford's gold foil experiment showed that atoms are mostly empty space with a small, dense nucleus at the center containing protons and neutrons. Electrons orbit the nucleus in specific energy levels. The Bohr model explained electrons in hydrogen having discrete energy levels. Atoms bond via ionic bonds between metals and nonmetals by transferring electrons, covalent bonds between nonmetals by sharing electrons, and metallic bonds between metals by delocalized electrons.
CBSE Class 9 Science Chapter 4- structure of atomAarthiSam
This document discusses the structure of atoms and the development of atomic models over time. It covers John Dalton's early atomic theory that atoms are indivisible and different for each element. J.J. Thomson's plum pudding model described atoms as electrically neutral spheres. Rutherford determined atoms have a small, dense nucleus through alpha particle scattering experiments. Niels Bohr incorporated quantum theory, proposing electrons orbit in discrete energy levels. The document also discusses atomic number, mass number, isotopes, electron configuration, and valency.
The document discusses the structure of atoms. It describes how early scientists like Dalton, Thomson, Rutherford, and Bohr contributed to developing models of the atom through experiments. Thomson's "plum pudding" model depicted electrons distributed uniformly in a positively charged sphere, but it could not explain atomic stability. Rutherford's gold foil experiment showed that the positive charge and mass of atoms are concentrated in a tiny nucleus, leading to his nuclear model. Bohr's model incorporated discrete electron orbits to explain atomic stability. The document also discusses subatomic particles like protons, neutrons, isotopes, isobars, and how electrons are arranged in shells according to their quantum numbers.
This document discusses Rutherford's atomic model and Bohr's model of the atom. It provides details of Rutherford's alpha particle scattering experiment which showed that atoms have a small, dense nucleus. This led Rutherford to propose a planetary model of the atom with electrons orbiting the nucleus. The document then discusses limitations of Rutherford's model and how Bohr proposed quantized electron orbits to explain atomic stability. It provides Bohr's key postulates and formulas for the hydrogen atom spectrum and energy levels.
The Fundamentals of Chemistry is an introduction to the Periodic Table, stoichiometry, chemical states, chemical equilibria, acid & base, oxidation & reduction reactions, chemical kinetics, inorganic nomenclature, and chemical bonding.
Atomic Structure Powerpoint Presentation by Computer CareersYaman Singhania
Powerpoint Presentation on Atomic Structure by Computer Careers.What is an Atom?ATOMIC STRUCTURE,There are two ways to represent the atomic structure of n element or compound,DOT & CROSS DIAGRAMS and many more ....
This document provides information about atomic structure. It discusses the basic parts of an atom including protons, neutrons, and electrons. Early atomic models proposed by Rutherford and Bohr are described, noting their limitations in explaining experimental observations. The modern quantum mechanical model represents electrons using quantum numbers and wave functions or "fuzzy clouds" to describe atomic orbitals. Electrons occupy different energy levels and sublevels based on their quantum numbers.
4.1 The Atomic Models of Thomson and Rutherford
4.2 Rutherford Scattering
4.3 The Classic Atomic Model
4.4 The Bohr Model of the Hydrogen Atom
4.5 Successes and Failures of the Bohr Model
4.6 Characteristic X-Ray Spectra and Atomic Number
4.7 Atomic Excitation by Electrons
This document discusses the electronic structure of atoms. It begins by reviewing early atomic models proposed by Thomson, Rutherford, and Chadwick that included a dense nucleus surrounded by electrons. The document then discusses how quantum mechanics provides a better model of electronic structure through the use of orbitals and quantum numbers to describe allowed electron configurations. Key points covered include the wave-particle duality of electrons, Schrodinger's wave equation describing orbital shape and orientation, and the four quantum numbers (n, l, ml, ms) that provide unique descriptions of electron states.
The document discusses the evolution of atomic models over time from Dalton's model to the current quantum mechanical model. It summarizes key developments including the plum pudding model, Bohr's model of electrons in orbits, and how the Schrodinger equation led to the quantum mechanical model where electrons occupy distinct energy levels and orbitals. The modern model describes electron probability clouds rather than set orbits and accounts for properties like electron spin and the Pauli exclusion principle.
The document discusses the historical development of atomic models from Dalton to Bohr and beyond. It introduces John Dalton's early model of atoms as indivisible particles with no internal structure in 1863. Later models incorporated the discoveries of the electron by J.J. Thomson in 1897 and the nuclear structure of atoms by Ernest Rutherford in 1911. Niels Bohr's 1913 model proposed that electrons orbit the nucleus in fixed, quantized energy levels. This laid the foundations for understanding atomic emission spectra and the quantum mechanical model that later replaced Bohr's model.
1) The document discusses the electronic structure of atoms, including the quantum mechanical model of the atom and how it explains experimental observations.
2) Key aspects covered include the wave-particle duality of electrons and light, the development of quantum numbers to describe electron orbitals and energies, and how the organization of electrons in atoms is reflected in the periodic table.
3) The document also notes some anomalies that arise when s and d orbitals are partially filled due to their similar energies.
The document summarizes Dalton's atomic theory and provides information about atomic structure and subatomic particles. It discusses Dalton's four main postulates, including that atoms are indivisible and atoms of different elements combine in whole number ratios. The document also outlines the discoveries of key subatomic particles like electrons, protons, and neutrons by scientists such as Thomson, Rutherford, and Chadwick. It describes Bohr's model of the atom and introduces concepts like orbitals, electron configuration, and quantum numbers.
This document provides an overview of atomic structure, bonding, and electron distribution. It begins by defining the basic subatomic particles that make up atoms. It then discusses several historical atomic models including Thomson's plum pudding model, Rutherford's nuclear model, and Bohr's early quantum model. The document introduces concepts like electron orbitals and quantum numbers. It also covers bonding theories such as ionic and covalent bonding as well as localized and delocalized bonding. Hybridization of atomic orbitals is discussed through examples like sp, sp2, and sp3 hybridization. The summary concludes with an introduction to molecular orbital theory.
The document provides information on the structure of atoms, including key experiments and models that helped reveal the internal structure of atoms. It discusses J.J. Thomson's cathode ray experiment that discovered electrons, Rutherford's alpha particle scattering experiment that showed atoms have a small, dense nucleus, and Bohr's model of electron orbits around the nucleus. It also covers topics like isotopes, mass number, atomic number, electron configuration, and valency.
ELECTRON-theory ppt industrials arts part2GalangRoxanne
Electron theory aims to explain the structure and properties of matter through its electronic structure. It states that all matter is comprised of molecules made up of atoms, which contain protons, neutrons, and electrons. Electrons play a key role in electricity as their movement constitutes electric current. Within atoms, electrons exist in specific energy levels or shells around the nucleus, and their behavior is described by quantum mechanics.
The document discusses atomic structure and bonding. It describes the structure of atoms including protons, neutrons, and electrons. It explains how atomic number determines the element and how isotopes have the same number of protons but different neutrons. Electron configuration and quantum numbers are also summarized. The three main types of bonds - ionic, covalent, and metallic - are introduced along with how they influence material properties.
The document discusses the development of atomic models from Dalton to Bohr and beyond. It describes Rutherford's discovery of the nucleus and Bohr's model of electrons in fixed orbits around the nucleus. Later, the quantum mechanical model was developed, restricting electrons to specific energy levels rather than exact orbits. This modern model determines the probability of finding electrons in different locations around the nucleus.
MODULE 1 ELETCRONIC STRUCTURE OF MATTER.pptxKathelynRuiz1
- The document discusses the history and development of atomic models from Dalton to Bohr. It describes key discoveries and experiments.
- Rutherford's gold foil experiment showed that the atom is mostly empty space, with a small, dense nucleus at the center containing positive charge. This led him to propose the nuclear model of the atom.
- Bohr incorporated Rutherford's nuclear model into his planetary model of the atom. He proposed that electrons orbit the nucleus in defined energy levels. This model had limitations when describing multi-electron atoms.
- The quantum mechanical model was developed based on discoveries that electrons can be described as waves using quantum physics principles like wave-particle duality and the uncertainty principle. This model uses atomic orbit
Atomic structure is composed of protons, neutrons, and electrons. Rutherford's gold foil experiment showed that atoms are mostly empty space with a small, dense nucleus at the center containing protons and neutrons. Electrons orbit the nucleus in specific energy levels. The Bohr model explained electrons in hydrogen having discrete energy levels. Atoms bond via ionic bonds between metals and nonmetals by transferring electrons, covalent bonds between nonmetals by sharing electrons, and metallic bonds between metals by delocalized electrons.
CBSE Class 9 Science Chapter 4- structure of atomAarthiSam
This document discusses the structure of atoms and the development of atomic models over time. It covers John Dalton's early atomic theory that atoms are indivisible and different for each element. J.J. Thomson's plum pudding model described atoms as electrically neutral spheres. Rutherford determined atoms have a small, dense nucleus through alpha particle scattering experiments. Niels Bohr incorporated quantum theory, proposing electrons orbit in discrete energy levels. The document also discusses atomic number, mass number, isotopes, electron configuration, and valency.
The document discusses the structure of atoms. It describes how early scientists like Dalton, Thomson, Rutherford, and Bohr contributed to developing models of the atom through experiments. Thomson's "plum pudding" model depicted electrons distributed uniformly in a positively charged sphere, but it could not explain atomic stability. Rutherford's gold foil experiment showed that the positive charge and mass of atoms are concentrated in a tiny nucleus, leading to his nuclear model. Bohr's model incorporated discrete electron orbits to explain atomic stability. The document also discusses subatomic particles like protons, neutrons, isotopes, isobars, and how electrons are arranged in shells according to their quantum numbers.
This document discusses Rutherford's atomic model and Bohr's model of the atom. It provides details of Rutherford's alpha particle scattering experiment which showed that atoms have a small, dense nucleus. This led Rutherford to propose a planetary model of the atom with electrons orbiting the nucleus. The document then discusses limitations of Rutherford's model and how Bohr proposed quantized electron orbits to explain atomic stability. It provides Bohr's key postulates and formulas for the hydrogen atom spectrum and energy levels.
The Fundamentals of Chemistry is an introduction to the Periodic Table, stoichiometry, chemical states, chemical equilibria, acid & base, oxidation & reduction reactions, chemical kinetics, inorganic nomenclature, and chemical bonding.
Atomic Structure Powerpoint Presentation by Computer CareersYaman Singhania
Powerpoint Presentation on Atomic Structure by Computer Careers.What is an Atom?ATOMIC STRUCTURE,There are two ways to represent the atomic structure of n element or compound,DOT & CROSS DIAGRAMS and many more ....
This document provides information about atomic structure. It discusses the basic parts of an atom including protons, neutrons, and electrons. Early atomic models proposed by Rutherford and Bohr are described, noting their limitations in explaining experimental observations. The modern quantum mechanical model represents electrons using quantum numbers and wave functions or "fuzzy clouds" to describe atomic orbitals. Electrons occupy different energy levels and sublevels based on their quantum numbers.
4.1 The Atomic Models of Thomson and Rutherford
4.2 Rutherford Scattering
4.3 The Classic Atomic Model
4.4 The Bohr Model of the Hydrogen Atom
4.5 Successes and Failures of the Bohr Model
4.6 Characteristic X-Ray Spectra and Atomic Number
4.7 Atomic Excitation by Electrons
This document discusses the electronic structure of atoms. It begins by reviewing early atomic models proposed by Thomson, Rutherford, and Chadwick that included a dense nucleus surrounded by electrons. The document then discusses how quantum mechanics provides a better model of electronic structure through the use of orbitals and quantum numbers to describe allowed electron configurations. Key points covered include the wave-particle duality of electrons, Schrodinger's wave equation describing orbital shape and orientation, and the four quantum numbers (n, l, ml, ms) that provide unique descriptions of electron states.
Physiology and chemistry of skin and pigmentation, hairs, scalp, lips and nail, Cleansing cream, Lotions, Face powders, Face packs, Lipsticks, Bath products, soaps and baby product,
Preparation and standardization of the following : Tonic, Bleaches, Dentifrices and Mouth washes & Tooth Pastes, Cosmetics for Nails.
Executive Directors Chat Leveraging AI for Diversity, Equity, and InclusionTechSoup
Let’s explore the intersection of technology and equity in the final session of our DEI series. Discover how AI tools, like ChatGPT, can be used to support and enhance your nonprofit's DEI initiatives. Participants will gain insights into practical AI applications and get tips for leveraging technology to advance their DEI goals.
it describes the bony anatomy including the femoral head , acetabulum, labrum . also discusses the capsule , ligaments . muscle that act on the hip joint and the range of motion are outlined. factors affecting hip joint stability and weight transmission through the joint are summarized.
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.
A review of the growth of the Israel Genealogy Research Association Database Collection for the last 12 months. Our collection is now passed the 3 million mark and still growing. See which archives have contributed the most. See the different types of records we have, and which years have had records added. You can also see what we have for the future.
This presentation includes basic of PCOS their pathology and treatment and also Ayurveda correlation of PCOS and Ayurvedic line of treatment mentioned in classics.
A Strategic Approach: GenAI in EducationPeter Windle
Artificial Intelligence (AI) technologies such as Generative AI, Image Generators and Large Language Models have had a dramatic impact on teaching, learning and assessment over the past 18 months. The most immediate threat AI posed was to Academic Integrity with Higher Education Institutes (HEIs) focusing their efforts on combating the use of GenAI in assessment. Guidelines were developed for staff and students, policies put in place too. Innovative educators have forged paths in the use of Generative AI for teaching, learning and assessments leading to pockets of transformation springing up across HEIs, often with little or no top-down guidance, support or direction.
This Gasta posits a strategic approach to integrating AI into HEIs to prepare staff, students and the curriculum for an evolving world and workplace. We will highlight the advantages of working with these technologies beyond the realm of teaching, learning and assessment by considering prompt engineering skills, industry impact, curriculum changes, and the need for staff upskilling. In contrast, not engaging strategically with Generative AI poses risks, including falling behind peers, missed opportunities and failing to ensure our graduates remain employable. The rapid evolution of AI technologies necessitates a proactive and strategic approach if we are to remain relevant.
This presentation was provided by Steph Pollock of The American Psychological Association’s Journals Program, and Damita Snow, of The American Society of Civil Engineers (ASCE), for the initial session of NISO's 2024 Training Series "DEIA in the Scholarly Landscape." Session One: 'Setting Expectations: a DEIA Primer,' was held June 6, 2024.
How to Add Chatter in the odoo 17 ERP ModuleCeline George
In Odoo, the chatter is like a chat tool that helps you work together on records. You can leave notes and track things, making it easier to talk with your team and partners. Inside chatter, all communication history, activity, and changes will be displayed.
1. Studocu is not sponsored or endorsed by any college or university
Atomic Theory Unit Test Review
Chemistry (High School - Canada)
Studocu is not sponsored or endorsed by any college or university
Atomic Theory Unit Test Review
Chemistry (High School - Canada)
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2. January 17th, 2019 Chloe Popov
Atomic Theory Unit Test Review
History of Atomic Theory
• Dalton: Matter’s composed of small parts called atoms.
• Thomson: Matter has positive and negative charged parts.
• Rutherford: Positive parts (nucleus) is the center, negative parts (electrons) are
in empty space, and the atom is mainly empty space.
• Planck: Matter emits EM radiation in quanta called photons.
• Einstein: The photons must have a threshold energy to remove an electron from
an atom. By colliding and removing electrons, photons act as particles.
• Bohr: Explained and calculated the size of hydrogens orbital radii (energy level),
and the energy of electrons at different orbitals (energy levels).
• De Brolie: Explained that electrons, like all matter, are particles that act like
waves.
• Schrodinger: Explained and calculated a region in space where an electron at a
specific energy level (orbital) can be found.
• Born: Calculated the probability of finding an electron in a region of space
(orbital).
• Heisenberg: Explained and calculated that for electrons, it’s difficult to know their
exact position and velocity at the same time.
Quantum Mechanics
• The current theory of atomic structure based on the wave properties of electrons.
• It’s currently unable to explain what electrons are and what they’re doing, and it
also has limits as to the precision with which we can measure and know things.
The 4 Quantum Numbers
• The Principal Quantum Number (n):
o Represents the energy level and relative size of the orbital. It can be any
integer and is the main energy level for an electron.
o To determine it, a line spectra of hydrogen was observed and it was
determined that electrons have specific energy levels, quanta, and that the
energy level describes the size of the orbital in which the electron can be
found.
• The Secondary Quantum Number (l):
o Represents the shape of the orbital and the energy sublevels for each
electron at a main level. They are the range of integers from 0 to n-1 and
are most commonly represented by the letters s (0), p (1), d (2), and f (3).
o To determine it, a higher quality spectra was observed where the initial
lines produced in the initial energy levels were thought to be made up from
electrons released at slightly different energies. Later, it was determined
that the different energy levels were due to the presence of different
shaped orbitals.
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3. January 17th, 2019 Chloe Popov
• The Magnetic Quantum Number (m):
o Represents the orientation or the orbital in space. Electrons with different
magnetic quantum numbers will have the same energy level but different
orientations. The number of possible magnetic quantum numbers is equal
to 2l+1 and the range are the integers from -l to +l.
o To determine it, spectra were observed under a magnetic field. The
energy released by the electrons was found to split again. The values
were later described as the orientation of the orbitals the electrons were
found.
• The Spin Quantum Number (ms):
o Represents the orientation which the electron spins on its axis and it’s
either +
1
2
or -
1
2
.
o To determine it, atoms/molecules were observed to have a weak magnetic
force. The difference in magnetic attraction was deemed due to the
movement of electrons in orbitals. Each orbital can hold a maximum of 2
electrons which must have opposite spins.
Equations
Planks Equation:
E= hf
E- minimum energy needed to remove an electron.
h- Planck’s constant which equals 6.6 x 10-34 J•s
f- frequency of light in Hz (s-1)
Rydberg Equation:
1
𝜆
= 𝑅𝐻(
1
𝑛𝑓
2 −
1
𝑛𝑖
2)
λ- wavelength of photon
RH- Rydberg constant which is 1.10 x 10-7 m-1
nf- final energy level
ni- initial energy level
Defintions:
Orbital: A region in space around the nucleus that the probability of finding an electron
is maximum. It represents the 3D motion of an electron and doesn’t specify a specific
path since the electron can be anywhere in the orbital region. This is part of the
uncertainty principle. They can’t accommodate more than 2 electrons.
Electron Energy Diagrams: A diagram which displays the characteristics of an atom’s
electrons and can help explain the behaviour of an atom.
Polarity: A measure of a molecule’s electrostatic difference within the molecule. To
determine polarity of a molecule, you have to find the polarity of it’s individual bonds.
Dipole: The separation of charge between 2 atoms within a molecule.
Diamagnetic: All electrons are paired; opposed/repelled by magnetic field.
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4. January 17th, 2019 Chloe Popov
Bohr’s Theory
• Bohr’s understanding of an atom was formed by the evidence that there’s a
periodicity of the physical and chemical properties of the elements, that there are
2 elements in the first period and 8 elements in the second period on the periodic
table, as well as the evidence that emission and absorption line spectra, but not
continuous spectra, exist for gaseous elements.
• Bohr believed that this evidence represented that electrons travel in the atom in
circular orbits with quantized energy, there’s a maximum number of electrons
allowed in each orbit, and electrons “jump” to a higher level when a photon is
absorbed and a photon is emitted when the electron “drops” to a lower level.
• Bohr’s first postulate was that electrons don’t radiate energy as they orbit the
nucleus. Each orbit corresponds to a state of constant energy, referred to as a
stationary state. Bohr’s second postulate was that electrons can change their
energy only by undergoing a transition from one stationary state to another.
• The main problem with his theory is that it worked well for only the spectrum for
hydrogen atoms or ions with only one electron. His theory can only explain the
results of experiments on atoms with single electrons.
• He thought that electrons orbited the nucleus in circular paths whereas in the
modernized concept, atomic electron structure is more alike 3D standing waves.
Electrons have probability distribution around the nucleus that form different
shaped orbitals.
Pauli Exclusion Principle
• No 2 electrons in an atom can have the same 4 quantum numbers.
• No 2 electrons in the same atomic orbital can have the same spin.
• Only 2 electrons with opposite spins can be in any one orbital.
Aufbau Principle
• Each electron is added to the lowest energy orbital available in an atom or ion.
They build up, filling orbitals from lowest to highest.
Hund’s Rule
• Each orbital will fill with one electron prior to the others filling with a second of the
opposite spin.
Electron Configuration Diagrams (Elements and Ions)
• Draw the atom first
• When losing electrons, they usually come from the s orbitals. The concentration
around the nucleus because of all of the overlapping s orbitals creates more
repulsion. There’s so much electron density near the nucleus because of the
overlapping s orbitals so there’s extreme repulsion, making electrons likely to
leave the highest s orbital to escape the extreme repulsion. They found that for
metals, electrons are usually lost from the highest s orbital. There’s more s orbital
density than d or p, so they usually come from the s orbital, even if it’s a lower
charge.
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VSEPR Theory
Predicting and explaining the shape of molecules.
1. Draw the Lewis structure including bonding and lone pairs of electrons.
2. All double and triple bonds are treated like single bonds.
3. Add the pairs of lone electrons and then bonding electrons.
4. Use the VSEPR table to find out the shape of the molecule (lone pairs repel the
bonding pairs forcing the atoms closer together).
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Polarity
Bond Polarity:
• To determine a bond’s polarity, you have to find the change in EN between the 2
bonding atoms.
Nonpolar Bond Polar Bond Ionic Bond
0→ 0.5→ 1.7→
Nonpolar Bonds
• Formed when there’s an equal sharing of electrons, so the electrons would have
a high chance of being found around both atoms’ nuclei.
Polar Bonds
• Formed when there’s an unequal sharing of electrons, so the electrons would
have a higher chance of being found around one atom’s nucleus than the other’s.
• When a polar bond is formed a dipole is produced within the molecule.
Ionic Bonds
• Formed when electrons are removed from one atom and move to another, so the
electrons have a high probability of being found around only one of the atoms’
nucleus.
Molecular Polarity:
• Based on 2 main factors, bond polarity and shape of the molecule. If a molecule
has no has no polar bonds, it would be a nonpolar molecule. If it does have polar
bonds, it’s polarity is based on the shape of the molecule and where the polar
bonds are present. If they’re in opposing directions to each other they cancel out.
If they’re in a similar direction they add up.
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Intermolecular Forces
• Forces between molecules. They’re much weaker than any intramolecular force
(ionic or covalent bonding).
• There are 2 types, hydrogen bonding and van der waals forces.
Hydrogen Bonding
• Exists when a molecule (not a ketone or aldehyde, because no oxygen and
hydrogen attached) has a highly electronegative atom such as O or N bonded to
a hydrogen atom. The high polarity between the hydrogen and O or N creates a
relatively strong bond between molecules.
• The strongest type of intermolecular forces.
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Van der Waals Forces
• London Forces: The weakest intermolecular force. The very weak attraction
between the positive particles (protons) that make up a molecule and the
negative particles (electrons) that make up a molecule. The more electrons
present, the stronger the London forces.
• Debye Forces: Second weakest intermolecular force. More important when
dealing with intermolecular forces between 2 different types of molecules
(solubility). Attraction of molecules with polarity to the positive and negative
particles of non-polar molecules.
• Keesom (dipole-dipole): Third weakest or second strongest intermolecular
force. If 2 polar molecules are present (same or different molecule) they would
have a stronger force of attraction.
Lewis Structures
1) Find electrons has (have) and needs (need).
2) If the compound is charged, add or subtract the number of electrons needed to
give it the charge (if negative then add, if positive then subtract).
3) Make a symmetrical arrangement of atoms.
4) Put the electrons needed on and then add the rest of the electrons (have).
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5) Find the formal charge of each atom by adding the total lone pairs plus half of the
bonding electrons around an atom and subtracting it from the valence electrons
that the atom started with.
6) Shift electrons so that the lowest magnitude of formal charges is produced and
the negative formal charge is on the more electronegative atom (whatever’s
closer to F has a higher electronegativity). Make sure there’s no formal charge on
as many atoms as possible. If it has to be charged, positive goes on the lower En
atom and negative goes on the higher EN atom.
7) If a charged compound, redraw with stick bonds, showing lone pairs, and
brackets with charge.
Ex, PO4-3
Valence Bond Theory
• Pauling’s theory that when 2 half-filled orbitals overlap, they bond. The electrons
would have opposing spins.
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The p orbitals of N have 1 electron each so the hydrogens, that also only have on
electron, bond covalently to share electrons.
• Hybrid orbitals are identical spontaneously formed orbitals that only form when
bonding happens. Each orbital has an equal energy level and are usually
represented by a petal shape.
Atomic Orbitals vs. Hybrid Orbitals
• Atomic Orbitals: s, p d, and f.
• Hybrid Orbitals: Orbitals of different shape and energy that form when
molecules form (they’re from different atomic orbitals).
Ex, CH4
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Multiple Bonds
• Valence bond theory is used to explain multiple bonds.
• There are 2 ways orbitals can overlap and that partial hybridization can happen
which leaves remaining orbitals as they are.
• The first way is end-to-end overlapping. This creates a sigma bond (an example
of this was NH3). They’re found in single, double, and triple bonds.
• The second way is side-by-side overlapping which creates a pi bond. They’re
found in both double and triple bonds.
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Atomic Spectra Lab Information
• Small transition (1 shell back)- lower energy of wavelength (red light).
• Big transition (2+ shells back)- higher energy of wavelength (violet light).
• When energy is added to a particle (i.e. gases) its electrons get “excited” and
jump to a higher energy level (further shell from the nucleus).
o The light we see is after this happens when the electrons transition back
down to a lower energy level (shell closer to nucleus). Light is released.
The closer to the nucleus, the higher the wavelength (violet-higher, blue-
lower).
• Only specific energy levels an electron can have, can’t have halfway.
• When electrons transition from an outer shell to the first are emitting UV light.
𝐻2 =
1
𝑠
= 𝑠−1
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Crystal Structures
• Ionic Crystals
o Solid structures made of ions/ionic compounds.
o Formed from metal and non-metal ions.
o At room temperature, they’re fairly hard but brittle.
o High melting points due to strong ionic bonds between ions.
o If dissolved in water, will conduct electricity but not as undissolved solids.
o If heated to point of melting, will conduct electricity.
• Metallic Crystals
o Solid structures made of metal atoms.
o At room temperature, range from soft to hard and are malleable.
o Differing ranges of melting points.
o Believed to form from indirect electron attraction to the nuclei of
surrounding atoms.
o Loosely held ‘cloud’ of electrons that surrounds the nuclei are believed to
be the reason why they’re shiny, flexible and able to conduct electricity.
• Molecular Crystals
o Formed from solid, covalently bonded molecules.
o Relatively soft and brittle.
o Relatively low melting points due to weak intermolecular bonds holding
together (van der Waals and hydrogen bonding).
o They don’t have freely moving electrons and aren’t ions so they don’t
conduct electricity.
• Covalent Network Crystals
o Formed from metalloids (B, Si, Ge, As, Sb, Te, Po) and crystals of carbon.
o Extremely hard because of atoms forming covalent bonds with
surrounding atoms.
o Brittle because they aren’t malleable, although difficult to break.
o Due to strong bonds, have very high melting points than the other 3
crystals.
o Don’t have free moving electrons and don’t conduct electricity.
o Most common include carbon, graphite and diamond, but other gemstones
would also be covalent networks too.
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