This 2 page summary discusses how to draw, name, and label the dot structures for three molecules, including two ions and one compound with a double bond.
An electron configuration is a shorthand description of how electrons are arranged around the nucleus of an atom. It helps predict chemical behavior by showing how elements will react and the type and strength of reactions. The electron configuration pattern in the periodic table shows that elements are organized by their atomic number, which is also the number of protons and electrons in a neutral atom. The standard electron configuration notation uses numbers and letters (s, p, d, f) with superscripts to represent the arrangement of electrons in orbitals.
The document discusses electronic configuration, which is the distribution of electrons in atomic or molecular orbitals. It explains that electrons are arranged in shells and subshells around the nucleus, with the subshells labeled s, p, d, and f. The order that electron subshells are filled is provided, with exceptions to a simple ordering. Examples of writing the electronic configuration of different elements are given by asking a series of questions.
Applied Chapter 3.3 : Electron ConfigurationChris Foltz
The document discusses electron configuration and the quantum models of the atom. It compares the Rutherford, Bohr, and quantum models. It explains the four quantum numbers - principal, angular momentum, magnetic, and spin quantum numbers. It describes how light emission spectra provide information about an atom's energy levels. Rules for writing electron configurations are given, including the Aufbau principle, Pauli exclusion principle, and Hund's rule. Orbital notation, electron configuration notation, and noble gas notation are defined. Examples are provided of writing the electron configuration for atoms with specific atomic numbers.
The document outlines objectives for learning about electron configuration including listing the number of electrons in each main energy level, stating the Aufbau principle, Pauli exclusion principle, and Hund's rule, and describing electron configurations using orbital, electron-configuration, and noble-gas notations. It then defines electron configuration and the ground state and provides rules for electron configuration including the Aufbau principle, Pauli exclusion principle, and Hund's rule.
Electron configuration represents the arrangement of electrons in an atom's orbital shells and subshells. There are three main rules that determine electron configuration: the Aufbau principle, which states that orbitals are filled in order of increasing energy; the Pauli exclusion principle, where no two electrons can have the same quantum numbers; and Hund's rule, where orbitals in a subshell are singly occupied with parallel spins before pairing. Electron configuration can be written out or represented through orbital diagrams, the periodic table, or Möller diagrams.
The document discusses the rules for writing electron configurations of elements:
- The aufbau principle states that electrons occupy the lowest energy orbitals first.
- The Pauli exclusion principle allows a maximum of two electrons per orbital with opposite spins.
- According to Hund's rule, electrons occupy degenerate orbitals with parallel spins before pairing.
This document discusses quantum numbers and their role in describing electron orbitals and configurations. It covers the principal (n), azimuthal (l), and magnetic (ml) quantum numbers, as well as electron spin (ms). The document defines orbitals for the first five energy levels, discusses how electrons fill orbitals based on the Aufbau principle and Hund's rule, and notes exceptions like chromium and copper. It asks the reader to write electron configurations and diagrams for chlorine, osmium, and cesium.
The document discusses electron configurations and how to write them. It explains that electron configurations describe how electrons are arranged in an atom's shells and subshells. It provides examples of writing configurations for different elements like oxygen, bromine, sulfur, rubidium, and barium. The document also introduces noble gas configurations, which provide a shorthand version of electron configurations by writing the closest noble gas in brackets followed by the remaining electrons.
An electron configuration is a shorthand description of how electrons are arranged around the nucleus of an atom. It helps predict chemical behavior by showing how elements will react and the type and strength of reactions. The electron configuration pattern in the periodic table shows that elements are organized by their atomic number, which is also the number of protons and electrons in a neutral atom. The standard electron configuration notation uses numbers and letters (s, p, d, f) with superscripts to represent the arrangement of electrons in orbitals.
The document discusses electronic configuration, which is the distribution of electrons in atomic or molecular orbitals. It explains that electrons are arranged in shells and subshells around the nucleus, with the subshells labeled s, p, d, and f. The order that electron subshells are filled is provided, with exceptions to a simple ordering. Examples of writing the electronic configuration of different elements are given by asking a series of questions.
Applied Chapter 3.3 : Electron ConfigurationChris Foltz
The document discusses electron configuration and the quantum models of the atom. It compares the Rutherford, Bohr, and quantum models. It explains the four quantum numbers - principal, angular momentum, magnetic, and spin quantum numbers. It describes how light emission spectra provide information about an atom's energy levels. Rules for writing electron configurations are given, including the Aufbau principle, Pauli exclusion principle, and Hund's rule. Orbital notation, electron configuration notation, and noble gas notation are defined. Examples are provided of writing the electron configuration for atoms with specific atomic numbers.
The document outlines objectives for learning about electron configuration including listing the number of electrons in each main energy level, stating the Aufbau principle, Pauli exclusion principle, and Hund's rule, and describing electron configurations using orbital, electron-configuration, and noble-gas notations. It then defines electron configuration and the ground state and provides rules for electron configuration including the Aufbau principle, Pauli exclusion principle, and Hund's rule.
Electron configuration represents the arrangement of electrons in an atom's orbital shells and subshells. There are three main rules that determine electron configuration: the Aufbau principle, which states that orbitals are filled in order of increasing energy; the Pauli exclusion principle, where no two electrons can have the same quantum numbers; and Hund's rule, where orbitals in a subshell are singly occupied with parallel spins before pairing. Electron configuration can be written out or represented through orbital diagrams, the periodic table, or Möller diagrams.
The document discusses the rules for writing electron configurations of elements:
- The aufbau principle states that electrons occupy the lowest energy orbitals first.
- The Pauli exclusion principle allows a maximum of two electrons per orbital with opposite spins.
- According to Hund's rule, electrons occupy degenerate orbitals with parallel spins before pairing.
This document discusses quantum numbers and their role in describing electron orbitals and configurations. It covers the principal (n), azimuthal (l), and magnetic (ml) quantum numbers, as well as electron spin (ms). The document defines orbitals for the first five energy levels, discusses how electrons fill orbitals based on the Aufbau principle and Hund's rule, and notes exceptions like chromium and copper. It asks the reader to write electron configurations and diagrams for chlorine, osmium, and cesium.
The document discusses electron configurations and how to write them. It explains that electron configurations describe how electrons are arranged in an atom's shells and subshells. It provides examples of writing configurations for different elements like oxygen, bromine, sulfur, rubidium, and barium. The document also introduces noble gas configurations, which provide a shorthand version of electron configurations by writing the closest noble gas in brackets followed by the remaining electrons.
The document discusses electronic configuration, which is the arrangement of electrons in an atom's orbitals. It is described using symbols that indicate the principal shell, subshell, and number of electrons. The Aufbau principle states that electrons fill the lowest available energy levels. Pauli's exclusion principle limits each orbital to two electrons with different quantum numbers. Hund's rule states that orbitals in a subshell will each have one electron before any are doubly filled, with parallel electron spins. Partial configurations, orbital diagrams, and number of inner electrons are provided for potassium, molybdenum, and lead as examples. Key terms like isoelectronic, valence electrons, and magnetic properties are also defined.
This document discusses electron configuration and the principles that govern how electrons are arranged in atoms. It explains that electrons exhibit wave-like properties and occupy regions of space called atomic orbitals. The distribution of electrons in atoms, known as electron configuration, follows three main principles: the Pauli Exclusion Principle limits orbitals to two electrons of opposite spin; the Aufbau Principle states that orbitals are filled from lowest to highest energy; and Hund's Rule specifies that orbitals are singly occupied with parallel electron spins before double occupation occurs.
- Electron configuration describes the arrangement of electrons in an atom's orbitals and energy levels according to 4 quantum numbers.
- The Pauli Exclusion Principle states that no two electrons can have the same set of 4 quantum numbers.
- Elements are arranged on the periodic table based on their electron configurations, with groups reflecting which orbitals are being filled.
- A noble gas configuration can be used to abbreviate long electron configurations by writing the nearest noble gas followed by the remaining electrons.
The document discusses electron configuration and energy levels. It introduces the historical models of the atom including Rutherford's plum pudding model and Bohr's model introducing energy levels. Modern models use orbitals that show the probability of finding electrons. Light can be described as both a wave and particle. Electrons absorb energy to move to excited states then emit light returning to ground states. Elements' electron configurations follow the aufbau principle filling lower energy orbitals first up to the periodic table.
The document discusses electronic configurations of atoms. It explains that the electron configuration represents the arrangement of electrons in an atom's orbital shells and subshells in its ground state, and can also represent ionized atoms. Many physical and chemical properties correlate to unique electron configurations, especially the valence electrons in the outermost shell. Electrons fill orbitals according to increasing energy levels and subshells in a set order. Orbital diagrams, spdf notation, and noble gas notation are used to represent electron configurations.
How to draw the lewis structure of carbon monoxidegengarl
The document provides instructions for drawing the Lewis structure of carbon monoxide (CO). It explains that carbon contributes 4 valence electrons and oxygen contributes 6, for a total of 10. A carbon-oxygen single bond is drawn by pairing 2 electrons between them. However, carbon then only has 6 electrons and needs 8 to complete its octet. Therefore, oxygen must donate its remaining 2 electrons to carbon to form a triple bond, completing the octet for both atoms.
Orbital notation represents the electron configuration of an element by showing the placement of each electron through arrows. It illustrates the Pauli exclusion principle, which states that no two electrons in the same atom can have the same quantum numbers, requiring opposite spin directions. The homework involves writing out electron configurations using orbital notation diagrams for various elements like sodium, boron, and titanium.
Formal charges provide a simple method to estimate charge distribution within ions and molecules using integer charges assigned to each atom. The procedure involves drawing Lewis structures, determining the group number, unshared electrons, and bonded electrons for each atom to calculate its formal charge. Common patterns show atoms of the same element in a structure usually have the same formal charge, and structures aim to minimize the number of atoms with formal charges other than 0, +1, or -1. While only an estimate, formal charges can predict reactivity and properties by indicating which atoms are most electron rich or deficient.
Electron configuration refers to the distribution of electrons in an atom's orbitals and energy levels. There are three main principles that determine electron configuration: 1) Electrons have higher energy further from the nucleus and in higher principal quantum shells, 2) Orbitals within a shell have set ordering of energies (s < p < d < f), and 3) Each orbital can hold a maximum of two electrons of opposite spin. The Aufbau principle, Pauli exclusion principle, and Hund's rule further define how electrons fill the orbitals in atoms.
This document defines and explains electronic configuration, which shows the distribution of electrons in an atom or molecule. It describes electron shells as the areas where electrons orbit, and atomic orbitals as specific regions that electrons can occupy according to set filling orders. Overlap orbitals occur when electrons share the same orbital space, as in the H2 molecule. The document also provides an example of how to write electronic configurations and use an electron configuration table to visualize the notation.
Electron configurations 1a presentationPaul Cummings
The document discusses electron configuration, which is the arrangement of electrons around an atom's nucleus. It describes the quantum mechanical model developed in the 1920s using quantum numbers like the principal quantum number n, angular momentum quantum number l, and magnetic quantum number m. Electrons occupy specific orbitals within energy sublevels based on these quantum numbers. There are rules for building electron configurations, including the Aufbau principle, Pauli exclusion principle, and Hund's rule. Electron configurations are written using standard or shorthand notation.
The document provides an overview of electron configuration, which is the arrangement of electrons in an atom. It explains the key concepts of principal quantum number (n), sublevels (s, p, d, f), orbitals, the Aufbau principle, Pauli exclusion principle, Hund's rule, and how to write out the electron configuration for different elements. Examples are given for elements such as hydrogen, helium, lithium, carbon, nitrogen, fluorine, aluminum, argon, iron, and lanthanum.
The document provides sample exercises for summarizing Lewis structures and predicting properties of ionic compounds based on their Lewis structures. It includes examples of determining formal charges, identifying preferred resonance structures, and predicting whether SO3 or SO32- will have shorter S-O bond lengths based on their Lewis structures. The sample exercises provide solutions and explanations for practicing drawing and analyzing Lewis structures.
This document provides information about electron configuration, which describes the arrangement of electrons in an atom. It discusses three main principles for determining electron configuration:
1. The Aufbau principle states that electrons occupy the lowest available energy orbitals. Energy levels are filled in order of increasing energy, with s orbitals having the lowest energy followed by p, d, and f orbitals.
2. Hund's rule states that electrons occupy different orbitals within the same energy level with parallel spin before pairing electrons with opposite spin in the same orbital.
3. The Pauli exclusion principle allows no more than two electrons to occupy the same orbital, and these electrons must have opposite spins.
The document uses an
This document discusses electron configuration and the rules that define how electrons are arranged in an atom's orbitals. It explains:
1) There are three main rules that define electron configuration: the Aufbau principle, Pauli exclusion principle, and Hund's rule.
2) Higher energy levels can hold more electrons than lower energy levels because they are associated with larger volumes that can contain more orbitals.
3) Electron configuration can be represented using orbital diagrams with arrows or electron configuration notation using symbols and superscripts.
This document discusses electron configuration and how electrons are arranged in atoms. It introduces the Bohr model of electron shells and energy levels, then describes more accurate quantum mechanical models of electron orbitals. The document explains the Aufbau principle of filling orbitals from lowest to highest energy, the Pauli exclusion principle of maximum two electrons per orbital, and Hund's rule of filling orbitals within the same level. It outlines the shapes and notations of s, p, d, and f orbitals and how they relate to the periodic table and electron configuration notation.
Chapter 8 electron configuration and periodicity (1)ElizabethAyala45
The document discusses electron configurations and periodic trends in atomic properties. It describes how electrons fill atomic orbitals according to the building-up principle and Hund's rule. Trends in atomic radius, ionization energy, and electron affinity across the periodic table are also explained, with atomic radius generally decreasing and ionization energy and electron affinity (becoming more negative) generally increasing within a period. Exceptions to trends are seen for some p-block elements.
- Thompson discovered electrons using cathode ray tubes. Rutherford discovered the nucleus and that it is positively charged through gold foil experiments.
- Chadwick discovered neutrons by observing particles in a beam that were not deflected by a magnetic field.
- Atoms are made up of protons, neutrons, and electrons. Protons and neutrons reside in the nucleus, while electrons orbit the nucleus. The number of protons defines the element.
The document discusses Lewis structures and the rules for drawing them. It explains that Lewis structures show how atoms bond via shared electron pairs to achieve stable noble gas configurations. It provides a 4-step process for drawing Lewis structures, covering counting electrons, identifying the central atom, adding lone pairs to complete octets, and checking that all electrons are accounted for. Exceptions to the octet rule and drawing structures for ions are also covered.
The document outlines 7 general rules for drawing Lewis structures:
1. Count the total valence electrons from the group numbers of each atom. Add/subtract electrons for ions.
2. Calculate the total electrons needed for each atom to have an octet or doublet.
3. Subtract the total valence electrons from the total electrons needed to get the bonding electrons.
4. Assign two electrons to each bond. Assign remaining electrons to double/triple bonds if possible.
5. Assign any electrons left as lone pairs to complete octets, except for hydrogen.
6. Calculate and note the formal charges to ensure they add up to the total charge on the molecule/ion.
The document discusses electronic configuration, which is the arrangement of electrons in an atom's orbitals. It is described using symbols that indicate the principal shell, subshell, and number of electrons. The Aufbau principle states that electrons fill the lowest available energy levels. Pauli's exclusion principle limits each orbital to two electrons with different quantum numbers. Hund's rule states that orbitals in a subshell will each have one electron before any are doubly filled, with parallel electron spins. Partial configurations, orbital diagrams, and number of inner electrons are provided for potassium, molybdenum, and lead as examples. Key terms like isoelectronic, valence electrons, and magnetic properties are also defined.
This document discusses electron configuration and the principles that govern how electrons are arranged in atoms. It explains that electrons exhibit wave-like properties and occupy regions of space called atomic orbitals. The distribution of electrons in atoms, known as electron configuration, follows three main principles: the Pauli Exclusion Principle limits orbitals to two electrons of opposite spin; the Aufbau Principle states that orbitals are filled from lowest to highest energy; and Hund's Rule specifies that orbitals are singly occupied with parallel electron spins before double occupation occurs.
- Electron configuration describes the arrangement of electrons in an atom's orbitals and energy levels according to 4 quantum numbers.
- The Pauli Exclusion Principle states that no two electrons can have the same set of 4 quantum numbers.
- Elements are arranged on the periodic table based on their electron configurations, with groups reflecting which orbitals are being filled.
- A noble gas configuration can be used to abbreviate long electron configurations by writing the nearest noble gas followed by the remaining electrons.
The document discusses electron configuration and energy levels. It introduces the historical models of the atom including Rutherford's plum pudding model and Bohr's model introducing energy levels. Modern models use orbitals that show the probability of finding electrons. Light can be described as both a wave and particle. Electrons absorb energy to move to excited states then emit light returning to ground states. Elements' electron configurations follow the aufbau principle filling lower energy orbitals first up to the periodic table.
The document discusses electronic configurations of atoms. It explains that the electron configuration represents the arrangement of electrons in an atom's orbital shells and subshells in its ground state, and can also represent ionized atoms. Many physical and chemical properties correlate to unique electron configurations, especially the valence electrons in the outermost shell. Electrons fill orbitals according to increasing energy levels and subshells in a set order. Orbital diagrams, spdf notation, and noble gas notation are used to represent electron configurations.
How to draw the lewis structure of carbon monoxidegengarl
The document provides instructions for drawing the Lewis structure of carbon monoxide (CO). It explains that carbon contributes 4 valence electrons and oxygen contributes 6, for a total of 10. A carbon-oxygen single bond is drawn by pairing 2 electrons between them. However, carbon then only has 6 electrons and needs 8 to complete its octet. Therefore, oxygen must donate its remaining 2 electrons to carbon to form a triple bond, completing the octet for both atoms.
Orbital notation represents the electron configuration of an element by showing the placement of each electron through arrows. It illustrates the Pauli exclusion principle, which states that no two electrons in the same atom can have the same quantum numbers, requiring opposite spin directions. The homework involves writing out electron configurations using orbital notation diagrams for various elements like sodium, boron, and titanium.
Formal charges provide a simple method to estimate charge distribution within ions and molecules using integer charges assigned to each atom. The procedure involves drawing Lewis structures, determining the group number, unshared electrons, and bonded electrons for each atom to calculate its formal charge. Common patterns show atoms of the same element in a structure usually have the same formal charge, and structures aim to minimize the number of atoms with formal charges other than 0, +1, or -1. While only an estimate, formal charges can predict reactivity and properties by indicating which atoms are most electron rich or deficient.
Electron configuration refers to the distribution of electrons in an atom's orbitals and energy levels. There are three main principles that determine electron configuration: 1) Electrons have higher energy further from the nucleus and in higher principal quantum shells, 2) Orbitals within a shell have set ordering of energies (s < p < d < f), and 3) Each orbital can hold a maximum of two electrons of opposite spin. The Aufbau principle, Pauli exclusion principle, and Hund's rule further define how electrons fill the orbitals in atoms.
This document defines and explains electronic configuration, which shows the distribution of electrons in an atom or molecule. It describes electron shells as the areas where electrons orbit, and atomic orbitals as specific regions that electrons can occupy according to set filling orders. Overlap orbitals occur when electrons share the same orbital space, as in the H2 molecule. The document also provides an example of how to write electronic configurations and use an electron configuration table to visualize the notation.
Electron configurations 1a presentationPaul Cummings
The document discusses electron configuration, which is the arrangement of electrons around an atom's nucleus. It describes the quantum mechanical model developed in the 1920s using quantum numbers like the principal quantum number n, angular momentum quantum number l, and magnetic quantum number m. Electrons occupy specific orbitals within energy sublevels based on these quantum numbers. There are rules for building electron configurations, including the Aufbau principle, Pauli exclusion principle, and Hund's rule. Electron configurations are written using standard or shorthand notation.
The document provides an overview of electron configuration, which is the arrangement of electrons in an atom. It explains the key concepts of principal quantum number (n), sublevels (s, p, d, f), orbitals, the Aufbau principle, Pauli exclusion principle, Hund's rule, and how to write out the electron configuration for different elements. Examples are given for elements such as hydrogen, helium, lithium, carbon, nitrogen, fluorine, aluminum, argon, iron, and lanthanum.
The document provides sample exercises for summarizing Lewis structures and predicting properties of ionic compounds based on their Lewis structures. It includes examples of determining formal charges, identifying preferred resonance structures, and predicting whether SO3 or SO32- will have shorter S-O bond lengths based on their Lewis structures. The sample exercises provide solutions and explanations for practicing drawing and analyzing Lewis structures.
This document provides information about electron configuration, which describes the arrangement of electrons in an atom. It discusses three main principles for determining electron configuration:
1. The Aufbau principle states that electrons occupy the lowest available energy orbitals. Energy levels are filled in order of increasing energy, with s orbitals having the lowest energy followed by p, d, and f orbitals.
2. Hund's rule states that electrons occupy different orbitals within the same energy level with parallel spin before pairing electrons with opposite spin in the same orbital.
3. The Pauli exclusion principle allows no more than two electrons to occupy the same orbital, and these electrons must have opposite spins.
The document uses an
This document discusses electron configuration and the rules that define how electrons are arranged in an atom's orbitals. It explains:
1) There are three main rules that define electron configuration: the Aufbau principle, Pauli exclusion principle, and Hund's rule.
2) Higher energy levels can hold more electrons than lower energy levels because they are associated with larger volumes that can contain more orbitals.
3) Electron configuration can be represented using orbital diagrams with arrows or electron configuration notation using symbols and superscripts.
This document discusses electron configuration and how electrons are arranged in atoms. It introduces the Bohr model of electron shells and energy levels, then describes more accurate quantum mechanical models of electron orbitals. The document explains the Aufbau principle of filling orbitals from lowest to highest energy, the Pauli exclusion principle of maximum two electrons per orbital, and Hund's rule of filling orbitals within the same level. It outlines the shapes and notations of s, p, d, and f orbitals and how they relate to the periodic table and electron configuration notation.
Chapter 8 electron configuration and periodicity (1)ElizabethAyala45
The document discusses electron configurations and periodic trends in atomic properties. It describes how electrons fill atomic orbitals according to the building-up principle and Hund's rule. Trends in atomic radius, ionization energy, and electron affinity across the periodic table are also explained, with atomic radius generally decreasing and ionization energy and electron affinity (becoming more negative) generally increasing within a period. Exceptions to trends are seen for some p-block elements.
- Thompson discovered electrons using cathode ray tubes. Rutherford discovered the nucleus and that it is positively charged through gold foil experiments.
- Chadwick discovered neutrons by observing particles in a beam that were not deflected by a magnetic field.
- Atoms are made up of protons, neutrons, and electrons. Protons and neutrons reside in the nucleus, while electrons orbit the nucleus. The number of protons defines the element.
The document discusses Lewis structures and the rules for drawing them. It explains that Lewis structures show how atoms bond via shared electron pairs to achieve stable noble gas configurations. It provides a 4-step process for drawing Lewis structures, covering counting electrons, identifying the central atom, adding lone pairs to complete octets, and checking that all electrons are accounted for. Exceptions to the octet rule and drawing structures for ions are also covered.
The document outlines 7 general rules for drawing Lewis structures:
1. Count the total valence electrons from the group numbers of each atom. Add/subtract electrons for ions.
2. Calculate the total electrons needed for each atom to have an octet or doublet.
3. Subtract the total valence electrons from the total electrons needed to get the bonding electrons.
4. Assign two electrons to each bond. Assign remaining electrons to double/triple bonds if possible.
5. Assign any electrons left as lone pairs to complete octets, except for hydrogen.
6. Calculate and note the formal charges to ensure they add up to the total charge on the molecule/ion.
Molecular orbital theory provides an approach to calculate molecular orbitals through a variational method. This involves taking linear combinations of atomic orbitals to form molecular orbitals. Electrons occupy these molecular orbitals according to certain rules. The molecular orbital theory can explain properties such as why some molecules are paramagnetic that valence bond theory cannot. Calculating molecular orbitals variationally involves using trial wave functions in the Schrodinger equation to find the lowest possible energy state.
This 4 step process for writing Lewis dot structures involves: 1) arranging atoms with the central atom having lower electronegativity, 2) adding up valence electrons and accounting for any charges, 3) drawing single bonds between atoms and removing 2 electrons for each, and 4) distributing any remaining electrons in pairs to satisfy the octet rule and complete the structure.
The document summarizes key information about atomic structure:
- The nucleus contains protons and neutrons and is positively charged. Electrons orbit around the outside in shells and are negatively charged.
- Electrons fill inner shells first and are arranged in set numbers per shell (2 in the first shell, 8 in the second, etc.).
- Bohr models represent atoms by drawing the nucleus and electrons in shells, with the number of protons and neutrons labeled in the nucleus.
The document provides information about atomic structure including:
- Atoms consist of a nucleus surrounded by electrons, with protons and neutrons in the nucleus.
- Electrons are arranged in shells, with the first shell holding up to 2 electrons and subsequent shells holding up to 8 electrons each.
- The number of protons equals the atomic number and number of electrons. Neutrons plus protons equals the mass number.
- Students are assigned tasks to learn about electron configurations, atomic shorthand, and properties of elements based on their position in the periodic table.
The document summarizes key information about atomic structure:
- The nucleus is positively charged and contains nearly all an atom's mass, while electrons are much smaller and negatively charged, orbiting in shells outside the nucleus.
- Electrons are arranged in shells (also called energy levels) around the nucleus, with the first shell holding up to 2 electrons and subsequent shells holding up to 8 electrons each.
- Atoms can be represented using Bohr models that show the nucleus and electrons arranged in shells, with the number of protons and neutrons indicated in the nucleus.
This document discusses covalent compounds and their formation through shared electron pairs between nonmetals. It covers the octet rule for achieving stable electron configurations, different types of covalent bonds, and how to draw Lewis structures by arranging electrons around atoms. Exceptions to the octet rule are presented. Guidelines for naming covalent compounds from their formulas and writing formulas from names are also provided, along with examples.
Lewis structures show the bonding between atoms in a molecule using dots to represent valence electrons. The octet rule states that atoms are most stable when their valence shells are filled with eight electrons. Valence shell electron pair repulsion (VSEPR) theory predicts molecular geometry based on minimizing electron pair repulsion around a central atom. VSEPR identifies five basic molecular geometries (linear, trigonal planar, tetrahedral, trigonal bipyramidal, and octahedral) depending on the number of electron pairs around the central atom. Lone pairs influence molecular geometry differently than bonding pairs.
The document discusses valence electrons and bonding. It introduces the 2-8-8 rule which states that the first energy level can hold up to 2 electrons, and the second and third levels can each hold up to 8 electrons. Atoms are stable once their energy levels are filled. Ionic bonds form when a metal atom transfers electrons to a nonmetal, becoming cations and anions. Covalent bonds form when atoms share electrons rather than transfer them. Metallic bonds form a "sea" of delocalized electrons between positively charged metal ions.
Ionic bonding occurs through the transfer of electrons between atoms to form ions, resulting in electrostatic attraction. Covalent bonding involves the sharing of electrons between atoms to form molecules. The shape of covalently bonded molecules can be predicted using VSEPR theory. Hydrogen bonding and van der Waals forces are weaker intermolecular forces that influence properties like boiling points. Bonding types exist on a continuum between purely ionic and purely covalent.
This document provides a summary of key concepts and steps for drawing Lewis structures of molecules and ions. It defines important terms like valence electrons, octet rule, and bonding vs. lone pairs. It outlines a 6-step process for drawing Lewis structures, including determining the number of valence electrons and arranging atoms to achieve full valence shells. Exceptions to the octet rule are noted for small atoms and those in period 3 or below. Mnemonics are provided to help remember electron configurations.
The document discusses chemical bonding and Lewis structures. It begins by defining a chemical bond as the force that holds atoms together, and discusses how atoms combine or share electrons to form ionic or covalent bonds. It then explains Lewis structures, showing how to draw the Lewis dot symbols and structures for various molecules by placing the atoms and distributing electrons to achieve full octets. Exceptions to the octet rule are also noted. Hybridization and theories of covalent bonding such as valence bond theory are introduced.
The document discusses chemical bonding and Lewis structures. It begins by defining a chemical bond as the force that holds atoms together, and discusses how atoms combine or share electrons to form ionic or covalent bonds. It then explains Lewis structures, showing how to draw the Lewis dot symbols and determine the hybridization of atoms. Examples are provided of writing Lewis structures for different molecules like CCl4 and NH4+. The document also discusses exceptions to the octet rule and theories of covalent bonding like valence bond theory and hybridization theory.
Lewis dot diagrams show the localization of valence electrons in atoms and molecules. They are drawn by placing the chemical symbol of each atom and then arranging electrons around the symbols to achieve full octets. For molecules, the least electronegative atom is placed in the center and single bonds are drawn between atoms before distributing remaining electrons. The octet rule is followed to maximize each atom's valence electrons to 8, with exceptions for hydrogen and some other elements. Practice examples are provided to demonstrate drawing Lewis dot diagrams for different atoms and molecules.
Class 11 Chapter 4 Chemical Bonding and Molecular Structure.pptxRajnishPrasadSarma
This document provides an overview of chemical bonding and molecular structure. It discusses topics such as octet rule, covalent bonds, limitations of the octet rule, ionic or electrovalent bonds, lattice enthalpy, bond parameters including bond length, bond angle, bond enthalpy and bond order. It also covers concepts of resonance, polar covalent bonds, dipole moment and covalent character in ionic bonds based on Fajans' rule. The document is presented as part of a Class XI chemistry curriculum on this unit.
This document discusses electron configuration and orbital diagrams. It explains how to write electron configurations, draw orbital diagrams, and the importance of understanding electron configuration. The key principles for determining electron configuration are explained, including the Aufbau principle for filling orbitals, Hund's rule for maximizing unpaired electrons, and Pauli's exclusion principle limiting each orbital to two electrons of opposite spin. Examples are provided and an activity asks the reader to write configurations and draw diagrams for elements.
The document discusses different types of bonds - metallic, ionic, and covalent - and provides details about each. Metallic bonds form between metal atoms and allow for electron transfer and conductivity. Ionic bonds form between ionic compounds through electrostatic attraction between oppositely charged ions. Covalent bonds form between nonmetal elements through the sharing of electron pairs to achieve stable noble gas configurations. The document also covers Lewis structures, exceptions to the octet rule, and periodic trends in ionization energy, electronegativity, atomic size, and ionic size.
Lewis structuresvsepr theory cheat sheetTimothy Welsh
1. The document provides an overview of key concepts for Lewis structures and VSEPR theory, including how to draw Lewis structures for covalent compounds and ions.
2. It explains valence shell electron pair repulsion theory (VSEPR) which is used to predict the 3D shape of molecules based on electron pairs around a central atom.
3. A table is included that lists common molecular geometries determined by VSEPR theory based on the number of electron regions around the central atom.
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Unlocking the mysteries of reproduction: Exploring fecundity and gonadosomati...AbdullaAlAsif1
The pygmy halfbeak Dermogenys colletei, is known for its viviparous nature, this presents an intriguing case of relatively low fecundity, raising questions about potential compensatory reproductive strategies employed by this species. Our study delves into the examination of fecundity and the Gonadosomatic Index (GSI) in the Pygmy Halfbeak, D. colletei (Meisner, 2001), an intriguing viviparous fish indigenous to Sarawak, Borneo. We hypothesize that the Pygmy halfbeak, D. colletei, may exhibit unique reproductive adaptations to offset its low fecundity, thus enhancing its survival and fitness. To address this, we conducted a comprehensive study utilizing 28 mature female specimens of D. colletei, carefully measuring fecundity and GSI to shed light on the reproductive adaptations of this species. Our findings reveal that D. colletei indeed exhibits low fecundity, with a mean of 16.76 ± 2.01, and a mean GSI of 12.83 ± 1.27, providing crucial insights into the reproductive mechanisms at play in this species. These results underscore the existence of unique reproductive strategies in D. colletei, enabling its adaptation and persistence in Borneo's diverse aquatic ecosystems, and call for further ecological research to elucidate these mechanisms. This study lends to a better understanding of viviparous fish in Borneo and contributes to the broader field of aquatic ecology, enhancing our knowledge of species adaptations to unique ecological challenges.
The use of Nauplii and metanauplii artemia in aquaculture (brine shrimp).pptxMAGOTI ERNEST
Although Artemia has been known to man for centuries, its use as a food for the culture of larval organisms apparently began only in the 1930s, when several investigators found that it made an excellent food for newly hatched fish larvae (Litvinenko et al., 2023). As aquaculture developed in the 1960s and ‘70s, the use of Artemia also became more widespread, due both to its convenience and to its nutritional value for larval organisms (Arenas-Pardo et al., 2024). The fact that Artemia dormant cysts can be stored for long periods in cans, and then used as an off-the-shelf food requiring only 24 h of incubation makes them the most convenient, least labor-intensive, live food available for aquaculture (Sorgeloos & Roubach, 2021). The nutritional value of Artemia, especially for marine organisms, is not constant, but varies both geographically and temporally. During the last decade, however, both the causes of Artemia nutritional variability and methods to improve poorquality Artemia have been identified (Loufi et al., 2024).
Brine shrimp (Artemia spp.) are used in marine aquaculture worldwide. Annually, more than 2,000 metric tons of dry cysts are used for cultivation of fish, crustacean, and shellfish larva. Brine shrimp are important to aquaculture because newly hatched brine shrimp nauplii (larvae) provide a food source for many fish fry (Mozanzadeh et al., 2021). Culture and harvesting of brine shrimp eggs represents another aspect of the aquaculture industry. Nauplii and metanauplii of Artemia, commonly known as brine shrimp, play a crucial role in aquaculture due to their nutritional value and suitability as live feed for many aquatic species, particularly in larval stages (Sorgeloos & Roubach, 2021).
Immersive Learning That Works: Research Grounding and Paths ForwardLeonel Morgado
We will metaverse into the essence of immersive learning, into its three dimensions and conceptual models. This approach encompasses elements from teaching methodologies to social involvement, through organizational concerns and technologies. Challenging the perception of learning as knowledge transfer, we introduce a 'Uses, Practices & Strategies' model operationalized by the 'Immersive Learning Brain' and ‘Immersion Cube’ frameworks. This approach offers a comprehensive guide through the intricacies of immersive educational experiences and spotlighting research frontiers, along the immersion dimensions of system, narrative, and agency. Our discourse extends to stakeholders beyond the academic sphere, addressing the interests of technologists, instructional designers, and policymakers. We span various contexts, from formal education to organizational transformation to the new horizon of an AI-pervasive society. This keynote aims to unite the iLRN community in a collaborative journey towards a future where immersive learning research and practice coalesce, paving the way for innovative educational research and practice landscapes.
Describing and Interpreting an Immersive Learning Case with the Immersion Cub...Leonel Morgado
Current descriptions of immersive learning cases are often difficult or impossible to compare. This is due to a myriad of different options on what details to include, which aspects are relevant, and on the descriptive approaches employed. Also, these aspects often combine very specific details with more general guidelines or indicate intents and rationales without clarifying their implementation. In this paper we provide a method to describe immersive learning cases that is structured to enable comparisons, yet flexible enough to allow researchers and practitioners to decide which aspects to include. This method leverages a taxonomy that classifies educational aspects at three levels (uses, practices, and strategies) and then utilizes two frameworks, the Immersive Learning Brain and the Immersion Cube, to enable a structured description and interpretation of immersive learning cases. The method is then demonstrated on a published immersive learning case on training for wind turbine maintenance using virtual reality. Applying the method results in a structured artifact, the Immersive Learning Case Sheet, that tags the case with its proximal uses, practices, and strategies, and refines the free text case description to ensure that matching details are included. This contribution is thus a case description method in support of future comparative research of immersive learning cases. We then discuss how the resulting description and interpretation can be leveraged to change immersion learning cases, by enriching them (considering low-effort changes or additions) or innovating (exploring more challenging avenues of transformation). The method holds significant promise to support better-grounded research in immersive learning.
The technology uses reclaimed CO₂ as the dyeing medium in a closed loop process. When pressurized, CO₂ becomes supercritical (SC-CO₂). In this state CO₂ has a very high solvent power, allowing the dye to dissolve easily.
The binding of cosmological structures by massless topological defectsSérgio Sacani
Assuming spherical symmetry and weak field, it is shown that if one solves the Poisson equation or the Einstein field
equations sourced by a topological defect, i.e. a singularity of a very specific form, the result is a localized gravitational
field capable of driving flat rotation (i.e. Keplerian circular orbits at a constant speed for all radii) of test masses on a thin
spherical shell without any underlying mass. Moreover, a large-scale structure which exploits this solution by assembling
concentrically a number of such topological defects can establish a flat stellar or galactic rotation curve, and can also deflect
light in the same manner as an equipotential (isothermal) sphere. Thus, the need for dark matter or modified gravity theory is
mitigated, at least in part.
The ability to recreate computational results with minimal effort and actionable metrics provides a solid foundation for scientific research and software development. When people can replicate an analysis at the touch of a button using open-source software, open data, and methods to assess and compare proposals, it significantly eases verification of results, engagement with a diverse range of contributors, and progress. However, we have yet to fully achieve this; there are still many sociotechnical frictions.
Inspired by David Donoho's vision, this talk aims to revisit the three crucial pillars of frictionless reproducibility (data sharing, code sharing, and competitive challenges) with the perspective of deep software variability.
Our observation is that multiple layers — hardware, operating systems, third-party libraries, software versions, input data, compile-time options, and parameters — are subject to variability that exacerbates frictions but is also essential for achieving robust, generalizable results and fostering innovation. I will first review the literature, providing evidence of how the complex variability interactions across these layers affect qualitative and quantitative software properties, thereby complicating the reproduction and replication of scientific studies in various fields.
I will then present some software engineering and AI techniques that can support the strategic exploration of variability spaces. These include the use of abstractions and models (e.g., feature models), sampling strategies (e.g., uniform, random), cost-effective measurements (e.g., incremental build of software configurations), and dimensionality reduction methods (e.g., transfer learning, feature selection, software debloating).
I will finally argue that deep variability is both the problem and solution of frictionless reproducibility, calling the software science community to develop new methods and tools to manage variability and foster reproducibility in software systems.
Exposé invité Journées Nationales du GDR GPL 2024
Phenomics assisted breeding in crop improvementIshaGoswami9
As the population is increasing and will reach about 9 billion upto 2050. Also due to climate change, it is difficult to meet the food requirement of such a large population. Facing the challenges presented by resource shortages, climate
change, and increasing global population, crop yield and quality need to be improved in a sustainable way over the coming decades. Genetic improvement by breeding is the best way to increase crop productivity. With the rapid progression of functional
genomics, an increasing number of crop genomes have been sequenced and dozens of genes influencing key agronomic traits have been identified. However, current genome sequence information has not been adequately exploited for understanding
the complex characteristics of multiple gene, owing to a lack of crop phenotypic data. Efficient, automatic, and accurate technologies and platforms that can capture phenotypic data that can
be linked to genomics information for crop improvement at all growth stages have become as important as genotyping. Thus,
high-throughput phenotyping has become the major bottleneck restricting crop breeding. Plant phenomics has been defined as the high-throughput, accurate acquisition and analysis of multi-dimensional phenotypes
during crop growing stages at the organism level, including the cell, tissue, organ, individual plant, plot, and field levels. With the rapid development of novel sensors, imaging technology,
and analysis methods, numerous infrastructure platforms have been developed for phenotyping.
20240520 Planning a Circuit Simulator in JavaScript.pptx
Lewis Dot Structures - A Memory Aid -- EELS B 8!
1. A Brief Tutorial on Drawing Lewis Dot Structures
We will use three molecules (CO2, CO3
2-
and NH4
+
) as our examples in this guided tour of a simple method for
drawing Lewis dot structures. While this algorithm may not work in all cases, it should be adequate the vast
majority of the time.
Memory Aid: EELS B 8 - Electrons Electronegativity Layout Structure Bonds 8 (Octet Rule)
Procedure for Neutral Molecules (CO2)
1. Decide how many valence (outer shell) electrons are possessed by each atom in the molecule.
Memory Aid: Electrons
2. If there is more than one atom type in the molecule, put the most metallic or least electronegative atom in the
center. Recall that electronegativity decreases as atoms move further away from fluorine on the periodic chart.
Memory Aid: Electronegativity
Arrangement of atoms in CO2:
3. Arrange the electrons so that each atom contributes one electron to a single bond between each atom.
Memory Aid: Layout
4. Count the electrons around each atom: are the octets complete? If so, your Lewis dot structure is complete.
Memory Aid: Structure
5. If the octets are incomplete, and more electrons remain to be shared, move one electron per bond per atom to
make another bond. Note that in some structures there will be open octets (example: the B of BF3) or atoms
which have 10/12 electrons (example: the S of SF5 or the S of SF6). B and S are exceptions to the octet rule.
Memory Aid: Bonds
6. Repeat steps 4 and 5 as needed until all octets are full.
Memory Aid: 8 (Octet rule)
7. Redraw the dots so that electrons on any given atom are in pairs wherever possible.
2. Procedure for Negatively Charged Ions (CO3
2-)
Use the same procedure as outlined above; then as a last step, add one electron per negative charge to fill
octets. Carbonate ion has a 2- charge, so we have two electrons available to fill octets. Using the procedure
above, we arrive at this structure:
The two singly-bonded oxygen atoms each have an open octet, so we add one electron to each so as to fill these
octets. The added electrons are shown with arrows. Don't forget to assign formal charges as well! The final
Lewis structure for carbonate ion is:
Procedure for Positively Charged Ions (NH4
+)
Use the same procedure as outlined above; then remove one electron per positive charge as needed to avoid
expanded octets. When using this procedure for positively charged ions, it may be necessary to have some
atoms with expanded octets (nitrogen in this example). For each unit of positive charge on the ion, remove one
electron from these expanded octets. If done correctly, your final structure should have no first or second
period elements with expanded octets.
Using the basic procedure outlined above, we arrive at a structure in which nitrogen has nine valence
electrons. (Electrons supplied by hydrogen are red; electrons supplied by nitrogen are black.) Removal of one
of these valence electrons to account for the 1+ charge of ammonium ion solves this octet rule violation.