The document discusses different models of the atom, including Bohr's model and the quantum mechanical model. Bohr's model proposed that electrons orbit the nucleus in fixed orbits, but this was later found to be incorrect. According to the quantum mechanical model, electrons move in regions of probability called orbitals around the nucleus. Orbitals are defined using quantum numbers, which describe the electron's energy level, orbital shape, and orientation in space. Electrons are arranged into discrete energy levels and can move between levels by absorbing or emitting energy.
Rutherford's model of the atom did not fully explain the properties of elements and compounds. Bohr proposed that electrons orbit in specific energy levels, with the closest level having the lowest energy. However, Bohr's model was incomplete. Schrodinger found that electrons instead move in regions of probability called orbitals. Quantum mechanics describes electrons as existing in different energy levels and orbitals around the nucleus.
This document summarizes key concepts from a chapter on the atomic structure and quantum mechanical model of the atom. It describes early atomic models proposed by Rutherford, Bohr, and Schrodinger, and how they led to the current quantum mechanical model. It discusses how electrons occupy specific energy levels and orbitals, and how transitions between these levels result in the emission of photons of light at characteristic frequencies, producing atomic emission spectra.
This document provides information on wave quantum mechanics and electron configurations. It discusses:
- Erwin Schrodinger's contributions to developing quantum mechanics and proposing the wave-like nature of electrons.
- How electrons occupy distinct energy levels and orbitals around the nucleus, rather than defined circular orbits. Electrons have wave-like properties.
- The shapes of s, p, d and f orbitals and how electrons fill these orbitals according to various principles like Aufbau and Hund's rule.
- Exceptions to the Aufbau principle seen in some elements.
- How to represent electron configurations using both energy level diagrams and shorthand notation.
This document provides information on wave quantum mechanics and electron configurations. It discusses:
- Erwin Schrodinger's contributions to developing quantum mechanics and proposing the wave-like nature of electrons.
- How electrons occupy distinct energy levels and orbitals around the nucleus, with specific shapes defined by Schrodinger's wave equation.
- Rules for building up electron configurations, including Hund's rule and the Aufbau principle for filling orbitals in order of increasing energy.
- Exceptions to the Aufbau principle seen in some transition metals where half or fully filled subshells are more stable.
- How electron configurations are written using shorthand notation based on noble gas cores.
This document provides an overview of atomic structure and quantum mechanics. It discusses early atomic models proposed by Rutherford and Bohr and limitations they faced. It then introduces the quantum mechanical model, which describes electrons as existing in distinct energy levels and orbitals. The document explains how electron configurations are written based on Aufbau principle, Hund's rule and Pauli exclusion principle. It also discusses atomic spectra and how light emitted during electron energy level transitions can be used to identify elements.
This document provides an overview of quantum mechanics concepts related to light and atomic structure. It discusses how light behaves as both a wave and particle, and introduces the electromagnetic spectrum. It then covers atomic structure concepts like electron configurations, energy levels, quantum numbers, and orbital shapes and filling diagrams. The document aims to explain how electrons are arranged in atoms and the underlying quantum mechanical principles.
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.
This document provides information about electron configuration. It begins by defining electron configuration as the arrangement of electrons in an atom's orbitals, which is described using quantum numbers. It then discusses the three main rules for writing electron configurations: 1) Aufbau principle, which states that electrons fill the lowest available energy levels first, 2) Pauli exclusion principle, which limits each orbital to two electrons of opposite spin, and 3) Hund's rule, which states that degenerate orbitals will fill with one electron each before pairing. The document provides examples of writing full and condensed electron configurations and drawing orbital diagrams for various elements. It includes an activity for students to practice these skills.
Rutherford's model of the atom did not fully explain the properties of elements and compounds. Bohr proposed that electrons orbit in specific energy levels, with the closest level having the lowest energy. However, Bohr's model was incomplete. Schrodinger found that electrons instead move in regions of probability called orbitals. Quantum mechanics describes electrons as existing in different energy levels and orbitals around the nucleus.
This document summarizes key concepts from a chapter on the atomic structure and quantum mechanical model of the atom. It describes early atomic models proposed by Rutherford, Bohr, and Schrodinger, and how they led to the current quantum mechanical model. It discusses how electrons occupy specific energy levels and orbitals, and how transitions between these levels result in the emission of photons of light at characteristic frequencies, producing atomic emission spectra.
This document provides information on wave quantum mechanics and electron configurations. It discusses:
- Erwin Schrodinger's contributions to developing quantum mechanics and proposing the wave-like nature of electrons.
- How electrons occupy distinct energy levels and orbitals around the nucleus, rather than defined circular orbits. Electrons have wave-like properties.
- The shapes of s, p, d and f orbitals and how electrons fill these orbitals according to various principles like Aufbau and Hund's rule.
- Exceptions to the Aufbau principle seen in some elements.
- How to represent electron configurations using both energy level diagrams and shorthand notation.
This document provides information on wave quantum mechanics and electron configurations. It discusses:
- Erwin Schrodinger's contributions to developing quantum mechanics and proposing the wave-like nature of electrons.
- How electrons occupy distinct energy levels and orbitals around the nucleus, with specific shapes defined by Schrodinger's wave equation.
- Rules for building up electron configurations, including Hund's rule and the Aufbau principle for filling orbitals in order of increasing energy.
- Exceptions to the Aufbau principle seen in some transition metals where half or fully filled subshells are more stable.
- How electron configurations are written using shorthand notation based on noble gas cores.
This document provides an overview of atomic structure and quantum mechanics. It discusses early atomic models proposed by Rutherford and Bohr and limitations they faced. It then introduces the quantum mechanical model, which describes electrons as existing in distinct energy levels and orbitals. The document explains how electron configurations are written based on Aufbau principle, Hund's rule and Pauli exclusion principle. It also discusses atomic spectra and how light emitted during electron energy level transitions can be used to identify elements.
This document provides an overview of quantum mechanics concepts related to light and atomic structure. It discusses how light behaves as both a wave and particle, and introduces the electromagnetic spectrum. It then covers atomic structure concepts like electron configurations, energy levels, quantum numbers, and orbital shapes and filling diagrams. The document aims to explain how electrons are arranged in atoms and the underlying quantum mechanical principles.
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.
This document provides information about electron configuration. It begins by defining electron configuration as the arrangement of electrons in an atom's orbitals, which is described using quantum numbers. It then discusses the three main rules for writing electron configurations: 1) Aufbau principle, which states that electrons fill the lowest available energy levels first, 2) Pauli exclusion principle, which limits each orbital to two electrons of opposite spin, and 3) Hund's rule, which states that degenerate orbitals will fill with one electron each before pairing. The document provides examples of writing full and condensed electron configurations and drawing orbital diagrams for various elements. It includes an activity for students to practice these skills.
The document discusses electron configuration and the rules for filling electron orbitals in atoms. It explains that electrons exist in energy levels called shells, with each shell able to hold a maximum number of electrons according to the formula 2n2, where n is the shell number. Within shells, electrons occupy specific atomic orbitals that have distinct shapes and are grouped into sublevels. The document outlines Hund's rule and the Aufbau principle for determining the order in which orbitals are filled with electrons. Several examples of deducing full electron configurations for different elements are also provided.
The document discusses orbitals and quantum numbers. It introduces the four quantum numbers - principal (n), azimuthal (l), magnetic (ml), and spin (ms) - that describe an electron's location and properties. The classical view of electrons orbiting the nucleus like planets gave way to the modern atomic model where electrons exist as probabilistic wave functions within orbitals. Orbitals are regions of high probability of finding an electron and come in s, p, d, and f shapes depending on the quantum numbers. Electron configurations use these orbitals to describe the arrangement of electrons in atoms and predict properties.
ELECTRONIC STRUCTURE OF ATOMS AND PERIODICITY (STU).pptxiftinanwafiyah
This document provides an overview of chapter 2 from a foundation chemistry textbook. It covers the following topics:
- Atomic structure and the development of atomic models from Thomson to Bohr, including Rutherford's nuclear model and the dual nature of electrons.
- Quantum mechanics models of the atom including quantized energy levels, orbitals, and the four quantum numbers (n, l, ml, ms) that describe electrons.
- The shapes and orientations of s, p, d and f orbitals.
- Rules for writing electron configurations including the Aufbau principle and Hund's rule. Orbital diagrams are used to represent electron configurations.
- Partial orbital diagrams and condensed electron configurations are introduced
1) Electrons in atoms occupy different energy levels rather than following classical orbital models. Higher energy levels are farther from the nucleus.
2) Energy levels are divided into sublevels which have different shapes designated by letters. Electrons fill these sublevels according to specific rules.
3) The aufbau principle and Pauli exclusion principle govern how electrons fill atomic orbitals based on energy and allowing no more than two electrons of opposite spin per orbital. Hund's rule favors occupying each orbital in a sublevel singly before pairing electrons.
The document summarizes the Rutherford, Bohr, and quantum mechanical models of the atom. The Rutherford model could not explain properties of atoms and treated electrons as particles. The Bohr model proposed electrons orbiting in specific energy levels and treated electrons as particles. The quantum mechanical model, based on Schrodinger's wave equation, describes electrons as probability waves and alleviated issues with earlier models. It is the modern description of electrons' locations and energies in an atom.
The document summarizes the Rutherford, Bohr, and quantum mechanical models of the atom. The Rutherford model could not explain spectral observations. The Bohr model proposed electrons orbiting in specific energy levels and orbits, but failed for atoms with multiple electrons. The quantum mechanical model, based on Schrodinger's wave equation, describes electrons as probability waves and alleviated issues with earlier models. It uses quantum numbers like principal and angular momentum to describe electron location and energy levels.
This document discusses quantum theory and the electronic structure of atoms. It introduces quantum numbers like principal, angular momentum, and electron spin quantum numbers used to describe atomic orbitals. Atomic orbitals like s, p, and d orbitals are described along with their shapes and orientations. Electron configurations follow rules like the Aufbau principle, Pauli exclusion principle, and Hund's rule. The document shows how electrons fill atomic orbitals in order of increasing energy to write electron configurations of elements, which are represented using noble gas cores. Exceptions to electron filling order are noted for some transition metals.
1. Electrons in atoms are arranged in shells, subshells, and orbitals according to their quantum numbers. Each orbital can contain a maximum of two electrons with opposing spins.
2. Atoms experience an effective nuclear charge that increases across a period, leading to higher ionization energies and smaller atomic and ionic sizes as more protons are exposed.
3. Trends in properties like ionization energy, atomic size, and electron affinity are explained by the changing effective nuclear charge experienced by valence electrons.
This document provides an overview of atomic structure and electron configurations in chemistry. It defines key terms like atoms, electrons, energy levels, subshells and orbitals. It explains the organization of electrons according to the Aufbau principle, Hund's rule and Pauli exclusion principle. Electron configurations are represented using boxes and arrows, spectroscopic notation and noble gas notation. The document also discusses ion formations and exceptions to the rules, along with quantum numbers that describe electron location.
1) The document discusses the quantum mechanical model of the atom, including atomic orbitals and electron configurations.
2) It describes how electrons can occupy different atomic orbitals based on their energy levels and how the electron configuration notation is used to show the filling of these orbitals.
3) It also mentions exceptions to the expected electron configuration filling order that result in more stable half-filled or filled orbitals.
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.
This document provides a summary of key concepts for electron configuration in high school chemistry, including:
1) Electrons fill subshells according to the aufbau principle to achieve lowest energy, with Hund's rule specifying that electrons occupy each orbital singly before pairing up.
2) The four subshells are s, p, d, and f, with set numbers of orbitals and maximum electrons in each. Valence electrons are in the outermost shell.
3) Electron configuration can be written using boxes and arrows, spectroscopic notation, or noble gas notation, with examples provided.
The document summarizes the development of atomic models from Rutherford to the current quantum mechanical model. It discusses inadequacies in Rutherford's model that could not explain atomic properties and emissions. Bohr proposed electrons orbit in distinct energy levels, but this failed to explain multi-electron atoms. The quantum mechanical model treats electrons as waves using Schrodinger's equation, describing electrons as probability distributions rather than particles. It introduces quantum numbers to characterize electron states and explains how orbitals are filled according to various rules.
The document discusses the quantum mechanical model of the atom. It describes how Bohr's model improved on Rutherford's model by considering electrons as particles orbiting the nucleus in specific circular paths. Later, de Broglie proposed that electrons could be considered as both particles and waves. Schrodinger and Heisenberg further developed these ideas into mathematical models using quantum numbers to describe electron location and energy levels. The quantum numbers are n (principal), l (angular momentum), ml (magnetic), and ms (electron spin). Electron configuration notes how electrons fill these energy levels and orbitals according to Aufbau principle, Hund's rule, and Pauli exclusion principle.
This document provides information about electron configuration and the principles that define how electrons are arranged in an atom's orbitals. It discusses the ground-state electron configuration as the most stable, lowest-energy arrangement of electrons in an atom. Three main rules that define electron configuration are described as the Aufbau principle, Pauli exclusion principle, and Hund's rule. The document also explains how electron configuration can be written using orbital diagrams or notation, and how the noble gas notation is used.
This document discusses the quantum mechanical model of the atom. It explains that electrons can only exist in specific orbits or energy levels, and cannot exist between levels. When electrons move between levels, they absorb or emit quanta of energy. The quantum mechanical model uses mathematics to determine allowed electron energies and locations around the nucleus. Electrons occupy regions of space called atomic orbitals, and their exact locations determine an atom's properties.
The document discusses the evolution of atomic models from Dalton to Bohr, including key discoveries by Thomson, Rutherford, and Bohr. It then explains electron configurations, describing the organization of electrons into energy levels, subshells, and orbitals according to quantum numbers and rules like the Aufbau principle and Hund's rule. The document provides examples of writing electron configurations using orbital, electron-configuration, and electron-dot notations.
This document provides an overview of electronic configuration and the rules for writing the electronic configurations of elements. It defines electronic configuration as describing how electrons are distributed in an atom's atomic orbitals. It discusses the quantum numbers that describe atomic orbitals and electrons, including the principal, azimuthal, magnetic, and spin quantum numbers. It also explains the Aufbau principle, Pauli exclusion principle, and Hund's rule that govern the filling of electrons in orbitals to write electronic configurations. The goal is for students to understand these concepts and be able to write the electronic configuration of different elements.
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 defines parts per million (ppm) as the number of units of mass of a contaminant per million units of total mass. It provides an example that 1 ppm in soils and sediments equals 1 mg of substance per kg of solid. The document also presents the formula for calculating ppm and works through two practice problems calculating ppm concentrations given the mass of solute and solution. It finds that the concentration of calcium ions is 76 ppm in the first example and the concentration of ethanol is 230000 ppm in the second example.
This document discusses greenhouse gases and global climate change. It defines greenhouse gases as gases that cause the greenhouse effect and trap heat in the lower atmosphere. It then defines global climate change as identifiable changes in Earth's climate that last for decades or longer, and are usually caused by either natural processes or human activities that release greenhouse gases. The document goes on to explain that current climate changes happening include warming oceans and atmospheres and melting ice, and that these changes are extremely likely to be caused by human-caused greenhouse gases according to the IPCC. The effects of continued climate change will include more extreme weather, sea level rise, damage to ecosystems and increased species extinctions.
The document discusses electron configuration and the rules for filling electron orbitals in atoms. It explains that electrons exist in energy levels called shells, with each shell able to hold a maximum number of electrons according to the formula 2n2, where n is the shell number. Within shells, electrons occupy specific atomic orbitals that have distinct shapes and are grouped into sublevels. The document outlines Hund's rule and the Aufbau principle for determining the order in which orbitals are filled with electrons. Several examples of deducing full electron configurations for different elements are also provided.
The document discusses orbitals and quantum numbers. It introduces the four quantum numbers - principal (n), azimuthal (l), magnetic (ml), and spin (ms) - that describe an electron's location and properties. The classical view of electrons orbiting the nucleus like planets gave way to the modern atomic model where electrons exist as probabilistic wave functions within orbitals. Orbitals are regions of high probability of finding an electron and come in s, p, d, and f shapes depending on the quantum numbers. Electron configurations use these orbitals to describe the arrangement of electrons in atoms and predict properties.
ELECTRONIC STRUCTURE OF ATOMS AND PERIODICITY (STU).pptxiftinanwafiyah
This document provides an overview of chapter 2 from a foundation chemistry textbook. It covers the following topics:
- Atomic structure and the development of atomic models from Thomson to Bohr, including Rutherford's nuclear model and the dual nature of electrons.
- Quantum mechanics models of the atom including quantized energy levels, orbitals, and the four quantum numbers (n, l, ml, ms) that describe electrons.
- The shapes and orientations of s, p, d and f orbitals.
- Rules for writing electron configurations including the Aufbau principle and Hund's rule. Orbital diagrams are used to represent electron configurations.
- Partial orbital diagrams and condensed electron configurations are introduced
1) Electrons in atoms occupy different energy levels rather than following classical orbital models. Higher energy levels are farther from the nucleus.
2) Energy levels are divided into sublevels which have different shapes designated by letters. Electrons fill these sublevels according to specific rules.
3) The aufbau principle and Pauli exclusion principle govern how electrons fill atomic orbitals based on energy and allowing no more than two electrons of opposite spin per orbital. Hund's rule favors occupying each orbital in a sublevel singly before pairing electrons.
The document summarizes the Rutherford, Bohr, and quantum mechanical models of the atom. The Rutherford model could not explain properties of atoms and treated electrons as particles. The Bohr model proposed electrons orbiting in specific energy levels and treated electrons as particles. The quantum mechanical model, based on Schrodinger's wave equation, describes electrons as probability waves and alleviated issues with earlier models. It is the modern description of electrons' locations and energies in an atom.
The document summarizes the Rutherford, Bohr, and quantum mechanical models of the atom. The Rutherford model could not explain spectral observations. The Bohr model proposed electrons orbiting in specific energy levels and orbits, but failed for atoms with multiple electrons. The quantum mechanical model, based on Schrodinger's wave equation, describes electrons as probability waves and alleviated issues with earlier models. It uses quantum numbers like principal and angular momentum to describe electron location and energy levels.
This document discusses quantum theory and the electronic structure of atoms. It introduces quantum numbers like principal, angular momentum, and electron spin quantum numbers used to describe atomic orbitals. Atomic orbitals like s, p, and d orbitals are described along with their shapes and orientations. Electron configurations follow rules like the Aufbau principle, Pauli exclusion principle, and Hund's rule. The document shows how electrons fill atomic orbitals in order of increasing energy to write electron configurations of elements, which are represented using noble gas cores. Exceptions to electron filling order are noted for some transition metals.
1. Electrons in atoms are arranged in shells, subshells, and orbitals according to their quantum numbers. Each orbital can contain a maximum of two electrons with opposing spins.
2. Atoms experience an effective nuclear charge that increases across a period, leading to higher ionization energies and smaller atomic and ionic sizes as more protons are exposed.
3. Trends in properties like ionization energy, atomic size, and electron affinity are explained by the changing effective nuclear charge experienced by valence electrons.
This document provides an overview of atomic structure and electron configurations in chemistry. It defines key terms like atoms, electrons, energy levels, subshells and orbitals. It explains the organization of electrons according to the Aufbau principle, Hund's rule and Pauli exclusion principle. Electron configurations are represented using boxes and arrows, spectroscopic notation and noble gas notation. The document also discusses ion formations and exceptions to the rules, along with quantum numbers that describe electron location.
1) The document discusses the quantum mechanical model of the atom, including atomic orbitals and electron configurations.
2) It describes how electrons can occupy different atomic orbitals based on their energy levels and how the electron configuration notation is used to show the filling of these orbitals.
3) It also mentions exceptions to the expected electron configuration filling order that result in more stable half-filled or filled orbitals.
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.
This document provides a summary of key concepts for electron configuration in high school chemistry, including:
1) Electrons fill subshells according to the aufbau principle to achieve lowest energy, with Hund's rule specifying that electrons occupy each orbital singly before pairing up.
2) The four subshells are s, p, d, and f, with set numbers of orbitals and maximum electrons in each. Valence electrons are in the outermost shell.
3) Electron configuration can be written using boxes and arrows, spectroscopic notation, or noble gas notation, with examples provided.
The document summarizes the development of atomic models from Rutherford to the current quantum mechanical model. It discusses inadequacies in Rutherford's model that could not explain atomic properties and emissions. Bohr proposed electrons orbit in distinct energy levels, but this failed to explain multi-electron atoms. The quantum mechanical model treats electrons as waves using Schrodinger's equation, describing electrons as probability distributions rather than particles. It introduces quantum numbers to characterize electron states and explains how orbitals are filled according to various rules.
The document discusses the quantum mechanical model of the atom. It describes how Bohr's model improved on Rutherford's model by considering electrons as particles orbiting the nucleus in specific circular paths. Later, de Broglie proposed that electrons could be considered as both particles and waves. Schrodinger and Heisenberg further developed these ideas into mathematical models using quantum numbers to describe electron location and energy levels. The quantum numbers are n (principal), l (angular momentum), ml (magnetic), and ms (electron spin). Electron configuration notes how electrons fill these energy levels and orbitals according to Aufbau principle, Hund's rule, and Pauli exclusion principle.
This document provides information about electron configuration and the principles that define how electrons are arranged in an atom's orbitals. It discusses the ground-state electron configuration as the most stable, lowest-energy arrangement of electrons in an atom. Three main rules that define electron configuration are described as the Aufbau principle, Pauli exclusion principle, and Hund's rule. The document also explains how electron configuration can be written using orbital diagrams or notation, and how the noble gas notation is used.
This document discusses the quantum mechanical model of the atom. It explains that electrons can only exist in specific orbits or energy levels, and cannot exist between levels. When electrons move between levels, they absorb or emit quanta of energy. The quantum mechanical model uses mathematics to determine allowed electron energies and locations around the nucleus. Electrons occupy regions of space called atomic orbitals, and their exact locations determine an atom's properties.
The document discusses the evolution of atomic models from Dalton to Bohr, including key discoveries by Thomson, Rutherford, and Bohr. It then explains electron configurations, describing the organization of electrons into energy levels, subshells, and orbitals according to quantum numbers and rules like the Aufbau principle and Hund's rule. The document provides examples of writing electron configurations using orbital, electron-configuration, and electron-dot notations.
This document provides an overview of electronic configuration and the rules for writing the electronic configurations of elements. It defines electronic configuration as describing how electrons are distributed in an atom's atomic orbitals. It discusses the quantum numbers that describe atomic orbitals and electrons, including the principal, azimuthal, magnetic, and spin quantum numbers. It also explains the Aufbau principle, Pauli exclusion principle, and Hund's rule that govern the filling of electrons in orbitals to write electronic configurations. The goal is for students to understand these concepts and be able to write the electronic configuration of different elements.
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 defines parts per million (ppm) as the number of units of mass of a contaminant per million units of total mass. It provides an example that 1 ppm in soils and sediments equals 1 mg of substance per kg of solid. The document also presents the formula for calculating ppm and works through two practice problems calculating ppm concentrations given the mass of solute and solution. It finds that the concentration of calcium ions is 76 ppm in the first example and the concentration of ethanol is 230000 ppm in the second example.
This document discusses greenhouse gases and global climate change. It defines greenhouse gases as gases that cause the greenhouse effect and trap heat in the lower atmosphere. It then defines global climate change as identifiable changes in Earth's climate that last for decades or longer, and are usually caused by either natural processes or human activities that release greenhouse gases. The document goes on to explain that current climate changes happening include warming oceans and atmospheres and melting ice, and that these changes are extremely likely to be caused by human-caused greenhouse gases according to the IPCC. The effects of continued climate change will include more extreme weather, sea level rise, damage to ecosystems and increased species extinctions.
The Kinetic Molecular Theory explains the properties of solids, liquids, and gases in terms of the motion and interaction of their particles. It states that all matter is made up of tiny particles that are constantly moving, with the speed and freedom of motion of the particles determining whether a substance is a solid, liquid, or gas. Changes in temperature can cause a substance to change states by altering the kinetic energy and interactions between its particles.
This document provides information about volcanoes and volcanic eruptions. It begins with a pre-assessment quiz about volcanic characteristics and eruptions. It then discusses the different types of volcanoes including shield, cinder cone, and composite volcanoes. The document outlines the primary factors that affect volcanic eruptive styles such as magma temperature, composition, and gas content. It also describes the different types of eruptions from phreatic to plinian. Finally, it discusses how volcanoes can be used as sources of geothermal energy and provides signs of an impending volcanic eruption.
Organic chemistry is the study of carbon-containing compounds. Hydrocarbons only contain carbon and hydrogen, forming chains and branches with single, double, or triple bonds between carbons. Functional groups include alcohols, acids, esters, ethers, amines, amides, ketones, and aldehydes. Organic reactions include substitution, addition, fermentation, saponification, polymerization, esterification, combustion, cracking, and fractional distillation. Hydrocarbons have low melting points and are nonpolar and nonconductive.
This chapter discusses organic compounds and carbon chemistry. It covers the properties of carbon that allow it to form large, complex molecules through catenation. The structures and classes of hydrocarbons like alkanes, alkenes, and alkynes are examined. Important classes of organic reactions such as addition, elimination, and substitution are described. The chapter also explores functional groups and how they determine a molecule's reactivity and properties. Alcohols are highlighted as one functional group and their naming conventions and intermolecular hydrogen bonding are discussed.
1. There are three classes of strong electrolytes: strong acids, strong bases, and most water soluble salts. Weak acids and bases only partially dissociate in water.
2. pH is a measure of the concentration of hydrogen ions [H+] in a solution. Low pH indicates high [H+] and an acidic solution, while high pH indicates low [H+] and a basic solution. Household substances like coffee, milk, and baking soda have different pH values.
3. The acid dissociation constant Ka and base dissociation constant Kb are equilibrium constants that indicate the strength of an acid or base. Strong acids and bases fully dissociate while weak acids and bases only partially dissociate,
There are 8 key characteristics that define life:
1. Living things are made of cells, either unicellular or multicellular.
2. Living things reproduce, either sexually requiring two parents or asexually with one parent.
3. Living things are based on a universal genetic code stored in DNA that is inherited from parents.
4. Living things grow and develop over their lifetime.
This document discusses chemical bonding and Lewis dot structures. It explains that atoms combine to achieve stable electron configurations, often those of noble gases. Ionic bonds form when electrons are transferred between atoms, creating ions. Covalent bonds form through electron sharing between nonmetals. Lewis dot structures represent valence electrons and can show electron transfers in ion formation. Different bond types - ionic, covalent, and metallic - depend on differences in electronegativity between atoms. Practice problems are provided to determine bond type based on electronegativity values.
This document discusses molecular polarity and how to determine if molecules are polar or nonpolar. It defines polar and nonpolar covalent bonds based on differences in electronegativity between atoms. A molecule is polar if it contains polar bonds in an asymmetrical arrangement, whereas a molecule is nonpolar if the polar bonds are symmetrical or if all bonds are nonpolar. Examples of polar molecules include HCl and H2O, while nonpolar examples include CO2 and CF4.
1. Early atomic models proposed by philosophers like Democritus were later developed into scientific theories by scientists like Dalton, Thomson, Rutherford, and Bohr based on new experimental evidence.
2. Dalton's early billiard ball model of atoms was revised after Thomson discovered the electron and Rutherford found evidence of a dense nucleus through his gold foil experiment.
3. The current quantum mechanical model sees electrons as existing as probabilistic clouds around the nucleus rather than distinct orbits, though early models by Bohr still explained much of chemistry.
12. Powering the Cell-Cellular Respiration.pptLiezlValiente1
Cellular respiration is the process by which living cells convert glucose into energy in the form of ATP. It occurs in three main stages: glycolysis, the Krebs cycle, and the electron transport chain. Glycolysis breaks down glucose and produces a small amount of ATP. In the presence of oxygen, the Krebs cycle and electron transport chain further break down pyruvate from glycolysis to produce much more ATP. Without oxygen, fermentation pathways like lactic acid fermentation or alcoholic fermentation produce ATP and regenerate NAD+ to allow glycolysis to continue. Aerobic respiration is much more efficient at producing ATP.
The document discusses sound and how it is produced, transmitted, and heard. It describes how sound is a longitudinal wave that travels through matter by causing compressions and rarefactions in molecules. The speed of sound depends on the temperature, density, and elasticity of the medium, being fastest in solids and slowest in gases. It also discusses how the human ear detects sound waves, with the outer ear funneling sound to the eardrum, middle ear bones vibrating the fluid-filled inner ear, and inner ear hair cells converting this into electrical signals sent to the brain.
The circulatory system transports oxygen, nutrients, hormones, and antibodies throughout the body while removing waste products such as carbon dioxide. It consists of the heart, blood vessels, and blood. The heart pumps blood through three types of circulation - systemic, coronary, and pulmonary. Blood travels from the heart through arteries, to capillaries where exchange occurs, and returns to the heart via veins. The circulatory system is vital for sustaining life.
Biogeochemical cycles describe the movement of elements and molecules through biotic and abiotic components of ecosystems. Key cycles include carbon, nitrogen, oxygen, phosphorus, and sulfur. In these cycles, matter is transferred between living organisms, non-living matter like soil and water, and long-term stores like fossil fuels and sedimentary rocks. The recycling of nutrients through biogeochemical cycles is essential for sustaining life on Earth.
This document provides an overview of magma, volcanoes, and volcanic eruptions. It discusses the following key points in 3 sentences:
Magma is molten rock beneath the Earth's surface that rises towards the surface through vents called volcanoes. There are different types of volcanoes that produce eruptions ranging from gentle flows to catastrophic explosions, depending on the viscosity and gas content of the magma. The composition and viscosity of magmas influence the type of eruption, whether nonexplosive eruptions producing lava flows or explosive eruptions ejecting tephra and forming eruption columns and pyroclastic flows.
Seismic waves from earthquakes and explosions allow scientists to map the interior of Earth. Layers are identified by how fast p-waves and s-waves travel through materials with different densities and states. The crust is thin and varies in thickness and composition between continents and oceans. The mantle below is hot and convects slowly. The outer core is liquid and the inner core is solid, and their rotation generates Earth's magnetic field.
The document outlines the schedule and topics for a mid-year in-service training for teachers from February 2-4, 2022. On the first day, there will be presentations on Covid-19 updates and health status, as well as Filipino orthography. The second day will involve workshops for crafting project proposals, action plans, and monitoring/evaluation tools. On the final day, teachers will present and critique their outputs and receive certificates. The training aims to provide professional development on various teaching-related topics over the course of three days.
This document discusses two early atomic models:
1) J.J. Thomson's "plum pudding" model from 1897 which viewed the atom as a uniform positively charged sphere with electrons embedded inside like raisins in a pudding.
2) Ernest Rutherford's gold foil experiment from 1908 which found that most alpha particles passed through a gold foil with little deflection, but some bounced off at sharp angles, indicating the positive charge of atoms must be concentrated in a small, dense nucleus.
3) Rutherford concluded atoms have a small, dense positively charged nucleus surrounded by electrons, overturning the plum pudding model - this came to be called the Rutherford model of the atom.
This document outlines a demonstration teaching session on electricity generation, transmission, and distribution that was presented at Bangui National High School - Main Campus in Bangui, Ilocos Norte, Philippines. The presentation covered how electrical energy is generated at power plants and transmitted through power towers, transmission lines, and step-up transformers before being distributed to neighborhoods through step-down transformers. It also engaged students in activities to identify different electrical sources like solar, wind, and hydro power plants and explain the components involved in transporting electricity from generators to consumers.
How to Build a Module in Odoo 17 Using the Scaffold MethodCeline George
Odoo provides an option for creating a module by using a single line command. By using this command the user can make a whole structure of a module. It is very easy for a beginner to make a module. There is no need to make each file manually. This slide will show how to create a module using the scaffold method.
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.
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Exploiting Artificial Intelligence for Empowering Researchers and Faculty, In...Dr. Vinod Kumar Kanvaria
Exploiting Artificial Intelligence for Empowering Researchers and Faculty,
International FDP on Fundamentals of Research in Social Sciences
at Integral University, Lucknow, 06.06.2024
By Dr. Vinod Kumar Kanvaria
The simplified electron and muon model, Oscillating Spacetime: The Foundation...RitikBhardwaj56
Discover the Simplified Electron and Muon Model: A New Wave-Based Approach to Understanding Particles delves into a groundbreaking theory that presents electrons and muons as rotating soliton waves within oscillating spacetime. Geared towards students, researchers, and science buffs, this book breaks down complex ideas into simple explanations. It covers topics such as electron waves, temporal dynamics, and the implications of this model on particle physics. With clear illustrations and easy-to-follow explanations, readers will gain a new outlook on the universe's fundamental nature.
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.
How to Manage Your Lost Opportunities in Odoo 17 CRMCeline George
Odoo 17 CRM allows us to track why we lose sales opportunities with "Lost Reasons." This helps analyze our sales process and identify areas for improvement. Here's how to configure lost reasons in Odoo 17 CRM
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 slide is special for master students (MIBS & MIFB) in UUM. Also useful for readers who are interested in the topic of contemporary Islamic banking.
Main Java[All of the Base Concepts}.docxadhitya5119
This is part 1 of my Java Learning Journey. This Contains Custom methods, classes, constructors, packages, multithreading , try- catch block, finally block and more.
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.
2. Section 7.5
The Quantum Mechanical Model of the Atom
Return to TOC
BOHR’S THEORY
Electrons are located at
specific energy levels
surrounding the nucleus
Each rung on the ladder
represents an energy
level
The higher the energy
level – the farther it is
from the nucleus
Bohr thought the electrons moved in
fixed ORBITS around the nucleus –
we know this is not true today
3. Section 7.5
The Quantum Mechanical Model of the Atom
Return to TOC
BOHR MODEL
First model of the electron structure
Gives levels where an electron is most likely to be found
Incorrect today, but a key in understanding the atom
3
4. Section 7.4
The Bohr Model
Return to TOC
4
• Bohr’s model gave hydrogen atom energy levels
consistent with the hydrogen emission spectrum.
• Ground state – lowest possible energy state (n = 1)
• Bohr’s model is incorrect. This model only works for
hydrogen.
• Electrons do not move around the nucleus in
circular orbits.
Electronic
Transitions in the
Bohr Model for the
Hydrogen Atom
a) An Energy-Level
Diagram for Electronic
Transitions
Electronic Transitions
in the Bohr Model for
the Hydrogen Atom
b) An Orbit-
Transition Diagram,
Which Accounts for
the Experimental
Spectrum
5. Section 7.5
The Quantum Mechanical Model of the Atom
Return to TOC
SCHRÖDINGER'S THEORY
He agreed that electrons
have a specific amount of
energy
He believed that the
distance between rungs on
the ladder were not
consistent – they get closer
together as you move
higher up
Quantum – the amount of
energy needed to move
from one energy level to
another
The electrons
move in regions of
probability
around the
nucleus called
ORBITALS
6. Section 7.5
The Quantum Mechanical Model of the Atom
Return to TOC
Quantum theory, also called wave mechanics,
describes the arrangement and space occupied
by electrons. Orbitals refers to the three-
dimensional regions in space where there is a high
probability of finding an electron around an atom.
6
7. Section 7.5
The Quantum Mechanical Model of the Atom
Return to TOC
CHARACTERISTICS OF ELECTRONS
Extremely small mass
Located outside the nucleus
Moving at extremely high speeds in a
sphere
Have specific energy levels
7
8. Section 7.5
The Quantum Mechanical Model of the Atom
Return to TOC
ENERGY OF ELECTRONS
When atoms are heated, bright lines appear
called line spectra
Electrons in atoms arranged in discrete
levels.
An electron absorbs energy to “jump” to a
higher energy level.
When an electron falls to a lower energy level,
energy is emitted.
8
9. Section 7.5
The Quantum Mechanical Model of the Atom
Return to TOC
LOSS AND GAIN OF ENERGY
9
G
a
I
n
L
o
s
s
10. Section 7.5
The Quantum Mechanical Model of the Atom
Return to TOC
LEARNING CHECK
Answer with
1) Energy absorbed 2) Energy emitted
3) No change in energy
A. What energy change takes place when an
electron in a hydrogen atom moves from the
first (n=1) to the second shell (n=2)?
B. What energy change takes place when the
electron moves from the third shell to the
second shell?
10
11. Section 7.5
The Quantum Mechanical Model of the Atom
Return to TOC
SOLUTION
A. 1) Energy absorbed
B. 2) Energy emitted
11
12. Section 7.5
The Quantum Mechanical Model of the Atom
Return to TOC
RELATIVE ORBITAL SIZE
Difficult to define precisely.
Orbital is a wave function.
Picture an orbital as a three-dimensional electron
density map.
Hydrogen 1s orbital:
Radius of the sphere that encloses 90% of the
total electron probability. 12
• We do not know the detailed pathway of an electron.
• The electrons move in regions of probability around the
nucleus called ORBITALS
13. Section 7.5
The Quantum Mechanical Model of the Atom
Return to TOC
THE ELECTRONS MOVE IN REGIONS OF PROBABILITY
AROUND THE NUCLEUS CALLED ORBITALS
DEFINING THESE ORBITALS:
Quantum Numbers are used to define:
The energy of the electron
The electron’s relative distance from the nucleus
The size and shape of the ORBITAL
The pairings of the electrons
14. Section 7.5
The Quantum Mechanical Model of the Atom
Return to TOC
QUANTUM NUMBERS
Principle Quantum Number (n) – define the
energy of the electron
n=1 is closest to the nucleus – low energy
n=2 is farther than n=1, slightly more energy
n=3 is farther than n=1 and n=2, still
increasing in energy
n=4 …..
Remember – The difference in energy
between energy levels decreases as “n”
increases
15. Section 7.5
The Quantum Mechanical Model of the Atom
Return to TOC
SUBLEVELS
Within each principle energy level (n) – there are
sublevel(s).
The larger the value of ‘n’, the more sublevels you
can have.
Sublevels – named by their shape
s – sphere p – pear
d- dumbbell f - fundamental
18. Section 7.5
The Quantum Mechanical Model of the Atom
Return to TOC
THERE IS A SET OF FIVE DIFFERENT D ORBITALS.
THERE IS A SET OF SEVEN F ORBITALS.
EACH ORBITAL REGARDLESS OF ITS SHAPE HOLDS 2 ELECTRONS.
19. Section 7.7
Orbital Shapes and Energies
Return to TOC
19
Two Representations
of the Hydrogen 1s,
2s, and 3s Orbitals
20. Section 7.7
Orbital Shapes and Energies
Return to TOC
20
The Boundary Surface Representations of All Three 2p Orbitals
21. Section 7.7
Orbital Shapes and Energies
Return to TOC
21
The Boundary Surfaces of All of the 3d Orbitals
22. Section 7.7
Orbital Shapes and Energies
Return to TOC
22
Representation of the 4f Orbitals in Terms of Their Boundary
Surfaces
23. Section 7.7
Orbital Shapes and Energies
Return to TOC
• As shown in Table, the s subshell has one lobe, the p
subshell has three lobes, the d subshell has five lobes,
and the f subshell has seven lobes. Each of these lobes
is labeled differently and is named depending on which
plane the lobe is resting in. If the lobe lies along the x
plane, then it is labeled with an x, as in 2px. If the lobe
lies along the xy plane, then it is labeled with a xy such
as dxy. Electrons are found within the lobes. The plane
(or planes) that the orbitals do not fill are called nodes.
These are regions in which there is a 0 probability
density of finding electrons.
23
25. Section 7.5
The Quantum Mechanical Model of the Atom
Return to TOC
SUBLEVELS
Principle Energy Level Sublevel
n= 1 s
n=2 s and p
n=3 s and p and d
n=4 s, p, d, and f
NOTICE: The value of ‘n’ tells you how many
sublevels are present in that energy level
26. Section 7.6
Quantum Numbers
Return to TOC
26
• Principal quantum number (n) – size and energy
of the orbital.
• Angular momentum quantum number (l) – shape
of atomic orbitals (sometimes called a subshell).
• Magnetic quantum number (ml) – orientation of
the orbital in space relative to the other orbitals
in the atom.
27. QUANTUM NUMBERS FOR THE FIRST FOUR LEVELS OF ORBITALS IN
THE HYDROGEN ATOM
27
28. Section 7.6
Quantum Numbers
Return to TOC
28
Exercise
For principal quantum level n = 3,
determine the number of allowed subshells
(different values of l), and give the
designation of each. (hint refer to previous
chart)
# of allowed subshells = 3
l = 0, 3s
l = 1, 3p
l = 2, 3d
29. Section 7.6
Quantum Numbers
Return to TOC
29
Exercise
For l = 2, determine the magnetic quantum
numbers (ml) and the number of orbitals.
(note refer to previous chart)
magnetic quantum numbers = –2, – 1, 0, 1, 2
number of orbitals = 5
30. Section 7.7
Orbital Shapes and Energies
Return to TOC
Locating these on the Periodic Table
Principle Energy Level (n) – is the period in the periodic
table
The Sublevels are located in specific regions – Color these
together
31. Section 7.7
Orbital Shapes and Energies
Return to TOC
31
• The periodic table is structured so that elements with the same type of valence electron
configuration are arranged in columns.
• The left-most columns include the alkali metals and the alkaline earth metals. In these elements
the valence s orbitals are being filled
• On the right hand side, the right-most block of six elements are those in which the valence p
orbitals are being filled
• In the middle is a block of ten columns that contain transition metals. These are elements in which
d orbitals are being filled
• Below this group are two rows with 14 columns. These are commonly referred to the f-block
metals. In these columns the f orbitals are being filled
32. Section 7.7
Orbital Shapes and Energies
Return to TOC
32
• The periodic table is structured so that elements with the same type of valence electron
configuration are arranged in columns.
Important facts to remember:
• 2, 6, 10 and 14 are the number of electrons that can fill the s, p, d and f subshells
(the l=0,1,2,3 azimuthal quantum number)
• The 1s subshell is the first s subshell, the 2p is the first p subshell
• 3d is the first d subshell, and the 4f is the first f subshell
33. Section 7.7
Orbital Shapes and Energies
Return to TOC
Naming the sublevels
1s
2s 2p
3s 3p 3d
4s 4p 4d 4f
34. Section 7.7
Orbital Shapes and Energies
Return to TOC
Orbitals
Orbitals are regions of probability – each orbital can hold a
maximum of 2 e-
The ‘s’ sublevel has 1 orbital
The ‘p’ sublevel has 3 orbitals
The ‘d’ sublevel has 5 orbitals
The ‘f’ sublevel has 7 orbitals
35. Section 7.7
Orbital Shapes and Energies
Return to TOC
Orbitals
Do you have to memorize this?
NO
Look at the sublevel regions that you colored in on your
periodic table.
36. Section 7.7
Orbital Shapes and Energies
Return to TOC
Orbitals
Count how many electrons are in the ‘s’ sublevel
2
This means that since there are two electrons, and each
orbital can hold two electrons, that there is only ONE
orbital.
37. Section 7.7
Orbital Shapes and Energies
Return to TOC
Orbitals
Count how many electrons are in the ‘p’ sublevel
6
This means that since there are six electrons, and each
orbital can hold two electrons, that there are THREE
orbitals.
38. Section 7.7
Orbital Shapes and Energies
Return to TOC
Orbitals
Count how many electrons are in the ‘d’ sublevel
10
This means that since there are ten electrons, and each
orbital can hold two electrons, that there are FIVE
orbitals.
39. ORBITALS
Count how many electrons are in the ‘f’ sublevel
14
This means that since there are fourteen electrons,
and each orbital can hold two electrons, that there
are SEVEN orbitals.