Chapter 5, extra notes
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The document discusses the Pauli exclusion principle and electron configurations. It explains that according to the Pauli exclusion principle, no two electrons can have the same four quantum numbers. This means that within one orbital, electrons must have opposite spin, and one orbital can hold a maximum of two electrons. The document then discusses how electrons fill atomic orbitals according to the aufbau principle, from lowest to highest energy in a specific order, and provides examples of writing electron configurations.
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How To Write Electron Configurations
1. Determine the total number of electrons by finding the element's atomic number on the periodic table.
2. Fill the lowest energy electron shells first, following the order of filling: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, etc.
3. Write the electron configuration by specifying the energy level, subshell, and number of electrons in each subshell using superscripts. The valence electrons are those in the highest occupied shell.
Electron configuration
The document discusses electron configurations, which describe the arrangement of electrons in an atom's orbitals. It notes that electrons occupy specific orbitals and that orbitals are grouped into shells. It then provides examples of the full and abbreviated electron configurations for the first 20 elements, showing how many electrons occupy each orbital. The document also outlines the steps for writing abbreviated electron configurations using the noble gas of the same period.
Chapter 8 electron configuration and periodicity (1)
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.
Echon (electron configuration)
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.
Orbital Notation
The document discusses orbital notation, which visually represents electron configuration by showing the placement and spin of each electron in an atom. It defines key terms like Pauli exclusion principle, Aufbau principle, and Hund's rule that govern how electrons fill atomic orbitals. Examples are given of writing the electron configurations of sodium, boron, titanium, and calcium using orbital notation with half arrows to symbolize individual electrons.
How to write the lewis structure of carbon disulfide
The document provides steps to write the Lewis structure for carbon disulfide (CS2):
1. Determine the number of valence electrons for each atom (C has 4, S has 6). Total valence electrons for CS2 is 10.
2. Draw an initial sketch showing the atoms. Place each atom's valence electrons around it.
3. Form single bonds between atoms by pairing their valence electrons. Additional electrons are needed to complete the octet of each sulfur atom.
4. The final Lewis structure shows each carbon and sulfur atom with 8 electrons, completing their octets.
2016 Topic 2: Electron Configuration
This document discusses electron configuration and how it is described using quantum numbers. It explains the main concepts like energy shells (n), subshells (s, p, d, f), orbitals and how electrons fill these according to the Aufbau principle and Hund's rule. Examples are provided to show the electron configurations of elements like hydrogen, helium, lithium and how the configurations change as the atomic number increases. Practice problems are included at the end to determine configurations of additional elements.
Periodic Table4
The document summarizes key concepts about the periodic table including its organization based on electron configuration and blocks (s, p, d, f), rules for filling orbitals like the Pauli exclusion principle and Hund's rule, and examples of writing electronic configurations for elements like sodium, neon, aluminum, argon, and potassium.
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How To Write Electron Configurations
1. Determine the total number of electrons by finding the element's atomic number on the periodic table.
2. Fill the lowest energy electron shells first, following the order of filling: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, etc.
3. Write the electron configuration by specifying the energy level, subshell, and number of electrons in each subshell using superscripts. The valence electrons are those in the highest occupied shell.
Electron configuration
The document discusses electron configurations, which describe the arrangement of electrons in an atom's orbitals. It notes that electrons occupy specific orbitals and that orbitals are grouped into shells. It then provides examples of the full and abbreviated electron configurations for the first 20 elements, showing how many electrons occupy each orbital. The document also outlines the steps for writing abbreviated electron configurations using the noble gas of the same period.
Chapter 8 electron configuration and periodicity (1)
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.
Echon (electron configuration)
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.
Orbital Notation
The document discusses orbital notation, which visually represents electron configuration by showing the placement and spin of each electron in an atom. It defines key terms like Pauli exclusion principle, Aufbau principle, and Hund's rule that govern how electrons fill atomic orbitals. Examples are given of writing the electron configurations of sodium, boron, titanium, and calcium using orbital notation with half arrows to symbolize individual electrons.
How to write the lewis structure of carbon disulfide
The document provides steps to write the Lewis structure for carbon disulfide (CS2):
1. Determine the number of valence electrons for each atom (C has 4, S has 6). Total valence electrons for CS2 is 10.
2. Draw an initial sketch showing the atoms. Place each atom's valence electrons around it.
3. Form single bonds between atoms by pairing their valence electrons. Additional electrons are needed to complete the octet of each sulfur atom.
4. The final Lewis structure shows each carbon and sulfur atom with 8 electrons, completing their octets.
2016 Topic 2: Electron Configuration
This document discusses electron configuration and how it is described using quantum numbers. It explains the main concepts like energy shells (n), subshells (s, p, d, f), orbitals and how electrons fill these according to the Aufbau principle and Hund's rule. Examples are provided to show the electron configurations of elements like hydrogen, helium, lithium and how the configurations change as the atomic number increases. Practice problems are included at the end to determine configurations of additional elements.
Periodic Table4
The document summarizes key concepts about the periodic table including its organization based on electron configuration and blocks (s, p, d, f), rules for filling orbitals like the Pauli exclusion principle and Hund's rule, and examples of writing electronic configurations for elements like sodium, neon, aluminum, argon, and potassium.
Electron configuration
Electron configurations provide information about the location of electrons in an atom's orbitals. There are three main rules for determining electron configurations: 1) electrons fill orbitals starting with the lowest energy levels and moving upwards, 2) the Pauli exclusion principle states that no two electrons can occupy the same orbital with the same spin, and 3) Hund's rule states that electrons will occupy orbitals singly before pairing up. Writing out electron configurations involves determining the element's number of electrons and following the Aufbau principle of filling orbitals in order of increasing energy.
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This document discusses electron configurations and how they are written. It explains that electron configurations tell us where electrons are located in orbitals, and outlines the Aufbau principle, Pauli exclusion principle, and Hund's rule for filling orbitals. It then provides examples of writing out full and abbreviated noble gas electron configurations, identifying valence electrons, writing electron dot formulas, and predicting ionic charges.
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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.
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Electron Configuration
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.
Electron configuration
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.
What is electronic configuration
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.
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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.
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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.
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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.
Ch. 7 electron configuration
- 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.
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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.
Electron configuration
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
Auger spectroscopy
1. Anger spectroscopy involves removing an electron from the K shell of an atom, creating a vacancy that is filled by an electron from the L shell. This releases energy that is used to eject an electron from the L shell.
2. The ejected electron from the L shell is the Auger electron, which provides information about the atom's electronic configuration.
3. Examples are provided of Auger spectroscopy involving thiosulfate and sulfur dioxide molecules, where the electronic transitions and resulting spectra are explained.
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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.
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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.
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This document discusses electron configurations and how they are written. It explains that electron configurations tell us where electrons are located in orbitals, and outlines the Aufbau principle, Pauli exclusion principle, and Hund's rule for filling orbitals. It then provides examples of writing out full and abbreviated noble gas electron configurations, identifying valence electrons, writing electron dot formulas, and predicting ionic charges.
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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.
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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.
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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.
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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.
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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.
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The document discusses electron configurations, which describe how electrons are distributed in atomic orbitals. It explains the Aufbau principle, which states that electrons fill lower energy orbitals first. The Pauli exclusion principle is described, stating that no more than two electrons can occupy any single orbital. Hund's rule is also covered, regarding the filling of degenerate orbitals. Examples are provided to illustrate these principles.
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Hund's rules describe the ground state of multielectron atoms and ions. Rule 1 states that the atomic state with the highest total spin S will have the lowest energy. Rule 2 states that for a given spin multiplicity, the term with the largest orbital angular momentum L has the lowest energy. Rule 3 considers energy shifts due to spin-orbit coupling, where the splitting is determined by the spin-orbit coupling constant λ and the quantum numbers J, L, and S. Examples are provided to illustrate the application of Hund's rules to determine the ground states of silicon and titanium.
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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.
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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.
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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.
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- 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.
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- A noble gas configuration can be used to abbreviate long electron configurations by writing the nearest noble gas followed by the remaining electrons.
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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.
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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
Auger spectroscopy
1. Anger spectroscopy involves removing an electron from the K shell of an atom, creating a vacancy that is filled by an electron from the L shell. This releases energy that is used to eject an electron from the L shell.
2. The ejected electron from the L shell is the Auger electron, which provides information about the atom's electronic configuration.
3. Examples are provided of Auger spectroscopy involving thiosulfate and sulfur dioxide molecules, where the electronic transitions and resulting spectra are explained.
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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.
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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.
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Electron Configurations in Science Education and Chemistry .ppt
You might have heard of the wise saying that
'experience is the best teacher". Each
year in your life ushers in a lot of experiences for you
to reflect on. You may not like everything
that happens to you in school or in the community,
especially those that bring you discomfort,
dificulty or detriment, but you have to bear with these
occurrences with a positive disposition.
You have to remember that you cannot prevent
circumstances from happening
especially those that might challenge your patience
determination and drive as a young
learner. lt's good to remember that experiences
whether in school or in the community, will
open opportunities for you to gain lessons which you
can utilize to help and inspire yourself
and others. Your negative or positive personal
experiences coupled with your coping skills can
serve as your stepping stones to academic successYou might have heard of the wise saying that
'experience is the best teacher". Each
year in your life ushers in a lot of experiences for you
to reflect on. You may not like everything
that happens to you in school or in the community,
especially those that bring you discomfort,
dificulty or detriment, but you have to bear with these
occurrences with a positive disposition.
You have to remember that you cannot prevent
circumstances from happening
especially those that might challenge your patience
determination and drive as a young
learner. lt's good to remember that experiences
whether in school or in the community, will
open opportunities for you to gain lessons which you
can utilize to help and inspire yourself
and others. Your negative or positive personal
experiences coupled with your coping skills can
serve as your stepping stones to academic success You might have heard of the wise saying that
'experience is the best teacher". Each
year in your life ushers in a lot of experiences for you
to reflect on. You may not like everything
that happens to you in school or in the community,
especially those that bring you discomfort,
dificulty or detriment, but you have to bear with these
occurrences with a positive disposition.
You have to remember that you cannot prevent
circumstances from happening
especially those that might challenge your patience
determination and drive as a young
learner. lt's good to remember that experiences
whether in school or in the community, will
open opportunities for you to gain lessons which you
can utilize to help and inspire yourself
and others. Your negative or positive personal
experiences coupled with your coping skills can
serve as your stepping stones to academic successYou might have heard of the wise saying that
'experience is the best teacher". Each
year in your life ushers in a lot of experiences for you
to reflect on. You may not like everything
that happens to you in school or in the community,
especially those that bring you discomfort,
dificulty or detriment, but you have to bear with these
occurrences with a positive disposition.
You have to remember h
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- 2. Pauli’s Exclusion Principle Pauli Exclusion Principle No two electrons have the same quantum number. (opposite spin) Maximum electrons in any orbital is two ()
- 3. The Pauli exclusion principle summarizes experimental observations that no two electrons in one atom can have the same four quantum numbers. That means that within one orbital, electrons must have opposite spin. It also means that one orbital can hold a maximum of two electrons (with opposite spin). 8 | 3
- 4. An s sublevel, with one orbital, can hold a maximum of 2 electrons. A p sublevel, with three orbitals, can hold a maximum of 6 electrons. A d sublevel, with five orbitals, can hold a maximum of 10 electrons. An f sublevel, with seven orbitals, can hold a maximum of 14 electrons. 8 | 4
- 5. The building-up principle (or aufbau principle) is a scheme used to reproduce the ground-state electron configurations by successively filling sublevels with electrons in a specific order (the building-up order). This order generally corresponds to filling the orbitals from lowest to highest energy. Note that these energies are the total energy of the atom rather than the energy of the sublevels alone. 8 | 5
- 6. Hund’s Rule Hund’s Rule When filling degenerate orbital's, electrons will fill an empty orbital before pairing up with another electron. RIGHT WRONG
- 7. Electron Configuration ______ ______ ______ ______ ______ Aufbau Diagram IncreasingEnergy Electron Spin 2p61s2 2s2 Electron Configuration Electron configuration for Neon 2 p 2 s 1 s
- 8. Electron Configuration Aufbau diagram shows each orbital O (atomic number 8) ____ ____ ____ ____ ____ 2s 2p 1s electron configuration 1s2 2s2 2p4
- 9. Write the electron configuration for Sulfur by using: S (atomic number 16) ____ ____ ____ ____ ____ ____ ____ ____ 3s 3p ____ 2s 2p 1s 1s2 2s2 2p6 3s2 3p4 How many unpaired electrons does sulfur have? 2 unpaired electrons! Aufbau Diagram Electron configuration
- 10. 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 7s 7p This results in the following order: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p
- 11. Electron Configuration Let’s Practice P (atomic number 15) 1s2 2s2 2p6 3s2 3p3 Ca (atomic number 20) 1s2 2s2 2p6 3s2 3p6 4s2 As (atomic number 33) 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p3 W (atomic number 74) 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14 5d4
- 12. Electron Configuration Your Turn N (atomic number 7) 1s2 2s2 2p3 Na (atomic number 11) 1s2 2s2 2p6 3s1 Sb (atomic number 51) 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p3 Cr (atomic number 24) 1s2 2s2 2p6 3s2 3p6 4s2 3d4
- 13. The lowest-energy configuration of an atom is called its ground state. Any other configuration represents an excited state. 8 | 13
- 14. Exceptions Copper Expect: 1s2 2s2 2p6 3s2 3p6 4s2 3d9 Actual: 1s2 2s2 2p6 3s2 3p6 4s1 3d10 Silver Expect: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d9 Actual: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s1 4d10 Chromium Expect: 1s2 2s2 2p6 3s2 3p6 4s2 3d4 Actual: 1s2 2s2 2p6 3s2 3p6 4s1 3d5 Molybdenum Expect: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d4 Actual: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s1 4d5 Exceptions are explained, but not predicted! Atoms are more stable with half full sublevel
- 15. Another way to learn the building-up order is to correlate each sublevel with a position on the periodic table. The principal quantum number, n, correlates with the period number. Groups IA and IIA correspond to the s sublevel; Groups IIIA through VIIIA correspond to the p sublevel; the “B” groups correspond to the d sublevel; and the bottom two rows correspond to the f sublevel. This is shown on the next slide. 8 | 15
- 16. 8 | 16 Energy level equals period number Energy level equals period number minus one Energy level equals period number minus two 1 2 3 4 5 6 7
- 17. Electron Filling in Periodic Table 1 2 3 4 5 6 7 s d p s f * W W *
- 18. Periodic Table and Electron Configuration Using the periodic table for the filling order of orbitals, by going in atomic number sequence until you use all the needed electrons in the element