Unit IV: The Electronic Structure of the Atom and the Periodic Table PDF

**Unit IV: The Electronic Structure of the Atom and the Periodic Table** **Introduction** Our significant means of understanding an atom was brought about in the 20th century, from the proposal of Ernest Rutherford planetary model to Niels Bohr's application of quantum theory and waves to the beha...

**Unit IV: The Electronic Structure of the Atom and the Periodic Table** **Introduction** Our significant means of understanding an atom was brought about in the 20th century, from the proposal of Ernest Rutherford planetary model to Niels Bohr's application of quantum theory and waves to the behavior of electrons. Electron configuration can be described as the arrangement of electrons within the orbitals shells and subshells of an atom. It is important to understand the behavior of an electron in atom in order to fully comprehend the electron configuration. An electron is a subatomic particle that is associated with a negative charge. Additionally, electrons are associated with energy, more specifically quantum energy, and show wave-like and particle-like characteristics. Fully understanding the principles relating to electron configuration will promote a better comprehension of how to arrange them and give us a better understanding of each element in the periodic table. The arrangement of the elements in the periodic table has an especial correlation with electron configuration. After studying the relationship between electron configuration and the period table, it was pointed out by Niels Bohr that electron configurations are similar for elements within the same group or family in the periodic table. Groups occupy the vertical rows as opposed to a period which is the horizontal rows in the table of elements. The periodic recurrence of similar properties of elements is one aspect of periodicity. Another aspect consists of the trends of in the properties of elements and their compounds.  **Learning Outcomes** At the end of this unit you should be able to: 1. Describe and explain the light emitted by the atoms 2. Describe the quantum theory of light 3. Describe the Bohr's Model of the Hydrogen Atom 4. Compare the Quantum Mechanical Theory of the atom with the Bohr Model 5. Cite the significance of quantum numbers 6. identify quantum numbers of electrons in an atom 7. Apply the different principles in creating the orbital diagram of the elements 8. Predict correctly the position of the elements in the Periodic Table using its electronic configuration 9. Trace the development of the Periodic Table 10. Explain the factors that affect the periodic trends as to atomic size, ionization energy, electronegativity and electron affinity of atoms 11. Determine the valence electrons and charge of the ion being formed 12. Compare and contrast the different properties of the elements in the Periodic Table 13. To know the relationship between atomic spectra and the electronic structure of atoms 14. Relate the electron configurations of the elements to the shape of the periodic table. 15. Determine the expected electron configuration of an element by its place on the periodic table. **Activating Prior Learning** How would you describe each of the following models of the atom: - Billiard ball model - Prout's hypothesis - Raisin bread model - Nuclear model - Planetary model A. **Atomic Spectra**  **[Learning Objectives:]** 1. **discuss the meaning of the term \'atomic spectrum\'** 2. **distinguish between the two main types of atomic spectra: absorption and emission spectra and describe how each forms** **[Let's Learn]** ***Figure 4-2.** Comparison of Wavelength* Electromagnetic waves always travel at the same speed (3x10^8^ m per second). This is one of their defining characteristics. In the electromagnetic spectrum there are many different types of waves with varying frequencies and wavelengths. They are all related by one important equation: **wavelength x frequency of oscillation = speed of light** λ ν = c where c = speed of light We can use this relationship to figure out the wavelength or frequency of any electromagnetic wave if we have the other measurement. Just divide the speed of light by whichever measurement you have and then you've got the other.  Frequency is measured in units of Hertz, which is also cycles per second (cps) or simply s^-1^. Thus 1 Hz = 1cps = s^-1^. If c is in ms^-1^ and is used in the above equation, then the wavelength must be in meters. However, other common units of wavelength are nanometers, nm and Angstrom, Å. The equivalents are given below: 1 nanometer, nm = 10^-9^ meter, m 1 Angstrom, Å = 10-10 meters, m You should also take note that wavelength and frequency are inversely related. The higher the frequency, the shorter the wavelength. Since all light waves move through a vacuum at the same speed, the number of wave crests passing by a given point in one second depends on the wavelength. All electromagnetic radiation is light, but we can only see a small portion of this radiation, the visible light (Figure 3). Cone-shaped cells in our eyes act as receivers tuned to the wavelengths in this narrow band of the spectrum. Typically, our eyes can detect wavelength from 380-700 nanometers. ***Figure 4-3.** The Electromagnetic Spectrum* Light is a form of energy and the relationship between energy, frequency and wavelength if expressed in the following equation: E = hν = [hc] λ where: h = Planck's constant = 6.626 x 10^-34^ J^-s^ You should understand that energy is directly proportional to frequency, but inversely proportional to wavelength. The greater the energy, the larger the frequency and the shorter (smaller) the wavelength. Given the relationship between wavelength and frequency --- the higher the frequency, the shorter the wavelength --- it follows that short wavelengths are more energetic than long wavelengths. **The electrons in an atom tend to be arranged in such a way that the energy of the atom is as low as possible. The ground state of an atom is the lowest energy state of the atom. When those atoms are given energy, the electrons absorb the energy and move to a higher energy level. These energy levels of the electrons in atoms are quantized, meaning again that the electron must move from one energy level to another in discrete steps rather than continuously. An excited state of an atom is a state where its potential energy is higher than the ground state. An atom in the excited state is not stable. When it returns back to the ground state, it releases the energy that it had previously gained in the form of electromagnetic radiation.** Then continuous spectrum, like in a rainbow, comes from white light, line spectrum is evident in colored compounds. Light spectrum only has a few wavelengths (not all) or lines. Atoms tend to absorb some wavelengths when electromagnetic radiation is passed through them which display only a few narrow absorption lines when recorded. The major difference between these two is that continuous spectra has all the wavelengths while line spectrum only contains some of the wavelengths. Line spectrum can also be generated in emission and absorption spectrum while continuous spectrum occurs when both absorption and emission spectra of a single species are put together. In other words, line spectrum can be in either emission spectrum or absorption spectrum. **The corresponding spectrum may exhibit a continuum, or may have superposed on the continuum bright lines (an *emission spectrum*) or dark lines (an *absorption spectrum*), as illustrated in Figure 4-4.**  ***Figure 4-4. Comparison of a continuous, emission and absorption spectrum*** **Thus, *emission spectrum* are produced by thin gases in which the atoms do not experience many collisions (because of the low density). The emission lines correspond to photons of discrete energies that are emitted when excited atomic states in the gas make transitions back to lower-lying levels.** **A *continuum spectrum* results when the gas pressures are higher, so that lines are broadened by collisions between the atoms until they are smeared into a continuum. We may view a continuum spectrum as an emission spectrum in which the lines overlap with each other and can no longer be distinguished as individual emission lines.** **An *absorption spectrum* occurs when light passes through a cold, dilute gas and atoms in the gas absorb at characteristic frequencies; since the re-emitted light is unlikely to be emitted in the same direction as the absorbed photon, this gives rise to dark lines (absence of light) in the spectrum.** The emitted light can be observed as a series of colored lines with dark spaces in between; this series of colored lines is called a **line **or **atomic spectra**. Each element produces a unique set of spectral lines. Since no two elements emit the same spectral lines, elements can be identified by their line spectrum. **[Feedback]** Answer the following questions: 1. A light wave is an electromagnetic wave that has both an electric and magnetic component associated with it. Electromagnetic waves are often distinguished from mechanical waves. The distinction is based on the fact that electromagnetic waves \_\_\_\_\_\_. a. can travel through materials and mechanical waves cannot b. come in a range of frequencies and mechanical waves exist with only certain frequencies c. can travel through a region void of matter and mechanical waves cannot d. electromagnetic waves cannot transport energy and mechanical waves can transport energy 2. Consider the electromagnetic spectrum as you answer these questions. a. Which region of the electromagnetic spectrum has the highest frequency? b. Which region of the electromagnetic spectrum has the longest wavelength? c. Which region of the electromagnetic spectrum will travel with the fastest speed? 3. Consider the visible light spectrum as you answer these questions. a. Which color of the visible light spectrum has the lowest frequency? b. Which color of the visible light spectrum has the shortest wavelength? c. Which color has a higher energy, yellow or red? B. **Bohr Model of the Hydrogen Atom**  **[Learning Objectives:]** 1. **To understand how the emission spectrum of hydrogen demonstrates the quantized nature of energy** 2. **To describe Bohr's model of the hydrogen atom** 3. **Illustrate energy state using the energy-level diagram** **[Let's Learn]**  **The great Danish physicist Niels Bohr (1885-1962) introduced the atomic Hydrogen model in 1913. He used the planetary model of the atom to explain the atomic spectrum and size of the hydrogen atom. He described it as a positively charged nucleus, comprised of protons and neutrons, surrounded by a negatively charged electron cloud. In the model, electrons revolve around the nucleus in atomic shells. The atom is held together by electrostatic forces between the positive nucleus and negative surroundings.** **Bohr, in an attempt to understand the structure of an atom better, combined classical theory with the early quantum concepts and suggested the following theories:** 1. **In an atom, electrons encircle the nucleus in specific allowable paths called orbits. Each orbit is assigned a quantum number, n, which is restricted to integer values: 1, 2, 3, 4, 5, 6..... When the electron is in one of these orbits, its energy is fixed. The ground state of the hydrogen atom, where its energy is lowest, is when the electron is in the orbit that is closest to the nucleus. The orbits that are further from the nucleus are all of successively greater energy.** 2. **The electron can remain in an orbit indefinitely.** 3. **He explained that electrons can be moved into different orbits with the addition of energy. When the energy is removed, the electrons return back to their ground state, emitting a corresponding amount of energy - a quantum of light, or photon. This was the basis for what later became known as quantum theory.** **Bohr derived the equations for the energy levels available to the electron in the hydrogen atom thus,** **E = [-2π^2^me^4^] [Z^2^]** **h^2^ n^2^** **where:** **m = mass of the electron** **e = charge of the electron** **h = Planck's constant** **Z = nuclear charge, +1 for hydrogen** **n = an integer which may have values starting with 1** **upon substitution of the values for m, e, and h, the equation becomes** **E = [2.178 x 10-18] Joules** **n^2^** **The highest possible value of E in the above equation is zero when the electron is at an infinite distance from the nucleus (n = ꝏ). The energy of an electron bound to the nucleus is negative with the lowest value when n = 1.** **The radius of the orbit corresponding to each level is given by the equation:** **r = [n^2^ h^2^ \_ ]** **4 [ π^2^me^4^]** **The first orbit, where n = 1, has a radius of 0.529x10^-11^ m. this usually referred to as the first Bohr radius and given the symbol a~0~ :** **r = n^2^ a~0~** **The n values, radii and energies of the six Bohr orbits are shown in Figure 5.** ***Figure 4-5.** The first six electron orbits in the hydrogen atom* Suppose the electron moves from n = 5 orbit to n = 2 orbit with energies E = - 0.08712x10^-18^ J E = - 0.5445x10^-18^ J The difference in energy, ∆E is computed as follows: ∆E = E~final~ - E~initial~ = E~2~ -- E~5~ = (- 0.5445x10^-18^ J) -- (- 0.08712x10^-18^ J) = - 0.45738x10^-18^ J The negative sign of ∆E means that energy is given off in the process. This difference in energy is the energy of the quantum of light emitted. The wavelength of the light corresponding to this transition may be computed as follows: ∆E = [hc] λ Solving for wavelength, λ = [hc] ∆E = [(6.626 x 10^-34^ J.sec) (3 x 10^8^ m s^-1^)] 0.45738 x 10^-18^ J = 4.346 x 10^-7^ m x [1nm ] 10^-9^m The correspondence between transitions of the electron and the major lines in the visible region of the hydrogen spectrum is shown in Figure 6. An empirical formula to describe the positions (wavelengths) λ of the hydrogen emission lines in the series was discovered in 1885 by Johann Balmer. It is known as the Balmer formula: The constant R~H~ = 1.09737×10^7^m^−1^ is called the Rydberg constant for hydrogen. In equation, the positive integer n takes on values n=3, 4, 5, 6 for the four visible lines in this series. The series of emission lines given by the Balmer formula is called the Balmer series for hydrogen. Other emission lines of hydrogen that were discovered in the twentieth century are described by the Rydberg formula, which summarizes all of the experimental data: When n*~f~* = 1, the series of spectral lines is called the Lyman series. When n*~f~* = 2, the series is called the Balmer series, and in this case, the Rydberg formula coincides with the Balmer formula. When n*~f~* = 3, the series is called the Paschen series. When n*~f~* = 4, the series is called the Brackett series. When n*~f~* = 5 , the series is called the Pfund series. When n*~f~* = 6 , we have the Humphreys series. As you may guess, there are infinitely many such spectral bands in the spectrum of hydrogen because n~f~n~f~ can be any positive integer number.  ***Figure 4-6. The Balmer series---the spectral lines in the visible region*** ***of hydrogen\'s emission spectrum*** **The movement of electrons between these energy levels produces a spectrum. The Balmer equation is used to describe the four different wavelengths of Hydrogen which are present in the visible light spectrum. These wavelengths are at 656, 486, 434, and 410nm. These correspond to the emission of photons as an electron in an excited state transitions down to energy level n=2. The Rydberg formula, below, generalizes the Balmer series for all energy level transitions. To get the Balmer lines, the Rydberg formula is used with an n~f ~of 2.** **[Feedback]** 1. An emission spectrum gives one of the lines in the Balmer series of the hydrogen atom at 410 nm. This wavelength results from a transition from an upper energy level to n=2. What is the principal quantum number of the upper level? 2. The Bohr model of the atom was able to explain the Balmer series because: a. larger orbits required electrons to have more negative energy in order to match the angular momentum. b. differences between the energy levels of the orbits matched the difference between energy levels of the line spectra. c. electrons were allowed to exist only in allowed orbits and nowhere else. d. none of the above 3. One reason the Bohr model of the atom failed was it did not explain why a. accelerating electrons do not emit electromagnetic radiation. b. moving electrons have a greater mass. c. electrons in the orbits of an atom have negative energies. d. electrons in greater orbits of an atom have greater velocities C. **Quantum Theory: Quantum Mechanical Model of the Atom** **[Learning Objectives:]** 1. **Describe the quantum theory of light;** 2. **Describe the quantum mechanical model of the atom and how it compares with the Bohr model;** 3. **Give significance of quantum numbers;** 4. **Identify quantum numbers of an electron in a given atom.** 5. **To understand how the electron's position is represented in the wave mechanical model** **[Let's Learn]** *[The Quantum Theory]* Quantum theory is the theoretical basis of modern physics that explains the nature and behavior of matter and energy on the atomic and subatomic level.  In 1900, physicist Max Planck presented his quantum theory to the German Physical Society. Planck had sought to discover the reason that [radiation](https://whatis.techtarget.com/definition/ionizing-radiation) from a glowing body changes in color from red, to orange, and, finally, to blue as its temperature rises. He found that by making the assumption that energy existed in individual units in the same way that matter does, rather than just as a constant [electromagnetic](https://whatis.techtarget.com/definition/electromagnetic-radiation-spectrum) wave - as had been formerly assumed - and was therefore *quantifiable.* The existence of these units became the first assumption of quantum theory. Light consists of packets or bundles of energy with each bundle containing a definite amount of energy. Planck called the bundles of energy *quanta*. The size of the packet or bundle is smaller for light of lower frequency. If light is of a single frequency, then the bundles or quanta are of the same energy content (monochromatic). However, if light is of different frequencies, then the bundles or quanta have different energy content (polychromatic). **The Development of Quantum Theory** - In 1900, Planck made the assumption that energy was made of individual units, or quanta. - In 1905, Albert Einstein theorized that not just the energy, but the radiation itself was *quantized* in the same manner. - In 1924, Louis de Broglie proposed that there is no fundamental difference in the makeup and behavior of energy and matter; on the atomic and subatomic level either may behave as if made of either particles or waves. This theory became known as the *principle of wave-particle duality*: elementary particles of both energy and matter behave, depending on the conditions, like either particles or waves. - In 1927, Werner Heisenberg proposed that precise, simultaneous measurement of two complementary values - such as the position and momentum of a subatomic particle - is impossible. Contrary to the principles of classical physics, their simultaneous measurement is inescapably flawed; the more precisely one value is measured, the more flawed will be the measurement of the other value. This theory became known as the [uncertainty principle](https://whatis.techtarget.com/definition/uncertainty-principle), which prompted Albert Einstein\'s famous comment, \"God does not play dice.\" Light can be viewed as stream of particles called *photons*. Each photon carries the quantum or packet of energy given by the equation E = hν. Since the constant h is very small, the bundles of energy are very small. Thus, the particulate nature of light is not obvious to us. *[The Quantum Mechanical Model of an Atom]* The theory of light has a dual nature: wavelike or continuous and particulate (consisting of photons). This was also applied to matter by Louis de Broglie (1892). He suggested that particles properties such as mass, velocity and momentum also have wavelike properties such as wavelength. The dual behavior of particles is not obvious in our day-to-day experience because the wavelength associated with a particle is detectable or measurable only for particles with very small masses, like the electron or ant atom or molecule. German Physicist, Erwin Schrodinger (1887-1961), applied the idea of an electron having wavelike properties. He derived the equation for the energy of an atom and found that by assuming wavelike behavior for the electron, the quantized nature of the energy of the electron becomes an inevitable consequence. A major problem with Bohr\'s model was that it treated electrons as particles that existed in precisely-defined orbits. Based on de Broglie\'s idea that particles could exhibit wavelike behavior, Austrian physicist Erwin Schrödinger theorized that the behavior of electrons within atoms could be explained by treating them mathematically as matter waves. This model, which is the basis of the modern understanding of the atom, is known as the *Quantum Mechanical or Wave Mechanical Model.* The Quantum Mechanical or Wave Model came from the following history: 1. Louis de Broglie: Electrons behave with **wave** and **particle** properties 2. Werner Heisenberg: It is impossible to know both the position and velocity of an electron simultaneously - **Heisenberg\'s Uncertainty Principle.** 3. Erwin Schrodinger: refined the wave-particle theory proposed by de Broglie. **The quantum mechanical model of the atom treats an electron like a wave.** *[Quantum Number]* The quantum mechanical model describes the probable location of electrons in atoms by the quantum numbers. The first three quantum numbers describe fully an orbital. Since an orbital is an energy state, **n, *l,*** *and* (***m~l~***) therefore the define an energy state of the atom. 1. Principal quantum number (**n**) describes the shell or energy level or better the average relative distance of an electron from the nucleus. It also indicates the relative size and energy of atomic orbitals. **n**=integers: n= 1, 2, 3, etc. As **n** increases: - orbital becomes larger - electron spends more time farther away from nucleus - atom\'s energy level increases 2. **Angular or azimuthal quantum number (*l***) describes the *shape* of the orbital or the region of space occupied by the electron. The allowed values ***l*** of  depend on the value of ***n*** and can range from 0 to ***n*** − 1: ***ɭ*** = 0, 1, 2,..., n−1 For example, if *n* = 1,  ***l*** can be only 0; if *n* = 2,  ***l*** can be 0 or 1; etc.. For a given atom, all wave functions that have the same values of both ***n*** and ***l*** form a subshell. The regions of space occupied by electrons in the same subshell usually have the same shape, but they are oriented differently in space. To simply so as not to be confused, the ***l*** values have been given letter symbols as follows: ***l*** (numerical value) 0 1 2 3... ***l*** (equivalent subshells) s p d f... Principal quantum numbers consist of sublevels or subshells. Sublevels define the orbital shapes (s, p, d, f) ***Figure 4-7.** The Orbitals in Subshells within the Shells of atoms* ***Figure 4-8.** The Shapes of the Orbitals* The main factor affecting an electron\'s probability distribution for s orbitals, which are spherical, is the distance from the nucleus (the radius). But, for other types of orbitals such as p, d, and f orbitals, the electron\'s angular position relative to the nucleus becomes a factor in the probability density. This will make you appreciate more interesting orbital shapes, such as the ones in image below. **Let us consider the number of orbitals in e**ach sublevel: ***Figure 4-9.** Schematics showing the general shapes of s, p, d, and f orbitals.* *Image credit: [[UCDavis Chemwiki]](http://chemwiki.ucdavis.edu/@api/deki/files/8855/Single_electron_orbitals.jpg), [[CC BY-NC-SA 3.0 US]](http://creativecommons.org/licenses/by-nc-sa/3.0/us/)* 3. Magnetic **quantum number** (***m~l~***) describes the orientation in space of a particular orbital. It is called the *magnetic* quantum number because the effect of different orientations of orbitals was first observed in the presence of a magnetic field. Consequently, its value depends on the orbital angular quantum number ***l***. Given a certain ***l***, ***m~l~*** is an interval ranging from --l to +l, so it can be zero, a negative integer, or a positive integer. ***m~l~*** = −***l***, (−l+1),(−l+2),...,−2,−1, 0, +1, +2,...( ***l*** --1), ( ***l*** --2), + ***l*** **Here is an example: If n=3, and *l* =2, then what are the possible values of *m~l~***?  4. **Spin quantum number (*m~s~*) designates the direction of the electron spin and may have a spin of +1/2, represented by↑, or --1/2, represented by ↓. This means that when *m~s~* is positive the electron has an upward spin, which can be referred to as \"spin up.\" When it is negative, the electron has a downward spin, so it is \"spin down.\" The significance of the electron spin quantum number is its determination of an atom\'s ability to generate a magnetic field or not.** **[Feedback]** A. **For each set of quantum numbers below, determine and identify whether the combination of the quantum numbers is possible or not. If not, state the reason.** 1. **n = 2 *l*** *= 1* ***m~l~*** = -2 2. **n = 3 *l*** *= 0* ***m~l~*** = +3 3. **n = 6 *l*** *= 0* ***m~l~*** = 0 4. **n = 5 *l*** *= 3* ***m~l~*** = -1 5. **n = 3 *l*** *= 4* ***m~l~*** = +3 B. How many possible orbitals are there for 1. n = 4 - \_\_\_\_\_\_\_\_\_\_\_ 2. n = 6 - \_\_\_\_\_\_\_\_\_\_\_ C. Identify the subshell in which electrons with the following quantum numbers are found: a. **n = 3 *l*** *= 1* \_\_\_\_\_\_\_\_\_\_\_\_ b. **n = 5 *l*** *= 3 \_\_\_\_\_\_\_\_\_\_\_\_* c. **n = 2 *l*** *=* 0 \_\_\_\_\_\_\_\_\_\_\_\_ D. Choose from the given choices the most appropriate answer for each question 1. The magnetic quantum number describes the: a. Number of electrons b. Average distance of the most electron dense regions from the nucleus c. Spatial orientation of the orbital d. Shape of the orbital 2. How many electrons can inhabit all of the n=4 orbitals? a. 14 b. 24 c. 32 d. 36 3. The principle quantum number is related to:  a. the shape of the orbital b. the spatial orientation of the orbital  c. the average distance of the most electron-dense regions from the nucleus  d. the number of electrons 4. Which of the following sets is not an acceptable set of quantum numbers? a. **n = 2 *l*** *= 1* ***m~l~*** = -1 b. **n = 7 *l*** *= 3* ***m~l~*** = +3 c. **n = 2 *l*** *= 1* ***m~l~*** = +1 d. **n = 3 *l*** *= 1* ***m~l~*** = -3 5. How many orbitals are in the 5s subshell?  a. 2 b. 5 c. 1 d. 3 6. How many orbitals are in the 4d subshell? a. 3 b. 6 c. 4 d. 5 7. How many total electrons can the 'p' orbitals hold?  a. 1 b. 6 c. 7 d. 3 8. What are the quantum numbers that describe a 3p orbital? a. **n = 3 *l*** *= 1* ***m~l~*** = +1 b. **n = 3 *l*** *= 1* ***m~l~*** = 0 c. **n = 3 *l*** *= 1* ***m~l~*** = -1 d. **All of the above** 9. Probabilities are used because of Heisenberg\'s Uncertainty Principle. This principle states that the \_\_\_\_\_\_\_\_\_ and \_\_\_\_\_\_\_\_ of an electron can be measured but not simultaneously. a. charge, velocity b. charge, position c. position, velocity d. charge, spin 10. **Why are electrons assigned quantum numbers?** a. **Quantum numbers are designed to find the possible location of electrons** b. **Quantum numbers are designed to find the possible velocity of electrons** c. **Quantum numbers are designed to find the possible spin of electrons** 11. **The quantum mechanical model describes electrons as:** a. **Particles** b. **Waves** c. **Particles with wave-like properties** d. **Small, hard spheres** D. **Electronic Configuration of the Element** **[Learning Objectives:]** **After going through this module, you are expected to:** 1. **write the predicted ground-state electron configurations of atoms** 2. **specify the shell and subshell symbols and their positions.** 3. **define the position of electrons in different shells of an atom.** 4. **use orbital filling diagrams to describe the locations of electrons in an atom.**  **[Let's Learn]** The electron configuration, also called electronic structure of an atom describes the arrangement of electrons distributed among the orbital shells and subshells. Usually, the electron configuration is used to describe the orbitals of an atom in its *ground state* (lowest energy state), but it can also be used to describe electron at higher energies called *excited state*. It can also be used to represent an atom that has ionized into a cation or anion by compensating with the loss of or gain of electrons in their subsequent orbitals. The various physical and chemical properties of elements can be correlated to their unique electron configurations. The valence electrons, electrons in the outermost shell, are the determining factor for the unique properties of the element. We will be learning how to write the arrangement of electrons in an atom in its ground state. When assigning electrons to orbitals, we must follow a set of three rules or principles: - the Aufbau Principle - the Pauli-Exclusion Principle - Hund\'s Rule *[Aufbau Principle]* The \'Aufbau\' is a German word for \'building up\'. [The Aufbau principle](https://chem.libretexts.org/Bookshelves/Physical_and_Theoretical_Chemistry_Textbook_Maps/Supplemental_Modules_(Physical_and_Theoretical_Chemistry)/Electronic_Structure_of_Atoms_and_Molecules/Electronic_Configurations/Aufbau_Principle), also called the building-up principle, states that electron\'s occupy orbitals in order of increasing energy. The order of is as follows: **1s\