Atomic Spectra PowerPoint PDF

Atomic Structure: -Emission spectra -Shapes of orbitals and Electronic Configuration CAPE Chemistry Grade: 12 (Lower sixth form) Ms. Lawrence  Arrangement of electrons in the atom Rutherford correctly placed the nucleus of an...

Atomic Structure: -Emission spectra -Shapes of orbitals and Electronic Configuration CAPE Chemistry Grade: 12 (Lower sixth form) Ms. Lawrence  Arrangement of electrons in the atom Rutherford correctly placed the nucleus of an atom at its center, but the placement of electrons posed a problem.  Issue 1: If electrons were stationary at a distance from the nucleus, the electrostatic attraction between the positively charged nucleus and the negatively charged electrons would pull them into the nucleus.  Issue 2: If the electrons moved (as predicted by classical physics), they would continuously radiate energy, causing them to spiral into the nucleus, which would lead to the collapse of the atom.  Arrangement of electrons in the atom  Thus, electrons cannot be stationary or continuously moving.  Instead, they occupy specific fixed energy levels outside the nucleus, known as quantized states.  In these states, electrons sit in discrete/specific energy levels (or orbits) around the nucleus without losing/radiating energy.  Arrangement of electrons in the atom  Ground State: Electrons prefer to be in the lowest energy levels possible for each of them. We say that the electrons are in their ground state. This is where the atom is most stable.  Excited State: When an electron absorbs a fixed amount of energy/radiation (quantum of radiation), it can jump/move to a higher energy level. The electron is then said to be in an excited state.  When an excited electron falls back to a lower energy level a quantum of radiation is given out.  This jump from one quantized state (energy level) to another with the absorption or emission of energy, the transition is called a quantum jump.  The smallest fixed amount of energy required to change an electron’s energy level is called a quantum of energy.  The Atomic Spectra  The best evidence for the fact that electrons in an atom surround the nucleus in certain allowed energy levels, or orbitals comes from a study of the atomic spectra. How is the Atomic spectrum formed?  When the electrons in an atom move between energy levels, they absorb or emit light (electromagnetic radiation).  An atomic spectrum is formed/produced when the electrons in an atom absorbs or emit electromagnetic radiation (light).  Each element produces a unique spectrum. Therefore, the atomic spectrum can be used to identify the element.  Electromagnetic radiation can be thought of as a stream of photons, which are particles with no mass that travel in a wave- like pattern at the speed of light, 3.00 × 10¹⁰ meters per second.  Each photon contains a certain amount of energy, and this energy is related to its wavelength (λ, the distance between two crests of a wave) and frequency (ν, the number of waves passing a point per second).  Visible light is a type of electromagnetic radiation, and like all electromagnetic radiation, light can behave as photons.  The shorter the wavelength of the light, the higher the energy of its photons, while the longer the wavelength, the lower the energy.  This relationship can be expressed through Planck’s equation: ΔE=hv Where: Δ E is the energy of the photon of light (J). h is Planck’s constant (6.63×10^-34Js) v (pronounced "nu") is the frequency of the light (in seconds inverse, s -1) The frequency and wavelength of electromagnetic radiation is related by the equation: c= frequency (v) × wavelength (λ) Where: where c is the speed of electromagnetic radiation in vacuo (3.00 × 10¹⁰ meters per second). Since: c = frequency (v) × wavelength (λ) Then: The energy can also be expressed as E = hc / wavelength, i.e. E=hc/ λ  The Electromagnetic Spectrum  The whole range of frequencies of electromagnetic radiation is called the electromagnetic spectrum.  The range of frequencies corresponds to all different forms of electromagnetic radiation in the universe.  It begins at the longest wavelengths and lowest frequency to the shortest wavelengths and highest frequency.  Electromagnetic Spectrum (EM) This spectrum includes many types of radiation: Radio waves in televisions and cell phones. Microwaves in satellites and microwave ovens Infrared in toaster ovens and night-vision X-rays in medical imaging Visible light is the portion of the electromagnetic spectrum we can see. It falls between wavelengths of 400 to 700 nanometers.  Continuous Spectrum  If a beam of white light is passed through a prism on to a screen, a spectrum of colours made up of all wavelengths of visible light is seen like in a rainbow.  This is called a continuous spectrum. There is no distinct division between the colours as they blend in from one to another. There are two main types of atomic spectra:  Absorption Spectra  Emission  Absorption Spectrum  If white light passes through a substance, the atoms in the substance can absorb light of certain/specific wavelength, causing dark lines to appear in the spectrum.  A line spectrum is formed which appears as distinct lines and not bands of colours. This line spectrum is an absorption spectrum.   When atoms are heated or supplied with electrical energy, they can absorb enough energy to become excited.  This means that their electrons move to higher energy levels.  As these excited electrons return to their ground state (their Emission original, more stable energy levels), they release energy in the form of light at specific wavelengths. Spectrum This release of energy produces an emission spectrum, which appears as bright, colored lines on a dark background.  Each of these lines corresponds to a particular wavelength of light emitted by the atom, and together they form a unique pattern for each element.  Emission Spectrum  Because only certain colors of light (photons of certain energy) can be emitted when the electron falls from a higher energy level to a lower energy level.  This provides evidence for discrete energy levels within the atom Emission spectrum of some elements  Emission Spectrum  When element like Neon is excited, it will often emit light of a characteristic red-orange colour. Another example:  Sodium vapor emits the yellow light characteristic of some modern streetlights.  Bohr’s Model  Bohr’s impact on the growing knowledge of atomic structure was the idea that energies in an atom are quantized. That is:  Electrons in an atom are found in orbits which exist only in specific, fixed energy levels.  The electron can change its energy by moving from one level to another but can never have an energy between adjacent levels.  Think of the energy levels as stair steps and the electron as a rubber ball. The ball can come to rest on one step or another, but never between two steps.  Likewise, the electron can only be in one orbit (energy level) at a time. It cannot reside between orbits, but it can move from one orbit to another. Bohr’s Model Energy is emitted or absorbed by the electrons only as the electron changes from one energy state to another. This energy is the energy emitted or absorbed by a photon (a particle of light). Bohr’s Model of the Hydrogen atom  Bohr suggested that the electron in the hydrogen atom circulates around the positively charged nucleus much like planets moving in orbits around the sun.  Bohr’s Model Cont’d  Bohr assigned the energy levels of the electron with the letter n, known as the principal quantum number (n). The values of n are positive, whole numbers, 1, 2, 3, etc.  The orbit closest to the nucleus, the n = 1 level, has the lowest energy and the smallest radius. The n = 2 orbit has a higher energy and larger radius.  As n gets larger, the energy and radius of the orbit increases.  Bohr’s Model of Hydrogen  The electron in a ground state hydrogen is in the first energy level (n = 1), and it cannot randomly change the orbit without external energy input.  If hydrogen gas is enclosed in a tube at low pressure and subjected to high voltage, it emits a pink light.  When this energy is added to the atom, the electron in the ground state absorbs a quantum of energy and moves to an orbit with a higher energy level which is further away from the nucleus.  Bohr’s Model of Hydrogen  This excited electron in the higher energy level can’t stay in the higher energy level for long.  It will eventually fall back down to a lower energy level (like returning down the stairs).  When it does this, it loses energy.  The energy that the electron loses is released as light.  This light has a specific color (or wavelength) that corresponds to the amount of energy lost.  Each color of light emitted represents a unique energy difference between the energy levels Emission spectrum of Hydrogen  Emission spectrum of Hydrogen When we look at the light from excited hydrogen atoms, we see a series of bright lines of color. These lines represent the emission spectrum of hydrogen. The lines occur because the electron transitions between specific energy levels. This provide evidence for discrete energy levels within the atom  For hydrogen, when the electron falls from higher levels (like 5, 4, or 3) to the second level, we see these colors.  This is called the Balmer series and is the only series we can see with our eyes. For example: A transition from the third energy level to the second emits red light (656nm)  A transition from the fourth energy level to the second emits blue- green (cyan) light (486 nm)  A transition from the n = 5 to n = 2 transition emits blue/ violet light (434 nm)  A transition from the n = 6 to n = 2 transition emits violet light (410 nm)  Did you notice that the transitions associated with larger n-level gaps correspond to emissions of photos with higher energy?  Now, if the energy difference between the orbits is too large, then the light emitted from the transition won’t fall into the visible region of the spectrum. For example, the transition from n = 6 to n = 1 level corresponds to UV radiation.  On the other hand, if the energy is too small then we have infrared (IR) or near IR radiation which we cannot see either. Non-visible series (UV or IR Region) Electron transitions to level n = 1 from any of the other orbits result in lines in the ultraviolet region of the spectrum and is called the Lyman series. The Paschen series arise from transitions from higher energy levels to principal quantum number 3. The Brackett series arise from transitions from higher energy levels to principal quantum number 4. The Pfund series arise from transitions from higher energy levels to principal quantum number 5.  Calculating The Wavelength of Electron Transition  The next question we need to address is how to calculate the wavelength of light for electron transitions in the hydrogen atom. For this, Johann Balmer (1825–1898) and Johannes Rydberg (1854–1919) derived a formula correlating the wavelength of emitted photons for the given n levels.  The Balmer – Rydberg equation was derived before the Bohr’s model for the hydrogen atom, and Bohr modified it to calculate the values of energy levels and their difference for the electron transitions.  Calculating The Wavelength of Electron Transition  Now, if the electron moves from a lower to a higher energy state, then energy is absorbed, and subtracting 1/ni2 from 1/nf2 will allow canceling of the negative signs, thus making ΔE positive.  On the other hand, when the electron falls from a higher energy state, a photon is released, and the negative sign will stay making ΔE negative meaning the system loses energy. For example, when the electron moves from n = 2 to n = 1, we will have a negative ΔE.  Bohr’s Model Cont’d  The energy difference between the lines in the series decreases as the energy increases (or the wavelength decreases).  Eventually the lines become so close that they form a continuous band called a continuum.  At this point and beyond it is impossible to distinguish between the lines and a convergence limit is said to be reached.  At the convergence limit, the electron will no longer experience the effect of the nuclear attraction.  In other words, it is ‘free’ from the influence of the nucleus and the atom that has lost the electron has become ionized. Quantum Mechanical Model of the Atom  The Bohr’s atomic model was one-dimensional and used one quantum number, n which explained the emission spectrum of hydrogen but failed to explain the spectra of more complex atoms.  With the development of Schroedinger’s wave equation in 1926 which assumes wave-particle duality of electrons, the position of an electron is described in terms of probability density which is the volume of space in which the electron is likely to be found.  The region where there is a high probability of finding an electron is called an orbital.  Schroedinger described an atomic model with electrons in three dimensions and with four quantum numbers. Quantum Mechanical Model of the Atom  The quantum mechanical model of the atom uses 4 quantum numbers: n, l, ml & ms to describe/ denote an orbital and the location of an electron in that orbital.  Electrons are NOT in circular orbits around nucleus.  Electrons are in a 3-D region around the nucleus called atomic orbitals.  Unlike orbits that are easy to visualize, orbitals have shapes that do not resemble the circular paths of orbits.  1. Principal quantum number (n)  The quantum number used to label each principal shell of an atom.  It is similar to the quantum number used to identify the orbits in Bohr’s model. But a principal shell is not an orbit.  The principal shells are the major energy levels within the atom.  Each principal shell is identified with a quantum number, n, a positive, whole number, 1 or greater. The principal quantum number indicates the:  Energy of the shell  Size of the shell  The n = 1 principal shell is closest to the nucleus and lowest in energy. As n increases, 2, 3, 4, and so on, the size and energy of the principal shell increases. 2. Subsidiary quantum number/ Angular momentum quantum number (l) Within each principal shell are subshells. l describes the subshells in n and the shape of the orbitals. l usually takes positive integer values from 0 to (n-1) The number of subshells in a principal shell equals the value of its quantum number, n. Each subshell is identified with a one-letter label: s, p, d, or f. Within known elements, the s-, p-, and d-subshells are encountered most often.  The n = 2 principal shell The n = 1 principal shell contains two subshells, an contains one subshell, an s-subshell and a p- Angular s-subshell. subshell. momentum It is identified as the 1s-subshell. They are identified as the 2s-subshell and the quantum 2p-subshell. number (l) The n = 3 principal shell The n = 4 principal shell contains four subshells, an Cont’d contains three subshells, an s-, p-, and a d-subshell. s-, p-, d-, and an f-subshell. They are known as the They are called the 3s- 4s-subshell, 4p- subshell, 3p-subshell, subshell, 4d-subshell, and the 3d-subshell. and the 4f-subshell.  3. The magnetic quantum number (ml)  The subshells are composed of one or more orbitals.  m, describes the number of orbitals within a sublevel.  s-subshell has 1 orbital and can hold a minimum of 2 electrons.  p-subshell has 3 orbitals and contain a maximum of six electrons, two in each of the three p-orbitals  d-subshell has 5 d-orbitals within each d-subshell. A d- subshell can contain a maximum of ten electrons, two in each of the five d-orbitals.  f-subshell has 7 orbitals 4. Spin magnetic quantum number (mElectrons  s ) much behave as if they are spinning on an axis as the earth spins on its axis.  A spinning electron acts like a very small bar magnet with north and south poles.  Small arrows pointing upward, ↑, or downward, ↓, are used to indicate the two orientations of spin.  Electron spin is important because two electrons in the same orbital must spin in opposite directions, ↑ ↓.  Ms can only have 2 values : either +½ (spin: up) or -½ (spin: down)  Practice Question 1. For an electron in the n=4 principal quantum number. List all possible values of l and ml.  Shape of Orbitals S orbitals  Each principal shell/energy level contains one s-orbital in an s-subshell.  The labels used to identify the s-subshells, 1s, 2s, 3s, etc., are also used to identify the s-orbital they contain.  For example: The 1s-orbital is in the 1s-subshell  An s-orbital can hold a minimum of 2 electrons.  Note that s-orbitals have spherical shapes. S orbitals  Shape of Orbitals Cont’d P orbitals  p-orbitals come in sets of three within each p-subshell.  Each p orbital is shaped like a dumbbell and is at right angles to each other.  The three 2p orbitals are referred to as the 2px, 2py and 2pz orbitals.  The labels used to identify the p-subshells, 2p, 3p, 4p, etc., also identify the sets of p- orbitals.  p-orbitals are always kept together as a set, and the three orbitals always have the same energy (they are degenerate).  A p-subshell can contain a maximum of six electrons, two in each of the three p-orbitals.  P orbitals Cont’d p-orbitals have two-lobe, dumbbell shapes. The nucleus is at the point where the two lobes meet.  Shape of Orbitals Cont’d D orbitals  d-orbitals come in sets of five within each d-subshell.  Since d-subshells first appear in the n = 3 principal shell, d-orbitals are first encountered in the n = 3 principal shell.  The d-orbitals in the 3d-subshell are referred to as the 3d-orbitals.  d-orbitals are always kept together as a set and they all have the same energy.  A d-subshell can contain a maximum of ten electrons, two in each of the five d-orbitals.  D orbitals Cont’d The 3d-orbitals have a four-lobe, four-leaf clover shape.  The symbols used to indicate the number of electrons in a subshell include the number of the principal shell (1, 2, 3...), the letter designation of the subshell (s, p, d, f), with a superscript number indicating the number of electrons. For example:  One electron in the s-orbital in the 2s-subshell: 2s1  Four electrons in a set of p-orbitals in the 2p-subshell: 2p4  Six electrons in a set of d-orbitals in the 3d-subshell: 3d6  Electronic Configuration  The electronic configuration represents the electronic structure of an atom.  The electronic configuration of an atom shows the way the electrons are distributed among its subshells and orbitals.  Though there are countless ways in which electrons could be arranged about the nucleus:  the most stable arrangement of electrons is the one in which the electrons are in the lowest energy subshells possible, and this arrangement is called the ground state electronic configuration of the atom.  Any other arrangement of electrons would be an “excited” state, one less stable than the ground state  Electronic Configuration  Within a given principal shell, the s-subshell is a bit more stable (lower energy) than the p-subshell, and the p-subshell is a bit more stable than the d-subshell.  The ground state configuration of electrons for a many-electron element is determined using a building-up process in which electrons are added to subshells in a specific sequence starting with the 1s-subshell, the most stable subshell, and continuing in order of increasing subshell energies.  The sequence of increasing energies is: 1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p → 5s → 4d, etc.  This building-up principle has been known by its German name, the Aufbau principle Electronic Configuration Two-dimensional array of the building up order of subshells  Rules to remember when writing Electronic configuration The Pauli principle:  No more than two electrons can be placed in a single orbital.  Two electrons in the same orbital must have spins in opposite directions, ↑ ↓.  Two electrons with opposite spins in an orbital are said to be “paired.  Rules to remember when writing Electronic configuration Hund’s rule:  Electrons entering a set of equal-energy orbitals will fill them singly before any electrons are paired. Hund’s rule concerns the way electrons fill the sets of orbitals in p-,and d subshells.  For example, the three orbitals in a p-subshell have the same energy. If three electrons are in a set of p-orbitals, they will arrange themselves so that each orbital has one electron.  If a fourth electron is added, it will “pair up” with one of the single electrons. It makes no difference which one.  Remember, orbitals of the same energy fill singly before any pairing occurs.  Rules to remember when writing Electronic configuration  The number of electrons in a neutral atom equals its atomic number.  The lowest energy subshell is filled first, then the next lowest, following the sequence of increasing subshell energies.  The increasing subshell energies was shown in the two-dimensional array.  No more than two electrons can be placed in an orbital.  Two electrons in one orbital must have their spins paired, ↑ ↓.  This is a requirement of the Pauli principle. In p-,and d-,subshells that have three or five orbitals, each orbital is filled singly before any orbital contains two electrons. This is Hund’s rule. Writing Electronic Configurations  Writing Electronic Configurations  The ground state electronic configuration of helium, 2He, completely fills the 1s-orbital with two electrons.  The spins of the electrons are “paired.”  The configuration for helium would be read as “one-s-two” (not one-s- squared).  Writing Electronic Configurations Lithium (atomic number 3), 3Li, has three electrons to distribute among the subshells.  Following the first arrow in the filling diagram, two electrons are added to the 1s-orbital, filling it completely.  The next higher arrow leads to the 2s-subshell. The third electron is placed in the 2s-orbital in the 2s-subshell.  The configuration of lithium would be read as “one-s-two, two-s-one.”  Writing Electronic Configurations Noble Gas Notation:  To simplify the electron configuration, we often use noble gas notation.  This involves writing the electron configuration of the nearest noble gas preceding the element in brackets, then continuing with the remaining electrons.  For example, the electron configuration of silicon can be started from the noble gas neon [Ne], which accounts for the first 10 electrons.  Writing Electronic Configurations Noble Gas Notation: Example: Silicon (Si) Atomic Number: Silicon has an atomic number of 14, meaning it has 14 electrons.  Using Noble Gas Notation: The electron configuration starts with [Ne], which accounts for the first 10 electrons.  Writing Electronic Configurations Noble Gas Notation:  Filling the Remaining Electrons: After [Ne], we need to add 4 more electrons: The next two electrons fill the 3s sub-shell: 3s2 The remaining two electrons go into the 3p sub-shell: 3p2. Final Configuration: Therefore, the complete electron configuration for silicon is: Si: [Ne]3s23p2  Ionization Energy  The ionization energy of an element can be defined as the energy required to remove one or more loosely bounded electrons from the isolated gaseous atom in the ground electronic state to form a positively charged ion, i.e., form a cation.  In this process, energy must be absorbed to remove electrons, and therefore it is an endothermic process.  The first ionization energy can be expressed as X(g) + energy →X+(g) + e-  Here, X can be denoted as any isolated gaseous atoms or molecule,  X+ is denoted as the resultant ion, and e- is the loosely bounded electron.  The energy required for this process can be expressed in kJ/mol or electron volts (eV).  First Ionization Energy The first ionization energy is the energy required for the removal of the first (outermost) electron from a neutral atom in its gaseous state. This is generally the lowest ionization energy, as each successive electron becomes harder to remove due to an increased effective nuclear charge on remaining electrons. The energy required to remove each successive electron is called the second, third etc. ionization energy. All ionization energies are positive because it requires energy to remove an electron. Factors Affecting First Ionization Energy of Elements Atomic Radius The atomic radius is the distance between the nucleus and the outermost electron. As the atomic radius increases, the outer electron is farther from the nucleus, making it easier to remove. Trend: Elements with a larger atomic radius (typically lower down a group) have lower first ionization energies because the nucleus has a weaker hold on the distant outer electron. Factors Affecting First Ionization Energy of Elements Electron Shielding/ The screening effect of the inner electrons.  Apart from electrons being attracted to the nucleus, they also experience electron repulsion by other electrons.  The outer electrons are shielded from the attraction of the nucleus by the repelling effect of the inner electrons. The screening effect of electrons in lower energy levels is more effective than electrons in higher energy levels.  Electrons in the same energy level exert negligible screening effect on each other. Therefore, as the screening effect of the inner electron becomes more effective, the ionization energy decreases.  Factors Affecting First Ionization Energy of Elements Nuclear Charge The nuclear charge is the total positive charge from all the protons in the nucleus of an atom. Each proton has a positive charge, so more protons mean a stronger positive charge. How Nuclear Charge Affects Electrons A higher nuclear charge (more protons) means a stronger attraction between the nucleus and the electrons. This stronger pull makes it harder to remove an electron from the atom, resulting in a higher first ionization energy.  Factors Affecting First Ionization Energy of Elements How Nuclear Charge Affects Electrons  As the nuclear charge increases, the attraction of the nucleus for the outer electron increases and the ionization energy increases.  It must be noted that the atomic radii and electron screening can outweigh the effect of the nuclear charge.  For example, in group 1 although Cs has a larger nuclear charge than Na, it loses an electron more readily than Na.  Factors Affecting First Ionization Energy of Elements Explanation for Sodium (Na) and Cesium (Cs): Nuclear Charge: o Cesium (Cs) has a greater nuclear charge than sodium (Na) because it has more protons. In theory, this higher nuclear charge should mean a stronger pull on electrons, which would suggest a higher ionisation energy for Cs. Atomic Radius: o Cesium, however, has a much larger atomic radius than sodium because it’s further down the group in the periodic table. o Each additional energy level increases the distance between the nucleus and the outermost electron, making Cs's outer electron much farther away than Na’s.  Factors Affecting First Ionization Energy of Elements Explanation for Sodium (Na) and Cesium (Cs): Electron Shielding: Cesium has more inner electron shells, which create a “shielding” effect. This shielding reduces the attractive force between the nucleus and the outermost electron. This shielding effect essentially “cancels out” some of the nuclear charge, making it easier for Cs to lose its outer electron despite having a higher nuclear charge than Na.  Ionization energy and sub-shells  Ionization energy data provides evidence for sub-shells by showing that removing electrons from certain energy levels requires noticeably different amounts of energy.  Ionization energy and sub-shells Evidence for Sub-shells from Ionization Energy Data Trends in Ionization Energy: o Ionization energy generally increases across a period (from left to right) due to increasing nuclear charge, which attracts electrons more strongly, making them harder to remove. o However, there are notable dips in ionization energy that correspond to changes in electron configuration related to sub-shells.  Ionization energy and sub-shells Explanation using Period 3 elements Consider the elements in Period 3 and their electron configuration as shown in the table below: Element Symbol Electron First Ionization Configuration Energy (IE₁) Sodium Na [Ne] 3s¹ 496 Magnesium Mg [Ne] 3s² 737 Aluminum Al [Ne] 3s² 3p¹ 577 Silicon Si [Ne] 3s² 3p² 786 Phosphorus P [Ne] 3s² 3p³ 1012 Sulfur S [Ne] 3s² 3p⁴ 999 Chlorine Cl [Ne] 3s² 3p⁵ 1251 Argon Ar [Ne] 3s² 3p⁶ 1521  Ionization energy and sub-shells Element Symbol Electron First First Ionization Energy: Configura Ionization  The first ionization energy increases from tion Energy Na to Mg due to the addition of protons (IE₁) and the effective nuclear charge Sodium Na [Ne] 3s¹ 496 experienced by the outer electrons. Magnesiu Mg [Ne] 3s² 737 m  When moving from Mg to Al, there is a Aluminu Al [Ne] 3s² 577 drop in ionization energy. m 3p¹ Silicon Si [Ne] 3s² 786  This drop occurs because Al has a 3p¹ 3p² electron, which is in a higher energy sub- Phosphor P [Ne] 3s² 1012 shell compared to the 3s electrons. us 3p³  The 3p electron is further from the Sulfur S [Ne] 3s² 999 nucleus and experiences more shielding 3p⁴ from the inner electrons, making it easier Chlorine Cl [Ne] 3s² 1251 to remove. 3p⁵ Argon Ar [Ne] 3s² 1521 3p⁶  Ionization energy and sub-shells Element Symbol Electron First  As we continue from Al to Si, the Configura Ionization ionization energy increases again. tion Energy (IE₁)  However, when we reach Phosphorus (P), Sodium Na [Ne] 3s¹ 496 there is a drop in ionization energy when Magnesiu Mg [Ne] 3s² 737 comparing P to S. This drop is attributed m to the pairing of electrons in the 3p sub- Aluminu Al [Ne] 3s² 577 shell. m 3p¹  Sulfur has paired electrons in the 3p⁴ Silicon Si [Ne] 3s² 786 configuration, which leads to increased 3p² electron-electron repulsion and makes it Phosphor P [Ne] 3s² 1012 easier to remove one of the paired us 3p³ electrons. Sulfur S [Ne] 3s² 999 3p⁴  The ionization energy of Argon is Chlorine Cl [Ne] 3s² 1251 significantly higher due to its complete 3p⁵ outer shell (3s² 3p⁶), indicating a stable Argon Ar [Ne] 3s² 1521 configuration that resists electron 3p⁶ removal.  Derive the electronic configuration of an element from data on successive ionisation energies Step 1: Let’s say you have the following ionization energies for an element, say Element X: Ionization Step Ionization Energy (kJ/mol) 1st IE 500 2nd IE 1000 3rd IE 1400 4th IE 3000 5th IE 4000  Derive the electronic configuration of an element from data on successive ionisation energies Step 2: Analyze the Data Ionization Step Ionization Energy (kJ/mol)  Look for large jumps in the ionization energy 1st IE 500 values. These jumps indicate when you are removing an electron from a different shell or 2nd IE 1000 subshell. 3rd IE 1400 For Element X, we can see: 4th IE 3000  The first three ionization energies (500, 1000, 1400 kJ/mol) are relatively low and close to each 5th IE 4000 other.  However, the jump from the 3rd (1400 kJ/mol) to the 4th ionization energy (3000 kJ/mol) is very large.  Derive the electronic configuration of an element from data on successive ionisation energies  Step 3: Interpret the Jumps Ionization Step Ionization Energy (kJ/mol) Jumps indicate stability: 1st IE 500  The first three electrons can be 2nd IE 1000 removed without a huge increase in energy, which means they are likely in 3rd IE 1400 the outer shell (valence electrons). 4th IE 3000  The large jump to the 4th ionization energy suggests that removing this 5th IE 4000 electron is much harder because it comes from a more stable, inner shell.  Derive the electronic configuration of an element from data on successive ionisation energies  Step 4: Deduce the Electronic Ionization Step Ionization Energy (kJ/mol) Configuration 1st IE 500 From the analysis: 2nd IE 1000  Since the first three ionization energies are relatively low, we can conclude that 3rd IE 1400 Element X has three valence electrons. 4th IE 3000  This means the element is likely in Group 13 of the periodic table (elements like 5th IE 4000 Aluminum), which have three valence electrons.  Derive the electronic configuration of an element from data on successive ionisation energies Ionization Step Ionization Example Electronic Configuration Energy (kJ/mol)  If Element X has three valence electrons, 1st IE 500 its electronic configuration might be: 2nd IE 1000  Ne 3s² 3p¹ (which means it has 2 3rd IE 1400 electrons in the 3s subshell and 1 electron in the 3p subshell). 4th IE 3000 5th IE 4000