IB Chemistry DP Atomic Structure PDF

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Head to savemyexams.co.uk for more awesome resources YOUR NOTES IB Chemistry DP  2. Atomic Structure CONTENTS 2.1 Atomic & Electronic Structure 2.1.1 The Nuclear Atom 2.1.2 Deducing Subatomic Particles 2.1.3 Relative Atomic Mass Calculations 2.1.4 The Electromagnetic Spectrum 2.1.5 Emission Spectra 2.1.6 Energy Levels & Sublevels 2.1.7 Sublevels & Orbitals 2.1.8 Writing Electron Configurations Page 1 of 32 © 2015-2023 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to savemyexams.co.uk for more awesome resources 2.1 Atomic & Electronic Structure YOUR NOTES  2.1.1 The Nuclear Atom Mass & Charge Distribution The mass of an atom is concentrated in the nucleus, because the nucleus contains the heaviest subatomic particles (the neutrons and protons) The mass of the electron is negligible The nucleus is also positively charged due to the protons Electrons orbit the nucleus of the atom, contributing very little to its overall mass, but creating a ‘cloud’ of negative charge The electrostatic attraction between the positive nucleus and negatively charged electrons orbiting around it is what holds an atom together The mass of the atom is concentrated in the positively charged nucleus which is attracted to the negatively charged electrons orbiting around it Page 2 of 32 © 2015-2023 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to savemyexams.co.uk for more awesome resources Types of Subatomic Particles YOUR NOTES The protons, neutrons and electrons that an atom is made up of are called subatomic  particles These subatomic particles are so small that it is not practical to measure their masses and charges using conventional units (such as grams or coulombs) Instead, their masses and charges are compared to each other, and so are called ‘relative atomic masses’ and ‘relative atomic charges’ These are not actual charges and masses, but rather charges and masses of particles relative to each other Protons and neutrons have a very similar mass, so each is assigned a relative mass of 1 Electrons are 1836 times smaller than a proton and neutron, and so their mass is often described as being negligible The relative mass and charge of the subatomic particles are: Relative Mass & Charge of Subatomic Particles Table  Exam Tip You can see from the table how the relative mass of an electron is almost negligibleThe charge of a single electron is -1.602189 x 10-19 coulombs, whereas the charge of a proton is +1.602189 x 10-19 coulombs. However, relative to each other, their charges are -1 and +1 respectively. This information can also been found in the IB Data Booklet Page 3 of 32 © 2015-2023 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to savemyexams.co.uk for more awesome resources 2.1.2 Deducing Subatomic Particles YOUR NOTES  Atoms: Key Terms The atomic number (or proton number) is the number of protons in the nucleus of an atom and has the symbol Z The atomic number is also equal to the number of electrons that are present in a neutral atom of an element E.g. the atomic number of lithium is 3, meaning that a neutral lithium atom has 3 protons and, therefore, also has 3 electrons The mass number (or nucleon number) is the total number of protons + neutrons in the nucleus of an atom, and has the symbol A The number of neutrons can be calculated by: Number of neutrons = mass number - atomic number Protons and neutrons are also called nucleons, because they are found in the nucleus  Exam Tip The mass (nucleon) and atomic (proton) number are given for each element in the Periodic Table Page 4 of 32 © 2015-2023 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to savemyexams.co.uk for more awesome resources Isotopes: Basics YOUR NOTES Isotopes are atoms of the same element that contain the same number of protons and  electrons but a different number of neutrons The way to represent an isotope is to write the chemical symbol (or the word) followed by a dash and then the mass number E.g. carbon-12 and carbon-14 are isotopes of carbon containing 6 and 8 neutrons respectively These isotopes could also be written as 12C or C-12, and 14C or C-14 respectively The atomic structure and symbols of the three isotopes of hydrogen Page 5 of 32 © 2015-2023 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to savemyexams.co.uk for more awesome resources Determining the Subatomic Structure of Atoms & Ions YOUR NOTES An atom is neutral and has no overall charge  Ions on the other hand have either gained or lost electrons causing them to become charged The number of subatomic particles in atoms and ions can be determined given their atomic (proton) number, mass (nucleon) number and charge Protons The atomic number of an atom and ion determines which element it is Therefore, all atoms and ions of the same element have the same number of protons (atomic number) in the nucleus E.g. lithium has an atomic number of 3 (three protons) whereas beryllium has atomic number of 4 (4 protons) The number of protons equals the atomic (proton) number The number of protons of an unknown element can be calculated by using its mass number and number of neutrons: Mass number = number of protons + number of neutrons Number of protons = mass number - number of neutrons  Worked Example Determine the number of protons of the following ions and atoms: 1. Mg2+ ion 2. Carbon atom 3. An unknown atom of element X with mass number 63 and 34 neutrons Answer: Answer 1: The atomic number of a magnesium atom is 12 suggesting that the number of protons in the magnesium element is 12 Therefore the number of protons in a Mg2+ ion is also 12 - the number of protons does not change when an ion is formed Answer 2: The atomic number of a carbon atom is 6 suggesting that a carbon atom has 6 protons in its nucleus Answer 3: Use the formula to calculate the number of protons Number of protons = mass number - number of neutrons Number of protons = 63 - 34 Number of protons = 29 Element X is therefore copper Page 6 of 32 © 2015-2023 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to savemyexams.co.uk for more awesome resources Electrons YOUR NOTES An atom is neutral and therefore has the same number of protons and electrons  Ions have a different number of electrons to the number of protons, depending on their charge A positively charged ion has lost electrons and therefore has fewer electrons than protons A negatively charged ion has gained electrons and therefore has more electrons than protons  Worked Example Determine the number of electrons of the following ions and atoms: 1. Mg2+ ion 2. Carbon atom 3. An unknown atom of element X with mass number 63 and 34 neutrons Answer: Answer 1: The atomic number of a magnesium atom is 12 suggesting that the number of protons in the neutral magnesium atom is 12 However, the 2+ charge in Mg2+ ion suggests it has lost two electrons It only has 10 electrons left now Answer 2: The atomic number of a carbon atom is 6 suggesting that the neutral carbon atom has 6 electrons orbiting around the nucleus Answer 3: The number of protons of element X can be calculated by: Number of protons = mass number - number of neutrons Number of protons = 63 - 34 Number of protons = 29 The neutral atom of element X therefore also has 29 electrons Neutrons The mass and atomic numbers can be used to find the number of neutrons in ions and atoms: Number of neutrons = mass number (A) - number of protons (Z) Page 7 of 32 © 2015-2023 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to savemyexams.co.uk for more awesome resources YOUR NOTES  Worked Example  Determine the number of neutrons of the following ions and atoms: 1. Mg2+ ion 2. Carbon atom 3. An unknown atom of element X with mass number 63 and 29 protons Answer: Answer 1: The atomic number of a magnesium atom is 12 and its mass number is 24 Number of neutrons = mass number (A) - number of protons (Z) Number of neutrons = 24 - 12 Number of neutrons = 12 The Mg2+ ion has 12 neutrons in its nucleus Answer 2: The atomic number of a carbon atom is 6 and its mass number is 12 Number of neutrons = mass number (A) - number of protons (Z) Number of neutrons = 12 - 6 Number of neutrons = 6 The carbon atom has 6 neutrons in its nucleus Answer 3: The atomic number of an element X atom is 29 and its mass number is 63 Number of neutrons = mass number (A) - number of protons (Z) Number of neutrons = 63 - 29 Number of neutrons = 34 The neutral atom of element X has 34 neutrons in its nucleus Page 8 of 32 © 2015-2023 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to savemyexams.co.uk for more awesome resources 2.1.3 Relative Atomic Mass Calculations YOUR NOTES  Relative Atomic Mass Calculations Isotopes are different atoms of the same element that contain the same number of protons and electrons but a different number of neutrons These are atoms of the same elements but with different mass numbers Because of this, the mass of an element is given as relative atomic mass (Ar) by using the average mass of all of the isotopes The relative atomic mass of an element can be calculated by using the percentage abundance values The percentage abundance of an isotope is either given or can be read off the mass spectrum Firstly, find the mass of 100 atoms by multiplying the percentage abundance by the mass of each isotope Secondly, divide by 100 to find the average atomic mass For example, if you have two isotopes A and B:  Worked Example A sample of oxygen contains the following isotopes What is the relative atomic mass of oxygen to 2 dp? A 16.00 B 17.18 C 16.09 D 17.00 Answer: The correct answer is A Total mass of 100 atoms = (99.76 x 16) + ( 0.04 x 17) + (0.20 x 18) = 1600.44 Page 9 of 32 © 2015-2023 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to savemyexams.co.uk for more awesome resources Mass of 1 atom = 1600.44 ÷ 100 = 16.0044 = 16.00 (2 dp) YOUR NOTES  Here is another example, but this time using a mass spectrum to obtain the information:  Worked Example Calculate the relative atomic mass of boron using its mass spectrum, to 2 dp: Answer: Total mass of 100 atoms = (19.9 x 10) + (80.1 x 11) = 1080.1 Mass of 1 atom = 1080.1 ÷ 100 = 10.801 = 10.80 (2 dp) Page 10 of 32 © 2015-2023 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to savemyexams.co.uk for more awesome resources 2.1.4 The Electromagnetic Spectrum YOUR NOTES  The Electromagnetic Spectrum The electromagnetic spectrum is a range of frequencies that covers all electromagnetic radiation and their respective wavelengths and energy It is divided into bands or regions, and is very important in analytical chemistry The spectrum shows the relationship between frequency, wavelength and energy Frequency is how many waves pass per second, and wavelength is the distance between two consecutive peaks on the wave Gamma rays, X-rays and UV radiation are all dangerous - you can see from that end of the spectrum that it is high frequency and high energy, which can be very damaging to your health All light waves travel at the same speed; what distinguishes them is their different frequencies The speed of light (symbol ‘c’) is constant and has a value of 3.00 x 108 ms-1 As you can see from the spectrum, frequency (symbol ‘ν') is inversely proportional to wavelength (symbol ‘λ') In other words, the higher the frequency, the shorter the wavelength The equation that links them is c = νλ Page 11 of 32 © 2015-2023 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to savemyexams.co.uk for more awesome resources Since c is constant you can use the formula to calculate the frequency of radiation given YOUR NOTES the wavelength, and vice versa  Continuous versus line spectrum A continuous spectrum in the visible region contains all the colours of the spectrum This is what you are seeing in a rainbow, which is formed by the refraction of white light through a prism or water droplets in rain A continuous spectrum shows all frequencies of light However, a line spectrum only shows certain frequencies The line spectrum of helium which shows only certain frequencies of light This tells us that the emitted light from atoms can only be certain fixed frequencies - it is quantised (quanta means 'little packet') Electrons can only possess certain amounts of energy - they cannot have any energy value  Exam Tip The formula that relates frequency and wavelength is printed in Section 1 of the IB Chemistry Data Booklet so you don’t need to learn itYou will also find the speed of light and other useful constants in Section 2 Page 12 of 32 © 2015-2023 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to savemyexams.co.uk for more awesome resources 2.1.5 Emission Spectra YOUR NOTES  Emission Spectra Electrons move rapidly around the nucleus in energy shells If their energy is increased, then they can jump to a higher energy level The process is reversible, so electrons can return to their original energy levels When this happens, they emit energy The frequency of energy is exactly the same, it is just being emitted rather than absorbed: The difference between absorption and emission depends on whether electrons are jumping from lower to higher energy levels or the other way around The energy they emit is a mixture of different frequencies This is thought to correspond to the many possibilities of electron jumps between energy shells If the emitted energy is in the visible region, it can be analysed by passing it through a diffraction grating The result is a line emission spectrum Line emission spectra The line emission (visible) spectrum of hydrogen Each line is a specific energy value This suggests that electrons can only possess a limited choice of allowed energies These packets of energy are called 'quanta' (plural quantum) What you should notice about this spectrum is that the lines get closer together towards the blue end of the spectrum This is called convergence and the set of lines is converging towards the higher energy end, so the electron is reaching a maximum amount of energy This maximum corresponds to the ionisation energy of the electron Page 13 of 32 © 2015-2023 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to savemyexams.co.uk for more awesome resources These lines were first observed by the Swiss school teacher Johannes Balmer, and they are YOUR NOTES named after him  We now know that these lines correspond to the electron jumping from higher levels down to the second or n = 2 energy level Page 14 of 32 © 2015-2023 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to savemyexams.co.uk for more awesome resources The Hydrogen Spectrum YOUR NOTES A larger version of the hydrogen spectrum from the infrared to the ultraviolet region looks  like this The full hydrogen spectrum In the spectrum, we can see sets or families of lines Balmer could not explain why the lines were formed - an explanation had to wait until the arrival of Planck's Quantum Theory in 1900 Niels Bohr applied the Quantum Theory to electrons in 1913, and proposed that electrons could only exist in fixed energy levels The line emission spectrum of hydrogen provided evidence of these energy levels and it was deduced that the families of lines corresponded to electrons jumping from higher levels to lower levels Page 15 of 32 © 2015-2023 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to savemyexams.co.uk for more awesome resources YOUR NOTES  Electron jumps in the hydrogen spectrum The findings helped scientists to understand how electrons work and provided the backbone to our knowledge of energy levels, sublevels and orbitals The jumps can be summarised as follows: Electron Jumps & Energy Table Page 16 of 32 © 2015-2023 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to savemyexams.co.uk for more awesome resources YOUR NOTES  Worked Example  Which electron transition in the hydrogen atom emits visible light? A. n = 1 to n = 2 B. n = 2 to n = 3 C. n = 2 to n = 1 D. n = 3 to n = 2 Answer: Option D is correct Emission in the visible region occurs for an electron jumping from any higher energy level to n = 2 Page 17 of 32 © 2015-2023 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to savemyexams.co.uk for more awesome resources 2.1.6 Energy Levels & Sublevels YOUR NOTES  Electron Energy Levels Shells The arrangement of electrons in an atom is called the electronic configuration Electrons are arranged around the nucleus in principal energy levels or principal quantum shells Principal quantum numbers (n) are used to number the energy levels or quantum shells The lower the principal quantum number, the closer the shell is to the nucleus The higher the principal quantum number, the lesser the energy of the shell Each principal quantum number has a fixed number of electrons it can hold n = 1 : up to 2 electrons n = 2 : up to 8 electrons n = 3 : up to 18 electrons n = 4 : up to 32 electrons There is a pattern here - the mathematical relationship between the number of electrons and the principal energy level is 2n2 So for example, in the third shell n = 3 and the number of electrons is 2 x (32 ) = 18 Electrons are arranged in principal quantum shells, which are numbered by principal quantum numbers Subshells Page 18 of 32 © 2015-2023 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to savemyexams.co.uk for more awesome resources The principal quantum shells are split into subshells which are given the letters s, p and d YOUR NOTES Elements with more than 57 electrons also have an f subshell  The energy of the electrons in the subshells increases in the order s < p < d The order of subshells overlap for the higher principal quantum shells as seen in the diagram below: Electrons are arranged in principal quantum shells, which are numbered by principal quantum numbers Orbitals The subshells contain one or more atomic orbitals Orbitals exist at specific energy levels and electrons can only be found at these specific levels, not in between Each atomic orbital can be occupied by a maximum of two electrons The orbitals have specific 3D shapes Page 19 of 32 © 2015-2023 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to savemyexams.co.uk for more awesome resources YOUR NOTES  Representation of orbitals (the dot represents the nucleus of the atom) showing spherical s orbitals (a), p orbitals containing ‘lobes’ along the x, y and z axis Note that the shape of the d orbitals is not required for IB Chemistry Page 20 of 32 © 2015-2023 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to savemyexams.co.uk for more awesome resources YOUR NOTES  An overview of the shells, subshells and orbitals in an atom Ground state The ground state is the most stable electronic configuration of an atom which has the lowest amount of energy This is achieved by filling the subshells of energy with the lowest energy first (1s) - this is called the Aufbau Principle The order of the subshells in terms of increasing energy does not follow a regular pattern at n= 3 and higher Page 21 of 32 © 2015-2023 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to savemyexams.co.uk for more awesome resources The Aufbau Principle - following the arrows gives you the filling order YOUR NOTES  Sublevels & Energy The principal quantum shells increase in energy with increasing principal quantum number Eg. n = 4 is higher in energy than n = 2 The subshells increase in energy as follows: s ＜ p ＜ d ＜ f The only exception to these rules is the 3d orbital which has slightly higher energy than the 4s orbital, so the 4s orbital is filled before the 3d orbital All the orbitals in the same subshell have the same energy and are said to be degenerate Eg. px, py and pz are all equal in energy Relative energies of the shells and subshells Page 22 of 32 © 2015-2023 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to savemyexams.co.uk for more awesome resources 2.1.7 Sublevels & Orbitals YOUR NOTES  Electron Orbitals Each shell can be divided further into subshells, labelled s, p, d and f Each subshell can hold a specific number of orbitals: s subshell : 1 orbital p subshell : 3 orbitals labelled px, py and pz d subshell : 5 orbitals f subshell : 7 orbitals Each orbital can hold a maximum number of 2 electrons so the maximum number of electrons in each subshell are as follows: s : 1 x 2 = total of 2 electrons p : 3 x 2 = total of 6 electrons d : 5 x 2 = total of 10 electrons f : 7 x 2 = total of 14 electrons In the ground state, orbitals in the same subshell have the same energy and are said to be degenerate, so the energy of a px orbital is the same as a py orbital Page 23 of 32 © 2015-2023 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to savemyexams.co.uk for more awesome resources YOUR NOTES  Shells are divided into subshells which are further divided into orbitals Summary of the Arrangement of Electrons in Atoms Table Page 24 of 32 © 2015-2023 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to savemyexams.co.uk for more awesome resources YOUR NOTES  Page 25 of 32 © 2015-2023 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to savemyexams.co.uk for more awesome resources The s & p Orbitals YOUR NOTES s orbitals  The s orbitals are spherical in shape The size of the s orbitals increases with increasing shell number E.g. the s orbital of the third quantum shell (n = 3) is bigger than the s orbital of the first quantum shell (n = 1) The s orbitals become larger with increasing principal quantum number p orbitals The p orbitals are dumbbell-shaped Every shell has three p orbitals except for the first one (n = 1) The p orbitals occupy the x, y and z axes and point at right angles to each other, so are oriented perpendicular to one another The lobes of the p orbitals become larger and longer with increasing shell number The p orbitals become larger and longer with increasing principal quantum number Page 26 of 32 © 2015-2023 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to savemyexams.co.uk for more awesome resources 2.1.8 Writing Electron Configurations YOUR NOTES  Electron Configurations: Basics The electron configuration gives information about the number of electrons in each shell, subshell and orbital of an atom The subshells are filled in order of increasing energy The electron configuration shows the number of electrons occupying a subshell in a specific shell Page 27 of 32 © 2015-2023 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to savemyexams.co.uk for more awesome resources Electron Configurations: Explained YOUR NOTES Electrons can be imagined as small spinning charges which rotate around their own axis in  either a clockwise or anticlockwise direction The spin of the electron is represented by its direction The spin creates a tiny magnetic field with N-S pole pointing up or down Electrons can spin either in a clockwise or anticlockwise direction around their own axis Electrons with the same spin repel each other which is also called spin-pair repulsion Therefore, electrons will occupy separate orbitals in the same subshell first to minimise this repulsion and have their spin in the same direction They will then pair up, with a second electron being added to the first p orbital, with its spin in the opposite direction This is known as Hund's Rule E.g. if there are three electrons in a p subshell, one electron will go into each px, py and pz orbital Electron configuration: three electrons in a p subshell The principal quantum number indicates the energy level of a particular shell but also indicates the energy of the electrons in that shell A 2p electron is in the second shell and therefore has an energy corresponding to n = 2 Page 28 of 32 © 2015-2023 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to savemyexams.co.uk for more awesome resources Even though there is repulsion between negatively charged electrons, they occupy the YOUR NOTES same region of space in orbitals  An orbital can only hold two electrons and they must have opposite spin - the is known as the Pauli Exclusion Principle This is because the energy required to jump to a higher empty orbital is greater than the inter-electron repulsion For this reason, they pair up and occupy the lower energy levels first Orbital Diagrams The electron configuration can also be represented using the orbital spin diagrams Each box represents an atomic orbital The boxes are arranged in order of increasing energy from lower to higher (i.e. starting from closest to the nucleus) The electrons are represented by opposite arrows to show the spin of the electrons E.g. the box notation for titanium is shown below The electrons in titanium are arranged in their orbitals as shown. Electrons occupy the lowest energy levels first before filling those with higher energy Page 29 of 32 © 2015-2023 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to savemyexams.co.uk for more awesome resources Determining Electronic Configurations YOUR NOTES Writing out the electronic configuration tells us how the electrons in an atom or ion are  arranged in their shells, subshells and orbitals This can be done using the full electron configuration or the shorthand version The full electron configuration describes the arrangement of all electrons from the 1s subshell up The shorthand electron configuration includes using the symbol of the nearest preceding noble gas to account for however many electrons are in that noble gas, followed by the rest of the electron configuration Ions are formed when atoms lose or gain electrons Negative ions are formed by adding electrons to the outer subshell Positive ions are formed by removing electrons from the outer subshell The transition metals fill the 4s subshell before the 3d subshell, but they also lose electrons from the 4s first rather than from the 3d subshell The Periodic Table is split up into four main blocks depending on their electronic configuration: s block elements (valence electron(s) in s orbital) p block elements (valence electron(s) in p orbital) d block elements (valence electron(s) in d orbital) f block elements (valence electron(s) in f orbital) The elements can be divided into four blocks according to their outer shell electron configuration Exceptions to the Aufbau Principle Chromium and copper have the following electron configurations: Cr is [Ar] 3d5 4s1 not [Ar] 3d4 4s2 Cu is [Ar] 3d10 4s1 not [Ar] 3d9 4s2 Page 30 of 32 © 2015-2023 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to savemyexams.co.uk for more awesome resources This is because the [Ar] 3d5 4s1 and [Ar] 3d10 4s1 configurations are energetically YOUR NOTES favourable  By promoting an electron from 4s to 3d, these atoms achieve a half full or full d-subshell, respectively  Worked Example Write down the full and shorthand electron configuration of the following elements: 1. Potassium 2. Calcium 3. Gallium 4. Ca2+ Answer: Answer 1: Potassium has 19 electrons so the full electronic configuration is: 1s2 2s2 2p6 3s2 3p6 4s1 The 4s orbital is lower in energy than the 3d subshell and is therefore filled first The nearest preceding noble gas to potassium is argon which accounts for 18 electrons so the shorthand electron configuration is: [Ar] 4s1 Answer 2: Calcium has 20 electrons so the full electronic configuration is: 1s2 2s2 2p6 3s2 3p6 4s2 The 4s orbital is lower in energy than the 3d subshell and is therefore filled first The shorthand version is [Ar] 4s2 since argon is the nearest preceding noble gas to calcium which accounts for 18 electrons Answer 3: Gallium has 31 electrons so the full electronic configuration is: [Ar] 3d10 4s2 4p1 Answer 4: If you ionise calcium and remove two of its outer electrons, the electronic configuration of the Ca2+ ion is identical to that of argon: Ca2+ is 1s2 2s2 2p6 3s2 3p6 Ar is also 1s2 2s2 2p6 3s2 3p6 Page 31 of 32 © 2015-2023 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to savemyexams.co.uk for more awesome resources Exam Tip YOUR NOTES  Orbital spin diagrams can be drawn horizontally or vertically, going up or down the  page - there is no hard and fast rule about this. The important thing is that you label the boxes and have the right number of electrons shown. The arrows you use for electrons can be full or half-headed arrows, but they must be in opposite directions in the same box. Page 32 of 32 © 2015-2023 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers