For a given value of n, the ns orbital is always lower in energy than the np orbitals, which are lower in energy than the nd orbitals, and so forth. Identify the location of the elements in the periodic table. Now recall you are back in the third row and your view is blocked by the people in the two rows in front of you. Although it is not possible to measure an ionic radius directly for the same reason it is not possible to directly measure an atom’s radius, it is possible to measure the distance between the nuclei of a cation and an adjacent anion in an ionic compound to determine the ionic radius (the radius of a cation or anion) of one or both. For more information contact us at info@libretexts.org or check out our status page at https://status.libretexts.org. We can think of effective nuclear charge as the positive charge felt by the outermost electrons in an atom. It isn't just about distance: there is also something else going on. Well, that didn't work! You can pretty much find anything here. In this section, we discuss how atomic and ion “sizes” are defined and obtained. And just like our people, these core electrons shield the outer electrons from the full nuclear charge. increases down a group, decreases across a period. Click here to let us know! The periodic table is the arrangement of elements in rows and columns according to atomic numbers. Again, principal shells with larger values of n lie at successively greater distances from the nucleus. Conversely, adding one or more electrons to a neutral atom causes electron–electron repulsions to increase and the effective nuclear charge to decrease, so the size of the probability region increases and the ion expands (compare F at 42 pm with F− at 133 pm). Trends across a period follow from the increasing number of protons in the nucleus and the decrease in radius. The designations cation or anion come from the early experiments with electricity which found that positively charged particles were attracted to the negative pole of a battery, the cathode, while negatively charged ones were attracted to the positive pole, the anode. Because elements in different columns tend to form ions with different charges, it is not possible to compare ions of the same charge across a row of the periodic table. If different numbers of electrons can be removed to produce ions with different charges, the ion with the greatest positive charge is the smallest (compare Fe2+ at 78 pm with Fe3+ at 64.5 pm). Effective Nuclear Charge (ENC) –effectiveness of the positive pull of the nucleus on the electrons of an atom: •Across a period: –no change in shielding –increase in # of positive charges (protons) in nucleus increases pulling force •Down a group: –increased shielding –increased # protons in nucleus counteracted by increased shielding 1 month ago. If, on the other hand, an electron is very close to the nucleus, then at any given moment most of the other electrons are farther from the nucleus and do not shield the nuclear charge. As we will discuss later on in the chapter, this phenomenon can explain the decrease in atomic radii we see as we go across the periodic table as electrons are held closer to the nucleus due to increase in number of protons and increase in effective nuclear charge. This means the simple approximation (Equation \ref{simple}) overestimates the shielding constant $$S$$. You rush down to watch him, but when you arrive, quite a few people are already there. 7th Grade World History: Enrichment Program, Algebra I Curriculum Resource & Lesson Plans, 10th Grade English: Homework Help Resource, DSST Technical Writing: Study Guide & Test Prep, Quiz & Worksheet - Evaluating Independence with Two-Way Tables, Quiz & Worksheet - Economic & Social Factors in City Land Use, Quiz & Worksheet - Population Density Around the World, Quiz & Worksheet - Distance, Time & Speed, Quiz & Worksheet - History of the 26th Amendment, Reapportionment & Redistricting for Congressional Constituencies: Definition & Process, Medal of Honor Recipient Sergeant Salvatore Giunta, English Language Learning Programs in California, How to Study for a Placement Test for College, Tech and Engineering - Questions & Answers, Health and Medicine - Questions & Answers, The first excited state of the nucleus of uranium-235 is 0.051 MeV above the ground state. Aluminum and silicon are both in the third row with aluminum lying to the left, so silicon is smaller than aluminum (Si < Al) because its effective nuclear charge is greater. Atomic radii decrease from left to right across a row because of the increase in effective nuclear charge due to poor electron screening by other electrons in the same principal shell. Each change in shell number is a new group; s and p subshells are in the same group but d and f orbitals are their own group. Ionic radii follow the same vertical trend as atomic radii; that is, for ions with the same charge, the ionic radius increases going down a column. All six of the ions contain 10 electrons in the 1s, 2s, and 2p orbitals, but the nuclear charge varies from +7 (N) to +13 (Al). Ionization energy generally increases across a period and decreases down a group. An ion is formed when either one or more electrons are removed from a neutral atom to form a positive ion (cation) or when additional electrons attach themselves to neutral atoms to form a negative one (anion). Each species has 10 electrons, and the number of nonvalence electrons is 2 (10 total electrons - 8 valence), but the effective nuclear charge varies because each has a different atomic number $$A$$. This is an application of Equations \ref{4} and \ref{2.6.0}. Determine which ions form an isoelectronic series. This also suggests that $$\mathrm{Na}^+$$ has the smallest radius of these species and that is correct. What is the effective nuclear charge on the Cl atom? a. atomic number b.principal quantum number c.effective nuclear charge d.number of electrons 8.6: Periodic Trends in the Size of Atoms and Effective Nuclear Charge, https://chem.libretexts.org/@app/auth/3/login?returnto=https%3A%2F%2Fchem.libretexts.org%2FBookshelves%2FGeneral_Chemistry%2FMap%253A_A_Molecular_Approach_(Tro)%2F08%253A_Periodic_Properties_of_the_Elements%2F8.06%253A_Periodic_Trends_in_the_Size_of_Atoms_and_Effective_Nuclear_Charge. How does overall nuclear charge change across a period? The atoms in the second row of the periodic table (Li through Ne) illustrate the effect of electron shielding. The sizes of the ions in this series decrease smoothly from N3− to Al3+. The further away the energy level is from the nucleus, the higher the energy. 1) Electrons in a 'p orbital' are more effective than those in other orbitals at shielding other electrons fro. Select who you are below, and we'll recommend a plan for you. 88 chapters | where Z is the actual nuclear charge (the atomic number) and Z e f f is the effective nuclear charge. The same dynamic is responsible for the steady increase in size observed as we go down the other columns of the periodic table. We can see from Equation \ref{4} that the effective nuclear charge of an atom increases as the number of protons in an atom increases (Figure $$\PageIndex{2}$$). The peak for the filled n = 1 shell occurs at successively shorter distances for neon (Z = 10) and argon (Z = 18) because, with a greater number of protons, their nuclei are more positively charged than that of helium. The shielding constant can be estimated by totaling the screening by all electrons ($$n$$) except the one in question. Anions are larger than their neutral atoms. The electron in the 2s energy level is further away from the nucleus and has highest energy. Consider the elements S, Sn, Si and Sr. The people in front of you are getting in your way. This understanding of the shielding effect can be used to predict periodic trends in both ionization energy and atomic radius. An atom consists of a positively charged nucleus, containing protons and neutrons, surrounded by negatively charged electrons. That force depends on the effective nuclear charge experienced by the the inner electrons. where, Zeff is the effective nuclear charge, Z is the actual nuclear charge, and σ is the shielding constant, where the shielding constant is greater than zero but smaller than Z. In this lesson, we learn about the effective nuclear charge, and its effect on the properties of atoms. A similar approach for measuring the size of ions is discussed later in this section. Electrons in the closest, or inner, filled energy levels are the core electrons. 2. The $$Z_{eff}$$ in Table $$\PageIndex{1}$$ for $$Z_\mathrm{eff}(\mathrm{Na}$$ is 10.63 and appreciables larger than the 8 estimated above. It has the electron configuration of neon (Ne): 3s^2 3p^5. In contrast, neon, with filled n = 1 and 2 principal shells, has two peaks. Argon, with filled n = 1, 2, and 3 principal shells, has three peaks. Figure $$\PageIndex{1}$$ also shows that there are distinct peaks in the total electron density at particular distances and that these peaks occur at different distances from the nucleus for each element. An anomaly to the general trend occurs in a period in going from the Group 1A element to the Group 2A element and in going from the Group 4A to Group 5A. This time the atomic number and number of electrons is increasing across the period. What do you think might happen if you move closer to the second row to try and see better? Determine the relative sizes of elements located in the same column from their principal quantum number. Chlorine has the atomic number of 17. Give the trend for atomic radius across a period and down a group, and explain each of these trends in terms of attraction, repulsion, and effective - 2848851 Jasamara2010 Jasamara2010 02/09/2017 Chemistry ... and effective nuclear charge. B Combining the two inequalities gives the overall order: C < Si < Al. Such a set of species is known as an isoelectronic series. This effect is called electron shielding. Rb, In, Sb, Sn, Sr, Answer the following according to increasing: A) increasing size: S, Se, Se^2-, S^2- B) increasing lattice energy: MgO, CaO, Al2O3, NaCl C) increasing 1st ionization energy: S, Si, Ar, Na, Mg, Calculate Z* of Ce[Xe] 6s^2 4f^1 5d^1 4f and 6s. Rank the following elements by effective nuclear charge,Zeff, for a valence electron from highest to lowest Zeff: Po, Rn, Ba, Bi, Pb, for 208 pb find the net electrical charge of the nucleus, Working Scholars® Bringing Tuition-Free College to the Community. Trends down a group follow from the increasing number of electron shells and the increased distance of the outer electrons from the nucleus. On the periodic table, first ionization energy generally decreases as you move down a group. We have already said that there are different energy levels in an atom, and electrons always go into the lowest energy level first until the energy level is filled. The radius of sodium in each of its three known oxidation states is given in Table $$\PageIndex{1}$$. Of those ions, predict their relative sizes based on their nuclear charges. Postby Chem_Mod » Thu Oct 26, 2017 1:01 am Effective nuclear charge Zeff decreases down a column because of the extra shielding of the previous energy level n. The d block electrons contribute to this shielding. Covalent atomic radii can be determined for most of the nonmetals, but how do chemists obtain atomic radii for elements that do not form covalent bonds? We assign half of this distance to each chlorine atom, giving chlorine a covalent atomic radius ($$r_{cov}$$), which is half the distance between the nuclei of two like atoms joined by a covalent bond in the same molecule, of 99 pm or 0.99 Å (Figure $$\PageIndex{2a}$$). Services. Cations are always smaller than the neutral atom and anions are always larger. An uncharged, or neutral, gold atom, has 79 protons and 79 electrons. This is somewhat difficult for helium which does not form a solid at any temperature. Of electrons in the inner shell keep on increasing, now the inner shell electron always repels the electron in the outter shell away from the nucleus and so as we go down the net charge faced by each valence shell electron after the battle between the repulsive and attractive forces,called “effective nuclear charge”, decreases as the number of electrons in inner orbits increase faster than the nuclear charge at centre which also results in increase in size down the group. Trend of Electronegativity Down A Group. K+, Cl−, and S2− form an isoelectronic series with the [Ar] closed-shell electron configuration; that is, all three ions contain 18 electrons but have different nuclear charges. The approximation in Equation \ref{simple} is a good first order description of electron shielding, but the actual $$Z_{eff}$$ experienced by an electron in a given orbital depends not only on the spatial distribution of the electron in that orbital but also on the distribution of all the other electrons present. Because K+ has the greatest nuclear charge (Z = 19), its radius is smallest, and S2− with Z = 16 has the largest radius. 1. the trend in _______ accounts for the increase in ionization energy across a period.? Quiz & Worksheet - Overview of Lewis Dot Structures, Flashcards - Real Estate Marketing Basics, Flashcards - Promotional Marketing in Real Estate. Penetration describes the proximity to which an electron can approach to the nucleus. At $$r ≈ 0$$, the positive charge experienced by an electron is approximately the full nuclear charge, or $$Z_{eff} ≈ Z$$. For example, a 1s electron (Figure $$\PageIndex{3}$$; purple curve) has greater electron density near the nucleus than a 2p electron (Figure $$\PageIndex{3}$$; red curve) and has a greater penetration. To calculate σ, we will write out all the orbitals in an atom, separating them into "groups". The Na+ ion is significantly smaller than the neutral Na atom because the 3s1 electron has been removed to give a closed shell with n = 2. size of an atom. 2 Electrons that are shielded from the full charge of the nucleus experience an effective nuclear charge ($$Z_{eff}$$) of the nucleus, which is some degree less than the full nuclear charge an electron would feel in a hydrogen atom or hydrogenlike ion. Because of the effects of shielding and the different radial distributions of orbitals with the same value of n but different values of l, the different subshells are not degenerate in a multielectron atom. The first energy level is closest to the nucleus, and just like the first row in our street performance, feels the greatest effect. In this way the 2s electron can "avoid" some of the shielding effect of the inner 1s electron. From Equations \ref{4} and \ref{2.6.0}, $$Z_{eff}$$ for a specific electron can be estimated is the shielding constants for that electron of all other electrons in species is known. Going down a group, distance and shielding increase. What Can You Do With a Master's in Education? the other trend occurs when you move from the top of the periodic table down (moving within a group These effects are the underlying basis for the periodic trends in elemental properties that we will explore in this chapter. The outer electron configuration in a group is the same, but as you go down the group there are more and more core electrons. Sciences, Culinary Arts and Personal Let's do a very quick review on electrons. The energy of the n = 1 shell also decreases tremendously (the filled 1s orbital becomes more stable) as the nuclear charge increases. Down a group the number of protons and therefore nuclear charge increases, BUT the number of shielding electrons increases more significantly, which more than counteracts the increasing nuclear charge... resulting in an overall decreasing effective nuclear charge down a group. These core electrons are held close to the nucleus and the shielding effect increases. We learned that effective nuclear charge is the positive charge felt by the outermost electrons in an atom. Irregularities can usually be explained by variations in effective nuclear charge. For example, we would predict a carbon–chlorine distance of 77 pm + 99 pm = 176 pm for a C–Cl bond, which is very close to the average value observed in many organochlorine compounds. The reason behind this behavioral trend of ionic radius can be attributed to the increase in effective nuclear charge on moving across the period. Atomic radii decrease from left to right across a row and increase from top to bottom down a column. The increased nuclear charge as you go down the group is offset by extra screening electrons. We also acknowledge previous National Science Foundation support under grant numbers 1246120, 1525057, and 1413739. This related to the shielding constants since the 1s electrons are closer to the nucleus than a 2p electron, hence the 1s screens a 2p electron almost perfectly ($$S=1$$. The atomic radius is decreasing. Among all the elements, fluorine is the most electronegative element. The atomic number increases and the amount of shells does not FQ: What can we learn about radius, ionization, and electronegativity from periodic trends? Data from E. Clementi and D. L. Raimondi; The Journal of Chemical Physics 38, 2686 (1963). Electrons that have greater penetration can get closer to the nucleus and effectively block out the charge from electrons that have less proximity. In general, the electronegativity decreases on going down the group. Both contributions can be explained by the change in effective nuclear charge. From where you are sitting, you can only see glimpses of the act. This is known as the shielding effect. It has 9 protons in the nucleus. Effective Nuclear Charge (Z eff) remains constant. Moreover, atomic radii increase from top to bottom down a column because the effective nuclear charge remains relatively constant as the principal quantum number increases. Use the simple approximation for shielding constants. Effective nuclear charge decreases as you move down the group due to addition of an extra shell due to which distance between valence electrons and nucleus decreases These methods produce sets of ionic radii that are internally consistent from one ionic compound to another, although each method gives slightly different values. Trends in atomic size result from differences in the effective nuclear charges ($$Z_{eff}$$) experienced by electrons in the outermost orbitals of the elements. This leads to large differences in $$Z_{eff}$$ for different elements, as shown in Figure $$\PageIndex{2}$$ for the elements of the first three rows of the periodic table. Asked for: arrange in order of increasing atomic radius. (More detailed calculations give a value of Zeff = +1.26 for Li.) The concepts of electron shielding, orbital penetration and effective nuclear charge were introduced above, but we did so in a qualitative manner (e.g., Equations \ref{better1} and \ref{better2}). Based on their positions in the periodic table, arrange these ions in order of increasing radius: Cl−, K+, S2−, and Se2−. The increase in atomic size going down a column is also due to electron shielding, but the situation is more complex because the principal quantum number n is not constant. Quiz & Worksheet - What is the Fairness Doctrine? © copyright 2003-2021 Study.com. Create your account. They are in the inner shells. Simply put, because the outer electron is not held so tightly, it is easier to remove. Carbon and silicon are both in group 14 with carbon lying above, so carbon is smaller than silicon (C < Si). Consequently, beryllium is significantly smaller than lithium. It is this number that gives the atom its unique identity. What is the effective attraction $$Z_{eff}$$ experienced by the valence electrons in the three isoelectronic species: the fluorine anion, the neutral neon atom, and sodium cation? The Na− ion is larger than the parent Na atom because the additional electron produces a 3s2 valence electron configuration, while the nuclear charge remains the same. The charge $$Z$$ of the nucleus of a fluorine atom is 9, but the valence electrons are screened appreciably by the core electrons (four electrons from the 1s and 2s orbitals) and partially by the 7 electrons in the 2p orbitals. Instead of rows of people, there are rows of energy levels of electrons. The two trends we will look at are atomic radius and ionization energy. The atomic number (Z) is the number of protons in the nucleus of an atom. Exercise $$\PageIndex{1}$$: Sodium Species. Enrolling in a course lets you earn progress by passing quizzes and exams. Ionic radii share the same vertical trend as atomic radii, but the horizontal trends differ due to differences in ionic charges. The atomic number also tells us the number of electrons in an uncharged atom. Tags. Try refreshing the page, or contact customer support. Number of valence electrons for group 2 atoms. Moreover, atomic radii increase from top to bottom down a column because the effective nuclear charge remains relatively constant as the principal quantum number increases. In the periodic table, atomic radii decrease from left to right across a row and increase from top to bottom down a column. But what is happening as you go across a period? We learned that core electrons shield outer electrons from the nuclear charge. credit by exam that is accepted by over 1,500 colleges and universities. Because the outer electron is held more tightly, it is more difficult to remove, and ionization energy increases across a period. Because the 1s2 shell is closest to the nucleus, its electrons are very poorly shielded by electrons in filled shells with larger values of n. Consequently, the two electrons in the n = 1 shell experience nearly the full nuclear charge, resulting in a strong electrostatic interaction between the electrons and the nucleus. As the distance between an electron and the nucleus approaches infinity, $$Z_{eff}$$ approaches a value of 1 because all the other ($$Z − 1$$) electrons in the neutral atom are, on the average, between it and the nucleus. The ionic radii of cations and anions are always smaller or larger, respectively, than the parent atom due to changes in electron–electron repulsions, and the trends in ionic radius parallel those in atomic size. A comparison of ionic radii with atomic radii (Figure $$\PageIndex{7}$$) shows that a cation, having lost an electron, is always smaller than its parent neutral atom, and an anion, having gained an electron, is always larger than the parent neutral atom. The size of the atom decreases because the outer electrons are held more tightly. Thus the single 2s electron in lithium experiences an effective nuclear charge of approximately +1 because the electrons in the filled 1s2 shell effectively neutralize two of the three positive charges in the nucleus. Shielding electrons are all non-valence electrons. In other words, penetration depends on the shell ($$n$$) and subshell ($$l$$). As we go down the column of the group 1 elements, the principal quantum number n increases from 2 to 6, but the nuclear charge increases from +3 to +55! We will finish by doing a quick example. Not sure what college you want to attend yet? In a similar approach, we can use the lengths of carbon–carbon single bonds in organic compounds, which are remarkably uniform at 154 pm, to assign a value of 77 pm as the covalent atomic radius for carbon. Effective nuclear charge is the attractive positive charge of nuclear protons acting on valence electrons. So let's review. Over 83,000 lessons in all major subjects, {{courseNav.course.mDynamicIntFields.lessonCount}}, Excited State in Chemistry: Definition & Overview, Diamagnetism & Paramagnetism: Definition & Explanation, Four Quantum Numbers: Principal, Angular Momentum, Magnetic & Spin, Electron Configurations in Atomic Energy Levels, Ground State Electron Configuration: Definition & Example, Electron Configurations in the s, p & d Orbitals, Energy & Momentum of a Photon: Equation & Calculations, Biological and Biomedical For an atom or an ion with only a single electron, we can calculate the potential energy by considering only the electrostatic attraction between the positively charged nucleus and the negatively charged electron. This happens to me all the time: I only get a good view, or feel the effect of the performance, if I am close to the performer. The same effect is also true of electrons and atoms. All have a filled 1s2 inner shell, but as we go from left to right across the row, the nuclear charge increases from +3 to +10. just create an account. For the same shell value ($$n$$) the penetrating power of an electron follows this trend in subshells (Figure $$\PageIndex{3}$$): for different values of shell (n) and subshell (l), penetrating power of an electron follows this trend: $\ce{1s > 2s > 2p > 3s > 3p > 4s > 3d > 4p > 5s > 4d > 5p > 6s > 4f ...} \label{better2}$. In contrast, the two 2s electrons in beryllium do not shield each other very well, although the filled 1s2 shell effectively neutralizes two of the four positive charges in the nucleus. The reason is the same as for atomic radii: shielding by filled inner shells produces little change in the effective nuclear charge felt by the outermost electrons. The trend on the periodic table is to increase across a period and increase down a group. Did you know… We have over 220 college This is due to the decrease in the effective nuclear charge with increase in the atomic number in a group. Atomic radius generally increases down a group. Because it is impossible to measure the sizes of both metallic and nonmetallic elements using any one method, chemists have developed a self-consistent way of calculating atomic radii using the quantum mechanical functions. The neon atom in this isoelectronic series is not listed in Table $$\PageIndex{3}$$, because neon forms no covalent or ionic compounds and hence its radius is difficult to measure. Because distances between the nuclei in pairs of covalently bonded atoms can be measured quite precisely, however, chemists use these distances as a basis for describing the approximate sizes of atoms. Consequently, we must use approximate methods to deal with the effect of electron-electron repulsions on orbital energies. An easy way to remember this formula is by using the following rhyme: Zebras Ecstatically Enjoy Zesty Moving Spacemen. However, the application of these rules is outside the scope of this text. In group 1, for example, the size of the atoms increases substantially going down the column. 882 lessons The trend is reversed. Once a semester I use Study.com to prepare for all my finals. When one or more electrons is removed from a neutral atom, two things happen: (1) repulsions between electrons in the same principal shell decrease because fewer electrons are present, and (2) the effective nuclear charge felt by the remaining electrons increases because there are fewer electrons to shield one another from the nucleus. flashcard sets, {{courseNav.course.topics.length}} chapters | For example, a fluorine atom has an electronic structure of 1s 2 2s 2 2p x 2 2p y 2 2p z 1. decreases down a group, increases across a period. An anomaly can also be found between rows 2 and 3 in going down Give an explanation for each of these anomalies. {{courseNav.course.mDynamicIntFields.lessonCount}} lessons A simple model is shown here. Let's first remind ourselves about the atom. This point is illustrated in Figure $$\PageIndex{1}$$ which shows a plot of total electron density for all occupied orbitals for three noble gases as a function of their distance from the nucleus. Unless otherwise noted, LibreTexts content is licensed by CC BY-NC-SA 3.0. When you create an account with Study.com, you get access to any resource you Both contributions can be explained by the change in effective nuclear charge. The size of the atom increases because the outer electrons are not held so tightly. You may also recall these levels are further divided into sub-levels, called s, p, d, and f. These sub-levels can take 2, 6, 10, and 14 electrons, respectively. The effective nuclear charge is the charge of the nucleus felt by the valence electron. But unlike core electrons, shielding between electrons in the same energy level is poor. On the basis of their positions in the periodic table, arrange these elements in order of increasing atomic radius: aluminum, carbon, and silicon. Going down a group, distance and shielding A These elements are not all in the same column or row, so we must use pairwise comparisons. One is the first energy level, then two, and so on. Fluorine has the atomic number of 9. All three species have a nuclear charge of +11, but they contain 10 (Na+), 11 (Na0), and 12 (Na−) electrons. To unlock this lesson you must be a Study.com Member. For similar reasons, the filled n = 2 shell in argon is located closer to the nucleus and has a lower energy than the n = 2 shell in neon. Get access risk-free for 30 days, Visit the MCAT Test: Practice and Study Guide page to learn more. This helps us predict periodic trends. An atom such as chlorine has both a covalent radius (the distance between the two atoms in a $$\ce{Cl2}$$ molecule) and a van der Waals radius (the distance between two Cl atoms in different molecules in, for example, $$\ce{Cl2(s)}$$ at low temperatures). To understand periodic trends in atomic radii. Nicky has a PhD in Physical Chemistry. The result is a steady increase in the effective nuclear charge and a steady decrease in atomic size (Figure $$\PageIndex{5}$$). flashcard set{{course.flashcardSetCoun > 1 ? All other trademarks and copyrights are the property of their respective owners. ionization energy trend. Although some people fall into the trap of visualizing atoms and ions as small, hard spheres similar to miniature table-tennis balls or marbles, the quantum mechanical model tells us that their shapes and boundaries are much less definite than those images suggest. Here you can see gold, and there are two numbers shown. Because helium has only one filled shell (n = 1), it shows only a single peak. For example, the electron configuration for the lithium atom is 1s^2 2s^1and the electron configuration for the Fluorine atom is 1s^2 2s^2 2p^5.
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