Electron configurations in which all of the electrons are in their lowest-energy configurations are known as ground state configurations. If an electron absorbs energy, it can move into a higher-energy orbital, producing an excited state configuration.
For atoms with a large number of electrons, the complete electron can be very cumbersome, and not very informative. For instance, the complete configuration of the element radium is. With a description like that, you'd be radioactive too!
As we go across the row from left to right, electrons are added to the 3 d subshell to neutralize the increase in the positive charge of the nucleus as the atomic number increases. Unexpectedly, however, chromium has a 4 s 1 3 d 5 electron configuration rather than the 4 s 2 3 d 4 configuration predicted by the aufbau principle, and copper is 4 s 1 3 d 10 rather than 4 s 2 3 d 9.
In Chapter 7 "The Periodic Table and Periodic Trends" , we attributed these anomalies to the extra stability associated with half-filled subshells. Table In the second-row transition metals, electron—electron repulsions within the 4 d subshell cause additional irregularities in electron configurations that are not easily predicted. For example, Nb and Tc, with atomic numbers 41 and 43, both have a half-filled 5 s subshell, with 5 s 1 4 d 4 and 5 s 1 4 d 6 valence electron configurations, respectively.
Further complications occur among the third-row transition metals, in which the 4 f , 5 d , and 6 s orbitals are extremely close in energy. Although La has a 6 s 2 5 d 1 valence electron configuration, the valence electron configuration of the next element—Ce—is 6 s 2 5 d 0 4 f 2. From this point through element 71, added electrons enter the 4 f subshell, giving rise to the 14 elements known as the lanthanides.
After the 4 f subshell is filled, the 5 d subshell is populated, producing the third row of the transition metals. Next comes the seventh period, where the actinides have three subshells 7 s , 6 d , and 5 f that are so similar in energy that their electron configurations are even more unpredictable. As we saw in the s -block and p -block elements, the size of neutral atoms of the d -block elements gradually decreases from left to right across a row, due to an increase in the effective nuclear charge Z eff with increasing atomic number.
In addition, the atomic radius increases down a group, just as it does in the s and p blocks. Because of the lanthanide contraction , however, the increase in size between the 3 d and 4 d metals is much greater than between the 4 d and 5 d metals Figure Figure Because of the lanthanide contraction, the second- and third-row transition metals are very similar in size.
Consequently, the ionization energies of these elements increase very slowly across a given row Figure 7. In addition, as we go from the top left to the bottom right corner of the d block, electronegativities generally increase, densities and electrical and thermal conductivities increase, and enthalpies of hydration of the metal cations decrease in magnitude, as summarized in Figure The relatively high ionization energies and electronegativities and relatively low enthalpies of hydration are all major factors in the noble character of metals such as Pt and Au.
The electronegativity of the elements increases, and the hydration energies of the metal cations decrease in magnitude from left to right and from top to bottom of the d block. The similarity in ionization energies and the relatively small increase in successive ionization energies lead to the formation of metal ions with the same charge for many of the transition metals.
This in turn results in extensive horizontal similarities in chemistry, which are most noticeable for the first-row transition metals and for the lanthanides and actinides. The relatively small increase in successive ionization energies causes most of the transition metals to exhibit multiple oxidation states separated by a single electron. Because of the slow but steady increase in ionization potentials across a row, high oxidation states become progressively less stable for the elements on the right side of the d block.
The occurrence of multiple oxidation states separated by a single electron causes many, if not most, compounds of the transition metals to be paramagnetic, with one to five unpaired electrons. This behavior is in sharp contrast to that of the p -block elements, where the occurrence of two oxidation states separated by two electrons is common, which makes virtually all compounds of the p -block elements diamagnetic.
Due to a small increase in successive ionization energies, most of the transition metals have multiple oxidation states separated by a single electron.
Most compounds of transition metals are paramagnetic, whereas virtually all compounds of the p -block elements are diamagnetic. Thus Sc is a rather active metal, whereas Cu is much less reactive. Exceptions to the overall trends are rather common, however, and in many cases, they are attributable to the stability associated with filled and half-filled subshells.
Consequently, all transition-metal cations possess d n valence electron configurations , as shown in Table The most common oxidation states of the first-row transition metals are shown in Table The second- and third-row transition metals behave similarly but with three important differences:.
The highest possible oxidation state, corresponding to the formal loss of all valence electrons, becomes increasingly less stable as we go from group 3 to group 8, and it is never observed in later groups. In the transition metals, the stability of higher oxidation states increases down a column. Binary transition-metal compounds, such as the oxides and sulfides, are usually written with idealized stoichiometries, such as FeO or FeS, but these compounds are usually cation deficient and almost never contain a cation:anion ratio.
Thus a substance such as ferrous oxide is actually a nonstoichiometric compound with a range of compositions. The acid—base character of transition-metal oxides depends strongly on the oxidation state of the metal and its ionic radius.
Conversely, oxides of metals in higher oxidation states are more covalent and tend to be acidic, often dissolving in strong base to form oxoanions. Identify these metals; predict the stoichiometry of the oxides; describe the general physical and chemical properties, type of bonding, and physical state of the oxides; and decide whether they are acidic or basic oxides.
Given: group 8 metals. Refer to the trends outlined in Figure Decide whether their oxides are covalent or ionic in character, and, based on this, predict the general physical and chemical properties of the oxides. Because the heavier transition metals tend to be stable in higher oxidation states, we expect Ru and Os to form the most stable tetroxides. Because oxides of metals in high oxidation states are generally covalent compounds, RuO 4 and OsO 4 should be volatile solids or liquids that consist of discrete MO 4 molecules, which the valence-shell electron-pair repulsion VSEPR model predicts to be tetrahedral.
Finally, because oxides of transition metals in high oxidation states are usually acidic, RuO 4 and OsO 4 should dissolve in strong aqueous base to form oxoanions. Predict the identity and stoichiometry of the stable group 9 bromide in which the metal has the lowest oxidation state and describe its chemical and physical properties.
Answer: Because the lightest element in the group is most likely to form stable compounds in lower oxidation states, the bromide will be CoBr 2. The transition metals are characterized by partially filled d subshells in the free elements and cations. In the second- and third-row transition metals, such irregularities can be difficult to predict, particularly for the third row, which has 4 f , 5 d , and 6 s orbitals that are very close in energy.
The increase in atomic radius is greater between the 3 d and 4 d metals than between the 4 d and 5 d metals because of the lanthanide contraction.
Ionization energies and electronegativities increase slowly across a row, as do densities and electrical and thermal conductivities, whereas enthalpies of hydration decrease. Anomalies can be explained by the increased stabilization of half-filled and filled subshells. Transition-metal cations are formed by the initial loss of ns electrons, and many metals can form cations in several oxidation states.
Higher oxidation states become progressively less stable across a row and more stable down a column. Oxides of small, highly charged metal ions tend to be acidic, whereas oxides of metals with a low charge-to-radius ratio are basic. The transition metals show significant horizontal similarities in chemistry in addition to their vertical similarities, whereas the same cannot be said of the s -block and p -block elements.
Explain why this is so. The energy of the d subshell does not change appreciably in a given period. What effect does this have on the ionization potentials of the transition metals? Standard reduction potentials vary across the first-row transition metals. What effect does this have on the chemical reactivity of the first-row transition metals? Which two elements in this period are more active than would be expected?
Many transition metals are paramagnetic have unpaired electrons. How does this affect electrical and thermal conductivities across the rows? What is the lanthanide contraction? What effect does it have on the radii of the transition metals of a given group?
What effect does it have on the chemistry of the elements in a group? Why are the atomic volumes of the transition elements low compared with the elements of groups 1 and 2? Ir has the highest density of any element in the periodic table Of the elements Ti, Ni, Cu, and Cd, which do you predict has the highest electrical conductivity?
The chemistry of As is most similar to the chemistry of which transition metal? Where in the periodic table do you find elements with chemistry similar to that of Ge? Explain your answers. The coinage metals group 11 have significant noble character. In fact, they are less reactive than the elements of group Explain why this is so, referring specifically to their reactivity with mineral acids, electronegativity, and ionization energies.
Why are the group 12 elements more reactive? Explain your reasoning. We turn now to a brief survey of the chemistry of the transition metals, beginning with group 3. As we shall see, the two heaviest members of each group usually exhibit substantial similarities in chemical behavior and are quite different from the lightest member. As shown in Table As expected based on periodic trends, these elements are highly electropositive metals and powerful reductants, with La and Ac being the most reactive.
In keeping with their highly electropositive character, the group 3 metals react with water to produce the metal hydroxide and hydrogen gas:. The group 3 metals react with nonmetals to form compounds that are primarily ionic in character. For example, reacting group 3 metals with the halogens produces the corresponding trihalides: MX 3. The trifluorides are insoluble in water because of their high lattice energies, but the other trihalides are very soluble in water and behave like typical ionic metal halides.
All group 3 elements react with air to form an oxide coating, and all burn in oxygen to form the so-called sesquioxides M 2 O 3 , which react with H2O or CO 2 to form the corresponding hydroxides or carbonates, respectively. They are isolated by initial conversion to the tetrachlorides, as shown for Ti:.
Only Ti has an extensive chemistry in lower oxidation states. In contrast to the elements of group 3, the group 4 elements have important applications. For example, friction with the air heats the skin of supersonic aircraft operating above Mach 2.
The corrosion resistance of titanium is increasingly exploited in architectural applications, as shown in the chapter-opening photo. Metallic zirconium is used in UO 2 -containing fuel rods in nuclear reactors, while hafnium is used in the control rods that modulate the output of high-power nuclear reactors, such as those in nuclear submarines.
Consistent with the periodic trends shown in Figure Unexpectedly, however, the atomic radius of Hf is slightly smaller than that of Zr due to the lanthanide contraction. Reaction of the group 4 metals with excess halogen forms the corresponding tetrahalides MX 4 , although titanium, the lightest element in the group, also forms dihalides and trihalides X is not F.
The covalent character of the titanium halides increases as the oxidation state of the metal increases because of increasing polarization of the anions by the cation as its charge-to-radius ratio increases. Thus TiCl 2 is an ionic salt, whereas TiCl 4 is a volatile liquid that contains tetrahedral molecules. All three metals react with excess oxygen or the heavier chalcogens Y to form the corresponding dioxides MO 2 and dichalcogenides MY 2.
Industrially, TiO 2 , which is used as a white pigment in paints, is prepared by reacting TiCl 4 with oxygen at high temperatures:. The group 4 dichalcogenides have unusual layered structures with no M—Y bonds holding adjacent sheets together, which makes them similar in some ways to graphite Figure The group 4 metals also react with hydrogen, nitrogen, carbon, and boron to form hydrides such as TiH 2 , nitrides such as TiN , carbides such as TiC , and borides such as TiB 2 , all of which are hard, high-melting solids.
Many of these binary compounds are nonstoichiometric and exhibit metallic conductivity. Each titanium atom is surrounded by an octahedral arrangement of six sulfur atoms that are shared to form extended layers of atoms. Because the layers are held together by only van der Waals forces between adjacent sulfur atoms, rather than covalent bonds, the layers slide past one another relatively easily when a mechanical stress is applied.
Because of the lanthanide contraction, the chemistry of Nb and Ta is so similar that these elements are usually found in the same ores. Three-fourths of the vanadium produced annually is used in the production of steel alloys for springs and high-speed cutting tools.
Adding a small amount of vanadium to steel results in the formation of small grains of V 4 C 3 , which greatly increase the strength and resilience of the metal, especially at high temperatures. The other major use of vanadium is as V 2 O 5 , an important catalyst for the industrial conversion of SO 2 to SO 3 in the contact process for the production of sulfuric acid. In contrast, Nb and Ta have only limited applications, and they are therefore produced in relatively small amounts. Although niobium is used as an additive in certain stainless steels, its primary application is in superconducting wires such as Nb 3 Zr and Nb 3 Ge, which are used in superconducting magnets for the magnetic resonance imaging of soft tissues.
Because tantalum is highly resistant to corrosion, it is used as a liner for chemical reactors, in missile parts, and as a biologically compatible material in screws and pins for repairing fractured bones.
As indicated in Table Oxides of these metals in lower oxidation states tend to be nonstoichiometric. Although group 5 metals react with the heavier chalcogens to form a complex set of binary chalcogenides, the most important are the dichalcogenides MY 2 , whose layered structures are similar to those of the group 4 dichalcogenides.
The elements of group 5 also form binary nitrides, carbides, borides, and hydrides, whose stoichiometries and properties are similar to those of the corresponding group 4 compounds. As an illustration of the trend toward increasing polarizability as we go from left to right across the d block, in group 6 we first encounter a metal Mo that occurs naturally as a sulfide ore rather than as an oxide.
Molybdenite MoS 2 is a soft black mineral that can be used for writing, like PbS and graphite. Because of this similarity, people long assumed that these substances were all the same. In addition, molybdenum is the only second- or third-row transition element that is essential for humans. Pure chromium can be obtained by dissolving Cr 2 O 3 in sulfuric acid followed by electrolytic reduction; a similar process is used for electroplating metal objects to give them a bright, shiny, protective surface layer.
Pure tungsten is obtained by first converting tungsten ores to WO 3 , which is then reduced with hydrogen to give the metal. Consistent with periodic trends, the group 6 metals are slightly less electropositive than those of the three preceding groups, and the two heaviest metals are essentially the same size because of the lanthanide contraction Table The metals become increasing polarizable across the d block. The chemistry of the lightest element Cr is dominated by lower oxidation states.
As observed in previous groups, the group 6 halides become more covalent as the oxidation state of the metal increases: their volatility increases, and their melting points decrease.
Recall that as the electronegativity of the halogens decreases from F to I, they are less able to stabilize high oxidation states; consequently, the maximum oxidation state of the corresponding metal halides decreases.
Chromium will form CrO 3 , which is a highly toxic compound that can react explosively with organic materials. Consistent with periodic trends, the sesquioxide of the lightest element in the group Cr 2 O 3 is amphoteric.
An isopolymolybdate cluster. Reacting molybdenum or tungsten with heavier chalcogens gives binary chalcogenide phases, most of which are nonstoichiometric and electrically conducting.
Consequently, both MoS 2 and WS 2 are used as lubricants in a variety of applications, including automobile engines. As in groups 4 and 5, the elements of group 6 form binary nitrides, carbides, and borides whose stoichiometries and properties are similar to those of the preceding groups.
Tungsten carbide WC , one of the hardest compounds known, is used to make the tips of drill bits. Continuing across the periodic table, we encounter the group 7 elements Table One group 7 metal Mn is usually combined with iron in an alloy called ferromanganese , which has been used since to improve the mechanical properties of steel by scavenging sulfur and oxygen impurities to form MnS and MnO. One isotope, 99 m Tc m for metastable , has become an important biomedical tool for imaging internal organs.
For more information on biomedical imaging, see Chapter 20 "Nuclear Chemistry" , Section Because of its scarcity, Re is one of the most expensive elements, and its applications are limited.
Once again, the lightest element exhibits multiple oxidation states. Because the electronegativity of Mn is anomalously low, elemental manganese is unusually reactive.
In contrast, the chemistry of Tc is similar to that of Re because of their similar size and electronegativity, again a result of the lanthanide contraction. It is difficult to generalize about other oxidation states for Tc and Re because their stability depends dramatically on the nature of the compound. Consistent with higher oxidation states being more stable for the heavier transition metals, reacting Mn with F 2 gives only MnF 3 , a high-melting, red-purple solid, whereas Re reacts with F 2 to give ReF 7 , a volatile, low-melting, yellow solid.
Again, reaction with the less oxidizing, heavier halogens produces halides in lower oxidation states.
In contrast, Tc and Re form high-valent oxides, the so-called heptoxides M 2 O 7 , consistent with the increased stability of higher oxidation states for the second and third rows of transition metals. Under forced conditions, manganese will form Mn 2 O 7 , an unstable, explosive, green liquid. Both Tc and Re form disulfides and diselenides with layered structures analogous to that of MoS 2 , as well as more complex heptasulfides M 2 S 7. As is typical of the transition metals, the group 7 metals form binary nitrides, carbides, and borides that are generally stable at high temperatures and exhibit metallic properties.
The chemistry of the group 7 metals Mn, Tc, and Re is dominated by lower oxidation states. In many older versions of the periodic table, groups 8, 9, and 10 were combined in a single group group VIII because the elements of these three groups exhibit many horizontal similarities in their chemistry, in addition to the similarities within each column. In part, these horizontal similarities are due to the fact that the ionization potentials of the elements, which increase slowly but steadily across the d block, have now become so large that the oxidation state corresponding to the formal loss of all valence electrons is encountered only rarely group 8 or not at all groups 9 and The heavier elements of these three groups are called precious metals because they are rather rare in nature and mostly chemically inert.
Ruthenium and osmium, on the other hand, are extremely rare elements, with terrestrial abundances of only about 0. The advanced techniques needed to work iron were first developed by the Hittite civilization in Asia Minor sometime before BC, and they remained a closely guarded secret that gave the Hittites military supremacy for almost a millennium.
With the collapse of the Hittite civilization around BC, the technology became widely distributed, however, leading to the Iron Age. Cobalt is one of the least abundant of the first-row transition metals. The heavier elements of group 9 are also rare, with terrestrial abundances of less than 1 ppb; they are generally found in combination with the heavier elements of groups 8 and 10 in Ni—Cu—S ores. Nickel silicates are easily processed; consequently, nickel has been known and used since antiquity.
In contrast to nickel, palladium and platinum are rare their terrestrial abundance is about 10—15 ppb , but they are at least an order of magnitude more abundant than the heavier elements of groups 8 and 9. Over years ago, the Bactrian civilization in Western Asia used a alloy of copper and nickel for its coins. A modern US nickel has the same composition, but a modern Canadian nickel is nickel-plated steel and contains only 2. Some properties of the elements in groups 8—10 are summarized in Table As in earlier groups, similarities in size and electronegativity between the two heaviest members of each group result in similarities in chemistry.
We are now at the point in the d block where there is no longer a clear correlation between the valence electron configuration and the preferred oxidation state. We stated that higher oxidation states become less stable as we go across the d -block elements and more stable as we go down a group. The hexafluorides of Rh and Ir are extraordinarily powerful oxidants, and Pt is the only element in group 10 that forms a hexafluoride.
Similar trends are observed among the oxides. As expected for compounds of metals in such high oxidation states, the latter are potent oxidants.
The tendency of the metals to form the higher oxides decreases rapidly as we go farther across the d block. Higher oxidation states become less stable across the d -block, but more stable down a group.
Reactivity with the heavier chalcogens is rather complex. This combination of highly charged cations and easily polarized anions results in substances that are not simple ionic compounds and have significant covalent character. The groups 8—10 metals form a range of binary nitrides, carbides, and borides. By far the most important of these is cementite Fe 3 C , which is used to strengthen steel.
At high temperatures, Fe 3 C is soluble in iron, but slow cooling causes the phases to separate and form particles of cementite, which gives a metal that retains much of its strength but is significantly less brittle than pure iron.
Palladium is unusual in that it forms a binary hydride with the approximate composition PdH 0. Because the H atoms in the metal lattice are highly mobile, thin sheets of Pd are highly permeable to H 2 but essentially impermeable to all other gases, including He. Consequently, diffusion of H 2 through Pd is an effective method for separating hydrogen from other gases. The coinage metals—copper, silver, and gold—occur naturally like the gold nugget shown here ; consequently, these were probably the first metals used by ancient humans.
For example, decorative gold artifacts dating from the late Stone Age are known, and some gold Egyptian coins are more than yr old. Copper is almost as ancient, with objects dating to about BC. Bronze, an alloy of copper and tin that is harder than either of its constituent metals, was used before BC, giving rise to the Bronze Age.
Deposits of silver are much less common than deposits of gold or copper, yet by BC, methods had been developed for recovering silver from its ores, which allowed silver coins to be widely used in ancient times. Deposits of gold and copper are widespread and numerous, and for many centuries it was relatively easy to obtain large amounts of the pure elements. For example, a single gold nugget discovered in Australia in weighed more than lb.
Copper is used primarily to manufacture electric wires, but large quantities are also used to produce bronze, brass, and alloys for coins. Much of the silver made today is obtained as a by-product of the manufacture of other metals, especially Cu, Pb, and Zn. For more information on button batteries, see Chapter 19 "Electrochemistry" , Section Gold is typically found either as tiny particles of the pure metal or as gold telluride AuTe 2. It is used as a currency reserve, in jewelry, in the electronics industry for corrosion-free contacts, and, in very thin layers, as a reflective window coating that minimizes heat transfer.
Each gigantic truck in the foreground and barely visible in the lower right center can hold metric tn , kg of copper ore. Some properties of the coinage metals are listed in Table The coinage metals have the highest electrical and thermal conductivities of all the metals, and they are also the most ductile and malleable. All three elements have significant electron affinities due to the half-filled ns orbital in the neutral atoms. All group 11 elements are relatively unreactive, and their reactivity decreases from Cu to Au.
Hence they are noble metals that are particularly well suited for use in coins and jewelry. The remaining two electrons occupy the 2 p subshell. We now have a choice of filling one of the 2 p orbitals and pairing the electrons or of leaving the electrons unpaired in two different, but degenerate, p orbitals. Thus, the two electrons in the carbon 2 p orbitals have identical n , l , and m s quantum numbers and differ in their m l quantum number in accord with the Pauli exclusion principle.
The electron configuration and orbital diagram for carbon are:. These three electrons have unpaired spins. Oxygen atomic number 8 has a pair of electrons in any one of the 2 p orbitals the electrons have opposite spins and a single electron in each of the other two. Fluorine atomic number 9 has only one 2 p orbital containing an unpaired electron. The electron configurations and orbital diagrams of these four elements are:. The alkali metal sodium atomic number 11 has one more electron than the neon atom.
This electron must go into the lowest-energy subshell available, the 3 s orbital, giving a 1 s 2 2 s 2 2 p 6 3 s 1 configuration. The electrons occupying the outermost shell orbital s highest value of n are called valence electrons , and those occupying the inner shell orbitals are called core electrons Figure. Since the core electron shells correspond to noble gas electron configurations, we can abbreviate electron configurations by writing the noble gas that matches the core electron configuration, along with the valence electrons in a condensed format.
For our sodium example, the symbol [Ne] represents core electrons, 1 s 2 2 s 2 2 p 6 and our abbreviated or condensed configuration is [Ne]3 s 1. Similarly, the abbreviated configuration of lithium can be represented as [He]2 s 1 , where [He] represents the configuration of the helium atom, which is identical to that of the filled inner shell of lithium.
Writing the configurations in this way emphasizes the similarity of the configurations of lithium and sodium. Both atoms, which are in the alkali metal family, have only one electron in a valence s subshell outside a filled set of inner shells. The alkaline earth metal magnesium atomic number 12 , with its 12 electrons in a [Ne]3 s 2 configuration, is analogous to its family member beryllium, [He]2 s 2.
Both atoms have a filled s subshell outside their filled inner shells. Aluminum atomic number 13 , with 13 electrons and the electron configuration [Ne]3 s 2 3 p 1 , is analogous to its family member boron, [He]2 s 2 2 p 1. Figure shows the lowest energy, or ground-state, electron configuration for these elements as well as that for atoms of each of the known elements. When we come to the next element in the periodic table, the alkali metal potassium atomic number 19 , we might expect that we would begin to add electrons to the 3 d subshell.
However, all available chemical and physical evidence indicates that potassium is like lithium and sodium, and that the next electron is not added to the 3 d level but is, instead, added to the 4 s level Figure.
As discussed previously, the 3 d orbital with no radial nodes is higher in energy because it is less penetrating and more shielded from the nucleus than the 4 s , which has three radial nodes. Thus, potassium has an electron configuration of [Ar]4 s 1. Hence, potassium corresponds to Li and Na in its valence shell configuration.
The next electron is added to complete the 4 s subshell and calcium has an electron configuration of [Ar]4 s 2. This gives calcium an outer-shell electron configuration corresponding to that of beryllium and magnesium. Beginning with the transition metal scandium atomic number 21 , additional electrons are added successively to the 3 d subshell. The 4 p subshell fills next. Note that for three series of elements, scandium Sc through copper Cu , yttrium Y through silver Ag , and lutetium Lu through gold Au , a total of 10 d electrons are successively added to the n — 1 shell next to the n shell to bring that n — 1 shell from 8 to 18 electrons.
Quantum Numbers and Electron Configurations What is the electron configuration and orbital diagram for a phosphorus atom? What are the four quantum numbers for the last electron added? Solution The atomic number of phosphorus is Thus, a phosphorus atom contains 15 electrons.
The 15 electrons of the phosphorus atom will fill up to the 3 p orbital, which will contain three electrons:. The last electron added is a 3 p electron. The three p orbitals are degenerate, so any of these m l values is correct. For unpaired electrons, convention assigns the value of for the spin quantum number; thus,.
Check Your Learning Identify the atoms from the electron configurations given:. The periodic table can be a powerful tool in predicting the electron configuration of an element.
However, we do find exceptions to the order of filling of orbitals that are shown in Figure or Figure. For instance, the electron configurations shown in Figure of the transition metals chromium Cr; atomic number 24 and copper Cu; atomic number 29 , among others, are not those we would expect.
In general, such exceptions involve subshells with very similar energy, and small effects can lead to changes in the order of filling. In the case of Cr and Cu, we find that half-filled and completely filled subshells apparently represent conditions of preferred stability. This stability is such that an electron shifts from the 4 s into the 3 d orbital to gain the extra stability of a half-filled 3 d subshell in Cr or a filled 3 d subshell in Cu.
Other exceptions also occur. For example, niobium Nb, atomic number 41 is predicted to have the electron configuration [Kr]5 s 2 4 d 3. Experimentally, we observe that its ground-state electron configuration is actually [Kr]5 s 1 4 d 4.
We can rationalize this observation by saying that the electron—electron repulsions experienced by pairing the electrons in the 5 s orbital are larger than the gap in energy between the 5 s and 4 d orbitals. What is the valence electron configuration of zirconium Zr? What is the full electron configuration for Zr? What is the electron configuration for germanium?
What is the chemical symbol for germanium? What is the electron configuration for CO? What is the electron configuration of Cu? Why does CU have a 2 charge? Why is the electron configuration of Cu different? How many 3d electrons are in CU? What is Hunds?
What violates Hunds? What is Aufbau rule in chemistry? Why does 3d orbital fill before 4s? Why do d orbitals start at 3?
0コメント