omparison of the physical properties and chemical properties between alkaline earth and alkali metals Just like their names, they don’t differ completely. The main difference is the electron configuration, which is ns2 for alkaline earth metals and ns1 for alkali metals. For the alkaline earth metals, there are two electrons that are available to for a metallic bond, and the nucleus contains an additional positive charge. Also, the elements of group 2A (alkaline earth) have much higher melting points and boiling points compared to those of group 1A (alkali metals). The alkali also have a softer and more lighweight figure whereas the alkaline earth metals are much harder and denser.
The second valence electron is very important when it comes to comparing chemical properties of the alkaline earth and the alkali metals. The second valence electron is in the same “sublevel” as the first valence electron. Therefore, the Zeff is much greater. This means that the elements of the group 2A contain a smaller atomic radius and much higher ionization energy than the group 1A. Even though the group 2A contains much higher ionization energy, they still form an ionic compound with 2+ cations. Beryllium, however, behaves differently. This is due to the fact that in order to remove two electrons from this particular atom, it requires significantly more energy. It never forms Be2+ and its bonds are polar covalent.
[edit] Beryllium As mentioned earlier, Be is “special”; it behaves differently. If the Be2+ ion did exist, it would polarize electron clouds that are near it very strongly and would cause extensive orbital overlap, since Be has a high charge density. All compounds that include Be have a covalent bond. Even the compound BeF2 , which is the most ionic Be compound, has a low melting point and a low electrical conductivity when melted.
[edit] Important reactions and compounds Reactions:
Note: E = elements that act as reducing agents
1. The metals reduce halogens to form ionic halides: E(s) + X2 → EX2 (s) where X = F, Cl, Br or I
2. The metals reduce O2 to form the oxides: 2E(s) + O2 → 2EO(s)
3. The larger metals react with water to produce hydrogen gas: E(s) + 2H2O(l) → E2+(aq) + 2OH-aq + H2 (g) where E = Ca, Sr or Ba
Compounds
1. Alkylmagnesium halides (RMgX where R = hydrocarbon group and X = halogen). They are used to synthetise organic compounds.
Here’s an example: 3RMgCl + SnCl4 → 3MgCl2 + R3SnCl
2. Magnesium oxide (MgO). It is used as a material to refract furnace brick and wire insulation (melting point of 2852°C).
3. Calcium carbonate (CaCO3). It’s mainly used in the construction industry. It’s responsible for making limestone, marble, chalk, and coral.
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The second valence electron is very important when it comes to comparing chemical properties of the alkaline earth and the alkali metals. The second valence electron is in the same “sublevel” as the first valence electron. Therefore, the Zeff is much greater. This means that the elements of the group 2A contain a smaller atomic radius and much higher ionization energy than the group 1A. Even though the group 2A contains much higher ionization energy, they still form an ionic compound with 2+ cations. Beryllium, however, behaves differently. This is due to the fact that in order to remove two electrons from this particular atom, it requires significantly more energy. It never forms Be2+ and its bonds are polar covalent.
[edit] Beryllium As mentioned earlier, Be is “special”; it behaves differently. If the Be2+ ion did exist, it would polarize electron clouds that are near it very strongly and would cause extensive orbital overlap, since Be has a high charge density. All compounds that include Be have a covalent bond. Even the compound BeF2 , which is the most ionic Be compound, has a low melting point and a low electrical conductivity when melted.
[edit] Important reactions and compounds Reactions:
Note: E = elements that act as reducing agents
1. The metals reduce halogens to form ionic halides: E(s) + X2 → EX2 (s) where X = F, Cl, Br or I
2. The metals reduce O2 to form the oxides: 2E(s) + O2 → 2EO(s)
3. The larger metals react with water to produce hydrogen gas: E(s) + 2H2O(l) → E2+(aq) + 2OH-aq + H2 (g) where E = Ca, Sr or Ba
Compounds
1. Alkylmagnesium halides (RMgX where R = hydrocarbon group and X = halogen). They are used to synthetise organic compounds.
Here’s an example: 3RMgCl + SnCl4 → 3MgCl2 + R3SnCl
2. Magnesium oxide (MgO). It is used as a material to refract furnace brick and wire insulation (melting point of 2852°C).
3. Calcium carbonate (CaCO3). It’s mainly used in the construction industry. It’s responsible for making limestone, marble, chalk, and coral.
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The term transition metal (sometimes also called a transition element) has two possible meanings:
Contents [hide] // [edit] Classification In the d-block the atoms of the elements have between 1 and 10 d electrons.
Group 3 4 5 6 7 8 9 10 11 12 Period 4 Sc 21 Ti 22 V 23 Cr 24 Mn 25 Fe 26 Co 27 Ni 28 Cu 29 Zn 30 Period 5 Y 39 Zr 40 Nb 41 Mo 42 Tc 43 Ru 44 Rh 45 Pd 46 Ag 47 Cd 48 Period 6 La 57 Hf 72 Ta 73 W 74 Re 75 Os 76 Ir 77 Pt 78 Au 79 Hg 80 Period 7 Ac 89 Rf 104 Db 105 Sg 106 Bh 107 Hs 108 Mt 109 Ds 110 Rg 111 Uub 112 Although atoms of scandium and yttrium have a single d electron in the outermost shell, these elements are not usually considered transition metals as all their compounds contain the ions Sc3+ and Y3+ in which there are no d electrons. Lanthanum is usually considered a lanthanide element and actinium an actinide element.[2]
The electronic structure of transition metal atoms can be written, with a few minor exceptions, as [ ]ns2(n-1)dm, as the inner d orbital has more energy than the valence-shell s orbital. In divalent and trivalent ions of the transition metals the situation is reversed so that the s electrons have higher energy. Consequently an ion such as Fe2+ has no s electrons, it has the electronic configuration [Ar]3d6 as compared with the configuration of the atom, [Ar]4s23d6.
Zinc, cadmium, and mercury are not transition metals.[3] as they have the electronic configuration [ ]d10s2, with no incomplete d shell.[4] In the oxidation state +2 the ions have the electronic configuration [ ] d10. While compounds in the +1 oxidation state, such as Hg22+, are known there are no unpaired electrons because of the formation of a covalent bond between the two atoms of the dimer. Zn, Cd and Hg may be classed as post-transition metals. However, it is often convenient to include these elements in a discussion of the transition elements. For example, when discussing the Crystal field stabilization energy of first-row transition elements, it is convenemt to include the non-transition elements calcium and zinc, as both Ca2+ and Zn2+ have a value of zero against which the value for other transtion metal ions may be compared. Another example occurs in the Irving-Williams series of stability constants of complexes.
There are a number of properties shared by the transition elements that are not found in other elements, which results from the partially filled d shell. These include
Main-group elements in groups 13 - 17 also exhibit multiple oxidation states. The "common" oxidation states of those elements differ by two. For example, compounds of gallium in oxidation states +1 and +3 exist in which there is a single gallium atom. No such compound of Ga(II) is known: any such compound would have an unpaired electron and would behave as a free radical and be destroyed rapidly. However, under the right conditions dimeric compounds such as [Ga2Cl6]2- can be made in which a Ga-Ga bond is formed from the unpaired electron on each Ga atom. [8] Thus the main difference, regarding oxidation states, between transition elements and other elements is that oxidation states are known in which there is a single atom of the element and one or more unpaired electrons.
The maximum oxidation state in the first row transition metals is equal to the number of valence electron from titanium (+4) up to manganese (+7), but decreases in the later elements. In the second and third rows the maximum occurs with ruthenium and osmium (+8). In compounds such as [MnO4]- and OsO4 the elements achieve a stable octet by forming four covalent bonds.
The lowest oxidation states are exhibited in such compounds as Cr(CO)6 (oxidation state zero) and [Fe(CO)4]2- (oxidation state -2) in which the 18-electron rule is obeyed. These complexes are also covalent.
Ionic compounds are mostly formed with oxidation states +2 and +3. In aqueous solution the ions are hydrated by (usually) six water molecules arranged octahedrally.
[edit] Magnetism Transition metal compounds are paramagnetic when they have one or more unpaired d electrons.[9] In octahedral complexes with between four and seven d electrons both high spin and low spin states are possible. Tetrahedral transition metal complexes such as [FeCl4]2- are high spin because the crystal field splitting is small so that the energy to be gained by virtue of the electrons being in lower energy orbitals is always less that the energy needed to pair up the spins. Some compounds are diamagnetic. These include octahedral, low-spin, d6 and square-planar d8 complexes. In these cases, crystal field splitting is such that all the electrons are paired up.
Ferromagnetism occurs when individual atoms are paramagnetic and the spin vectors are aligned parallel to each other in a crystalline material. Metallic iron and the alloy alnico are examples of ferromagnetic materials involving transition metals. Antiferromagnetism is another example of a magnetic property arising from a particular alignment of individual spins in the solid state.
- In the past it referred to any element in the d-block of the periodic table, which includes groups 3 to 12 on the periodic table. All elements in the d-block are metals.
- The modern, IUPAC definition[1] states that a transition metal is "an element whose atom has an incomplete d sub-shell, or which can give rise to cations with an incomplete d sub-shell." Group 12 elements are not transition metals in this definition.
Contents [hide] // [edit] Classification In the d-block the atoms of the elements have between 1 and 10 d electrons.
Group 3 4 5 6 7 8 9 10 11 12 Period 4 Sc 21 Ti 22 V 23 Cr 24 Mn 25 Fe 26 Co 27 Ni 28 Cu 29 Zn 30 Period 5 Y 39 Zr 40 Nb 41 Mo 42 Tc 43 Ru 44 Rh 45 Pd 46 Ag 47 Cd 48 Period 6 La 57 Hf 72 Ta 73 W 74 Re 75 Os 76 Ir 77 Pt 78 Au 79 Hg 80 Period 7 Ac 89 Rf 104 Db 105 Sg 106 Bh 107 Hs 108 Mt 109 Ds 110 Rg 111 Uub 112 Although atoms of scandium and yttrium have a single d electron in the outermost shell, these elements are not usually considered transition metals as all their compounds contain the ions Sc3+ and Y3+ in which there are no d electrons. Lanthanum is usually considered a lanthanide element and actinium an actinide element.[2]
The electronic structure of transition metal atoms can be written, with a few minor exceptions, as [ ]ns2(n-1)dm, as the inner d orbital has more energy than the valence-shell s orbital. In divalent and trivalent ions of the transition metals the situation is reversed so that the s electrons have higher energy. Consequently an ion such as Fe2+ has no s electrons, it has the electronic configuration [Ar]3d6 as compared with the configuration of the atom, [Ar]4s23d6.
Zinc, cadmium, and mercury are not transition metals.[3] as they have the electronic configuration [ ]d10s2, with no incomplete d shell.[4] In the oxidation state +2 the ions have the electronic configuration [ ] d10. While compounds in the +1 oxidation state, such as Hg22+, are known there are no unpaired electrons because of the formation of a covalent bond between the two atoms of the dimer. Zn, Cd and Hg may be classed as post-transition metals. However, it is often convenient to include these elements in a discussion of the transition elements. For example, when discussing the Crystal field stabilization energy of first-row transition elements, it is convenemt to include the non-transition elements calcium and zinc, as both Ca2+ and Zn2+ have a value of zero against which the value for other transtion metal ions may be compared. Another example occurs in the Irving-Williams series of stability constants of complexes.
There are a number of properties shared by the transition elements that are not found in other elements, which results from the partially filled d shell. These include
- the formation of compounds whose colour is due to d - d electronic transitions
- the formation of compounds in many oxidation states, due to the relatively low reactivity of unpaired d electrons.[5]
- the formation of many paramagnetic compounds due to the presence of unpaired d electrons. A few compounds of main group elements are paramagnetic (e.g. nitric oxide, oxygen)
- charge transfer transitions. An electron may jump from a predominantly ligand orbital to a predominantly metal orbital, giving rise to a ligand-to-metal charge-transfer (LMCT) transition. These can most easily occur when the metal is in a high oxidation state. The colour of chromate, dichromate and permanganate ions is due to LMCT transitions. A metal-to ligand charge transfer (MLCT) transition will be most likely when the metal is in a low oxidation state and the ligand is easily oxidised. Mercuric iodide, HgI2, is red because of a MLCT transition. As this example shows, charge transfer transitions are not restricted to transition metals.[6]
- d-d transitions. An electron jumps from one d-orbital to another. In complexes of the transition metals the d orbitals do not all have the same energy. The pattern of splitting of the d orbitals can be calculated using crystal field theory. The extent of the splitting depends on the particular metal, its oxidation state and the nature of the ligands. The actual energy levels are shown on Tanabe-Sugano diagrams.
Main-group elements in groups 13 - 17 also exhibit multiple oxidation states. The "common" oxidation states of those elements differ by two. For example, compounds of gallium in oxidation states +1 and +3 exist in which there is a single gallium atom. No such compound of Ga(II) is known: any such compound would have an unpaired electron and would behave as a free radical and be destroyed rapidly. However, under the right conditions dimeric compounds such as [Ga2Cl6]2- can be made in which a Ga-Ga bond is formed from the unpaired electron on each Ga atom. [8] Thus the main difference, regarding oxidation states, between transition elements and other elements is that oxidation states are known in which there is a single atom of the element and one or more unpaired electrons.
The maximum oxidation state in the first row transition metals is equal to the number of valence electron from titanium (+4) up to manganese (+7), but decreases in the later elements. In the second and third rows the maximum occurs with ruthenium and osmium (+8). In compounds such as [MnO4]- and OsO4 the elements achieve a stable octet by forming four covalent bonds.
The lowest oxidation states are exhibited in such compounds as Cr(CO)6 (oxidation state zero) and [Fe(CO)4]2- (oxidation state -2) in which the 18-electron rule is obeyed. These complexes are also covalent.
Ionic compounds are mostly formed with oxidation states +2 and +3. In aqueous solution the ions are hydrated by (usually) six water molecules arranged octahedrally.
[edit] Magnetism Transition metal compounds are paramagnetic when they have one or more unpaired d electrons.[9] In octahedral complexes with between four and seven d electrons both high spin and low spin states are possible. Tetrahedral transition metal complexes such as [FeCl4]2- are high spin because the crystal field splitting is small so that the energy to be gained by virtue of the electrons being in lower energy orbitals is always less that the energy needed to pair up the spins. Some compounds are diamagnetic. These include octahedral, low-spin, d6 and square-planar d8 complexes. In these cases, crystal field splitting is such that all the electrons are paired up.
Ferromagnetism occurs when individual atoms are paramagnetic and the spin vectors are aligned parallel to each other in a crystalline material. Metallic iron and the alloy alnico are examples of ferromagnetic materials involving transition metals. Antiferromagnetism is another example of a magnetic property arising from a particular alignment of individual spins in the solid state.
