1.3.2 (a) Redox Reactions of Group 2 Metals. It cannot be said that by moving down the group these metals burn more vigorously. As I said earlier, they are powerful reducing age… The LibreTexts libraries are Powered by MindTouch® and are supported by the Department of Education Open Textbook Pilot Project, the UC Davis Office of the Provost, the UC Davis Library, the California State University Affordable Learning Solutions Program, and Merlot. The lattice energy is greatest if the ions are small and highly charged - the ions will be close together with very strong attractions. You might possibly be able to imagine a trace of very pale greenish colour surrounding the white flame in the third video, but to my eye, they all count as a white flame. "X" in the equation can represent any of the metals in the Group. There is an increase in the tendency to form the peroxide as you go down the Group. The size of the lattice energy depends on the attractions between the ions. (h) trend in general reactivity of Group 1 and Group 2 metals; Northern Ireland. If this is the first set of questions you have done, please read the introductory page before you start. It explains why it is difficult to observe many tidy patterns. M = Mg, Ca, Sr,Ba --> I will be using 'M' as the general symbol for a Group II element in this topic. 2:09 know the approximate percentages by volume of the four most abundant gases in dry air The strontium equation would look just the same. The equations for the reactions: Only in lithium's case is enough energy released to compensate for the energy needed to ionize the metal and the nitrogen - and so produce an exothermic reaction overall. For example, the familiar white ash you get when you burn magnesium ribbon in air is a mixture of magnesium oxide and magnesium nitride. Group 2 reactions Reactivity of group 2 metals increases down the group Mg will also react slowly with oxygen without a flame. Mg + 2 H2O Mg(OH) 2 + H2 This is a much slower reaction than the reaction with steam and there is no flame. This leads to lower activation energies, and therefore faster reactions. 2.11 Group II elements and their compounds. The speed is controlled by factors like the presence of surface coatings on the metal and the size of the activation energy. When something like magnesium nitride forms, you have to supply all the energy needed to form the magnesium ions as well as breaking the nitrogen-nitrogen bonds and then forming N3- ions. As group 2 in the periodic table comprises of metals, the reactivity of group 2 elements towards chlorine increases when working our way down the group 2 metals. 2Li(s) + Cl 2 (g) → 2LiCl(s) A similar reaction takes place with the other elements of group 7. In each case, you will get a mixture of the metal oxide and the metal nitride. Beryllium is reluctant to burn unless it is in the form of dust or powder. Ca(s) + H2O(l) → Ca(OH)2(aq) + H2(g) REACTIONS OF THE GROUP 2 ELEMENTS WITH AIR OR OXYGEN. Unless otherwise noted, LibreTexts content is licensed by CC BY-NC-SA 3.0. 2Mg + O2 2MgO This needs to be cleaned off by emery paper before doing reactions with Mg ribbon. For example, Magnesium reacts with Oxygen to form Magnesium Oxide the formula for which is: 2Mg (s) + O 2 (g) 2MgO (s) This is a redox reaction. My best guess would be the same sort of silvery sparkles that magnesium or aluminium powder burn with if they are scattered into a flame - but I don't know that for sure. Beryllium has a very strong (but very thin) layer of beryllium oxide on its surface, and this prevents any new oxygen getting at the underlying beryllium to react with it. The size of the lattice energy depends on the attractions between the ions. On the whole, the metals burn in oxygen to form a simple metal oxide. Now imagine bringing a small 2+ ion close to the peroxide ion. Oxygen: All of the elements in group 2 react vigorously with Oxygen, the product of which is an ionic oxide. Reaction of Group-2 Metals with Cl 2 : All Gr-2 metals except Be react with chlorine to give ionic chlorides whereas Be reacts with chlorine to give covalent chloride . REACTIONS OF THE GROUP 2 ELEMENTS WITH COMMON ACIDS This page looks at the reactions of the Group 2 elements - beryllium, magnesium, calcium, strontium and barium - with common acids. Reaction with oxygen. This energy has to be recovered from somewhere to give an overall exothermic reaction - if the energy can't be recovered, the overall change will be endothermic and will not happen. In each case, you will get a mixture of the metal oxide and the metal nitride. The general trend in acidity in oxides of the Period 3 elements as we go across the period from left (Group 1) to right (Group 17): basic oxides (Group 1, 2) → amphoteric oxide (Al 2 O 3) → acidic oxides (oxyacids) The same trend can be seen in each period of the Periodic table and we have: Bases react with acids such is HCl: Barium peroxide can form because the barium ion is so large that it doesn't have such a devastating effect on the peroxide ions as the metals further up the Group. This energy has to be recovered from somewhere to give an overall exothermic reaction - if the energy can't be recovered, the overall change will be endothermic and won't happen. Discusses trends in atomic radius, ionisation energy, electronegativity and melting point of the Group 2 elements. Anything else that I could find in a short clip from YouTube involved a flame test for a barium compound, irrespective of how it was described in the video. . Beryllium, magnesium and calcium don't form peroxides when heated in oxygen, but strontium and barium do. The Facts. reactivity trend down group 2 with water. Why do some metals form peroxides on heating in oxygen? In this case, though, the effect of the fall in the activation energy is masked by other factors - for example, the presence of existing oxide layers on the metals, and the impossibility of controlling precisely how much heat you are supplying to the metal in order to get it to start burning. The covalent bond between the two oxygen atoms is relatively weak. Describe the trend in the reactivity of group 2 elements with chlorine as you descend down the group. The activation energy will fall because the ionisation energies of the metals fall. (b) Relative Reactivities of the Group 2 elements Mg → Ba shown by their redox reactions with: (i) Oxygen (ii) Water (iii) Dilute acids {Reactions with acids will be limited to those producing a salt and Hydrogen.} But how reactive a metal seems to be depends on how fast the reaction happens - not the overall amount of heat evolved. Mg ribbon will often have a thin layer of magnesium oxide on it formed by reaction with oxygen. The general formula for this reaction is MO (where M is the group 2 element). Reactions with oxygen. Reactions. It would be tempting to say that the reactions get more vigorous as you go down the Group, but it isn't true. Strontium forms this if it is heated in oxygen under high pressures, but barium forms barium peroxide just on normal heating in oxygen. $Ba_{(s)} + O_{2(s)} \rightarrow BaO_{2(s)}$. As a whole, metals when burns with the oxygen form a simple metal oxide. The group 2 metals will burn in oxygen. Why do these metals form nitrides on heating in air? Beryllium has a very strong (but very thin) layer of beryllium oxide on its surface, and this prevents any new oxygen getting at the underlying beryllium to react with it. When the crystal lattices form, so much energy is released that it more than compensates for the energy needed to produce the various ions in the first place. 2.11.3 investigate and describe the reactions of the elements with oxygen, water and dilute acids; Their ions only carry one positive charge, and so the lattice energies of their nitrides will be much less. On the whole, the metals burn in oxygen to form a simple metal oxide. information contact us at info@libretexts.org, status page at https://status.libretexts.org. Beryllium is reluctant to burn unless in the form of powder or dust. This is in contrast to what happens in Group 1 of the Periodic Table (lithium, sodium, potassium, rubidium and caesium). . a) Virtually no reaction occurs between magnesium and cold water. Representative reactions of alkaline earth metals. Lithium is the only metal in Group 1 to form a nitride. Unit AS 2: Further Physical and inorganic Chemistry and an Introdution to Organic Chemistry. This Module addressed why it is difficult to observe a tidy pattern of this reactivity. The reactions of the Group 2 metals with air rather than oxygen is complicated by the fact that they all react with nitrogen to produce nitrides. Calcium is quite reluctant to start burning, but then bursts dramatically into flame, burning with an intense white flame with a tinge of red at the end. You will need to use the BACK BUTTON on your browser to come back here afterwards. However, in a reaction with steam it forms magnesium oxide and hydrogen. The Group II elements are powerful reducing agents. At room temperature, oxygen reacts with the surface of the metal. Magnesium, of course, burns with a typical intense white flame. Energy is evolved when the ions come together to produce the crystal lattice (lattice energy or enthalpy). $2X_{(s)} + O_{2(g)} \rightarrow 2XO_{(s)}$. This works best if the positive ion is small and highly charged - if it has a high charge density. Beryllium, magnesium and calcium don't form peroxides when heated in oxygen, but strontium and barium do. What the metals look like when they burn is a bit problematical! Group 2 have 2 outer electrons which are less easily lost compared to group 1 At the top of group 2 ionisation energies are often too high for the electrons to be removed so they're relatively unreactive, reactivity increases down group 2 also. oxygen, to forma metal oxide with the formula MO where M is the metal and O is oxygen ... (OH)2 is only sparingly soluble. All of these processes absorb energy. As you go down the Group and the positive ions get bigger, they don't have so much effect on the peroxide ion. It would be quite untrue to say that they burn more vigorously as you go down the Group. Ions of the metals at the top of the Group have such a high charge density (because they are so small) that any peroxide ion near them falls to pieces to give an oxide and oxygen. . This is then well on the way to forming a simple oxide ion if the right-hand oxygen atom (as drawn below) breaks off. Strontium and barium will also react with oxygen to form strontium or barium peroxide. $2Mg_{(s)} + O_{2(g)} \rightarrow 2MgO_{(s)}$, $3Mg_{(s)} + N_{2(g)} \rightarrow Mg_3N_{2(s)}$. questions on the reactions of Group 2 elements with air or oxygen, © Jim Clark 2002 (last modified February 2015), reactions of these metals with water (or steam). This is then well on the way to forming a simple oxide ion if the right-hand oxygen atom (as drawn below) breaks off. We also acknowledge previous National Science Foundation support under grant numbers 1246120, 1525057, and 1413739. When the crystal lattices form, so much energy is released that it more than compensates for the energy needed to produce the various ions in the first place. The group 2 metals (M (s)) react with oxygen gas (O 2(g)) at room temperature and pressure to form oxides with the general formula MO as shown in the balanced chemical reactions below: 2Be (s) O 2(g) (3 Marks) (c) Draw The Molecular Orbital Diagram For Peroxide Ion. Energy is evolved when the ions come together to produce the crystal lattice. CaO(s) + H2O(l) ——> Ca(OH)2(s) Hydroxides • basic strength also increases down group • this is because the solubility increases • the metal ions get larger so charge density decreases • there is a lower attraction between the OH¯ ions and larger dipositive ions It is almost impossible to find any trend in the way the metals react with oxygen. As you go down the Group and the positive ions get bigger, they don't have so much effect on the peroxide ion. The peroxide ion, O22- looks like this: The covalent bond between the two oxygen atoms is relatively weak. Once started, the reactions with Oxygen and Chlorine are vigorous: 2Mg(s) + O 2 (g) è2MgO(s) Ca(s) + Cl 2 (g) è CaCl 2 (s) All the metals except Beryllium form oxides in air at room temperature which dulls the surface of the metal. Those reactions don't happen, and the nitrides of sodium and the rest are not formed. REACTIONS OF THE GROUP 2 ELEMENTS WITH AIR OR OXYGEN This page looks at the reactions of the Group 2 elements - beryllium, magnesium, calcium, strontium and barium - with air or oxygen. (3 Marks) (d) Heating Group 2 Carbonates, Such As CaCO3 Leads To Decomposition. Acid-Base reactions are not Redox reactions because there are no changes in Oxidation number. In all the other cases in Group 1, the overall reaction would be endothermic. The excess energy evolved makes the overall process exothermic. The excess energy evolved makes the overall process exothermic. 2:07 understand how displacement reactions involving halogens and halides provide evidence for the trend in reactivity in Group 7; 2:08 (Triple only) explain the trend in reactivity in Group 7 in terms of electronic configurations (c) Gases in the atmosphere. This works best if the positive ion is small and highly charged - if it has a high charge density. Mg ribbon will often have a thin layer of magnesium oxide on it formed by reaction with oxygen. Group 2 reactions Reactivity of group 2 metals increases down the group Mg will also react slowly with oxygen without a flame. Strontium and barium will also react with oxygen to form strontium or barium peroxide. Nitrogen is fairly unreactive because of the very large amount of energy is required to break the triple bond joining the two atoms in the nitrogen molecule, N2. To be able to make any sensible comparison, you would have to have pieces of metal which were all equally free of oxide coating, with exactly the same surface area and shape, exactly the same flow of oxygen around them, and heated to exactly the same extent to get them started. The overall amount of heat evolved when one mole of oxide is produced from the metal and oxygen shows no simple pattern: If anything, there is a slight tendency for the amount of heat evolved to get less as you go down the Group. with $$X$$ representing any group 2 metal. Reactions with water . 2Mg + O2 2MgO This needs to be cleaned off by emery paper before doing reactions with Mg ribbon. Strontium: I have only seen this burn on video. Reactivity increases down the group. Beryllium: I can't find a reference anywhere (text books or internet) to the colour of the flame that beryllium burns with. The Reactions with Air. It would be quite untrue to say that they burn more vigorously as you go down the Group. This is in contrast to what happens in Group 1 of the Periodic Table (lithium, sodium, potassium, rubidium and cesium). The group 1 elements react with oxygen from the air to make metal oxides. You could argue that the activation energy will fall as you go down the Group and that will make the reaction go faster. Reactions with dilute hydrochloric acid All the metals react with dilute hydrochloric acid to give bubbles of hydrogen and a colourless solution of the metal chloride. It can't be done! It would obviously be totally misleading to say that magnesium is more reactive than potassium on the evidence of the bright flame. Beryllium is reluctant to burn unless it is in the form of dust or powder. We say that the positive ion polarizes the negative ion. The strontium equation would look just the same. A/AS level. It explains why it is difficult to observe many tidy patterns. Group 2 elements (beryllium, magnesium, calcium, strontium and barium) react oxygen. (a) describe the redox reactions of the Group 2 elements Mg to Ba: (i) with oxygen, (ii) with water; (b) explain the trend in reactivity of Group 2 elements down the group due to the increasing ease of forming cations, in terms of atomic size, shielding and nuclear attraction; Reactions of Group 2 compounds It explains why it is difficult to observe many tidy patterns. This is compared to non-metals when the reactivity decreases working down a non-metal group such as group 7. It is easier for group 2 elements to lose 2 electrons the further away the electrons are from the nucleus ( as you go down there are more shells), hence the trend is as you go down the group 2 elements the reactivity with oxygen increases. What the metals look like when they burn is a bit problematical! A high charge density simply means that you have a lot of charge packed into a small volume. increases down the group because it becomes more easy to lose the two electrons. This page looks at the reactions of the Group 2 elements - beryllium, magnesium, calcium, strontium and barium - with air or oxygen. The Facts The reactions with oxygen Formation of simple oxides In this case, though, the effect of the fall in the activation energy is masked by other factors - for example, the presence of existing oxide layers on the metals, and the impossibility of controlling precisely how much heat you are supplying to the metal in order to get it to start burning. The group 2 elements react vigorously with oxygen in a redox reaction, forming an oxide with the general formula where is the group 2 element. In all the other Group 1 elements, the overall reaction would be endothermic. Nitrogen is often thought of as being fairly unreactive, and yet all these metals combine with it to produce nitrides, X3N2, containing X2+ and N3- ions. Ca + Cl 2 → CaCl 2. There are no simple patterns. MgO + 2HCl MgCl 2 + H 2O Reactions with oxygen. Missed the LibreFest? It is almost impossible to find any trend in the way the metals react with oxygen. The general equation for the Group is: $3X_{(s)} + N_{2(g)} \rightarrow X_3N_{2(s)}$. The alkali metals react with oxygen. Now imagine bringing a small 2+ ion close to the peroxide ion. The reactions of the Group 2 elements proceed more readily as the energy needed to form positive ions falls. For more information contact us at info@libretexts.org or check out our status page at https://status.libretexts.org. to generate metal oxides. In each case, you will get a mixture of the metal oxide and the metal nitride. Those reactions don't happen, and the nitrides of sodium and the rest aren't formed. This is mainly due to a decrease in ionization energy down the group. The lattice energy is greatest if the ions are small and highly charged - the ions will be close together with very strong attractions. There are no simple patterns in the way the metals burn. The has been reduced from 0 to -2. Nitrogen is often thought of as being fairly unreactive, and yet all these metals combine with it to produce nitrides, X3N2, containing X2+ and N3- ions. An example reaction is shown below: In this reaction, the is oxidised from 0 to +2. For example, Barium peroxide can form because the barium ion is so large that it doesn't have such a devastating effect on the peroxide ions as the metals further up the Group. The reactions of the alkaline earth metals with oxygen are less complex than those of the alkali metals. This is because the less electronegative sodium has a weak Na-O bond and the oxygen is more easily given up to react with H+. Ions of the metals at the top of the Group have such a high charge density (because they are so small) that any peroxide ion near them falls to pieces to give an oxide and oxygen. Lithium has by far the smallest ion in the Group, and so lithium nitride has the largest lattice energy of any possible Group 1 nitride. It is then so hot that it produces the typical intense white flame. Strontium forms this if it is heated in oxygen under high pressures, but barium forms barium peroxide just on normal heating in oxygen. Reaction with halogens. This page looks at the reactions of the Group 2 elements - beryllium, magnesium, calcium, strontium and barium - with air or oxygen. This energy is known as lattice energy or lattice enthalpy. Only in lithium's case is enough energy released to compensate for the energy needed to ionise the metal and the nitrogen - and so produce an exothermic reaction overall. Electrons in the peroxide ion will be strongly attracted towards the positive ion. Electrons in the peroxide ion will be strongly attracted towards the positive ion. . The speed is controlled by factors like the presence of surface coatings on the metal and the size of the activation energy. 2M(s) + O The reaction of Group II Elements with Oxygen. Why do these metals form nitrides on heating in air? The reactions with oxygen. Mg burns with a bright white flame. Their ions only carry one positive charge, and so the lattice energies of their nitrides will be much less. in the air. The overall amount of heat evolved when one mole of oxide is produced from the metal and oxygen also shows no simple pattern: If anything, there is a slight tendency for the amount of heat evolved to decrease as you go down the Group. It reacts with cold water to produce an alkaline solution of calcium hydroxide and hydrogen gas is released. This property is known as deliquescence. While it would be tempting to say that the reactions get more vigorous as you go down the Group, but it is not true. You could argue that the activation energy will fall as you go down the Group and that will make the reaction go faster. In the whole of Group 2, the attractions between the 2+ metal ions and the 3- nitride ions are big enough to produce very high lattice energies. 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