Beryllium is reluctant to burn unless in the form of powder or dust. (3 Marks) (d) Heating Group 2 Carbonates, Such As CaCO3 Leads To Decomposition. 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. It would be quite untrue to say that they burn more vigorously as you go down the Group. You haven't had to heat them by the same amount to get the reactions happening. Beryllium: I can't find a reference anywhere (text books or internet) to the colour of the flame that beryllium burns with. Mixtures of barium oxide and barium peroxide will be produced. $2X_{(s)} + O_{2(g)} \rightarrow 2XO_{(s)}$. Formation of simple oxides. This property is known as deliquescence. For example, the familiar white ash you get when you burn magnesium ribbon in air is a mixture of magnesium oxide and magnesium nitride. 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. The activation energy is much higher. You could argue that the activation energy will fall as you go down the Group and that will make the reaction go faster. What the metals look like when they burn is a bit problematical! This is in contrast to what happens in Group 1 of the Periodic Table (lithium, sodium, potassium, rubidium and caesium). (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.} As I said earlier, they are powerful reducing age… with $$X$$ representing any group 2 metal. to generate metal oxides. (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 Nitrogen is fairly unreactive because of the very large amount of energy needed to break the triple bond joining the two atoms in the nitrogen molecule, N2. 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. This works best if the positive ion is small and highly charged - if it has a high charge density. There are no simple patterns. 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 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. You will notice that the trend in the pH of the solutions formed goes from alkaline to acidic. There is an increase in the tendency to form the peroxide as you go down the Group. This Module addressed why it is difficult to observe a tidy pattern of this reactivity. Trying to pick out patterns in the way the metals burn. This forms a white oxide, which covers the surface. The alkali metals react with oxygen. 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. The activation energy will fall because the ionisation energies of the metals fall. Mg burns with a bright white flame. Unit AS 2: Further Physical and inorganic Chemistry and an Introdution to Organic Chemistry. Anhydrous calcium chloride is a hygroscopic substance that is used as a desiccant. M = Mg, Ca, Sr,Ba --> I will be using 'M' as the general symbol for a Group II element in this topic. 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. Water: Why do these metals form nitrides on heating in air? in the air. 2.11 Group II elements and their compounds. $Ba_{(s)} + O_{2(s)} \rightarrow BaO_{2(s)}$. Reactions of Group 2 Elements with Oxygen, [ "article:topic", "Oxygen", "authorname:clarkj", "barium", "Magnesium", "strontium", "calcium", "Beryllium", "showtoc:no", "Air", "simple oxides", "metal oxides", "Peroxides", "polarizes", "Nitrides" ], https://chem.libretexts.org/@app/auth/2/login?returnto=https%3A%2F%2Fchem.libretexts.org%2FBookshelves%2FInorganic_Chemistry%2FModules_and_Websites_(Inorganic_Chemistry)%2FDescriptive_Chemistry%2FElements_Organized_by_Block%2F1_s-Block_Elements%2FGroup__2_Elements%253A_The_Alkaline_Earth_Metals%2F1Group_2%253A_Chemical_Reactions_of_Alkali_Earth_Metals%2FReactions_of_Group_2_Elements_with_Oxygen, Former Head of Chemistry and Head of Science. Ca(s) + H2O(l) → Ca(OH)2(aq) + H2(g) 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. Mg ribbon will often have a thin layer of magnesium oxide on it formed by reaction with oxygen. a) Virtually no reaction occurs between magnesium and cold water. 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. Their ions only carry one positive charge, and so the lattice energies of their nitrides will be much less. increases down the group because it becomes more easy to lose the two electrons. This is clearly seen if we observe the reactions of magnesium and calcium in water. Those reactions don't happen, and the nitrides of sodium and the rest aren't formed. 1.3.2 (a) Redox Reactions of Group 2 Metals. Reactions with oxygen … At room temperature, oxygen reacts with the surface of the metal. This is in contrast to what happens in Group 1 of the Periodic Table (lithium, sodium, potassium, rubidium and cesium). There are also problems with surface coatings. 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. Question: (a) Write Chemical Equations For The Reactions Of Oxygen With Group 1 Metals And Group 2 Metals Respectiv (4 Marks) (b) Discuss The Trend Of Thermal Stability Of Group 1 And Group 2 Peroxides. 2:09 know the approximate percentages by volume of the four most abundant gases in dry air Reactions with oxygen. Exposed to air, it will absorb water vapour from the air, forming a solution. Reactivity increases down the group. 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. Group 2 elements (beryllium, magnesium, calcium, strontium and barium) react oxygen. MgO + 2HCl MgCl 2 + H 2O Reactions with oxygen. 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. 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+. Describe the trend in the reactivity of group 2 elements with chlorine as you descend down the group. This leads to lower activation energies, and therefore faster reactions. Further along though, a strong S-O bond keeps this together and more H+ is generated. The has been reduced from 0 to -2. The excess energy evolved makes the overall process exothermic. They react violently in pure oxygen producing a white ionic oxide. The reaction of Group II Elements with Oxygen. PERIODIC TABLE GROUP 2 MENU . 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. 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: 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 strontium equation would look just the same. All Group 2 elements tarnish in air to form a coating of the metal oxide. (h) trend in general reactivity of Group 1 and Group 2 metals; Northern Ireland. The speed is controlled by factors like the presence of surface coatings on the metal and the size of the activation energy. In all the other Group 1 elements, the overall reaction would be endothermic. strontium and water reaction. 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 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. The excess energy evolved makes the overall process exothermic. Beryllium, magnesium and calcium don't form peroxides when heated in oxygen, but strontium and barium do. Mixtures of barium oxide and barium peroxide will be produced. 2Mg + O2 2MgO This needs to be cleaned off by emery paper before doing reactions with Mg ribbon. Mg(s) + H2O(g) → MgO(s) + H2(g) b) Calcium is more reactive. Ca + Cl 2 → CaCl 2. The speed is controlled by factors like the presence of surface coatings on the metal and the size of the activation energy. The size of the lattice energy depends on the attractions between the ions. Mg ribbon will often have a thin layer of magnesium oxide on it formed by reaction with oxygen. 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 size of the lattice energy depends on the attractions between the ions. An example reaction is shown below: In this reaction, the is oxidised from 0 to +2. Energy is evolved when the ions come together to produce the crystal lattice (lattice energy or enthalpy). . Strontium: I have only seen this burn on video. The Group II elements are powerful reducing agents. 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. In each case, you will get a mixture of the metal oxide and the metal nitride. The chemical properties of Group2 elements are dominated by the strong reducing power of the metals. The Facts. Beryllium is reluctant to burn unless it is in the form of dust or powder. CCEA Chemistry. 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. The general equation for the Group is: $3X_{(s)} + N_{2(g)} \rightarrow X_3N_{2(s)}$ This is then well on the way to forming a simple oxide ion if the right-hand oxygen atom (as drawn below) breaks off. 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. As you go down the Group and the positive ions get bigger, they don't have so much effect on the peroxide ion. 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. 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. Now imagine bringing a small 2+ ion close to the peroxide ion. The familiar white ash you get when you burn magnesium ribbon in air is a mixture of magnesium oxide and magnesium nitride (despite what you might have been told when you were first learning Chemistry!). On the whole, the metals burn in oxygen to form a simple metal oxide. Why do these metals form nitrides on heating in air? While it would be tempting to say that the reactions get more vigorous as you go down the Group, but it is not true. REACTIONS OF THE GROUP 2 ELEMENTS WITH AIR OR OXYGEN. When these metals (M) are heated in oxygen they burn vigorously to produce a white ionic oxide, M2+O2-. Those reactions don't happen, and the nitrides of sodium and the rest are not formed. 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. 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. The Facts The reactions with oxygen Formation of simple oxides In all the other cases in Group 1, the overall reaction would be endothermic. The equations for the reactions: Beryllium, magnesium and calcium don't form peroxides when heated in oxygen, but strontium and barium do. Group 2 reactions Reactivity of group 2 metals increases down the group Mg will also react slowly with oxygen without a flame. The group 2 metals will burn in 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. Reaction with oxygen. All group 2 elements want to lose 2 electrons and all group 6 elements (oxygen) want to gain 2 electrons. The Reactions with Air. Mg + 2 H2O Mg(OH) 2 + H2 This is a much slower reaction than the reaction with steam and there is no flame. Why do some metals form peroxides on heating in oxygen? The lattice energy is greatest if the ions are small and highly charged - the ions will be close together with very strong attractions. 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. 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. It would be tempting to say that the reactions get more vigorous as you go down the Group, but it isn't true. e.g. It can't be done! We say that the positive ion polarizes the negative ion. It explains why it is difficult to observe many tidy patterns. Reactions with water . But how reactive a metal seems to be depends on how fast the reaction happens - not the overall amount of heat evolved. For more information contact us at info@libretexts.org or check out our status page at https://status.libretexts.org. As you go down the Group and the positive ions get bigger, they don't have so much effect on the peroxide ion. 2Mg + O 2MgO Mg will also react with warm water, giving a different magnesium hydroxide product. This page looks at the reactions of the Group 2 elements - beryllium, magnesium, calcium, strontium and barium - with air or oxygen. This is mainly due to a decrease in ionization energy down the group. Magnesium, on the other hand, has to be heated to quite a high temperature before it will start to react. 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. 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. A reducing agent is the compound that gets oxidised in the reaction and, therefore, loses electrons. What the metals look like when they burn is a bit problematical! It would be quite untrue to say that they burn more vigorously as you go down the Group. A high charge density simply means that you have a lot of charge packed into a small volume. All of these processes absorb energy. The reactions with oxygen. 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) It explains why it is difficult to observe many tidy patterns. Strontium forms this if it is heated in oxygen under high pressures, but barium forms barium peroxide just on normal heating in oxygen. We say that the positive ion polarises the negative ion. The strontium equation would look just the same. It is then so hot that it produces the typical intense white flame. We also acknowledge previous National Science Foundation support under grant numbers 1246120, 1525057, and 1413739. 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. Group 2 reactions Reactivity of group 2 metals increases down the group Mg will also react slowly with oxygen without a flame. Oxygen: All of the elements in group 2 react vigorously with Oxygen, the product of which is an ionic oxide. You could argue that the activation energy will fall as you go down the Group and that will make the reaction go faster. Watch the recordings here on Youtube! There are no simple patterns in the way the metals burn. If this is the first set of questions you have done, please read the introductory page before you start. Reaction with water Most Group II oxides react with water to produce the hydroxide e.g. Strontium and barium will also react with oxygen to form strontium or barium peroxide. The covalent bond between the two oxygen atoms is relatively weak. Discusses trends in atomic radius, ionisation energy, electronegativity and melting point of the Group 2 elements. The reactions of the Group 2 elements proceed more readily as the energy needed to form positive ions falls. This is compared to non-metals when the reactivity decreases working down a non-metal group such as group 7. Acid-Base reactions are not Redox reactions because there are no changes in Oxidation number. 2Li(s) + Cl 2 (g) → 2LiCl(s) A similar reaction takes place with the other elements of group 7. The lattice energy is greatest if the ions are small and highly charged - the ions will be close together with very strong attractions. Strontium forms this if it is heated in oxygen under high pressures, but barium forms barium peroxide just on normal heating in oxygen. Why do some metals form peroxides on heating in oxygen? The group 1 elements react with oxygen from the air to make metal oxides. Beryllium is reluctant to burn unless it is in the form of dust or powder. All of these processes absorb energy. It cannot be said that by moving down the group these metals burn more vigorously. There is an increase in the tendency to form the peroxide as you go down the Group. The general equation for the Group is: $3X_{(s)} + N_{2(g)} \rightarrow X_3N_{2(s)}$. . 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. 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