Chemistry Pack Level 10 - Semester 1 PDF

Chemistry Pack Level 10 Ms Tiziana Grech Chemistry Level 10 Pack Topic 1: The Alkali Metals.......................................................................................

Chemistry Pack Level 10 Ms Tiziana Grech Chemistry Level 10 Pack Topic 1: The Alkali Metals....................................................................................................... 2 Topic 2: The Alkaline Earth Metals......................................................................................... 21 Topic 3: The Limestone Cycle in the Industry......................................................................... 39 Topic 4: Water Hardness...................................................................................................... 46 Topic 5: The Chemistry of The Halogens................................................................................. 57 Topic 6: The Chemistry of Transition Metals.......................................................................... 75 Topic 7: Reaction Rates....................................................................................................... 102 Page 1 of 122 Chemistry Pack Level 10 Ms Tiziana Grech Topic 1: The Alkali Metals The Periodic Table.............................................................................................................. 3 Electronic configuration and the Periodic table................................................................. 4 The Properties and Trends of Group 1 Metals....................................................................... 5 Physical properties of Group 1 Metals.............................................................................. 5 Why are these physical changes occurring down the group?.............................................. 6 The Chemical properties of Group 1 Metals.......................................................................... 7 Why is reactivity increasing down Group I?...................................................................... 7 Reactions of Group 1 Metals with Oxygen............................................................................. 8 Reactions of the Group 1 Metals with Water........................................................................ 9 Reactions of Group I metals with chlorine gas.................................................................... 10 Classwork 1.1 – Group 1 Metals..................................................................................... 11 Classwork 1.2 – Properties of Group 1 Metals................................................................. 11 Homework 1.1 – Chemical and Physical Properties of Group 1 Metals............................... 13 School-Based Assessment: Group 1 Metals Poster Project................................................... 15 Checkpoint....................................................................................................................... 18 Assessment Checkpoint..................................................................................................... 20 Page 2 of 122 Chemistry Pack Level 10 Ms Tiziana Grech The Periodic Table Elements are arranged on the Periodic Table in order of increasing atomic number, where each element has one proton more than the element preceding it. The table is arranged in vertical columns called Groups numbered 1 – 8 and in rows called Periods. Period: these are the horizontal rows that show the number of shells of electrons an atom has. Example: elements in Period 2 have two electron shells, elements in Period 3 have three electron shells. Group: these are the vertical columns that show how many outer electrons each atom has. Example: Group 4 elements have atoms with 4 electrons in the outermost shell. Page 3 of 122 Chemistry Pack Level 10 Ms Tiziana Grech Electronic configuration and the Periodic table The electronic configuration is the arrangement of electrons into shells for an atom (the electronic configuration of carbon is 2, 4). There is a link between the electronic configuration of the elements and their position on the Periodic table. The number of notations in the electronic configuration will show the number of shells of electrons the atom has, showing the Period. The last notation shows the number of outer electrons the atom has, showing the Group number. Example: Electron configuration of Chlorine: Period: The red numbers at the bottom show the number of notations which is 3, showing that a chlorine atom has 3 shells of electrons. Group: The last notation, which is 7, showing that a chlorine atom has 7 outer electrons in its outer shell therefore it must be found in group 7. Chemical properties of elements in the same group Elements in the same Group in the Periodic table have similar chemical properties. When atoms collide and react, it is the outermost electrons that interact. The similarity in their chemical properties stems from having the same number of electrons in their outer shell. For example, both lithium and sodium are in Group 1 and can react with elements in Group 7 to form an ionic compound (charges of Group 1 ions are +1, charges of Group 7 ions are -1). Page 4 of 122 Chemistry Pack Level 10 Ms Tiziana Grech The Properties and Trends of Group 1 Metals The Group 1 metals are also called the alkali metals as they form alkaline solutions with high pH values when reacted with water. Group 1 metals are lithium, sodium, potassium, rubidium, caesium, and francium. They all contain just one electron in their outer shell. Physical properties of Group 1 Metals. 1) Are soft and easy to cut, getting softer and denser as you move down the Group. a) IMPORTANT: sodium and potassium do not follow the trend in density, potassium is less dense than sodium because despite potassium having a greater atomic mass than sodium, the increase in volume (due to the larger atomic size) is more significant. This results in a lower density because density is mass per unit volume, and the increased volume outweighs the mass difference. 2) Have shiny silvery white surfaces when freshly cut. 3) Conduct heat and electricity. 4) They all have low melting points and low boiling points, which decreases as you move down the Group. The values in this table are not for the exam. Page 5 of 122 Chemistry Pack Level 10 Ms Tiziana Grech Why are these physical changes occurring down the group? Ongoing down group 1: 1. The atomic radius (radius of the atom, the distance from the centre of the nucleus to the circumference of the outermost shell ) is increasing. The atom is becoming larger because there are more shells. 2. The electron shielding is increasing (inner electrons repel outer electrons due to their negative charge, partially counteracting the pull of the nucleus on the outer electrons ). 3. The electrostatic force of attraction between the positive metal cations and delocalised electrons becomes weaker. 4. Therefore, the weaker the electrostatic force the weaker the metallic bond therefore the softer and the lower the melting and boiling points (less heat energy is needed to break the metallic bond). Page 6 of 122 Chemistry Pack Level 10 Ms Tiziana Grech The Chemical properties of Group 1 Metals Ongoing down group 1, reactivity increases. This means that potassium is more reactive than sodium. Why is reactivity increasing down Group I? 1) As you go down Group 1, the atomic radius and electron shielding increases. 2) This means that the outer electron (valence electron) is further away from the nucleus so there are weaker electrostatic forces of attraction pulling it towards the nucleus. 3) When group 1 metals react, they react by losing the valence electron, ongoing down group 1 the atomic radius increases and electron shielding increases making the electrostatic forces weaker. This means that the valence electron can be lost easier, hence the metal reacts quicker. 4) Therefore, potassium is more reactive than sodium and so on. Page 7 of 122 Chemistry Pack Level 10 Ms Tiziana Grech Reactions of Group 1 Metals with Oxygen Sodium 4Na(s) + O2(g) → 2Na2O(s) Sodium burns with a luminous yellow flame. A white solid is produced (sodium oxide). Potassium 4K(s) + O2(g) → 2K2O(s) Potassium burns with a lilac flame. A white solid is produced (potassium oxide). Page 8 of 122 Chemistry Pack Level 10 Ms Tiziana Grech Reactions of the Group 1 Metals with Water Sodium 2Na(s) + 2H2O(l) → 2NaOH(aq) + H2(g) Fizzing/Bubbling: Sodium reacts vigorously, producing hydrogen gas, which causes fizzing or bubbling on the water's surface. Movement: The sodium piece often moves rapidly across the surface of the water as the reaction occurs. Heat Production: The reaction is exothermic, releasing heat that may melt the sodium into a small, silvery ball. Hydroxide Formation: A colourless solution of sodium hydroxide (NaOH) forms, which is basic and can turn indicators like phenolphthalein pink. Possible Flames: In some cases, the hydrogen gas produced can ignite, burning with a yellow-orange flame. Potassium 2K(s) + 2H2O(l) → 2KOH(aq) + H2(g) Violent Reaction: The reaction is more vigorous than that of sodium with water, often occurring explosively. Fizzing/Bubbling: Potassium produces hydrogen gas rapidly, leading to intense fizzing or bubbling. Movement: The potassium piece darts around quickly on the water's surface. Heat and Flames: The reaction is highly exothermic, often igniting the hydrogen gas and producing a lilac-colored flame. Melting: Potassium usually melts into a small, silvery ball due to the heat generated. Hydroxide Formation: A colourless, basic solution of potassium hydroxide (KOH) is formed, which can turn indicators like phenolphthalein pink. Page 9 of 122 Chemistry Pack Level 10 Ms Tiziana Grech Reactions of Group I metals with chlorine gas Sodium 2Na(s) + Cl2(g) → 2NaCl(s) Sodium catches fire and burns with a luminous yellow flame in the green chlorine gas. A white solid is produced (sodium chloride). Potassium 2K(s) + Cl2(g) → 2KCl(s) Potassium catches fire and burns with a lilac flame in the green chlorine gas. A white solid is produced (potassium chloride). Page 10 of 122 Chemistry Pack Level 10 Ms Tiziana Grech Classwork 1.1 – Group 1 Metals. May 2022 paper 1 Q. 9 Classwork 1.2 – Properties of Group 1 Metals. 1. What elements are found in Group 1 of the periodic table? (2) 2. Why are Group 1 elements called alkali metals? (2) 3. How does the reactivity of Group 1 elements change as you move down the group? (2) 4. What is the electron configuration for the first 3 metals in Group 1? (3) Page 11 of 122 Chemistry Pack Level 10 Ms Tiziana Grech 5. Explain why alkali metals have low melting and boiling points. (2) 6. How does the density of alkali metals change as you move down the group? (2) 7. Describe the reaction between sodium and water, including the products formed. (2) 8. Why do Group 1 elements have only one valence electron? (2) 9. What happens to the atomic radius of alkali metals as you go down the group? (2) 10. Compare and explain the reactivity of sodium with that of potassium when reacting with water. (3) Total: 22 marks Page 12 of 122 Chemistry Pack Level 10 Ms Tiziana Grech Homework 1.1 – Chemical and Physical Properties of Group 1 Metals. 1. Why do alkali metals tend to form +1 ions? (2) 2. How does the energy needed to lose the valance electron change as you move down the group? (2) 3. What safety precautions are necessary when handling alkali metals in the laboratory? (2) 4. Describe the flame test colours for lithium, sodium, and potassium. (2) 5. What is electron shielding, and how does it affect the reactivity of alkali metals? (3) 6. How do alkali metals react with halogens, and what type of compounds do they form? (2) 7. Why are alkali metals stored under oil? (2) Page 13 of 122 Chemistry Pack Level 10 Ms Tiziana Grech 8. Explain why francium is the least stable of the alkali metals. (2) 9. List three observations that one can notice when sodium reacts with water. (3) 10. Which of the group 1 metals is the most dangerous and why? (2) Total: 22 marks Page 14 of 122 Chemistry Pack Level 10 Ms Tiziana Grech School-Based Assessment: Group 1 Metals Poster Project Title: Exploring the Chemistry of Group 1 Metals (Alkali Metals) Objective: To assess students' understanding of the properties, trends, and uses of Group 1 metals by creating an informative and visually engaging poster. Assessment Overview: Students will design a poster that highlights the key aspects of ONE of the metals in Group 1, including its physical and chemical properties, trends within the group, real- world applications, and safety considerations. The poster should be both educational and visually appealing, suitable for display in the classroom. Instructions: Students are required to choose one metal from group 1 and create a poster that covers the following topics: 1. Catchy Title! Title: Include a clear and catchy title for your poster. 2. Physical and Chemical Properties: Physical Properties: Include details about the physical properties of the metal you chose, such as colour, density, melting points, and softness. Use visuals or icons to represent these properties. Chemical Properties: Describe the key chemical properties, such as high reactivity, the formation of +1 ions, and reactions with water. Illustrate a simple chemical equation showing the reaction of the metal chosen with water. 3. Reactivity and Trends: Reactivity Trend: Visually represent how the reactivity of alkali metals increases as you move down the group. Include a brief explanation of why this trend occurs, focusing on factors like atomic size and electron shielding. Flame Colours: Provide examples of the flame colour for the metal chosen. Consider using coloured graphics or drawings. 4. Real-World Applications: Uses of Alkali Metals: Highlight at least three real-world applications of the metal chosen. Page 15 of 122 Chemistry Pack Level 10 Ms Tiziana Grech 5. Safety Considerations: Safety Precautions: List the safety measures that should be taken when handling the metal chosen, both in the laboratory and in industrial settings. Use warning symbols or visuals to emphasize these points. Storage and Handling: Briefly explain why the metal chosen is stored under oil and the dangers associated with their reactivity. 6. Creative and Visual Elements: Design: Ensure that the poster is visually appealing with a clear layout. Use headings, bullet points, and sections to organize the information. Graphics and Diagrams: Incorporate relevant images, diagrams, and charts to enhance understanding and engagement. Colour Scheme: Use a consistent and attractive colour scheme that highlights key information without overwhelming the viewer. 7. Conclusion: Summary: Include a brief concluding section summarizing the importance of understanding Group 1 metals in both chemistry and everyday life. Reflection: Add a personal reflection on what you found most interesting about the metal chosen. Submission: Posters should be A2 size. Cardboard will be given to you in class to work together. Include your names somewhere on the poster. We are all proud of the work we do. Posters will be displayed in the classroom, and students will give a brief (2-3 minute) presentation explaining their work. After which the poster can be hang elsewhere in the school. Page 16 of 122 Chemistry Pack Level 10 Ms Tiziana Grech Assessment Criteria: Criteria Marks The poster should provide accurate and relevant Content Accuracy /6 information. The design should be engaging and easy to read, Visual Appeal with effective use of colour, /6 images, and layout. The poster should reflect creativity in presenting Creativity /6 information in a unique and visually compelling way. Information should be well- organized, with clear Clarity and Organization /6 headings and logical flow. Demonstrates an understanding of the trends, Critical Thinking properties, and applications /6 of alkali metals. Remarks Page 17 of 122 Chemistry Pack Level 10 Ms Tiziana Grech Checkpoint Alkali Metals: The elements in Group 1 of the periodic table, which include lithium (Li), sodium (Na), potassium (K), rubidium (Rb), cesium (Cs), and francium (Fr). Reactivity: The tendency of a substance to undergo chemical reactions. In Group 1 metals, reactivity increases as you move down the group. Electron Configuration: The arrangement of electrons around the nucleus of an atom. Group 1 metals have a single electron in their outermost shell, leading to their high reactivity. Atomic Radius: The distance from the nucleus to the outermost electron. The atomic radius increases as you move down Group 1 due to the addition of electron shells. Flame Test: A qualitative test used to identify metal ions based on the colour they produce when heated in a flame. Group 1 metals show characteristic flame colours (e.g., sodium gives a yellow flame). Electron Shielding: The phenomenon where inner electrons block the effective nuclear charge from reaching outer electrons. In Group 1 metals, increased shielding contributes to greater reactivity down the group. Density: The mass per unit volume of a substance. Group 1 metals have low densities, with density increasing slightly as you move down the group. Melting Point: The temperature at which a solid turns into a liquid. Group 1 metals have low melting points, which decrease as you move down the group. Hydroxides: Compounds formed when Group 1 metals react with water, producing a metal hydroxide (e.g., sodium hydroxide, NaOH) and hydrogen gas. Exothermic Reaction: A reaction that releases heat. The reaction of Group 1 metals with water is highly exothermic. Basicity: The property of a substance that determines its ability to act as a base. Group 1 metal hydroxides are strong bases. Alkaline Solutions: Solutions that are basic in nature, often resulting from the dissolution of Group 1 metal hydroxides in water. Metallic Bonding: The type of bonding found in metals where electrons are free to move throughout the structure. Group 1 metals exhibit metallic bonding, which contributes to their conductivity and malleability. Page 18 of 122 Chemistry Pack Level 10 Ms Tiziana Grech Soft Metals: A characteristic of Group 1 metals, which are softer and can be cut with a knife, especially as you move down the group. Storage Under Oil: A safety practice where alkali metals are stored under oil to prevent them from reacting with moisture or oxygen in the air. Page 19 of 122 Chemistry Pack Level 10 Ms Tiziana Grech Assessment Checkpoint Assignment Mark Remarks Excellent Job Classwork 1.1 – Group 1 Metals. Good Progress Better Effort Needs Improvement Excellent Job Classwork 1.2 – Properties of Group 1 Good Progress Metals. Better Effort Needs Improvement Excellent Job Homework 1.1 – Chemical and Physical Good Progress Properties of Group 1 Metals. Better Effort Needs Improvement Excellent Job School-Based Assessment: Group 1 Metals Good Progress Poster Project Better Effort Needs Improvement Excellent Job Test 1: The Alkali Metals Good Progress Better Effort Needs Improvement Page 20 of 122 Chemistry Pack Level 10 Ms Tiziana Grech Topic 2: The Alkaline Earth Metals Physical properties of The Alkaline Earth Metals................................................................. 22 Why do these physical changes occur down Group 2?...................................................... 23 Chemical properties of the Alkaline Earth Metals................................................................ 24 Reactions of The Alkaline Earth Metals with Oxygen........................................................... 25 Reactions of The Alkaline Earth Metals with Water............................................................. 26 Reactions of The Alkaline Earth Metals with Acids.............................................................. 27 Classwork 2.1: The Alkaline Earth Metals....................................................................... 28 Classwork 2.2: Chemical and Physical Properties of The Alkaline Earth Metals................. 29 Homework 2.1: The Alkaline Earth Metals and Their Properties – SEC Past Paper Questions..................................................................................................................................... 30 School-Based Assessment: Reactivity of Group 2 Metals (Magnesium and Calcium).............. 33 Checkpoint....................................................................................................................... 37 Assessment Checkpoint..................................................................................................... 38 Page 21 of 122 Chemistry Pack Level 10 Ms Tiziana Grech Physical properties of The Alkaline Earth Metals The Group 2 metals are called the alkaline earth metals as they form most of the Earth’s crust and they form alkaline solution with high pH values when reacted with water. The metals in Group 2 are beryllium, magnesium, calcium, strontium, barium, and radium. They all have two electrons in their outer shell. The following are the common physical properties of these metals: 1) These metals are harder and not easy to cut when compared to group 1 metals. 2) They get softer and denser as you move down the Group. 3) They have shiny grey surfaces when freshly cut. 4) They conduct heat and electricity. 5) They all have higher melting points and higher densities when compared with Group 1 metals. 6) The melting point decreases ongoing down Group 2. Page 22 of 122 Chemistry Pack Level 10 Ms Tiziana Grech Why do these physical changes occur down Group 2? Ongoing down Group 2: 1. The atomic radius (radius of the atom, the distance from the centre of the nucleus to the circumference of the outermost shell ) is increasing. The atom is becoming larger because there are more shells. 2. The electron shielding is increasing (inner electrons repel outer electrons due to their negative charge, partially counteracting the pull of the nucleus on the outer electrons ). 3. The electrostatic force of attraction between the positive metal cations and delocalised electrons becomes weaker. 4. Therefore, the weaker the electrostatic force the weaker the metallic bond therefore the softer and the lower the melting and boiling points (less heat energy is needed to break the metallic bond). Page 23 of 122 Chemistry Pack Level 10 Ms Tiziana Grech Chemical properties of the Alkaline Earth Metals Ongoing down Group 2, the reactivity of the metals increases. 1) As you go down Group 2, the atomic radius and electron shielding increases. 2) This means that the outer electron (valence electron) is further away from the nucleus so there are weaker electrostatic forces of attraction pulling it towards the nucleus. 3) When group 2 metals react, they react by losing their valence electrons, ongoing down group 2 the atomic radius increases and electron shielding increases making the electrostatic forces weaker. This means that the valence electrons can be lost easier, hence the metal reacts quicker. 4) Therefore, calcium is more reactive than magnesium and so on. Calcium reacting with water Page 24 of 122 Chemistry Pack Level 10 Ms Tiziana Grech Reactions of The Alkaline Earth Metals with Oxygen They react with oxygen and water vapour in air, in fact calcium is usually kept under oil to stop it from reacting. 1) Magnesium catches fire and burns with a bright white flame. 2Mg(s) + O2(g) → 2MgO(s) Magnesium oxide is a strong base. 2) Calcium catches fire and burns with a brick-red flame. 2Ca(s) + O2(g) → 2CaO(s) Calcium oxide is a strong base. Important! When Group 1 and Group 2 metals burn in air, they also react with nitrogen to produce the metal nitride. The product is a metal nitride. 3Mg(s) + N2(g) → Mg3N2(s) Page 25 of 122 Chemistry Pack Level 10 Ms Tiziana Grech Reactions of The Alkaline Earth Metals with Water Group 2 metals, namely magnesium and calcium react differently with water. 1) Magnesium reacts extremely slowly with cold water (it seems like it is not reacting at all). Mg(s) + H2O(l) → no reaction 2) Magnesium can react vigorously with steam, which is water in the gaseous state. Steam has a higher kinetic energy than water and therefore the rate of reaction is significantly increased. Mg(s) + H2O(g) → MgO(s) + H2(g) 1) Calcium on the other hand reacts vigorously with cold water to produce an alkaline solution of calcium hydroxide and a colourless, odourless gas, hydrogen. Ca(s) + 2H2O(l) → Ca (OH)2(aq) + H2(g) Some of the calcium hydroxide produces sediments at the bottom of the beaker because it is very soluble in water. Page 26 of 122 Chemistry Pack Level 10 Ms Tiziana Grech Reactions of The Alkaline Earth Metals with Acids Group 2 metals react with acids to produce salt, and hydrogen gas. 1. Magnesium reacts rapidly with acids to produce salt and hydrogen gas. Mg(s) + 2HCl(aq) → MgCl2(aq) + H2(g) 2. Calcium reacts vigorously with acids to produce salt and hydrogen gas. Ca(s) + 2HCl(aq) → CaCl2(aq) + H2(g) Page 27 of 122 Chemistry Pack Level 10 Ms Tiziana Grech Classwork 2.1: The Alkaline Earth Metals 1) List all the elements in Group 2 of the periodic table. (2) 2) Explain why Group 2 elements are called alkaline earth metals. (1) 3) Describe the trend in reactivity of Group 2 elements as you move down the group. (2) 4) What is the electron configuration of the first three metals in Group 2? (2) 5) Explain why the atomic radius of Group 2 elements increases as you move down the group. (2) 6) How does the solubility of Group 2 hydroxides in water change as you move down the group? (2) 7) Describe the trend in the melting points of Group 2 elements as you move down the group. (2) 8) Explain why Group 2 elements typically form +2 ions. (2) Page 28 of 122 Chemistry Pack Level 10 Ms Tiziana Grech 9) What is the flame test colour for calcium? How can this be used to identify calcium in a compound? (2) Total: 17 marks Classwork 2.2: Chemical and Physical Properties of The Alkaline Earth Metals. 1) Write a balanced chemical equation for the reaction of magnesium with oxygen. (2) 2) Describe the reaction of calcium with water. (2) 3) Write the balanced chemical equation for this reaction. (2) 4) Write a balanced chemical equation for the reaction of barium with hydrochloric acid. (2) 5) Predict the product(s) and write a balanced chemical equation for the reaction between beryllium and chlorine. (3) 6) Predict and write a balanced chemical equation for the reaction between calcium and nitrogen. (2) Total: 12 marks Page 29 of 122 Chemistry Pack Level 10 Ms Tiziana Grech Homework 2.1: The Alkaline Earth Metals and Their Properties – SEC Past Paper Questions. May 2021 paper 1 Q. 1 May 2021 paper 1 Q. 2 Page 30 of 122 Chemistry Pack Level 10 Ms Tiziana Grech May 2019 paper 2A Q. 14 Page 31 of 122 Chemistry Pack Level 10 Ms Tiziana Grech Total: 16 marks Page 32 of 122 Chemistry Pack Level 10 Ms Tiziana Grech School-Based Assessment: Reactivity of Group 2 Metals (Magnesium and Calcium) Title: Comparing the Reactivity of Magnesium and Calcium with Water and Hydrochloric Acid Objective: To assess students' understanding of the reactivity of Group 2 metals by comparing the reactions of magnesium and calcium with water and hydrochloric acid. Assessment Overview: Students will conduct an experiment to observe and compare the reactivity of magnesium and calcium when they react with water and hydrochloric acid. They will record their observations, write balanced chemical equations, analyse the data, and draw conclusions about the relative reactivity of these metals. Materials: Magnesium ribbon Calcium granules Distilled water Hydrochloric acid Test tubes Test tube rack Measuring cylinder Thermometer pH indicator – phenolphthalein Safety goggles and gloves Safety Precautions: 1) Wear safety goggles and gloves at all times during the experiment. 2) Handle hydrochloric acid with care, avoiding skin contact. 3) Conduct the experiment in a well-ventilated area or under a fume hood. Experimental Procedure: Part 1: Reaction with Water 1) Setup: a) Fill two test tubes with 10 mL of distilled water. b) Add a small piece of magnesium ribbon to the first test tube. c) Add a few granules of calcium to the second test tube. Page 33 of 122 Chemistry Pack Level 10 Ms Tiziana Grech 2) Observation: a) Observe the reactions in each test tube. Record any bubbles, temperature change, or other noticeable changes in the Observations Table. b) Test the resulting solution with a pH indicator. Part 2: Reaction with Hydrochloric Acid 1) Setup: a) Fill two test tubes with 10 mL of hydrochloric acid. b) Add a small piece of magnesium ribbon to the first test tube. c) Add a few granules of calcium to the second test tube. 2) Observation: a) Observe the reactions in each test tube, focusing on the rate of gas production (bubbling) and any temperature change in the Observations Table. b) Capture any gas produced using the boiling tube, by placing it upside down on the test- tube. c) Test for the gas collected by using a burning wooden splint. Place the burning wooden splint in boiling tube. VERY important to leave the tube upside down. Page 34 of 122 Chemistry Pack Level 10 Ms Tiziana Grech Data Collection: Magnesium with water Calcium with water Magnesium with hydrochloric acid Calcium with hydrochloric acid Analysis Questions: 1) Write the balanced chemical equations for the reactions of magnesium and calcium with water. (4) 2) Write the balanced chemical equations for the reactions of magnesium and calcium with hydrochloric acid. (4) Page 35 of 122 Chemistry Pack Level 10 Ms Tiziana Grech 3) Compare the reactivity of magnesium and calcium with water. Which metal is more reactive and why? Your reason should reflect the observations in this experiment. (2) 4) Compare the reactivity of magnesium and calcium with hydrochloric acid. Which metal is more reactive and why? Your reason should reflect the observations in this experiment. (2) 5) Explain the trend in reactivity within Group 2 metals based on your observations. (2) 6) Discuss any differences in the nature of the reactions when comparing water and hydrochloric acid as reactants. (2) 7) How does the pH of the solution change after each reaction? What does this indicate about the nature of the products formed? (2) Total: 18 marks Page 36 of 122 Chemistry Pack Level 10 Ms Tiziana Grech Checkpoint Alkaline Earth Metals: The elements in Group 2 of the periodic table, which include beryllium (Be), magnesium (Mg), calcium (Ca), strontium (Sr), barium (Ba), and radium (Ra). Reactivity: The tendency of a substance to undergo chemical reactions. For Group 2 metals, reactivity increases as you move down the group. Electron Configuration: The arrangement of electrons in an atom. Group 2 metals have two electrons in their outermost shell. Atomic Radius: The distance from the nucleus to the outermost electron. The atomic radius increases as you move down Group 2 due to the addition of electron shells. Flame Test: A qualitative test used to identify metal ions based on the colour they produce when heated in a flame. Group 2 metals produce characteristic flame colours (e.g., calcium gives an orange-red flame). Density: The mass per unit volume of a substance. Group 2 metals generally have higher densities compared to Group 1 metals, with density increasing down the group. Melting Point: The temperature at which a solid turns into a liquid. Group 2 metals typically have higher melting points than Group 1 metals, though the melting point decreases down the group. Hydroxides: Compounds formed when Group 2 metals react with water, producing a metal hydroxide (e.g., calcium hydroxide, Ca(OH)₂) and hydrogen gas. These hydroxides are generally basic in nature. Oxides: Compounds formed when Group 2 metals react with oxygen, resulting in metal oxides (e.g., magnesium oxide, MgO). These oxides are basic and react with water to form hydroxides. Basicity: The property of a substance to act as a base. Group 2 metal hydroxides are strong bases and increase in basicity as you move down the group. Reacting with Water: Group 2 metals react with water to form hydroxides and hydrogen gas. The reaction becomes more vigorous as you move down the group. Page 37 of 122 Chemistry Pack Level 10 Ms Tiziana Grech Assessment Checkpoint Assignment Mark Remarks Excellent Job Classwork 2.1: The Alkaline Earth Metals Good Progress Better Effort Needs Improvement Excellent Job Homework 2.1: Chemical and Physical Good Progress Properties of The Alkaline Earth Metals. Better Effort Needs Improvement Homework 2.1: The Alkaline Earth Metals Excellent Job and Their Properties – SEC Past Paper Good Progress Questions. Better Effort Needs Improvement Excellent Job School-Based Assessment: Reactivity of Good Progress Group 2 Metals (Magnesium and Calcium) Better Effort Needs Improvement Excellent Job Test 2: The Alkaline Earth Metals Good Progress Better Effort Needs Improvement Page 38 of 122 Chemistry Pack Level 10 Ms Tiziana Grech Topic 3: The Limestone Cycle in the Industry The limestone cycle........................................................................................................... 40 Importance of the Limestone Cycle in Industry.................................................................. 41 Classwork 3.1: The Limestone Cycle............................................................................... 42 Homework 3.1: The Limestone Cycle – SEC Past Paper Questions..................................... 43 Checkpoint....................................................................................................................... 44 Assessment Checkpoint..................................................................................................... 45 Page 39 of 122 Chemistry Pack Level 10 Ms Tiziana Grech The limestone cycle The limestone cycle describes the chemical reactions involved in the transformation of limestone (calcium carbonate, CaCO₃) through various stages and its eventual return to its original form. This cycle is fundamental in industrial processes, especially in the construction and manufacturing industries. Steps in the Limestone Cycle: 1. Calcination (Heating of Limestone): o Limestone is heated at high temperatures (around 900°C) in a kiln. o Chemical Reaction: CaCO3 → CaO + CO2 o Product: Calcium oxide (quicklime) and carbon dioxide gas. 2. Slaking of Quicklime: o Quicklime (CaO) reacts with water in an exothermic reaction to form calcium hydroxide (slaked lime). Reaction is dangerous therefore safety gear must be worn. o Chemical Reaction: CaO+H2O→Ca(OH)2 o Product: Calcium hydroxide (slaked lime). 3. Carbonation: o Calcium hydroxide reacts with carbon dioxide in the air to form calcium carbonate again, completing the cycle. o Chemical Reaction: Ca(OH)2+CO2→CaCO3+H2O o Product: Calcium carbonate (limestone) and water. Page 40 of 122 Chemistry Pack Level 10 Ms Tiziana Grech Importance of the Limestone Cycle in Industry 1. Cement and Concrete Production: The limestone cycle is crucial in the production of cement, where limestone is heated to produce quicklime. This quicklime is then mixed with other materials to produce clinker, the primary component of cement. Concrete, a mixture of cement, sand, and aggregates, is essential in construction. 2. Construction Materials: Slaked lime (calcium hydroxide) is used in making mortar and plaster. Its ability to set by reacting with carbon dioxide to form calcium carbonate gives it strength and durability, which are essential properties for building materials. 3. Soil Stabilization: In road construction and soil improvement, quicklime is used to stabilize and strengthen the soil by reducing moisture content and binding particles together. 4. Water Treatment: The cycle's products are used to treat water, neutralizing acidity and removing impurities. Calcium hydroxide (slaked lime) is particularly important in softening water and treating wastewater. 5. Environmental Applications: The cycle plays a role in flue gas desulfurization, where calcium carbonate or calcium hydroxide is used to remove sulfur dioxide from industrial emissions, reducing air pollution. 6. Agriculture: Calcium oxide and calcium hydroxide are used to neutralize acidic soils, improving crop yields and soil health. Page 41 of 122 Chemistry Pack Level 10 Ms Tiziana Grech Classwork 3.1: The Limestone Cycle May 2021 paper 1 Q. 9 Page 42 of 122 Chemistry Pack Level 10 Ms Tiziana Grech Homework 3.1: The Limestone Cycle – SEC Past Paper Questions September 2021 paper 2B Q. 5 Page 43 of 122 Chemistry Pack Level 10 Ms Tiziana Grech Checkpoint Limestone (Calcium Carbonate, CaCO₃): A naturally occurring sedimentary rock composed primarily of calcium carbonate. It is the starting material in the limestone cycle and is widely used in construction, agriculture, and industrial processes. Calcination: The process of heating limestone (calcium carbonate) at high temperatures (around 900°C) to produce calcium oxide (quicklime) and carbon dioxide gas. This is the first step in the limestone cycle. Quicklime (Calcium Oxide, CaO): A white, caustic, alkaline solid produced by the thermal decomposition of limestone during calcination. Quicklime is used in construction, steelmaking, and environmental applications. Slaking: The process in which quicklime (calcium oxide) reacts with water to produce calcium hydroxide (slaked lime). This reaction is highly exothermic. Slaked Lime (Calcium Hydroxide, Ca(OH)₂): A soft, white powder or putty formed by the reaction of quicklime with water. Slaked lime is used in making mortar, plaster, and for water treatment. Carbonation: The process by which calcium hydroxide reacts with carbon dioxide from the air to form calcium carbonate, completing the limestone cycle. This process is important in the setting of lime-based building materials. Thermal Decomposition: A chemical reaction in which a compound breaks down into simpler substances when heated. In the limestone cycle, it refers to the decomposition of calcium carbonate into calcium oxide and carbon dioxide. Soil Stabilization: The use of quicklime to improve the properties of soil, making it more stable and suitable for construction. Quicklime reacts with soil moisture, reducing water content and binding particles together. pH Neutralization: The process of adjusting the pH of acidic substances, such as soil or water, to a neutral level. Calcium hydroxide (slaked lime) is often used for this purpose. Hydration: The chemical reaction of a substance with water. In the limestone cycle, it refers to the reaction of quicklime with water to form slaked lime. Page 44 of 122 Chemistry Pack Level 10 Ms Tiziana Grech Assessment Checkpoint Assignment Mark Remarks Excellent Job Classwork 3.1: The Limestone Cycle Good Progress Better Effort Needs Improvement Excellent Job Homework 3.1: The Limestone Cycle – SEC Good Progress Past Paper Questions Better Effort Needs Improvement Excellent Job Good Progress Test 3: The Limestone Cycle Better Effort Needs Improvement Page 45 of 122 Chemistry Pack Level 10 Ms Tiziana Grech Topic 4: Water Hardness Water Hardness............................................................................................................... 47 Temporary Hardness......................................................................................................... 48 Permanent Hardness......................................................................................................... 49 Advantages and disadvantages of Water Hardness............................................................. 50 Soap Test......................................................................................................................... 52 Classwork 4.1 – Water Hardness................................................................................... 53 Checkpoint....................................................................................................................... 54 Assessment Checkpoint..................................................................................................... 56 Page 46 of 122 Chemistry Pack Level 10 Ms Tiziana Grech Water Hardness Water hardness is primarily caused by the presence of dissolved minerals, specifically calcium (Ca²⁺) and magnesium (Mg²⁺) ions. These ions enter the water when it flows over or through rocks, such as limestone (calcium carbonate) or dolomite (calcium magnesium carbonate), which dissolve and release these minerals into the water. Water hardness is classified into two types: temporary hardness and permanent hardness. Both types are due to the presence of dissolved minerals in the water, but they differ in their chemical composition and how they can be removed. Page 47 of 122 Chemistry Pack Level 10 Ms Tiziana Grech Temporary Hardness Cause: Temporary hardness is primarily caused by the presence of dissolved bicarbonate minerals, specifically calcium bicarbonate (Ca(HCO₃)₂) and magnesium bicarbonate (Mg(HCO₃)₂). Formation: As water flows over or through limestone (calcium carbonate, CaCO₃) or dolomite (a mix of calcium and magnesium carbonates), these rocks dissolve slightly in the water, particularly in the presence of carbon dioxide (CO₂), forming calcium and magnesium bicarbonates. CaCO3 + CO2 + H2O → Ca(HCO3)2 Removal: Temporary hardness can be removed by boiling the water. When water is boiled, the bicarbonates decompose into carbonates, releasing carbon dioxide and water, and forming insoluble calcium carbonate (CaCO₃) or magnesium carbonate (MgCO₃), which precipitates out. Ca(HCO3)2 →CaCO3+CO2+H2O Page 48 of 122 Chemistry Pack Level 10 Ms Tiziana Grech Permanent Hardness Cause: Permanent hardness is caused by the presence of calcium sulfate (CaSO₄), magnesium sulfate (MgSO₄), and sometimes chlorides of calcium and magnesium (CaCl₂ and MgCl₂). Unlike bicarbonates, these compounds do not break down upon heating. Formation: These salts can dissolve directly into water from gypsum (CaSO₄.2H₂O) or other sulfate- and chloride-containing minerals in the Earth's crust. When water passes over such minerals, calcium and magnesium ions dissolve, leading to hardness. CaSO4→Ca2+ +SO4-2 Removal: Permanent hardness cannot be removed by boiling. Instead, it requires chemical treatment methods, such as adding washing soda (sodium carbonate, Na₂CO₃) to precipitate out the calcium and magnesium ions as their respective carbonates or using ion-exchange resins that replace calcium and magnesium ions with sodium ions. Chemical equation for the ion-exchange resin: 2RNa + Ca2+ → R2Ca + 2Na+ Chemical equation for the washing soda: Ca2+ + Na2CO3 → CaCO3 + 2Na+ Page 49 of 122 Chemistry Pack Level 10 Ms Tiziana Grech Advantages and disadvantages of Water Hardness Advantages of Water Hardness: 1. Health Benefits: o Mineral Content: Hard water contains essential minerals like calcium and magnesium, which are beneficial for bone health, cardiovascular health, and overall wellness. 2. Taste: o Improved Flavour: Some people prefer the taste of hard water due to the presence of minerals, which can enhance the flavour of drinking water. 3. Reduced Metal Toxicity: o Protective Layer Formation: The minerals in hard water can form a protective scale on the inside of pipes, reducing the leaching of toxic metals like lead from old pipes. 4. Industrial Uses: o Certain Applications: In some industrial processes, hard water is preferred for tasks like brewing beer or making certain types of concrete, where the mineral content can be beneficial. Page 50 of 122 Chemistry Pack Level 10 Ms Tiziana Grech Disadvantages of Water Hardness: 1. Scale Formation: o Clogged Pipes and Appliances: Hard water causes the buildup of limescale in pipes, boilers, and household appliances like kettles and washing machines, leading to reduced efficiency, higher energy costs, and potentially expensive repairs or replacements. 2. Reduced Soap Efficiency: o Soap Scum: Hard water reacts with soap to form an insoluble substance known as soap scum, which reduces the cleaning effectiveness of soap and can leave residue on skin, hair, and surfaces. 3. Laundry Issues: o Stiff Fabrics: Hard water can make laundry feel stiff and rough because it prevents detergents from fully dissolving and interacting with fabrics. It can also cause colours to fade more quickly. 4. Increased Energy Costs: o Less Efficient Heating: The scale that forms inside water heaters and boilers acts as an insulator, making these appliances work harder to heat water, which increases energy consumption and costs. 5. Skin and Hair Problems: o Dryness and Irritation: Hard water can leave a residue on skin and hair, causing dryness, irritation, and a feeling of being unclean. It can also exacerbate conditions like eczema. 6. Difficult Cleaning: o More Cleaning Products Required: The formation of soap scum and scale means that more cleaning products and effort are needed to achieve the same level of cleanliness, leading to increased costs and time spent on cleaning tasks. Page 51 of 122 Chemistry Pack Level 10 Ms Tiziana Grech Soap Test Objective: Determine the presence of hardness by observing soap lathering. Procedure: 1. Fill a conical flask: Take a conical flask and fill it with a sample of the water to be tested. 2. Add Soap: Add a few drops of liquid soap or dishwashing detergent. 3. Shake: Shake the container vigorously for about 10 seconds. 4. Observe: If the water forms a lot of lather easily and quickly, it is likely soft. If the water produces little lather or forms a soapy scum, it is likely hard. Explanation: Hard water contains calcium and magnesium ions that react with soap, reducing its effectiveness and forming soap scum. Soft water allows soap to lather easily. Reaction with Calcium Ions (Ca²⁺): When soap (sodium stearate) reacts with calcium ions in hard water, the reaction produces calcium stearate, which is an insoluble soap scum. 2RCOONa + Ca2+ → (RCOO)2Ca + 2Na+ Where: RCOONa: Sodium stearate (a common soap component). Ca²⁺: Calcium ions from hard water. (RCOO)₂Ca: Calcium stearate, which precipitates as soap scum. Na⁺: Sodium ions remain in the solution. Page 52 of 122 Chemistry Pack Level 10 Ms Tiziana Grech Classwork 4.1 – Water Hardness May 2019 paper 1 Q. 7 Page 53 of 122 Chemistry Pack Level 10 Ms Tiziana Grech Checkpoint 1. Water Hardness: Definition: The amount of dissolved calcium (Ca²⁺) and magnesium (Mg²⁺) ions in water. Cause: Minerals like calcium carbonate (CaCO₃), magnesium sulfate (MgSO₄), and calcium sulfate (CaSO₄). 2. Temporary Hardness: Definition: Hardness that can be removed by boiling. Cause: Presence of dissolved bicarbonate minerals like calcium hydrogen carbonate [Ca(HCO₃)₂] and magnesium hydrogen carbonate [Mg(HCO₃)₂]. Removal: When heated, the calcium hydrogen carbonates, and magnesium hydrogen carbonate decompose into the metal carbonate and releasing carbon dioxide. 3. Permanent Hardness: Definition: Hardness that cannot be removed by boiling. Cause: Presence of calcium sulfate (CaSO₄) and magnesium sulfate (MgSO₄). Removal: Requires chemical treatment, such as the addition of washing soda (sodium carbonate) or ion-exchange methods. 4. Testing for Hardness Using Soap: Soap Test: Soap reacts with calcium and magnesium ions to form an insoluble scum. Procedure: Soap solution is added to the water sample; lather formation is observed. Less lather means harder water. 6. Advantages of Water Hardness: Taste: Hard water can taste better due to the minerals present. Health: It provides essential minerals like calcium and magnesium, which can be beneficial for bone health. Pipe Coating: Hard water forms a layer of scale inside pipes, which can protect against corrosion. Page 54 of 122 Chemistry Pack Level 10 Ms Tiziana Grech 7. Disadvantages of Water Hardness: Soap Wastage: More soap is needed to get a lather in hard water, leading to wastage. Scaling: Hard water causes scaling in kettles, boilers, and pipes, which can reduce efficiency and increase maintenance costs. Laundry Issues: Hard water can make clothes feel harsh or appear dull because the minerals reduce the effectiveness of detergents. Skin and Hair: Hard water can make skin feel dry and hair less smooth due to soap scum residue. 8. Softening Methods: Boiling: Effective for removing temporary hardness but not permanent hardness. Ion-exchange: Swaps calcium and magnesium ions with sodium or potassium ions using a water softener. Chemical Additives: Use of substances like washing soda (sodium carbonate) to precipitate calcium and magnesium. Page 55 of 122 Chemistry Pack Level 10 Ms Tiziana Grech Assessment Checkpoint Assignment Mark Remarks Excellent Job Classwork 4.1: Water Hardness Good Progress Better Effort Needs Improvement Excellent Job Good Progress Test 4: Water Hardness Better Effort Needs Improvement Page 56 of 122 Chemistry Pack Level 10 Ms Tiziana Grech Topic 5: The Chemistry of The Halogens Properties and Trends of Group 7 Elements........................................................................ 58 The Difference in Reactivity of the Halogens...................................................................... 59 Displacement Reactions.................................................................................................... 60 Classwork 5.1: Displacement reaction............................................................................ 62 Laboratory preparation of Chlorine Gas............................................................................. 63 Testing for Chlorine gas.................................................................................................... 65 Classwork 5.2: Preparation of Chlorine Gas.................................................................... 66 Uses of Chlorine in the Industry or Households.................................................................. 68 Dilute Hydrochloric acid.................................................................................................... 69 Homework 5.1: Chemistry of the Halogens...................................................................... 70 Checkpoint....................................................................................................................... 73 Assessment Checkpoint..................................................................................................... 74 Page 57 of 122 Chemistry Pack Level 10 Ms Tiziana Grech Properties and Trends of Group 7 Elements The halogens These are the elements found in Group 7, including fluorine, chlorine, bromine, iodine, and astatine. The halogens are diatomic. All halogens have seven electrons in their outer shell. They form halide ions by gaining one electron to complete their outer shells. Colour The halogens become darker as you go down the group. Chlorine is pale green, bromine is red-brown, and iodine is black (purple when in the gaseous state). Physical state of the halogens at room temperature At room temperature (20 °C), the physical state of the halogens are as follows: Chlorine is a gas. Bromine is a liquid Iodine is a solid. This means that the lower you go down the group the bigger the atomic radius is therefore the more intermolecular forces there are in between the diatomic molecules. Iodine has the highest melting point when compared to the other two halogens, this is because iodine has the biggest atomic radius with many intermolecular forces between the diatomic molecules. Page 58 of 122 Chemistry Pack Level 10 Ms Tiziana Grech The Difference in Reactivity of the Halogens Reactivity of Group 7 elements decreases ongoing down Group 7. Ongoing down Group 7, the number of electron - filled shells increases. This means that the outer shell (valance shell) is further away from the nucleus, and more electron shielding is present. This means that there are weaker electrostatic forces. These electrostatic forces of attraction are important since they are used to attract the needed electron from other atoms. For example, when chlorine reacts with sodium, the electrostatic forces of chlorine are very strong therefore the electron in the outer shell of sodium is pulled into the outer shell of chlorine. This means that the smaller the atomic radius the stronger the attractive forces, therefore the faster the element will react. Fluorine is the most reactive, while Astatine is the least reactive. Order of reactivity (most reactive first): Fluorine → Chlorine → Bromine → Iodine → Astatine Page 59 of 122 Chemistry Pack Level 10 Ms Tiziana Grech Displacement Reactions A halogen displacement reaction occurs when a more reactive halogen displaces a less reactive halogen from an aqueous solution of its halide. For example, chlorine is more reactive than bromine, therefore when chlorine gas is bubbled through a solution of potassium bromide (containing the bromide ions), the chlorine will displace the bromide ions. Chlorine becomes part of the solution as chloride ions, while the bromide ions are displaced into bromine liquid. The colour of the aqueous solutions of the halogens: Halogen displacement reactions 1. Chlorine reacting with bromide ion solution. a. When chlorine water (solution) is added to colourless potassium bromide solution, the solution becomes orange as bromine liquid forms, from the displacement reaction taking place. b. Chlorine is more reactive than bromine therefore it will displace the bromide ions from the solution. Chemical equation representing this reaction Chlorine + potassium bromide → potassium chloride + bromine Cl2 (l) + 2KBr (aq) → 2KCl (aq) + Br2(l) Page 60 of 122 Chemistry Pack Level 10 Ms Tiziana Grech 2. Bromine reacting with iodide ion solution. a. Bromine is more reactive than iodine therefore it will displace the iodide ions from the solution. b. Important: If excess bromine is added a black solid will be observed forming, this is iodine precipitating. Chemical equation representing this reaction Bromine + potassium iodide → potassium bromide + iodine Br2 (l) + 2KI (aq) → 2KBr (aq) + I2 (aq or s) 3. Iodine reacted with bromide ion solution. a. Iodine is less reactive than bromine and chlorine therefore it will not displace these halide ions from solution. Chemical equation representing this reaction Iodine + potassium bromide → no reaction I2 (l) + KBr (aq) → X The diagram below illustrates the results obtained from an experiment carried out to investigate displacement reactions in halogens. 1. No change means there was no reaction 2. The grey boxes mean no reaction was expected since the halogen is reacted with itself. 3. Where there is a change describe that means a displacement reaction took place. Page 61 of 122 Chemistry Pack Level 10 Ms Tiziana Grech Classwork 5.1: Displacement reaction. May 2021 paper 2A Q. 6 Page 62 of 122 Chemistry Pack Level 10 Ms Tiziana Grech Laboratory preparation of Chlorine Gas Chlorine gas can be prepared, dried, and collected in the laboratory using the following materials and set-up of apparatus. Materials: 1. Concentrated hydrochloric acid, HCl (l). 2. Manganese (IV) oxide powder, MnO2 (s). 3. Drying agent: concentrated sulfuric acid, H2SO4 (l). 4. Water: to remove gaseous hydrogen chloride gas (a by-product). Diagram: Setup of apparatus used in the laboratory preparation, drying, and collection of chlorine gas. Method: 1. Powdered manganese (IV) oxide was placed in the round bottomed flask, and the concentrated hydrochloric acid poured onto the it through the thistle funnel. 2. The mixture was heated, and any gases formed were first bubbled through water, (to remove any hydrogen chloride gas that might have evaporated). 3. The remaining mixture of gases was then bubbled through concentrated sulfuric acid, (to remove any water present, and to dry the chlorine gas). 4. The dry chlorine gas was then collected by the downward delivery method in a gas jar (since it is heavier than air). Page 63 of 122 Chemistry Pack Level 10 Ms Tiziana Grech Safety considerations: 1. The experiment was carried out in a fume cupboard since chlorine gas and hydrogen chloride gas are poisonous. 2. Lab coat, gloves and safety specs were worn to avoid any harm from the corrosive and toxic chemicals. 3. Hair was tied up to avoid it catching on fire while using the Bunsen burner. Precautions: 1. All apparatus must be washed three times with distilled water, to eliminate any impurities. 2. The tip of the delivery tube delivering the gases into the water, and concentrated sulfuric acid, must be submerged, to force the gases to come in contact with the water and concentrated sulfuric acid. 3. The tip of the thistle funnel must be submerged in the concentrated hydrochloric acid used or else any chlorine produced would escape through the thistle funnel. Important chemical equations 1. Manganese (IV) oxide reacted with concentrated hydrochloric acid to produce manganese (II) chloride, chlorine gas, and water. MnO2 (s) + 4HCl (l) → MnCl2 (aq) + Cl2 (g) + 2H2O (l) 2. Any hydrogen chloride gas that evaporated from the round bottomed flask while heating, is then bubbled through water, so it can dissolve in water. (Equation not important). HCl (g) → HCl (aq) 3. Any water vapour present with the chlorine gas is absorbed by the drying agent – concentrated sulfuric acid. (Equation not important). xH2O (g) + H2SO4 (l) → H2SO4. xH2O Page 64 of 122 Chemistry Pack Level 10 Ms Tiziana Grech Testing for Chlorine gas Chlorine gas has a characteristic yellowish – green colour and it can be tested – using damp blue litmus paper. Chlorine gas first dissolved in the water present on the damp blue litmus paper to produce an acid and a bleaching agent. The acid is hydrochloric acid, and the bleaching agent is hypochlorous acid. The acid will change the blue litmus paper to red, while the bleaching agent bleaches the litmus paper from red to white. Chemical equation Chlorine + water → hydrochloric acid + hypochlorous acid Cl2 (g) + H2O (l) → HCl (aq) + HClO (aq) Page 65 of 122 Chemistry Pack Level 10 Ms Tiziana Grech Classwork 5.2: Preparation of Chlorine Gas. May 2022 paper 2A Q. 14 Page 66 of 122 Chemistry Pack Level 10 Ms Tiziana Grech Page 67 of 122 Chemistry Pack Level 10 Ms Tiziana Grech Uses of Chlorine in the Industry or Households 1. Chlorine is used as a bleaching agent in the manufacture of paper and cloth. 2. Chlorine is used as a pesticide (insect killer). 3. Chlorine is used in the industrial manufacturing of rubber. 4. Chlorine is used in the production of various solvents. 5. Chlorine is used in the purification of drinking water, since it kills any microorganisms present in the water. Page 68 of 122 Chemistry Pack Level 10 Ms Tiziana Grech Dilute Hydrochloric acid Hydrochloric acid can be prepared by dissolving chlorine gas in ice cold water. Solubility of gases tend to increase as the temperature of the water decreases. Hydrochloric acid is a strong acid, since it dissociates completely into ions when in solution (remember that strong acids are considered strong because they can break apart and give a lot of hydrogen cations when in solution). The more H+ cations are present in solution the lower the pH, meaning the more acidic the solution becomes. Testing solubility and dissociation of hydrogen chloride gas in different solvents. To test the degree of solubility and dissociation of hydrogen chloride gas, one needs to bubble the gas in a beaker filled with distilled water (inorganic solvent), and another beaker filled with methylbenzene (organic solvent). Test both solutions with a blue litmus paper. The hydrogen chloride gas in water will change the blue litmus to red, indicating that the hydrogen chloride gas dissolved and dissociated in water – producing hydrochloric acid. The one using methylbenzene will not change the colour of the blue litmus, since although hydrogen chloride dissolves, it DOES NOT DISSOCIATE, therefore there are no H+ cations to change the pH. Page 69 of 122 Chemistry Pack Level 10 Ms Tiziana Grech Homework 5.1: Chemistry of the Halogens. May 2019 paper 2A no. 13 Page 70 of 122 Chemistry Pack Level 10 Ms Tiziana Grech Page 71 of 122 Chemistry Pack Level 10 Ms Tiziana Grech Page 72 of 122 Chemistry Pack Level 10 Ms Tiziana Grech Checkpoint 1. Halogens Definition: A group of non-metal elements found in Group 7 of the periodic table. They include fluorine (F), chlorine (Cl), bromine (Br), iodine (I), and astatine (At). Properties: Halogens are highly reactive, particularly with alkali metals and alkaline earth metals, forming salts. They exist as diatomic molecules (e.g., Cl₂) and become less reactive down the group. 2. Displacement Reactions of Halogens Definition: A type of reaction where a more reactive halogen displaces a less reactive halogen from its compound (usually from an aqueous solution of a halide salt). Example: Chlorine (Cl₂) can displace bromide ions (Br⁻) from potassium bromide (KBr) because chlorine is more reactive than bromine: Cl2+2KBr→2KCl+Br2 3. Preparation of Chlorine Gas in the Laboratory Method: Chlorine gas can be prepared by reacting concentrated hydrochloric acid (HCl) with manganese (IV) oxide (MnO₂) in a fume cupboard: 4HCl+MnO2→MnCl2+2H2O+Cl2 Safety Note: Chlorine gas is toxic and should be handled with care, in well-ventilated areas or under a fume hood. 4. Testing for Chlorine Gas Litmus Paper Test: Chlorine gas turns moist blue litmus paper red and then bleaches it white. Reason: Chlorine is an acidic gas and a strong bleaching agent, which causes the color change and bleaching effect. 5. Solubility of Hydrogen Chloride (HCl) Gas in Water HCl in Water: Hydrogen chloride gas is highly soluble in water, forming hydrochloric acid (HCl). 6. Solubility of Hydrogen Chloride (HCl) Gas in Methylbenzene HCl in Methylbenzene: Hydrogen chloride gas dissolves in methylbenzene but does not ionize, meaning it does not behave as an acid. Page 73 of 122 Chemistry Pack Level 10 Ms Tiziana Grech Assessment Checkpoint Assignment Mark Remarks Excellent Job Classwork 5.1: Displacement reaction. Good Progress Better Effort Needs Improvement Excellent Job Classwork 5.2: Preparation of Chlorine Gas. Good Progress Better Effort Needs Improvement Excellent Job Homework 5.1: Chemistry of the Halogens. Good Progress Better Effort Needs Improvement Test 5: Chemistry of Halogens. Excellent Job Good Progress Better Effort Needs Improvement Page 74 of 122 Chemistry Pack Level 10 Ms Tiziana Grech Topic 6: The Chemistry of Transition Metals General properties of the Transition Metals....................................................................... 76 Characteristic properties of Transition Metals................................................................... 77 Transition Metals and Their Different Valences.................................................................. 78 Reactions of Iron metal..................................................................................................... 79 The Rusting of Iron........................................................................................................... 81 Rust Prevention Methods.................................................................................................. 82 Barrier methods............................................................................................................ 82 Sacrificial Protection..................................................................................................... 83 Galvanising................................................................................................................... 83 Classwork 6.1: Rusting of Iron....................................................................................... 84 Homework 6.1: Rusting Process and Prevention.............................................................. 86 Reactions of Copper Metal................................................................................................ 87 Reactions of Copper (II) oxide with Acids.......................................................................... 88 Reaction of Copper (II) Oxide with Hydrogen..................................................................... 90 Classwork 6.2: Chemistry of Copper............................................................................... 91 Thermal Decomposition Reactions of Copper (II) Compounds............................................. 92 Classwork 6.3: Chemistry of Iron and Copper................................................................. 96 Checkpoint....................................................................................................................... 98 Assessment Checkpoint................................................................................................... 101 Page 75 of 122 Chemistry Pack Level 10 Ms Tiziana Grech General properties of the Transition Metals The following are general properties that most metallic elements have: 1. They are very hard and strong metals 2. Good conductors of heat and electricity 3. They have very high melting points 4. Highly dense metals Page 76 of 122 Chemistry Pack Level 10 Ms Tiziana Grech Characteristic properties of Transition Metals The following are unique characteristics (important) that only transition metals possess: 1. Transition metals form-coloured compounds. 2. Transition metals have more than one valence (also called oxidation state/ number). 3. Transition metals are often used as catalysts. General information (not for the exam) Although scandium and zinc are in the transition metal area of the Periodic table, they are not considered transition elements as they do not form coloured compounds and have only one oxidation state. Page 77 of 122 Chemistry Pack Level 10 Ms Tiziana Grech Transition Metals and Their Different Valences Why do transition metals have different valences (oxidation number / states)? The transition elements have more than one oxidation state, as they can lose a different number of electrons, depending on the chemical environment they are in. 1. Iron (Fe) for example can lose two electrons to form Fe2+ or three electrons to form Fe3+ depending on what it is reacting with. 2. Copper (Cu) for example can lose one electron to form Cu+, or two electrons to form Cu2+ depending on what it is reacting with. 3. When writing the name of a transition metal compound that has a valence of +2 one must write it in roman number form, for example, Iron (II) chloride, is iron with a valence of +2 bonded with chloride ions. Imagen below is illustrating sediment of iron (II) carbonate on natural stone. Page 78 of 122 Chemistry Pack Level 10 Ms Tiziana Grech Reactions of Iron metal Iron is a grey shiny metal which can lose 2 or 3 electrons when reacting, depending on what it is reacting with. There are 3 reactions of iron to be discussed in this topic: 1) Iron reacting with steam. 2) Iron reacting with hydrochloric acid. 3) Iron reacting with chlorine gas. Chemical reactions of iron 1) Iron does not react with water, but it reacts with steam to produce a BLACK mixed oxide iron (II, III,) oxide and hydrogen gas. Mixed oxide of iron – both iron (II) oxide, FeO, and iron (III) oxide, Fe2O3, are produced in this reaction but it is written as one compound – Fe3O4. Iron + steam → iron (II, III) oxide + hydrogen 3Fe (s) + 4H2O (g) → Fe3O4 (s) + 4H2 (g) Page 79 of 122 Chemistry Pack Level 10 Ms Tiziana Grech 2) Iron reacts with hydrogen chloride gas or hydrochloric acid, to produce, iron (II) chloride, FeCl2 and hydrogen gas, H2. Fe + 2HCl → FeCl2 + H2 3) Iron reacts with chlorine gas, to produce, iron (III) chloride, FeCl3. 2Fe (s) + 3Cl2 (g) → 2FeCl3 (s) Page 80 of 122 Chemistry Pack Level 10 Ms Tiziana Grech The Rusting of Iron Rust is a chemical reaction between iron, water and oxygen to form the compound hydrated iron (III) oxide, Fe2O3. H2O (rust). Oxygen and water must be present for the rusting process to occur. Iron + water + oxygen → hydrated iron (III) oxide 4Fe (s) + 3O2 (g) + 2H2O (l) → 2Fe2O3. H2O (s) Investigating rusting 1) To investigate the conditions required for rusting, prepare three test tubes as shown in the diagram 2) The oil in the 2nd tube keeps out air and the water has been boiled so that no air is left in it. 3) The calcium chloride in the 3rd tube is used to remove any moisture in the air. 4) After a few days, the iron nail in the 1st tube will be the only nail to show signs of rust Page 81 of 122 Chemistry Pack Level 10 Ms Tiziana Grech Rust Prevention Methods Barrier methods Rust can be prevented by coating iron with barriers that prevent the iron from coming into contact with water and oxygen. However, if the coatings are washed away or scratched, the iron is once again exposed to water and oxygen and will rust Common important barrier methods include: 1) Greasing the object. 2) Oiling the object – common in cars. 3) Painting the object. 4) Covering in plastic. Page 82 of 122 Chemistry Pack Level 10 Ms Tiziana Grech Sacrificial Protection A more reactive metal can be attached to less reactive metal. The more reactive metal will oxidise and therefore corrode first, protecting the less reactive metal from corrosion. For example, using zinc bars on the side of steel ships. Zinc is more reactive than iron therefore will react instead of iron. For continued protection, the zinc bars have to be replaced before they completely corrode. Galvanising Galvanising is a process where the iron to be protected is coated with a layer of zinc. This can be done by electroplating or dipping it into molten zinc. Zinc (II) carbonate ZnCO3 is formed when zinc reacts with oxygen and carbon dioxide in the air and protects the iron by the barrier method. If the coating is damaged or scratched, the iron is still protected from rusting by sacrificial protection. Page 83 of 122 Chemistry Pack Level 10 Ms Tiziana Grech Classwork 6.1: Rusting of Iron September 2016 paper 2B Q. 12 Page 84 of 122 Chemistry Pack Level 10 Ms Tiziana Grech Page 85 of 122 Chemistry Pack Level 10 Ms Tiziana Grech Homework 6.1: Rusting Process and Prevention September 2018 paper 1 Q. 9 Page 86 of 122 Chemistry Pack Level 10 Ms Tiziana Grech Reactions of Copper Metal Copper is an unreactive pink-orange shiny metal. Copper does not react with acids or water, but it reacts with oxygen when heated strongly. Copper is considered an unreactive metal. Copper + oxygen → coper (II) oxide 2Cu (s) + O2 (g) → 2CuO (s) Page 87 of 122 Chemistry Pack Level 10 Ms Tiziana Grech Reactions of Copper (II) oxide with Acids Copper (II) oxide, CuO(s), is a basic oxide (alkaline oxide) therefore it can react with acids to neutralise them and produce salt and water. 1) Copper (II) oxide, CuO, reacts with hydrochloric acid to produce, GREEN copper (II) chloride solution, CuCl2 (aq), and water, H2O (l). CuO (s) + 2HCl (aq) → CuCl2 (aq) + H2O (l) 2) Copper (II) oxide, CuO, reacts with sulfuric acid, H2SO4 (aq) to produce, PALE BLUE copper (II) sulfate solution, CuSO4 (aq), and water, H2O (l). CuO (s) + H2SO4 (aq) → CuSO4 (aq) + H2O (l) Page 88 of 122 Chemistry Pack Level 10 Ms Tiziana Grech 3) Copper (II) oxide, CuO, reacts with nitric acid, HNO3 (aq), to produce, BLUE copper (II) nitrate, Cu(NO3)2 (aq), and water, H2O (l). CuO (s) + HNO3 (aq) → Cu(NO3)2 (aq) + H2O (l) Page 89 of 122 Chemistry Pack Level 10 Ms Tiziana Grech Reaction of Copper (II) Oxide with Hydrogen Copper (II) oxide can be reduced (removal of oxygen) to copper metal by heating copper (II) oxide powder, CuO(s), with hydrogen gas (or any other reducing agent: methane, or ammonia). CuO(s) + H2 (g) → Cu(s) + H2O (g) 3CuO (s) + 2NH3 (g) → 3Cu (s) + N2 + 3H2O (g) 2CuO (s) + CH4 (g) → Cu (s) + CO2 (g) + 2H2O (g) Page 90 of 122 Chemistry Pack Level 10 Ms Tiziana Grech Classwork 6.2: Chemistry of Copper September 2019 paper 1 Q. 8 Page 91 of 122 Chemistry Pack Level 10 Ms Tiziana Grech Thermal Decomposition Reactions of Copper (II) Compounds 1) Copper (II) nitrate decomposes into copper (II) oxide, CuO(s), Nitrogen dioxide, NO 2(g) and Oxygen, O2(g), when heated strongly. 2Cu(NO3)2(s) → 2CuO(s) + 4NO2(g) + O2(g) Important observations: 1) Blue solid turns black. 2) Brown gas is produced. 3) Colourless droplets form at the neck of the boiling tube (water condensing). Page 92 of 122 Chemistry Pack Level 10 Ms Tiziana Grech 2) Copper (II) carbonate, CuCO3(s), decomposes into copper (II) oxide, CuO(s), and Carbon dioxide, CO2(g). CuCO3(s) → CuO(s) + CO2(g) Observations: 1) Green solid turns black. Page 93 of 122 Chemistry Pack Level 10 Ms Tiziana Grech May 2014 paper 2A Q. 12. Page 94 of