08 Empirical Crystallisation
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Calculations worksheet aimed at practicing empirical formulae calculations. The questions involve the formulae of hydrated salts.
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Transition metals such as manganese, iron, and copper can exist in multiple oxidation states. Manganese commonly exists as Mn2+, Mn4+, and Mn7+. Potassium manganate (VII), KMnO4, is a powerful oxidizing agent. Iron exists as Fe2+ and Fe3+ and acts as a catalyst in the Haber process. Copper exists as Cu+ and Cu2+. Compounds of Cu+ are generally colorless and insoluble while compounds of Cu2+ are blue and soluble, forming complexes with ligands like OH-, NH3, and CO32-.
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The document discusses different types of metal compounds including oxides, hydroxides, carbonates, nitrates, and chlorides. It describes methods of preparing these compounds such as direct combination of metals with oxygen or other reactants, or reactions of metal salts with bases or acids. The properties, reactions and uses of these compounds are also outlined. For example, metal oxides can be basic, acidic, or amphoteric and are used to form salts or in manufacturing. Hydroxides vary in solubility depending on the metal's reactivity and react with acids to form salts. Carbonates and nitrates similarly react with acids.
Ch4 z5e reactions
This document summarizes key concepts about aqueous solutions and types of reactions from Chapter 4. It defines terms like solvent, solute, solubility, electrolytes, and molarity. It describes how water is a polar solvent that can dissolve ionic compounds by hydrating ions. It explains types of solutions like strong/weak electrolytes and acids/bases. Precipitation, acid-base, and other reactions in solutions are discussed along with molecular, complete ionic, and net ionic equations. Solubility rules and stoichiometry of precipitation reactions are also covered.
4. Use your solubility table to predict whether the following compoun.pdf
4. Use your solubility table to predict whether the following compounds will be soluble in water
at 25°C: (5 points) CompoundSoluble/Insoluble CaSO4 K2SO4 Ba(OH)2 BeS CuCl2
Solubleinsolüble Compound Mg(NO3)2 CaCl2 Al2S3 (NH4)3PO4 SrCO3
Solution
First see the rules of solubility:1)Rule for phosphates says all phosphates are insoluble except for
the alkali metals and NH4+.
2) Rule for nitrates: All nitrates are soluble.
3) Rule for sulfides: All sulfides are insoluble except those of the alkali metals, alkali earths, and
NH4+.
4) Rule for Chlorides: All chlorides are soluble except AgCl, Hg2Cl2, and PbCl2
Mg(NO3) Soluble in water
CaCl2 Soluble in water
Al2S3: Aluminum sulfide actually reacts with water to form aluminum hydroxide and hydrogen
sulfide gas.Thus it decomposes in water instead of dissolving.
(NH4)3PO4: Soluble in water
SrCO3: not soluble in water. t dissolves in water
caso4 : yes Soluble in water
K2SO4 :yes Soluble in water
Ba(OH)2 : yes Soluble in water
BeS: yes Soluble in water
CuCl2: yes Soluble in water.
Chemical reactions and equations activity based question 10th
The document contains questions and answers related to chemical reactions and equations. Some key points:
- Hydrogen gas is evolved when zinc reacts with dilute sulfuric acid. Copper sulfate crystals change color from blue to white on heating due to loss of water of crystallization.
- When iron is added to copper sulfate solution, a displacement reaction occurs forming a brown coating of copper on the iron. Barium sulfate precipitate forms when sodium sulfate solution is added to barium chloride.
- Zinc hydroxide precipitate forms when sodium hydroxide is added to zinc sulfate solution. Lead nitrate decomposes on heating with a crackling sound, producing nitrogen dioxide, oxygen and lead oxide.
Acids, bases and salt
This document discusses acids, bases, and salts. It defines acids as substances that produce hydrogen ions in solution and turn litmus blue. Bases are defined as substances that produce hydroxide ions in solution and turn litmus red. Strong acids fully ionize in solution, while weak acids only partially ionize. Salts are formed from the reaction of acids and bases. Properties of acids, bases, and salts are discussed such as taste, effect on indicators, and solubility. Methods of preparing different types of salts are also outlined.
Hydrogen
This document discusses hydrogen, including its position in the periodic table, isotopes, methods of preparation, properties, and uses. Key points include:
1. Hydrogen is the lightest element with an atomic number of 1. It exists as diatomic molecules (H2) and has an anomalous position in the periodic table.
2. Methods of preparing hydrogen include the electrolysis of water and reactions of metals like zinc with acids. Hydrogen has uses as a fuel and in reducing reactions.
3. Hydrogen peroxide is another topic discussed, with methods of preparation like the anthraquinone process. It is a strong oxidizing agent with many uses in industry and medicine.
Determination of dissolved oxygen.pdf
1. The Winkler's method is used to determine the amount of dissolved oxygen (DO) in water samples. In this method, DO oxidizes manganese sulfate to form brown hydrated manganese dioxide precipitate in an alkaline medium.
2. The manganese dioxide then dissolves and oxidizes iodide ions to form iodine. The iodine is then titrated with a standardized sodium thiosulfate solution.
3. The volume of thiosulfate solution used corresponds to the amount of DO originally present in the water sample. Calculations are done to determine the concentration of DO in the sample water in units of mg/L or g/L.
Chapter 8-salts1
This document provides information about salts in chemistry. It defines a salt as an ionic substance formed by the replacement of hydrogen ions in an acid by metal ions or ammonium ions. Salts consist of cation and anion parts from the base and acid respectively. The document also includes tables and diagrams showing solubility rules for common salts and their reactions with heat. It describes methods for preparing and purifying soluble and insoluble salts, as well as qualitative analysis of salts through observation of physical properties and chemical tests.
IB Chemistry on Redox Titration, Biological Oxygen Demand and Redox.
This document discusses titration methods including acid-base titration and redox titration. It provides details on common primary standard acids and bases used in titration as well as indicators. It also discusses the principles and reactions involved in acid-base titration and redox titration. Examples are given of various redox titrations to determine concentrations of substances like copper, iron, chlorine, vitamin C, and more. Procedures and calculations for determining percentage compositions of substances from redox titrations are outlined.
Inorganic chem - chemical reactions
This document provides an overview of chemical equations and reactions. It discusses:
- Chemical equations, reactants, products, and how atoms rearrange during reactions.
- Balancing chemical equations by ensuring equal numbers of each atom on both sides.
- Information that can be obtained from a balanced chemical equation, such as moles of substances.
- Four main types of chemical reactions: combination, decomposition, displacement, and double displacement. Examples of each type are provided along with general reaction equations.
Types of chemical reactions - Laboratory Activity
These are synthesis, decomposition, combustion, single replacement and double replacement.
Laboratory Report.
University of Makati, Philippines.
III- BSE General Science
#Biochemistry #GeneralChemistry
To estimate the amount of Fe as Fe2O3 in the given solution of ferric chloride
1) The document provides instructions for estimating the amount of iron (Fe) in a solution as Fe2O3 through a gravimetric analysis.
2) Ferric ions in a solution containing barium chloride and hydrochloric acid are separated by precipitating barium sulfate, then iron is precipitated as ferric hydroxide and converted to ferric oxide.
3) The procedure involves filtering, washing, igniting, and weighing the ferric oxide precipitate to determine the amount of iron present originally.
#13 Key
This document provides sample problems and questions for a general chemistry exam. It includes sample balanced equations for neutralization reactions and definitions of key terms like oxidation, reduction, concentration, and indicators. It also asks students to determine oxidation states, identify redox reactions, calculate molarity and amounts of substances in solutions, and determine concentrations of ions after mixing solutions.
K2Cr2O7 & KMnO4
K2Cr2O7 has the following key properties:
1. It forms orange crystals that melt at 669 K and are moderately soluble in cold water but less soluble in hot water.
2. It is a powerful oxidizing agent that oxidizes substances like I-, Fe2+, H2S, and SO2 in the presence of dilute sulfuric acid.
3. Its solutions change color from orange to yellow upon addition of an alkali due to the equilibrium between dichromate and chromate ions, and the color changes back to orange with an acid.
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4. Use your solubility table to predict whether the following compoun.pdf
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Chemical reactions and equations activity based question 10th
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IB Chemistry on Redox Titration, Biological Oxygen Demand and Redox.
IB Chemistry on Redox Titration, Biological Oxygen Demand and Redox.
To estimate the amount of Fe as Fe2O3 in the given solution of ferric chloride
To estimate the amount of Fe as Fe2O3 in the given solution of ferric chloride
More from Anthony Hardwicke
13 Reacting Masses
This document discusses calculating reacting masses in chemical reactions from balanced chemical equations. It provides examples of calculating the mass of products formed given the mass of reactants or vice versa. The key steps are: 1) calculating moles of reactants, 2) using the mole ratio from the balanced equation to determine moles of products, 3) calculating mass of products from moles and molar mass. Several example calculations are shown and questions are provided for practice.
09 Percentage By Mass
Worksheet for working out the Percentage by Mass of various compounds. Pupils will need a data sheet or a list of relative atomic masses to be able to complete the questions.
06 Empirical Formulae
Worksheet for practicing working out empirical formulae. Pupils will need a data sheet or a list of relative atomic masses to be able to complete the questions.
04 Numbers of Atoms
Worksheet looking at working out the number of atoms in a given mass. Pupils will need a data sheet or a list of relative atomic masses to be able to complete the questions.
03 Moles Compounds
Worksheet for working out the number of moles of compounds. Pupils will need a data sheet or a list of relative atomic masses to be able to complete the questions.
01 Relative Molecular Mass
This document discusses relative molecular mass (RMM) and provides examples of calculating RMM for various molecules and compounds. It explains that RMM is calculated by adding up the relative atomic masses (RAM) of all the atoms in a molecule or compound. Several practice problems are worked out, including calculating RMM for covalent compounds like O2, H2S, and C6H12O6 as well as ionic compounds such as NaCl, MgO, and Fe2O3.3H2O. Key information about relative atomic masses and how they are typically rounded to whole numbers is also summarized.
Helicopters
This document provides instructions for conducting an experiment to investigate how the time it takes for a paper helicopter to fall is affected by changing variables. Students are asked to make a paper helicopter and then drop it, timing how long it takes to fall. They should record 3 trial times. Then they should cut 1cm off each wing and repeat the timing trials. This process should be repeated, cutting 1cm off the wings each time, until there is not enough wing left. The goal is to see how reducing the wing size affects the falling time.
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08 Empirical Crystallisation
- 1. 8. Empirical Formulae and Water of Crystallisation If you leave a saturated solution of copper(II) sulphate in a crystallising dish, some of the water evaporates and beautiful blue diamond-shaped crystals form. These hydrated crystals contain a fixed number of trapped water molecules. This is known as water of crystallisation.In the case of hydrated copper(II) sulphate, there are exactly five moles of water for every mole of copper(II) sulphate, so the formula is written as CuSO4⋅5H2O. A crystal that contains water of crystallisation is described as hydrated. If there is no water of crystallisation, the salt is anhydrous. CuSO4⋅5H2O(s) → CuSO4(s) + 5H2O(l) anhydrous copper(II) hydrated copper(II) sulphate (white sulphate (blue crystals) powder) Example 40g of hydrated copper(II) sulphate crystals were heated, driving off the water of crystallisation, leaving 25.6g of anhydrous copper(II) sulphate. Show that this is consistent with hydrated copper(II) sulphate having the formula CuSO4⋅5H2O. These headings don’t always have to be elements. Here we are working out the ratio of moles of two compounds; copper(II) sulphate and water. CuSO4 H2O mass 25.6g 14.4g 25.6 14.4 moles = 0.16 mol = 0.8 mol 160 18 ratio of moles 1 : 5 empirical formula CuSO4⋅5H2O 1. A student found that a sample of 30.5g to 26.0g. Work out the formula of hydrated cobalt(II) chloride contained the hydrated salt. 5.2g of anhydrous cobalt chloride and 4.32g of water. Calculate the formula of 4. Hydrated sodium carbonate, hydrated cobalt(II) chloride. Na2CO3⋅xH2O was found to contain 62.9% water by mass. Calculate the value of x. 2. 13.9g of hydrated iron(II) sulphate Answers crystals were heated, driving off the water 5. Potash alum has the 1. CoCl2⋅6H2O of crystallisation, leaving 7.6g of formula 2. FeSO4⋅7H2O anhydrous iron(II) sulphate. Calculate the K2SO4⋅Al2(SO4)3⋅yH2O. 3. BaCl2⋅2H2O 4. x = 10 formula of hydrated iron(II) sulphate. Calculate y, the number 5. y = 24 of moles of water of 3. A sample of hydrated barium chloride crystallisation, given that the hydrated was heated to drive off water of salt contains 45.57% water by mass. crystallisation. Its mass decreased from