When solutions of ionic compounds are mixed, precipitation reactions can occur if the mixing produces an insoluble ionic compound. For example, when solutions of silver nitrate and sodium chloride are mixed, a precipitate of insoluble silver chloride will form. Only the ions that react (Ag+ and Cl-) are shown in the balanced net ionic equation, with any spectator ions (Na+ and NO3-) that do not react omitted. The precipitate can be isolated by filtering the mixture.
This document discusses various concepts related to aqueous reactions and solution stoichiometry including:
- Ionic compounds dissociate into ions when dissolved in water, allowing them to conduct electricity.
- Strong electrolytes fully dissociate into ions while weak electrolytes only partially dissociate.
- Chemical equilibria exist for reactions that proceed in both the forward and backward directions.
- Precipitation reactions form insoluble products that drop out of solution.
- Acid-base reactions form water and salts by neutralizing hydrogen and hydroxide ions.
- Redox reactions involve the transfer of electrons and a change in oxidation states.
The document discusses several theories of acids and bases that developed over time:
- Lavoisier's original oxygen theory defined acids as containing oxygen (1776). This was disproven by Davy in 1810.
- Liebig proposed in 1838 that acids contain replaceable hydrogen.
- Arrhenius' 1884 definition defined acids as producing hydrogen ions (H+) and bases as producing hydroxide ions (OH-) in aqueous solutions, which became the standard definition.
- Lewis in 1923 expanded the definition to electron pair transfers between any acids and bases, not just involving hydrogen.
- Other theories such as Lux-Flood's oxygen theory of 1939 and Pearson's hard/soft acid base principle
This document discusses ionic equilibrium and electrolytes. It defines electrolytes as substances that conduct electricity when dissolved in water by dissociating into ions. Electrolytes are divided into strong and weak based on their degree of ionization. Strong electrolytes almost completely ionize while weak electrolytes ionize only partially. The document discusses Arrhenius theory of electrolytic dissociation and factors that affect the degree of ionization like concentration, temperature, and presence of a common ion. It also defines concepts like isohydric solutions and dissociation constants.
This document provides a summary of key concepts in oxidation-reduction (redox) reactions:
1) Redox reactions involve the transfer of electrons between chemical species, either through the complete transfer of electrons to form ionic bonds, or partial transfer to form covalent bonds. Oxidation is the loss of electrons and reduction is the gain of electrons.
2) Redox pairs are couples of oxidized and reduced forms of elements that differ in their oxidation state. Common redox pairs include Fe3+/Fe2+, O2/H2O, and MnO4-/MnO2.
3) The standard electrode potential (E°) indicates the tendency of half-reactions to occur.
This document summarizes elements and compounds that are essential or hazardous to the human body. It discusses:
1) Essential macroelements and microelements that are regularly found in the human body like carbon, hydrogen, oxygen, nitrogen, phosphorus, sulfur, calcium, magnesium, sodium, potassium, chloride, iron, zinc, and copper.
2) Potentially essential elements like cobalt, chromium, molybdenum, and manganese which are components of enzymes.
3) Hazardous substances like reactive oxygen species and very toxic inorganic compounds like white phosphorus, arsenic, cyanides, and fluorides.
4) Organic compounds can be classified by functional groups and the number of primary,
Chemistry & Physics UNIT 4 discusses key concepts in chemical bonding including:
1) Ionic bonding occurs through the transfer of electrons between atoms to form ions that are attracted via electrostatic forces. Covalent bonding involves the sharing of electron pairs between atoms.
2) Molecules are formed when two or more atoms combine via chemical bonds. Their formulas represent the elements present, with molecular formulas showing actual atom ratios.
3) Hydrogen bonding between water molecules gives rise to water's unique properties and ability to dissolve many substances. It also allows for the three-dimensional structures of proteins and DNA.
This document discusses various concepts related to aqueous reactions and solution stoichiometry including:
- Ionic compounds dissociate into ions when dissolved in water, allowing them to conduct electricity.
- Strong electrolytes fully dissociate into ions while weak electrolytes only partially dissociate.
- Chemical equilibria exist for reactions that proceed in both the forward and backward directions.
- Precipitation reactions form insoluble products that drop out of solution.
- Acid-base reactions form water and salts by neutralizing hydrogen and hydroxide ions.
- Redox reactions involve the transfer of electrons and a change in oxidation states.
The document discusses several theories of acids and bases that developed over time:
- Lavoisier's original oxygen theory defined acids as containing oxygen (1776). This was disproven by Davy in 1810.
- Liebig proposed in 1838 that acids contain replaceable hydrogen.
- Arrhenius' 1884 definition defined acids as producing hydrogen ions (H+) and bases as producing hydroxide ions (OH-) in aqueous solutions, which became the standard definition.
- Lewis in 1923 expanded the definition to electron pair transfers between any acids and bases, not just involving hydrogen.
- Other theories such as Lux-Flood's oxygen theory of 1939 and Pearson's hard/soft acid base principle
This document discusses ionic equilibrium and electrolytes. It defines electrolytes as substances that conduct electricity when dissolved in water by dissociating into ions. Electrolytes are divided into strong and weak based on their degree of ionization. Strong electrolytes almost completely ionize while weak electrolytes ionize only partially. The document discusses Arrhenius theory of electrolytic dissociation and factors that affect the degree of ionization like concentration, temperature, and presence of a common ion. It also defines concepts like isohydric solutions and dissociation constants.
This document provides a summary of key concepts in oxidation-reduction (redox) reactions:
1) Redox reactions involve the transfer of electrons between chemical species, either through the complete transfer of electrons to form ionic bonds, or partial transfer to form covalent bonds. Oxidation is the loss of electrons and reduction is the gain of electrons.
2) Redox pairs are couples of oxidized and reduced forms of elements that differ in their oxidation state. Common redox pairs include Fe3+/Fe2+, O2/H2O, and MnO4-/MnO2.
3) The standard electrode potential (E°) indicates the tendency of half-reactions to occur.
This document summarizes elements and compounds that are essential or hazardous to the human body. It discusses:
1) Essential macroelements and microelements that are regularly found in the human body like carbon, hydrogen, oxygen, nitrogen, phosphorus, sulfur, calcium, magnesium, sodium, potassium, chloride, iron, zinc, and copper.
2) Potentially essential elements like cobalt, chromium, molybdenum, and manganese which are components of enzymes.
3) Hazardous substances like reactive oxygen species and very toxic inorganic compounds like white phosphorus, arsenic, cyanides, and fluorides.
4) Organic compounds can be classified by functional groups and the number of primary,
Chemistry & Physics UNIT 4 discusses key concepts in chemical bonding including:
1) Ionic bonding occurs through the transfer of electrons between atoms to form ions that are attracted via electrostatic forces. Covalent bonding involves the sharing of electron pairs between atoms.
2) Molecules are formed when two or more atoms combine via chemical bonds. Their formulas represent the elements present, with molecular formulas showing actual atom ratios.
3) Hydrogen bonding between water molecules gives rise to water's unique properties and ability to dissolve many substances. It also allows for the three-dimensional structures of proteins and DNA.
The document discusses the language and terminology used in chemistry, including:
- Atoms, elements, compounds, and molecules and their representations using symbols and formulas.
- Ionic and covalent bonds between atoms and the structures they form.
- The periodic table and how it organizes elements and predicts their properties.
- Types of chemical reactions like synthesis, decomposition, and acid-base.
- Naming conventions for binary ionic compounds, acids, and other substances.
Concept of oxidation and reduction, redox reactions, oxidation number, balancing redox reactions, loss and gain of electrons, Balancing redox reactions, Half reaction method, Types of redox reaction- direct and indirect method, Electrochemical cell, Classification of redox reactions.
Chemical reactions and equations activity based question 10thBharathbabu68
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.
This document provides an overview of inorganic chemical nomenclature. It discusses the naming of elements, ions, binary compounds including hydrides, oxides, peroxides, halides, hydroxides, oxoacids and oxosalts. Different naming systems are covered, including compositional, stoichiometric, oxidation number, traditional and IUPAC nomenclature. Key concepts covered include oxidation numbers, types of ions, writing formulas for binary compounds, and distinguishing features of different compound classes.
The document discusses the nomenclature of inorganic compounds according to IUPAC rules. It covers naming conventions for binary ionic compounds such as metal oxides, salts, metal hydrides, non-metal halides, and binary acids. It also discusses determining oxidation states and writing chemical formulas. Key points include naming metal oxides as "stock name of metal cation" + "oxide", salts as "stock name of metal cation" + "root name of non-metal anion" + "-ide", and distinguishing between ionic and molecular compounds in nomenclature.
The document discusses various types of chemical reactions including oxidation-reduction, thermal dissociation, corrosion, and rancidity. It provides examples of oxidation-reduction reactions and explains that they involve the loss or gain of oxygen or hydrogen. Thermal dissociation reactions are described as reactions where compounds split into simpler substances when heated. Corrosion and rancidity are discussed as oxidation processes that damage iron structures and cause fats and oils to smell foul, respectively.
The document discusses redox reactions and assigning oxidation numbers. It provides examples of assigning oxidation numbers to elements in various compounds. It also justifies whether certain reactions are redox reactions by analyzing the changes in oxidation numbers of elements. Finally, it discusses the structures of some compounds and common ranges of oxidation numbers for some elements.
This document provides an overview of chemical bonding, including ions, ionic bonds, covalent bonds, and metallic bonds. It discusses the formation of cations and anions and how ions combine to form ionic compounds. It also covers covalent bonding, including how single, double and triple bonds are formed via orbital overlap. The document explains how to write formulas for ionic and covalent compounds using oxidation states to determine subscripts. Key properties of ionic and covalent compounds are also summarized.
Chemistry - Chp 11 - Chemical Reactions - PowerPointMr. Walajtys
This document summarizes key points from a chemistry textbook chapter on chemical reactions. It describes the five major types of chemical reactions: combination, decomposition, single replacement, double replacement, and combustion. It explains how to identify the type of reaction based on the reactants and how to write and balance chemical equations. It also discusses precipitation reactions in double replacement reactions in aqueous solution and the activity series for predicting single replacement reactions.
ICSE Class IX Chemistry The Language of Chemistry- TopperLearningAlok Singh
The document discusses the language of chemistry, including chemical symbols, elements, ions, valency, molecular formulas, and chemical equations. It provides details on:
1. How chemical symbols represent elements and are derived from their names.
2. The development of the modern periodic table and symbols assigned by IUPAC.
3. Definitions and examples of valency, ions, and molecular formulas.
4. How chemical equations are used to represent chemical reactions through balanced equations.
Class 10 chemical reactions and equationssarunkumar31
Types of reactions, Redox reactions, Reaction between acid and metal, Types of decomposition reaction, corrosion and rancidity.Acidic and basic nature of oxide, prevention method of corrosion.
The document discusses chemical functions and inorganic compounds. It introduces simple substances like elemental gases and halogens. It also covers compound substances and how elements combine in specific proportions based on their valence capacities. It provides rules for determining oxidation numbers and examples of oxidation numbers for various elements. Finally, it discusses topics like chemical symbols, formulas, nomenclature systems, and binary compounds like hydrides.
This document discusses reactions that occur in aqueous solutions. It explains that ionic compounds dissociate in water into their component ions due to water's polar nature. Conductivity is defined as a solution's ability to conduct electricity, with electrolytes able to conduct due to containing free ions. The three main types of reactions that occur in solutions are precipitation reactions, acid-base reactions, and redox reactions. Precipitation reactions form an insoluble precipitate when ions react, while acid-base reactions involve an acid and base reacting to form a salt and water. Redox reactions involve electron transfer between substances.
8 ppt v2_chemistry_10_regularity of changes in acid-base properties of compou...Galymzhan Tuleushov
The document summarizes the acid-base properties of period 3 oxides and their reactions. Sodium and magnesium oxides are basic and will neutralize acids to form salts and water. Aluminum oxide is amphoteric and can react with both acids and bases. Silicon, phosphorus, and sulfur oxides are acidic and will neutralize bases to form salts and water. There is a transition from basic to acidic character going across the period, reflecting the transition from metallic to non-metallic character.
Carbon exists in several allotropes with unique properties. Graphite has layered structures that allow for easy sliding of layers and is used as lubricant and pencil lead. Diamond has a tetrahedral structure and is the hardest material. Fullerenes like buckminsterfullerene have soccer ball shapes. Carbon also forms many inorganic compounds including carbon monoxide, carbon dioxide, carbonates, bicarbonates, carbides and cyanides that have various applications. Organic chemistry is the study of carbon compounds.
Carbon and its compounds ncert shashikumar b sghsykhalli
This document provides an overview of carbon and its compounds. It discusses carbon's position in the periodic table, its tetravalency, and ability to form catenated structures through covalent bonds. It describes the different types of isomers and classifications of hydrocarbons. Key carbon compounds like ethanol, ethanoic acid and their reactions are explained. Soaps and detergents are discussed in terms of their structures and cleansing mechanisms. Functional groups and their impact on naming conventions are also summarized.
This document summarizes information about carbon and its compounds. It discusses that carbon is a non-metal element that forms the basis of all living things. It exists in three allotropes - diamond, graphite, and buckminsterfullerene. It then describes the structures of diamond and graphite. The document further discusses that carbon can form many compounds due to its ability to form chains and bonds with four other atoms. It provides examples of organic compounds like hydrocarbons, alcohols, and carboxylic acids. In particular, it summarizes the types and properties of saturated and unsaturated hydrocarbons.
1) Water is a polar molecule due to the unequal sharing of electrons between the oxygen and hydrogen atoms. This gives water a partial negative charge on the oxygen side and partial positive charges on the hydrogen sides.
2) When ionic compounds dissolve in water, they dissociate into their constituent ions. The ions are then able to move freely within the water, allowing the solution to conduct electricity and making them electrolytes.
3) Double displacement precipitation reactions occur when two aqueous ionic solutions are mixed. The cations and anions are swapped between the reactants to form new ionic compounds. One of the products may have limited solubility and precipitate out of solution as an insoluble solid.
This document discusses how to write ionic equations for dissolution and precipitation reactions. It explains that dissolution reactions involve a solid salt breaking into its constituent ions in aqueous solution. Precipitation reactions involve mixing ions that form an insoluble solid product. The document provides examples of writing balanced equations for calcium chloride dissolving, potassium nitrate dissolving with nitrate staying together as a polyatomic ion, and silver and chloride ions forming an insoluble silver chloride precipitate.
The document discusses the language and terminology used in chemistry, including:
- Atoms, elements, compounds, and molecules and their representations using symbols and formulas.
- Ionic and covalent bonds between atoms and the structures they form.
- The periodic table and how it organizes elements and predicts their properties.
- Types of chemical reactions like synthesis, decomposition, and acid-base.
- Naming conventions for binary ionic compounds, acids, and other substances.
Concept of oxidation and reduction, redox reactions, oxidation number, balancing redox reactions, loss and gain of electrons, Balancing redox reactions, Half reaction method, Types of redox reaction- direct and indirect method, Electrochemical cell, Classification of redox reactions.
Chemical reactions and equations activity based question 10thBharathbabu68
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.
This document provides an overview of inorganic chemical nomenclature. It discusses the naming of elements, ions, binary compounds including hydrides, oxides, peroxides, halides, hydroxides, oxoacids and oxosalts. Different naming systems are covered, including compositional, stoichiometric, oxidation number, traditional and IUPAC nomenclature. Key concepts covered include oxidation numbers, types of ions, writing formulas for binary compounds, and distinguishing features of different compound classes.
The document discusses the nomenclature of inorganic compounds according to IUPAC rules. It covers naming conventions for binary ionic compounds such as metal oxides, salts, metal hydrides, non-metal halides, and binary acids. It also discusses determining oxidation states and writing chemical formulas. Key points include naming metal oxides as "stock name of metal cation" + "oxide", salts as "stock name of metal cation" + "root name of non-metal anion" + "-ide", and distinguishing between ionic and molecular compounds in nomenclature.
The document discusses various types of chemical reactions including oxidation-reduction, thermal dissociation, corrosion, and rancidity. It provides examples of oxidation-reduction reactions and explains that they involve the loss or gain of oxygen or hydrogen. Thermal dissociation reactions are described as reactions where compounds split into simpler substances when heated. Corrosion and rancidity are discussed as oxidation processes that damage iron structures and cause fats and oils to smell foul, respectively.
The document discusses redox reactions and assigning oxidation numbers. It provides examples of assigning oxidation numbers to elements in various compounds. It also justifies whether certain reactions are redox reactions by analyzing the changes in oxidation numbers of elements. Finally, it discusses the structures of some compounds and common ranges of oxidation numbers for some elements.
This document provides an overview of chemical bonding, including ions, ionic bonds, covalent bonds, and metallic bonds. It discusses the formation of cations and anions and how ions combine to form ionic compounds. It also covers covalent bonding, including how single, double and triple bonds are formed via orbital overlap. The document explains how to write formulas for ionic and covalent compounds using oxidation states to determine subscripts. Key properties of ionic and covalent compounds are also summarized.
Chemistry - Chp 11 - Chemical Reactions - PowerPointMr. Walajtys
This document summarizes key points from a chemistry textbook chapter on chemical reactions. It describes the five major types of chemical reactions: combination, decomposition, single replacement, double replacement, and combustion. It explains how to identify the type of reaction based on the reactants and how to write and balance chemical equations. It also discusses precipitation reactions in double replacement reactions in aqueous solution and the activity series for predicting single replacement reactions.
ICSE Class IX Chemistry The Language of Chemistry- TopperLearningAlok Singh
The document discusses the language of chemistry, including chemical symbols, elements, ions, valency, molecular formulas, and chemical equations. It provides details on:
1. How chemical symbols represent elements and are derived from their names.
2. The development of the modern periodic table and symbols assigned by IUPAC.
3. Definitions and examples of valency, ions, and molecular formulas.
4. How chemical equations are used to represent chemical reactions through balanced equations.
Class 10 chemical reactions and equationssarunkumar31
Types of reactions, Redox reactions, Reaction between acid and metal, Types of decomposition reaction, corrosion and rancidity.Acidic and basic nature of oxide, prevention method of corrosion.
The document discusses chemical functions and inorganic compounds. It introduces simple substances like elemental gases and halogens. It also covers compound substances and how elements combine in specific proportions based on their valence capacities. It provides rules for determining oxidation numbers and examples of oxidation numbers for various elements. Finally, it discusses topics like chemical symbols, formulas, nomenclature systems, and binary compounds like hydrides.
This document discusses reactions that occur in aqueous solutions. It explains that ionic compounds dissociate in water into their component ions due to water's polar nature. Conductivity is defined as a solution's ability to conduct electricity, with electrolytes able to conduct due to containing free ions. The three main types of reactions that occur in solutions are precipitation reactions, acid-base reactions, and redox reactions. Precipitation reactions form an insoluble precipitate when ions react, while acid-base reactions involve an acid and base reacting to form a salt and water. Redox reactions involve electron transfer between substances.
8 ppt v2_chemistry_10_regularity of changes in acid-base properties of compou...Galymzhan Tuleushov
The document summarizes the acid-base properties of period 3 oxides and their reactions. Sodium and magnesium oxides are basic and will neutralize acids to form salts and water. Aluminum oxide is amphoteric and can react with both acids and bases. Silicon, phosphorus, and sulfur oxides are acidic and will neutralize bases to form salts and water. There is a transition from basic to acidic character going across the period, reflecting the transition from metallic to non-metallic character.
Carbon exists in several allotropes with unique properties. Graphite has layered structures that allow for easy sliding of layers and is used as lubricant and pencil lead. Diamond has a tetrahedral structure and is the hardest material. Fullerenes like buckminsterfullerene have soccer ball shapes. Carbon also forms many inorganic compounds including carbon monoxide, carbon dioxide, carbonates, bicarbonates, carbides and cyanides that have various applications. Organic chemistry is the study of carbon compounds.
Carbon and its compounds ncert shashikumar b sghsykhalli
This document provides an overview of carbon and its compounds. It discusses carbon's position in the periodic table, its tetravalency, and ability to form catenated structures through covalent bonds. It describes the different types of isomers and classifications of hydrocarbons. Key carbon compounds like ethanol, ethanoic acid and their reactions are explained. Soaps and detergents are discussed in terms of their structures and cleansing mechanisms. Functional groups and their impact on naming conventions are also summarized.
This document summarizes information about carbon and its compounds. It discusses that carbon is a non-metal element that forms the basis of all living things. It exists in three allotropes - diamond, graphite, and buckminsterfullerene. It then describes the structures of diamond and graphite. The document further discusses that carbon can form many compounds due to its ability to form chains and bonds with four other atoms. It provides examples of organic compounds like hydrocarbons, alcohols, and carboxylic acids. In particular, it summarizes the types and properties of saturated and unsaturated hydrocarbons.
1) Water is a polar molecule due to the unequal sharing of electrons between the oxygen and hydrogen atoms. This gives water a partial negative charge on the oxygen side and partial positive charges on the hydrogen sides.
2) When ionic compounds dissolve in water, they dissociate into their constituent ions. The ions are then able to move freely within the water, allowing the solution to conduct electricity and making them electrolytes.
3) Double displacement precipitation reactions occur when two aqueous ionic solutions are mixed. The cations and anions are swapped between the reactants to form new ionic compounds. One of the products may have limited solubility and precipitate out of solution as an insoluble solid.
This document discusses how to write ionic equations for dissolution and precipitation reactions. It explains that dissolution reactions involve a solid salt breaking into its constituent ions in aqueous solution. Precipitation reactions involve mixing ions that form an insoluble solid product. The document provides examples of writing balanced equations for calcium chloride dissolving, potassium nitrate dissolving with nitrate staying together as a polyatomic ion, and silver and chloride ions forming an insoluble silver chloride precipitate.
Investigation Of The Thermal Decomposition Of Copper...Alexis Naranjo
This molecular dynamics simulation examines the indentation response of an aluminum-amorphous silicon core-shell nanostructure. The study investigates the deformation behavior of the amorphous silicon shell and aluminum core under spherical indentation. It also explores how the density of the amorphous silicon, indenter radius size, and core/shell ratio size affect the structural deformation of the nanostructure. The simulation aims to provide insights into optimizing the properties of core-shell nanostructures for applications.
This document discusses key concepts related to solutions and acid-base chemistry, including:
- Definitions of solutions, solvents, solutes, electrolytes, and nonelectrolytes. Common examples are provided.
- Dissociation and hydration of ionic compounds in solution. Equations are given to represent these processes.
- Precipitation reactions and how to write balanced molecular, ionic, and net ionic equations.
- Arrhenius definitions of acids and bases. Examples of strong and weak acids and bases ionizing in water are given through chemical equations.
- Key reactions include neutralization and acid-base reactions where examples of writing balanced equations are provided.
This document provides information on various chemistry concepts related to solutions and reactions in aqueous solutions. It defines key terms like electrolytes, nonelectrolytes, dissociation, and precipitation reactions. It also discusses acid-base reactions and neutralization reactions. Oxidation-reduction reactions and displacement reactions are introduced. Molarity is defined as a way to quantify concentration in solutions.
This document provides information on chemical reactions involving aqueous solutions and precipitation. It discusses:
- How ionic compounds dissolve in water to form ions, and how this allows them to conduct electricity. Electrolytes fully or partially dissociate, while non-electrolytes do not.
- When solutions of two ionic compounds are mixed, double displacement occurs as ions exchange to form new ionic compounds. If one product is insoluble, it precipitates out as a solid.
- The process of predicting products of precipitation reactions using solubility rules to determine which compound will precipitate.
- How to write balanced molecular, complete ionic, and net ionic equations for precipitation reactions. Examples
The document provides an overview of chemistry concepts related to life, including atomic structure, molecules, compounds, chemical bonds, and organic chemistry. It discusses the building blocks of organic molecules like carbohydrates, lipids, proteins, and nucleic acids. Key points include that atoms bond to form molecules through ionic or covalent bonds, and that compounds have fixed ratios of elements that take on new properties.
Answer A] Water solventates any substances that .pdfanilart346
Answer: A] Water solventates any substances that may dissociate into charged ions.
Taking salt for example (NaCl), the water will allow the solid salt to dissociate into Na+ and Cl-
ions. How does it do this chemically? Water molecules are polar. Hence, they surround the ions
in something called a hydration shell. Water is H2O, the hydrogens have a positive dipole
whereas the oxygen has a negative dipole. The negative sides of the water molecule (oxygen)
will all surround and face the positive charged Na+ ions. The positive sides of the water
molecule (hydrogens) will surround the negatively charged Cl- ions. B] A hydrogen bond is
formed between an hydrogen atom on one molecule and an electronegative atom on another
molecule (oxygen, flourine, nitrogen, etc.). The hydrogen atom on the first molecule must be
covalently binded to an electronegative atom as well. Overall hydrogen bond: O-H...O (dotted
line is the hydrogen bond). In water, these bonds are rapidly formed and destroyed. It is not a
covalent bond, but sort of an interaction force between water molecules. Physically, this may be
seen with ice. Ice forms crystal-like structures due to the neat organization of these hydrogen
interactions. C] Water will tend to isolate hydrophobic substances. Water is polar and hence
likes to mix with other polar substances. Since hydrophobic substances are non-polar, the water
will physically separate out the hydrophobic substances in solution. In chemical terms, the water
will just form clatherates (organized shells of water) around any hydrophobic substance. This can
be seen if you add oil to water. It just doesn\'t mix.
Solution
Answer: A] Water solventates any substances that may dissociate into charged ions.
Taking salt for example (NaCl), the water will allow the solid salt to dissociate into Na+ and Cl-
ions. How does it do this chemically? Water molecules are polar. Hence, they surround the ions
in something called a hydration shell. Water is H2O, the hydrogens have a positive dipole
whereas the oxygen has a negative dipole. The negative sides of the water molecule (oxygen)
will all surround and face the positive charged Na+ ions. The positive sides of the water
molecule (hydrogens) will surround the negatively charged Cl- ions. B] A hydrogen bond is
formed between an hydrogen atom on one molecule and an electronegative atom on another
molecule (oxygen, flourine, nitrogen, etc.). The hydrogen atom on the first molecule must be
covalently binded to an electronegative atom as well. Overall hydrogen bond: O-H...O (dotted
line is the hydrogen bond). In water, these bonds are rapidly formed and destroyed. It is not a
covalent bond, but sort of an interaction force between water molecules. Physically, this may be
seen with ice. Ice forms crystal-like structures due to the neat organization of these hydrogen
interactions. C] Water will tend to isolate hydrophobic substances. Water is polar and hence
likes to mix with other polar substances. Since hydr.
This document provides information on types of chemical reactions and solution stoichiometry. It discusses the common water molecule and how water acts as a solvent. It describes different types of solutions, including strong and weak electrolytes, and nonelectrolytes. Precipitation, acid-base, and oxidation-reduction reactions are introduced. Methods for writing balanced and net ionic equations are presented. Finally, techniques for performing stoichiometric calculations on reactions in solution, including determining limiting reactants and amounts of products formed, are covered.
This document summarizes key concepts in solution chemistry and stoichiometry, including:
1) Solutions, electrolytes, dissociation, and precipitation reactions are discussed. Strong and weak electrolytes are defined.
2) Acid-base reactions such as neutralization and gas-forming reactions are covered. Oxidation-reduction reactions and displacement reactions are also summarized.
3) Concepts including molarity, dilution, and titration reactions are introduced for quantitative chemical calculations.
This document summarizes key concepts in solution chemistry and stoichiometry, including:
1) Solutions, electrolytes, dissociation, and precipitation reactions are discussed. Strong and weak electrolytes are defined.
2) Acid-base reactions such as neutralization and gas-forming reactions are covered. Oxidation-reduction reactions and oxidation numbers are also introduced.
3) Concepts like molarity, dilution, and titration are explained as methods to quantify concentrations in solutions and chemical reactions.
This document discusses various concepts related to ionic equilibrium in solution including strong and weak electrolytes, acid-base theories of Arrhenius, Bronsted-Lowry, and Lewis. It defines strong electrolytes as completely dissociating in water and weak electrolytes as achieving an equilibrium between dissociated and undissociated molecules. Acids are defined as proton donors and bases as proton acceptors under the Bronsted-Lowry theory. The Lewis theory further defines acids as electron pair acceptors and bases as electron pair donors. Dissociation constants and factors affecting acid strength are also covered.
The document discusses ligand field theory, which examines how ligands affect the energies of d orbitals in metal complexes. It explains that for an octahedral complex:
1) Ligands raise the dx2-y2 and dz2 orbitals significantly due to strong interactions, forming an "eg" set of orbitals.
2) Ligands interact weakly with the dxy, dxz, and dyz orbitals, forming a lower "t2g" set of orbitals.
3) The splitting of orbitals into eg and t2g sets can lead to high-spin or low-spin electron configurations, influencing magnetic and chemical properties.
This document provides an overview of electrolytes, acids, and bases according to different theories. It defines strong and weak electrolytes, and lists examples of each. It then summarizes the Arrhenius, Bronsted-Lowry, and Lewis theories of acids and bases. The Arrhenius theory defines acids as substances that produce hydrogen ions in water and bases as those that produce hydroxide ions. The Bronsted-Lowry theory broadened this to include proton donors and acceptors. The Lewis theory further expanded the definition to any electron pair acceptor or donor.
The document discusses the theory of coordination compounds. It defines coordination compounds as products of Lewis acid-base reactions where ligands bond to a central metal atom via coordinate covalent bonds. It describes coordination bonds and terminology like ligands, central metal ion, oxidation state, and coordination number. The document also discusses Werner's theory of coordination complexes, valence bond theory, types of ligands, nomenclature of coordination compounds, and limitations of Werner's theory.
1. Ions are formed when atoms gain or lose electrons to form electrically charged particles. Ionic compounds have high melting points and conduct electricity when molten or dissolved in water.
2. Transition metal hydroxides are insoluble in water and form precipitates when a soluble transition metal compound is mixed with sodium hydroxide solution.
3. Spectroscopy studies the patterns of light emitted from heated samples to identify elements and discover new elements like rubidium and caesium.
Why does salt dissolve in waterThe positive charges of sodium (Na.pdflibowskymcinnisell69
Why does salt dissolve in water?
The positive charges of sodium (Na) and the negative charges of chloride (Cl) interact with water
molecules
The NaCl molecules form covalent bonds with the surrounding water molecules
The water molecules are attracted to the nonpolar bonds of the sodium and chloride ions
The ionic bond between the sodium ion and chloride ion is strong
None of the above
Why does salt dissolve in water?A.
The positive charges of sodium (Na) and the negative charges of chloride (Cl) interact with
water moleculesB.
The NaCl molecules form covalent bonds with the surrounding water moleculesC.
The water molecules are attracted to the nonpolar bonds of the sodium and chloride ionsD.
The ionic bond between the sodium ion and chloride ion is strongE.
None of the above
Solution
Answer from the above options is A PART
When we try to dissolve an ionic compound by stirring it in water, the positive poles of the water
molecules are attracted to the anions, while the negative poles of other water molecules are
attracted to the cations, so the polar water molecules \"pull\" the ions out of a crystal. As a result
of these interactions, the ionic bonds eventually break and ions are released into the water. When
the salt is dissolved, every ion is surrounded by water molecules, creating a kind of concentric
shells of polar water molecules centered around the ions. (They also look like petals of a daisy).
This process is called \"hydration.\"
The hydration of its ions causes a salt crystal to break apart (dissolve) in the water. Therefore,
when table salt (NaCl) is dissolved, two hydrated ions appear in the water: a positively cation
Na+ and a negative charged anion Cl-. Each is surrounded with a shell of closely attracted water
molecules that prevent the ions from reconciling into a crystal again.
These watery shells prevent the ions from getting together again into a solid. Hydrated ions
jostles around the solution, dispersing evenly in the solution and causing the particles to
disappear from view as they are spreading within the solution. The hydrated ions are so small
that we cannot see them even with a strong microscope. and therefore we see the process as a
disappearance of particles from view as the dissolving progresses.
Different ions can attract different number of water molecules into the \"watery shell\" because
they can have different strength (density) of their charge. The density of the charge depends on
the ratio of charge to surface area of an ion; the larger the ratio, the larger the hydration number
will be..
This document discusses different types of chemical reactions including decomposition, synthesis, combustion, double replacement, and single replacement reactions. It provides examples of each type of reaction and explains the key features that define them. Double replacement reactions are highlighted, where a metal replaces a metal in a compound and a nonmetal replaces a nonmetal, forming a precipitate if one of the products is insoluble. Guidelines are provided for writing molecular, total ionic, and net ionic equations for double replacement reactions.
The document discusses the history and modern understanding of the periodic table. It covers how elements are arranged based on proton number and how this explains trends in properties within groups. Specific groups like alkali metals, halogens, and transition metals are examined in terms of their structures, properties, and reactions. Common acid-base reactions and quantitative chemical calculations are also summarized.
This document summarizes key concepts about aqueous solutions, including:
- Solutions contain a solvent and one or more solutes
- Ionic compounds generally dissolve well in water due to ion separation through solvation
- "Like dissolves like" - polar substances dissolve in water and nonpolar substances do not
- Electrolytes conduct electricity in solution while nonelectrolytes do not
- Hydrates contain water within their crystalline structure.
Deserts are expanding across large parts of the world due to various human and environmental factors. Overgrazing of livestock and deforestation have degraded soils, reducing their ability to retain water. Climate change has exacerbated droughts in many regions, reducing rainfall and causing desertification. As deserts expand, they threaten the livelihoods of those living in increasingly arid lands and could contribute to conflicts if populations are forced to migrate. Understanding and addressing the root causes, such as sustainable land management practices and reducing greenhouse gas emissions, is essential to slow the spread of deserts.
This friendship is rare and special. The friends have laughed together, cried together, and helped each other through difficult times. They share secrets without shame and tell each other the truth, even when at fault. The thought of the friendship ending makes one sad, but they take comfort in their memories. Though distance may separate them, their close bond of friendship remains in their hearts.
This friendship is rare and special. The friends have laughed together, cried together, and helped each other through difficult times. They share secrets without shame and tell each other the truth, even when at fault. The thought of the friendship ending makes one sad, but they take comfort in their memories. Though distance may separate them, their close bond of friendship remains in their hearts.
This document contains the same URL, www.communication4all.co.uk, repeated six times without any other text or context. The URL www.communication4all.co.uk is mentioned six consecutive times in the document.
This document contains the same URL, www.communication4all.co.uk, repeated six times without any other text or context. The URL www.communication4all.co.uk is mentioned six consecutive times in the document.
This document is a 13-page exam for the International General Certificate of Secondary Education in Biology. It contains 7 multiple choice questions testing knowledge of topics including food webs, human nutrition, plant and animal physiology, environmental issues, and genetics. The exam is designed to be completed in 1 hour and 15 minutes by writing answers directly on the question paper, with no additional materials allowed.
This document provides instructions for candidates taking the International General Certificate of Secondary Education Biology exam. It specifies that candidates should write their identification information on all work, use blue or black pen or pencil for diagrams, not use staples or correction fluid, answer all questions, and fasten all work together at the end. The exam consists of 13 printed pages, 3 blank pages, and 6 questions testing knowledge of biology topics including tissues, classification of organisms, response and control systems in humans and plants, hormones and disease, and population ecology.
This document is a 16-page exam for the International General Certificate of Secondary Education in Chemistry. It contains 7 multi-part chemistry questions testing knowledge of topics including the periodic table, atmospheric pollutants, acid-base reactions, extraction of metals, combustion reactions, organic compounds, and properties of group 1 and transition metals. It also includes a copy of the periodic table.
This document consists of a chemistry exam paper containing multiple choice and short answer questions testing knowledge of chemical apparatus, reactions, and experimental procedures. The questions cover topics such as identifying experimental set ups, describing chemical reactions and tests, planning investigations, analyzing results, and drawing conclusions from experiments.
This document consists of a 14 page chemistry exam with multiple choice and free response questions covering topics like the halogens, redox reactions, organic chemistry, acid-base reactions, and metals/alloys. It includes diagrams of lab setups and reagents/products. The exam provides space for students to show their work and includes a periodic table reference.
This document consists of three paragraphs summarizing the content of a 16-page biology exam. The exam contains multiple choice and short answer questions about fungi reproduction, acid rain, and sensitivity. It provides context about classifying fungi species, the effects of acid rain, and defining voluntary vs involuntary actions. Tables and figures are referenced to support analyzing trends in sulfur dioxide and sulfur concentrations over time.
This document consists of instructions and questions for a biology exam. It contains 12 pages, with the first 9 pages consisting of exam questions and the last 3 pages being blank. The exam has two sections - Section A contains short answer questions and Section B requires longer answers to two out of three essay questions. The questions cover topics in biology such as human reproduction, cell structure, genetics, the water cycle, osmosis, nutrition, and plant growth.
This document is a 19-page exam for the International General Certificate of Secondary Education (IGCSE) Biology exam. It contains 10 multiple choice and short answer questions covering topics like the respiratory system, photosynthesis, nutrient cycling, food webs, cell structure, genetics, and inherited conditions. Students are instructed to answer all questions directly on the exam paper and work is to be completed in 1 hour and 15 minutes.
1. This document consists of 17 printed pages and 3 blank pages for the Cambridge International Level 1/2 Certificate in Physics.
2. The document is approved for use in England, Wales and Northern Ireland and contains a multiple choice exam from October/November 2014 with 40 questions on topics related to physics.
3. Students are instructed to choose the correct answer for each question and record their choice on an answer sheet within 45 minutes, with each correct answer scoring one mark.
This document consists of a 16 page multiple choice exam for physics. It contains 40 multiple choice questions testing various concepts in physics such as motion, forces, energy, electricity, waves, and nuclear physics. The questions are accompanied by diagrams, graphs, and short paragraphs of information as context for the questions.
This document is a multiple choice exam for physics that contains 40 questions. It covers topics like mechanics, energy, waves, electricity, atomic physics, and more. The questions require analyzing diagrams, graphs, and scenarios to choose the best answer from four options (A, B, C, or D).
This document consists of a 20-page exam for the International General Certificate of Secondary Education in Physics. The exam contains 40 multiple choice questions testing various concepts in physics, including kinematics, forces, energy, waves, electricity, and radioactivity. The questions are presented over two columns on each page with answer choices A, B, C, or D provided.
This document is a 20 page exam for the International General Certificate of Secondary Education Physics exam. It contains 40 multiple choice questions testing various concepts in physics. The questions cover topics such as measurement, motion, forces, energy, electricity, waves, and radioactivity. Students have 45 minutes to complete the exam.
1. This document consists of 18 printed pages and 2 blank pages for a Cambridge International Level 1/Level 2 Certificate in Physics.
2. The document contains a multiple choice exam with 40 questions on topics related to physics. For each question there are four possible answers (A, B, C, D) and students must choose the one they consider correct.
3. The exam covers various concepts in physics including measurement, motion, energy, electricity, magnetism, waves, and radioactivity. Diagrams and tables are provided with some questions.
हिंदी वर्णमाला पीपीटी, hindi alphabet PPT presentation, hindi varnamala PPT, Hindi Varnamala pdf, हिंदी स्वर, हिंदी व्यंजन, sikhiye hindi varnmala, dr. mulla adam ali, hindi language and literature, hindi alphabet with drawing, hindi alphabet pdf, hindi varnamala for childrens, hindi language, hindi varnamala practice for kids, https://www.drmullaadamali.com
How to Build a Module in Odoo 17 Using the Scaffold MethodCeline George
Odoo provides an option for creating a module by using a single line command. By using this command the user can make a whole structure of a module. It is very easy for a beginner to make a module. There is no need to make each file manually. This slide will show how to create a module using the scaffold method.
LAND USE LAND COVER AND NDVI OF MIRZAPUR DISTRICT, UPRAHUL
This Dissertation explores the particular circumstances of Mirzapur, a region located in the
core of India. Mirzapur, with its varied terrains and abundant biodiversity, offers an optimal
environment for investigating the changes in vegetation cover dynamics. Our study utilizes
advanced technologies such as GIS (Geographic Information Systems) and Remote sensing to
analyze the transformations that have taken place over the course of a decade.
The complex relationship between human activities and the environment has been the focus
of extensive research and worry. As the global community grapples with swift urbanization,
population expansion, and economic progress, the effects on natural ecosystems are becoming
more evident. A crucial element of this impact is the alteration of vegetation cover, which plays a
significant role in maintaining the ecological equilibrium of our planet.Land serves as the foundation for all human activities and provides the necessary materials for
these activities. As the most crucial natural resource, its utilization by humans results in different
'Land uses,' which are determined by both human activities and the physical characteristics of the
land.
The utilization of land is impacted by human needs and environmental factors. In countries
like India, rapid population growth and the emphasis on extensive resource exploitation can lead
to significant land degradation, adversely affecting the region's land cover.
Therefore, human intervention has significantly influenced land use patterns over many
centuries, evolving its structure over time and space. In the present era, these changes have
accelerated due to factors such as agriculture and urbanization. Information regarding land use and
cover is essential for various planning and management tasks related to the Earth's surface,
providing crucial environmental data for scientific, resource management, policy purposes, and
diverse human activities.
Accurate understanding of land use and cover is imperative for the development planning
of any area. Consequently, a wide range of professionals, including earth system scientists, land
and water managers, and urban planners, are interested in obtaining data on land use and cover
changes, conversion trends, and other related patterns. The spatial dimensions of land use and
cover support policymakers and scientists in making well-informed decisions, as alterations in
these patterns indicate shifts in economic and social conditions. Monitoring such changes with the
help of Advanced technologies like Remote Sensing and Geographic Information Systems is
crucial for coordinated efforts across different administrative levels. Advanced technologies like
Remote Sensing and Geographic Information Systems
9
Changes in vegetation cover refer to variations in the distribution, composition, and overall
structure of plant communities across different temporal and spatial scales. These changes can
occur natural.
A workshop hosted by the South African Journal of Science aimed at postgraduate students and early career researchers with little or no experience in writing and publishing journal articles.
How to Fix the Import Error in the Odoo 17Celine George
An import error occurs when a program fails to import a module or library, disrupting its execution. In languages like Python, this issue arises when the specified module cannot be found or accessed, hindering the program's functionality. Resolving import errors is crucial for maintaining smooth software operation and uninterrupted development processes.
A review of the growth of the Israel Genealogy Research Association Database Collection for the last 12 months. Our collection is now passed the 3 million mark and still growing. See which archives have contributed the most. See the different types of records we have, and which years have had records added. You can also see what we have for the future.
Pengantar Penggunaan Flutter - Dart programming language1.pptx
11 bctopic6
1. TOPIC 6.
CHEMICAL REACTIONS AND IONIC EQUATIONS.
Reactions involving ionic compounds.
As discussed earlier, ionically bonded compounds consist of large aggregations of
cations and anions which pack together in crystal lattices in such a way that the
electrostatic attractions between oppositely charged ions are maximised and
repulsions between like charged ions minimised. When an ionic crystal is placed
in water, in many cases the solid dissolves, releasing the component ions into the
SOLVENT to form a SOLUTION. Such compounds are said to be SOLUBLE
and the substance that dissolves is called the SOLUTE. A well known example of
a soluble ionic compound is table salt or sodium chloride. The process of
dissolving can be best represented by an equation which is slightly different from
the formula equations used in Topic 5. Instead, an IONIC EQUATION is used to
show the ions released into the solution as follows.
NaCl(s)
v
Na+ + Cl–
The same rules apply as for formula equations in that all species shown on the left
must also be present on the right hand side of the equation. In addition, notice that
the electrical charge present on both sides of the equation also balances.
Thus the equation for another ionic solid, barium chloride, dissolving in water
would be as follows
BaCl2(s)
v
Ba2+ + Cl– + Cl–
which is usually written as
BaCl2(s)
v
Ba2+ + 2Cl–
Because the Cl– ions are separate individual species, they are represented as 2Cl–
and not as Cl22– , which would mean two Cl atoms bonded together and bearing a 2
negative charge, (Cl!Cl)2–.
When ions are released into water solution, they all experience attractions to water
molecules which form spheres around them, as illustrated for NaCl below.
VI - 1
2. VI - 2
The reason for this attractive force between water molecules and ions is the ability
of the oxygen atoms in water molecules to attract the electrons in their O–H bonds
to a greater extent than do the hydrogen atoms. The O atom is said to be more
ELECTRONEGATIVE than the H atom. This results in a slight negative charge
on the oxygen atom and a slight positive charge on each hydrogen atom. The O–H
bond is an example of a POLAR BOND and the water molecule, being angular in
shape, has a non-symmetric distribution of charge and is a POLAR MOLECULE.
In the dissolution of ionic solids such as sodium chloride, the oxygen atoms of
water molecules are attracted to the positive charge on cations (Na+ in this
example) while its hydrogen atoms are attracted to the anions (Cl– in this example).
This results in considerable energy being released and the subsequent breaking up
of the crystal lattice of the ionic solid to form individual ions surrounded by the
attracted water molecules.
Consequently, ions in water solution are said to be AQUATED and sometimes the
suffix (aq) is used to emphasise this point. Thus the two equations above might also
be written as
NaCl(s) v Na+(aq) + Cl–(aq)
BaCl2(s)
v
Ba2+(aq) + 2Cl–(aq)
Initially the (aq) suffix will be used here, but later it will be assumed that all ions in
water solution are aquated and the (aq) suffix will be omitted. Some other
examples of ionic equations for ionic solids dissolving follow.
K2CO 3(s)
v
2K+(aq) + CO 32–(aq)
(NH4)3PO4(s)
v
3NH 4+(aq) + PO 43–(aq)
3. VI - 3
Many text books use the (aq) symbolism (incorrectly) to indicate a solution of a
substance in water by attaching (aq) to the formula of that solid. Clearly, there
can be no such species as an aquated ionic substance because if it has dissolved,
it is totally present as ions. Therefore equations showing species such as
NaCl(aq) are most misleading and should be ignored.
Not all ionic solids will dissolve in water to a significant extent. For example, the
amount of the ionic solid silver chloride, AgCl, which will dissolve in water is so
small that it is classed as insoluble. Insoluble ionic compounds of common metals
include three chlorides, about five sulfates, most carbonates, most phosphates and
most sulfides.
How does one know if a given ionic compound is soluble?
A table classifying the solubilities of ionic compounds in water is given at the end
of this Topic. This table lists the ionic compounds of the common cations as
soluble, insoluble or slightly soluble. This information probably will be needed to
be committed to memory at a later stage of your course, but for the present you
should consult it if needed for the ensuing exercises. The contents of the table are
more easily remembered in terms of the general situation for each anion. For
example: all nitrates are soluble; most hydroxides are insoluble except for those of
the first family members Na and K, plus Ba; most carbonates are insoluble except
those of the first family members Na and K. It is also useful to remember that all
of the compounds of the elements Na and K from the first family are soluble.
Check your understanding of this section:
Why do ionic compounds when soluble, dissolve best in water?
Write an ionic equation for the dissolution of iron(III) chloride in water.
Account for the water molecule being polar.
What is the source of the energy required to break up an ionic crystal when it
dissolves?
How can a soluble ionic compound be obtained back from a solution?
The aquated ions released into a solution when an ionic compound dissolves and its
ionic bonds broken are free to move about throughout the solution. Ionic bonding
is only present in the solid state. However, if the solution is boiled, the water is
driven off as a gas (VOLATILE) but the ions remain in the solution (NONVOLATILE). Volatile substances have sufficiently weak forces of attraction
between their constituent entities to allow them to escape to the vapour phase when
enough heat energy is supplied. In non-volatile substances, the attractive forces
operating are much stronger and much higher temperatures would be required to
vaporise them. When sufficient water has been evaporated, the cations and anions
are deprived of their surrounding water molecules and they can then recombine to
form the ionically bonded solid again. In this way, any solution of an ionic
compound can be evaporated sufficiently for crystals to form. The equation for the
evaporation process would simply be the reverse of the equation for the dissolution,
for example
Na+(aq) + Cl– (aq)
v
NaCl(s)
To emphasise that the solution is being evaporated, it may be helpful to write
"evap" or similar over the arrow.
4. VI - 4
Precipitation reactions.
Silver chloride is seen from the solubility table to be insoluble in water. Silver
nitrate and sodium chloride are both soluble compounds and in water-solution
would release their component ions as shown in the following ionic equations.
AgNO3(s)
v
Ag+(aq) + NO 3–(aq)
NaCl(s)
v
Na+(aq) + Cl–(aq)
If a solution of silver nitrate were mixed with a solution of sodium chloride, then
the Ag+ ions would react with the Cl– ions to form a PRECIPITATE of the
insoluble salt silver chloride, AgCl, as shown in the following ionic equation.
Ag+ (aq) + NO3– (aq) + Na+ (aq) + Cl– (aq)
v
AgCl(s) + NO 3–(aq) + Na+(aq)
Note that only the Ag+ and the Cl– ions have reacted, leaving the Na+ and NO 3– ions
free in the solution. As these last two ions have not in fact entered into a reaction,
they can be deleted from the equation in much the same way as common terms are
cancelled from both sides of a mathematical equation. Such ions are called
SPECTATOR IONS. Initially it may be helpful to write down all the ions which
are being mixed together in order to establish whether any combination can form
an insoluble salt, and then cancel out the spectator species. However, with practice
you will be able to delete this step and write the final equation in one step. For this
reaction it would be
Ag+(aq) + Cl–(aq) v AgCl(s)
Being a solid, the silver chloride could be obtained by filtering the mixture. The
precipitate would be retained in the filter paper and the solution containing the
spectator ions (called the FILTRATE) would pass through the filter paper. It
would contain only the ions Na+ and NO 3–, provided exactly equal amounts of the
Ag+ and Cl– ions were in the solutions which were mixed originally. By
evaporating the water, crystals of sodium nitrate could be isolated from the filtrate.
Na+ (aq) +
NO 3– (aq)
evap
v
NaNO 3(s)
From a knowledge of the solubilities of ionic compounds, one can predict whether
any combination of solutions of soluble compounds will lead to a precipitation
reaction. The following examples show how to establish whether a precipitation
reaction occurs when solutions are mixed.
Example 1.
If a water solution of potassium chloride were added to a water solution of
copper(II) nitrate, would any reaction occur and if so, what would be the product?
To answer this, consider all possible combinations of the four ions which are to be
mixed - K+, Cl– , Cu2+ and NO3– . The formulas of the possible products are KNO 3
and CuCl2. By consulting the solubility table given, it is seen that both of these
compounds are soluble so there would be no reaction. If the solution formed by
mixing this combination were to be evaporated, the resulting solid would be a
mixture of all four possible compounds, KNO 3, CuCl2, Cu(NO 3) 2 and KCl.
5. VI - 5
Example 2.
Water solutions of lead(II) nitrate and sodium sulfate are mixed. What, if any,
reaction occurs?
The possible combinations of ions could produce the compounds of formulas
PbSO4 and NaNO3. Of these two, lead(II) sulfate is insoluble and sodium nitrate is
soluble. Therefore, a precipitate of PbSO 4 would form according to the equation
Pb2+ (aq) + SO 42–(aq) v PbSO 4(s)
In this example, only one of the possible products is insoluble (PbSO 4) but if both
possible products were insoluble, then a mixture of both compounds would form.
Example 3.
Water solutions of sodium carbonate and barium chloride are mixed. What, if any,
reaction occurs?
The possible combinations of ions could produce the compounds NaCl and BaCO 3.
Of these two, barium carbonate is insoluble while sodium chloride is soluble.
Therefore, a precipitate of BaCO 3 would form according to the equation
Ba2+ (aq) + CO 32–(aq) v BaCO 3(s)
Example 4.
Water solutions of iron(II) sulfate and potassium phosphate are mixed. What, if
any, reaction occurs?
The possible combinations of ions could produce the compounds K 2SO 4 and
Fe3(PO 4)2. Of these two, iron(II) phosphate is insoluble while potassium sulfate is
soluble. Therefore, a precipitate of Fe 3(PO 4) 2 would form according to the equation
3Fe2+ (aq) + 2PO 43–(aq) v Fe 3(PO 4) 2(s)
Check your understanding of this section:
Given that the formation of an ionic solid from its component ions involves the
release of energy, why is it necessary to supply heat to a solution of sodium
chloride in order to reclaim the solid from solution?
What advantages are there in using an ionic equation to represent the
precipitation of silver chloride from a mixture of sodium chloride and silver
nitrate solutions?
In the previous reaction, which are the spectator ions?
Reactions of acids.
Apart from the cations from metals and the polyatomic cation NH 4+, another cation
frequently encountered in reactions is the hydrogen ion, H+, which is also
sometimes called the oxonium ion or the hydronium ion. Recall from Topic 2 that
the H atom consists of just one proton for the nucleus, surrounded by a single
orbiting electron. If this electron were removed, the H+ ion would be formed and it
would be just a free proton. A proton is extremely small and the charge density on
6. VI - 6
it would therefore be very concentrated. Consequently, the H+ ion does not have an
independent existence in solution. Instead it associates with water molecules by
joining on to them using one of the lone pairs of electrons on the O atom, and is
more correctly represented as H+ (aq), or frequently as H 3O+ . These are all equally
acceptable ways of representing the hydrogen ion. Hydrogen ions are supplied in
water solution by compounds called ACIDS. The three most common acids
encountered are:
nitric acid, HNO 3:
contains H+(aq) and NO 3–(aq) ions
sulfuric acid, H 2SO 4:
which contains H+(aq), HSO 4–(aq) and SO 42–(aq) ions
hydrochloric acid:
a water solution of the gas hydrogen chloride, HCl(g),
which ionizes to H+(aq) and Cl–(aq) ions in the water.
Acids such as these provide H+ ions as well as the anion from the acid in solution.
Both the H+ ions and the anions can participate in reactions. Thus addition of
hydrochloric acid to silver nitrate solution produces silver chloride precipitate just
like the previous example when solutions of silver nitrate and sodium chloride were
mixed. Remaining in solution would be the spectator ions, H+ and NO 3–, which is a
solution of nitric acid. The ionic equation for the reaction is
H+(aq) + Cl–(aq) + Ag+(aq) + NO3–(aq) v AgCl(s) + H+(aq) + NO 3–(aq)
When the spectator ions are deleted, it reduces to
Cl– (aq) + Ag+(aq)
v
AgCl(s)
which is identical to the ionic equation given previously for the reaction between
silver nitrate solution and sodium chloride solution.
From this example, it can be seen that any soluble silver salt and any soluble
chloride or hydrochloric acid, would react when mixed together in a water solution
to precipitate silver chloride. One advantage of writing ionic equations is that only
the actual species reacting are shown.
Reactions involving the H + ion.
In the precipitation reactions involving acids discussed above, it is the anion from
the acid which is reacting and the H+ is the spectator ion. However, there are many
reactions of acids where the important ion is H+ while the anion from the acid is
merely a spectator ion.
All acids supply hydrogen ions in solution. This ion can participate in a number of
common reaction types, recognition of which makes the writing of chemical
equations much simpler. These reaction types are:
1. Acids with reactive metals.
Many metals react with acids to form the cation of the metal and produce hydrogen
gas, H2, from the H+ ions. These metals are sometimes called "reactive metals", as
distinct from the "coinage metals" such as silver, gold, copper, and platinum which
are inert to most acids. In particular, hydrochloric acid and sulfuric acid behave in
this way with reactive metals, and in the process, will leave the chloride or sulfate
ions in the solution. Evaporation of the water from the resulting solution will
produce the ionic compound of the metal and the acid used. Ionic compounds are
7. VI - 7
frequently called "SALTS" which is a general term and, as pointed out earlier, is
not restricted to common table salt, NaCl.
For example, hydrochloric acid added to magnesium metal produces hydrogen
which is evolved as a gas and leaves in solution magnesium ions, Mg2+ and
chloride ions, Cl– . Evaporation of this solution results in isolation of the salt,
magnesium chloride.
Mg + 2HCl v MgCl 2 + H 2
Formula equation:
Ionic equation: Mg(s) + 2H+ (aq) + 2Cl– (aq) v Mg2+ (aq) + H 2(g) + 2Cl–(aq)
or, deleting the Cl– spectator ions,
Mg(s) + 2H+(aq) v Mg2+(aq) + H 2(g)
an acid + a reactive metal forms a salt and hydrogen gas
2. Acids with oxides of metals.
All ionic oxide compounds, regardless of whether they are water soluble, react with
acids to release the metal ion from the oxide into the solution and also to form
water.
As an example, consider the equation for the reaction of hydrochloric acid with
calcium oxide.
CaO + 2HCl v CaCl 2 + H 2O
Formula equation:
Ionic equation:
CaO(s) + 2H+(aq) + 2Cl–(aq) v Ca2+(aq) + 2Cl–(aq) + H 2O(l)
If the water were evaporated from the solution, the Ca2+ and Cl– ions would
crystallise to form the solid compound, calcium chloride.
an acid + an oxide forms a salt and water.
Note that in the equation above, the chloride ion is a spectator ion. This ion should
be omitted, to give for the correct equation
CaO(s) + 2H+(aq)
v
Ca2+(aq) + H 2O(l)
3. Acids with hydroxides of metals.
In the same way as the H+ ions of acids react with metal oxides to form a salt and
water, so too can acids react with hydroxide compounds of metals, regardless of
whether they are water soluble, to also form a salt and water. As in the case of
oxides, if the salt formed is soluble in water, then the resulting solution must be
evaporated in order to isolate the solid compound.
For example, nitric acid reacts with copper(II) hydroxide to form water and
copper(II) nitrate in solution.
8. VI - 8
Formula equation:
Ionic:
Cu(OH) 2 + 2HNO 3 v Cu(NO 3) 2 + 2H 2O
Cu(OH)2(s) + 2H+ (aq) + 2NO 3– (aq) v 2H 2O(l) + Cu2+ (aq) + 2NO 3–(aq)
Again, the NO3– (aq) ion is a spectator ion and should be deleted from both sides of
the equation, as the H+(aq) is the only component of the acid which reacts with the
Cu(OH)2(s).
Cu(OH)2(s) + 2H+ (aq)
v
Cu2+(aq) + 2H 2O(l)
Evaporation of this solution would produce the salt, Cu(NO 3) 2, according to the
equation
evap
Cu2+ (aq) + 2NO 3–(aq)
v Cu(NO 3) 2(s)
If the hydroxide reacting with the acid is soluble in water and a solution of that
hydroxide is specified, then the cation of the hydroxide is also a spectator ion and
should be deleted along with the anion from the acid. For example, hydrochloric
acid reacting with a solution of sodium hydroxide would be represented by the
ionic equation
H+(aq) + OH–(aq) v
H 2O(l)
The salt NaCl(s) could be obtained by evaporating the solution.
an acid + a hydroxide forms a salt and water
or, combining this with the previous section,
an acid + an oxide or hydroxide forms a salt and water.
4. Acids with carbonates.
Another common reaction of an acid leads to the evolution of a gas when the acid
is mixed with the metal salts of certain anions. This type of reaction is typified by
the well-known case of acids reacting with carbonates to produce carbon dioxide
gas, water and a salt. Again, the reaction occurs regardless of whether the
compound reacting with the acid is soluble or insoluble in water. For example,
sulfuric acid reacts with solid magnesium carbonate to produce carbon dioxide gas
and magnesium sulfate in solution as follows:
Formula equation:
Ionic:
MgCO 3 + H 2SO 4 v MgSO 4 + CO 2 + H 2O
MgCO3(s) + 2H+(aq) v Mg2+(aq) + CO 2(g) + H 2O(l)
Note that the spectator ion, SO42– (aq) has been deleted from both sides of the ionic
equation.
If the carbonate compound used were soluble and was already in solution, then the
ionic equation would not include the metal ion because it would be a spectator ion,
9. VI - 9
as in the following example of adding nitric acid to a solution of sodium carbonate:
2Na+(aq) + CO 32–(aq) + 2H+(aq) + 2NO 3–(aq)
v
2Na+(aq) + 2NO 3–(aq) + H 2O(l) + CO 2(g)
which simply becomes, after deleting spectator ions,
CO 32– (aq) + 2H+(aq)
v
CO 2(g) + H 2O(l)
an acid + a carbonate forms carbon dioxide gas, a salt and water
Check your understanding of this section:
By what criterion would you decide if a species is an acid?
What is the expected reaction when an acid is mixed with a reactive metal?
What happens when an acid is placed with an insoluble oxide or hydroxide?
When an acid is placed on a compound such as calcium carbonate, what is
observed?
Objectives of this Topic.
When you have completed this Topic, including the tutorial questions, you should
have achieved the following goals:
1.
Know the meaning of the terms solvent; solute; solution; aquated ions;
volatile; non-volatile; spectator ions; precipitate; filtrate; electronegative
atom; polar bond; polar molecule; non-polar bond; non-polar molecule.
2.
Understand the concept of polarity in covalent molecules, especially as it
relates to the role of water as a solvent for ionic compounds.
3.
Be able to write ionic equations for any dissolution process or any
crystallisation of a salt.
4.
Know how to find out if any given ionic compound is soluble.
5.
Be able to write ionic equations for any precipitation of a salt.
6.
By use of solubility tables, be able to predict whether any given combination
of solutions of salts will lead to formation of a precipitate.
7.
Recognise acids as a source of hydrogen ions.
8.
Know the names and formulas for several common acids.
9.
Recognise acids as being able to precipitate insoluble salts by supplying the
relevant anion.
10. VI - 10
10.
Know the reactions of acids with (i) reactive metals; (ii) oxides and
hydroxides; (iii) carbonates.
11.
Be able to write ionic equations for each of the above reaction types.
SOLUBILITY TABLE OF SOME COMMON SALTS
BEHAVIOUR WITH COMMON CATIONS
ANION
USUAL
EXCEPTIONAL
SLIGHTLY
SOLUBLE
F–
soluble
except Mg2+, Ca2+, Sr2+, Ba2+, Mn2+, Pb2+
Al3+
Cl–
soluble
except Ag+, Hg 22+
Pb2+
Br–
soluble
except Ag+, Hg 22+
Hg2+, Pb2+
I–
soluble
except Ag+, Hg 22+, Hg2+, Pb2+
SO42–
soluble
except Sr2+, Ba2+, Hg 22+, Pb2+
NO3–
soluble
CH3CO2–
soluble
OH–
SO32–
insoluble except Na+, K +, Ba2+
except Na+, K+, NH 4+, Mg2+
insoluble
PO43–
insoluble except Na+, K+
Ca2+, Sr2+
insoluble except Na+, K+, NH 4+
C2O42–
Ag+, Hg22+
insoluble except Na+, K+, NH 4+
CO 32–
Ca2+, Ag+
S2–
NH 4+
insoluble except Na+, K+, NH 4+, Mg2+, Ca2+, Sr2+, Ba2+
The ammonium ion, NH 4+, is not an element but is a polyatomic cation which also
forms salts like any normal metal cation.
11. VI - 11
SUMMARY
Ionic compounds consist of very large aggregates of cations and anions packed into
a crystal lattice so as to maximise the attractive forces between cations and anions
and minimise the repulsive cation/cation and anion/anion forces. Despite the
strength of ionic bonds, many ionic compounds can dissolve in solvents such as
water to form a solution, a process best represented by an ionic equation. Ionic
equations show the species involved in the actual form in which they are present, so
ions in solution are shown as such. Ionic equations require not only that the mass
should balance on both sides but also the charge. Otherwise, ionic equations follow
the same rules as for formula equations.
The reason why ionic compounds may dissolve in water is that water is a polar
molecule in which the oxygen atoms attract more than their share of the bonding
electrons (said to be more electronegative) than the hydrogen atoms, leaving a slight
negative charge on the O atoms and a slight positive charge on the H atoms. This
uneven charge distribution allows water molecules to be attracted to both cations
and anions in the ionic crystal, releasing much energy in the process. If enough
energy is released through these attractions, the ions may leave the crystal lattice
and go into solution as aquated ions, each surrounded by a sheath of water
molecules. These aquated ions are free to move around throughout the solution.
Once dissolved, ionic bonds cease to exist as this form of bonding is really only
present in the solid state. Not all ionic compounds are soluble and solubility tables
listing this data are available.
If the solvent is boiled off from a solution of an ionic compound, the ions, stripped
of their surrounding water molecules, are non-volatile and remain to recombine as
crystals of the solid compound.
If cations and anions from two different salts are mixed in the same solution and
any combination of them corresponds to an insoluble salt, then that compound will
be precipitated from the solution. Other cations and anions present will remain in
solution and are termed spectator ions. The precipitated solid can be collected by
filtration and the remaining solution (the filtrate) can be boiled down to collect any
soluble compounds. This type of reaction is best represented by an ionic equation
in which only the reacting ions are shown and the spectator ions are deleted as they
are present on both sides of the reaction equation.
Acids are compounds that release the H+ ion in solution. This ion would strictly be
just a proton with no electron, but this species would have too large a charge on too
small a volume to be stable, so it associates with water molecules via a lone pair of
electrons on the O atom and thus can also be represented as H+(aq) or H 3O+. Apart
from the anion from the acid in solution, reactions of the H+ ion are conveniently
grouped into several classes.
Acids on reactive metals: form a salt and hydrogen gas.
Acids on oxides of metals: form a salt and water.
Acids on hydroxides of metals: form a salt and water.
Acids on a carbonate: form a salt and water and carbon dioxide gas.
In each case, the anion in the salt formed is the anion that belonged to the acid, for
example, chloride from hydrochloric acid or nitrate from nitric acid.
Each of these reaction types are conveniently represented by ionic equations which
more clearly show the process which is taking place than overall formula equations.
12. VI - 12
TUTORIAL QUESTIONS - TOPIC 6.
1. Explain the meaning of the following terms:
solvent; solution; precipitate; filtrate; formula equation; ionic equation;
volatile substance.
2. Give the missing formula or name for each entry in the following table.
FORMULA
NAME
cobalt(II) chloride
CuBr2
magnesium sulfate
BaSO4
NH4I
Mn(NO3)2
AgF
NaCl
Fe3(PO4)2
FePO4
ZnCO3
K2CO 3
Ca(CH3CO 2)2
Zn(HCO3)2
LiOH
CuSO4
Pb(NO3)2
Pb(NO2)2
Li2SO3
lead(II) carbonate
strontium sulfite
aluminium nitrate
silver oxalate
tin(II) acetate
iron(II) bromide
sodium nitrite
iron(III) bromide
nickel(II) iodide
lithium carbonate
13. VI - 13
3. Write equations for the dissolution of the following ionic compounds in water.
cobalt(II) sulfate
ammonium nitrate
iron(III) chloride
sodium phosphate
aluminium nitrate
potassium
carbonate
iron(II) bromide
4. Reactions of acids with reactive metals.
Write (i) the formula equation and (ii) the ionic equation for the reaction that occurs
in each of the following:
(a) Hydrochloric acid is added to magnesium.
(i) Formula equation
(ii) Ionic equation
(b) Hydrochloric acid is added to aluminium.
(i) Formula equation
(ii) Ionic equation
(c) Sulfuric acid is added to zinc.
(i) Formula equation
(ii) Ionic equation
14. VI - 14
5. Reactions of acids with oxides and hydroxides of metals.
Write (i) the formula equation and (ii) the ionic equation for the reaction that occurs
in each of the following:
(a) Hydrochloric acid is added to magnesium oxide.
(i) Formula equation
(ii) Ionic equation
(b) Sulfuric acid is added to solid sodium hydroxide.
(i) Formula equation
(ii) Ionic equation
(c) Hydrochloric acid is added to zinc oxide.
(i) Formula equation
(ii) Ionic equation
6. Reactions of acids with carbonates of metals.
Write (i) the formula equation and (ii) the ionic equation for the reaction that occurs
in each of the following:
(a) Hydrochloric acid is added to solid sodium carbonate.
(i) Formula equation
(ii) Ionic equation
(b) Sulfuric acid is added to copper(II) carbonate.
(i) Formula equation
(ii) Ionic equation
15. VI - 15
(c) Nitric acid is added to silver carbonate.
(i) Formula equation
(ii) Ionic equation
7. Precipitation reactions.
Write (i) the formula equation and (ii) the ionic equation for any reaction that occurs
in each of the following when the specified solutions are mixed:
(a) a water solution of sodium sulfate and a water solution of barium chloride.
(i) Formula equation
(ii) Ionic equation
(b) a water solution of silver nitrate and a water solution of sodium chloride.
(i) Formula equation
(ii) Ionic equation
(c) a water solution of potassium iodide and a water solution of lead(II) nitrate.
(i) Formula equation
(ii) Ionic equation
(d) a water solution of potassium iodide and a water solution of sodium chloride.
(i) Formula equation
(ii) Ionic equation
(e) water solution of copper(II) nitrate and a water solution sodium carbonate.
(i) Formula equation
(ii) Ionic equation
16. VI - 16
8. Write ionic equations for the reaction of magnesium carbonate with each of the
acids - hydrochloric acid, nitric acid, sulfuric acid, hydrobromic acid, hydriodic
acid. What do you notice about the equations? Explain your observation.
9. What is a "salt"? Which of the following compounds are salts and how did you
decide in each case?
Na3PO4, SO3, BaBr2, NH 3, Mg(NO 3)2, CO 2, H 2O, AgCl, CCl 4, H 2SO 4, H 2S, HNO 3
17. VI - 17
ANSWERS TO TUTORIAL TOPIC 6
1. Solvent: Substance in which another substance (solute) dissolves, the
combination being termed a solution. Solvents are usually but not necessarily
liquids.
Solution: A homogeneous combination of two or more substances which retain
their separate identities.
Precipitate: A solid formed in a chemical reaction.
Filtrate: The liquid phase that remains after any solids previously present have
been removed, for example by filtering.
Formula equation: A representation of a chemical reaction in which all the
reactants and products are represented by their entire formulas.
Ionic equation: A representation of a chemical reaction in which any species
which are present as ions are shown as such. Only the species actually reacting
are shown and any other ions are called spectator ions because they play no part
in the actual reaction.
Volatile substance: Element or compound that can be readily converted to the
gaseous state.
2. Formulas:
CoCl 2; MgSO 4; PbCO 3; SrSO 3; Al(NO 3) 3; Ag 2C 2O 4;
Sn(CH 3CO 2)2; FeBr2; NaNO 2; FeBr3; NiI2; Li2CO 3.
Names:
copper(II) bromide; barium sulfate; ammonium iodide;
manganese(II) nitrate; silver(I) fluoride; sodium chloride;
iron(II) phosphate; iron(III) phosphate; zinc carbonate;
potassium carbonate; calcium acetate; zinc hydrogencarbonate;
lithium hydroxide; copper(II) sulfate; lead(II) nitrate;
lead(II) nitrite; lithium sulfite.
Note the use of Roman numerals in conjunction with those cations that can
occur with more than one cationic charge, e.g. copper
v
Co2+(aq)
+
SO 42–(aq)
NH4NO3(s) v
NH 4+(aq)
+
NO 3–(aq)
3. CoSO4(s)
FeCl3(s)
v
Fe3+(aq)
Na3PO4(s)
v
3Na+(aq)
+
+
3Cl–(aq)
(note 3Cl– required to balance)
PO 43–(aq) (note 3Na+ required to balance)
Al(NO3)3(s) v
Al3+ (aq)
+
3NO 3–(aq) (note 3NO 3– required to balance)
K2CO3(s)
v
2K+ (aq)
+
CO 32– (aq)
(note 2K+ required to balance)
FeBr2(s)
v
Fe2+ (aq)
+
2Br– (aq)
(note 2Br– required to balance)
18. VI - 18
4. When writing ionic equations, make sure that you have given the correct
number of each ion on both sides of the equation for any compound that
contains more than one of a given ion - all equations must balance in the same
way as formula equations.
(a)
(i) Mg + 2HCl
v MgCl 2 + H 2
(ii) Mg(s) + 2H+(aq)
(b)
(i) 2Al + 6HCl
v
2AlCl3 + 3H 2
(ii) 2Al(s) + 6H+(aq)
(c)
(i) Zn + H 2SO 4
v
ZnSO 4 + H 2
MgCl 2 + H 2O
(ii) MgO(s) + 2H+(aq)
(b)
v
(i) 2NaOH + H 2SO 4
(i) ZnO + 2HCl
v
(ii) ZnO(s) + 2H+(aq)
6. (a)
(i) Na2CO 3 + 2HCl
Mg2+(aq) + H 2O(l)
v
(ii) NaOH(s) + H+(aq)
(c)
Zn2+(aq) + H 2(g)
v
v
(i) MgO + 2HCl
2Al3+(aq) + 3H 2(g)
v
(ii) Zn(s) + 2H+(aq)
5. (a)
Mg2+(aq) + H 2(g)
v
v
Na 2SO 4 +
2H 2O
Na+(aq) + H 2O(l)
ZnCl2 + H 2O
v
v
(ii) Na2CO 3(s) + 2H+(aq)
Zn2+(aq) + H 2O(l)
2NaCl + H 2O + CO 2
v
2Na+(aq) + H 2O(l) + CO 2(g)
19. VI - 19
(b)
(i) CuCO 3 + 2HCl
v
CuCl 2 + H 2O + CO 2
(ii) CuCO 3(s) + 2H+(aq)
(c)
(i) Ag2CO 3 + 2HCl
v
2AgCl + H 2O + CO 2
(ii) Ag2CO 3(s) + 2H+(aq)
7. (a)
(i) Na2SO 4 + BaCl2
(ii) Ba2+ (aq)
+
Cu2+(aq) + H 2O(l) + CO 2(g)
v
2Ag+(aq) + H 2O(l) + CO 2(g)
v
v
BaSO 4 + 2NaCl
SO 42–(aq)
v
BaSO 4(s)
[Consider the formulas of the two possible compounds that might form, BaSO 4
and NaCl. Only barium sulfate is insoluble, so it will precipitate while sodium
ions and chloride ions remain in solution.]
(b)
(i) AgNO 3 + NaCl
(ii) Ag+(aq)
+
v
Cl–(aq)
AgCl + NaNO 3
v
AgCl(s)
[The combination of silver(I) ions and chloride ions results in an insoluble
compound, silver(I) chloride, being precipitated while the sodium and nitrate
ions remain in solution.]
(c)
(i) 2KI + Pb(NO 3) 2
(ii) I!(aq) + Pb2+(aq)
v
v
2KNO 3 + PbI 2
PbI2(s)
[The potassium iodide provides K+ (aq) and I–(aq) ions in the solution and
Pb2+ (aq) and NO 3! ions come from the lead(II) nitrate solution. The
combination of Pb2+ and I– ions results in the formation of lead(II) iodide
precipitate. The potassium and nitrate ions remain in solution.]
(d)
no reaction
[The ions in the mixture of the two solutions are K+(aq), I!(aq), Na+(aq) and
Cl! (aq). There is no combination of these cations and anions that form an
insoluble compound.
20. VI - 20
(e)
(i) Cu(NO 3)2 + Na 2CO 3
(ii) Cu2+(aq) + CO 32!(aq)
v
v
CuCO 3 + 2NaNO 3
CuCO 3(s)
[The combination of copper(II) ions and carbonate ions results in an insoluble
compound, copper(II) carbonate, being precipitated while the sodium and
nitrate ions remain in solution.]
8. MgCO3(s) + 2H+ (aq)
v Mg2+ (aq) + CO 2(g) + H 2O(l)
MgCO3(s) + 2H+ (aq)
v Mg2+ (aq) + CO 2(g) + H 2O(l)
MgCO3(s) + 2H+ (aq)
v Mg2+ (aq) + CO 2(g) + H 2O(l)
MgCO3(s) + 2H+ (aq)
v Mg2+ (aq) + CO 2(g) + H 2O(l)
MgCO3(s) + 2H+ (aq)
v Mg2+ (aq) + CO 2(g) + H 2O(l)
The equations are identical because the reaction does not involve the chloride,
nitrate, sulfate, bromide or iodide ions which are merely spectator ions in each
case.
9. A salt is any ionic compound in which the cation is not H+ or the anion is not
OH– . The compounds Na 3PO 4, BaBr2, Mg(NO 3)2, AgCl are all salts
because they meet these criteria. The other compounds are either covalently
bonded compounds involving only non-metals (SO 3, NH 3, CO 2, H 2O, CCl 4,
H2S) or acids (HNO 3, H 2SO 4).