1) A chemical reaction can be differentiated from a phase change or nuclear change based on whether it involves a change in chemical composition, state of matter, or type of element.
2) In a chemical reaction, reactants yield products through the rearrangement of atoms. Chemical equations use symbols to represent the reactants and products.
3) Evidence that a chemical reaction occurred includes the release or absorption of energy, production of a gas, formation of a precipitate, or a color change.
This document provides an overview of chapter 4 from a chemistry textbook, which covers chemical reactions. It begins by defining electrolytes, non-electrolytes, and discussing the properties of aqueous solutions. It then covers the three main types of reactions that occur in aqueous solutions: precipitation reactions, acid-base reactions involving proton transfer, and redox reactions involving electron transfer. Specific examples of each reaction type are provided. Key concepts around oxidation, reduction, and oxidation numbers are also explained.
This document contains lecture notes on quantitative analysis in chemistry. It discusses gravimetric analysis, which determines the amount of a substance by converting it into a product that can be isolated and weighed. An example is given of determining the amount of lead in water by precipitating lead sulfate, filtering and weighing the precipitate. A practice problem demonstrates calculating the mass of lead from the mass of lead sulfate precipitate obtained.
This document discusses different types of chemical reactions including precipitation, acid-base, and oxidation-reduction reactions. It provides examples of strong acids, weak acids, and bases. Neutralization reactions between acids and bases are described which produce water and a salt. Oxidation and reduction are defined as the loss or gain of electrons respectively. Half-reaction and oxidation state rules are outlined for balancing redox reactions.
The document discusses various chemistry concepts including:
1. Avogadro's number defines the number of particles in one mole of a substance as 6.02x1023.
2. Molar mass is the mass of one mole of a pure substance and is used to convert between grams and moles.
3. Empirical and molecular formulas can be determined from percent composition data using mole ratios and molar mass.
4. Chemical equations represent chemical reactions and must be balanced to satisfy the law of conservation of mass.
The document discusses redox reactions and concepts of oxidation and reduction:
1) It provides examples of redox reactions and identifies which species are oxidized and reduced as well as the oxidizing agent.
2) It explains that metals are oxidized when they form cations and that cations are reduced when they gain electrons to form elemental metals.
3) Combustion reactions are identified as redox reactions where carbon and hydrogen are oxidized and oxygen is reduced.
4) Rules for assigning oxidation numbers to elements in compounds are outlined and examples are provided to illustrate oxidation and reduction based on changes in oxidation numbers.
The document discusses chemical reactions and equations. It defines indicators of chemical reactions, such as gas formation or color change. It explains that reactants turn into products in chemical reactions, and describes ways to write chemical equations, including using symbols to indicate state or energy changes. It also summarizes the main types of chemical reactions: synthesis, decomposition, single replacement, double replacement, and acid-base reactions.
This document provides information on stoichiometry, which involves using mole ratios from balanced chemical equations to calculate mass relationships between substances in a chemical reaction. It outlines the steps to solve stoichiometry problems, which include writing a balanced equation, identifying known and unknown quantities, setting up mole ratio conversion factors between moles of reactants and products, and checking the answer. Key concepts discussed include the mole ratio from coefficients in a balanced equation, molar mass to convert between moles and grams, and the molar volume used to calculate liters of gas at standard temperature and pressure.
This document provides an overview of chapter 4 from a chemistry textbook, which covers chemical reactions. It begins by defining electrolytes, non-electrolytes, and discussing the properties of aqueous solutions. It then covers the three main types of reactions that occur in aqueous solutions: precipitation reactions, acid-base reactions involving proton transfer, and redox reactions involving electron transfer. Specific examples of each reaction type are provided. Key concepts around oxidation, reduction, and oxidation numbers are also explained.
This document contains lecture notes on quantitative analysis in chemistry. It discusses gravimetric analysis, which determines the amount of a substance by converting it into a product that can be isolated and weighed. An example is given of determining the amount of lead in water by precipitating lead sulfate, filtering and weighing the precipitate. A practice problem demonstrates calculating the mass of lead from the mass of lead sulfate precipitate obtained.
This document discusses different types of chemical reactions including precipitation, acid-base, and oxidation-reduction reactions. It provides examples of strong acids, weak acids, and bases. Neutralization reactions between acids and bases are described which produce water and a salt. Oxidation and reduction are defined as the loss or gain of electrons respectively. Half-reaction and oxidation state rules are outlined for balancing redox reactions.
The document discusses various chemistry concepts including:
1. Avogadro's number defines the number of particles in one mole of a substance as 6.02x1023.
2. Molar mass is the mass of one mole of a pure substance and is used to convert between grams and moles.
3. Empirical and molecular formulas can be determined from percent composition data using mole ratios and molar mass.
4. Chemical equations represent chemical reactions and must be balanced to satisfy the law of conservation of mass.
The document discusses redox reactions and concepts of oxidation and reduction:
1) It provides examples of redox reactions and identifies which species are oxidized and reduced as well as the oxidizing agent.
2) It explains that metals are oxidized when they form cations and that cations are reduced when they gain electrons to form elemental metals.
3) Combustion reactions are identified as redox reactions where carbon and hydrogen are oxidized and oxygen is reduced.
4) Rules for assigning oxidation numbers to elements in compounds are outlined and examples are provided to illustrate oxidation and reduction based on changes in oxidation numbers.
The document discusses chemical reactions and equations. It defines indicators of chemical reactions, such as gas formation or color change. It explains that reactants turn into products in chemical reactions, and describes ways to write chemical equations, including using symbols to indicate state or energy changes. It also summarizes the main types of chemical reactions: synthesis, decomposition, single replacement, double replacement, and acid-base reactions.
This document provides information on stoichiometry, which involves using mole ratios from balanced chemical equations to calculate mass relationships between substances in a chemical reaction. It outlines the steps to solve stoichiometry problems, which include writing a balanced equation, identifying known and unknown quantities, setting up mole ratio conversion factors between moles of reactants and products, and checking the answer. Key concepts discussed include the mole ratio from coefficients in a balanced equation, molar mass to convert between moles and grams, and the molar volume used to calculate liters of gas at standard temperature and pressure.
Periodic Properties Of Elements In The Periodic Tablejuanjose
The document summarizes periodic properties of elements in the periodic table, including periodic trends in atomic radius, ionization energy, electronegativity, and melting points. It also discusses periodic trends in chemical properties such as formulas of hydrides, oxides, and chlorides as well as their hydrolytic behaviors.
1. The half-reaction method consists of 7 steps to balance redox equations. This includes writing oxidation and reduction half-reactions, balancing atoms and charge, and combining the half-reactions.
2. The reaction of potassium permanganate with iron(II) sulfate is used as an example. The oxidation half-reaction involves manganese and the reduction half-reaction involves iron.
3. The balanced equation for the reaction is: 10FeSO4 + 2KMnO4 + 8H2SO4 → 5Fe3(SO4)2 + 2MnSO4 + K2SO4 + 8H2O.
The document discusses various topics related to reactions in aqueous solutions including:
- Solutions, solvents, solutes, electrolytes, and nonelectrolytes
- Strong and weak electrolytes and their ionization
- Acid-base theories and classifications of acids and bases
- Oxidation-reduction reactions and oxidation numbers
- Precipitation, acid-base, and redox reactions
- Solution stoichiometry, dilution, titrations, and gravimetric analysis
The document provides information on chemical formulae, equations, calculations involving moles and molar mass/volume. It also covers the chemical properties and reactions of group 1 and 17 elements, as well as properties of salts such as solubility, color, and the effects of heating on different salts such as carbonates and nitrates.
This document discusses three tables that are part of an electrochemical series. Table 1 provides an alphabetical list of reduction reactions and their standard reduction potentials (E° values) in volts. Table 2 lists reactions that have positive E° values relative to the standard hydrogen electrode, in order of increasing potential. Table 3 lists reactions with negative E° values relative to the standard hydrogen electrode, in order of decreasing potential. The document notes that the reliability of the E° values varies and is sometimes subject to revision.
This document contains a chemistry exam paper with multiple choice and long answer questions. It tests knowledge of topics including:
1. Cation and anion radii and lattice energies of ionic compounds.
2. Enthalpy calculations involving formation reactions of molecules like dinitrogen tetroxide.
3. Determining standard electrode potentials and writing redox half reactions.
4. Calculating standard enthalpies of solution and hydration using calorimetry data.
5. Trends in first ionization energies across the third period and acid-base properties of oxides.
This document summarizes different types of chemical reactions:
1) Combustion reactions involve the burning of a substance in oxygen to produce heat and oxides.
2) Synthesis reactions combine two or more substances to form a single product compound.
3) Decomposition reactions involve a single reactant breaking down into simpler substances.
4) Single replacement reactions involve one element replacing another in a compound.
1. The document provides information on the rates of chemical reactions including definitions and methods for calculating rates from graphs.
2. It also summarizes several common chemical reactions like precipitation, combustion, substitution, addition, oxidation and reduction reactions.
3. Details are given for important industrial processes like the Haber, Contact, and Ostwald processes for producing ammonia, sulfuric acid, and nitric acid respectively.
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.
Chemistry deals with the composition, properties, and interactions of matter, as well as energy changes. There are four main topics in chemistry - what matter is made of, how its particles are arranged, what characteristics it exhibits, and how it undergoes transformations through different types of changes including phase changes, chemical reactions, and nuclear reactions. Scientists make observations, which can be qualitative descriptions or quantitative measurements, to understand the physical and chemical properties of substances.
There are three types of changes in matter:
1) Phase changes involve a change in state without a change in chemical composition, such as water changing between solid, liquid, and gas forms. Energy is absorbed or released during these changes.
2) Chemical changes involve a change in chemical composition, as seen through color changes, formation of solids/gases, or energy absorption/release. An example is the reaction of hydrogen and oxygen gases to form water.
3) Nuclear changes involve altering the composition of an element's nucleus through processes like fission, fusion, and radioactive decay. Energy is released or absorbed during these changes.
- Mendeleev created the periodic table to organize elements based on similarities in their chemical reactions and properties.
- The periodic table is arranged in 18 vertical columns called groups and horizontal rows called periods.
- Elements are classified as metals, nonmetals, metalloids, or noble gases based on their physical properties such as appearance, conductivity, state of matter, and reactivity.
This document defines and classifies different types of matter. It describes elements as substances made of single atoms and compounds as bonded groups of elements. Pure substances include elements and compounds, while mixtures contain two or more substances mixed without chemical bonding. Mixtures can be homogeneous, appearing uniform, or heterogeneous, with distinct phases. Special types of mixtures include colloids, ores, alloys, and plated metals. The document provides examples to illustrate each classification of matter.
This chapter discusses different types of chemical reactions in aqueous solutions. It introduces driving forces that cause reactions, such as formation of a solid, water, or gas. It explains how to predict products using solubility rules and oxidation-reduction reactions when metals react with nonmetals. Reactions are classified into double displacement, acid-base, single replacement, combustion, synthesis, or decomposition reactions based on their driving forces.
Chemical reactions involve reactants transforming into products. There are several types of chemical reactions that can be identified by their reactants and products. These include synthesis reactions where two reactants combine to form one product, decomposition reactions where one reactant breaks down into multiple products, and displacement reactions where one element replaces another in a compound. Balancing chemical equations ensures the same number and type of atoms are on both sides of the reaction.
The document discusses balancing redox reactions using the half-reaction method. It provides several examples of writing and balancing half-reactions and using them to derive the overall balanced redox equation. Key steps include separating the reaction into oxidation and reduction half-reactions, balancing all elements except H and O, adding H2O to balance O, adding H+ or OH- to balance H, and adding electrons to balance charge.
The document discusses different types of chemical reactions including synthesis, decomposition, single displacement, and double displacement reactions. It provides examples of each type of reaction along with balanced chemical equations. It also describes the activity series of metals and how reactivity determines whether single displacement reactions will occur.
ch.-1 chemical reactions and equations .pptxRajat Sardana
Here are some potential next steps based on the findings:
- Further research different types of chemical reactions like displacement, double displacement, oxidation-reduction, etc. to gain a deeper understanding of each.
- Practice writing balanced chemical equations for various reactions to strengthen skills in this area. Identifying the reactants, products, and coefficients is an important part of understanding chemical changes.
- Explore applications of chemical reactions in different contexts like cooking, medicine, industry, environment, etc. to see the relevance of chemistry concepts.
- Design experiments to observe chemical reactions firsthand in the lab. Seeing reactions occur can help cement conceptual understanding. Safety should always be the top priority for any experiments.
- Teach others about
Chemical formulae, equations, calculations, and reactions are summarized. Molar mass, moles, volume, and molarity calculations are explained for gases, solids, liquids, and solutions. Common cationic and anionic symbols are listed. Formulae for common compounds are provided. Group 1 and 17 elements and their reactions are summarized. Electrochemistry concepts like electrolytes, ionization, and the discharge of ions are condensed. Acid-base reactions and properties are highlighted at a high level. Solubility, preparation, color, and the effect of heating on salts are briefly touched upon.
Periodic Properties Of Elements In The Periodic Tablejuanjose
The document summarizes periodic properties of elements in the periodic table, including periodic trends in atomic radius, ionization energy, electronegativity, and melting points. It also discusses periodic trends in chemical properties such as formulas of hydrides, oxides, and chlorides as well as their hydrolytic behaviors.
1. The half-reaction method consists of 7 steps to balance redox equations. This includes writing oxidation and reduction half-reactions, balancing atoms and charge, and combining the half-reactions.
2. The reaction of potassium permanganate with iron(II) sulfate is used as an example. The oxidation half-reaction involves manganese and the reduction half-reaction involves iron.
3. The balanced equation for the reaction is: 10FeSO4 + 2KMnO4 + 8H2SO4 → 5Fe3(SO4)2 + 2MnSO4 + K2SO4 + 8H2O.
The document discusses various topics related to reactions in aqueous solutions including:
- Solutions, solvents, solutes, electrolytes, and nonelectrolytes
- Strong and weak electrolytes and their ionization
- Acid-base theories and classifications of acids and bases
- Oxidation-reduction reactions and oxidation numbers
- Precipitation, acid-base, and redox reactions
- Solution stoichiometry, dilution, titrations, and gravimetric analysis
The document provides information on chemical formulae, equations, calculations involving moles and molar mass/volume. It also covers the chemical properties and reactions of group 1 and 17 elements, as well as properties of salts such as solubility, color, and the effects of heating on different salts such as carbonates and nitrates.
This document discusses three tables that are part of an electrochemical series. Table 1 provides an alphabetical list of reduction reactions and their standard reduction potentials (E° values) in volts. Table 2 lists reactions that have positive E° values relative to the standard hydrogen electrode, in order of increasing potential. Table 3 lists reactions with negative E° values relative to the standard hydrogen electrode, in order of decreasing potential. The document notes that the reliability of the E° values varies and is sometimes subject to revision.
This document contains a chemistry exam paper with multiple choice and long answer questions. It tests knowledge of topics including:
1. Cation and anion radii and lattice energies of ionic compounds.
2. Enthalpy calculations involving formation reactions of molecules like dinitrogen tetroxide.
3. Determining standard electrode potentials and writing redox half reactions.
4. Calculating standard enthalpies of solution and hydration using calorimetry data.
5. Trends in first ionization energies across the third period and acid-base properties of oxides.
This document summarizes different types of chemical reactions:
1) Combustion reactions involve the burning of a substance in oxygen to produce heat and oxides.
2) Synthesis reactions combine two or more substances to form a single product compound.
3) Decomposition reactions involve a single reactant breaking down into simpler substances.
4) Single replacement reactions involve one element replacing another in a compound.
1. The document provides information on the rates of chemical reactions including definitions and methods for calculating rates from graphs.
2. It also summarizes several common chemical reactions like precipitation, combustion, substitution, addition, oxidation and reduction reactions.
3. Details are given for important industrial processes like the Haber, Contact, and Ostwald processes for producing ammonia, sulfuric acid, and nitric acid respectively.
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.
Chemistry deals with the composition, properties, and interactions of matter, as well as energy changes. There are four main topics in chemistry - what matter is made of, how its particles are arranged, what characteristics it exhibits, and how it undergoes transformations through different types of changes including phase changes, chemical reactions, and nuclear reactions. Scientists make observations, which can be qualitative descriptions or quantitative measurements, to understand the physical and chemical properties of substances.
There are three types of changes in matter:
1) Phase changes involve a change in state without a change in chemical composition, such as water changing between solid, liquid, and gas forms. Energy is absorbed or released during these changes.
2) Chemical changes involve a change in chemical composition, as seen through color changes, formation of solids/gases, or energy absorption/release. An example is the reaction of hydrogen and oxygen gases to form water.
3) Nuclear changes involve altering the composition of an element's nucleus through processes like fission, fusion, and radioactive decay. Energy is released or absorbed during these changes.
- Mendeleev created the periodic table to organize elements based on similarities in their chemical reactions and properties.
- The periodic table is arranged in 18 vertical columns called groups and horizontal rows called periods.
- Elements are classified as metals, nonmetals, metalloids, or noble gases based on their physical properties such as appearance, conductivity, state of matter, and reactivity.
This document defines and classifies different types of matter. It describes elements as substances made of single atoms and compounds as bonded groups of elements. Pure substances include elements and compounds, while mixtures contain two or more substances mixed without chemical bonding. Mixtures can be homogeneous, appearing uniform, or heterogeneous, with distinct phases. Special types of mixtures include colloids, ores, alloys, and plated metals. The document provides examples to illustrate each classification of matter.
This chapter discusses different types of chemical reactions in aqueous solutions. It introduces driving forces that cause reactions, such as formation of a solid, water, or gas. It explains how to predict products using solubility rules and oxidation-reduction reactions when metals react with nonmetals. Reactions are classified into double displacement, acid-base, single replacement, combustion, synthesis, or decomposition reactions based on their driving forces.
Chemical reactions involve reactants transforming into products. There are several types of chemical reactions that can be identified by their reactants and products. These include synthesis reactions where two reactants combine to form one product, decomposition reactions where one reactant breaks down into multiple products, and displacement reactions where one element replaces another in a compound. Balancing chemical equations ensures the same number and type of atoms are on both sides of the reaction.
The document discusses balancing redox reactions using the half-reaction method. It provides several examples of writing and balancing half-reactions and using them to derive the overall balanced redox equation. Key steps include separating the reaction into oxidation and reduction half-reactions, balancing all elements except H and O, adding H2O to balance O, adding H+ or OH- to balance H, and adding electrons to balance charge.
The document discusses different types of chemical reactions including synthesis, decomposition, single displacement, and double displacement reactions. It provides examples of each type of reaction along with balanced chemical equations. It also describes the activity series of metals and how reactivity determines whether single displacement reactions will occur.
ch.-1 chemical reactions and equations .pptxRajat Sardana
Here are some potential next steps based on the findings:
- Further research different types of chemical reactions like displacement, double displacement, oxidation-reduction, etc. to gain a deeper understanding of each.
- Practice writing balanced chemical equations for various reactions to strengthen skills in this area. Identifying the reactants, products, and coefficients is an important part of understanding chemical changes.
- Explore applications of chemical reactions in different contexts like cooking, medicine, industry, environment, etc. to see the relevance of chemistry concepts.
- Design experiments to observe chemical reactions firsthand in the lab. Seeing reactions occur can help cement conceptual understanding. Safety should always be the top priority for any experiments.
- Teach others about
Chemical formulae, equations, calculations, and reactions are summarized. Molar mass, moles, volume, and molarity calculations are explained for gases, solids, liquids, and solutions. Common cationic and anionic symbols are listed. Formulae for common compounds are provided. Group 1 and 17 elements and their reactions are summarized. Electrochemistry concepts like electrolytes, ionization, and the discharge of ions are condensed. Acid-base reactions and properties are highlighted at a high level. Solubility, preparation, color, and the effect of heating on salts are briefly touched upon.
Chemical formulae, equations, calculations, and reactions are summarized. Molar mass, moles, volume, and molarity calculations are explained for gases, solids, liquids, and solutions. Common cationic and anionic symbols are listed. Formulae for molecules and ions are provided. Periodic trends and reactions of Groups 1 and 17 are summarized. Electrochemistry principles of electrolytes, discharge reactions, and test observations are condensed. Characteristics of acids, bases, and ionization are highlighted. Solubility, preparation, color, and effects of heating for various salts are summarized concisely.
This document discusses different types of chemical reactions including synthesis, decomposition, single displacement, and double displacement reactions. It provides examples of each type of reaction using chemical equations. It also describes the activity series of metals and how reactivity determines which metals can displace other metals in single displacement reactions.
The document discusses different types of chemical reactions including synthesis, decomposition, single replacement, double replacement, and combustion reactions. It provides examples of each type of reaction and how to write and balance chemical equations to properly represent these reactions. Key aspects covered include reactants and products, word equations, chemical formulas and symbols, and balancing equations so the number of atoms of each element are equal on both sides.
The document discusses chemical reactions and equations. It defines chemical reactions as chemical changes where bonds in reactants break and products form. It also lists several types of chemical reactions including synthesis, decomposition, single replacement, and combustion. The document emphasizes the importance of balancing chemical equations so that mass is conserved in reactions according to the law of conservation of mass.
The document discusses chemical reactions and equations. It defines chemical reactions as chemical changes where bonds in reactants break and products form. It also lists several types of chemical reactions including synthesis, decomposition, single replacement, double replacement, and combustion reactions. The document emphasizes the importance of balancing chemical equations so that mass is conserved in reactions according to the law of conservation of mass.
The document discusses chemical reactions and equations. It explains that a chemical reaction involves the breaking and forming of bonds between reactants and products. It also describes various types of chemical reactions including synthesis, decomposition, single replacement, and combustion. Additionally, it discusses how to write and balance chemical equations and identifies factors like states of matter and coefficients.
This document discusses oxidation and reduction reactions. It begins by defining oxidation as a reaction where substances combine with oxygen and reduction as a reaction where a substance "gave up" oxygen. It then explains that oxidation and reduction actually refer to the gain or loss of electrons in a chemical reaction, regardless of whether oxygen is present. Oxidation involves the loss of electrons, while reduction involves the gain of electrons. Redox reactions always involve both oxidation and reduction occurring together through the transfer of electrons. The document provides examples of how to identify the oxidizing agent, reducing agent, and what is being oxidized and reduced in redox reactions. It also discusses how to balance redox reactions through half-reactions and the role of acid and
This document discusses the main types of chemical reactions: synthesis, decomposition, single displacement, double displacement, and combustion. It provides examples of each type and teaches how to identify and write balanced chemical equations to predict products from reactants. Key aspects covered include oxidation/reduction, use of phase labels (s), (g), (aq), and how double displacement reactions involve switching the outermost ions in the reactants. Mixed practice problems are provided to allow identification and writing of different reaction types.
The document discusses solubility and writing ionic equations. It defines a soluble substance as one that dissociates into ions in aqueous solution. It provides a 5-step process for writing ionic equations: 1) Write the balanced molecular equation, 2) Add state symbols, 3) Write out aqueous ions, 4) Cancel spectator ions on both sides, 5) Write the net ionic equation with only reactants and products. An example uses this process to derive the net ionic equation for the reaction of lead nitrate and potassium iodide.
This document describes different types of chemical reactions:
1) Synthesis reactions where two reactants combine to form one new product
2) Decomposition reactions where one reactant breaks down into multiple products
3) Single displacement reactions where a metal displaces another in a compound
4) Double displacement reactions where elements in reactants exchange places
It also provides examples of each type of reaction.
This document provides information on acid-base reactions and oxidation-reduction (redox) reactions. It defines acids and bases, and explains that in acid-base reactions, acids donate protons to bases. Neutralization reactions between acids and bases produce water and a salt. The document also discusses how to determine oxidation states of elements in compounds and identify the oxidized and reduced substances in redox reactions. It provides steps for balancing redox equations, including dividing the reaction into partial equations and adding electrons to balance charges. Examples of assigning oxidation states and balancing redox reactions are included.
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1. Chemistry 20 – Review Unit 2 Chemical Reactions - A chemical reaction can be differentiated from a phase change or a nuclear change in the following manner: state of matter C 25 H 52(s) C 25 H 52(l) chemical composition C 25 H 52(s) + O 2(g) CO 2(g) + H 2 O (g) type of element * Type of Change What Changes? Example Phase Chemical Nuclear
2. - In all chemical reactions, you start with one or more substances called reactants and produce substances called products . - The plus symbol ( + ) is used to separate lists of chemicals and is interpreted as the word “ and ”. - Reactants and products are separated by an arrow ( ) and is interpreted as the word “ yields ”. - How do chemists recognize that a chemical reaction has taken place? What is the evidence for a chemical reaction ? *
3. 1) Energy Exchange – 1) Energy Exchange – heat is given off ( exothermic reaction) or heat is absorbed by the reactants ( endothermic reaction) Examples: - the carbon in wood burns, releasing energy in the form of light and heat, as products are made - C (s) + O 2(g) CO 2(g) + heat + light - exothermic reaction - electricity ( energy ) is added to water in a process called electrolysis - 2 H 2 O (l) + energy 2 H 2(g) + O 2(g) - endothermic reaction *
4. 2) Gas ( bubbles ) Produced – sometimes bubbling can be easily noticed ( fizzing ); at other times a vapour is produced ( smoke ) Examples: - an acid is poured into a metal - Zn (s) + HCl (aq) ZnCl 2(aq) + H 2(g) - baking soda fizzes in water - NaHCO 3(s) + H 2 O (l) CO 2(g) + NaOH (aq) 3) Formation of a Precipitate – a precipitate is a solid , usually formed by mixing two clear solutions Examples: - two clear solutions are mixed and a yellow solid forms - 2 NaI (aq) + Pb(NO 3 ) 2(aq) 2 NaNO 3(aq) + PbI 2(s) - iron rusts when left out in the rain - 2 Fe (s) + 2 H 2 O (l) + O 2(g) 2 Fe(OH) 2(s) *
5. 4) Colour Change – a substance formed may exhibit a new characteristic colour Examples: - when nitric acid reacts with copper metal, a dark brown gas is produced and the resulting solution turns blue - Cu (s) + HNO 3(aq) NO 2(g) + Cu(NO 3 ) 2(aq) + H 2 O (l) Note: This equation is not balanced. Technically, the equation is written differently in Chemistry 30, but for illustrative purposes, we’ll use it in this form. *
6. Types of Chemical Reactions Simple Composition - Just as the name implies, this reaction is the simplest of all. It implies that a product is composed from the simplest substances . - A generic equation can be used to describe a simple composition reaction. - Remember that reactants are placed on the left side of the equation and the products are found on the right side. - element + element compound * any number of elements only one compound formed
7. hydrogen and oxygen yields water vapour H 2(g) + O 2(g) H 2 O (g) - Is this equation correct? What is missing? - Whenever we write out a chemical formula, we must make sure that the equation is balanced . - The number of atoms of each element must be the same number on both reactant and product sides. - To finish balancing the equation: H 2(g) + O 2(g) H 2 O (g) - There are 2 hydrogens on the left and 2 hydrogens on the right, therefore the hydrogens are balanced . * Examples 1) Word Equation Chemical Equation (Formula)
8. H 2(g) + O 2(g) H 2 O (g) - There are 2 oxygens on the left but 1 oxygen on the right. - To balance the oxygens we place a number in front of water to make the oxygens balance. The number being placed is called a coefficient . - We never place a subscript on atoms to balance, since subscripts would change the chemical composition of a reactant or product. 2 H 2(g) + O 2(g) 2 H 2 O (g) and not H 2(g) + O 2(g) H 2 O 2(g) * Note that the chemical composition of water has been changed so that we do not have water any more but another chemical called hydrogen peroxide.
9. calcium and chlorine yield calcium chloride Ca (s) + Cl 2(g) CaCl 2(s) magnesium and oxygen yield magnesium oxide Mg (s) + O 2(g) MgO (s) Mg (s) + O 2(g) 2 MgO (s) 2 Mg (s) + O 2(g) 2 MgO (s) aluminum and sulfur yield aluminum sulfide Al (s) + S 8(s) Al 2 S 3(s) Al (s) + S 8(s) 8 Al 2 S 3(s) 16 Al (s) + S 8(s) 8 Al 2 S 3(s) 16 Al (s) + 3 S 8(s) 8 Al 2 S 3(s) * 2) Word Equation Chemical Formula 3) Word Equation Chemical Formula 4) Word Equation Chemical Formula
10. Simple Decomposition - Simple decomposition is exactly the same as simple composition except that it is the reverse equation. - The generic equation can be stated as follows: compound element + element + … Examples 1) Potassium permanganate is decomposed. KMnO 4(s) K (s) + Mn (s) + O 2(g) KMnO 4(s) K (s) + Mn (s) + 2 O 2(g) 2) Ammonia is decomposed. NH 3(g) N 2(g) + H 2(g) NH 3(g) N 2(g) + 3 H 2(g) 2 NH 3(g) N 2(g) + 3 H 2(g) 3) Calcium carbonate is decomposed. CaCO 3(s) Ca (s) + C (s) + O 2(g) CaCO 3(s) Ca (s) + C (s) + 3 O 2(g) 2 CaCO 3(s) Ca (s) + C (s) + 3 O 2(g) *
11. 4) Sucrose is decomposed. C 12 H 22 O 11(s) C (s) + H 2(g) + O 2(g) C 12 H 22 O 11(s) 12 C (s) + H 2(g) + O 2(g) C 12 H 22 O 11(s) 12 C (s) + 11 H 2(g) + O 2(g) C 12 H 22 O 11(s) 12 C (s) + 11 H 2(g) + 5.5 O 2(g) 2 C 12 H 22 O 11(s) 24 C (s) + 22 H 2(g) + 11 O 2(g) *
12. Single Replacement - As the name implies, something that is single is replaced . The generic equation is: element + compound element new + compound new Na (s) Na (s) + LiCl (aq) Na (s) + LiCl (aq) Li (s) Na (s) + LiCl (aq) Li (s) + NaCl (aq) * Example 1) sodium metal and a lithium chloride solution yield lithium metal and a sodium chloride solution
13. How does this work? - Break up each atom to show what their ions look like. - Replace positive or negative atoms (switch) and reform new products. Na (s) + LiCl (aq) ? + ? Na + Na + Li + Cl Na + Li + Cl Li + Na + Cl Na (s) + Na (s) + LiCl (aq) Na (s) + LiCl (aq) Li (s) + Na (s) + LiCl (aq) Li (s) + NaCl (aq) *
14. Double Replacement - In double replacement reactions, two compounds are reacted to form two new compounds . - Generic formula: compound 1 + compound 2 compound 3 + compound 4 Examples 1) Sodium chloride solution and silver nitrate solutions are mixed. NaCl (aq) + NaCl (aq) + AgNO 3(aq) NaCl (aq) + AgNO 3(aq) ? + ? Na + Cl Na + Cl Ag + Na + Cl Ag + NO 3 Na + Cl Ag + NO 3 forms new products Na + Cl Ag + NO 3 Na + NO 3 + Na + Cl Ag + NO 3 Na + NO 3 + Ag + Cl - NaCl (aq) + AgNO 3(aq) NaNO 3(aq) + AgCl (s) * join join
15. 2) Solutions of lead (II) nitrate and sodium iodide are mixed. Pb ( NO 3 ) 2(aq) + Pb ( NO 3 ) 2(aq) + NaI (aq) Pb ( NO 3 ) 2(aq) + NaI (aq) ? + ? Pb 2+ NO 3 – + Pb 2+ NO 3 – + Na + I – Pb 2+ NO 3 – + Na + I – Pb 2+ I – + Pb 2+ NO 3 – + Na + I – Pb 2+ I – + Na + NO 3 – Pb ( NO 3 ) 2(aq) + Pb ( NO 3 ) 2(aq) + NaI (aq) Pb ( NO 3 ) 2(aq) + NaI (aq) PbI 2(s) + Pb ( NO 3 ) 2(aq) + NaI (aq) PbI 2(s) + NaNO 3(aq) Pb ( NO 3 ) 2(aq) + NaI (aq) PbI 2(s) + 2 NaNO 3(aq) Pb ( NO 3 ) 2(aq) + 2 NaI (aq) PbI 2(s) + 2 NaNO 3(aq) *
16. 3) Solutions of gallium chloride and hydrosulfuric acid are mixed. GaCl 3(aq) + GaCl 3(aq) + H 2 S (aq) ? + ? Ga 3+ Cl – + Ga 3+ Cl – + H + S 2– Ga 3+ Cl – + H + S 2– Ga 3+ S 2– + Ga 3+ Cl – + H + S 2– Ga 3+ S 2– + H + Cl – GaCl 3(aq) + GaCl 3(aq) + H 2 S (aq) GaCl 3(aq) + H 2 S (aq) Ga 2 S 3(aq) + GaCl 3(aq) + H 2 S (aq) Ga 2 S 3(aq) + HCl (aq) _ GaCl 3(aq) + _ H 2 S (aq) _ Ga 2 S 3(aq) + _ HCl (aq) _ GaCl 3(aq) + 3 H 2 S (aq) _ Ga 2 S 3(aq) + _ HCl (aq) _ GaCl 3(aq) + _ H 2 S (aq) _ Ga 2 S 3(aq) + 6 HCl (aq) 2 GaCl 3(aq) + _ H 2 S (aq) _ Ga 2 S 3(aq) + _ HCl (aq) _ GaCl 3(aq) + _ H 2 S (aq) 1 Ga 2 S 3(aq) + _ HCl (aq) *
17. 4) Solutions of calcium nitrate and sodium phosphate are mixed. Ca(NO 3 ) 2(aq) + Na 3 PO 4(aq) ? + ? Ca 2+ NO 3 – + Ca 2+ NO 3 – + Na + PO 4 3– Ca 2+ NO 3 – + Na + PO 4 3– Ca 2+ PO 4 3– + Ca 2+ NO 3 – + Na + PO 4 3– Ca 2+ PO 4 3– + Na + NO 3 – Ca ( NO 3 ) 2(aq) + Ca ( NO 3 ) 2(aq) + Na 3 PO 4(aq) Ca ( NO 3 ) 2(aq) + Na 3 PO 4(aq) Ca 3 ( PO 4 ) 2(aq) + Ca ( NO 3 ) 2(aq) + Na 3 PO 4(aq) Ca 3 ( PO 4 ) 2(aq) + NaNO 3(aq) _ Ca ( NO 3 ) 2(aq) + _ Na 3 PO 4(aq) _ Ca 3 ( PO 4 ) 2(s) + _ NaNO 3(aq) 3 Ca ( NO 3 ) 2(aq) + _ Na 3 PO 4(aq) _ Ca 3 ( PO 4 ) 2(s) + _ NaNO 3(aq) _ Ca ( NO 3 ) 2(aq) + _ Na 3 PO 4(aq) _ Ca 3 ( PO 4 ) 2(s) + 6 NaNO 3(aq) _ Ca ( NO 3 ) 2(aq) + 2 Na 3 PO 4(aq) _ Ca 3 ( PO 4 ) 2(s) + _ NaNO 3(aq) _ Ca ( NO 3 ) 2(aq) + _ Na 3 PO 4(aq) 1 Ca 3 ( PO 4 ) 2(s) + _ NaNO 3(aq) *
18. Hydrocarbon Combustion In hydrocarbon combustion, a fuel made up of at least hydrogen and carbon , is burned ( combustion ) in the presence of oxygen . Generic Formula: hydrocarbon + hydrocarbon + O 2(g) hydrocarbon + O 2(g) CO 2(g) + hydrocarbon + O 2(g) CO 2(g) + H 2 O (g) Examples 1) Methane is burned to heat our homes in winter. CH 4(g) + CH 4(g) + O 2(g) CH 4(g) + O 2(g) CO 2(g) + CH 4(g) + O 2(g) CO 2(g) + H 2 O (g) _ CH 4(g) + _ O 2(g) _ CO 2(g) + _ H 2 O (g) _ CH 4(g) + _ O 2(g) 1 CO 2(g) + _ H 2 O (g) _ CH 4(g) + _ O 2(g) _ CO 2(g) + 2 H 2 O (g) _ CH 4(g) + 2 O 2(g) _ CO 2(g) + _ H 2 O (g) *
19. 2) Gasoline is combusted in a car engine. C 8 H 18(l) + C 8 H 18(l) + O 2(g) CO 2(g) + H 2 O (g) _ C 8 H 18(l) + _ O 2(g) _ CO 2(g) + _ H 2 O (g) _ C 8 H 18(l) + _ O 2(g) 8 CO 2(g) + _ H 2 O (g) _ C 8 H 18(l) + _ O 2(g) _ CO 2(g) + 9 H 2 O (g) 2 C 8 H 18(l) + 25 O 2(g) 16 CO 2(g) + 18 H 2 O (g) 3) The burning of wood alcohol produces a flame. CH 3 OH (l) + CH 3 OH (l) + O 2(g) CO 2(g) + H 2 O (g) _ CH 3 OH (l) + _ O 2(g) _ CO 2(g) + _ H 2 O (g) _ CH 3 OH (l) + _ O 2(g) 1 CO 2(g) + _ H 2 O (g) _ CH 3 OH (l) + _ O 2(g) _ CO 2(g) + 2 H 2 O (g) 2 CH 3 OH (l) + 3 O 2(g) 2 CO 2(g) + 4 H 2 O (g) *
20. 4) We burn glucose in our body cells. C 6 H 12 O 6(s) + C 6 H 12 O 6(s) + O 2(g) CO 2(g) + H 2 O (g) _ C 6 H 12 O 6(s) + _ O 2(g) _ CO 2(g) + _ H 2 O (g) _ C 6 H 12 O 6(s) + _ O 2(g) 6 CO 2(g) + _ H 2 O (g) _ C 6 H 12 O 6(s) + _ O 2(g) _ CO 2(g) + 6 H 2 O (g) _ C 6 H 12 O 6(s) + 6 O 2(g) _ CO 2(g) + _ H 2 O (g) Other Some chemical reactions do not fit in any of the previously studied chemical reactions. These reactions will then be placed in a category called “ other ”. Examples 1) sulfur trioxide and water sulfuric acid SO 3(g) + SO 3(g) + H 2 O (l) SO 3(g) + H 2 O (l) H 2 SO 4(aq) *
21. 2) Nitrogen monoxide gas reacts with oxygen to produce nitrogen dioxide gas. NO (g) + NO (g) + O 2(g) NO (g) + O 2(g) NO 2(g) _ NO (g) + _ O 2(g) _ NO 2(g) 2 NO (g) + _ O 2(g) _ NO 2(g) _ NO (g) + _ O 2(g) 2 NO 2(g) 3) A sodium bicarbonate solution decomposes into a sodium carbonate solution, carbon dioxide and water. NaHCO 3(aq) NaHCO 3(aq) Na 2 CO 3(aq) + NaHCO 3(aq) Na 2 CO 3(aq) + CO 2(g) + NaHCO 3(aq) Na 2 CO 3(aq) + CO 2(g) + H 2 O (l) _ NaHCO 3(aq) _ Na 2 CO 3(aq) + _ CO 2(g) + _ H 2 O (l) 2 NaHCO 3(aq) _ Na 2 CO 3(aq) + _ CO 2(g) + _ H 2 O (l) *
22. Determining Whether a Substance Dissolves Easily In Water or Forms a Precipitate (Solid) - We are told that ionic substances are solids at room temperature and molecular substances can be any state depending on which ones we are talking about. - What happens if we try to dissolve them in water? Will they easily dissolve or will they tend not to dissolve, thereby remaining a solid ? - In chemical reactions where solutions are involved, it is necessary to know if a substance dissolves, so that we can place ( aq ) after its chemical formula. - To determine the “ solubility ” of a molecular substance we will need to memorize which ones easily dissolve and which ones don’t. - To determine the “ solubility ” of an ionic substance we will use the “ Solubility Table ” found at the bottom of the Periodic Table of Ions Sheet. *
23. - First, find the ions that make up the compound. [ Na + and Cl – ] - Next, locate one of those ions in the top row. [ Na + is a Gr IA ion ] - What does “ All ” and “ None ” mean? - Any Gr IA ion combined with “All” other negative ions will be “ Soluble ”, therefore we we put ( aq ) after that compound’s chemical formula. [ NaCl (aq) ] Which of the following ionic compounds will be soluble? a) ammonium hydroxide b) calcium phosphate c) potassium hydroxide d) strontium carbonate e) silver chloride NH 4 OH (aq) Ca 3 (PO 4 ) 2(s) KOH (aq) SrCO 3(s) AgCl (s) * How does this work? - Suppose you want to know if an ionic substance such as salt, NaCl (s) , dissolves in water.
24. _ NaI (aq) + _ NaI (aq) + _ Pb(NO 3 ) 2(aq) _ NaI (aq) + _ Pb(NO 3 ) 2(aq) _ NaNO 3(aq) + _ NaI (aq) + _ Pb(NO 3 ) 2(aq) _ NaNO 3(aq) + _ PbI 2(s) 2 NaI (aq) + _ Pb(NO 3 ) 2(aq) _ NaNO 3(aq) + _ PbI 2(s) _ NaI (aq) + _ Pb(NO 3 ) 2(aq) 2 NaNO 3(aq) + _ PbI 2(s) 2) Crystals of silver nitrate and sodium chloride are placed in separate beakers containing distilled water. Then their contents are mixed. Write out the balanced chemical equation. _ AgNO 3(aq) + _ AgNO 3(aq) + _ NaCl (aq) _ AgNO 3(aq) + _ NaCl (aq) _ AgCl (s) + _ AgNO 3(aq) + _ NaCl (aq) _ AgCl (s) + _ NaNO 3(aq) * Predict the states of matter for each of the following chemical reactions. 1) Crystals of sodium iodide and lead (II) nitrate are placed in separate beakers containing distilled water. Then their contents are mixed. Write out the balanced chemical equation.
25. Solubility Table For Ionic Substances * Ion H + Gr IA NH 4 + NO 3 – CH 3 COO – Cl – Br – I – SO 4 2– S 2– OH – PO 4 3 – SO 3 2 – CO 3 2 – Soluble (aq) All All All All All Most Most Gr IA Gr IIA NH 4 + Gr IA NH 4 + Sr 2+ Ba 2+ Gr IA NH 4 + Low Solubility (s) None None None None None Ag + Pb 2+ Hg + Cu + Ag + Pb 2+ Ca 2+ Ba 2+ Sr 2+ Most Most Most
26. The Mole Before we discuss the concept of the “mole”, it is important to be able to work with numbers. Scientific Notation Example: Convert 4638600000 into scientific notation 4.638600000 x 10 9 Example: Convert 0.00000475 into scientific notation 4.75 x 10 –6 Note: A positive exponent indicates a number greater than 1 and a negative exponent indicates a number smaller than 1. Example: Calculate the following. a) (4.50 x 10 3 ) (9.25 x 10 5 ) = b) (3.36 x 10 –7 ) (5.50 x 10 4 ) = 4.16 x 10 9 1.85 x 10 –2 9.38 c)
27. 6.92 x 10 –4 4.00 x 10 9 4.30 x 10 19 Significant Digits - The number 354 has _ significant digits. - The number 354 has 3 significant digits. - The number 35.4 has _ significant digits. - The number 35.4 has 3 significant digits. - The number 52 has _ significant digits. - The number 52 has 2 significant digits. - The number 3.54 x 10 4 has _ significant digits. - The number 3.54 x 10 4 has 3 significant digits. - The number 0.354 has _ significant digits. - The number 0.354 has 3 significant digits. - The number 0.00354 has _ significant digits. - The number 0.00354 has 3 significant digits. - The number 35400 has _ significant digits. - The number 35400 has 5 significant digits. - The number 0.03540 has _ significant digits. - The number 0.03540 has 4 significant digits. - The number 0.00350040 has _ significant digits. - The number 0.00350040 has 6 significant digits. d) e) = f) =
28. Rules For Multiplying and Dividing - When multiplying or dividing, the answer must contain the least number of significant digits that are found in the numbers being calculated . - Example: 3.3 x 0.134 = - The answer is 0.4422 but since the least number of significant digits being used is “ _ ”, the answer must be rounded off to “ _ ” digits. - The answer is 0.4422 but since the least number of significant digits being used is “ 2 ”, the answer must be rounded off to “ 2 ” digits. Answer: Answer: 0.44 - Example: 3746 x 0.120 = - The calculator answer is 449.52 , but since the least number of significant digits is “ _ ”, the answer is rounded off to _ digits. - The calculator answer is 449.52 , but since the least number of significant digits is “ 3 ”, the answer is rounded off to 3 digits. - Example: 3746 x 120 = - The calculator answer is 449520 , but since the least number of significant digits is “ _ ”, the answer is rounded off to _ digits. - The calculator answer is 449520 , but since the least number of significant digits is “ 3 ”, the answer is rounded off to 3 digits. - The answer must be about 450000, but that number has _ digits. - The answer must be about 450000, but that number has 6 digits. - In this case, use scientific notation . Answer: Answer: 4.50 x 10 5 Answer: Answer: 450
29. Rules for Adding and Subtracting - Perform the calculation in your calculator. - Determine the least number of place holding there are after the decimal in each entry . - Your answer must contain the least number of place holdings after the decimal . - Round off the answer. Example: 34.65 + 4.235 = - Calculator answer is - Calculator answer is 38.885 . - The least number of place holding after the decimal is “ _ ”. - The least number of place holding after the decimal is “ 2 ”. Answer: Answer: 38.89 - Example: 1356.245 + 245.33 – 0.0001 = 1601.5749 = 1601.5749 = 1601.57 Note: When calculating long series of numbers, round off only at the end .
30. The Mole - Atoms are extremely small so it becomes very difficult to make measurements when only a few atoms or molecules are used. - If we wanted to take the mass of 1 atom, it would be appropriate to find the mass of say, 1000 atoms and then divide by 1000, but 1000 atoms still gives us a very small mass to measure . - What number of atoms or molecules would be large enough? 1 dozen? 1 gross (144)? 1 million? - In the early 1800’s, a scientist named Amedeo Avogadro devised a way of measuring the masses of very small particles of matter. - Avogadro suggested that if you took exactly 12.00 g of the carbon-12 atom you would have a very large number of carbon atoms. Just as words like “dozen”, “gross” or “couple” have numerical meaning, he needed a word to describe how many atoms there would be in 12.00 g of carbon-12. - This word is now known as the “ mole ”.
31. - One dozen atoms would be 12 atoms. - One gross of atoms would be 144 atoms. - One mole of atoms is 6.02 x 10 23 atoms , now known as Avogadro’s number .
32. The Green Pea Analogy (from Alchem 2000) - One hundred (10 2 ) green peas have a volume of about 20 cm 3 or 20 mL in a graduated cylinder. - One million (10 6 ) green peas have enough volume to fill an ordinary refrigerator. - One billion (10 9 ) green peas can fill an average 3 bedroom house. - One trillion (10 12 ) green peas can fill 1000 average homes. - One quadrillion (10 15 ) green peas can fill all the building in a larger city such as Edmonton. - Imagine all of Alberta covered one metre deep in green peas. Now you got one quintillion (10 18 ) green peas. - Imagine that all the continents are now covered 1 metre deep with green peas. Now you have one sextillion (10 21 ). - If you froze all the oceans and totally covered the whole earth with green peas and had 250 planets just like earth, all covered with green peas, you would now have a mole’s worth ! - 250 000 planets like earth, all covered one metre deep in green peas would give you a cotillion (10 27 ), the number of atoms in your body!
33.
34. Calculating Number of Particles In Given Number of Moles Examples: 1. How many atoms are there in 2.50 mol of lithium? Since 1 mol of Li (s) = Since 1 mol of Li (s) = 6.02 x 10 23 atoms Since 1 mol of Li (s) = 6.02 x 10 23 atoms then 2.50 mol of Li (s) should have 2.50 mol x 6.02 x 10 23 atoms = Since 1 mol of Li (s) = 6.02 x 10 23 atoms then 2.50 mol of Li (s) should have 2.50 mol x 6.02 x 10 23 atoms = 1.51 x 10 24 atoms or
35. 2. How many molecules are there in 25.0 mol of water? 3. How many moles are there in 4.55 x 10 15 atoms of zinc?
36. 4. How many moles are there in 2.75 x 10 24 molecules of ammonia?
37. Calculating the Molar Mass of Any Substance Find the molar masses for each of the following examples: 1) Cu (s) Since the formula suggests that copper is made up only one type of atom, we simply look up the molar mass value from the periodic table. Cu (s) = Cu (s) = 63.55 g in one mole 2) H 2 O (l) There are 2 hydrogens and 1 oxygen involved in making the water molecule we need to add the molar masses of all atoms. 2 – H = 2 – H = 2 x 1.01 = 2 – H = 2 x 1.01 = 2.02 g 1 – O = 1 – O = 1 x 16.00 = 1 – O = 1 x 16.00 = 16.00 g H 2 O (l) = = 63.55 H 2 O (l) = 18.02
38. 3) CaSO 4(s) 1 – Ca = 1 – Ca = 1 x 40.08 = 1 – Ca = 1 x 40.08 = 40.08 g 1 – S = 1 – S = 1 x 32.06 = 1 – S = 1 x 32.06 = 32.06 g 4 – O = 4 – O = 4 x 16.00 = 4 – O = 4 x 16.00 = 64.00 g CaSO 4(s) = 4) Al(OH) 3(s) 1 – Al = 1 – Al = 1 x 26.98 = 1 – Al = 1 x 26.98 = 26.98 g 3 – O = 3 – O = 3 x 16.00 = 3 – O = 3 x 16.00 = 48.00 g 3 – H = 3 – H = 3 x 1.01 = 3 – H = 3 x 1.01 = 3.03 g Al(OH) 3(s) = CaSO 4(s) = 136.14 Al(OH) 3(s) = 78.01
39. 5) Na 2 H 2 PO 4(s) 2 – Na = 2 – Na = 2 x 22.99 = 2 – Na = 2 x 22.99 = 45.98 g 2 – H = 2 – H = 2 x 1.01 = 2 – H = 2 x 1.01 = 2.02 g 1 – P = 1 – P = 1 x 30.97 = 1 – P = 1 x 30.97 = 30.97 g 4 – O = 4 – O = 4 x 16.00 = 4 – O = 4 x 16.00 = 64.00 g Na 2 H 2 PO 4(s) = Na 2 H 2 PO 4(s) = 142.97
40. Calculations Converting Mass to Moles We can use a simple formula for this type of calculation: - molar mass of water = Examples 1. How many moles are there in 75.0 g of water?
41. 2. How many moles are there in 100 g of sucrose? - molar mass of sucrose = Calculations Converting Moles to Mass We can use a simple formula for this type of calculation:
42. 2. What is the mass of 25.0 mol of ammonia? Examples 1. What mass do you have in 2.83 mol of table salt?