Okay, here are the steps:
1) Convert the mass percentages to grams of each element in 100 g of the compound:
K: 24.75% = 24.75 g
Mn: 34.77% = 34.77 g
O: 40.51% = 40.51 g
2) Calculate the moles of each element:
K: 24.75 g / 39.10 g/mol (molar mass of K) = 0.634 mol
Mn: 34.77 g / 54.94 g/mol (molar mass of Mn) = 0.634 mol
O: 40.51 g / 16.00 g/mol (molar mass of O)
C03 relative masses of atoms and moleculesdean dundas
This document discusses relative atomic and molecular masses. It defines relative atomic mass as the average mass of an atom compared to 1/12 the mass of one carbon-12 atom. Relative molecular mass is defined similarly as the average mass of a molecule compared to 1/12 the mass of one carbon-12 atom. The document provides examples of calculating relative atomic masses from the periodic table and relative molecular masses by adding atomic masses. It also discusses calculating the percentage composition and purity of compounds.
This chapter discusses chemical formulae and equations. It defines relative atomic mass and relative molecular mass, and explains how to calculate these values. It also describes the mole concept and relationships between number of moles, particles, and mass or volume. It discusses writing chemical formulae for elements, compounds, ions and ionic compounds. The chapter also explains how to determine empirical and molecular formulae, and how to balance chemical equations.
The document defines atomic mass, molar mass, molecular mass, and formula mass. It explains that atomic mass is the mass of an atom measured in atomic mass units (amu) and is based on carbon-12. Molar mass is the mass of one mole of a substance in grams. One mole contains 6.022x1023 elementary units. Molecular mass is the sum of atomic masses in a molecule. Formula mass is the sum of atomic masses in a formula unit of an ionic compound. The document provides examples of calculating molar mass, molecular mass, and formula mass from atomic masses on the periodic table.
This document discusses mass relationships in chemical reactions, including:
1) Atomic mass, molecular mass, molar mass, and formula mass. It defines the mole and Avogadro's number.
2) Chemical equations and how they are used to represent chemical reactions by balancing the atoms on each side.
3) Calculations involving the amounts of reactants and products in chemical reactions, including limiting reagents.
The document provides information about atoms and their structure. It defines key terms like protons, neutrons, electrons, nucleus and isotopes. It explains that the number of protons determines the element and distinguishes one atom from another. The mole is also defined as 6.02x10^23 particles and is used to measure amounts of substances on a macroscopic scale. Formulas are given to calculate molar mass and empirical formulas.
The correct formula for an ionic compound must contain positive and negative ions in a ratio to achieve an overall neutral charge. CO2 is a molecular compound and does not contain ions, so the answer is A.
Okay, here are the steps:
1) Convert the mass percentages to grams of each element in 100 g of the compound:
K: 24.75% = 24.75 g
Mn: 34.77% = 34.77 g
O: 40.51% = 40.51 g
2) Calculate the moles of each element:
K: 24.75 g / 39.10 g/mol (molar mass of K) = 0.634 mol
Mn: 34.77 g / 54.94 g/mol (molar mass of Mn) = 0.634 mol
O: 40.51 g / 16.00 g/mol (molar mass of O)
C03 relative masses of atoms and moleculesdean dundas
This document discusses relative atomic and molecular masses. It defines relative atomic mass as the average mass of an atom compared to 1/12 the mass of one carbon-12 atom. Relative molecular mass is defined similarly as the average mass of a molecule compared to 1/12 the mass of one carbon-12 atom. The document provides examples of calculating relative atomic masses from the periodic table and relative molecular masses by adding atomic masses. It also discusses calculating the percentage composition and purity of compounds.
This chapter discusses chemical formulae and equations. It defines relative atomic mass and relative molecular mass, and explains how to calculate these values. It also describes the mole concept and relationships between number of moles, particles, and mass or volume. It discusses writing chemical formulae for elements, compounds, ions and ionic compounds. The chapter also explains how to determine empirical and molecular formulae, and how to balance chemical equations.
The document defines atomic mass, molar mass, molecular mass, and formula mass. It explains that atomic mass is the mass of an atom measured in atomic mass units (amu) and is based on carbon-12. Molar mass is the mass of one mole of a substance in grams. One mole contains 6.022x1023 elementary units. Molecular mass is the sum of atomic masses in a molecule. Formula mass is the sum of atomic masses in a formula unit of an ionic compound. The document provides examples of calculating molar mass, molecular mass, and formula mass from atomic masses on the periodic table.
This document discusses mass relationships in chemical reactions, including:
1) Atomic mass, molecular mass, molar mass, and formula mass. It defines the mole and Avogadro's number.
2) Chemical equations and how they are used to represent chemical reactions by balancing the atoms on each side.
3) Calculations involving the amounts of reactants and products in chemical reactions, including limiting reagents.
The document provides information about atoms and their structure. It defines key terms like protons, neutrons, electrons, nucleus and isotopes. It explains that the number of protons determines the element and distinguishes one atom from another. The mole is also defined as 6.02x10^23 particles and is used to measure amounts of substances on a macroscopic scale. Formulas are given to calculate molar mass and empirical formulas.
The correct formula for an ionic compound must contain positive and negative ions in a ratio to achieve an overall neutral charge. CO2 is a molecular compound and does not contain ions, so the answer is A.
This document discusses chemical formulae and equations. It defines relative atomic mass and relative molecular mass, which are used to calculate the mass of elements and compounds from their chemical formulae. The mole concept is introduced, relating the Avogadro constant to the number of particles in a given number of moles. Relationships are shown between moles, mass, particles and volume. Empirical and molecular formulae are distinguished. Ionic compounds have formulae showing cation and anion combinations. Examples of writing and balancing chemical equations are provided.
This document provides an overview of key concepts related to the mole concept in chemistry. It defines the mole as the number of atoms or molecules in 1 gram of hydrogen or 12 grams of carbon. The mole concept allows chemists to relate mass, number of particles, and volume of gases. It discusses how to calculate empirical and molecular formulas, Avogadro's constant, molar mass, limiting reactants, and other mole-related calculations and applications. Worked examples are provided to demonstrate how to use the mole concept to find formulas of compounds from percentage composition data and other information.
The document discusses organic chemistry topics including:
- Classes of organic compounds such as aliphatic hydrocarbons, aromatic hydrocarbons, and functional groups.
- Alkanes have the general formula CnH2n+2 and contain only single bonds. Cycloalkanes contain carbon atoms joined in rings.
- Alkenes contain carbon-carbon double bonds and have the general formula CnH2n. Alkynes contain carbon-carbon triple bonds and have the formula CnH2n-2.
- Aromatic compounds contain benzene rings, and their naming involves indicating substituted groups on the ring.
This document discusses the mole concept in chemistry. It defines the mole as the amount of substance containing 6.02x1023 particles. A mole of any substance has a mass in grams equal to its molar mass. The document explains how to determine empirical and molecular formulas from percentage composition data using mole calculations. It also discusses limiting reactants and using moles to calculate gas volumes based on Avogadro's Law. Several examples are provided to demonstrate determining formulas from mass or molar mass data.
1) The document discusses chemical formulas and equations, including relative atomic mass (RAM), relative molecular mass (RMM), moles, and molar mass.
2) RAM is the average mass of an atom of an element compared to 1/12 the mass of one carbon-12 atom. RMM is the average mass of a molecule compared to 1/12 the mass of one carbon-12 atom.
3) A mole is defined as the amount of substance containing as many elementary entities (atoms, molecules, ions, etc) as there are atoms in exactly 12 grams of carbon-12, which is a quantity known as Avogadro's number (6.02x1023).
The document provides information about chemical formulas and equations. It defines empirical and molecular formulas, and explains how to determine them through calculation of moles and mass ratios of elements in a compound. It also describes writing and balancing chemical equations, naming ionic compounds based on their constituent ions, and using formulas and equations to solve stoichiometric problems. Key topics covered include determining formulas from experimental data, relating formulas to molecular structure and mass, and representing chemical reactions systematically.
This document discusses chemical equations and reaction stoichiometry. It begins by defining chemical equations and explaining what information they provide, such as reactants, products, and relative quantities. It then discusses the law of conservation of matter and provides examples of balancing chemical equations. The document concludes by explaining how to perform calculations based on chemical equations, such as determining quantities in moles, grams, or other units using mole-to-mole conversions. It provides several examples of these types of stoichiometry calculations.
The document discusses the mole concept in chemistry. Some key points:
- A mole is equal to 6.022x10^23 particles and can refer to atoms, molecules, etc.
- The molar mass of an element is its atomic mass in grams and the molar mass of a compound is the sum of the atomic masses of its elements.
- One mole of an ideal gas occupies 22.4 liters at STP.
- Questions at the end calculate things like moles, mass, volume using molar mass and mole ratios in chemical equations.
This document provides an overview of chemistry concepts including:
1. Chemistry is the study of matter and the changes it undergoes. The scientific method uses a systematic approach involving hypotheses, experiments, and analysis.
2. Matter can exist as elements, compounds, mixtures, and in three main states - solids, liquids, and gases. Chemical and physical changes alter substances in different ways.
3. The study of chemistry incorporates macroscopic observations and measurements as well as analysis at the microscopic level of atoms and molecules. Significant figures, units, and mathematical representations are important tools in chemistry.
Gases exist as individual molecules that are in constant random motion. The kinetic molecular theory describes gases as composed of molecules that are separated by large distances and move rapidly in random directions, frequently colliding with one another. The theory states that the average kinetic energy of gas molecules is proportional to the absolute temperature of the gas. Higher temperatures cause molecules to move faster on average with more molecules possessing higher speeds.
The document discusses chemical formulae and how they are used to represent elements and compounds. It provides examples of how to determine the empirical and molecular formula of substances based on experimental data like mass percentages and ratios of elements. Empirical formulae show the simplest whole number ratio of atoms in a compound, while molecular formulae show the actual number of each atom in a molecule of a substance. Percentage composition can also be calculated from the relative atomic masses and molecular formula.
This document discusses aqueous solutions and their properties. It defines key terms including solute, solvent, solution, electrolyte, and nonelectrolyte. It explains that solutions can be solid, liquid, or gas and describes different types of aqueous solutions. Common examples like sea water, vinegar, and sugar water are provided. The document also discusses solubility, dissociation, hydration, and precipitation reactions.
The document discusses different types of stoichiometry calculations encountered in AS Organic Chemistry, including:
1) Percentage yield calculations from chemical reactions.
2) Determining empirical formulas from elemental analysis data involving masses, percentages, or combustion analysis.
3) Calculating molecular formulas using empirical formulas and experimentally determined molar masses.
This chapter discusses the mole concept, including defining the mole, deriving empirical and molecular formulas, stating Avogadro's Law, and applying the mole concept to ionic and molecular equations. It introduces the mole as the amount of substance containing 6x1023 particles. It provides examples of how to determine the empirical formula, molecular formula, and formula of a compound from composition data. It also discusses molar volume of gases and limiting reactants. Worked examples are included for many of these concepts.
This document discusses stoichiometry, which is the quantitative relationships between reactants and products in chemical reactions. Some key points covered include:
- Stoichiometric amounts refer to the exact molar amounts of reactants and products in a balanced chemical equation.
- Atomic masses are given in atomic mass units (amu) which is based on carbon-12 having a mass of exactly 12 amu.
- Average atomic masses take into account the natural abundances of isotopes.
- The mole is the unit used to express amounts of substances in chemistry and 1 mole contains 6.022x1023 elementary entities.
- Molar mass is the mass in grams of 1 mole of a substance and is equal to the sum of the
1. A mole is defined as the amount of substance containing 6.02 x 1023 particles and can refer to atoms, molecules, or ions.
2. The mole is the unit for quantifying amount of substance, with the symbol "mol".
3. The number of particles in a given number of moles can be calculated by multiplying the moles by Avogadro's number, and the moles for a given number of particles is calculated by dividing the particles by Avogadro's number.
This document provides an overview of stoichiometry concepts including:
- Stoichiometry deals with quantitative relationships in chemical reactions and compositions.
- Reaction stoichiometry uses mole ratios from balanced equations to convert between amounts of reactants and products.
- There are four main types of stoichiometry problems depending on what is given and unknown.
- Limiting reactants determine the maximum amount of product that can be formed in a reaction.
- Percent yield compares the actual product amount to the theoretical maximum.
This document summarizes key concepts relating to gases, including:
- The gas laws of Boyle, Charles, Avogadro, and combined gas law and how they relate pressure, volume, temperature, and moles of gas
- The ideal gas equation and how it incorporates these variables
- Standard temperature and pressure conditions for reporting gas measurements
- Gas density calculations using molar mass and gas stoichiometry problems
- Dalton's law of partial pressures which states that in a mixture of gases, the total pressure is equal to the sum of the partial pressures of the individual gases.
The document discusses atomic structure and the building blocks of matter. It defines key terms like atoms, subatomic particles, isotopes, and moles. It explains that atoms of the same element are distinguished by their number of protons. The mole is introduced as a unit used to measure amount of substance equal to 6.02x1023 particles. Examples are given for calculating molar mass and empirical formulas from elemental composition data.
This document provides an introduction and table of contents for a chemistry course book on Cambridge International AS and A Level Chemistry. It covers topics like the mass of atoms and molecules, relative atomic masses, isotopic masses, amount of substance, mole calculation, chemical formulae, solutions, gas volume calculations, and more. The document gives definitions and examples for these concepts. It also provides sample problems and homework questions related to chemical calculations involving moles, masses, and chemical equations.
This document provides an outline and objectives for a unit on formulas and equations. The unit covers calculating atomic mass, the mole concept including molar mass and conversions between moles and mass, determining empirical and molecular formulas through combustion analysis, and stoichiometry including writing and balancing chemical equations, limiting reactants, theoretical and percent yields. Example problems are provided to illustrate key concepts like calculating atomic mass, determining moles of atoms from mass, finding empirical and molecular formulas, and stoichiometry calculations.
This document provides an overview of stoichiometry concepts including:
- Calculating molecular and molar masses
- Determining percent composition of compounds
- Balancing chemical equations
- Using the mole concept to convert between masses, moles, and numbers of particles
- Calculating limiting reactants and determining amounts of excess reactants
The document uses examples and practice problems to illustrate these stoichiometry concepts and calculations.
This document discusses chemical formulae and equations. It defines relative atomic mass and relative molecular mass, which are used to calculate the mass of elements and compounds from their chemical formulae. The mole concept is introduced, relating the Avogadro constant to the number of particles in a given number of moles. Relationships are shown between moles, mass, particles and volume. Empirical and molecular formulae are distinguished. Ionic compounds have formulae showing cation and anion combinations. Examples of writing and balancing chemical equations are provided.
This document provides an overview of key concepts related to the mole concept in chemistry. It defines the mole as the number of atoms or molecules in 1 gram of hydrogen or 12 grams of carbon. The mole concept allows chemists to relate mass, number of particles, and volume of gases. It discusses how to calculate empirical and molecular formulas, Avogadro's constant, molar mass, limiting reactants, and other mole-related calculations and applications. Worked examples are provided to demonstrate how to use the mole concept to find formulas of compounds from percentage composition data and other information.
The document discusses organic chemistry topics including:
- Classes of organic compounds such as aliphatic hydrocarbons, aromatic hydrocarbons, and functional groups.
- Alkanes have the general formula CnH2n+2 and contain only single bonds. Cycloalkanes contain carbon atoms joined in rings.
- Alkenes contain carbon-carbon double bonds and have the general formula CnH2n. Alkynes contain carbon-carbon triple bonds and have the formula CnH2n-2.
- Aromatic compounds contain benzene rings, and their naming involves indicating substituted groups on the ring.
This document discusses the mole concept in chemistry. It defines the mole as the amount of substance containing 6.02x1023 particles. A mole of any substance has a mass in grams equal to its molar mass. The document explains how to determine empirical and molecular formulas from percentage composition data using mole calculations. It also discusses limiting reactants and using moles to calculate gas volumes based on Avogadro's Law. Several examples are provided to demonstrate determining formulas from mass or molar mass data.
1) The document discusses chemical formulas and equations, including relative atomic mass (RAM), relative molecular mass (RMM), moles, and molar mass.
2) RAM is the average mass of an atom of an element compared to 1/12 the mass of one carbon-12 atom. RMM is the average mass of a molecule compared to 1/12 the mass of one carbon-12 atom.
3) A mole is defined as the amount of substance containing as many elementary entities (atoms, molecules, ions, etc) as there are atoms in exactly 12 grams of carbon-12, which is a quantity known as Avogadro's number (6.02x1023).
The document provides information about chemical formulas and equations. It defines empirical and molecular formulas, and explains how to determine them through calculation of moles and mass ratios of elements in a compound. It also describes writing and balancing chemical equations, naming ionic compounds based on their constituent ions, and using formulas and equations to solve stoichiometric problems. Key topics covered include determining formulas from experimental data, relating formulas to molecular structure and mass, and representing chemical reactions systematically.
This document discusses chemical equations and reaction stoichiometry. It begins by defining chemical equations and explaining what information they provide, such as reactants, products, and relative quantities. It then discusses the law of conservation of matter and provides examples of balancing chemical equations. The document concludes by explaining how to perform calculations based on chemical equations, such as determining quantities in moles, grams, or other units using mole-to-mole conversions. It provides several examples of these types of stoichiometry calculations.
The document discusses the mole concept in chemistry. Some key points:
- A mole is equal to 6.022x10^23 particles and can refer to atoms, molecules, etc.
- The molar mass of an element is its atomic mass in grams and the molar mass of a compound is the sum of the atomic masses of its elements.
- One mole of an ideal gas occupies 22.4 liters at STP.
- Questions at the end calculate things like moles, mass, volume using molar mass and mole ratios in chemical equations.
This document provides an overview of chemistry concepts including:
1. Chemistry is the study of matter and the changes it undergoes. The scientific method uses a systematic approach involving hypotheses, experiments, and analysis.
2. Matter can exist as elements, compounds, mixtures, and in three main states - solids, liquids, and gases. Chemical and physical changes alter substances in different ways.
3. The study of chemistry incorporates macroscopic observations and measurements as well as analysis at the microscopic level of atoms and molecules. Significant figures, units, and mathematical representations are important tools in chemistry.
Gases exist as individual molecules that are in constant random motion. The kinetic molecular theory describes gases as composed of molecules that are separated by large distances and move rapidly in random directions, frequently colliding with one another. The theory states that the average kinetic energy of gas molecules is proportional to the absolute temperature of the gas. Higher temperatures cause molecules to move faster on average with more molecules possessing higher speeds.
The document discusses chemical formulae and how they are used to represent elements and compounds. It provides examples of how to determine the empirical and molecular formula of substances based on experimental data like mass percentages and ratios of elements. Empirical formulae show the simplest whole number ratio of atoms in a compound, while molecular formulae show the actual number of each atom in a molecule of a substance. Percentage composition can also be calculated from the relative atomic masses and molecular formula.
This document discusses aqueous solutions and their properties. It defines key terms including solute, solvent, solution, electrolyte, and nonelectrolyte. It explains that solutions can be solid, liquid, or gas and describes different types of aqueous solutions. Common examples like sea water, vinegar, and sugar water are provided. The document also discusses solubility, dissociation, hydration, and precipitation reactions.
The document discusses different types of stoichiometry calculations encountered in AS Organic Chemistry, including:
1) Percentage yield calculations from chemical reactions.
2) Determining empirical formulas from elemental analysis data involving masses, percentages, or combustion analysis.
3) Calculating molecular formulas using empirical formulas and experimentally determined molar masses.
This chapter discusses the mole concept, including defining the mole, deriving empirical and molecular formulas, stating Avogadro's Law, and applying the mole concept to ionic and molecular equations. It introduces the mole as the amount of substance containing 6x1023 particles. It provides examples of how to determine the empirical formula, molecular formula, and formula of a compound from composition data. It also discusses molar volume of gases and limiting reactants. Worked examples are included for many of these concepts.
This document discusses stoichiometry, which is the quantitative relationships between reactants and products in chemical reactions. Some key points covered include:
- Stoichiometric amounts refer to the exact molar amounts of reactants and products in a balanced chemical equation.
- Atomic masses are given in atomic mass units (amu) which is based on carbon-12 having a mass of exactly 12 amu.
- Average atomic masses take into account the natural abundances of isotopes.
- The mole is the unit used to express amounts of substances in chemistry and 1 mole contains 6.022x1023 elementary entities.
- Molar mass is the mass in grams of 1 mole of a substance and is equal to the sum of the
1. A mole is defined as the amount of substance containing 6.02 x 1023 particles and can refer to atoms, molecules, or ions.
2. The mole is the unit for quantifying amount of substance, with the symbol "mol".
3. The number of particles in a given number of moles can be calculated by multiplying the moles by Avogadro's number, and the moles for a given number of particles is calculated by dividing the particles by Avogadro's number.
This document provides an overview of stoichiometry concepts including:
- Stoichiometry deals with quantitative relationships in chemical reactions and compositions.
- Reaction stoichiometry uses mole ratios from balanced equations to convert between amounts of reactants and products.
- There are four main types of stoichiometry problems depending on what is given and unknown.
- Limiting reactants determine the maximum amount of product that can be formed in a reaction.
- Percent yield compares the actual product amount to the theoretical maximum.
This document summarizes key concepts relating to gases, including:
- The gas laws of Boyle, Charles, Avogadro, and combined gas law and how they relate pressure, volume, temperature, and moles of gas
- The ideal gas equation and how it incorporates these variables
- Standard temperature and pressure conditions for reporting gas measurements
- Gas density calculations using molar mass and gas stoichiometry problems
- Dalton's law of partial pressures which states that in a mixture of gases, the total pressure is equal to the sum of the partial pressures of the individual gases.
The document discusses atomic structure and the building blocks of matter. It defines key terms like atoms, subatomic particles, isotopes, and moles. It explains that atoms of the same element are distinguished by their number of protons. The mole is introduced as a unit used to measure amount of substance equal to 6.02x1023 particles. Examples are given for calculating molar mass and empirical formulas from elemental composition data.
This document provides an introduction and table of contents for a chemistry course book on Cambridge International AS and A Level Chemistry. It covers topics like the mass of atoms and molecules, relative atomic masses, isotopic masses, amount of substance, mole calculation, chemical formulae, solutions, gas volume calculations, and more. The document gives definitions and examples for these concepts. It also provides sample problems and homework questions related to chemical calculations involving moles, masses, and chemical equations.
This document provides an outline and objectives for a unit on formulas and equations. The unit covers calculating atomic mass, the mole concept including molar mass and conversions between moles and mass, determining empirical and molecular formulas through combustion analysis, and stoichiometry including writing and balancing chemical equations, limiting reactants, theoretical and percent yields. Example problems are provided to illustrate key concepts like calculating atomic mass, determining moles of atoms from mass, finding empirical and molecular formulas, and stoichiometry calculations.
This document provides an overview of stoichiometry concepts including:
- Calculating molecular and molar masses
- Determining percent composition of compounds
- Balancing chemical equations
- Using the mole concept to convert between masses, moles, and numbers of particles
- Calculating limiting reactants and determining amounts of excess reactants
The document uses examples and practice problems to illustrate these stoichiometry concepts and calculations.
This document provides an overview of chemical reactions, including the key parts of a chemical equation, types of reactions such as synthesis, decomposition, and combustion reactions, and how to write and balance chemical equations. It also discusses writing total ionic and net ionic equations by identifying soluble and insoluble reactants and products. The document aims to teach students the essential concepts and skills for understanding and working with chemical reactions.
This document provides an overview of chemical reactions and equations. It begins with defining chemical reactions and their components like reactants, products, and coefficients. It then covers the parts of a chemical equation like physical states. Several types of chemical reactions are described including synthesis, decomposition, single displacement, double displacement, and combustion reactions. The document explains how to write and balance chemical equations as well as identify reaction types. It introduces total ionic and net ionic equations and discusses solubility. The overall summary is that this document outlines the key concepts for writing and classifying chemical equations.
This document provides an overview of basic chemistry concepts and calculations. It begins by explaining that chemistry is involved in everyday life through materials like polymers, drugs, and alloys. Matter can be classified based on physical state as solid, liquid, or gas, and chemically as pure substances or mixtures. Elements are composed of single atom types, while compounds contain two or more different atom types bonded together. The mole concept allows chemists to quantify amounts of substances down to the molecular level. Chemical calculations can be performed based on molar mass, stoichiometry from balanced equations, and other relationships between amounts of reactants and products.
This document provides an overview of basic chemistry concepts and calculations. It begins by explaining that chemistry is involved in everyday life through materials like polymers, drugs, and alloys. Matter can be classified based on physical state as solid, liquid, or gas, and chemically as pure substances or mixtures. Elements are composed of single atom types that can be monatomic or polyatomic. Compounds contain two or more elements. The mole concept is introduced to quantify atoms and molecules using Avogadro's number. Empirical and molecular formulas are distinguished. Methods for determining formulas from elemental analysis or molar mass are presented. Equivalent mass and stoichiometric calculations are also covered.
This document provides an overview of basic chemistry concepts and calculations. It begins by explaining that chemistry is involved in everyday life through materials like polymers, drugs, and alloys. Matter can be classified based on physical state as solid, liquid, or gas, and chemically as pure substances or mixtures. Elements are composed of single atom types that can be monatomic or polyatomic. Compounds contain two or more elements. The mole concept is introduced to quantify atoms and molecules using Avogadro's number of 6.022x1023 particles. Empirical and molecular formulas are distinguished. Methods are described for determining empirical formulas from elemental analysis data and calculating molecular formulas. Stoichiometry allows determining quantitative relationships between reactants and products in chemical equations
This document discusses the mole concept and stoichiometry calculations. It defines key terms like the mole, molar mass, empirical and molecular formulas, and limiting reagents. It also describes how to use balanced chemical equations and molar relationships to calculate amounts of reactants and products involved in chemical reactions. Experimental procedures are provided for determining relative molecular masses and identifying unknown metals based on stoichiometric calculations.
C03 relative masses of atoms and moleculesChemrcwss
The document discusses relative atomic mass and relative molecular mass. It defines relative atomic mass as the average mass of an atom compared to 1/12 the mass of one carbon-12 atom. Relative molecular mass is defined similarly on a molecular level. Examples are provided for calculating relative atomic masses from the periodic table and relative molecular masses by adding atomic masses. Percentage composition, yield, and purity calculations involving relative masses are also illustrated.
- The mole is used to indirectly count particles of matter and relates to Avogadro's number of 6.022x10^23 particles.
- A mole of any pure substance, whether an element or compound, has a mass in grams equal to its molar mass.
- Molar mass can be used to convert between moles of a substance and its mass in grams.
This document discusses relative atomic mass and relative molecular mass. It defines these terms and explains how they are calculated by comparing the mass of an atom or molecule to 1/12 the mass of one carbon-12 atom. The key points covered are:
- Relative atomic mass is the average mass of an atom of an element compared to 1/12 the mass of one carbon-12 atom.
- Relative molecular mass is the average mass of a molecule compared to 1/12 the mass of one carbon-12 atom. It is calculated by adding the relative atomic masses of the atoms in the molecule.
- Examples are provided to demonstrate calculating relative atomic masses from the periodic table and relative molecular masses by adding atomic masses
Chemical reactions involve the breaking and forming of chemical bonds to create new substances. There are several types of chemical reactions including synthesis, decomposition, and combustion. Chemical equations are used to represent reactions and must be balanced to show that the same number of atoms are on both sides. Reactions can be exothermic, releasing energy as heat, or endothermic, absorbing energy from their surroundings. Living things and industry rely on chemical reactions like photosynthesis and respiration.
Topic 1.3 chemical reactions and related calculationsJimiCarter
This document provides information on chemical reactions and calculations. It defines chemical reactions as when elements join to form compounds. It discusses relative formula mass and how to calculate it by summing atomic masses. Examples of relative formula mass calculations are provided. The document also discusses moles, Avogadro's number, and using molar mass to convert between moles and mass. Balancing chemical equations and identifying common ions are explained. Methods for calculating percentage composition and reacting masses are presented. Irreversible and reversible reactions are defined. Revision materials are suggested on relevant topics.
Chemistry zimsec chapter 2 atoms, molecules and stoichiometryalproelearning
This document provides an overview of Chapter 2 in a chemistry textbook, which covers topics including:
- The mass of atoms and molecules, including relative atomic mass and molecular mass
- Using a mass spectrometer to determine relative isotopic masses and abundances
- The mole concept and amount of substance in relation to mass, volume of gases, and concentration of solutions
- Calculating empirical formulas from combustion data or elemental composition by mass and deducing molecular formulas
- Stoichiometry, including writing balanced chemical equations and ionic equations
The document defines relative atomic mass and relative molecular mass, and explains how they are used to calculate the average mass of atoms and molecules compared to carbon-12. It provides examples of calculating relative atomic masses from the periodic table, and relative molecular masses by adding atomic masses. The document also discusses calculating percentage composition and yield of elements in compounds.
Chemical Reactions and Reaction StoichiometryPatriciaBabin1
This document discusses chemical reactions and stoichiometry. It begins by explaining that zinc and silver typically form single ionic charges of Zn2+ and Ag+ due to their electron configurations. It then defines stoichiometry as the study of mass relationships in chemistry based on the law of conservation of mass. The document goes on to explain what is included in a chemical equation such as reactants, products, physical states, and coefficients. It discusses various scenarios for balancing chemical equations including using the least common denominator method, keeping polyatomic ions together, balancing oxygen last in combustion reactions, and using decimal coefficients. The final sections cover concepts such as formula weight, molecular weight, percent composition, Avogadro's number, moles, and
The document discusses key concepts in stoichiometry including:
- Stoichiometry deals with the relative quantities of reactants and products in chemical reactions.
- Chemical reactions proceed according to mole ratios derived from balanced chemical equations.
- There are specific steps to solve mole-to-mole, mole-to-mass, and mass-to-mass stoichiometry problems.
- Limiting reactants determine the amount of product formed in a reaction. Excess reactants remain after the reaction completes.
- Percent yield calculations compare the actual and theoretical yields of a chemical reaction.
This document discusses suffixes and terminology used in medicine. It begins by listing common combining forms used to build medical terms and their meanings. It then defines several noun, adjective, and shorter suffixes and provides their meanings. Examples are given of medical terms built using combining forms and suffixes. The document also examines specific medical concepts in more depth, such as hernias, blood cells, acromegaly, splenomegaly, and laparoscopy.
The document is a chapter from a medical textbook that discusses anatomical terminology pertaining to the body as a whole. It defines the structural organization of the body from cells to tissues to organs to systems. It also describes the body cavities and identifies the major organs contained within each cavity, as well as anatomical divisions of the abdomen and back.
This document is from a textbook on medical terminology. It discusses the basic structure of medical words and how they are built from prefixes, suffixes, and combining forms. Some key points:
- Medical terms are made up of elements including roots, suffixes, prefixes, and combining vowels. Understanding these elements is important for analyzing terms.
- Common prefixes include hypo-, epi-, and cis-. Common suffixes include -itis, -algia, and -ectomy.
- Dozens of combining forms are provided, such as gastro- meaning stomach, cardi- meaning heart, and aden- meaning gland.
- Rules are provided for analyzing terms, such as reading from the suffix backward and dropping combining vowels before suffixes starting with vowels
This document is the copyright information for Chapter 25 on Cancer from the 6th edition of the textbook Molecular Cell Biology published in 2008 by W. H. Freeman and Company. The chapter was authored by a team that includes Lodish, Berk, Kaiser, Krieger, Scott, Bretscher, Ploegh, and Matsudaira.
This document is the copyright information for Chapter 24 on Immunology from the 6th edition of the textbook Molecular Cell Biology published in 2008 by W. H. Freeman and Company. The chapter was authored by Lodish, Berk, Kaiser, Krieger, Scott, Bretscher, Ploegh, and Matsudaira.
Nerve cells, also known as neurons, are highly specialized cells that process and transmit information through electrical and chemical signals. This chapter discusses the structure and function of neurons, how they communicate with each other via synapses, and how signals are propagated along neurons through changes in their membrane potentials. Neurons play a vital role in the nervous system by allowing organisms to process information and coordinate their responses.
This document is the copyright information for Chapter 22 from the 6th edition of the textbook "Molecular Cell Biology" published in 2008 by W. H. Freeman and Company. The chapter is titled "The Molecular Cell Biology of Development" and is authored by Lodish, Berk, Kaiser, Krieger, Scott, Bretscher, Ploegh, and Matsudaira.
This document is the copyright information for Chapter 21 from the sixth edition of the textbook "Molecular Cell Biology" published in 2008 by W. H. Freeman and Company. The chapter is titled "Cell Birth, Lineage, and Death" and is authored by Lodish, Berk, Kaiser, Krieger, Scott, Bretscher, Ploegh, and Matsudaira.
This document is the copyright page for Chapter 20 from the 6th edition of the textbook "Molecular Cell Biology" published in 2008 by W. H. Freeman and Company. The chapter is titled "Regulating the Eukaryotic Cell Cycle" and is authored by a group of scientists including Lodish, Berk, Kaiser, Krieger, Scott, Bretscher, Ploegh, and Matsudaira.
This document is the copyright information for Chapter 19 from the 6th edition textbook "Molecular Cell Biology" published in 2008 by W. H. Freeman and Company. The chapter is titled "Integrating Cells into Tissues" and is authored by Lodish, Berk, Kaiser, Krieger, Scott, Bretscher, Ploegh, and Matsudaira.
This chapter discusses microtubules and intermediate filaments, which are types of cytoskeletal filaments that help organize and move cellular components. Microtubules are involved in processes like cell division and intracellular transport, while intermediate filaments provide mechanical strength and help integrate the nucleus with the cytoplasm. Together, these filaments play important structural and functional roles in eukaryotic cells.
This chapter discusses microfilaments, which are one of the three main types of cytoskeletal filaments found in eukaryotic cells. Microfilaments are composed of actin filaments and play important roles in cell motility, structure, and intracellular transport. They allow cells to change shape and to move by contracting or extending parts of the cell surface.
This document is the copyright page for Chapter 16 from the 6th edition of the textbook "Molecular Cell Biology" published in 2008 by W. H. Freeman and Company. The chapter is titled "Signaling Pathways that Control Gene Activity" and is authored by a group of scientists including Lodish, Berk, Kaiser, Krieger, Scott, Bretscher, Ploegh and Matsudaira.
This document is the copyright page for Chapter 15 of the 6th edition textbook "Molecular Cell Biology" by Lodish, Berk, Kaiser, Krieger, Scott, Bretscher, Ploegh, and Matsudaira. It provides the chapter title "Cell Signaling I: Signal Transduction and Short-Term Cellular Responses" and notes the copyright is held by W. H. Freeman and Company in 2008.
This document is the copyright page for Chapter 14 from the 6th edition textbook "Molecular Cell Biology" published in 2008 by W. H. Freeman and Company. The chapter is titled "Vesicular Traffic, Secretion, and Endocytosis" and is authored by a group of scientists including Lodish, Berk, Kaiser, Krieger, Scott, Bretscher, Ploegh and Matsudaira.
This chapter discusses how proteins are transported into membranes and organelles within cells. Proteins destined for membranes or organelles have targeting signals that are recognized by transport systems. The transport systems then direct the proteins to their proper destinations, such as inserting membrane proteins into membranes or delivering soluble proteins into organelles.
This document is the copyright information for Chapter 12 from the sixth edition of the textbook "Molecular Cell Biology" published in 2008 by W. H. Freeman and Company. The chapter is titled "Cellular Energetics" and is authored by Lodish, Berk, Kaiser, Krieger, Scott, Bretscher, Ploegh, and Matsudaira.
This chapter discusses the transmembrane transport of ions and small molecules across cell membranes. It covers topics such as passive transport through membrane channels and pumps, as well as active transport using ATP. The chapter is from the 6th edition of the textbook Molecular Cell Biology and is copyrighted by W. H. Freeman and Company in 2008.
This document is the copyright information for Chapter 10, titled "Biomembrane Structure", from the sixth edition of the textbook "Molecular Cell Biology" published in 2008 by W. H. Freeman and Company. The chapter was written by a team of authors including Lodish, Berk, Kaiser, Krieger, Scott, Bretscher, Ploegh and Matsudaira.
This document is the copyright information for Chapter 9 from the 6th edition of the textbook "Molecular Cell Biology" published in 2008 by W. H. Freeman and Company. The chapter is titled "Visualizing, Fractionating, and Culturing Cells" and is authored by Lodish, Berk, Kaiser, Krieger, Scott, Bretscher, Ploegh, and Matsudaira.
How to Add Chatter in the odoo 17 ERP ModuleCeline George
In Odoo, the chatter is like a chat tool that helps you work together on records. You can leave notes and track things, making it easier to talk with your team and partners. Inside chatter, all communication history, activity, and changes will be displayed.
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.
Main Java[All of the Base Concepts}.docxadhitya5119
This is part 1 of my Java Learning Journey. This Contains Custom methods, classes, constructors, packages, multithreading , try- catch block, finally block and more.
This slide is special for master students (MIBS & MIFB) in UUM. Also useful for readers who are interested in the topic of contemporary Islamic banking.
A Strategic Approach: GenAI in EducationPeter Windle
Artificial Intelligence (AI) technologies such as Generative AI, Image Generators and Large Language Models have had a dramatic impact on teaching, learning and assessment over the past 18 months. The most immediate threat AI posed was to Academic Integrity with Higher Education Institutes (HEIs) focusing their efforts on combating the use of GenAI in assessment. Guidelines were developed for staff and students, policies put in place too. Innovative educators have forged paths in the use of Generative AI for teaching, learning and assessments leading to pockets of transformation springing up across HEIs, often with little or no top-down guidance, support or direction.
This Gasta posits a strategic approach to integrating AI into HEIs to prepare staff, students and the curriculum for an evolving world and workplace. We will highlight the advantages of working with these technologies beyond the realm of teaching, learning and assessment by considering prompt engineering skills, industry impact, curriculum changes, and the need for staff upskilling. In contrast, not engaging strategically with Generative AI poses risks, including falling behind peers, missed opportunities and failing to ensure our graduates remain employable. The rapid evolution of AI technologies necessitates a proactive and strategic approach if we are to remain relevant.
বাংলাদেশের অর্থনৈতিক সমীক্ষা ২০২৪ [Bangladesh Economic Review 2024 Bangla.pdf] কম্পিউটার , ট্যাব ও স্মার্ট ফোন ভার্সন সহ সম্পূর্ণ বাংলা ই-বুক বা pdf বই " সুচিপত্র ...বুকমার্ক মেনু 🔖 ও হাইপার লিংক মেনু 📝👆 যুক্ত ..
আমাদের সবার জন্য খুব খুব গুরুত্বপূর্ণ একটি বই ..বিসিএস, ব্যাংক, ইউনিভার্সিটি ভর্তি ও যে কোন প্রতিযোগিতা মূলক পরীক্ষার জন্য এর খুব ইম্পরট্যান্ট একটি বিষয় ...তাছাড়া বাংলাদেশের সাম্প্রতিক যে কোন ডাটা বা তথ্য এই বইতে পাবেন ...
তাই একজন নাগরিক হিসাবে এই তথ্য গুলো আপনার জানা প্রয়োজন ...।
বিসিএস ও ব্যাংক এর লিখিত পরীক্ষা ...+এছাড়া মাধ্যমিক ও উচ্চমাধ্যমিকের স্টুডেন্টদের জন্য অনেক কাজে আসবে ...
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.
Macroeconomics- Movie Location
This will be used as part of your Personal Professional Portfolio once graded.
Objective:
Prepare a presentation or a paper using research, basic comparative analysis, data organization and application of economic information. You will make an informed assessment of an economic climate outside of the United States to accomplish an entertainment industry objective.
2. 4.1 The Mole Concept and Atoms
• Atoms are exceedingly small
– Unit of measurement for mass of an atom is
atomic mass unit (amu) – unit of measure for
the mass of atoms
• carbon-12 assigned the mass of exactly 12 amu
• 1 amu = 1.66 x 10-24
g
• Periodic table gives atomic weights in amu
3. • What is the atomic weight of one atom of
fluorine? Answer: 19.00 amu
• What would be the mass of this one atom
in grams?
• Chemists usually work with much larger
quantities
– It is more convenient to work with grams
than amu when using larger quantities
Mass of Atoms
4.1TheMoleConceptand
Atoms
atomF
Fg10156.3
Famu1
g101.661
atomF
Famu19.00 23-24 −
×
=
×
×
4. • A practical unit for defining a collection
of atoms is the mole
1 mole of atoms = 6.022 x 1023
atoms
• This is called Avogadro’s number
– This has provided the basis for the concept
of the mole
The Mole and Avogadro’s
Number
4.1TheMoleConceptand
Atoms
5. 4.1TheMoleConceptand
Atoms
The Mole
• To make this connection we must define
the mole as a counting unit
– The mole is abbreviated mol
• A mole is simply a unit that defines an
amount of something
– Dozen defines 12
– Gross defines 144
6. Fmol1
Fatom10022.6
Famu1
Fg1066.1
Fatom1
Famu00.19 2324
×
×
×
×
−
=19.00 g F/mol F or 19.00 g/mol F
4.1TheMoleConceptand
Atoms
Atomic Mass
• The atomic mass of one atom of an element
corresponds to:
– The average mass of a single atom in amu
– The mass of a mole of atoms in grams
– 1 atom of F is 19.00 amu 19.00 amu/atom F
– 1 mole of F is 19.00 g 19.00 g/mole F
7. 4.1TheMoleConceptand
Atoms
Molar Mass
• Molar mass - The mass in grams of 1 mole of
atoms
• What is the molar mass of carbon?
12.01 g/mol C
• This means counting out a mole of Carbon
atoms (i.e., 6.022 x 1023
) they would have a mass
of 12.01 g
• One mole of any element contains the same
number of atoms, 6.022 x 1023
, Avogadro’s number
8. 4.1TheMoleConceptand
Atoms
Calculating Atoms, Moles, and Mass
• We use the following conversion factors:
• Density converts grams – milliliters
• Atomic mass unit converts amu –
grams
• Avogadro’s number converts moles –
number of atoms
• Molar mass converts grams – moles
9. 4.1TheMoleConceptand
Atoms
Strategy for Calculations
• Map out a pattern for the required
conversion
• Given a number of grams and asked for
number of atoms
• Two conversions are required
• Convert grams to moles
1 mol S/32.06 g S OR 32.06 g S/1 mol S
• Convert moles to atoms
mol S x (6.022 x 1023
atoms S) / 1 mol S
10. 4.1TheMoleConceptand
Atoms
Practice Calculations
1. Calculate the number of atoms in 1.7
moles of boron.
2. Find the mass in grams of 2.5 mol Na
(sodium).
3. Calculate the number of atoms in 5.0 g
aluminum.
4. Calculate the mass of 5,000,000 atoms
of Au (gold)
12. 4.2 The Chemical Formula,
Formula Weight, and Molar
Mass
• Chemical formula - a combination of
symbols of the various elements that make up
the compound
• Formula unit - the smallest collection of
atoms that provide two important pieces of
information
– The identity of the atoms
– The relative number of each type of atom
13. 4.2TheChemicalFormula,
FormulaWeightandMolarMass Chemical Formula
Consider the following formulas:
• H2 – 2 atoms of hydrogen are chemically bonded
forming diatomic hydrogen, subscript 2
• H2O – 2 atoms of hydrogen and 1 atom of
oxygen, lack of subscript means one atom
• NaCl – 1 atom each of sodium and chlorine
• Ca(OH)2 – 1 atom of calcium and 2 atoms each
of oxygen and hydrogen, subscript outside
parentheses applies to all atoms inside
14. 4.2TheChemicalFormula,
FormulaWeightandMolarMass Chemical Formula
Consider the following formulas:
• (NH4)3SO4 – 2 ammonium ions and 1 sulfate ion
– Ammonium ion contains 1 nitrogen and 4 hydrogen
– Sulfate ion contains 1 sulfur and 4 oxygen
– Compound contains 2 N, 8 H, 1 S, and 4 O
• CuSO4
.
5H2O
– This is an example of a hydrate - compounds containing
one or more water molecules as an integral part of their
structure
– 5 units of water with 1 CuSO4
15. Comparison of Hydrated and
Anhydrous Copper Sulfate
Hydrated copper sulfate Anhydrous copper sulfate
4.2TheChemicalFormula,
FormulaWeightandMolarMass
Marked color difference illustrates the fact
that these are different compounds
16. Formula Weight and Molar Mass
• Formula weight - the sum of the atomic weights
of all atoms in the compound as represented by its
correct formula
– expressed in amu
• What is the formula weight of H2O?
– 16.00 amu + 2(1.008 amu) = 18.02 amu
• Molar mass – mass of a mole of compound in
grams / mole
– Numerically equal to the formula weight in amu
• What is the molar mass of H2O?
– 18.02 g/mol H2O
4.2TheChemicalFormula,
FormulaWeightandMolarMass
17. Formula Unit
• Formula unit – smallest
collection of atoms from which
the formula of a compound can
be established
• When calculating the formula
weight (or molar mass) of an
ionic compound, the smallest
unit of the crystal is used
4.2TheChemicalFormula,Formula
WeightandMolarMass
What is the molar mass of (NH4)3PO4?
3(N amu) + 12(H amu) + P amu + 4(O amu)=
3(14.01) + 12(1.008) + 30.97 + 4(16.00)=
149.10 g/mol (NH4)3PO4
18. 4.3 The Chemical Equation and the
Information It Conveys
A Recipe For Chemical Change
• Chemical equation - shorthand notation of a
chemical reaction
– Describes all of the substances that react and all
the products that form, physical states, and
experimental conditions
– Reactants – (starting materials) – the substances
that undergo change in the reaction
– Products – substances produced by the reaction
19. 4.3TheChemicalEquation
andtheInformationItConveys
Features of a Chemical Equation
1. Identity of products and reactants must
be specified using chemical symbols
2. Reactants are written to the left of the
reaction arrow and products are written
to the right
3. Physical states of reactants and products
may be shown in parentheses
4. Symbol ∆ over the reaction arrow
means that energy is necessary for the
reaction to occur
5. Equation must be balanced
20. )(O)2Hg()2HgO( 2 gls +→∆
ProductsProducts – written on the right
Reactants – written on the left of arrow
Products and reactants must be
specified using chemical symbols
Physical states are shown in parentheses
∆ – energy is needed
4.3TheChemicalEquation
andtheInformationItConveys
Features of a Chemical Equation
22. 4.3TheChemicalEquation
andtheInformationItConveysEvidence of a Reaction Occurring
The following can be visual evidence of a reaction:
•Release of a gas
– CO2 is released when acid is placed in a solution
containing CO3
2-
ions
•Formation of a solid (precipitate)
– A solution containing Ag+
ions mixed with a solution
containing Cl-
ions
•Heat is produced or absorbed
– Acid and base are mixed together
•Color changes
24. Writing Chemical Reactions
• We will learn to identify the following
patterns of chemical reactions:
– combination
– decomposition
– single-replacement
– double-replacement
• Recognizing the pattern will help you
write and understand reactions
4.3TheChemicalEquation
andtheInformationItConveys
25. A + B → AB
• Examples:
2Na(s) + Cl2(g) → 2NaCl(s)
MgO(s) + CO2(g) → MgCO3(s)
4.3TheChemicalEquation
andtheInformationItConveys
Combination Reactions
• The joining of two or more elements or
compounds, producing a product of
different composition
26. 4.3TheChemicalEquation
andtheInformationItConveys
Types of Combination
Reactions
1. Combination of a metal and a nonmetal
to form a salt
2. Combination of hydrogen and chlorine
molecules to produce hydrogen chloride
3. Formation of water from hydrogen and
oxygen molecules
4. Reaction of magnesium oxide and
carbon dioxide to produce magnesium
carbonate
27. AB → A + B
• Examples:
2HgO(s) → 2Hg(l) + O2(g)
CaCO3(s) → CaO(s) + CO2(g)
4.3TheChemicalEquation
andtheInformationItConveys
Decomposition Reactions
• Produce two or more products from a
single reactant
• Reverse of a combination reaction
29. 1. Single-replacement
• One atom replaces another in the
compound producing a new compound
• Examples:
Cu(s)+2AgNO3(aq) → 2Ag(s)+Cu(NO3)2(aq)
2Na(s) + 2H2O(l) → 2NaOH(aq) + H2(g)
A + BC → B + AC
4.3TheChemicalEquation
andtheInformationItConveys Replacement Reactions
30. 1. Replacement of copper by zinc in
copper sulfate
2. Replacement of aluminum by
sodium in aluminum nitrate
4.3TheChemicalEquation
andtheInformationItConveys
Types of Replacement
Reactions
31. 2. Double-replacement
• Two compounds undergo a “change
of partners”
• Two compounds react by
exchanging atoms to produce two
new compounds
4.3TheChemicalEquation
andtheInformationItConveys Replacement Reactions
AB + CD → AD + CB
32. AB + CD → AD + CB
4.3TheChemicalEquation
andtheInformationItConveys
Types of Double-Replacement
• Reaction of an acid with a base to
produce water and salt
HCl(aq)+NaOH(aq) →NaCl(aq)+H2O(l)
• Formation of solid lead chloride from
lead nitrate and sodium chloride
Pb(NO3)2(aq) + 2NaCl(aq) →
PbCl2(s) + 2NaNO3(aq)
33. Types of Chemical Reactions
Precipitation Reactions
• Chemical change in a solution that
results in one or more insoluble products
• To predict if a precipitation reaction can
occur it is helpful to know the
solubilities of ionic compounds
4.3TheChemicalEquation
andtheInformationItConveys
34. Predicting Whether Precipitation
Will Occur
• Recombine the ionic compounds to
have them exchange partners
• Examine the new compounds formed
and determine if any are insoluble
according to the rules in Table 4.1
• Any insoluble salt will be the precipitate
Pb(NO3)2(aq) + NaCl(aq) →
PbCl2 (?) + NaNO3 ( ?)(s) (aq)
4.3TheChemicalEquation
andtheInformationItConveys
37. The H+
on HCl was
transferred to the oxygen
in OH-
, giving H2O
4.3TheChemicalEquation
andtheInformationItConveys
Acid-Base Reactions
• These reactions involve the transfer of a
hydrogen ion (H+
) from one reactant (acid)
to another (base)
HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l)
38. Two electrons are
transferred from Zn to
Cu2+
4.3TheChemicalEquation
andtheInformationItConveys
Oxidation-Reduction Reactions
• Reaction involves the transfer of one or
more electrons from one reactant to
another
Zn(s) + Cu2+
(aq)→ Cu(s) + Zn2+
(aq)
39. Writing Chemical Reactions
Consider the following reaction:
hydrogen reacts with oxygen to produce water
• Write the above reaction as a chemical
equation
H2 + O2 → H2O
• Don’t forget the diatomic elements
4.3TheChemicalEquation
andtheInformationItConveys
40. 4.3TheChemicalEquation
andtheInformationItConveys Law of Conservation of Mass
• Law of conservation of mass - matter
cannot be either gained or lost in the
process of a chemical reaction
– The total mass of the products must equal
the total mass of the reactants
41. A Visual Example of the Law of
Conservation of Mass
4.4BalancingChemical
Equations
42. 4.4 Balancing Chemical Equations
• A chemical equation shows the molar
quantity of reactants needed to produce a
particular molar quantity of products
• The relative number of moles of each
product and reactant is indicated by
placing a whole-number coefficient
before the formula of each substance in
the chemical equation
43. )(O)2Hg()2HgO( 2 gls +→∆
Coefficient - how many of that substance
are in the reaction
4.4BalancingChemical
Equations
Balancing
• The equation must be balanced
– All the atoms of every reactant must also
appear in the products
• Number of Hg on left? 2
– on right 2
• Number of O on left? 2
– on right 2
44. 4.4BalancingChemical
Equations
Examine the Equation
H2 + O2 → H2O
• Is the law of conservation of mass obeyed
as written? NO
• Balancing chemical equations uses coefficients
to ensure that the law of conservation of mass is
obeyed
• You may never change subscripts!
• WRONG: H2 + O2 → H2O2
45. Step 1. Count the number of moles of
atoms of each element on both
product and reactant sides
Reactants Products
2 mol H 2 mol H
2 mol O 1 mol O
4.4BalancingChemical
Equations
Steps in Equation Balancing
The steps to balancing:
H2 + O2 → H2O
46. H2 + O2 → H2O
H2 + O2 → 2H2O
This balances oxygen, but is hydrogen
still balanced?
4.4BalancingChemical
Equations
Step 2. Determine which elements are not
balanced – do not have same number on
both sides of the equation
– Oxygen is not balanced
Step 3. Balance one element at a time by
changing the coefficients
Steps in Equation Balancing
47. Reactants Products
4 mol H 4 mol H
2 mol O 2 mol O
4.4BalancingChemical
Equations
Steps in Equation Balancing
H2 + O2 → 2H2O
How will we balance hydrogen?
2H2 + O2 → 2H2O
Step 4. Check! Make sure the law of
conservation of mass is obeyed
50. 4.5 Calculations Using the
Chemical Equation
• Calculation quantities of reactants and
products in a chemical reaction has many
applications
• Need a balanced chemical equation for the
reaction of interest
• The coefficients represent the number of
moles of each substance in the equation
51. 4.5CalculationsUsingthe
ChemicalEquation
General Principles
1. Chemical formulas of all reactants and
products must be known
2. Equation must be balanced to obey the
law of conservation of mass
• Calculations of an unbalanced equation
are meaningless
1. Calculations are performed in terms of
moles
• Coefficients in the balanced equation
represent the relative number of moles of
products and reactants
52. Using the Chemical Equation
• Examine the reaction:
2H2 + O2 → 2H2O
• Coefficients tell us?
– 2 mol H2 reacts with 1 mol O2 to produce 2
mol H2O
• What if 4 moles of H2 reacts with 2 moles of
O2?
– It yields 4 moles of H2O
4.5CalculationsUsingthe
ChemicalEquation
53. 2H2 + O2 → 2H2O
4.5CalculationsUsingthe
ChemicalEquation
Using the Chemical Equation
• The coefficients of the balanced equation
are used to convert between moles of
substances
• How many moles of O2 are needed to
react with 4.26 moles of H2?
• Use the factor-label method to perform
this calculation
54. 2H2 + O2 → 2H2O
=×
2
2
2
Hmol__
O__mol
Hmol26.4 1
2
2.13 mol O2
4.5CalculationsUsingthe
ChemicalEquation
Use of Conversion Factors
• Digits in the conversion factor come
from the balanced equation
55. 4.5CalculationsUsingthe
ChemicalEquation
Conversion Between Moles
and Grams
• Requires only the formula weight
• Convert 1.00 mol O2 to grams
– Plan the path
– Find the molar mass of oxygen
• 32.0 g O2 = 1 mol O2
– Set up the equation
– Cancel units 1.00 mol O2 x 32.0 g O2
1 mol O2
– Solve equation 1.00 x 32.0 g O2 = 32.0 g O2
moles of
Oxygen
grams of
Oxygen
56. 4.5CalculationsUsingthe
ChemicalEquation
Conversion of Mole Reactants to
Mole Products
• Use a balanced equation
• C3H8(g) + 5O2(g) 3CO2(g) + 4H2O(g)
• 1 mol C3H8results in:
– 5 mol O2 consumed 1 mol C3H8 /5 mol O2
– 3 mol CO2formed 1 mol C3H8/3 mol CO2
– 4 mol H2O formed 1 mol C3H8/4 mol H2O
• This can be rewritten as conversion
factors
57. 4.5CalculationsUsingthe
ChemicalEquationCalculating Reacting Quantities
• Calculate grams O2 reacting with 1.00 mol C3H8
• Use 2 conversion factors
– Moles C3H8 to moles O2
– Moles of O2 to grams O2
– Set up the equation and cancel units
– 1.00 mol C3H8 x 5 mol O2x 32.0 g O2 =
1 mol C3H8 1 mol O2
– 1.00 x 5 x 32.0 g O = 1.60 x 102
g O
moles
Oxygen
grams
Oxygen
moles
C3H8
58. 4.5CalculationsUsingthe
ChemicalEquation
Calculating Grams of Product
from Moles of Reactant
• Calculate grams CO2 from combustion of 1.00
mol C3H8
• Use 2 conversion factors
– Moles C3H8 to moles CO2
– Moles of CO2 to grams CO2
– Set up the equation and cancel units
– 1.00 mol C3H8 x 3 mol CO2x 44.0 g CO2 =
1 mol C3H8 1 mol CO2
– 1.00 x 3 x 44.0 g CO = 1.32 x 102
g CO
moles
CO2
grams
CO2
moles
C3H8
59. 4.5CalculationsUsingthe
ChemicalEquation
Relating Masses of Reactants
and Products
• Calculate grams C3H8 required to produce
36.0 grams of H2O
• Use 3 conversion factors
– Grams H2O to moles H2O
– Moles H2O to moles C3H8
– Moles of C3H8 to grams C3H8
– Set up the equation and cancel units
36.0 g H2O x 1 mol H2O x 1 mol C3H8x 44.0 g C3H8
18.0 g H2O 4 mol H2O 1 mol C3H8
– 36.0 x [1/18.0] x [1/4] x 44.0 g C3H8 = 22.0 g C3H8
moles
H2O
grams
C3H8
moles
C3H8
grams
H2O
60. 4.5CalculationsUsingthe
ChemicalEquation
Calculating a Quantity of Reactant
• Ca(OH)2 neutralizes HCl
• Calculate grams HCl neutralized by 0.500 mol
Ca(OH)2
– Write chemical equation and balance
• Ca(OH)2(s) + 2HCl(aq) CaCl2(s) + 2H2O(l)
– Plan the path
– Set up the equation and cancel units
0.500 mol Ca(OH)2 x 2 mol HCl x 36.5 g HCl
1 mol Ca(OH)2 1 mol HCl
Solve equation 0.500 x [2/1] x 36.5 g HCl = 36.5 g HCl
moles
Ca(OH)2
grams
HCl
moles
HCl
62. Na + Cl2 → NaCl
4.5CalculationsUsingthe
ChemicalEquation
Sample Calculation
1. Balance the equation
2. Calculate the moles Cl2 reacting with 5.00
mol Na
3. Calculate the grams NaCl produced when
5.00 mol Na reacts with an excess of Cl2
4. Calculate the grams Na reacting with
5.00 g Cl2
2Na + Cl2 → 2NaCl
64. 4.5CalculationsUsingthe
ChemicalEquation
Sample Calculation
If the theoretical yield of iron was 30.0 g
and actual yield was 25.0 g, calculate the
percent yield:
2 Al(s) + Fe2O3(s) → Al2O3(aq) + 2Fe(aq)
• [25.0 g / 30.0 g] x 100% = 83.3%
• Calculate the % yield if 26.8 grams iron
was collected in the same reaction