Tro3 lecture 06
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1) A mole is defined as 6.022x10^23 atoms or molecules and can be used to convert between the number of particles and mass of a sample. 2) The molar mass of an element or compound allows conversion between moles and grams. For elements, molar mass equals atomic mass in grams. 3) Chemical formulas indicate the number of moles of each element in a mole of compound using subscripts. This allows calculation of mass or number of atoms of one constituent from the formula unit.
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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)
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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.
Chapter 10 - Chemical Quantities
- The mole is a unit used to measure very large numbers of small particles like atoms or molecules. One mole equals 6.02 x 10^23 particles.
- The mass of one mole of an element is called its gram atomic mass (gam). The mass of one mole of a compound is called its gram molecular mass (gmm) or gram formula mass (gfm).
- At standard temperature and pressure, one mole of any gas occupies a volume of 22.4 L. This volume is known as the molar volume.
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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.
Mole concept
This presentation introduces the mole concept, which defines a mole as the amount of a substance that contains the same number of elementary entities (atoms, molecules, etc.) as there are atoms in exactly 12 grams of carbon-12. Key points covered include:
1. A mole is used to quantify the number of atoms or molecules in a given mass of a substance. The mass of one mole of a substance in grams is equal to its molar mass.
2. Calculations involving moles allow for determining amounts of substances in chemical reactions based on molar ratios and the law of conservation of mass.
3. Important formulas covered are the definitions of molar mass, relative molecular mass, and concentration of solutions
Chapter 3.powerpoint
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.
Class XI Chemistry - Mole Concept
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.
C05 the mole concept
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.
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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)
C05 the mole concept
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.
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- The mole is a unit used to measure very large numbers of small particles like atoms or molecules. One mole equals 6.02 x 10^23 particles.
- The mass of one mole of an element is called its gram atomic mass (gam). The mass of one mole of a compound is called its gram molecular mass (gmm) or gram formula mass (gfm).
- At standard temperature and pressure, one mole of any gas occupies a volume of 22.4 L. This volume is known as the molar volume.
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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.
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This presentation introduces the mole concept, which defines a mole as the amount of a substance that contains the same number of elementary entities (atoms, molecules, etc.) as there are atoms in exactly 12 grams of carbon-12. Key points covered include:
1. A mole is used to quantify the number of atoms or molecules in a given mass of a substance. The mass of one mole of a substance in grams is equal to its molar mass.
2. Calculations involving moles allow for determining amounts of substances in chemical reactions based on molar ratios and the law of conservation of mass.
3. Important formulas covered are the definitions of molar mass, relative molecular mass, and concentration of solutions
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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.
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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.
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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.
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This document provides an overview of key concepts in chemical quantities including:
1) The mole is a unit used to measure very large amounts of substances and is defined as 6.02x10^23 items. It can represent atoms, molecules, ions, or formula units.
2) Molar mass is the mass of one mole of a substance in grams and can be used to determine the mass of any amount of a substance.
3) Molar conversions allow calculations between the number of moles, mass in grams, and number of particles using molar mass and Avogadro's number.
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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.
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- Standard temperature and pressure conditions for reporting gas measurements
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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.
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The document discusses methods for converting between moles, mass, and number of particles for chemical substances:
1) Equations are provided to convert between moles and mass using molar mass, as well as to find the number of particles from moles using Avogadro's number.
2) Worked examples demonstrate using the formulas to calculate moles from mass and vice versa, as well as converting moles to number of atoms or molecules.
3) Determining empirical and molecular formulas from percentage composition by mass is explained. The empirical formula is used to calculate the molecular formula by considering the mass and moles of the empirical formula unit.
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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.
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The document provides learning objectives and concepts for a chemistry chapter, including:
a) Proving the law of conservation of mass through experiments.
b) Using experimental data to prove laws such as Gay-Lussac's law and Proust's law of definite proportions.
c) Explaining concepts such as moles, molar mass, stoichiometry, and reaction stoichiometry.
It outlines key fundamental laws of chemistry and concepts students are expected to understand after completing the chapter.
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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.
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1. The document discusses moles, molar mass, and calculations involving moles. It provides examples of calculating the number of particles, ions, or molecules in 1 mole of various substances like iron, carbon dioxide, sodium chloride, and magnesium chloride.
2. It also explains how to calculate molar mass using relative atomic mass and gives examples for iron, carbon dioxide, sodium chloride, and magnesium chloride.
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This document discusses relative atomic mass and relative molecular mass. It provides examples to calculate these values.
The key points are:
1. Relative atomic mass (Ar) is the average mass of a single atom of an element compared to 1/12 the mass of one carbon-12 atom.
2. The relative molecular mass (Mr) of a molecule is the sum of the relative atomic masses of all the atoms in the molecule.
3. Examples are provided to calculate relative atomic masses and relative molecular masses using atomic mass values and molecular formulas. Formulas, atomic masses, and molecular masses are compared to calculate unknown values.
Ib quan chemistry(mole concept)
1. The relative formula mass of calcium carbonate (CaCO3) is 100 g/mol.
2. One mole of calcium carbonate will react with 2 moles of hydrochloric acid (HCl).
3. Therefore, the mass of hydrochloric acid that will react with 1 mole (100 g) of calcium carbonate is 2 x 36.5 g = 73 g, since the molar mass of HCl is 36.5 g/mol.
MOLECULAR MASS
Molecular mass is the sum of the atomic masses of all the atoms in a molecule. Formula unit mass is the sum of atomic masses of all atoms in a formula unit of a compound, such as sodium chloride (NaCl) which has a formula unit mass of 58.5 u. The mole concept refers to one mole being the amount of substance that contains as many elementary entities (atoms, molecules, ions or particles) as there are atoms in exactly 12 grams of carbon-12, where one mole contains 6.022x10^23 elementary entities and the mass in grams of one mole of a substance is equal to its molar mass.
3.3 the mole
The document defines key concepts related to the mole including:
- The mole is a unit that represents 6.022x10^23 particles and allows chemists to conveniently work with substances at a macroscopic scale.
- One mole of any substance has a mass in grams numerically equal to its molar mass.
- Common calculations involve determining moles, mass, particles, or volume of a substance using its molar mass and amount in moles.
- The molar volume of a gas is 22.4L at STP allowing for volume calculations.
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The document discusses the concept of the mole in chemistry. It defines a mole as 6.02 x 1023 representative particles, which can be atoms, molecules, or formula units. It provides examples of calculating the number of moles and mass of different substances. It also explains that 1 mole of any gas occupies 22.4 liters at standard temperature and pressure. Key terms discussed include molar mass, percent composition, empirical formula, and molecular formula.
Mole concept ok1294991357
The document discusses atomic and molecular masses, moles, molar masses, and calculating empirical and molecular formulas. It provides examples of calculating moles from masses and vice versa using molar masses. It also discusses calculating percentage compositions and determining molecular formulas from empirical formulas using molar masses.
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This document provides an overview of key concepts in chemical quantities including:
1) The mole is a unit used to measure very large amounts of substances and is defined as 6.02x10^23 items. It can represent atoms, molecules, ions, or formula units.
2) Molar mass is the mass of one mole of a substance in grams and can be used to determine the mass of any amount of a substance.
3) Molar conversions allow calculations between the number of moles, mass in grams, and number of particles using molar mass and Avogadro's number.
Chemical formulae
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.
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This document summarizes key concepts relating to gases, including:
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- 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.
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This document discusses the mole concept in chemistry. It defines key terms like mole, molar mass, relative atomic mass, and Avogadro's number. The main points are:
- A mole is the unit used to measure the number of elementary particles in a substance and represents 6.022x10^23 particles.
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The document discusses the mole, which relates the number of particles in a substance to its mass in grams. It defines one mole as 6.02 x 10^23 particles, known as Avogadro's number. It provides examples of calculating moles, mass, and number of particles using molar mass and unit conversion with moles. Key relationships discussed are mass=moles×molar mass and number of particles=moles×Avogadro's number.
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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.
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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.
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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.
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The document discusses methods for converting between moles, mass, and number of particles for chemical substances:
1) Equations are provided to convert between moles and mass using molar mass, as well as to find the number of particles from moles using Avogadro's number.
2) Worked examples demonstrate using the formulas to calculate moles from mass and vice versa, as well as converting moles to number of atoms or molecules.
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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.
Stoikiometri reaksi
The document provides learning objectives and concepts for a chemistry chapter, including:
a) Proving the law of conservation of mass through experiments.
b) Using experimental data to prove laws such as Gay-Lussac's law and Proust's law of definite proportions.
c) Explaining concepts such as moles, molar mass, stoichiometry, and reaction stoichiometry.
It outlines key fundamental laws of chemistry and concepts students are expected to understand after completing the chapter.
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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.
IB Chemistry on Mole Concept
1. The document discusses moles, molar mass, and calculations involving moles. It provides examples of calculating the number of particles, ions, or molecules in 1 mole of various substances like iron, carbon dioxide, sodium chloride, and magnesium chloride.
2. It also explains how to calculate molar mass using relative atomic mass and gives examples for iron, carbon dioxide, sodium chloride, and magnesium chloride.
3. Key concepts discussed include the definition of 1 mole as 6.02 x 1023 particles and its relationship to the Avogadro constant, and how molar mass is used to express the mass of 1 mole of a substance.
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This document discusses relative atomic mass and relative molecular mass. It provides examples to calculate these values.
The key points are:
1. Relative atomic mass (Ar) is the average mass of a single atom of an element compared to 1/12 the mass of one carbon-12 atom.
2. The relative molecular mass (Mr) of a molecule is the sum of the relative atomic masses of all the atoms in the molecule.
3. Examples are provided to calculate relative atomic masses and relative molecular masses using atomic mass values and molecular formulas. Formulas, atomic masses, and molecular masses are compared to calculate unknown values.
Ib quan chemistry(mole concept)
1. The relative formula mass of calcium carbonate (CaCO3) is 100 g/mol.
2. One mole of calcium carbonate will react with 2 moles of hydrochloric acid (HCl).
3. Therefore, the mass of hydrochloric acid that will react with 1 mole (100 g) of calcium carbonate is 2 x 36.5 g = 73 g, since the molar mass of HCl is 36.5 g/mol.
MOLECULAR MASS
Molecular mass is the sum of the atomic masses of all the atoms in a molecule. Formula unit mass is the sum of atomic masses of all atoms in a formula unit of a compound, such as sodium chloride (NaCl) which has a formula unit mass of 58.5 u. The mole concept refers to one mole being the amount of substance that contains as many elementary entities (atoms, molecules, ions or particles) as there are atoms in exactly 12 grams of carbon-12, where one mole contains 6.022x10^23 elementary entities and the mass in grams of one mole of a substance is equal to its molar mass.
3.3 the mole
The document defines key concepts related to the mole including:
- The mole is a unit that represents 6.022x10^23 particles and allows chemists to conveniently work with substances at a macroscopic scale.
- One mole of any substance has a mass in grams numerically equal to its molar mass.
- Common calculations involve determining moles, mass, particles, or volume of a substance using its molar mass and amount in moles.
- The molar volume of a gas is 22.4L at STP allowing for volume calculations.
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The module aims to teach students how to use moles to relate mass, number of particles, and chemical equations.
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The Avogadro constant, represents the number of carbon-12 atoms in exactly 12 g of pure carbon-12. the value of Avogadro constant is 6.0221421 푥 10^23. How many atoms of K-40 (Radioactive isotope) are present in 225 mL of whole milk containing 1.65 mg K/mL?
Final Review
1. The document discusses various stoichiometry concepts including mole-to-mole, mole-to-mass, and mass-to-mass calculations using balanced chemical equations.
2. Examples are provided for determining moles, grams, and numbers of particles from given values through step-by-step working.
3. Key stoichiometry topics covered are mole ratios, molar mass, percent composition, and solving various stoichiometric problems.
Stoichiometry
1. The concentrations given are: 1.0 E -9, 2.4 E -3, 6.7 E -7, 4.4 E -5.
2. The document discusses stoichiometry, mole conversions, molarity, and percent composition in chemistry.
3. Stoichiometry is the study of quantitative relationships in chemical reactions using molar ratios and mole conversions.
Zumdahl Chapter 8.ppt
This document discusses concepts related to atomic masses and moles. It defines key terms like:
- Atomic mass unit (amu)
- Molar mass, which is the mass in grams of one mole of a compound
- Moles, which provide a way to count particles using Avogadro's number of 6.022x1023 particles per mole
- How to use atomic masses, molar masses, and moles to convert between masses and numbers of atoms/molecules in examples.
Ch3 stoichiometry
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
Introduction to stocihiometry milestone
Stoichiometry is the calculation of quantities in balanced chemical reactions. It allows chemists to determine the amounts of reactants needed and products expected based on balanced equations, similar to recipes. The document reviews key stoichiometry concepts like moles, molar mass, and mole ratios that allow conversions between masses, moles of substances, and atoms. It provides examples of using mole ratios from balanced equations to solve stoichiometry problems involving amounts in grams, moles, or particles.
Ch3 z53 stoich
This document discusses atomic mass and isotopes. It begins by explaining that an atomic mass unit (amu) is used to discuss the mass of atoms, where 1 amu is 1/12 the mass of a carbon-12 atom. Atomic masses listed in the periodic table are in amu. Isotopes have different atomic masses that result in the average atomic mass not being a whole number. Examples are provided to demonstrate calculating average atomic masses from the masses and abundances of isotopes.
Moles molar mass_avonumb pt1
1) The document discusses how Avogadro's number allows chemists to determine the number of atoms or molecules in a given number of moles of a substance.
2) It provides examples of using molar mass to convert between mass and moles, and using Avogadro's number to convert between moles and number of particles.
3) Key concepts covered include the mole as a unit for counting particles, molar mass, and conversion factors involving moles, mass, and number of particles.
Chemistry - Chp 12 - Stoichiometry - PowerPoint
The document summarizes key concepts from Chapter 12 of a chemistry textbook on stoichiometry. It discusses how to interpret and work with balanced chemical equations, including in terms of particles, moles, mass, and gas volume. It explains the concepts of limiting reagents, theoretical yield, actual yield and percent yield. Key steps for solving stoichiometry problems are outlined.
Mole
The document discusses several chemistry concepts including:
- It would take 19 million years to spend a mole of $1 coins (6.02 x 10^23 coins) if spent at a rate of 1 billion per second.
- Atomic masses on the periodic table represent relative masses of atoms and are useful for determining ratios and masses in grams of elements and compounds.
- A mole is a number (6.02 x 10^23) that allows easy conversion between atomic/molar masses and numbers of atoms/molecules.
Class 11 sbcc part VII
This document discusses empirical and molecular formulas. It provides examples to explain the process of determining empirical formulas from percentage composition data. The key steps are: 1) calculating grams of each element, 2) converting to moles, 3) dividing by the smallest number of moles to get ratios, 4) multiplying ratios to get near-whole numbers, 5) rounding to whole numbers, and 6) writing the empirical formula. A simple example finds the empirical formula of a compound containing 54.55% C, 9.09% H, and 36.36% O to be C2H4O. The document also discusses molecular formulas and how they differ from empirical formulas.
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Chemistry - Chp 10 - Chemical Quantities - PowerPoint
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Tro3 lecture 03[1]
Matter is anything that occupies space and has mass. It is composed of tiny particles called atoms and molecules. Matter exists in three physical states - solid, liquid, and gas - which are distinguished by their properties of shape, volume, compressibility, and flow. The composition and structure of matter can be further classified as pure substances or mixtures based on whether the matter's makeup is uniform or variable. Chemical and physical changes in matter are distinguished by whether they alter the substance's composition.
Tro3 lecture 02[1]
1) 0.987 + 125.1 - 1.22
2) 0.987 has 3 significant figures, 125.1 has 3 significant figures, 1.22 has 3 significant figures
3) The calculation should have 3 significant figures
4) 0.987 + 125.1 - 1.22 = 124.967
5) Rounded to 3 significant figures is 125
The final answer is 125
Tro3 lecture 01[1]
Chemistry seeks to understand the relationships between the particle structure of matter and its observable properties by studying atoms and molecules. The scientific method involves making observations, formulating hypotheses and experiments to test ideas, and developing laws and theories. Hypotheses are tentative explanations that can be tested and falsified through experimentation, which validates or invalidates the hypotheses. Laws summarize past observations into general predictive statements that can also be tested experimentally.
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The document discusses chemical reactions and equations. It provides examples of writing balanced chemical equations, including combustion reactions. Key steps in writing equations are identifying reactants and products, determining the chemical formulas, counting atoms on each side, and adding coefficients to balance the equation. Examples show balancing equations for reactions of metals with oxygen and other common reactions like burning of fuels.
Tro3 lecture 05
The document discusses the concepts of constant composition of pure substances and the molecular basis for this observation. It specifically addresses:
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Tro3 lecture 04
The document summarizes key concepts from Chapter 4 of Nivaldo Tro's "Introductory Chemistry" textbook, including:
1) John Dalton proposed atoms as tiny, indivisible particles that combine in whole number ratios to form compounds.
2) Atoms are composed of protons, neutrons, and electrons, with protons and electrons determining an element's identity and charge.
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Tro3 lecture 06
- 8. Practice—Calculate the Number of Atoms in 2.45 Mol of Copper. Tro's "Introductory Chemistry", Chapter 6
- 9. Practice—Calculate the Number of Atoms in 2.45 Mol of Copper, Continued. Tro's "Introductory Chemistry", Chapter 6 Since atoms are small, the large number of atoms makes sense. 1 mol = 6.022 x 10 23 atoms 2.45 mol Cu atoms Cu Check: Solution: Solution Map: Relationships: Given: Find: mol Cu atoms Cu
- 11. Example 6.2—Calculate the Moles of Sulfur in 57.8 G of Sulfur. Tro's "Introductory Chemistry", Chapter 6 Since the given amount is much less than 1 mol S, the number makes sense. 1 mol S = 32.07 g 57.8 g S mol S Check: Solution: Solution Map: Relationships: Given: Find: g S mol S
- 19. Practice—Calculate the Moles of Carbon in 0.0265 g of Pencil Lead. Tro's "Introductory Chemistry", Chapter 6
- 20. Practice—Calculate the Moles of Carbon in 0.0265 g of Pencil Lead, Continued. Tro's "Introductory Chemistry", Chapter 6 Since the given amount is much less than 1 mol C, the number makes sense. 1 mol C = 12.01 g 0.0265 g C mol C Check: Solution: Solution Map: Relationships: Given: Find: g C mol C
- 28. Practice—How Many Moles Are in 50.0 g of PbO 2 ? (Pb = 207.2, O = 16.00) Tro's "Introductory Chemistry", Chapter 6
- 29. Practice—How Many Moles Are in 50.0 G of PbO 2 ? (Pb = 207.2, O = 16.00), Continued Tro's "Introductory Chemistry", Chapter 6 Since the given amount is less than 239.2 g, the moles being < 1 makes sense. 1 mol PbO 2 = 239.2 g 50.0 g mol PbO 2 moles PbO 2 Check: Solution: Solution Map: Relationships: Given: Find: g PbO 2 mol PbO 2
- 38. Practice—Find the Mass of Sodium in 6.2 g of NaCl (MM: Na = 22.99, Cl = 35.45) Tro's "Introductory Chemistry", Chapter 6
- 39. Practice—Find the Mass of Sodium in 6.2 g of NaCl, Continued Tro's "Introductory Chemistry", Chapter 6 1 mol NaCl = 58.44 g, 1 mol Na = 22.99 g, 1 mol Na : 1 mol NaCl Since the amount of Na is less than the amount of NaCl, the answer makes sense. 6.2 g NaCl g Na Check: Solution: Solution Map: Relationships: Given: Find: g NaCl mol NaCl mol Na g Na
- 41. Example 6.9—Find the Mass Percent of Cl in C 2 Cl 4 F 2 . Since the percentage is less than 100 and Cl is much heavier than the other atoms, the number makes sense. C 2 Cl 4 F 2 % Cl by mass Check: Solution: Solution Map: Relationships: Given: Find:
- 46. Empirical Formulas, Continued Benzene Molecular formula = C 6 H 6 Empirical formula = CH Glucose Molecular formula = C 6 H 12 O 6 Empirical formula = CH 2 O Hydrogen Peroxide Molecular formula = H 2 O 2 Empirical formula = HO
- 47. Practice—Determine the Empirical Formula of Benzopyrene, C 20 H 12 . Tro's "Introductory Chemistry", Chapter 6
- 57. Example 6.12—Finding an Empirical Formula from Experimental Data Tro's "Introductory Chemistry", Chapter 6
- 65. Practice—Determine the Empirical Formula of Stannous Fluoride, which Contains 75.7% Sn (118.70) and the Rest Fluorine (19.00). Tro's "Introductory Chemistry", Chapter 6
- 66. Practice—Determine the Empirical Formula of Stannous Fluoride, which Contains 75.7% Sn (118.70) and the Rest Fluorine (19.00), Continued. Tro's "Introductory Chemistry", Chapter 6 Given: 75.7% Sn, (100 – 75.3) = 24.3% F in 100 g stannous fluoride there are 75.7 g Sn and 24.3 g F. Find: Sn x F y Conversion Factors: 1 mol Sn = 118.70 g; 1 mol F = 19.00 g Solution Map: g Sn mol Sn g F mol F pseudo- formula empirical formula mole ratio whole number ratio
- 67. Practice—Determine the Empirical Formula of Stannous Fluoride, which Contains 75.7% Sn (118.70) and the Rest Fluorine (19.00), Continued. Tro's "Introductory Chemistry", Chapter 6 Apply solution map: Sn 0.638 F 1.28 SnF 2
- 68. All These Molecules Have the Same Empirical Formula. How Are the Molecules Different? Tro's "Introductory Chemistry", Chapter 6 Name Molecular Formula Empirical Formula Glyceraldehyde C 3 H 6 O 3 CH 2 O Erythrose C 4 H 8 O 4 CH 2 O Arabinose C 5 H 10 O 5 CH 2 O Glucose C 6 H 12 O 6 CH 2 O
- 69. All These Molecules Have the Same Empirical Formula. How Are the Molecules Different?, Continued Tro's "Introductory Chemistry", Chapter 6 Name Molecular Formula Empirical Formula Molar Mass, g Glyceraldehyde C 3 H 6 O 3 CH 2 O 90 Erythrose C 4 H 8 O 4 CH 2 O 120 Arabinose C 5 H 10 O 5 CH 2 O 150 Glucose C 6 H 12 O 6 CH 2 O 180
Editor's Notes
- Formula is C15H24