Worksheet for practicing working out empirical formulae. Pupils will need a data sheet or a list of relative atomic masses to be able to complete the questions.
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.
The mole is a unit used in chemistry to express amounts of substances. It represents 6.022x10^23 elementary entities, such as atoms, molecules, ions or other particles of a substance. This number is known as Avogadro's constant, after scientist Amedeo Avogadro who proposed that equal volumes of gases at the same temperature and pressure contain equal numbers of molecules. The mass of one mole of a substance, known as its molar mass, can be used to determine the number of moles in a given mass of that substance and vice versa through calculation.
The document discusses how to determine the formula of a compound through experimentation. It provides an example of finding the formula of magnesium oxide. Key steps include measuring the masses of reactants and products, calculating moles of each substance, and determining the ratio of elements in the compound based on the mole ratios. The formula of magnesium oxide in this example is MgO.
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
1) Dalton's Atomic Theory from 1808 proposed that all matter is composed of extremely small indivisible particles called atoms, atoms of a given element are identical, and atoms of different elements have different properties.
2) J.J. Thomson discovered the electron in 1897 and proposed the "plum pudding" model of the atom, which had a positive charge spread throughout the atom with electrons embedded in it.
3) Rutherford's gold foil experiments from 1909 showed that the positive charge of the atom is concentrated in a small, dense nucleus with electrons orbiting the outside, overturning Thomson's model.
Protons and neutrons make up the tiny, dense nucleus at the center of the atom, accounting for nearly all of its mass. Electrons orbit rapidly around the nucleus and take up nearly the entire volume of the atom. The number of protons determines the identity of an element, while neutrons distinguish between isotopes of that element. Chemical properties depend on the number of protons and electrons. The mass and radioactive properties depend on the total number of protons and neutrons in the nucleus.
Chemistry - Chp 10 - Chemical Quantities - PowerPointMr. Walajtys
This chapter discusses the mole as a unit for measuring amounts of substances. It defines key terms like the mole, Avogadro's number, molar mass, and representative particles. It explains how to use molar mass to convert between mass and moles of a substance. The chapter also covers calculations involving chemical formulas, percent composition, and determining empirical and molecular formulas from experimental data.
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.
The mole is a unit used in chemistry to express amounts of substances. It represents 6.022x10^23 elementary entities, such as atoms, molecules, ions or other particles of a substance. This number is known as Avogadro's constant, after scientist Amedeo Avogadro who proposed that equal volumes of gases at the same temperature and pressure contain equal numbers of molecules. The mass of one mole of a substance, known as its molar mass, can be used to determine the number of moles in a given mass of that substance and vice versa through calculation.
The document discusses how to determine the formula of a compound through experimentation. It provides an example of finding the formula of magnesium oxide. Key steps include measuring the masses of reactants and products, calculating moles of each substance, and determining the ratio of elements in the compound based on the mole ratios. The formula of magnesium oxide in this example is MgO.
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
1) Dalton's Atomic Theory from 1808 proposed that all matter is composed of extremely small indivisible particles called atoms, atoms of a given element are identical, and atoms of different elements have different properties.
2) J.J. Thomson discovered the electron in 1897 and proposed the "plum pudding" model of the atom, which had a positive charge spread throughout the atom with electrons embedded in it.
3) Rutherford's gold foil experiments from 1909 showed that the positive charge of the atom is concentrated in a small, dense nucleus with electrons orbiting the outside, overturning Thomson's model.
Protons and neutrons make up the tiny, dense nucleus at the center of the atom, accounting for nearly all of its mass. Electrons orbit rapidly around the nucleus and take up nearly the entire volume of the atom. The number of protons determines the identity of an element, while neutrons distinguish between isotopes of that element. Chemical properties depend on the number of protons and electrons. The mass and radioactive properties depend on the total number of protons and neutrons in the nucleus.
Chemistry - Chp 10 - Chemical Quantities - PowerPointMr. Walajtys
This chapter discusses the mole as a unit for measuring amounts of substances. It defines key terms like the mole, Avogadro's number, molar mass, and representative particles. It explains how to use molar mass to convert between mass and moles of a substance. The chapter also covers calculations involving chemical formulas, percent composition, and determining empirical and molecular formulas from experimental data.
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.
The document provides an introduction to stoichiometry and the mole concept. It discusses key topics including:
1. The mole is a unit used to describe the amount of substance in chemistry and is equal to 6.022x1023 particles.
2. The molar mass of an element or compound is the mass in grams of one mole and can be used to calculate amounts in chemical reactions.
3. Conversions can be made between moles, particles, masses, and volumes using the molar mass and molar relationships like moles = mass/molar mass.
4. Solution concentration is expressed in molarity, which is the moles of solute per liter of solution. M
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.
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.
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.
The document discusses atoms, molecules, atomic number, atomic mass, and mole concepts. It provides examples of chemistry problems and their solutions involving calculations of moles, atoms, molecules, ions, and chemical compounds present in mixtures and solutions. Specifically, it discusses:
1) Calculations to determine moles of phosphoric acid, phosphorus atoms, and phosphate ions in an acid solution.
2) Calculations of aspirin mass, molecules, and carbon atoms consumed by students taking pain relievers.
3) Calculations of density, carbon atoms, and hydrogen molecules in a water-ethanol mixture.
4) Calculation of gold atoms in a ring given dimensions and alloy composition.
5) Calcul
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.
- 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.
Chemical Reactions & Mole Concept 10th Std ChemistryBabu Appat
This is prepared to impart an improved awareness to the 10th std students on their chemistry lesson two namely Chemical reactions- Mole concept. The standard theories, bosons, fermions, antimatter, quarks etc. are discussed in detail. Then the mole concept is exemplified in the light of these facts. These slides are prepared for the use of 10th standard students, SCERT , based on their chemistry lesson 2.
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.
2011 topic 01 lecture 1 - the mole and avogadro's constantDavid Young
This document provides an overview of quantitative chemistry concepts including:
- The mole is a unit used to quantify amounts of substances and represents 6.02x1023 particles.
- Pure substances can exist as individual atoms, molecules, ions or molecular units.
- Chemical formulas represent the types and ratios of elements in compounds.
- Molar mass allows conversion between mass of a sample and number of moles.
- Stoichiometry problems use molar mass and the mole concept to interconvert between moles, atoms, molecules and mass.
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.
The document discusses the mole concept in chemistry. Some key points:
- A mole is a number (6.022x1023) that represents a specific number of particles like atoms or molecules.
- 1 mole of any substance contains Avogadro's number of particles.
- The mass of 1 mole of a substance in grams is the molar mass.
- Calculations can be done to convert between moles, mass, number of particles, and molar mass.
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.
This chapter discusses methods for measuring quantities of substances. It introduces the mole as a unit for measuring amounts of matter equal to Avogadro's number of representative particles. The mole allows for easy conversion between mass and number of particles. Compound formulas allow calculating formula mass in grams per mole. Relationships between moles, mass, and volume at standard temperature and pressure enable interconversion between these quantities.
This document discusses relative molecular mass (RMM) and provides examples of calculating RMM for various molecules and compounds. It explains that RMM is calculated by adding up the relative atomic masses (RAM) of all the atoms in a molecule or compound. Several practice problems are worked out, including calculating RMM for covalent compounds like O2, H2S, and C6H12O6 as well as ionic compounds such as NaCl, MgO, and Fe2O3.3H2O. Key information about relative atomic masses and how they are typically rounded to whole numbers is also summarized.
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 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.
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
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.
The document provides an introduction to stoichiometry and the mole concept. It discusses key topics including:
1. The mole is a unit used to describe the amount of substance in chemistry and is equal to 6.022x1023 particles.
2. The molar mass of an element or compound is the mass in grams of one mole and can be used to calculate amounts in chemical reactions.
3. Conversions can be made between moles, particles, masses, and volumes using the molar mass and molar relationships like moles = mass/molar mass.
4. Solution concentration is expressed in molarity, which is the moles of solute per liter of solution. M
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.
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.
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.
The document discusses atoms, molecules, atomic number, atomic mass, and mole concepts. It provides examples of chemistry problems and their solutions involving calculations of moles, atoms, molecules, ions, and chemical compounds present in mixtures and solutions. Specifically, it discusses:
1) Calculations to determine moles of phosphoric acid, phosphorus atoms, and phosphate ions in an acid solution.
2) Calculations of aspirin mass, molecules, and carbon atoms consumed by students taking pain relievers.
3) Calculations of density, carbon atoms, and hydrogen molecules in a water-ethanol mixture.
4) Calculation of gold atoms in a ring given dimensions and alloy composition.
5) Calcul
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.
- 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.
Chemical Reactions & Mole Concept 10th Std ChemistryBabu Appat
This is prepared to impart an improved awareness to the 10th std students on their chemistry lesson two namely Chemical reactions- Mole concept. The standard theories, bosons, fermions, antimatter, quarks etc. are discussed in detail. Then the mole concept is exemplified in the light of these facts. These slides are prepared for the use of 10th standard students, SCERT , based on their chemistry lesson 2.
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.
2011 topic 01 lecture 1 - the mole and avogadro's constantDavid Young
This document provides an overview of quantitative chemistry concepts including:
- The mole is a unit used to quantify amounts of substances and represents 6.02x1023 particles.
- Pure substances can exist as individual atoms, molecules, ions or molecular units.
- Chemical formulas represent the types and ratios of elements in compounds.
- Molar mass allows conversion between mass of a sample and number of moles.
- Stoichiometry problems use molar mass and the mole concept to interconvert between moles, atoms, molecules and mass.
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.
The document discusses the mole concept in chemistry. Some key points:
- A mole is a number (6.022x1023) that represents a specific number of particles like atoms or molecules.
- 1 mole of any substance contains Avogadro's number of particles.
- The mass of 1 mole of a substance in grams is the molar mass.
- Calculations can be done to convert between moles, mass, number of particles, and molar mass.
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.
This chapter discusses methods for measuring quantities of substances. It introduces the mole as a unit for measuring amounts of matter equal to Avogadro's number of representative particles. The mole allows for easy conversion between mass and number of particles. Compound formulas allow calculating formula mass in grams per mole. Relationships between moles, mass, and volume at standard temperature and pressure enable interconversion between these quantities.
This document discusses relative molecular mass (RMM) and provides examples of calculating RMM for various molecules and compounds. It explains that RMM is calculated by adding up the relative atomic masses (RAM) of all the atoms in a molecule or compound. Several practice problems are worked out, including calculating RMM for covalent compounds like O2, H2S, and C6H12O6 as well as ionic compounds such as NaCl, MgO, and Fe2O3.3H2O. Key information about relative atomic masses and how they are typically rounded to whole numbers is also summarized.
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 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.
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
The document provides information about moles, Avogadro's number, molar mass, and stoichiometry calculations. It defines the mole as the unit for counting particles, explains that one mole contains 6.02 x 1023 particles, and how this relates to molar mass. It gives examples of calculating moles, grams, and atoms/molecules using molar mass and stoichiometric conversions through mole ratios.
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.
The document discusses calculating quantities of substances involved in chemical reactions using moles. It defines terms like relative atomic mass and explains that 1 mole of any substance has a mass in grams equal to its relative formula/molecular mass. Examples demonstrate using chemical equations and molar masses to determine the mass of products formed from given masses of reactants.
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.
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.
To find the molar mass of an element, look up its atomic mass on the periodic table. For compounds, add up the molar masses of each element present according to the chemical formula, multiplying by the subscript if there is one. For example, the molar mass of NaCl is 58.5 g/mol, calculated from the molar masses of one Na atom and one Cl atom.
A) Molecular formula of butane is C4H10
Molecular weight = (4 x 12) + (10 x 1) = 58 g/mol
B) 0.25 mol x 58 g/mol = 14.5 g
C) 2.9 g x 1 mol/58 g = 0.05 mol = 5 x 1022 molecules
D) 1.5 x 1022 molecules x 58 g/mol = 87 g
This document discusses several key concepts related to moles, grams, and liters in chemistry:
1) The gram atomic mass is the mass of one mole of an element in grams, and the gram formula mass is the mass of one mole of a compound by adding the gram atomic masses of its elements.
2) Conversions between moles and grams use the gram atomic/formula mass - moles are multiplied by mass to get grams, and grams are divided by mass to get moles.
3) One mole of any gas occupies 22.4 liters at standard temperature and pressure (0°C and 1 atmosphere). Volume can be converted to moles by dividing liters by 22.
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.
Worksheet looking at working out the number of atoms in a given mass. Pupils will need a data sheet or a list of relative atomic masses to be able to complete the questions.
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.
The document provides examples of calculations involving atomic structure including:
- Calculating the number of atoms that could fit across a penny based on atomic diameters
- Writing chemical symbols for ions and isotopes
- Predicting ionic charges
- Writing formulas for ionic and molecular compounds from element names or vice versa
- Naming acids based on their formulas
The examples illustrate various concepts and calculations involving atomic and molecular structure, isotopes, ions, and naming chemical compounds.
1. This document discusses writing chemical formulae and balancing chemical equations. It explains how relative atomic mass is used to compare the masses of different atoms by using carbon-12 as the standard.
2. Isotopes are discussed, which are atoms of the same element that have different numbers of neutrons and masses. The average atomic mass of an element takes into account the different isotopes and their abundances.
3. It also explains how the mole concept relates grams of a substance to numbers of atoms or molecules using Avogadro's number. Empirical formulae are derived from determining the ratios of elements that combine in a compound.
The document provides information on stoichiometry, including:
- Relative atomic mass is the average mass of atoms of an element taking into account isotopes, measured on a scale where carbon-12 is 12.
- Relative formula mass is the sum of the relative atomic masses of all the atoms in a chemical formula.
- A mole is the amount of a substance containing 6.02x10^23 particles like atoms or molecules. This allows for easy calculation of amounts in chemical reactions.
- Stoichiometry uses molar ratios from balanced chemical equations to calculate amounts of reactants and products in terms of moles and masses. The mole concept and molar ratios allow for determining reacting masses and dedu
A workshop hosted by the South African Journal of Science aimed at postgraduate students and early career researchers with little or no experience in writing and publishing journal articles.
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 ...।
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it describes the bony anatomy including the femoral head , acetabulum, labrum . also discusses the capsule , ligaments . muscle that act on the hip joint and the range of motion are outlined. factors affecting hip joint stability and weight transmission through the joint are summarized.
How to Manage Your Lost Opportunities in Odoo 17 CRMCeline George
Odoo 17 CRM allows us to track why we lose sales opportunities with "Lost Reasons." This helps analyze our sales process and identify areas for improvement. Here's how to configure lost reasons in Odoo 17 CRM
LAND USE LAND COVER AND NDVI OF MIRZAPUR DISTRICT, UPRAHUL
This Dissertation explores the particular circumstances of Mirzapur, a region located in the
core of India. Mirzapur, with its varied terrains and abundant biodiversity, offers an optimal
environment for investigating the changes in vegetation cover dynamics. Our study utilizes
advanced technologies such as GIS (Geographic Information Systems) and Remote sensing to
analyze the transformations that have taken place over the course of a decade.
The complex relationship between human activities and the environment has been the focus
of extensive research and worry. As the global community grapples with swift urbanization,
population expansion, and economic progress, the effects on natural ecosystems are becoming
more evident. A crucial element of this impact is the alteration of vegetation cover, which plays a
significant role in maintaining the ecological equilibrium of our planet.Land serves as the foundation for all human activities and provides the necessary materials for
these activities. As the most crucial natural resource, its utilization by humans results in different
'Land uses,' which are determined by both human activities and the physical characteristics of the
land.
The utilization of land is impacted by human needs and environmental factors. In countries
like India, rapid population growth and the emphasis on extensive resource exploitation can lead
to significant land degradation, adversely affecting the region's land cover.
Therefore, human intervention has significantly influenced land use patterns over many
centuries, evolving its structure over time and space. In the present era, these changes have
accelerated due to factors such as agriculture and urbanization. Information regarding land use and
cover is essential for various planning and management tasks related to the Earth's surface,
providing crucial environmental data for scientific, resource management, policy purposes, and
diverse human activities.
Accurate understanding of land use and cover is imperative for the development planning
of any area. Consequently, a wide range of professionals, including earth system scientists, land
and water managers, and urban planners, are interested in obtaining data on land use and cover
changes, conversion trends, and other related patterns. The spatial dimensions of land use and
cover support policymakers and scientists in making well-informed decisions, as alterations in
these patterns indicate shifts in economic and social conditions. Monitoring such changes with the
help of Advanced technologies like Remote Sensing and Geographic Information Systems is
crucial for coordinated efforts across different administrative levels. Advanced technologies like
Remote Sensing and Geographic Information Systems
9
Changes in vegetation cover refer to variations in the distribution, composition, and overall
structure of plant communities across different temporal and spatial scales. These changes can
occur natural.
The simplified electron and muon model, Oscillating Spacetime: The Foundation...RitikBhardwaj56
Discover the Simplified Electron and Muon Model: A New Wave-Based Approach to Understanding Particles delves into a groundbreaking theory that presents electrons and muons as rotating soliton waves within oscillating spacetime. Geared towards students, researchers, and science buffs, this book breaks down complex ideas into simple explanations. It covers topics such as electron waves, temporal dynamics, and the implications of this model on particle physics. With clear illustrations and easy-to-follow explanations, readers will gain a new outlook on the universe's fundamental nature.
How to Fix the Import Error in the Odoo 17Celine George
An import error occurs when a program fails to import a module or library, disrupting its execution. In languages like Python, this issue arises when the specified module cannot be found or accessed, hindering the program's functionality. Resolving import errors is crucial for maintaining smooth software operation and uninterrupted development processes.
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.
How to Make a Field Mandatory in Odoo 17Celine George
In Odoo, making a field required can be done through both Python code and XML views. When you set the required attribute to True in Python code, it makes the field required across all views where it's used. Conversely, when you set the required attribute in XML views, it makes the field required only in the context of that particular view.
1. 6. Empirical Formulae
What is an Empirical Formula?
You already know that the molecular formula of a
molecule represents the number of each kind of atom
in the molecule. So, for example, the molecular
formula of ethene is C2H4, because there are two
carbon atoms and four hydrogen atoms in the
molecule. But the empirical formula of a substance is
defined as the simplest ratio of atoms in a compound,
so the empirical formula of ethene is just CH2.
1. Write down the molecular and empirical formulae of the following molecules:
molecular
formula
empirical
formula
There are not very many cases of molecules in which the molecular formula and the empirical
formula are different. There is one in the table above.
2. Can you think of any more examples? (Clue: think of some hydrocarbons)
Example One What is the empirical formula of a compound containing 0.15 mol of zinc and
0.30 mol of chlorine?
Zn Cl
moles 0.15 mol 0.30 mol
ratio of moles 1 : 2
empirical formula ZnCl2
3. Write down the empirical formula of the compounds obtained when:
a) 0.2 mol of carbon and 0.8 mol of hydrogen are combined
b) 2 mol of oxygen and 2 mol of hydrogen are combined
c) 0.65 mol of oxygen and 1.30 mol of hydrogen are combined
d) 4.4 mol of barium and 8.8 mol of chlorine are combined
e) 1.8 mol of aluminium and 2.7 mol of oxygen are combined
Quantitative Analysis
You are probably wondering what empirical formulae are for. They are interesting from a
historical point of view. From 1860, when a list of relative atomic masses was agreed upon,
to about 1960, when NMR spectroscopy became widely available, chemists used empirical
formulae calculations to work out the formula of unknown substances. Modern chemists still
2. need to know how to do empirical formulae calculations, even if there are easier ways to
determine formulae experimentally.
The conventional Example 2 Thomas Edison weighed out 3.84g of copper and found that it
way to set out
empirical
reacted with 0.48g of oxygen. What was the empirical formula
formulae is as a of the oxide of copper that he was investigating? Notice that we have
divided by 16, which
table. The table
has a column for is the RAM of oxygen
atoms (not the RMM
each element
and we write the
Cu O of O2 molecules).
This is because we
symbol and then
the mass of each mass 3.83g 0.48g need to find out the
ratio of copper
one.
moles 3.83 0.48
= 0.0598 mol = 0.03 mol
atoms to oxygen
We then divide 64 16 atoms, in order to
obtain the empirical
the mass of each
elements by the ratio of moles 2 : 1 formula.
element’s RAM.
We are essentially empirical formula Cu2O Finally we have to write the two amounts as a
simple ratio. If you can’t ‘see’ the simple ratio
working out the
number of moles straight away, choose the smallest of the two
of each element, Don’t be put off if the ratio isn’t exact. It amounts and divide both amounts by this
because moles = is OK to round up or down, if your ratio smallest amount. This doesn’t always get you
mass / RAM. is close to a whole number. straight to the answer, but it often helps.
4. Find the empirical formulae of the d) 0.60g of magnesium was heated with
compounds formed when the following sulfur until further reaction occurred.
quantities react: The mass of magnesium sulfide formed
a) 4.32g of silver and 1.42g of chlorine was 1.40g.
b) 2.5g of calcium and 1.5g of carbon e) 10.0g of carbon was reacted with
c) 0.3g of carbon and 0.8g of oxygen hydrogen to give a compound that
d) 0.7g of iron and 3.0g of bromine weighed 12.5g.
e) 1.6g of calcium and 2.84g of chlorine f) 2.07g of sodium burnt completely in
f) 3.2g of oxygen and 5.5g of manganese oxygen to give 2.79g of sodium oxide.
g) 2.54g of iodine and 0.8g of oxygen g) 44.0g of manganese reacted completely
h) 2.1g of silicon and 2.66g of chlorine to give 88.8g of manganese oxide.
i) 5.0g of arsenic and 0.2g of hydrogen h) After reacting 1.5g of aluminium with
j) 4.00g of magnesium and 1.56g of fluorine gas, 4.67g of aluminium fluoride
nitrogen was produced.
i) 5.00g of chromium underwent complete
5. Calculate the empirical formulae of the combustion in oxygen to give an oxide
compounds formed in the following that weighed 7.31g
experiments: j) 1.54g of an oxide of nitrogen was found
a) 3.60g of magnesium was heated with to contain 0.40g of nitrogen.
bromine until further reaction
occurred. The mass of magnesium Answers
1. NH3, NH3, C3H8, C3H8, C2H6O, C2H6O, C4H10, C2H5
bromide formed was 27.60g.
b) 8.00g of a gaseous hydrocarbon was 3. a) CH4, b) HO, c) H2O, d) BaCl2, e) Al2O3
reduced to give 6.00g of carbon.
4. a) AgCl, b) CaC2, c) CO2, d) FeBr3, e) CaCl2, f) MnO2, g)
c) When 15.61g of an oxide of lead was I2O5, h) SiCl4, i) AsH3, j) Mg2N3
completely reduced, 14.49g of lead was
produced. 5. a) MgBr2, b) CH4, c) PbO, d) MgS, e) CH3, f) Na2O, g)
Mn2O7, h) AlF3, i) Cr2O3, j) N2O5