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Transcript
Unit 3/4 – Moles / Stoichiometry Chemistry Review Unit 3/4 – Moles / Stoichiometry Formula Writing, Naming & Writing Chemical Compound Formulas, Chemical Equations, Mole Interpretation, Stoichiometry Moles and Stoichiometry 1. A compound is a substance composed of two or more different elements that are chemically combined in a fixed proportion. A chemical compound can only be broken down by chemical means. 2. Chemical compounds can be represented by a specific formula and assigned a name based on the IUPAC system. 3. Types of chemical formulas include empirical, molecular, and structural. Empirical formulas show elements in their simplest whole number ratios. This may or may not be the same as the molecular formula. Molecular formulas show the actual number of atoms per element in a single molecule. Structural formulas show the number of each type of atom as well as their physical arrangement. 4. All chemical reactions show a conservation of mass, energy and charge. 5. A balanced chemical equation represents conservation of atoms. 6. The coefficients in a balanced chemical equation can be used to determine mole ratios in the reaction. 7. The formula mass of a substance is the sum of the atomic masses of its atoms. The molar mass (gram formula mass) equals the mass of one mole of that substance. 8. The percent composition by mass of each element in a compound can be calculated mathematically. 9. Types of chemical reactions include synthesis, decomposition single replacement, and double replacement. -1- Unit 3/4 – Moles / Stoichiometry June 2013 January 2013 June 2012 -2- Unit 3/4 – Moles / Stoichiometry 10: 3.3a In all chemical reactions there is a conservation of mass, energy, and charge. 17: 3.1cc A compound is a substance composed of two or more different elements that are chemically combined in a fixed proportion. A chemical compound can be broken down by chemical means. A chemical compound can be represented by a specific chemical formula and assigned a name based on the IUPAC system 37: 3.3viii Calculate the formula mass and gram-formula mass 51: 3.1ppThe concentration of a solution may be expressed in molarity (M), percent by volume, percent by mass, or parts per million (ppm) 52: 3.2b Types of chemical reactions include synthesis, decomposition, single replacement, and double replacement 53: 3.3c A balanced chemical equation represents conservation of atoms. The coefficients in a balanced chemical equation can be used to determine mole ratios in the reaction. 54: 3.2a A physical change results in the rearrangement of existing particles in a substance. A chemical change results in the formation of different substances with changed properties. 55: 3.3ix Determine the number of moles of a substance, given its mass 3.3iv Calculate simple mole-mole stoichiometry problems, given a balanced equation 68: 3.3f The percent composition by mass of each element in a compound can be calculated mathematically. -3- Unit 3/4 – Moles / Stoichiometry January 2012 -4- Unit 3/4 – Moles / Stoichiometry 17: 3.4e Equal volumes of gases at the same temperature and pressure contain an equal number of particles. 33: 3.2b Types of chemical reactions include synthesis, decomposition, single replacement, and double replacement 36: M1.1C – Use algebraic and geometric representations to describe and compare data by recognizing and converting various scales of measurement 37: 4.1d Energy released or absorbed during a chemical reaction (heat of reaction) is equal to the difference between the potential energy of the products and potential energy of the reactants 66: 3.1cc A compound is a substance composed of two or more different elements that are chemically combined in a fixed proportion. A chemical compound can be broken down by chemical means. A chemical compound can be represented by a specific chemical formula and assigned a name based on the IUPAC system. 74: 3.3c A balanced chemical equation represents conservation of atoms. The coefficients in a balanced chemical equation can be used to determine mole ratios in the reaction. 76: 3.3ix Determine the number of moles of a substance, given its mass 77: 3.3f The percent composition by mass of each element in a compound can be calculated mathematically. 78: 3.3d The empirical formula of a compound is the simplest whole-number ratio of atoms of the elements in a compound. It may be different from the molecular formula, which is the actual ratio of atoms in a molecule of that compound. June 2011 -5- Unit 3/4 – Moles / Stoichiometry 7: 3.3a In all chemical reactions there is a conservation of mass, energy, and charge. 9: 3.3e The formula mass of a substance is the sum of the atomic masses of its atoms. The molar mass (gram-formula mass) of a substance equals one mole of that substance. 34: 3.1cc A compound is a substance composed of two or more different elements that are chemically combined in a fixed -6- Unit 3/4 – Moles / Stoichiometry proportion. A chemical compound can be broken down by chemical means. A chemical compound can be represented by a specific chemical formula and assigned a name based on the IUPAC system. 35: 3.3f The percent composition by mass of each element in a compound can be calculated mathematically 37: 3.3iv Calculate simple mole-mole stoichiometry problems, given a balanced equation 3.3vi Determine the mass of a given number of moles of a substance 3.3viii Calculate the formula mass and gram-formula mass 3.3ix Determine the number of moles of a substance, given its mass 38: 3.3d The empirical formula of a compound is the simplest whole-number ratio of atoms of the elements in a compound. It may be different from the molecular formula, which is the actual ratio of atoms in a molecule of that compound. 39: 3.2b Types of chemical reactions include synthesis, decomposition, single replacement, and double replacement 41: 3.2a A physical change results in the rearrangement of existing particles in a substance. A chemical change results in the formation of different substances with changed properties. 46: 3.2d An oxidation-reduction (redox) reaction involves the transfer of electrons (e-). 49: 5.3c Energy released during nuclear reactions is much greater than the energy released during chemical reactions. 66: 3.3c A balanced chemical equation represents conservation of atoms. The coefficients in a balanced chemical equation can be used to determine mole ratios in the reaction. January 2011 -7- Unit 3/4 – Moles / Stoichiometry 10: 3.3e The formula mass of a substance is the sum of the atomic masses of its atoms. The molar mass (gram-formula mass) of a substance equals one mole of that substance. 39: 3.2b Types of chemical reactions include synthesis, decomposition, single replacement, and double replacement 47: 3.3c A balanced chemical equation represents conservation of atoms. The coefficients in a balanced chemical equation can be used to determine mole ratios in the reaction 54: 3.3iv Calculate simple mole-mole stoichiometry problems, given a balanced equation 3.3vi Determine the mass of a given number of moles of a substance 3.3viii Calculate the formula mass and gram-formula mass 3.3ix Determine the number of moles of a substance, given its mass 55: Types of chemical reactions include synthesis, decomposition, single replacement, and double replacement 69: Types of chemical reactions include synthesis, decomposition, single replacement, and double replacement 71: 3.3iv Calculate simple mole-mole stoichiometry problems, given a balanced equation 76: 3.2i Oxidation numbers (states) can be assigned to atoms and ions. Changes in oxidation numbers indicate that oxidation and reduction have occurred. August 2010 -8- Unit 3/4 – Moles / Stoichiometry 7: 3.3c A balanced chemical equation represents conservation of atoms. The coefficients in a balanced chemical equation can be used to determine mole ratios in the reaction. 9: 3.3f The percent composition by mass of each element in a compound can be calculated mathematically. 18: 3.4e Equal volumes of gases at the same temperature and pressure contain an equal number of particles 36: 3.3iv Calculate simple mole-mole stoichiometry problems, given a balanced equation 3.3vi Determine the mass of a given number of moles of a substance 3.3viii Calculate the formula mass and gram-formula mass 3.3ix Determine the number of moles of a substance, given its mass 70: 3.2b Types of chemical reactions include synthesis, decomposition, single replacement, and double replacement 71: 3.3iv Calculate simple mole-mole stoichiometry problems, given a balanced equation 3.3vi Determine the mass of a given number of moles of a substance 3.3viii Calculate the formula mass and gram-formula mass 3.3ix Determine the number of moles of a substance, given its mass 72: 3.3viii Calculate the formula mass and gram-formula mass June 2010 -9- Unit 3/4 – Moles / Stoichiometry 17: 3.4e Equal volumes of gases at the same temperature and pressure contain an equal number of particles. 34: 36: 3.3d The empirical formula of a compound is the simplest whole-number ratio of atoms of the elements in a compound. It may be different from the molecular formula, which is the actual ratio of atoms in a molecule of that compound 56: M1.1B – Use algebraic and geometric representations to describe and compare data by measuring and recording experimental data and use data in calculations such as using appropriate equations and significant digits 73: 3.2b Types of chemical reactions include synthesis, decomposition, single replacement, and double replacement 75: 3.3iv Calculate simple mole-mole stoichiometry problems, given a balanced equation 3.3vi Determine the mass of a given number of moles of a substance 3.3viii Calculate the formula mass and gram-formula mass 3.3ix Determine the number of moles of a substance, given its mass - 10 - Unit 3/4 – Moles / Stoichiometry January 2010 9: 3.3e The formula mass of a substance is the sum of the atomic masses of its atoms. The molar mass (gram-formula mass) of a substance equals one mole of that substance 10: 3.2b Types of chemical reactions include synthesis, decomposition, single replacement, and double replacement 22: 3.4e Equal volumes of gases at the same temperature and pressure contain an equal number of particles 23: 5.2m Intermolecular forces created by the unequal distribution of charge result in varying degrees of attraction between molecules. Hydrogen bonding is an example of a strong intermolecular force 36: 3.3iv Calculate simple mole-mole stoichiometry problems, given a balanced equation 3.3vi Determine the mass of a given number of moles of a substance 3.3viii Calculate the formula mass and gram-formula mass 3.3ix Determine the number of moles of a substance, given its mass 37: 3.3d The empirical formula of a compound is the simplest whole-number ratio of atoms of the elements in a - 11 - Unit 3/4 – Moles / Stoichiometry compound. It may be different from the molecular formula, which is the actual ratio of atoms in a molecule of that compound 68: 3.3f The percent composition by mass of each element in a compound can be calculated mathematically August 2009 - 12 - Unit 3/4 – Moles / Stoichiometry 35: 3.2b Types of chemical reactions include synthesis, decomposition, single replacement, and double replacement 36: 3.3d The empirical formula of a compound is the simplest whole-number ratio of atoms of the elements in a compound. It may be different from the molecular formula, which is the actual ratio of atoms in a molecule of that compound 40: 3.1ppThe concentration of a solution may be expressed in molarity (M), percent by volume, percent by mass, or parts per million (ppm) 52: 3.3iv Calculate simple mole-mole stoichiometry problems, given a balanced equation 55: 3.3vi Determine the mass of a given number of moles of a substance 56: 3.3a In all chemical reactions there is a conservation of mass, energy, and charge 57: 3.3c A balanced chemical equation represents conservation of atoms. The coefficients in a balanced chemical equation can be used to determine mole ratios in the reaction 79: 3.3f The percent composition by mass of each element in a compound can be calculated mathematically June 2009 - 13 - Unit 3/4 – Moles / Stoichiometry 7: 3.3vi Determine the mass of a given number of moles of a substance 33: 3.3d The empirical formula of a compound is the simplest whole-number ratio of atoms of the elements in a compound. It may be different from the molecular formula, which is the actual ratio of atoms in a molecule of that compound 34: 3.2b Types of chemical reactions include synthesis, decomposition, single replacement, and double replacement 74: 3.3c A balanced chemical equation represents conservation of atoms. The coefficients in a balanced chemical equation can be used to determine mole ratios in the reaction 79: M1.1C – Use algebraic and geometric representations to describe and compare data by recognizing and converting various scales of measurement 80: 3.3iv Calculate simple mole-mole stoichiometry problems, given a balanced equation 3.3vi Determine the mass of a given number of moles of a substance 3.3viii Calculate the formula mass and gram-formula mass 3.3ix Determine the number of moles of a substance, given its mass 81: 3.2b Types of chemical reactions include synthesis, decomposition, single replacement, and double replacement 83: 3.3f The percent composition by mass of each element in a compound can be calculated mathematically January 2009 - 14 - Unit 3/4 – Moles / Stoichiometry - 15 - Unit 3/4 – Moles / Stoichiometry - 16 - Unit 3/4 – Moles / Stoichiometry 11: 3.3a In all chemical reactions there is a conservation of mass, energy, and charge 17: 3.4e Equal volumes of gases at the same temperature and pressure contain an equal number of particles 22: 3.3b In a redox reaction the number of electrons lost is equal to the number of electrons gained 36: 3.3d The empirical formula of a compound is the simplest whole-number ratio of atoms of the elements in a compound. It may be different from the molecular formula, which is the actual ratio of atoms in a molecule of that compound 54: 3.3iv Calculate simple mole-mole stoichiometry problems, given a balanced equation 3.3vi Determine the mass of a given number of moles of a substance 3.3viii Calculate the formula mass and gram-formula mass 3.3ix Determine the number of moles of a substance, given its mass 69: 3.3c A balanced chemical equation represents conservation of atoms. The coefficients in a balanced chemical equation can be used to determine mole ratios in the reaction 71: 3.3iv Calculate simple mole-mole stoichiometry problems, given a balanced equation 3.3vi Determine the mass of a given number of moles of a substance - 17 - Unit 3/4 – Moles / Stoichiometry 3.3viii Calculate the formula mass and gram-formula mass 3.3ix Determine the number of moles of a substance, given its mass 76: 3.3f The percent composition by mass of each element in a compound can be calculated mathematically 78: 3.3f The percent composition by mass of each element in a compound can be calculated mathematically August 2008 34: 3.3a In all chemical reactions there is a conservation of mass, energy, and charge. 3.3c A balanced chemical equation represents conservation of atoms. The coefficients in a balanced chemical equation can be used to determine mole ratios in the reaction - 18 - Unit 3/4 – Moles / Stoichiometry 35: 3.3f The percent composition by mass of each element in a compound can be calculated mathematically 63: 3.3c A balanced chemical equation represents conservation of atoms. The coefficients in a balanced chemical equation can be used to determine mole ratios in the reaction 74: 3.3vi Determine the mass of a given number of moles of a substance 75: 3.3iv Calculate simple mole-mole stoichiometry problems, given a balanced equation 76: 3.3d The empirical formula of a compound is the simplest whole-number ratio of atoms of the elements in a compound. It may be different from the molecular formula, which is the actual ratio of atoms in a molecule of that compound June 2008 - 19 - Unit 3/4 – Moles / Stoichiometry 7: 3.3a In all chemical reactions there is a conservation of mass, energy, and charge 19: 3.4e Equal volumes of gases at the same temperature and pressure contain an equal number of particles 35: M1.1B – Use algebraic and geometric representations to describe and compare data by measuring and recording experimental data and use data in calculations such as calculating percent error 36: 3.3iv Calculate simple mole-mole stoichiometry problems, given a balanced equation 3.3vi Determine the mass of a given number of moles of a substance 3.3viii Calculate the formula mass and gram-formula mass 3.3ix Determine the number of moles of a substance, given its mass 42: 3.3iv Calculate simple mole-mole stoichiometry problems, given a balanced equation 47: 3.3d The empirical formula of a compound is the simplest whole-number ratio of atoms of the elements in a compound. It may be different from the molecular formula, which is the actual ratio of atoms in a molecule of that - 20 - Unit 3/4 – Moles / Stoichiometry compound 56: 3.3iv Calculate simple mole-mole stoichiometry problems, given a balanced equation 3.3vi Determine the mass of a given number of moles of a substance 3.3viii Calculate the formula mass and gram-formula mass 3.3ix Determine the number of moles of a substance, given its mass 67: 3.3iv Calculate simple mole-mole stoichiometry problems, given a balanced equation 3.3vi Determine the mass of a given number of moles of a substance 3.3viii Calculate the formula mass and gram-formula mass 3.3ix Determine the number of moles of a substance, given its mass 74: 3.3c A balanced chemical equation represents conservation of atoms. The coefficients in a balanced chemical equation can be used to determine mole ratios in the reaction January 2008 34: 3.3viii Calculate the formula mass and gram-formula mass 54: 3.3d The empirical formula of a compound is the simplest whole-number ratio of atoms of the elements in a compound. It may be different from the molecular formula, which is the actual ratio of atoms in a molecule of that compound 77: 3.3f The percent composition by mass of each element in a compound can be calculated mathematically 78: 3.2b Types of chemical reactions include synthesis, decomposition, single replacement, and double replacement 79: 3.3iv Calculate simple mole-mole stoichiometry problems, given a balanced equation - 21 - Unit 3/4 – Moles / Stoichiometry August 2007 10: 3.3c A balanced chemical equation represents conservation of atoms. The coefficients in a balanced chemical equation can be used to determine mole ratios in the reaction 12: 3.1cc A compound is a substance composed of two or more different elements that are chemically combined in a fixed proportion. A chemical compound can be broken down by chemical means. A chemical compound can be represented by a specific chemical formula and assigned a name based on the IUPAC system 17: 3.1cc A compound is a substance composed of two or more different elements that are chemically combined in a fixed proportion. A chemical compound can be broken down by chemical means. A chemical compound can be represented by a specific chemical formula and assigned a name based on the IUPAC system. 35: 3.3d The empirical formula of a compound is the simplest whole-number ratio of atoms of the elements in a compound. It may be different from the molecular formula, which is the actual ratio of atoms in a molecule of that compound 70: 3.2b Types of chemical reactions include synthesis, decomposition, single replacement, and double replacement 71: 3.3viii Calculate the formula mass and gram-formula mass 72: 3.1cc A compound is a substance composed of two or more different elements that are chemically combined in a fixed proportion. A chemical compound can be broken down by chemical means. A chemical compound can be represented by a specific chemical formula and assigned a name based on the IUPAC system - 22 - Unit 3/4 – Moles / Stoichiometry June 2007 7: 3.1cc A compound is a substance composed of two or more different elements that are chemically combined in a fixed proportion. A chemical compound can be broken down by chemical means. A chemical compound can be represented by a specific chemical formula and assigned a name based on the IUPAC system 9: 3.1cc A compound is a substance composed of two or more different elements that are chemically combined in a fixed proportion. A chemical compound can be broken down by chemical means. A chemical compound can be represented by a specific chemical formula and assigned a name based on the IUPAC system 36: 3.1cc A compound is a substance composed of two or more different elements that are chemically combined in a fixed proportion. A chemical compound can be broken down by chemical means. A chemical compound can be represented by a specific chemical formula and assigned a name based on the IUPAC system 38: 3.3d The empirical formula of a compound is the simplest whole-number ratio of atoms of the elements in a compound. It may be different from the molecular formula, which is the actual ratio of atoms in a molecule of that compound 55: 3.3f The percent composition by mass of each element in a compound can be calculated mathematically 59: 3.3c A balanced chemical equation represents conservation of atoms. The coefficients in a balanced chemical equation can be used to determine mole ratios in the reaction 60: 3.2i Oxidation numbers (states) can be assigned to atoms and ions. Changes in oxidation numbers indicate that oxidation and reduction have occurred. - 23 - Unit 3/4 – Moles / Stoichiometry January 2007 10: 3.3c A balanced chemical equation represents conservation of atoms. The coefficients in a balanced chemical equation can be used to determine mole ratios in the reaction 17: 3.1cc A compound is a substance composed of two or more different elements that are chemically combined in a fixed proportion. A chemical compound can be broken down by chemical means. A chemical compound can be represented by a specific chemical formula and assigned a name based on the IUPAC system 18: 3.1t The proportions of components in a mixture can be varied. Each component in a mixture retains its original properties. 35: 3.3viii Calculate the formula mass and gram-formula mass 36: 3.3iv Calculate simple mole-mole stoichiometry problems, given a balanced equation 54: M1.1B – Use algebraic and geometric representations to describe and compare data by measuring and recording experimental data and use data in calculations such as calculating percent error - 24 - Unit 3/4 – Moles / Stoichiometry August 2006 - 25 - Unit 3/4 – Moles / Stoichiometry - 26 - Unit 3/4 – Moles / Stoichiometry 6: 3.1cc A compound is a substance composed of two or more different elements that are chemically combined in a fixed proportion. A chemical compound can be broken down by chemical means. A chemical compound can be represented by a specific chemical formula and assigned a name based on the IUPAC system. 9: 5.2d Electron-dot diagrams (Lewis structures) can represent the valence electron arrangement in elements, compounds, and ions 37: 3.3f The percent composition by mass of each element in a compound can be calculated mathematically 38: 3.3iv Calculate simple mole-mole stoichiometry problems, given a balanced equation 44: 4.1b Chemical and physical changes can be exothermic or endothermic. 82: 3.3c A balanced chemical equation represents conservation of atoms. The coefficients in a balanced chemical equation can be used to determine mole ratios in the reaction 83: 3.2b Types of chemical reactions include synthesis, decomposition, single replacement, and double replacement 84: 3.3ix Determine the number of moles of a substance, given its mass 85: 3.2x Use an activity series to determine whether a redox reaction is spontaneous June 2006 - 27 - Unit 3/4 – Moles / Stoichiometry 6: 3.1cc A compound is a substance composed of two or more different elements that are chemically combined in a fixed proportion. A chemical compound can be broken down by chemical means. A chemical compound can be represented by a specific chemical formula and assigned a name based on the IUPAC system 7: 3.3c A balanced chemical equation represents conservation of atoms. The coefficients in a balanced chemical equation can be used to determine mole ratios in the reaction - 28 - Unit 3/4 – Moles / Stoichiometry 33: 3.3d The empirical formula of a compound is the simplest whole-number ratio of atoms of the elements in a compound. It may be different from the molecular formula, which is the actual ratio of atoms in a molecule of that compound 35: 3.3f The percent composition by mass of each element in a compound can be calculated mathematically 51: 3.1cc A compound is a substance composed of two or more different elements that are chemically combined in a fixed proportion. A chemical compound can be broken down by chemical means. A chemical compound can be represented by a specific chemical formula and assigned a name based on the IUPAC system 52: 3.2x Use an activity series to determine whether a redox reaction is spontaneous 69: 3.3c A balanced chemical equation represents conservation of atoms. The coefficients in a balanced chemical equation can be used to determine mole ratios in the reaction 70: 3.4ii Solve problems, using the combined gas laws January 2006 8: 3.1cc A compound is a substance composed of two or more different elements that are chemically combined in a fixed proportion. A chemical compound can be broken down by chemical means. A chemical compound can be represented by a specific chemical formula and assigned a name based on the IUPAC system 10: 3.3c A balanced chemical equation represents conservation of atoms. The coefficients in a balanced chemical equation can be used to determine mole ratios in the reaction 33: 3.3f The percent composition by mass of each element in a compound can be calculated mathematically 36: 3.3viii Calculate the formula mass and gram-formula mass 57: 3.3ix Determine the number of moles of a substance, given its mass - 29 - Unit 3/4 – Moles / Stoichiometry 75: 3.3viii Calculate the formula mass and gram-formula mass 76: 3.1ppThe concentration of a solution may be expressed in molarity (M), percent by volume, percent by mass, or parts per million (ppm) 77: 3.1xxx Describe the preparation of a solution, given the molarity August 2005 9: 3.1cc A compound is a substance composed of two or more different elements that are chemically combined in a fixed proportion. A chemical compound can be broken down by chemical means. A chemical compound can be represented by a specific chemical formula and assigned a name based on the IUPAC system 31: 3.3f The percent composition by mass of each element in a compound can be calculated mathematically - 30 - Unit 3/4 – Moles / Stoichiometry 36: 3.3iv Calculate simple mole-mole stoichiometry problems, given a balanced equation 38: 3.2b Types of chemical reactions include synthesis, decomposition, single replacement, and double replacement 48: 3.2b Types of chemical reactions include synthesis, decomposition, single replacement, and double replacement 51: 3.3viii Calculate the formula mass and gram-formula mass 52: 3.3d The empirical formula of a compound is the simplest whole-number ratio of atoms of the elements in a compound. It may be different from the molecular formula, which is the actual ratio of atoms in a molecule of that compound 68: 3.3c A balanced chemical equation represents conservation of atoms. The coefficients in a balanced chemical equation can be used to determine mole ratios in the reaction 69: 3.3ix Determine the number of moles of a substance, given its mass 70: Density equation – General math standard June 2005 9: 3.1cc A compound is a substance composed of two or more different elements that are chemically combined in a fixed proportion. A chemical compound can be broken down by chemical means. A chemical compound can be represented by a specific chemical formula and assigned a name based on the IUPAC system 36: 3.1cc A compound is a substance composed of two or more different elements that are chemically combined in a fixed proportion. A chemical compound can be broken down by chemical means. A chemical compound can be represented by a specific chemical formula and assigned a name based on the IUPAC system 37: 3.3d The empirical formula of a compound is the simplest whole-number ratio of atoms of the elements in a compound. It may be different from the molecular formula, which is the actual ratio of atoms in a molecule of that compound 39: 3.3f The percent composition by mass of each element in a compound can be calculated mathematically 54: 3.3iv Calculate simple mole-mole stoichiometry problems, given a balanced equation January 2005 - 31 - Unit 3/4 – Moles / Stoichiometry 8: 3.1cc A compound is a substance composed of two or more different elements that are chemically combined in a fixed proportion. A chemical compound can be broken down by chemical means. A chemical compound can be represented by a specific chemical formula and assigned a name based on the IUPAC system 9: 3.3f The percent composition by mass of each element in a compound can be calculated mathematically 19: 3.4e Equal volumes of gases at the same temperature and pressure contain an equal number of particles 34: 3.3d The empirical formula of a compound is the simplest whole-number ratio of atoms of the elements in a compound. It may be different from the molecular formula, which is the actual ratio of atoms in a molecule of that compound 35: 3.3a In all chemical reactions there is a conservation of mass, energy, and charge. 3.3c A balanced chemical equation represents conservation of atoms. The coefficients in a balanced chemical equation can be used to determine mole ratios in the reaction 68: 3.1cc A compound is a substance composed of two or more different elements that are chemically combined in a fixed proportion. A chemical compound can be broken down by chemical means. A chemical compound can be represented by a specific chemical formula and assigned a name based on the IUPAC system 69: 3.3f The percent composition by mass of each element in a compound can be calculated mathematically August 2004 - 32 - Unit 3/4 – Moles / Stoichiometry - 33 - Unit 3/4 – Moles / Stoichiometry 8: 3.3a In all chemical reactions there is a conservation of mass, energy, and charge 29: 3.1zz Titration is a laboratory process in which a volume of a solution of known concentration is used to determine the concentration of another solution. 37: 3.3d The empirical formula of a compound is the simplest whole-number ratio of atoms of the elements in a compound. It may be different from the molecular formula, which is the actual ratio of atoms in a molecule of that compound 50: M1.1B – Use algebraic and geometric representations to describe and compare data by measuring and recording experimental data and use data in calculations such as using appropriate equations and significant digits 52: 3.3iv Calculate simple mole-mole stoichiometry problems, given a balanced equation 53: 3.1ppThe concentration of a solution may be expressed in molarity (M), percent by volume, percent by mass, or parts per million (ppm) 61: 3.3viii Calculate the formula mass and gram-formula mass 62: 3.3vi Determine the mass of a given number of moles of a substance 66: 3.3c A balanced chemical equation represents conservation of atoms. The coefficients in a balanced chemical equation can be used to determine mole ratios in the reaction 67: 4.1c Energy released or absorbed during a chemical reaction can be represented by a potential energy diagram - 34 - Unit 3/4 – Moles / Stoichiometry 68: 3.2i Oxidation numbers (states) can be assigned to atoms and ions. Changes in oxidation numbers indicate that oxidation and reduction have occurred 69: M1.1C – Use algebraic and geometric representations to describe and compare data by recognizing and converting various scales of measurement June 2004 7: 3.3c A balanced chemical equation represents conservation of atoms. The coefficients in a balanced chemical equation can be used to determine mole ratios in the reaction 8: 3.3f The percent composition by mass of each element in a compound can be calculated mathematically 38: 3.3d The empirical formula of a compound is the simplest whole-number ratio of atoms of the elements in a compound. It may be different from the molecular formula, which is the actual ratio of atoms in a molecule of that compound 51: 3.2b Types of chemical reactions include synthesis, decomposition, single replacement, and double replacement - 35 - Unit 3/4 – Moles / Stoichiometry 52: 3.3a In all chemical reactions there is a conservation of mass, energy, and charge. 3.3c A balanced chemical equation represents conservation of atoms. The coefficients in a balanced chemical equation can be used to determine mole ratios in the reaction 53: 3.3iv Calculate simple mole-mole stoichiometry problems, given a balanced equation 81: 5.2d Electron-dot diagrams (Lewis structures) can represent the valence electron arrangement in elements, compounds, and ions 82: 3.1cc A compound is a substance composed of two or more different elements that are chemically combined in a fixed proportion. A chemical compound can be broken down by chemical means. A chemical compound can be represented by a specific chemical formula and assigned a name based on the IUPAC system 83: 3.1xxix Calculate solution concentration in molarity (M), percent mass, and parts per million (ppm) 84: S3.4 Using results of the test and through public discussion, revise the explanation and contemplate additional research January 2004 - 36 - Unit 3/4 – Moles / Stoichiometry 6: 3.1cc A compound is a substance composed of two or more different elements that are chemically combined in a fixed proportion. A chemical compound can be broken down by chemical means. A chemical compound can be represented by a specific chemical formula and assigned a name based on the IUPAC system 36: 3.3d The empirical formula of a compound is the simplest whole-number ratio of atoms of the elements in a compound. It may be different from the molecular formula, which is the actual ratio of atoms in a molecule of that compound 39: 3.3a In all chemical reactions there is a conservation of mass, energy, and charge. 3.3c A balanced chemical equation represents conservation of atoms. The coefficients in a balanced chemical equation can be used to determine mole ratios in the reaction 51: 3.3c A balanced chemical equation represents conservation of atoms. The coefficients in a balanced chemical equation can be used to determine mole ratios in the reaction 52: 3.3iv Calculate simple mole-mole stoichiometry problems, given a balanced equation 75: 3.3viii Calculate the formula mass and gram-formula mass 76: 3.3f The percent composition by mass of each element in a compound can be calculated mathematically - 37 - Unit 3/4 – Moles / Stoichiometry August 2003 6: 3.1cc A compound is a substance composed of two or more different elements that are chemically combined in a fixed proportion. A chemical compound can be broken down by chemical means. A chemical compound can be represented by a specific chemical formula and assigned a name based on the IUPAC system 8: 3.3f The percent composition by mass of each element in a compound can be calculated mathematically 10: 3.3a In all chemical reactions there is a conservation of mass, energy, and charge 39: 3.3iv Calculate simple mole-mole stoichiometry problems, given a balanced equation 42: 3.3d The empirical formula of a compound is the simplest whole-number ratio of atoms of the elements in a compound. It may be different from the molecular formula, which is the actual ratio of atoms in a molecule of that compound June 2003 - 38 - Unit 3/4 – Moles / Stoichiometry 8: 3.3d The empirical formula of a compound is the simplest whole-number ratio of atoms of the elements in a compound. It may be different from the molecular formula, which is the actual ratio of atoms in a molecule of that compound 10: 3.3f The percent composition by mass of each element in a compound can be calculated mathematically 19: 3.1cc A compound is a substance composed of two or more different elements that are chemically combined in a fixed proportion. A chemical compound can be broken down by chemical means. A chemical compound can be represented by a specific chemical formula and assigned a name based on the IUPAC system 20: 3.3iv Calculate simple mole-mole stoichiometry problems, given a balanced equation 59: 3.3a In all chemical reactions there is a conservation of mass, energy, and charge January 2003 - 39 - Unit 3/4 – Moles / Stoichiometry 8: 3.3f The percent composition by mass of each element in a compound can be calculated mathematically 13: 3.3d The empirical formula of a compound is the simplest whole-number ratio of atoms of the elements in a compound. It may be different from the molecular formula, which is the actual ratio of atoms in a molecule of that compound 22: 3.3f The percent composition by mass of each element in a compound can be calculated mathematically 42: 3.3iv Calculate simple mole-mole stoichiometry problems, given a balanced equation 43: M1.1B – Use algebraic and geometric representations to describe and compare data by measuring and recording experimental data and use data in calculations such as calculating percent error 48: 3.3c A balanced chemical equation represents conservation of atoms. The coefficients in a balanced chemical equation can be used to determine mole ratios in the reaction - 40 - Unit 3/4 – Moles / Stoichiometry 1) How many moles of sodium atoms correspond to 1.56x1021 atoms of sodium? 2) How many moles of Al atoms are needed to combine with 1.58 mol of O atoms to make aluminum oxide, Al2O3? 3) How many moles of Al are in 2.16 mol of Al2O3? 4) Aluminum sulfate, Al2(SO4)3, is a compound used in sewage treatment plants. a. Construct a pair of conversion factors that relate moles of aluminum to moles of sulfur for this compound b. Construct a pair of conversion factors that relate moles of sulfur to moles of Al2(SO4)3 c. How many moles of Al are in a sample of this compound if the sample also contains 0.900 mol S? d. How many moles of S are in 1.16 mol Al2(SO4)3? 5) How many moles of H2 and N2 can be formed by the decomposition of 0.145 mol of ammonia, NH3? 6) What is the total number of atoms in 0.260 mol of glucose, C6H12O6? 7) What is the mass of 1.00 mol of each of the following elements? a. Sodium b. Sulfur c. Chlorine 8) Determine the mass in grams of each of the following: a. 1.35 mol Fe b. 24.5 mol O c. 0.876 mol Ca d. 1.25 mol Ca3(PO4)2 e. 0.625 mol Fe(NO3)3 f. 0.600 mol C4H10 g. 1.45 mol (NH4)2CO3 9) Calculate the number of moles of each compound: a. 21.5 g CaCO3 b. 1.56 g NH3 c. 16.8 g Sr(NO3)2 d. 6.98 g Na2CrO4 41 Unit 3/4 – Moles / Stoichiometry Percent composition and empirical formulas 10) Calculate the percentage composition by mass of each element in the following compounds: a. NaH2PO4 b. NH4H2PO4 c. (CH3)2CO 11) Phencyclidine is C17H25N. A sample suspected of being this illicit drug was found to have a percentage composition of 83.71% C, 10.42% H, and 5.61% N. Do these data acceptably match the theoretical data for phencyclidine? 12) How many grams of O are combined with 7.14x1021 atoms of N in the compound N2O5? 13) Quantitative analysis of a sample of sodium pertechnetate with a mass of 0.896g found 0.111g Na and 0.477g technetium (Tc). The remainder was oxygen. Calculate the empirical formula of sodium pertechnetate, NaxTcyOz. 14) A substance was found to be composed of 22.9% Na, 21.5% B, and 55.7% O. What is the empirical formula of this compound? 15) When 0.684 g of an organic compound containing only C, H, and O was burned in oxygen 1.312g CO2 and 0.805g H2O were obtained. What is the empirical formula of the compound? Balancing equations 16) Write the equation that expresses in acceptable chemical shorthand the following statement: “Iron can be made to react with molecular oxygen (O2) to give iron oxide with the formula Fe2O3” 17) Balance the following reactions: a. Ca(OH)2 + HCl CaCl2 + H2O b. AgNO3 + CaCl2 Ca(NO3)2 +AgCl c. Fe2O3 + C Fe + CO3 d. NaHCO3 + H2SO4 Na2SO4 + H2O + CO2 e. C4H10 + O2 CO2 +H2O f. Mg(OH)2 + HBr MgBr2 + H2O g. Al2O3 + H2SO4 Al2(SO4)3 + H2O h. KHCO3 + H3PO4 K2HPO4 + H2O + CO2 i. C9H10O + O2 CO2 + H2O Stoichiometry/limiting reactants 18) Chlorine is used by textile manufacturers to bleach cloth. Excess chlorine is destroyed by its reaction with sodium thiosulfate, Na2S2O3: Na2S2O3(aq) + 4Cl2(g) + 5H2O(aq) 2NaHSO4(aq) + 8HCl(aq) 42 Unit 3/4 – Moles / Stoichiometry a. b. c. d. How many moles of Na2S2O3 are needed to react with 0.12mol of Cl2? How many moles of HCl can form from 0.12mol of Cl2? How many moles of H2O are required for the reaction of 0.12mol of Cl2? How many moles of H2O react if 0.24mol HCl is formed? 19) The incandescent white of a fireworks display is caused by the reaction of phosphorous with O2 to give P4O10. a. Write the balanced chemical equation for the reaction. b. How many grams of O2 are needed to combine with 6.85g of P? c. How many grams of P4O10 can be made from 8.00g of O2? d. How many grams of P are needed to make 7.46g P4O10? 20) In dilute nitric acid, HNO3, copper metal dissolves according to the following equation: 3Cu(s) + 8HNO3(aq) 3Cu(NO3)2(aq) + 2NO(g) + 4H2O(aq) How many grams of HNO3 are needed to dissolve 11.45g of Cu? 21) The reaction of powdered aluminum and iron(II)oxide, 2Al(s) + Fe2O3(s) Al2O3(s) + 2Fe(l) produces so much heat the iron that forms is molten. Because of this, railroads use the reaction to provide molten steel to weld steel rails together when laying track. Suppose that in one batch of reactants 4.20mol Al was mixed with 1.75mol Fe2O3. a. Which reactant, if either, was the limiting reactant? b. Calculate the mass of iron (in grams) that can be formed from this mixture of reactants. 22) Silver nitrate, AgNO3, reacts with iron(III) chloride, FeCl3, to give silver chloride, AgCl, and iron(III) nitrate, Fe(NO3)3. A solution containing 18.0g AgNO3 was mixed with a solution containing 32.4g FeCl3. How many grams of which reactant remains after the reaction is over? Theoretical and percent yield 23) Barium sulfate, BaSO4, is made by the following reaction: Ba(NO3)2(aq) + Na2SO4(aq) BaSO4(s) + 2NaNO3(aq) An experiment was begun with 75.00g of Ba(NO3)2 and an excess of Na2SO4. After collecting and drying the product, 63.45g BaSO4 was obtained. Calculate the theoretical yield and percent yield of BaSO4. 24) Aluminum sulfate can be made by the following reaction: 2AlCl3(aq) + 3H2SO4(aq) Al2(SO4)3(aq) + 6HCl(aq) It is quite soluble in water, so to isolate it the solution has to be evaporated to dryness. This drives off the volatile HCl, but the residual solid has to be treated to a little over 200C to drive off all the water. In one experiment, 25.0g of AlCl3 was mixed with 30.0g H2SO4. Eventually, 28.46g of pure Al2(SO4)3 was isolated. Calculate the percent yield. 43 Unit 3/4 – Moles / Stoichiometry Answers 2.59x103mol Na atoms 1.05mol Al 4.32mol Al a. 2mol Al/3mol S b. 3mol S/1mol Al2(SO4)3 c. 0.600mol Al d. 3.48mol S 5) 0.0725mol N2 and 0.218mol H2 6) 3.76x1024 atoms 7) a. 23.0g Na b. 32.1g S c. 35.3g Cl 8) a. 75.4g Fe b. 392g O c. 35.1g Ca d. 388g Ca3(PO4)2 e. 151g Fe(NO3)2 f. 34.9g C4H10 g. 139g (NH4)2CO3 9) a. 0.215mol b. 0.0916mol c. 0.0794mol d. 8 4.31x10 mol 10) a. 19.2% Na, 1.68% H, 25.8% P, 53.3% O b. 12.2% N, 5.26% H, 26.9% P, 55.6%O c. 62.0% C, 10.4% H, 27.6% O 11) Theoretical data (83.89% C, 10.35% H, 5.76% N) are consistent with experimental results. 12) 0.474g O 13) NaTcO4 14) Na2B4O7 15) C2H6O 16) 4Fe + 3O2 2Fe2O3 17) a. Ca(OH)2 + 2HCl CaCl2 + 2H2O b. 2AgNO3 + CaCl2 Ca(NO3)2 + 2AgCl c. 2Fe2O3 + 3C 4Fe + 3CO3 d. 2NaHCO3 + H2SO4 Na2SO4 + 2H2O + 2CO2 e. 2C4H10 + 13O2 8CO2 + 10H2O f. Mg(OH)2 + 2HBr MgBr2 + 2H2O g. Al2O3 + 3H2SO4 Al2(SO4)3 + 3H2O h. 2KHCO3 + H3PO4 K2HPO4 + 2H2O + 2CO2 i. C9H10O + 14O2 9CO2 + 10H2O 18) a. 0.030mol Na2S2O3 b. 0.24mol HCl c. 0.15mol H2O d. 0.15mol H2O 19) a. 4P + 5O2 P4O10 b. 8.85g O2 c. 14.2g P4O10 d. 3.26g P 20) 30.31g HNO3 21) a. limiting reactant is Fe2O3 b. 195g Fe is formed 22) 26.7g of FeCl3 are left over 23) theoretical yield = 66.98g BaSO4, % yield = 94.73% 24) % yield = 88.74% 1) 2) 3) 4) 1. Which unit is most analogous to a mole? 44 Unit 3/4 – Moles / Stoichiometry a fluid ounce a dozen a mile a degree Fahrenheit a pound 2. In order to balance the equation C2H6 + O2 ---> H2O + CO2, you should change the subscript of O in water to 2 to help balance the O. add O2 to the product side to help balance the O in the equation. change the coefficients. add H2 to the products to balance H. 3. Balance the following equation: B10H18 + O2 4. B10H18 + 7O2 5B2O3 + 9H2O B10H18 + 9O2 5B2O3 + 9H2O B10H18 + 19O2 10B2O3 + 9H2O B10H18 + 12O2 5B2O3 + 9H2O Balance the following equation and indicate whether it is a combustion, combination, or decomposition reaction. H2O2(l) 5. B2O3 + H2O H2O(l) + O2(g) 2H2O2(l) 2H2O(l) + O2(g), decomposition reaction H2O2(l) H2O(l) + O2(g), decomposition reaction H2O2(l) H2O(l) + (1/2)O2(g), decomposition reaction H2O2(l) H2O(l) + (1/2)O2(g), combination reaction H2O2(l) H2O(l) + O2(g), combustion reaction Convert the following description into a balanced equation: 45 Unit 3/4 – Moles / Stoichiometry When ammonia gas, NH3(g), is passed over hot sodium, hydrogen gas is released and sodium amide, NaNH2, is formed as a solid product. Be sure to indicate the state of each element or compound. 2NH3 + Na NH3 + Na 2NH3(g) + 2Na(s) 2NH3 + 2Na 2NH3(g) + 2Na(s) 6. 2NaNH2 + H2 NaNH2 + H2 2NaNH2(s) + H2(g) 2NaNH2 + H2 2NaNH2(s) + H2(l) How many molecules of CH2O are in 30.0 g of CH2O? 1.00 6.02 x 1023 1.81 x 1025 5.32 x 10-23 7. Calculate the mass in mg of Na+ in 10.0 g of sodium carbonate. 4.34 x 103 mg 46.0 mg 2.77 mg 4.34 mg 0.23 mg 8. How many carbon atoms are there in 200 molecules of C3H8O? 1.20 x 1026 200 3.61 x 1026 600 9. A sample of vitamin A, C20H30O, contains 4.0 x 1022 atoms of carbon. How many atoms of hydrogen and how many molecules of vitamin A does it contain? 46 Unit 3/4 – Moles / Stoichiometry 6.0 x 1022 atoms of H, 8.0 x 1023 molecules of vitamin A 4.0 x 1022 atoms of H, 4.0 x 1022 molecules of vitamin A 6.0 x 1022 atoms of H, 4.0 x 1022 molecules of vitamin A 6.0 x 1022 atoms of H, 2.0 x 1021 molecules of vitamin A 10 . The element zinc consists of five isotopes with masses 63.929, 65.926, 66.927, 67.925, and 69.925 amu. The relative abundances of these five isotopes are 48.89, 27.81, 4.110, 18.57, and 0.62 percent, respectively. From these data calculate the average atomic mass of zinc. 66.93 amu 65.39 amu 66.927 amu 65.389 amu 63.93 amu 11 . What is the mass in grams of 0.257 mol of sucrose, C12H22O11? 12.5 g 88.0 g 7.51 x 10-4 g 8.80 g 342 g 12 . Determine the approximate formula weight of the following: Ca(C2H3O2)2 69 158 152 99 94 Problem Solving Center Homework 2 1. What numbers would properly balance the reaction ___ C12H26 + ___ O2 ---> ___ CO2 + ___ H2O ? 47 Unit 3/4 – Moles / Stoichiometry 1, 25, 12, 13 2, 37, 24, 26 2, 37, 24, 13 1, 38, 12, 26 2. Complete the following statement: _____________ are in 10.0 moles of C10H8. 6.022 x 1024 atoms of C 10.0 moles of C 8.00 moles of H 4.818 x 1024 atoms of H 4.818 x 1025 atoms of H 3. Balance the following equation: Al + Cr2O3 Al2O3 + Cr 2Al + Cr2O3 Al2O3 + 2Cr 2Al + Cr2O3 Al2O3 + Cr Al + Cr2O3 Al2O3 + 2Cr 4Al + 2Cr2O3 4. Balance the following equation: C6H14O + O2 2C6H14O + 18O2 2C6H14O + 4O2 C6H14O + (19/2)O2 C6H14O + 9O2 2C6H14O + 19O2 5. 2Al2O3 + 4Cr CO2 + H2O 12CO2 + 14H2O 2CO2 + 2H2O 6CO2 + 7H2O 6CO2 + 7H2O 12CO2 + 14H2O Balance the following equation and indicate whether it is a combustion, combination, or decomposition reaction. 48 Unit 3/4 – Moles / Stoichiometry H2O2 + SO2 H2SO4 H2O2 + SO2 H2SO4, decomposition reaction 2H2O2 + SO2 H2SO4, decomposition reaction 2H2O2 + SO2 H2SO4, combination reaction H2O2 + SO2 6. H2SO4, combination reaction Calculate the number of molecules in 6.2 g of formaldehyde, CH2O. 3.7 x 1024 1.2 x 1023 6.0 x 1023 2.4 x 1023 7. Calculate the mass in grams of 0.0112 mol of -fructose, C6H12O6. 180 g 1.12 g 2.02 g 0.0112 g 8. A sample of glucose, C6H12O6, contains 4.0 x 1022 atoms of carbon. How many atoms of hydrogen and how many molecules of glucose does it contain? 8.0 x 1022 atoms of H, 6.7 x 1021 molecules of glucose 8.0 x 1022 atoms of H, 8.0 x 1022 molecules of glucose 8.0 x 1022 atoms of H, 2.4 x 1023 molecules of glucose 4.0 x 1022 atoms of H, 4.0 x 1022 molecules of glucose 8.0 x 1022 atoms of H, 4.0 x 1022 molecules of glucose 9. The element oxygen consists of three isotopes with masses 15.994915, 16.999133, and 17.99916. The relative abundances of these three isotopes are 99.7587, 0.0374, and 0.2039, respectively. From these data calculate the average atomic mass of oxygen. 49 Unit 3/4 – Moles / Stoichiometry 15.9930 15.999377 15.994915 16.0 15.9994 15.9563 10 . Calculate the number of atoms in 48.0 g glucose, C6H12O6. 1.60 x 1023 9.24 x 1025 3.85 x 1024 2.89 x 1025 1.64 x 1016 11 . Calculate the molecular weight of xenon tetrafluoride, XeF4, a colorless, crystalline compound at room temperature. 601.2 150.3 76 169.3 207.3 12 . How many iron ions (Fe3+) are present in 43.6 g FeCl3? 2.63 x 1025 4.86 x 1023 0.807 3.72 1.62 x 1023 Problem Solving Center Homework 3 50 Unit 3/4 – Moles / Stoichiometry 1. Which of these samples contains the most atoms? a gram of germanium a gram of francium a gram of americium a gram of gallium a gram of europium 2. What is the percent yield of CaO in the reaction CaCO 3 ---> CaO + CO2 if 5.33 g of CaO are obtained when 10.0 g of CaCO3 are used? 5.60 percent 64.7 percent 53.3 percent 5.33 percent 95.2 percent 3. Calculate the percentage of carbon present in cadaverine, C5H14N2, a compound present in rotting meat. 58.8 percent C 51 Unit 3/4 – Moles / Stoichiometry 68.2 percent C 67.4 percent C 51.7 percent C 4. What is the empirical formula of a compound that contains 7.989 g of carbon and 2.011 g of hydrogen? C2H5 C8H2 CH3 C3H C2H6 5. Give the empirical formula of the following compound if a sample contains 40.0 percent C, 6.7 percent H, and 53.3 percent O by mass. C2H4O2 CH20 C3H6O3 C6HO8 6. Determine the empirical formula of a compound that contains 52.9 percent aluminum and 47.1 percent oxygen. AlO Al2O3 Al4O6 Al0.53O0.47 Al3O2 7. In making H2O from hydrogen and oxygen, if we start with 4.6 mol of hydrogen and 3.1 mol of oxygen, how many moles of water can be produced and what remains unreacted? 2.3 mol of water would be produced, with 0.8 mol of O2 remaining. 4.6 mol of water would be produced, with 0.8 mol of O2 remaining. 7.7 mol of water would be produced, with 0.0 mol of O2 remaining. 4.6 mol of water would be produced, with 0.0 mol of O2 remaining. 3.1 mol of water would be produced, with 1.5 mol of O2 remaining. 52 Unit 3/4 – Moles / Stoichiometry 8. Automotive airbags inflate when sodium azide, NaN 3, rapidly decomposes to its component elements via the reaction 2NaN3 2Na + 3N2. How many grams of sodium azide are required to form 5.00 g of nitrogen gas? 3.33 g 11.61 g 15.48 g 7.74 g 9. What is the molecular formula of the following compound? empirical formula C2H3, molar mass 54 g/mol C2H3 C6H9 C4H6 C8H12 10 . Aluminum and bromine react vigorously according to the following equation: 2Al(s) + 3Br2(l) 2AlBr3(s) What mass of product can be made by reacting 5.0 g of aluminum and 25 g of bromine? 11 g 28 g 62 g 49 g 42 g 11 . The alcohol in "gasohol" burns according to the following equation: C2H5OH + 3O2 2CO2 + 3H2O How many grams of CO2 are produced when 3.00 g of C2H5OH are burned in this way? 53 Unit 3/4 – Moles / Stoichiometry 5.74 g 88.0 g 6.00 g 2.87 g 0.130 g 12 . How many moles of H2O are produced when 2.5 mol of O2 react according to the following equation? C3H8 + 5O2 3CO2 + 4H2O 3.0 2.5 4.0 2.0 Problem Solving Center Homework 4 1. You are setting up a reaction between two chemicals that react according to the equation 3 A + 4 B ---> products. If you start with 1.00 mole each of both A and B, which chemical will be in excess at the end, and by how much (assuming the reaction goes to completion)? Neither A nor B is in excess, because the reaction "goes to completion." B is in excess by 0.250 mol. B is in excess by 0.333 mol. A is in excess by 0.333 mol. A is in excess by 0.250 mol. 2. What is the molecular formula of the following compound? empirical formula CH, molar mass 78 g/mol C4H4 C3H3 C6H6 CH 54 Unit 3/4 – Moles / Stoichiometry C2H2 3. Give the empirical formula of the following compound if a sample contains 57.8 percent C, 3.6 percent H, and 38.6 percent O by mass. C4H3O2 C12H9O6 C2HO C8H6O4 4. Based on the following structural formula, calculate the percentage of carbon present. (CH2CO)2C6H3(COOH) 67.37 percent 66.67 percent 76.73 percent 64.70 percent 5. Calculate the mass percent of nitrogen in HNO3. 45.2 percent 22.2 percent 20.0 percent 25.0 percent none of these 6. Which of the following cannot be an empirical formula? C3H6 NO2 CH H2N CO2 7. A manufacturer of bicycles has 5350 wheels, 3023 frames, and 2655 handlebars. How many bicycles can be manufactured using these parts? 55 Unit 3/4 – Moles / Stoichiometry 2675 bicycles 5350 bicycles 3023 bicycles 2655 bicycles 8. CO2 exhaled by astronauts is removed from the spaceship atmosphere by reaction with KOH: CO2 + 2KOH K2CO3 + H2O How many kg of CO2 can be removed with 1.00 kg of KOH? 0.392 kg 0.784 kg 1.57 kg 0.500 kg 9. Aluminum and oxygen react according to the following equation: 4Al(s) + 3O2(g) 2Al2O3(s) In a certain experiment 4.6 g Al was reacted with excess oxygen and 6.8 g of product was obtained. What was the percent yield of the reaction? 63 percent 78 percent 74 percent 134 percent 68 percent 10 . For the reaction 3NO2 + H2O 2HNO3 + NO, how many grams of HNO3 can form when 1.00 g of NO2 and 2.25 g of H2O are allowed to react? 1.37 g 0.913 g 0.667 g 56 Unit 3/4 – Moles / Stoichiometry 15.7 g 11 . For the reaction Fe(CO)5 + 2PF3 + H2 Fe(CO)2(PF3)2(H)2 + 3CO, how many moles of CO are produced from a mixture of 5.0 mol Fe(CO)5, 8.0 mol PF3, and 6.0 mol H2? 18 mol 5.0 mol 6.0 mol 12 mol 15 mol 12 . If 4.0 moles of Li and 2.0 moles of O2 are used in the reaction 4Li + O2 ---> 2Li2O, then the limiting reactant is _________ and the theoretical yield of Li2O is ____________ g. oxygen, 6.0 x 10 oxygen, 1.2 x 102 lithium, 3.0 x 10 lithium, 6.0 x 10 Problem Solving Center Quiz 1 1. Write the balanced equation for the reaction that occurs when solid potassium nitrate is heated and decomposes to form solid potassium nitrite and oxygen gas. 2KNO3 2. 2KNO2 + O2 2KNO4(s) 2KNO3(s) + O2(g) 2KNO3(s) 2KNO2(s) + O2(g) KNO3(s) KNO2(s) + (1/2)O2(g) What is the formula weight of (NH4)2SO4? 118 amu 100 amu 116 amu 132 amu 3. Balance the following equation and indicate whether it is a combustion, combination, or decomposition reaction: 57 Unit 3/4 – Moles / Stoichiometry Li + N2 Li3N Li + N2 3Li3N, decomposition reaction 6Li + N2 2Li3N, decomposition reaction 6Li + N2 2Li3N, combination reaction Li + N2 6Li + N2 4. 3Li3N, combination reaction 2Li3N, combustion reaction Suppose you are setting up a reaction that requires an iodide salt and are planning to use sodium iodide. However, at the last minute you find that you are out of sodium iodide, so you must use potassium iodide instead. Will you need to weigh out more, less, or the same mass of potassium iodide in order to get the same number of moles of iodide ions? less same more 5. Convert the following to a balanced chemical reaction: Gaseous hydrogen reacts with carbon monoxide to form methanol, CH3OH. 4H + CO CH3OH H2 + CO CH3OH 2H2 + CO2 CH3OH 58 Unit 3/4 – Moles / Stoichiometry 2H2 + CO 6. Balance the following equation: Mg3N2 + H2O 7. CH3OH Mg(OH)2 + NH3 Mg3N2 + 6H2O 3Mg(OH)2 + NH3 Mg3N2 + 6H2O 3Mg(OH)2 + 2NH3 Mg3N2 + 2H2O Mg(OH)2 + NH3 Mg3N2 + 3H2O 3Mg(OH)2 + 2NH3 Calculate the number of molecules in a tablespoon of table sugar, C12H22O11, weighing 10.5 g. 2.22 x 1023 3.01 x 1023 1.85 x 1022 6.32 x 1024 6.02 x 1023 8. Potassium sulfate contains 44.9 percent potassium by mass. In a 50.0-g sample of potassium sulfate, the number of moles of potassium is 1.28 mol. 2.00 mol. 1.74 mol. 0.287 mol. 0.574 mol. 9. The reaction C7H8 + 3HNO3 ---> C7H5N3O6 + 3H2O can be used to make TNT. How many grams of HNO3 are required to react with 10.0 g of C7H8? 20.5 g 6.81 g 30.0 g 2.28 g 10.1 g 10 . Calculate the number of moles of water present in a 10.0-kg sample. 1.80 x 102 mol 55.5 mol 555 mol 1.80 x 105 mol none of these 11 . How many F- ions are present in 2.50 mol of BaF2? 5.00 3.01 x 1024 59 Unit 3/4 – Moles / Stoichiometry 1.51 x 1024 2.50 8.31 x 10-24 12 . What mass of silver chloride can be made from the reaction of 4.22 g of silver nitrate with 7.73 g of aluminum chloride? (Be sure to balance the reaction.) AgNO3 + AlCl3 Al(NO3)3 + AgCl 24.9 g 3.56 g 10.7 g 12.7 g Practice Questions Write the balanced equations for the following reactions. 1. C2H6 + O2→ CO2 + H2O 2. Na + H2O → NaOH + H2 3. Ammonium nitrate decomposes to yield dinitrogen monoxide and water. 4. Ammonia reacts with oxygen gas to form nitrogen monoxide and water. 5. Iron (III) oxide reacts with carbon (C) to yield iron metal and carbon monoxide. 6. Hydrogen gas reacts with carbon monoxide to yield methanol (CH3OH). How many grams of methanol are formed when 15.6 g of hydrogen react with excess carbon monoxide? 7. How many moles of carbon dioxide are formed in the fermentation of 75 g of glucose? 8. The thermite reaction (Fe2O3 + Al → Fe + Al2O3) can be used to ignite solid-fuel rockets or bombs. How much aluminum is needed to react with 10.0 g of Fe2O3? 9. Identify the limiting reactant and how much ammonia gas can be produced when 7.2 g of nitrogen gas react with 1.5 g of hydrogen gas by the use of the Haber process: 3H2 + N2→ 2NH3. 10. Identify the limiting reactant and how much carbon dioxide gas can be produced when 15.2 g of methane react with 18.5 g of oxygen gas to produce water and carbon dioxide. 11. Identify the limiting reactant and how much nitric acid can be produced when 60.0 g of nitrogen dioxide react with 18.5 g of water to produce nitric acid and nitrogen monoxide. 60 Unit 3/4 – Moles / Stoichiometry 12. Identify the limiting reactant and how much aspirin (C9H8O4) can be produced when 52.3 g of salicylic acid (C8H6O3) react with 25.0 g of acetic acid (CH3CO2H): C8H6O3 + CH3CO2H → C9H8O4 + H2O. Calculate the percent yield for practice problems 9 through 12 if 13. 6.3 g of ammonia were produced from problem 9. 14. 12.4 g of carbon dioxide were produced from problem 10. 15. 51 g of nitric acid were produced from problem 11. 16. 31.0 g of aspirin were produced from problem 12. Answers 1. 2C2H6 + 7O2 → 4CO2 + 6H2O 2. 2Na + 2H2O → 2NaOH + H2 3. NH4NO3 → N2O + 2H2O 4. 4NH3 + 5O2 → 4NO + 6H2O 5. Fe2O3 + 3C → 2Fe + 3CO 6. 124 g 7. 37 g 8. 3.38 g 9. Limiting reactant = H2; theoretical yield = 8.4 g 10. Limiting reactant = O2; theoretical yield = 12.7 g 11. Limiting reactant = NO2; theoretical yield = 55 g 12. Limiting reactant = salicylic acid; theoretical yield = 62.8 g 13. 75% 14. 97.6% 15. 93% 61 Unit 3/4 – Moles / Stoichiometry 16. 56% - Match the Following 62 Unit 3/4 – Moles / Stoichiometry 63 Unit 3/4 – Moles / Stoichiometry 64 Unit 3/4 – Moles / Stoichiometry 65 Unit 3/4 – Moles / Stoichiometry 66 Unit 3/4 – Moles / Stoichiometry Chemistry: Chapter 9 Stoichiometry and Baking Soda (NaHCO3) as the Limiting Reactant Activity Purposes: 1. Calculate theoretical mass of NaCl based on a known mass of NaHCO 3. 2. Experimentally determine the actual mass of NaCl produced. 3. Calculate the percent yield for your experiment. Reaction Equation: NaHCO3(s) + HCl(aq) NaCl(s) + CO2(g) + H2O(l) Materials: safety glasses HCl and dropper evaporating dish and matches watch glass baking soda (NaHCO3) concentrated ring stand with ring bunsen burner wire gauze tongs Procedure: 1. Find the mass of the evaporating dish and watch glass. Record this mass in the Data Table. 2. Add 1/3 of a teaspoon of baking soda to the evaporating dish, and record the total mass in the Data Table. 3. Cover the evaporating dish with the watch glass so that only the spout of the evaporating dish is exposed. 4. Use the dropper to drip HCl down the spout and into the dish. Add HCl until the fizzing ceases. 5. Leaving the watch glass in place, boil off the liquid until only table salt (NaCl) remains in the dish. 6. Let the dish cool for five minutes, then weigh it again and record the mass in the Data Table. 7. Clean up by rinsing your equipment with water and wiping dry with a paper towel. Data Table: (please include units) Quantity Measured Mass evaporating dish, watch glass evaporating dish, watch glass, NaHCO3 evaporating dish, watch glass, NaCl 67 Unit 3/4 – Moles / Stoichiometry Calculations: 1. Find the theoretical mass of NaCl that would be produced if your experiment were perfect. 2. Find the actual mass of NaCl that you obtained. 3. Find the percent yield for your experiment. If your percent yield is greater than 100%, provide at least one possible source of error that might have caused you to get more than 100% yield. Chemistry: Ch.9 Complete Stoichiometry Review 1. How many moles of O2 should be supplied to burn 1 mol of C3H8 (propane) molecules in a camping stove? 2. How many moles of O2 molecules should be supplied to burn 1 mol of CH4 molecules in a domestic furnace? 3. Sodium thiosulfate (Na2S2O3), photographer’s “hypo” reacts with unexposed silver bromide in the film emulsion to form sodium bromide and a compound of formula Na5[Ag(S2O3) 3]. How many moles of Na2S2O3 formula units are needed to make 0.10 mol of AgBr soluble? 4. Calculate the mass of alumina (Al2O3) produced when 100 g of aluminum burns in oxygen. 5. “Slaked lime,” Ca(OH) 2, is formed from “quick-lime” (CaO) by adding water. What mass of water is needed to convert 10 kg of quicklime to slaked lime? What mass of slaked lime is produced? 6. Camels store the fat tristearin (C57H110O6) in the hump. As well as being a source of energy, the fat is a source of water, because when it is used the reaction 2 C57H110O6(s) + 163 O2(g) 114 CO2(g) + 110 H2O(l) takes place. What mass of water is available from 1.0 kg of fat? 7. The compound diborane (B2H6) was at one time considered for use as a rocket fuel. How many grams of liquid oxygen would a rocket have to carry to burn 10 kg of diborane completely? (The products of the combustion are B2O3 and H2O.) 8. Given the balanced chemical equation Br2 + 2 NaI 2 NaBr + I2 How many moles of sodium bromide (NaBr) could be produced from 0.172 mol of bromine (Br 2)? 68 Unit 3/4 – Moles / Stoichiometry 9. How many formula units of calcium oxide (CaO) can be produced from 4.9 x 105 molecules of oxygen gas (O2) that react with calcium (Ca) according to this balanced chemical equation? 2 Ca(s) + O2 (g) 2 CaO(s) 10. Aluminum metal (Al) reacts with sulfur (S) to produce aluminum sulfide (Al2S3) according to this balanced chemical equation: 2 Al(s) + 3 S(s) Al2S3(s) How many atoms of aluminum will react completely with 1.33 x 10 24 atoms of sulfur? Name _____________________________ Hr ___ LIMITING REAGENTS 11. What is the maximum mass of methane (CH 4) that can be burned if only 1.0 g of oxygen is available? 12. What is the maximum mass of glucose (C6H12O6) that can be burned in 10 g of oxygen? 13. The solid fuel in the booster stage of the space shuttle is a mixture of ammonium perchlorate and aluminum powder, which react as follows: 6 NH4ClO4(s) + 10 Al(s) 5 Al2O3(s) + 3 N2(g) + 6 HCl(g) + 9 H2O(g) What mass of aluminum should be mixed with 5.0 x 103 kg of ammonium perchlorate, if the reaction proceeds as stated? 14. A solution containing 5.0 g of silver nitrate was mixed with another containing 5.0 g of potassium chloride. Which was the limiting reagent for the precipitation of silver chloride? 15. Given the balanced chemical equation 2 Ag + I2 2 AgI How many atoms of silver metal (Ag) are required to react completely with 531.8 g of iodine (I 2) to produce silver iodide (AgI)? 16. The theoretical yield of ammonia in an industrial synthesis was 550 tons, but only 480 tons was obtained. What was the percentage yield of the reaction? 17. Calculate the volume occupied by 16.3 moles of nitrogen gas (N2) at STP. 18. How many moles of fluorine gas (F2) are 69 Unit 3/4 – Moles / Stoichiometry contained in 0.269 dm3 container at STP? 19. Assuming that the gases are all at STP, find the volume of nitrogen dioxide gas (NO2) that could be produced from 71.11 dm 3 of nitrogen gas (N2) according to this balanced chemical equation. N2(g) + 2 O2(g) 2 NO2(g) 20. How many moles of oxygen (O2) would be needed to produce 79.60 moles of sulfur trioxide (SO3) according to the following balanced chemical equation? 2 SO2 + O2 2 SO3 21. How many grams of water will be produced from 50 g hydrogen reacting with 50 g oxygen? Think Critically 22. The reaction of 1 mol of C to form carbon monoxide in the reaction 2 C(s) + O 2(g) 2 CO(g) releases 113 kJ of heat. How much heat will be released by the combustion of 100 g of C according the the above information? 23. According to the balanced chemical equation; how many atoms of silver will be produced from combining 100 g of copper with 200 g of silver nitrate? Cu(s) + 2 AgNO3(aq) Cu(NO3) 2(aq) + 2 Ag(s) 24. According to the balanced chemical equation; how many moles of SO2(g) will be produced when 1.5 x 108 molecules of zinc sulfide react with 1000 dm3 of oxygen gas? Assume a 75% yield. 2 ZnS(s) + 3 O2 (g) 2 ZnO(s) + 2 SO2(g) 25. I need to produce 500 g of lithium oxide(Li2O) a) how many grams of Lithium AND b) how many liters of oxygen do I need The balanced equation is: Li + O2 LiO2 26. How many grams of water will be produce from 50 g hydrogen reacting with 50 g oxygen? Think Critically: 22. The reaction of 1 mol of C to form carbon monoxide in the reaction 2 C(s) + O2(g) 2 CO(g) releases 113 kJ of heat. How much heat will be released by the combustion of 100 g of C according the the above information? 23. According to the balanced chemical equation; how many atoms of silver will be produced from combining 100 g of copper with 200 g of silver nitrate? Cu(s) + 2 AgNO3(aq) Cu(NO3) 2(aq) + 2 Ag(s) 70 Unit 3/4 – Moles / Stoichiometry 24. According to the balanced chemical equation; how many moles of SO2(g) will be produced when 1.5 x 108 molecules of zinc sulfide react with 1000 dm3 of oxygen gas? Assume a 75% yield. 2 ZnS(s) + 3 O2 (g) 2 ZnO(s) + 2 SO2(g) 25. I need to produce 500 g of lithium oxide(Li2O) a. how many grams of Lithium AND b. how many liters of oxygen do I need The balanced equation is: Li + O2 LiO2 26. How many grams of water will be produced from 50 g hydrogen reacting with 50 g oxygen? Think Critically: 22. The reaction of 1 mol of C to form carbon monoxide in the reaction 2 C(s) + O 2(g) 2 CO(g) releases 113 kJ of heat. How much heat will be released by the combustion of 100 g of C according the the above information? 23. According to the balanced chemical equation; how many atoms of silver will be produced from combining 100 g of copper with 200 g of silver nitrate? Cu(s) + 2 AgNO3(aq) Cu(NO3) 2(aq) + 2 Ag(s) 24. According to the balanced chemical equation; how many moles of SO2(g) will be produced when 1.5 x 108 molecules of zinc sulfide react with 1000 dm3 of oxygen gas? Assume a 75% yield. 2 ZnS(s) + 3 O2 (g) 2 ZnO(s) + 2 SO2(g) 25. I need to produce 500 g of lithium oxide(Li2O) a. how many grams of Lithium AND b. how many liters of oxygen do I need The balanced equation is: Li + O2 LiO2 26. How many grams of water will be produced from 50 g hydrogen reacting with 50 g oxygen? Think Critically: 22. The reaction of 1 mol of C to form carbon monoxide in the reaction 2 C(s) + O 2(g) 2 CO(g) releases 113 kJ of heat. How much heat will be released by the combustion of 100 g of C according the the above information? 23. According to the balanced chemical equation; how many atoms of silver will be produced from combining 100 g of copper with 200 g of silver nitrate? Cu(s) + 2 AgNO3(aq) Cu(NO3) 2(aq) + 2 Ag(s) 71 Unit 3/4 – Moles / Stoichiometry 24. According to the balanced chemical equation; how many moles of SO2(g) will be produced when 1.5 x 108 molecules of zinc sulfide react with 1000 dm3 of oxygen gas? Assume a 75% yield. 2 ZnS(s) + 3 O2 (g) 2 ZnO(s) + 2 SO2(g) 25. I need to produce 500 g of lithium oxide(Li2O) a. how many grams of Lithium AND b. how many liters of oxygen do I need The balanced equation is: Li + O2 LiO2 26. How many grams of water will be produced from 50 g hydrogen reacting with 50 g oxygen? Think Critically 22. The reaction of 1 mol of C to form carbon monoxide in the reaction 2 C(s) + O 2(g) 2 CO(g) releases 113 kJ of heat. How much heat will be released by the combustion of 100 g of C according the the above information? 23. According to the balanced chemical equation; how many atoms of silver will be produced from combining 100 g of copper with 200 g of silver nitrate? Cu(s) + 2 AgNO3(aq) Cu(NO3) 2(aq) + 2 Ag(s) 24. According to the balanced chemical equation; how many moles of SO2(g) will be produced when 1.5 x 108 molecules of zinc sulfide react with 1000 dm3 of oxygen gas? Assume a 75% yield. 2 ZnS(s) + 3 O2 (g) 2 ZnO(s) + 2 SO2(g) 25. I need to produce 500 g of lithium oxide(Li2O) a) how many grams of Lithium AND b) how many liters of oxygen do I need The balanced equation is: Li + O2 LiO2 26. A tin ore contains 3.5% SnO2. How much tin is produced by reducing 2.0 kg of the ore with carbon? SnO2 + C Sn + CO2 27. If 36.5 g of HCl and 73 g of Zn are put together: 2 HCl + Zn ZnCl2 + H2 a. b. c. d. Determine which reactant is the limiting reactant, Find the mass of ZnCl2 formed, Find the volume of H2 (@ STP) formed, Determine which reactant is in excess and by how much. 28. Many plants synthesize glucose by photosynthesis as follows: CO2(g) + H2O(l) + energy C6H12O6(s) + O2(g) a. Write a balanced equation for this process, 72 Unit 3/4 – Moles / Stoichiometry b. How many molecules of water are needed to make one molecule of glucose? c. How many liters of oxygen (@STP) are given off when 2.50 mol of glucose is synthesized? d. How many moles of CO2 are needed for a plant to make 2.50 mole of glucose? e. How many carbon atoms are used to produce 2.50 mole of glucose? f. How many dm3 of oxygen gas are produced from 9.32 dm 3 of CO2 (all @ STP)? 29. Assume that the human body requires daily energy that comes from metabolizing 816 g of sucrose, C12H22O11, using the following reaction: C12H22O11(s) + 12 O2(g) 12 CO2(g) + 11 H2O(l) + energy How many dm3 of pure oxygen (@ STP) is consumed by a human being in 24 hours? 30. A student has a mixture of KClO3, K2CO3, and KCl. She heats 50 g of the mixture and determines that 5 g O2 and 7 g CO2 are produced by these reactions: 2 KClO3(s) 2 KCl(s) + 3 O2(g) K2CO3(s) K2O(s) + CO2(g) KCl is not affected by the heat. What is the percent composition of the original mixture? ANSWERS: 1. 5 mol O2 2. 2 mol O2 3. 0.3 mol Na2S2O3 4. 189 g Al2O3 5. 3214 g H2O and 13.2 kg slaked lime [Ca(OH) 2] 6. 998 g water 7. 34,783 g O2 8. 0.344 mol NaBr 9. 9.8 x 105 molecules CaO 10. 8.9 x 1023 atoms Al 11. 0.25 g CH4 12. 9.375 g C6H12O6 13. 1915 kg Al 14. silver nitrate 15. 2.5 x 1024 atoms Ag 16. 87.3 % yield 17. 365 L N2 18. 0.012 mol F2 19. 142 L NO2 20. 39.8 mol O2 21. 56.25 g H2O 22. 942,000 J 23. 7.1 x 1023 atoms Ag 24. 1.9 x 10-16mol (NOT 2.5 x 10-16mol: 75% Yield) 25a. g Li b. L O2 26. 27a. b. c. d. 28a. b. 6 c. 336 d. 15 e. 9 x 1024 f. 9.32 29. 641 L O2 30. 15.2 g KCl 73 Unit 3/4 – Moles / Stoichiometry Name: ________________________ Hour: ____ Date: ___________ Chemistry: Percent Yield Directions: Solve each of the following problems. Show your work, including proper units, to earn full credit. 1. “Slaked lime,” Ca(OH)2, is produced when water reacts with “quick lime,” CaO. If you start with 2 400 g of quick lime, add excess water, and produce 2 060 g of slaked lime, what is the percent yield of the reaction? 2. Some underwater welding is done via the thermite reaction, in which rust (Fe 2O3) reacts with aluminum to produce iron and aluminum oxide (Al2O3). In one such reaction, 258 g of aluminum and excess rust produced 464 g of iron. What was the percent yield of the reaction? 3. Use the balanced equation to find out how many liters of sulfur dioxide are actually produced at STP if 1.5 x 1027 molecules of zinc sulfide are reacted with excess oxygen and the percent yield is 75%. 2 ZnS(s) + 3 O2(g) 2 ZnO(s) + 2 SO2(g) 4. The Haber process is the conversion of nitrogen and hydrogen at high pressure into ammonia, as follows: N2(g) + 3 H2(g) 2 NH3(g) 74 Unit 3/4 – Moles / Stoichiometry If you must produce 700 g of ammonia, what mass of nitrogen should you use in the reaction, assuming that the percent yield of this reaction is 70%? Answers: 1. 65% 4. 824 g N2 2. 87% 3. 4.19 x 104 L SO2 75 Unit 3/4 – Moles / Stoichiometry 76 Unit 3/4 – Moles / Stoichiometry 77 Unit 3/4 – Moles / Stoichiometry 78 Unit 3/4 – Moles / Stoichiometry 79 Unit 3/4 – Moles / Stoichiometry 80 Unit 3/4 – Moles / Stoichiometry 81 Unit 3/4 – Moles / Stoichiometry 82 Unit 3/4 – Moles / Stoichiometry 83 Unit 3/4 – Moles / Stoichiometry 84 Unit 3/4 – Moles / Stoichiometry 85 Unit 3/4 – Moles / Stoichiometry 86 Unit 3/4 – Moles / Stoichiometry CHEMICAL REACTIONS OF COPPER AND PERCENT YIELD Objective To gain familiarity with basic laboratory procedures, some chemistry of a typical transition element, and the concept of percent yield. Apparatus and Chemicals 0.5 g piece of no. 16 or no. 18 copper wire evaporating dish 250 mL beaker (2) weighing paper concentrated HNO3 (4 – 6 mL) 6.0 M H2SO4 (15 mL) graduated cylinder granular zinc 3.0 M NaOH (30 mL) methanol carborundum boiling chips acetone stirring rod towel iron ring and ring stand balance wire gauze aluminum foil cut in 1-inch squares Bunsen burner concentrated HCl (drops) Discussion Most chemical synthesis involves separation and purification of the desired product from unwanted side products. Some methods of separation, such as filtration, sedimentation, decantation, extraction, and sublimation were discussed earlier. This experiment is designed as a quantitative evaluation of your individual laboratory skills in carrying out some of these operations. At the same time you will become more acquainted with two fundamental types of chemical reactions -- redox reactions and metathesis (double-displacement) reactions. By means of these reactions, you will finally recover the copper sample with maximum efficiency. The chemical reactions involved are the following. Cu(s) + 4 HNO3(aq) -----> Redox [1] Cu(NO3)2(aq) + 2 NaOH(aq) Metathesis [2] Cu(OH)2(s) -----> Dehydration [3] CuO(s) + Metathesis CuO(s) H2SO4(aq) [4] Cu(NO3)2(aq) + 2 NO2(g) + 2 H2O(l) -----> + -----> Cu(OH)2(s) + 2 NaNO3(aq) H2O(g) CuSO4(aq) + H2O(l) 87 Unit 3/4 – Moles / Stoichiometry CuSO4(aq) Redox + Zn(s) [5] -----> ZnSO4(aq) + Cu(s) Each of these reactions proceeds to completion. Metathesis reactions proceed to completion whenever one of the components is removed from the solution, such as in the formation of a gas or an insoluble precipitate (driving forces). This is the case for reaction [1], [2], and [3], where in reactions [1] and [3] a gas and in reaction [2] an insoluble precipitate are formed. Reaction [5] proceeds to completion because zinc has a lower ionization energy or oxidation potential that copper. The objective in this experiment is to recover all of the copper you begin with, in analytically pure form. This is the test of your laboratory skills. The percent yield of the copper can be expressed as the ratio of the recovered weight to initial weight, multiplied by 100: recovered wt of Cu % yield = x 100 initial wt of Co Procedure Weight approximately 0.500 g of no. 16 or no. 18 copper wire (1) to the nearest 0.0001 g and place it in a 250 mL beaker. Add 4-5 mL of concentrated HNO3 to the beaker, IN THE HOOD. After the reaction is complete, add 100 mL distilled H2O. Describe the reaction (6) as to color change, evolution of gas, and change in temperature (exothermic or endothermic) in the report sheet. Add 30 mL of 3.0 M NaOH to the solution in your beaker and describe the reaction (7). Add two or three boiling chips and carefully heat the solution -- while stirring with a glass stirring rod -- just to the boiling point. Describe the reaction on your report sheet (8). Remove the boiling chips. Allow the black CuO to settle; then decant the supernantant liquid. Add about 200 mL of very hot distilled water and allow the CuO to settle. Decant once more. What are you removing by washing and decanting (9)? Add 15 mL of 6.0 M H2SO4. What copper compound is present in the beaker now (10)? Your teacher will tell you whether you should use Zn or Al for the reduction of Cu (II) in the following step. A. Zinc In the hood, add 2.0 g of 30-mesh zinc metal all at once and stir until the supernatant liquid is colorless. Describe the reaction on your report sheet (11). What is present in solution (12)? When gas evolution has become very slow, heat the solution gently (but do not boil) and allow it to cool. What gas is formed in this reaction (13)? How do you know (14)? B. Aluminum 88 Unit 3/4 – Moles / Stoichiometry In the hood, add several 1-inch squares of aluminum foil and a few drops of concentrated HCl. Continue to add pieces of aluminum until the supernatant liquid is colorless. Describe the reaction on your report sheet (11). What is present in solution (12)? What gas is formed in this reaction (13)? How do you know (14)? When gas evolution has ceased, decant the solution and transfer the precipitate to a preweighed porcelain evaporating dish (3). Wash the precipitated copper with about 5 mL of distilled water, allow it to settle, decant the solution, and repeat the process. What are you removing by washing (15)? Wash the precipitate with about 5 mL of methanol (KEEP THE METHANOL AWAY FROM FLAMES _ IT IS FLAMMABLE!) Allow the precipitate to settle, and decant the methanol. (METHANOL IS ALSO EXTREMELY TOXIC: AVOID BREATHING THE VAPORS AS MUCH AS POSSIBLE.) Finally, wash the precipitate with about 5 mL of acetone (KEEP THE ACETONE AWAY FROM FLAMES - IT IS EXTREMELY FLAMMABLE!), allow the precipitate to settle, and decant the acetone from the precipitate. Prepare a steam bath as illustrated and dry the product on your steam bath for at least 5 minutes. Wipe the bottom of the evaporating dish with a towel, remove the boiling chips and weigh the evaporating dish plus copper (2). Calculate the final a\weight of copper (4). Compare the weight with your initial weight and calculate the percent yield (5). What color is your copper sample (16)? Is it uniform in appearance (17)? Suggest possible sources of error in this experiment (18). Chemical Reactions of Copper and Percent Yield Pre-lab (Review Questions) 1. Give an example, other than the ones listed in this experiment, of redox and metathesis reactions. 2. When will reactions proceed to completion? 3. Define percent yield in general terms. 4. Name six methods of separating materials. 89 Unit 3/4 – Moles / Stoichiometry 5. Give criteria in terms of temperature changes for exothermic and endothermic reactions. 6. If 1.65 g of Cu(NO3)2 are obtained from allowing 0.93 g of Cu to react with excess HNO 3, what is the percent yield of the reaction? Prelab – Chemical Reactions of Copper and Percent Yield 7. What is the maximum percent yield in any reaction? 8. What is meant by the terms decantation and filtration? 9. When Cu(OH)2(s) is heated, Copper (II) oxide and water are formed. Write a balanced equation for the reaction. 90 Unit 3/4 – Moles / Stoichiometry 10. When sulfuric acid and copper (II) oxide are allowed to react, copper (II) sulfate and water are formed. Write a balanced equation for this reaction. 11. When copper (II) sulfate and aluminum are allowed to react, aluminum sulfate and copper are formed. What kind of reaction is this? Write a balanced equation for this reaction. REPORT SHEET Chemical Reactions of Copper and Percent Yield 1. Weight copper initial _______________ 2. Weight of copper and evaporating dish _______________ 3. Weight of evaporating dish _______________ 4. Weight of copper final _______________ 5. % Yield (show calculations) _______________ 6. Describe the reaction Cu(s) + HNO3(aq) --> 7. Describe the reaction Cu(NO3)2(aq) + NaOH(aq) --> 8. Describe the reaction Cu(OH)2(s) --> 9. What are you removing by this washing (be specific)? 91 Unit 3/4 – Moles / Stoichiometry 10. What copper compound is present in the beaker? 11. Describe the reaction CuSO4(aq) + Zn(s), or CuSO4(aq) + Al(s) 12. What is present in solution (aqueous)? REPORT SHEET – page 2 Chemical Reactions of Copper and Percent Yield 13. What is the gas? Hint: Where did the HCl(aq) go? Why did you have to do this reaction in the fume hood? 14. How do you know? 15. What are you removing by washing? 16. What color is your copper sample? 17. Is it uniform in appearance? 18. Suggest possible sources of error in this experiment. POST LAB QUESTIONS 1. If your percent yield of copper was greater than 100%, what are two plausible errors you may have made? 92 Unit 3/4 – Moles / Stoichiometry 2. Consider the combustion of methane, CH4: CH4(g) + 2 O2(g) -----> CO2(g) + 2 H2O(g) Suppose 2 mole of methane is allowed to react with 3 mol of oxygen. a) What is the limiting reagent? (show work) b) How many moles of CO2 can be made from this mixture? How many grams of CO2? 3. Suppose 8.00 g of CH4 is allowed to burn in the presence of 6.00 g of oxygen. How much (in grams) CH4, O2, CO2, and H2O remain after the reaction is complete? 4. How many milliliters of 6.0 M H2SO4 are required to react with 0.80 g of CuO according to Equation [4]? 5. If 2.00 g of Zn is allowed to react with 1.75 g of CuSO 4 according to Equation [5], how many grams of Zn will remain after the reaction is complete? 93 Unit 3/4 – Moles / Stoichiometry 6. What is meant by the term limiting reagent? Explain 94 Unit 3/4 – Moles / Stoichiometry 95 Unit 3/4 – Moles / Stoichiometry 96 Unit 3/4 – Moles / Stoichiometry 97 Unit 3/4 – Moles / Stoichiometry 98 Unit 3/4 – Moles / Stoichiometry 99 Unit 3/4 – Moles / Stoichiometry Show your work for all calculations. 1. What volume of carbon dioxide can be produced by reacting 10.5 g baking soda (NaHCO 3) with excess hydrochloric acid at STP according to the reaction below? NaHCO3(s) + HCl(aq) → NaCl(aq) + CO2(g) + H2O(l) 2. Potassium chlorate decomposes to produce oxygen gas by the reaction shown below. What volume of oxygen gas can be produced by heating 5.89 g KClO 3 at a pressure of 0.50 atmospheres and a temperature of 79 °C? 2 KClO3(s) → 2KCl(s) + 3O2(g) 3. Cetyl alcohol (C16H34O) is a flammable solid at room temperature. If 15.8 g of the alcohol are combusted, what volume of water vapor will be produced at a pressure of 1.26 atmospheres and a temperature of 110 °C? (Watch out — this equation isn't balanced!) C16H34O(s) + O2(g) → CO2(g) + H2O(g) Conclusion Now, address the Focus Question: How much metal does it take to float a blimp? Given the information below, you should be able to calculate an answer. 100 Unit 3/4 – Moles / Stoichiometry Zinc reacts with sulfuric acid by the reaction shown below. If 759 kg zinc were to react with excess sulfuric acid, how large would a blimp be if it were filled with the hydrogen gas produced from this reaction at a pressure of 1.10 atm and a temperature of 22 °C? Zn(s) + H2SO4(aq) → ZnSO4(aq) + H2(g) Compound Table 2 Formula Mass (amu) MgO K2S C3H8O3 Ca(NO3)2 Mg3(PO4)2 C6H2CH3(NO2)3 Pt(NH3)2Cl2 Fe(ClO4)2 Na2CO3 • 10H2O MgSO4 • 7H2O Conclusion Now address the Focus Question: Which has more mass—salt (NaCl) or sugar (C12H22O11)? Show calculations to support your answer. A comparison of two domestic fuels - pollution This sheet compares the amount of pollution produced by two domestic fuels - anthracite, the purest form of coal, and propane, sold in cylinders as liquefied petroleum gas, LPG. Both anthracite and propane produce carbon dioxide when they burn. 1. What environmental problem does carbon dioxide contribute to? .......................................................................................................................................... 2. Give two practical consequences of this problem. .......................................................................................................................................... 101 Unit 3/4 – Moles / Stoichiometry .......................................................................................................................................... .......................................................................................................................................... .......................................................................................................................................... Anthracite contains between 89% and 98% pure carbon. When carbon burns, carbon dioxide is formed and heat energy is produced. C(s) + O2(g) ’ CO2(g) DH = -394 kJ mol-1 3. Calculate the mass of 1 mole of carbon dioxide. (RAMs: C = 12; O = 16) .......................................................................................................................................... 4. Calculate the mass of carbon dioxide formed per kJ of heat energy produced. .......................................................................................................................................... .......................................................................................................................................... Propane burns according to the equation: C3H8(g) + 5O2(g) ’ 3CO2(g) + 4H2O(l) DH = -2220 kJ mol-1 5. Calculate the mass of carbon dioxide formed per kJ of heat energy produced when propane burns. .......................................................................................................................................... .......................................................................................................................................... .......................................................................................................................................... 6. Which is the more environmentally friendly fuel in terms of carbon dioxide emission? .......................................................................................................................................... Sulphur dioxide production Sulphur dioxide is formed when sulphur and sulphur-containing compounds burn. 7. What environmental problem does sulphur dioxide contribute to? 102 Unit 3/4 – Moles / Stoichiometry .......................................................................................................................................... 8. Give two practical consequences of this problem. .......................................................................................................................................... .......................................................................................................................................... .......................................................................................................................................... Coal often contains sulphur compounds, and these are converted into sulphur dioxide when the coal burns. At the moment there is no way of removing sulphur dioxide from the gases produced by homes. Propane is freed of sulphur-containing compounds at the refinery. The sulphur compounds are converted into sulphur dioxide which can then be used to manufacture a useful chemical. 9. Which chemical? .......................................................................................................................................... 10. Which is the more environmentally friendly fuel in terms of sulphur dioxide emission? .......................................................................................................................................... 103 Unit 3/4 – Moles / Stoichiometry 1. Copper reacts with nitric acid according to the following reaction: 3 Cu (s) + 8 HNO3 (aq) 3 Cu(NO3)2 (aq) + 2 NO (g) + 4 H2O (l) If an pre 1982 copper penny contains 3.10 grams of copper, what volume of 8.00 M nitric acid is required to exactly consume it? What volume of nitrogen monoxide gas measured at STP would be produced? (Remember that 1 mole of any gas measured at STP = 22.4 dm3) 2. Hydrogen peroxide decomposes slowly to produce oxygen gas and water. Write a balanced equation for the decomposition of hydrogen peroxide. What mass of hydrogen peroxide (H2O2) must decompose to produce 0.77g of water? 3. Carbon monoxide is an air pollutant found in automobile exhausts. It slowly reacts with oxygen in the atmosphere to produce carbon dioxide. Write a balanced equation for this process. What mass of carbon dioxide would be produced from 10.00 grams of carbon monoxide? What mass of oxygen is required to convert the carbon monoxide to carbon dioxide? 4. Suppose that 50.0 cm3 of 2.00 M hydrochloric acid are added to 4.00 grams of zinc metal. What it the limiting reagent? How many grams of zinc are actually consumed? What is the concentration of the hydrochloric acid solution after the reaction has occurred? What mass of hydrogen will be produced? Zn (s) + 2 HCl (aq) ZnCl2 (aq) + H2 (g) 104 Unit 3/4 – Moles / Stoichiometry 5. Consider the reaction: 2 VO + 3 Fe2O3 6 FeO + V2O5 If 609.5 grams of vanadium (II) oxide, VO, and 832 grams of iron(III) oxide, Fe2O3, are put into a container and allowed to react according to the equation above, which substance(s) and how many grams of each would be present in the container after the reaction is complete? 6. Methyl alcohol (wood alcohol), CH3OH, is produced via the reaction CO(g) + 2 H2(g) CH3OH(l) A mixture of 1.20 g H2(g) and 7.45 g CO(g) are allowed to react. (a) Which reagent is the limiting reagent? (b) What is the yield of CH3OH? [Assume theoretical yield in g is what is wanted here.] (c) How much of the reagent present in excess is left over? (d) Suppose the actual yield is 7.52 g of CH3OH. What is the % yield? 7. A 0.32570 gram sample of NaClO3 is decomposed to form oxygen gas according to the following reaction: 2 NaClO3 (s) 2 NaCl (s) + 3 O2 (g) The gas is collected over water at a total pressure of 742 torr at 23.0 oC. What volume of oxygen will be collected? (The vapor pressure of water at 23.0 oC is 20.65 torr) 8. Hydrogen is produced when zinc reacts with sulfuric acid according to the following reaction Zn (s) + H2SO4 (aq) ZnSO4 (aq) + H2 (g) If 159 cm3 of wet hydrogen is collected over water at 23.0o C and a total pressure of 738 torr, how many grams of zinc are consumed? (The vapor pressure of water at 23.0 oC is 20.65 torr) 105 Unit 3/4 – Moles / Stoichiometry Stoichiometry Challenge 1. (pg 129, #118) Nitric acid is produced commercially by the Ostwald process, represented by the following equations: a. 4HN3(g) +5O2(g) 4NO(g) + 6H2O(g) b. 2NO(g) + O2(g) 2NO2(g) c. 3NO2(g) + H2O(l) 2HNO3(aq) + NO(g) What mass of NH3 must be used to produce 1.o x 106 kg HNO3 by the Ostwald process? Assume 100% yield in each reaction and assume that the NO produced in the third step is not recycled. 2. (pg 129, #120) The aspirin substitute, acetaminophen (C8H9O2N), is produced by the following three-step synthesis: a. C6H5O3N(s) + 3H2(g) + HCl(aq) C6H8ONCl(s) + 2H2O(l) b. C6H8ONCl(s) + NaOH(aq) C6H7ON(s) + H2O(l) + NaCl(aq) c. C6H7ON(s) + C4H6O3(l) C8H9O2N(s) + HC2H3O2(l) The first two reactions have percent yields of 87% and 98% by mass, respectively. The overall reaction yields 3 mol of acetaminophen product for every 4 mol CH5O3N reacted. a. What is the percent yield by mass for the overall process? b. What is the percent yield by mass of step III? - 106 -
