This chapter tell you about the reduction in the Oxidation reaction there he is revolutions their transfer of ions and also about the oxidizing agent in the reducing agent
Concept of oxidation and reduction, redox reactions, oxidation number, balancing redox reactions, loss and gain of electrons, Balancing redox reactions, Half reaction method, Types of redox reaction- direct and indirect method, Electrochemical cell, Classification of redox reactions.
1) The document discusses classical ideas of oxidation and reduction reactions by defining them as addition or removal of oxygen, hydrogen, or electronegative/electropositive elements.
2) It then moves to discussing redox reactions in terms of electron transfer, defining oxidation as loss of electrons and reduction as gain of electrons.
3) Rules for calculating oxidation numbers are provided, including that the sum of oxidation numbers in a compound or ion must equal the overall charge. Stock notation is also introduced for representing oxidation states.
4) Examples are given of identifying oxidizing and reducing agents, balancing redox reactions using the oxidation number method, and classifying reactions as redox based on changes in oxidation numbers.
1) Corrosion is the reaction of a metal with its environment that causes it to convert to a metal compound. This occurs as the metal loses electrons and forms cations that combine with anions.
2) Redox reactions involve the transfer of electrons from one substance to another, causing a change in oxidation states. Reduction occurs when an atom gains electrons and is reduced, while oxidation occurs when an atom loses electrons and is oxidized.
3) Ions are formed when atoms gain or lose electrons, becoming cations if positively charged or anions if negatively charged. Oxidation numbers indicate the charge of an atom in a compound.
The document discusses the concept of redox reactions and oxidation numbers. It defines oxidation numbers as the imaginary charge left on an atom when other atoms are removed from a compound. Rules are provided for assigning oxidation numbers, such as elements having an oxidation number of 0 when uncombined, monatomic ions taking the charge of the ion, and the sum of oxidation numbers in a molecule or ion equaling the overall charge. Examples are given to illustrate calculating oxidation numbers using these rules. The key points are that oxidation involves losing electrons while reduction involves gaining electrons.
This document discusses oxidation-reduction (redox) reactions and concepts including definitions of oxidation and reduction in terms of gaining or losing electrons, oxygen, and hydrogen. It provides examples of redox reactions and identifies the oxidizing agent and reducing agent in reactions. It also discusses oxidation numbers and how to balance redox equations using the oxidation number change method. Finally, it discusses redox titrations and the specific methods of iodimetry and iodometry which involve the use of iodine as the titrant or analyte.
This document discusses oxidation-reduction (redox) reactions and concepts related to them. It defines oxidation as the loss of electrons or an increase in oxidation number, and reduction as the gain of electrons or a decrease in oxidation number. Redox reactions involve both an oxidation and a reduction occurring simultaneously. Oxidizing agents undergo reduction, reducing electrons and undergoing oxidation. Methods for determining oxidation numbers and balancing redox equations are also outlined.
This document summarizes key concepts from Chapter 20 on oxidation-reduction (redox) reactions. It defines oxidation as losing electrons and reduction as gaining electrons. Redox reactions involve the transfer of electrons between reactants. The species donating electrons is the reducing agent and is oxidized, while the species accepting electrons is the oxidizing agent and is reduced. Oxidation numbers are assigned to track electron transfers and identify redox reactions. There are two methods for balancing redox equations: using oxidation number changes or splitting the reaction into oxidation and reduction half-reactions.
The document discusses oxidation-reduction (redox) reactions, where there is a transfer of electrons between reactants. It defines oxidation as the loss of electrons and reduction as the gain of electrons. An example redox reaction and its net ionic form are provided. The document explains how to determine the oxidation states of elements and identifies common oxidation states of nonmetals. It describes how to write half-reactions by separating a redox reaction into its oxidation and reduction components.
Concept of oxidation and reduction, redox reactions, oxidation number, balancing redox reactions, loss and gain of electrons, Balancing redox reactions, Half reaction method, Types of redox reaction- direct and indirect method, Electrochemical cell, Classification of redox reactions.
1) The document discusses classical ideas of oxidation and reduction reactions by defining them as addition or removal of oxygen, hydrogen, or electronegative/electropositive elements.
2) It then moves to discussing redox reactions in terms of electron transfer, defining oxidation as loss of electrons and reduction as gain of electrons.
3) Rules for calculating oxidation numbers are provided, including that the sum of oxidation numbers in a compound or ion must equal the overall charge. Stock notation is also introduced for representing oxidation states.
4) Examples are given of identifying oxidizing and reducing agents, balancing redox reactions using the oxidation number method, and classifying reactions as redox based on changes in oxidation numbers.
1) Corrosion is the reaction of a metal with its environment that causes it to convert to a metal compound. This occurs as the metal loses electrons and forms cations that combine with anions.
2) Redox reactions involve the transfer of electrons from one substance to another, causing a change in oxidation states. Reduction occurs when an atom gains electrons and is reduced, while oxidation occurs when an atom loses electrons and is oxidized.
3) Ions are formed when atoms gain or lose electrons, becoming cations if positively charged or anions if negatively charged. Oxidation numbers indicate the charge of an atom in a compound.
The document discusses the concept of redox reactions and oxidation numbers. It defines oxidation numbers as the imaginary charge left on an atom when other atoms are removed from a compound. Rules are provided for assigning oxidation numbers, such as elements having an oxidation number of 0 when uncombined, monatomic ions taking the charge of the ion, and the sum of oxidation numbers in a molecule or ion equaling the overall charge. Examples are given to illustrate calculating oxidation numbers using these rules. The key points are that oxidation involves losing electrons while reduction involves gaining electrons.
This document discusses oxidation-reduction (redox) reactions and concepts including definitions of oxidation and reduction in terms of gaining or losing electrons, oxygen, and hydrogen. It provides examples of redox reactions and identifies the oxidizing agent and reducing agent in reactions. It also discusses oxidation numbers and how to balance redox equations using the oxidation number change method. Finally, it discusses redox titrations and the specific methods of iodimetry and iodometry which involve the use of iodine as the titrant or analyte.
This document discusses oxidation-reduction (redox) reactions and concepts related to them. It defines oxidation as the loss of electrons or an increase in oxidation number, and reduction as the gain of electrons or a decrease in oxidation number. Redox reactions involve both an oxidation and a reduction occurring simultaneously. Oxidizing agents undergo reduction, reducing electrons and undergoing oxidation. Methods for determining oxidation numbers and balancing redox equations are also outlined.
This document summarizes key concepts from Chapter 20 on oxidation-reduction (redox) reactions. It defines oxidation as losing electrons and reduction as gaining electrons. Redox reactions involve the transfer of electrons between reactants. The species donating electrons is the reducing agent and is oxidized, while the species accepting electrons is the oxidizing agent and is reduced. Oxidation numbers are assigned to track electron transfers and identify redox reactions. There are two methods for balancing redox equations: using oxidation number changes or splitting the reaction into oxidation and reduction half-reactions.
The document discusses oxidation-reduction (redox) reactions, where there is a transfer of electrons between reactants. It defines oxidation as the loss of electrons and reduction as the gain of electrons. An example redox reaction and its net ionic form are provided. The document explains how to determine the oxidation states of elements and identifies common oxidation states of nonmetals. It describes how to write half-reactions by separating a redox reaction into its oxidation and reduction components.
This document discusses electrochemistry and redox reactions. It defines oxidation and reduction, and explains that they occur together in redox reactions. Redox reactions involve the transfer of electrons from one species to another, changing their oxidation states. Oxidizing agents accept electrons from reducing agents. Common oxidizing agents are hydrogen peroxide and chlorine, while common reducing agents are carbon and hydrogen. Redox reactions can be identified by changes in oxidation states and through color changes using indicators like potassium manganate or iodide.
The document discusses oxidation-reduction (redox) reactions. It defines oxidation as the loss of electrons and reduction as the gain of electrons. Redox reactions always involve both an oxidation and a reduction process. The substance undergoing oxidation is the reducing agent as it loses electrons, while the substance undergoing reduction is the oxidizing agent as it gains electrons. Assigning oxidation numbers to atoms allows identification of which substances are oxidized and reduced in a reaction.
Sulfur has an oxidation number of +4 in SO2 and Na2SO4. Carbon has an oxidation number of +4 in CO32-. Oxygen has an oxidation number of -2 in all cases. Sodium has an oxidation number of +1 in Na2SO4. Nitrogen has an oxidation number of -3 and hydrogen has an oxidation number of +1 in (NH4)2S.
This document discusses oxidation-reduction (redox) reactions. It defines oxidation as the gain of oxygen, loss of hydrogen, or loss of electrons, while reduction is defined as the loss of oxygen, gain of hydrogen, or gain of electrons. Redox reactions involve both oxidation and reduction occurring simultaneously. Oxidation numbers are assigned to elements to indicate degree of oxidation. Common oxidizing agents are listed that can oxidize Fe2+ to Fe3+, like potassium manganate, chlorine, and nitric acid. Reducing agents like sodium sulfite can reduce Fe3+ back to Fe2+. The document also covers IUPAC nomenclature of inorganic compounds and rules for assigning oxidation numbers.
1. Electrochemistry involves redox reactions where one element is oxidized and another is reduced. Oxidation is the loss of electrons and an increase in oxidation number, while reduction is the gain of electrons and a decrease in oxidation number.
2. Electrolysis is the passage of an electric current through an ionic substance to cause a non-spontaneous redox reaction. Oxidation occurs at the anode and reduction at the cathode.
3. Aluminum is extracted from bauxite via electrolysis. Bauxite is dissolved in molten cryolite to lower its melting point, then electrolysis separates aluminum ions at the cathode.
Redox reactions involve the transfer of electrons from one reactant to another, resulting in oxidation and reduction. Oxidation is the loss of electrons and reduction is the gain of electrons. Common redox reactions include photosynthesis, respiration, combustion, and production and use of fertilizers.
Includes a discussion of Voltaic and electrolytic cells, the Nernst equation and the relationship between electrochemical processes, chemical equilibrium and free energy.
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This document provides an overview of redox reactions and electrochemistry applications. It discusses oxidation-reduction concepts like oxidation states and the OIL RIG mnemonic. Examples of redox reactions and electrochemistry applications are given, including galvanic cells, corrosion, electrolysis, and batteries. Key concepts covered include cell potential, the Nernst equation, and how concentration affects cell potential. Diagrams illustrate galvanic cells and how they function.
This document provides an overview of redox (oxidation-reduction) chemistry. It defines oxidation as the loss of electrons and reduction as the gain of electrons. Oxidation and reduction always occur simultaneously in redox reactions. The document discusses identifying oxidizing and reducing agents, balancing redox reactions using the half-reaction method, standard reduction potentials, galvanic (voltaic) cells that produce electricity from spontaneous redox reactions, and electrolytic cells that use electricity to drive nonspontaneous reactions.
This document provides information on chemical equations and oxidation-reduction reactions. It defines key concepts like oxidation, reduction, oxidizing agents, reducing agents and oxidation numbers. Examples of different types of chemical reactions like combination, decomposition, displacement and combustion are outlined. Steps for writing and balancing chemical equations are described. Oxidation-reduction reactions are explained along with biological examples of electron transfer processes. Specific equations are given and identified as oxidation or reduction reactions.
Redox Reaction and Electrochemical Cell (Reaksi Redoks dan Sel Elektrokimia)DindaKamaliya
An electrochemical cell converts chemical energy into electrical energy through spontaneous redox reactions. It consists of two half-cells separated by a salt bridge. In the cathode half-cell, reduction occurs as oxidized species gain electrons. In the anode half-cell, oxidation occurs as reduced species lose electrons. Electrons flow through an external circuit from the anode to the cathode. The standard electrode potential of each half-reaction predicts the cell's voltage under standard conditions.
This document contains lecture notes on quantitative analysis in chemistry. It discusses gravimetric analysis, which determines the amount of a substance by converting it into a product that can be isolated and weighed. An example is given of determining the amount of lead in water by precipitating lead sulfate, filtering and weighing the precipitate. A practice problem demonstrates calculating the mass of lead from the mass of lead sulfate precipitate obtained.
This document provides an overview of redox reactions including:
- Redox reactions involve the transfer of electrons between chemical species, resulting in oxidation and reduction.
- Oxidizing agents gain electrons and are reduced, while reducing agents lose electrons and are oxidized.
- Latimer, Frost, and Pourbaix diagrams can be used to predict and understand redox reactions in aqueous solutions by showing the thermodynamic stability of different oxidation states.
- Key concepts like disproportionation, oxidizing/reducing abilities, and stable/unstable species can be determined from these types of diagrams.
This document provides an overview of oxidation-reduction (redox) reactions including:
1. Defining oxidation as the loss of electrons and reduction as the gain of electrons.
2. Explaining that the reducing agent is oxidized and the oxidizing agent is reduced in redox reactions.
3. Detailing that the first step in understanding redox reactions is determining the oxidation number of each element in a compound.
This document provides an overview of redox (reduction-oxidation) reactions, including definitions of key terms like oxidation, reduction, oxidizing agents, reducing agents, and disproportionation reactions. It discusses identifying oxidation and reduction based on changes in oxygen, hydrogen, or electron content. Methods for determining oxidation states and balancing redox reactions using the half-reaction method are also described. Real-world examples of redox processes like corrosion and the blue bottle experiment are mentioned.
This document provides a summary of key concepts in oxidation-reduction (redox) reactions:
1) Redox reactions involve the transfer of electrons between chemical species, either through the complete transfer of electrons to form ionic bonds, or partial transfer to form covalent bonds. Oxidation is the loss of electrons and reduction is the gain of electrons.
2) Redox pairs are couples of oxidized and reduced forms of elements that differ in their oxidation state. Common redox pairs include Fe3+/Fe2+, O2/H2O, and MnO4-/MnO2.
3) The standard electrode potential (E°) indicates the tendency of half-reactions to occur.
The document discusses oxidation and reduction reactions. It provides examples of oxidation as the loss of electrons and reduction as the gain of electrons. Corrosion is described as an oxidation reaction where metals react with substances like oxygen, water and acids. Electroplating is introduced as a process where a thin coating of a metal like copper or nickel is deposited onto a surface through an oxidation-reduction reaction.
This document discusses oxidation numbers, which indicate the degree of oxidation or reduction of atoms in chemical compounds. Oxidation numbers are assigned according to rules such as atoms in their elemental form having an oxidation number of 0, monoatomic ions having the same charge as their oxidation number, and the sum of oxidation numbers in a neutral molecule being 0. Oxidation numbers can be fractional and are indicated in Roman numerals after the symbol of a metal in its compound. An example of a fractional oxidation number is given for sulfur atoms in the tetrathionate ion.
The document discusses oxidation-reduction (redox) reactions. It defines oxidation as loss of electrons and reduction as gain of electrons. Redox reactions involve the transfer of electrons between reactants. Oxidation numbers are used to identify oxidized and reduced species in redox reactions. Redox equations can be balanced using two methods - comparing oxidation number changes or using half-reactions. Electrochemical cells use spontaneous redox reactions to generate electricity, with oxidation occurring at the anode and reduction at the cathode.
This document discusses concepts related to motion including position, relative and absolute position, distance and displacement, speed and velocity, uniform and non-uniform motion, and uniformly accelerated motion. It defines key terms and concepts and provides examples to illustrate them. Position can be relative or absolute depending on the reference point used. Distance refers to the total path travelled, while displacement refers to the net or direct distance between two points. Speed is the rate of change of distance and velocity is the rate of change of displacement, making velocity a vector quantity. Uniform motion involves equal distances in equal times while non-uniform motion does not. Uniformly accelerated motion follows three equations of motion and has constant acceleration.
This document provides information about atoms, molecules, ions and chemical formulas. It discusses key concepts such as:
- Dalton's atomic theory which states that matter is made of tiny indivisible particles called atoms that combine in small whole number ratios.
- Atoms have symbols to represent them and an atomic mass that is measured relative to carbon-12. Molecules are groups of atoms that are chemically bonded.
- Chemical formulas show the types and numbers of atoms or ions that make up a compound. Formulas are written with the cation written first followed by the anion.
This document discusses electrochemistry and redox reactions. It defines oxidation and reduction, and explains that they occur together in redox reactions. Redox reactions involve the transfer of electrons from one species to another, changing their oxidation states. Oxidizing agents accept electrons from reducing agents. Common oxidizing agents are hydrogen peroxide and chlorine, while common reducing agents are carbon and hydrogen. Redox reactions can be identified by changes in oxidation states and through color changes using indicators like potassium manganate or iodide.
The document discusses oxidation-reduction (redox) reactions. It defines oxidation as the loss of electrons and reduction as the gain of electrons. Redox reactions always involve both an oxidation and a reduction process. The substance undergoing oxidation is the reducing agent as it loses electrons, while the substance undergoing reduction is the oxidizing agent as it gains electrons. Assigning oxidation numbers to atoms allows identification of which substances are oxidized and reduced in a reaction.
Sulfur has an oxidation number of +4 in SO2 and Na2SO4. Carbon has an oxidation number of +4 in CO32-. Oxygen has an oxidation number of -2 in all cases. Sodium has an oxidation number of +1 in Na2SO4. Nitrogen has an oxidation number of -3 and hydrogen has an oxidation number of +1 in (NH4)2S.
This document discusses oxidation-reduction (redox) reactions. It defines oxidation as the gain of oxygen, loss of hydrogen, or loss of electrons, while reduction is defined as the loss of oxygen, gain of hydrogen, or gain of electrons. Redox reactions involve both oxidation and reduction occurring simultaneously. Oxidation numbers are assigned to elements to indicate degree of oxidation. Common oxidizing agents are listed that can oxidize Fe2+ to Fe3+, like potassium manganate, chlorine, and nitric acid. Reducing agents like sodium sulfite can reduce Fe3+ back to Fe2+. The document also covers IUPAC nomenclature of inorganic compounds and rules for assigning oxidation numbers.
1. Electrochemistry involves redox reactions where one element is oxidized and another is reduced. Oxidation is the loss of electrons and an increase in oxidation number, while reduction is the gain of electrons and a decrease in oxidation number.
2. Electrolysis is the passage of an electric current through an ionic substance to cause a non-spontaneous redox reaction. Oxidation occurs at the anode and reduction at the cathode.
3. Aluminum is extracted from bauxite via electrolysis. Bauxite is dissolved in molten cryolite to lower its melting point, then electrolysis separates aluminum ions at the cathode.
Redox reactions involve the transfer of electrons from one reactant to another, resulting in oxidation and reduction. Oxidation is the loss of electrons and reduction is the gain of electrons. Common redox reactions include photosynthesis, respiration, combustion, and production and use of fertilizers.
Includes a discussion of Voltaic and electrolytic cells, the Nernst equation and the relationship between electrochemical processes, chemical equilibrium and free energy.
**More good stuff available at:
www.wsautter.com
and
http://www.youtube.com/results?search_query=wnsautter&aq=f
This document provides an overview of redox reactions and electrochemistry applications. It discusses oxidation-reduction concepts like oxidation states and the OIL RIG mnemonic. Examples of redox reactions and electrochemistry applications are given, including galvanic cells, corrosion, electrolysis, and batteries. Key concepts covered include cell potential, the Nernst equation, and how concentration affects cell potential. Diagrams illustrate galvanic cells and how they function.
This document provides an overview of redox (oxidation-reduction) chemistry. It defines oxidation as the loss of electrons and reduction as the gain of electrons. Oxidation and reduction always occur simultaneously in redox reactions. The document discusses identifying oxidizing and reducing agents, balancing redox reactions using the half-reaction method, standard reduction potentials, galvanic (voltaic) cells that produce electricity from spontaneous redox reactions, and electrolytic cells that use electricity to drive nonspontaneous reactions.
This document provides information on chemical equations and oxidation-reduction reactions. It defines key concepts like oxidation, reduction, oxidizing agents, reducing agents and oxidation numbers. Examples of different types of chemical reactions like combination, decomposition, displacement and combustion are outlined. Steps for writing and balancing chemical equations are described. Oxidation-reduction reactions are explained along with biological examples of electron transfer processes. Specific equations are given and identified as oxidation or reduction reactions.
Redox Reaction and Electrochemical Cell (Reaksi Redoks dan Sel Elektrokimia)DindaKamaliya
An electrochemical cell converts chemical energy into electrical energy through spontaneous redox reactions. It consists of two half-cells separated by a salt bridge. In the cathode half-cell, reduction occurs as oxidized species gain electrons. In the anode half-cell, oxidation occurs as reduced species lose electrons. Electrons flow through an external circuit from the anode to the cathode. The standard electrode potential of each half-reaction predicts the cell's voltage under standard conditions.
This document contains lecture notes on quantitative analysis in chemistry. It discusses gravimetric analysis, which determines the amount of a substance by converting it into a product that can be isolated and weighed. An example is given of determining the amount of lead in water by precipitating lead sulfate, filtering and weighing the precipitate. A practice problem demonstrates calculating the mass of lead from the mass of lead sulfate precipitate obtained.
This document provides an overview of redox reactions including:
- Redox reactions involve the transfer of electrons between chemical species, resulting in oxidation and reduction.
- Oxidizing agents gain electrons and are reduced, while reducing agents lose electrons and are oxidized.
- Latimer, Frost, and Pourbaix diagrams can be used to predict and understand redox reactions in aqueous solutions by showing the thermodynamic stability of different oxidation states.
- Key concepts like disproportionation, oxidizing/reducing abilities, and stable/unstable species can be determined from these types of diagrams.
This document provides an overview of oxidation-reduction (redox) reactions including:
1. Defining oxidation as the loss of electrons and reduction as the gain of electrons.
2. Explaining that the reducing agent is oxidized and the oxidizing agent is reduced in redox reactions.
3. Detailing that the first step in understanding redox reactions is determining the oxidation number of each element in a compound.
This document provides an overview of redox (reduction-oxidation) reactions, including definitions of key terms like oxidation, reduction, oxidizing agents, reducing agents, and disproportionation reactions. It discusses identifying oxidation and reduction based on changes in oxygen, hydrogen, or electron content. Methods for determining oxidation states and balancing redox reactions using the half-reaction method are also described. Real-world examples of redox processes like corrosion and the blue bottle experiment are mentioned.
This document provides a summary of key concepts in oxidation-reduction (redox) reactions:
1) Redox reactions involve the transfer of electrons between chemical species, either through the complete transfer of electrons to form ionic bonds, or partial transfer to form covalent bonds. Oxidation is the loss of electrons and reduction is the gain of electrons.
2) Redox pairs are couples of oxidized and reduced forms of elements that differ in their oxidation state. Common redox pairs include Fe3+/Fe2+, O2/H2O, and MnO4-/MnO2.
3) The standard electrode potential (E°) indicates the tendency of half-reactions to occur.
The document discusses oxidation and reduction reactions. It provides examples of oxidation as the loss of electrons and reduction as the gain of electrons. Corrosion is described as an oxidation reaction where metals react with substances like oxygen, water and acids. Electroplating is introduced as a process where a thin coating of a metal like copper or nickel is deposited onto a surface through an oxidation-reduction reaction.
This document discusses oxidation numbers, which indicate the degree of oxidation or reduction of atoms in chemical compounds. Oxidation numbers are assigned according to rules such as atoms in their elemental form having an oxidation number of 0, monoatomic ions having the same charge as their oxidation number, and the sum of oxidation numbers in a neutral molecule being 0. Oxidation numbers can be fractional and are indicated in Roman numerals after the symbol of a metal in its compound. An example of a fractional oxidation number is given for sulfur atoms in the tetrathionate ion.
The document discusses oxidation-reduction (redox) reactions. It defines oxidation as loss of electrons and reduction as gain of electrons. Redox reactions involve the transfer of electrons between reactants. Oxidation numbers are used to identify oxidized and reduced species in redox reactions. Redox equations can be balanced using two methods - comparing oxidation number changes or using half-reactions. Electrochemical cells use spontaneous redox reactions to generate electricity, with oxidation occurring at the anode and reduction at the cathode.
This document discusses concepts related to motion including position, relative and absolute position, distance and displacement, speed and velocity, uniform and non-uniform motion, and uniformly accelerated motion. It defines key terms and concepts and provides examples to illustrate them. Position can be relative or absolute depending on the reference point used. Distance refers to the total path travelled, while displacement refers to the net or direct distance between two points. Speed is the rate of change of distance and velocity is the rate of change of displacement, making velocity a vector quantity. Uniform motion involves equal distances in equal times while non-uniform motion does not. Uniformly accelerated motion follows three equations of motion and has constant acceleration.
This document provides information about atoms, molecules, ions and chemical formulas. It discusses key concepts such as:
- Dalton's atomic theory which states that matter is made of tiny indivisible particles called atoms that combine in small whole number ratios.
- Atoms have symbols to represent them and an atomic mass that is measured relative to carbon-12. Molecules are groups of atoms that are chemically bonded.
- Chemical formulas show the types and numbers of atoms or ions that make up a compound. Formulas are written with the cation written first followed by the anion.
1. The document discusses concepts related to gravitation including Newton's Universal Law of Gravitation, Kepler's Laws of Planetary Motion, gravitational force, and acceleration due to gravity.
2. Key points covered include Newton's inverse square law formula for gravitational force, Kepler's three laws of planetary motion, and definitions of free fall and acceleration due to gravity.
3. The document also discusses properties of gravitational force and provides examples to illustrate concepts like why objects on Earth do not continuously accelerate towards each other due to gravitational attraction.
This document defines key concepts in matter, including:
- Matter is anything that has mass and occupies space, and can be made of elements, molecules, or compounds.
- Elements are the simplest forms of matter and cannot be broken down further into simpler substances. Atoms are the smallest particle of an element.
- Atoms bond together to form molecules or bond with different elements to form compounds like salt or water.
- All matter is made of atoms in constant motion according to kinetic theory. Atoms are arranged in the periodic table by atomic number.
- Atoms contain a nucleus of protons and neutrons, with electrons orbiting the nucleus in shells. The number of protons determines the element.
This document defines key concepts in matter, including:
- Matter is anything that has mass and occupies space, and can be made of elements, molecules, or compounds.
- Elements are the simplest forms of matter and cannot be broken down further into simpler substances. Atoms are the smallest particle of an element.
- Atoms bond together to form molecules or bond with different elements to form compounds like salt or water.
- All matter is made of atoms that are always in motion according to kinetic theory. Heavier atoms move slower than lighter ones.
- The periodic table organizes all known elements by their atomic structure.
This document provides an overview of the classification of living organisms. It discusses the need for classification, the basis used for classification including cell structure and nutrition type, and the hierarchical system used from kingdoms down to species. The five kingdom system is described, including Monera, Protista, Fungi, Plantae, and Animalia. Details are given on the classification of plants into five groups and animals into ten groups, with examples provided. The classification of living organisms arranges organisms into taxonomic groups based on similarities to allow for organized study.
Sound is a form of energy that propagates as longitudinal waves, requiring a medium. It is produced by vibrating objects and transmitted through compression and rarefaction variations in the medium. The human ear can detect sounds between 20 Hz to 20 kHz. Ultrasound with frequencies above this range has applications like medical imaging and industrial cleaning, while infrasound below 20 Hz is used for communication by some animals. Sonar uses ultrasound for underwater detection of objects.
The document discusses natural resources and the four main spheres of Earth - the lithosphere, hydrosphere, atmosphere, and biosphere. It describes each sphere and their composition. The document then discusses various natural resources like air, water, soil minerals, and their importance. It also discusses pollution of these resources and processes like the water cycle and biogeochemical cycles.
This document discusses factors related to health and disease. It defines health as a state of complete physical, mental and social well-being. Important characteristics of good health include being free from sickness, anxiety, and tensions. Health can fail due to poor physical/social environments, economic conditions, or lack of social equality. The document distinguishes between "healthy" and "disease-free," and outlines different types of diseases including acute, chronic, infectious/communicable, and non-infectious/non-communicable diseases. It describes causes of disease such as pathogens, genetic disorders, pollution and malnutrition. Means of disease transmission include air, water, food, vectors, contact and sexual contact. Principles of treatment are to
Here are the answers:
a) Disease is an abnormal condition affecting the body or mind that impairs normal functioning and causes discomfort.
b) The two major categories of human diseases are:
1. Infectious diseases - caused by pathogens like bacteria, viruses, fungi or parasites.
Examples: Malaria (caused by protozoan Plasmodium), Tuberculosis (caused by bacteria Mycobacterium tuberculosis)
2. Non-infectious diseases - not caused by pathogens. Develop due to genetic reasons, unhealthy lifestyle or environmental factors.
Examples: Cancer (uncontrolled cell growth), Heart disease (caused by risk factors like hypertension, smoking, obesity)
Here are the answers:
a) Disease is an abnormal condition affecting the body or mind that impairs normal functioning and causes discomfort.
b) The two major categories of human diseases are:
1. Infectious diseases - caused by pathogens like bacteria, viruses, fungi or parasites.
Examples: Malaria (caused by a protozoan parasite), Tuberculosis (caused by Mycobacterium tuberculosis bacteria)
2. Non-infectious diseases - not caused by pathogens. These include genetic diseases, cancer, heart diseases, mental illnesses etc.
Examples: Diabetes (caused due to malfunctioning of pancreas), Asthma (caused due to hypersensitivity of airways
Here are the key causes of cancer:
- Genetic factors - Some people inherit gene mutations from their parents that increase their risk of certain cancers.
- Tobacco use - Smoking or chewing tobacco is linked to cancers of the lung, esophagus, larynx, mouth, bladder, kidney, liver, stomach, pancreas, and colon/rectum. Tobacco contains chemicals that can damage DNA.
- Diet and obesity - A diet high in red/processed meats and low in fruits and vegetables increases cancer risk. Obesity is linked to several cancers. Excess weight increases hormone levels and inflammation.
- Radiation - Both natural sources like radon and man-made sources like X-rays can damage
Here are the key causes of cancer:
- Genetic factors - Some people inherit gene mutations from their parents that increase their risk of certain cancers.
- Tobacco use - Smoking or chewing tobacco is linked to cancers of the lung, esophagus, larynx, mouth, bladder, kidney, liver, stomach, pancreas, colon and rectum, and acute myeloid leukemia.
- Diet and obesity - A diet high in red/processed meats and low in fruits and vegetables increases the risk of several cancers. Obesity is linked to increased risk of multiple cancers.
- Alcohol use - Heavy drinking is linked to cancers of the mouth, esophagus, throat, liver and breast.
-
Sound is a form of energy that propagates as longitudinal waves, requiring a medium. It is produced by vibrating objects and transmitted through compression and rarefaction variations in the medium. The human ear can detect sounds between 20 Hz to 20 kHz. Ultrasound with frequencies above this range has applications like medical imaging and material cleaning, while infrasound below 20 Hz is used by some animals. Sonar also uses ultrasound for underwater object detection.
1. The document discusses Heinrich Hertz's experiments with sound and how it is produced through vibration and propagates as longitudinal waves through a medium like air.
2. Key experiments shown include using a vibrating tuning fork to produce compressions and rarefactions in air, demonstrating that sound needs a material medium to travel, and that the speed of sound depends on the medium and temperature.
3. Applications of sound reflection, resonance, infrasound, ultrasound, and SONAR are also summarized.
This document describes various concepts related to motion including:
1. Motion is defined as the change in position of a body over time. Distance moved is the total path travelled, while displacement is the shortest distance between initial and final positions.
2. Uniform motion means equal distances are travelled in equal times, while non-uniform motion means unequal distances in equal times. Examples of each are given.
3. Speed, velocity, average speed, average velocity, acceleration, and equations of motion relating these quantities are defined and explained with examples. Distance-time graphs and their use in representing motion are also described.
This document provides an overview of the classification of living organisms. It discusses the need for classification, the basis used for classification including cell structure and nutrition type, and the hierarchical system used from kingdoms down to species. The five kingdom system is described, including Monera, Protista, Fungi, Plantae, and Animalia. An overview of the classification of plants and animals is also provided, outlining the main groups within each kingdom.
1. Biodiversity refers to the variety of living organisms on Earth, including plants, animals, and microorganisms. 2. Taxonomy is the science of classifying organisms using their similarities and differences. A key aspect is assigning each organism a unique scientific name. 3. The binomial system of nomenclature assigns every organism a genus and species name, allowing for uniform identification worldwide.
Plants and animals are made of different types of tissues due to differences in their structure and function. In multicellular organisms, cells are grouped together into tissues to efficiently perform specialized functions. In plants, tissues include epidermis, a protective outer layer of flat cells covered with a waxy cuticle. Meristematic tissues contain actively dividing cells and produce permanent tissues through differentiation. Permanent tissues include parenchyma, collenchyma, and sclerenchyma, which provide structure and support to plants. Stomata in the leaf epidermis allow for gas exchange and transpiration. As plants age, cork replaces the epidermis and protects the bark.
ISO/IEC 27001, ISO/IEC 42001, and GDPR: Best Practices for Implementation and...PECB
Denis is a dynamic and results-driven Chief Information Officer (CIO) with a distinguished career spanning information systems analysis and technical project management. With a proven track record of spearheading the design and delivery of cutting-edge Information Management solutions, he has consistently elevated business operations, streamlined reporting functions, and maximized process efficiency.
Certified as an ISO/IEC 27001: Information Security Management Systems (ISMS) Lead Implementer, Data Protection Officer, and Cyber Risks Analyst, Denis brings a heightened focus on data security, privacy, and cyber resilience to every endeavor.
His expertise extends across a diverse spectrum of reporting, database, and web development applications, underpinned by an exceptional grasp of data storage and virtualization technologies. His proficiency in application testing, database administration, and data cleansing ensures seamless execution of complex projects.
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Throughout his career, he has taken on multifaceted roles, from leading technical project management teams to owning solutions that drive operational excellence. His conscientious and proactive approach is unwavering, whether he is working independently or collaboratively within a team. His ability to connect with colleagues on a personal level underscores his commitment to fostering a harmonious and productive workplace environment.
Date: May 29, 2024
Tags: Information Security, ISO/IEC 27001, ISO/IEC 42001, Artificial Intelligence, GDPR
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How to Setup Warehouse & Location in Odoo 17 InventoryCeline George
In this slide, we'll explore how to set up warehouses and locations in Odoo 17 Inventory. This will help us manage our stock effectively, track inventory levels, and streamline warehouse operations.
This presentation includes basic of PCOS their pathology and treatment and also Ayurveda correlation of PCOS and Ayurvedic line of treatment mentioned in classics.
Chapter wise All Notes of First year Basic Civil Engineering.pptxDenish Jangid
Chapter wise All Notes of First year Basic Civil Engineering
Syllabus
Chapter-1
Introduction to objective, scope and outcome the subject
Chapter 2
Introduction: Scope and Specialization of Civil Engineering, Role of civil Engineer in Society, Impact of infrastructural development on economy of country.
Chapter 3
Surveying: Object Principles & Types of Surveying; Site Plans, Plans & Maps; Scales & Unit of different Measurements.
Linear Measurements: Instruments used. Linear Measurement by Tape, Ranging out Survey Lines and overcoming Obstructions; Measurements on sloping ground; Tape corrections, conventional symbols. Angular Measurements: Instruments used; Introduction to Compass Surveying, Bearings and Longitude & Latitude of a Line, Introduction to total station.
Levelling: Instrument used Object of levelling, Methods of levelling in brief, and Contour maps.
Chapter 4
Buildings: Selection of site for Buildings, Layout of Building Plan, Types of buildings, Plinth area, carpet area, floor space index, Introduction to building byelaws, concept of sun light & ventilation. Components of Buildings & their functions, Basic concept of R.C.C., Introduction to types of foundation
Chapter 5
Transportation: Introduction to Transportation Engineering; Traffic and Road Safety: Types and Characteristics of Various Modes of Transportation; Various Road Traffic Signs, Causes of Accidents and Road Safety Measures.
Chapter 6
Environmental Engineering: Environmental Pollution, Environmental Acts and Regulations, Functional Concepts of Ecology, Basics of Species, Biodiversity, Ecosystem, Hydrological Cycle; Chemical Cycles: Carbon, Nitrogen & Phosphorus; Energy Flow in Ecosystems.
Water Pollution: Water Quality standards, Introduction to Treatment & Disposal of Waste Water. Reuse and Saving of Water, Rain Water Harvesting. Solid Waste Management: Classification of Solid Waste, Collection, Transportation and Disposal of Solid. Recycling of Solid Waste: Energy Recovery, Sanitary Landfill, On-Site Sanitation. Air & Noise Pollution: Primary and Secondary air pollutants, Harmful effects of Air Pollution, Control of Air Pollution. . Noise Pollution Harmful Effects of noise pollution, control of noise pollution, Global warming & Climate Change, Ozone depletion, Greenhouse effect
Text Books:
1. Palancharmy, Basic Civil Engineering, McGraw Hill publishers.
2. Satheesh Gopi, Basic Civil Engineering, Pearson Publishers.
3. Ketki Rangwala Dalal, Essentials of Civil Engineering, Charotar Publishing House.
4. BCP, Surveying volume 1
A workshop hosted by the South African Journal of Science aimed at postgraduate students and early career researchers with little or no experience in writing and publishing journal articles.
Strategies for Effective Upskilling is a presentation by Chinwendu Peace in a Your Skill Boost Masterclass organisation by the Excellence Foundation for South Sudan on 08th and 09th June 2024 from 1 PM to 3 PM on each day.
Leveraging Generative AI to Drive Nonprofit InnovationTechSoup
In this webinar, participants learned how to utilize Generative AI to streamline operations and elevate member engagement. Amazon Web Service experts provided a customer specific use cases and dived into low/no-code tools that are quick and easy to deploy through Amazon Web Service (AWS.)
This presentation was provided by Steph Pollock of The American Psychological Association’s Journals Program, and Damita Snow, of The American Society of Civil Engineers (ASCE), for the initial session of NISO's 2024 Training Series "DEIA in the Scholarly Landscape." Session One: 'Setting Expectations: a DEIA Primer,' was held June 6, 2024.
it describes the bony anatomy including the femoral head , acetabulum, labrum . also discusses the capsule , ligaments . muscle that act on the hip joint and the range of motion are outlined. factors affecting hip joint stability and weight transmission through the joint are summarized.
2. Tro - Chapter 16 2
Oxidation-Reduction Reactions
• oxidation-reduction reactions are also called redox reactions
• all redox reactions involve the transfer of electrons from one
atom to another
• spontaneous redox reactions are generally exothermic, and we
can use their released energy as a source of energy for other
applications
3. Development of oxidation and reduction reaction concept
1. Reaction of reduction oxidation based on releasing (lossing) and
gaining of oxygen
a. Oxidation reaction
Oxidation reaction is a reaction of gaining (capturing) of oxygen
by a substance
Example :
CH4(g)
+ 2O2(g)
CO2(g)
+ 2H2
Og)
b. Reduction reaction
Reduction reaction is a reaction of releasing (losing) of oxygen from a
oxide compound
Example:
Fe2
O3(s)
+ 3CO(g)
2Fe(s)
+ 3CO2(g)
P4(s)
+ 5O2(g)
2P2
O5(s)
CuO(s)
+ H2(g)
Cu(s)
+ H2
O(g)
4. Development of oxidation and reduction reaction concept
1. Reaction of reduction oxidation based on releasing (lossing) and
gaining of oxygen
a. Oxidation reaction
Oxidation reaction is a reaction of gaining (capturing) of oxygen
by a substance
Example :
CH4(g)
+ 2O2(g)
CO2(g)
+ 2H2
Og)
b. Reduction reaction
Reduction reaction is a reaction of releasing (lossing) of oxygen from a
oxide compound
Example:
Fe2
O3(s)
+ 3CO(g)
2Fe(s)
+ 3CO2(g)
P4(s)
+ 5O2(g)
2P2
O5(s)
CuO(s)
+ H2(g)
Cu(s)
+ H2
O(g)
5. What do you mean by oxidation
and reduction ?
• Oxidation can be defined as addition of
oxygen/electronegative element to a substance or
removal of hydrogen/ electropositive element from a
substance.
• Reduction can be defined as removal of
oxygen/electronegative element from a substance or
addition of hydrogen/ electropositive element to a
substance.
5
6. 6
Oxidation and ReductionAnother Def
• in order to convert a free element into an ion, the
atoms must gain or lose electrons
of course, if one atom loses electrons, another must
accept them
• reactions where electrons are transferred from one
atom to another are redox reactions
• atoms that lose electrons are being oxidized, atoms
that gain electrons are being reduced
2 Na(s) + Cl2(g) → 2 Na+
Cl–
(s)
Na → Na+
+ 1 e–
oxidation
Cl2 + 2 e–
→ 2 Cl–
reduction
7. What is an oxidizing and
reducing agent ?
• Oxidising agent: a reagent which
increases the oxidation number of an
element of a given substance. These
reagents are called oxidants.
• Reducing agent: a reagent that lowers
the oxidation number of a given
element . These reagents are also
called reductants.
7
8. 8
Oxidation–Reduction
• oxidation and reduction must occur simultaneously
if an atom loses electrons another atom must take them
• the reactant that reduces an element in another reactant
is called the reducing agent
the reducing agent contains the element that is oxidized
• the reactant that oxidizes an element in another reactant
is called the oxidizing agent
the oxidizing agent contains the element that is reduced
2 Na(s) + Cl2(g) → 2 Na+
Cl–
(s)
Na is oxidized, Cl is reduced
Na is the reducing agent, Cl2 is the oxidizing agent
9. 9
Oxidation and Reduction
A Better Definition
• oxidation occurs when an atom’s oxidation
state increases during a reaction
• reduction occurs when an atom’s oxidation
state decreases during a reaction
CH4 + 2 O2 → CO2 + 2 H2O
-4 +1 0 +4 –2 +1 -2
oxidation
reduction
10. 10
Will a Reaction Take Place?
• reactions that are energetically favorable are
said to be spontaneous
they can happen, but the activation energy may
be so large that the rate is very slow
• the relative reactivity of metals can be used
to determine if some redox reactions are
spontaneous
11. Electron transfer reactions
• Place a strip of metallic zinc in an aqueous solution of copper nitrate , for about one
hour. You may notice that the strip becomes coated with reddish metallic copper
and the blue colour of the solution disappears. Formation of Zn2+ ions among the
products can easily be judged when the blue colour of the solution due to Cu2+ has
disappeared. If hydrogen sulphide gas is passed through the colourless solution
containing Zn2+ ions, appearance of white zinc sulphide, ZnS can be seen on
making the solution alkaline with ammonia.
• The reaction between metallic zinc and the aqueous solution of copper nitrate is :-
• In this reaction , zinc has lost electrons to form Zn2+and, therefore, zinc is oxidised.
Evidently, now if zinc is oxidised, releasing electrons , copper ions is reduced by
gaining electrons from zinc.
11
12. 12
At this stage we may investigate the state of equilibrium for the reaction
represented by equation . For this purpose, let us place a strip of metallic
copper in a zinc sulphate solution. No visible reaction is noticed and
attempt to detect the presence of Cu2+ ions by passing H2S gas through
the solution to produce black colour cupric sulhpide. CuS, does not
succeed. Cupric sulphide has such a low solubility that this is an
extremely sensitive test. Cu2+ cannot be detected. Hence the equilibrium
For the reaction favours the products over the reactants.
13. • This suggests that we might develop a table in which metals and
their ions are listed on the basis of their tendency to release electrons
just as we do in the case of acids to indicate the strength of the acids.
As a matter of fact we have already made certain comparisons. By
comparison we have come to know that zinc releases electrons to
copper and copper releases electrons to silver and therefore electron
releasing tendency is in the order Zn>Cu>Ag .
13
14. 2. Reduction oxidation reaction based on electron transfer
a. Oxidation reaction
Oxidation reaction is a reaction of electron releasing (lossing) from a substance.
Example:
b. Reduction reaction
Reduction reaction is a reaction of electron gaining by a substance.
Example:
Na Na+
+ e−
Mg Mg2+
+ 2 e−
Cu Cu2+
+ 2 e−
Cl2
+ 2e−
2Cl −
S + 2 e−
S2−
15. Stock notation
• Stock notation is the notation used where the
oxidation state of the element is represented by
roman numerals.
Tro - Chapter 16 15
16. IUPAC Nomenclature
The compound that is formed by the elements have more than one type of
oxidation number , its name diferentiated by the Roman number writing in
the bracket in the back of that element name. The Roman number shows
the value of oxidation number of that element.
The compound that is formed by the element only has one type of oxidation
number, the Roman number does not need writen.
This IUPAC nomenclature applies in both ionic and covalent
compounds.
Examples IUPAC name of binary covalent compound:
CO : carbon(II) oxide (oxidation number of C = +2)
CO2
: carbon(IV) oxide (oxidation number of C = +4)
P2
O3
: phosphorus(III) oxide (oxidation number of P = +3)
N2O5 : nitrogen(V) oxide (oxidation number of N = +5)
Cl2O7 : chlorine(VII) oxide (oxidation number of Cl = +7)
17. 17
Oxidation state
• For reactions that are not metal + nonmetal, or do
not involve O2, we need a method for determining
how the electrons are transferred
• chemists assign a number to each element in a
reaction called an oxidation state that allows them
to determine the electron flow in the reaction
• Basically oxidation number denotes the
oxidation state of the element in a compound
according to a set of rules formulated on the
basis that electron fair in a covalent bond
belongs entirely to the more electrovalent bond.
18. Oxidation Numbers
• Oxidation is the loss of electrons;
Reduction is the gain of electrons
• Oxidation and reduction go together.
Whenever a substance loses electrons
and another substance gains electrons
• Oxidation Numbers are a system that
we can use to keep track of electron
transfers
19. Common Oxidation States
Chemical species Oxidation state and remarks
Any element eg Fe, O2 , S8 zero
Oxygen in any compound -2 except in peroxides example H2O2 or Na2O2 then oxygen atom has
oxidation state of -1 or in F2O , then oxygen atom has oxidation state
of +2
Fluorine in any compound -1 being most electronegative
Hydrogen in any compound +1 except in metal hydrides example NaH then hydrogen atom has
oxidation state of -1 as metals have a greater tendency to lose
electrons
Chlorine, bromine, iodine -ve oxidation state if bonded to less electronegative element eg
NaCl; then Cl = -1.
+ve oxidation state if bonded to more electronegative element eg
ClO-
, then Cl = +1; ClO3
-
, then Cl = +5
20. Oxidation Number
Oxdidation number is a number that states electrical charge possessed by each
one element atom in the molecular compound or the ion.
In the molecules of ionic compound, electrical charge contained element atom can
be raised by transfering of electrons.
In the formation of ionic bond:
-Metal atom losses electron to form the positive ion.
-Nonmetal atom gains electron to form the negative ion.
In the molecule of MgF2
, consist of Mg2+
ion with charge of 2+ dan F-
ion with
charge of 1−
Said that in the molecule of MgF2
, oxidation number of Mg is +2, and oxidation
number of F is -1.
In the molecule of covalent compound, the raising of the electrical charge each
element atom is caused by its existence the difference of electronegativity of
element, so that occur polarization covalent bond.
In the polar covalent compound, the more electronegative atom become more
negative charge and the other atom become more positive charge.
In the polar covalent compound of H2
O, H contain 1+ and O contain 2−
21. Determining Oxidation Numbers of Elements
The oxidation number of an element in the molecule or in the ion, by use
the rules of oxidation numbers can be determined.
•Write down the molecular or ionic formula which will be determined
oxidation number its element and between one atom of element and the
others, given enough space.
•Write each oxidation number of elements in below it and write x for
element that will be determined its oxidation number.
•Use the rules of oxidation number, that is rule of number 7 or 8, for
determine x value.
Determine the following element oxidation number
a. S in molecule of H2
SO4
b. Cr in ion of Cr2
O7
2−
Given : Molecule of H2
SO4
Ion of Cr2
O7
2−
Find : a. oxidation number of S in H2
SO4
b. oxidation number of Cr in Cr2
O7
2−
Solution :
Example:
22. Rules for determining oxidation number
• (1) In elements in the free or the uncombined
state each atoms bears an oxidation number of
zero.
• (2)For ions composed of only 1 atom the oxidation
number is equal to the charge on the ion.
• (3)For oxygen in the case of superoxide's and
peroxides oxidation state is assigned to oxygen as
-1 or –(½).
09/25/15 22
23. • (4)The second exception with oxygen is with the
fluorides and di-fluorides here the oxygen has an
oxidation state of +2 and+1.
• (5)The number assigned to oxygen will depend upon
the bonding state of oxygen but this will have a
positive number.
• (6)the oxidation state of hydrogen is +1, except when
it is bonded with elements with binary compounds.
When it is bonded with lithium , beryllium it has the
oxidation state of -1.
23
24. • (7)In all its compounds fluorine has an
oxidation state of -1.other halogens like
chlorine , bromine and iodine have also the
oxidation state as -1.except oxoanions and
oxoacids.
• (8)The algebraic sum of the oxidation number
of all the atoms in a compound must be zero. In
polyatomic ions the algebraic sum of all the
oxidation numbers of atoms of the ion must be
equal to the charge on the ion.
24
25. Oxidation Number basic Rules
1. Oxidation number of free elements
Free elements (include molecular elements: H2
, O2
, O3
, N2
, F2
, Cl2
, Br2
, I2
,
P4
, S8
) have oxidation number of 0 (zero).
2. Oxidation number of fluorine
In its compounds, oxidation number of F always –1.
3. Oxidation number of hydrogen
In its compounds, oxidation number of H always +1.
Except, hydrogen in the hydride compounds (compound of H with metal),
oxidation number of H, is –1
Example:
In the compound of H2
O, NH3
, H2
S, HCl, HNO3,
H2
SO4,
oxidation
number of H, is +1
In the hydride compound, like LiH, NaH, MgH2
, oxidation number
of H, is –1
26. 26
Rules for Assigning Oxidation States
5. in their compounds, nonmetals have oxidation
states according to the table below
nonmetals higher on the table take priority
Nonmetal Oxidation State Example
F -1 CF4
H +1 CH4
O -2 CO2
Group 7A -1 CCl4
Group 6A -2 CS2
Group 5A -3 NH3
27. Sn4+
(aq)
Sn2+
(aq)
Cl2(g)
2 Cl-
(g)
A. Reduction oxidation reaction based on oxidation
number change
a. Oxidation reaction
Oxidation reaction is a chemical reaction which is accompanied by increasing of
oxidation number.
Example:
b. Reduction reaction
Reduction reaction is a chemical reaction which is
accompanied by decreasing of oxidation number.
Al(s)
Al3+
(aq)
S2-
(aq)
S(s)
Example:
28. Table 8.1.
Oxidation numbers of several elements of group B
Elements of group B
Oxidation numbers
Name Symbol
Zink Zn +2
Silver Ag +1
Copper Cu +1, +2
Gold Au +1, +3
Iron Fe +2, +3
Lead Pb +2, +4
b. Generally, metallic elements of group B has oxidation number more than
one type.
Example:
6. Oxidation number of monoatomic ion
Oxidation number of mono atomic ions is equal to the charge on that ion
Example:
Na+
ion has oxidation number of +1
Ba2+
ion has oxidation number of +2
Fe3+
ion has oxidation number of +3
Cl−
ion has oxidation number of –1
S2−
ion has oxidation number of –2
29. 7. The sum of oxidation number of element atoms in a compound
molecule is equal to 0 (zero)
∑ o. n. of element in compound molecule = 0
Example: H2O
(o.n. of H x 2) + (o.n. of O x 1) = 0
{(+1) x2} + {(-2) x 1} = 0
{+2} + {-2} = 0
8. The sum of oxidation number of element atoms in a polyatomic ion
is equal to the charge on that ion.
∑ o. n. of element in ion = charge of ion
Example: OH−
(o.n. of O x 1) + (o.n. of H x 1) = -1
{(-2) x 1} + {(+1) x 1} = -1
{-2} + {+1} = -1
30. a. H2
SO4
∑ o. n. element in molecule = 0
( 2 x o. n. H) + (1 x o.n. S) + (4 x o.n. O) = 0
{ 2 x (+1 ) } + { 1 x (x ) } + { 4 x (–2) = 0
( +2) + (x) + (–8) = 0
x = +8 – 2 x = +6
The oxidation number of S in H2
SO4
is +6
o. n. H = +1, o. n. O = –2, o. n. S = x
H2 S O4
+1 –2
x
b. Cr2
O7
2−
∑ o. n. of element in ion = charge of ion
( 2 x o. n. Cr ) + ( 7 x o.n. O ) = –2
{ 2 x (x) } + { 7 x (–2) } = –2
( 2x ) + ( –14 ) = –2 2x = +14 - 2
x = + 6
The oxidation number of Cr in CrO4
2–
is +6
o. n. O = –2, o. n. Cr = x
O7 )
–2
x
( Cr2
2–
x = 2
12
+
31. Example 1
Let the oxidation state of Mn be x.
Thus, in MnO4
-
, x + 4(-2) = -1
x = +7
• Manganese is reduced from oxidation state of
+7 in MnO4
-
to +2 in Mn2+
, while iron is
oxidised from oxidation state of +2 in Fe2+
to
3+
O
H
Fe
Mn
H
Fe
MnO 2
3
2
2
4 4
5
8
5 +
+
→
+
+ +
+
+
+
−
+7 +2 +2 +3
32. Limitations of oxidation number
• The main limitation of oxidation number is
that oxidation number cannot be assigned a
particular species.
• The secondary limitation is that in recent
past it has been found out that the
oxidation process is visualized as a
decrease in electron density and reduction
process as an increase in electron density
around the atom(s) involved in the reaction.
09/25/15 32
33. Paradox of fractional oxidation
number
• Sometimes we come across compounds having
fractional oxidation number.
• Examples C3O2 where carbon has the oxidation
state of 4/3.
Br3O8 where bromine has a oxidation state of 16/3.
Na2S4O6 where sodium has an oxidation state of 2.5 .
33
34. • Fractional oxidation states are often used to represent the average
oxidation states of several atoms of the same element in a structure.
• Br3O8 has a oxidation state of 16/3 whereas it actually possess a
oxidation state of +4 and +6.
• Similarly thiosulphate ion exhibits oxidation state of +5 and 0 and
hence the average or fractional oxidation state becomes 2.5.( in reality
it possess +5 and +5 oxidation state )!
• Similarly carbon suboxide experiences a fractional oxidation state of
4/3 whereas each carbon has a oxidation state of +2 and +2.
34
35. Redox Reaction
In the chemical reaction, oxidation reaction and reduction reaction always occur
together, it is called oxidation reduction reaction abreviated as redox reaction.
In the redox reaction occurs transfering of electrons from the substance that
undergo oxidation to the substance that undergo reduction. Therefore, redox
reaction is also called reaction of transfering electrons
Special charateristic redox reaxtion is the oxidation number change.
Oxidation : lossing electron, increasing oxidation number.
Reduction : gaining electron, decreasing oxidation number.
The chemical reaction that does not espoused oxidation number change
(increasing or decreasing in oxidation number) called non-redox reaction.
36. (red)
o. n. of Cu decreases from
+2 to 0
Changing in o.n. of Cu is –
2
Example:
1. Redox reaction
(ox)
o. n. of H increases from 0
to +1
Total changing in o.n. of H
is +2
Reaction of copper(II) oxide with hydrogen gas to form copper and water vapor
In the redox reaction:
total number of increasing in oxidation number in oxidation
reaction = total number of decreasing in oxidation number in
reduction reaction.
CuO(s) + H2(g) Cu(s) + H2O(g) (redox)
+2 0 0 +1
37. Example problem :
Given a redox reaction:
3S(s)
+ 2KClO3(s)
3SO2(g)
+ 2KCl(s)
a. Identify and under line, element atoms of reactants undergo change
in oxidation number.
b. Determine the reactants that undergo reduction - oxidation
include their product, and calculate its oxidation number change
c. Determine the reactant behaves as oxidant and reductant.
Answer:
a. In the redox reaction:
3 S(s)
+ 2 KClO3(s)
3 SO2(g)
+ 2 KCl(s)
0 (+5) (+4) (-1)
Element atoms undergo change in oxidation number is:
- S : oxidation number of S increases from 0 to +4
- Cl : oxidation number of Cl element atom in KClO3
decreases
from +5 to -1
38. b. In the redox reaction:
3 S(s)
+ 2 KClO3(s)
3 SO2(g)
+ 2 KCl(s)
0 (+5) (+4) (-1)
(0) (+12)
(+10) (-2)
The total increasing o.n. of S (three atoms) is +12
The total decreasing o.n. of Cl (two atoms) is -12
S is oxidized into SO2
KClO3
is reduced into KCl
c. In the redox reaction:
3 S(s)
+ 2 KClO3(s)
3 SO2(g)
+ 2 KCl(s)
0 (+5) (+4) (-1)
S undergoes oxidation
KClO3
undergoes reduction
The element of S is reducing agent
The compound of KClO3
is oxidizing agent
(Ox)
(Red)
39. Types of redox reactions
• There are 5 types of redox reactions :-
Combination reactions
Decomposition reaction
Displacement reactions
Double displacement reactions
Disproportion reactions
39
40. Non-redox reactions
• The oxidation states of the elements
remained unchanged in the following
reactions:
• Neutralisation reactions:
O
H
NaCl
HCl
NaOH 2
+
→
+
O
H
CuSO
SO
H
CuO 2
4
4
2 +
→
+
41. Non-redox reactions
• The oxidation states of the elements
remained unchanged in the following
reactions:
• Precipitation reactions:
)
(
4
2
)
(
2
)
(
)
(
4 )
(
2 aq
s
aq
aq SO
Na
OH
Cu
NaOH
CuSO +
→
+
)
(
3
)
(
2
)
(
2
3
)
( 2
)
(
2 aq
s
aq
aq KNO
PbI
NO
Pb
KI +
→
+
42. Non-redox reactions
• The oxidation states of the elements
remained unchanged in the following
reactions:
• Complex formation:
( )
[ ] +
+
→
+
2
)
(
4
3
)
(
3
)
(
2
4 aq
aq
aq NH
Cu
NH
Cu
Tetraammine
copper(II) complex
(deep blue solution)
ligand
43. Disproportionation reactions
• These are a special type of reactions where an element in one
oxidation state is simultaneously oxidised and reduced.
• One of the reacting substances in a disproportion reaction always
contains an element that can exist in at least 3 oxidation states.
• The element in the form of reacting substance is in the
intermediate oxidation state.
43
44. 44
• Hypochlorite ion formed in a disproportion reaction
oxidises the colour bearing stains of the substances
to colourless compounds.
• Fluorine is the most electronegative element and
hence it cannot exhibit any positive oxidation state.
• Fluorine does not show a disproportion tendency.
45. Auto Redox Reaction (Disproportionation)
Auto redox reaction is a reaction of reduction and oxidation that occur in the
same substance (reactant).
Example of auto redox reaction:
Reaction of chlorine gas with sodium hydroxide solution
Cl2(g)
+ 2 NaOH(aq)
Na Cl(aq)
+ Na Cl O(aq)
+ H2
O(l)
0 –1 +1
(reduction)
o. n. of Cl decreases from 0 into –1
(oxidation)
o. n. of Cl increases from 0 into + 1
46. Disproportionation Reaction
• Example:
Is this a disproportionation reaction?
This is NOT a disproportionation reaction
• Disproportionation requires that the same
atom is both oxidised and reduced
simultaneously.
• In this case, different atoms (of nitrogen) are
oxidised and reduced.
-3 +5 +1
O
H
O
N
NO
NH 2
2
3
4 2
+
→
47. A Special Redox reaction:
Disproportionation
• Example
• Is it possible
• Chlorine is simultaneously reduced from oxidation
state of 0 in Cl2 to -1 in Cl-
, and oxidised from oxidation
state of 0 in Cl2 to +1 in ClO-
.
O
H
Cl
ClO
OH
Cl 2
2 2 +
+
→
+ −
−
−
0
+1 -1
48. Balancing Redox Reactions thorugh
half reaction method
There are several basic steps
1. Assign oxidation numbers to the species in the
reaction
2. Find the substance oxidized and the substance
reduced
3. Write half reactions for the oxidation and reduction
4. Balance the atoms that change in the half reaction
5. Determine the electrons transferred and balance the
electrons between the half reactions
6. Combine the half reactions and balance the remaining
atoms
7. Check your work. Make sure that both the atoms and
charges balance
49. 49
Identify the Oxidizing and Reducing
Agents in Each of the Following
3 H2S + 2 NO3
–
+ 2 H+
→ 3 S + 2 NO + 4 H2O
MnO2 + 4 HBr → MnBr2 + Br2 + 2 H2O
50. 50
Identify the Oxidizing and Reducing
Agents in Each of the Following
3 H2S + 2 NO3
–
+ 2 H+
→ 3 S + 2 NO + 4 H2O
MnO2 + 4 HBr → MnBr2 + Br2 + 2 H2O
+1 -2 +5 -2 +1 0 +2 -2 +1 -2
ox ag
red ag
+4 -2 +1 -1 +2 -1 0 +1 -2
oxidation
reduction
oxidation
reduction
red ag
ox ag
51. 51
Balancing Redox Reactions
1) assign oxidation states and
determine element oxidized
and element reduced
2) separate into oxidation &
reduction half-reactions
3) balance half-reactions by mass
a) first balance atoms other than O and
H
b) then balance O by adding H2O to
side that lacks O
c) finally balance H by adding H+
to
side that lacks H
Fe2+
+ MnO4
–
→ Fe3+
+ Mn2+
+2 +7 -2 +3 +2
oxid
red
Fe2+
→ Fe3+
MnO4
–
→ Mn2+
Fe2+
→ Fe3+
MnO4
–
→ Mn2+
MnO4
–
→ Mn2+
+ 4H2O
MnO4
–
+ 8H+
→ Mn2+
+ 4H2O
52. 52
Balancing Redox Reactions
4) balance each half-reaction
with respect to charge by
adjusting the numbers of
electrons
a) electrons on product side for oxid.
b) electrons on reactant side for red.
4) balance electrons between
half-reactions
5) add half-reactions, canceling
electrons and common
species
6) Check
Fe2+
→ Fe3+
+ 1 e-
MnO4
–
+ 8H+
+ 5 e-
→ Mn2+
+ 4H2O
MnO4
–
+ 8H+
→ Mn2+
+ 4H2O
+7 +2
Fe2+
→ Fe3+
+ 1 e-
MnO4
–
+ 8H+
+ 5 e-
→ Mn2+
+ 4H2O
} x 5
5 Fe2+
→ 5 Fe3+
+ 5 e-
MnO4
–
+ 8H+
+ 5 e-
→ Mn2+
+ 4H2O
5 Fe2+
+ MnO4
–
+ 8H+
→ Mn2+
+ 4H2O + 5 Fe3+
56. Redox reactions as the basis for
titrations
• In one situation, the reagent itself is intensely coloured. Here in
this case the permanganate ion it acts as a self indicatorhere the
end point is reached after the vlast of the reductant is oxidised
and the first lasting tinge of pink colour appears at low
concentration.
• This ensures minimal overshoot in colour beyond the equivalence
point , the point where the reductant and the oxidant are equal in
terms of their mole stoichiometry.
• if there is no dramatic auto-colour change there are indicators
which are oxidised immediately after the last bit of the reactant
is consumed, producing a dramatic colour change. The best
example is afforded by
• Dichromate salt , which is not a self indicator, but oxidises the
indicator substance diphenylamine just after the equivalence
point to produce an intense blue colour, thus signaling the end
point.
• There is yet another method which is interesting and quite 56
57. • For example,
This method relies on the facts iodine gives an intense blue
colour starch and has very specific reaction with thiosuphate
ion which is too a redox reaction.
Iodide remains a solution containing KI or KI3
On addition of starch after the liberation of iodide from the
reaction of Cu 2+ ions on iodide ions , an intense blue colour
appears , this colour disappears as soon as iodine is consumed
by thiosuphate ions . Thus the end point can be tracked easily
by stoichiometric calculations
57
59. 59
Mg is above
Cu on the
Activity Series
Mg will react with
Cu2+
to form Mg2+
and Cu metal
but Cu will not
react with Mg2+
60. 60
Electrochemical Cells
• electrochemistry is the study of redox reactions that
produce or require an electric current
• the conversion between chemical energy and electrical
energy is carried out in an electrochemical cell
• spontaneous redox reactions take place in a voltaic cell
also known as galvanic cells
batteries are voltaic cells
• nonspontaneous redox reactions can be made to occur
in an electrolytic cell by the addition of electrical
energy
61. 61
Electrodes
• Anode
electrode where oxidation occurs
anions attracted to it
connected to positive end of battery in electrolytic
cell
loses weight in electrolytic cell
• Cathode
electrode where reduction occurs
cations attracted to it
connected to negative end of battery in electrolytic
cell
gains weight in electrolytic cell
electrode where plating takes place in electroplating
63. Redox reactions and electrode process
• Redox couple:-
It is defined as having together the oxidised and
reduced forms of a substance taking part in an
oxidation or reduction half reaction
This is represented by separating the oxidised
form from the reduced form by a vertical line
showing for e.g. solid/solution interface.
63
64. • Example:- Zn2+
/Zn ,Cu2+
/Cu.
In both cases oxidised form is put before the reduced
form.
Experiment-daniell’s cell
Now we put the beaker containing copper sulphate solution and the
beaker containing zinc sulphate solution side by side . We connect
solutions in two beakers by a salt bridge (a U-tube containing a
solution of potassium chloride or ammonium nitrate usually
solidified by boiling with agar agar and later cooling to a jelly like
substance.)
This provides an electric contact between the two solutions without
allowing them to mix with each other. The zinc and copper rods
are connected by a metallic wire with a provision for an ammeter
and a switch. The set-up is known as Daniell cell. When the switch
is in the o position, no reaction takes place in either of the
beakers and no current flows through the metallic wire.
64
65. • As soon as the switch is on we get the
following observations:-
1. The transfer of electrons now does not take place directly
from Zn to Cu2+
but through the metallic wire connecting
the two rods as is apparent from the arrow which indicates
the flow of current.
2. The electricity from solution in one beaker to solution in
the other beaker flows by the migration of ions through the
salt bridge. We know that the flow of current is possible
only if there is a potential difference between the copper
and zinc rods known as electrodes here.
65
67. Types of Electrochemical Cells
Voltaic (or galvanic) cell: uses a spontaneous reaction (∆G < 0)
to generate electrical energy.
Electrolytic cell: uses electrical energy to drive a non-spontaneous
reaction (∆G > 0).
Contain two electrodes (anode and cathode) dipped into an
aqueous electrolyte solution.
The oxidation half-reaction occurs at the anode; the reduction
half-reaction occurs at the cathode.