IGCSE 11.1.10
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Most chemicals will decompose if heated strongly enough, requiring energy to break chemical bonds in an endothermic process. Bond-breaking being endothermic means that bond-making is exothermic, releasing energy when chemical bonds are formed.
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IGCSE 11.1.13
Activation energy (Ea) is the minimum amount of energy needed to start a chemical reaction by breaking bonds between reactants. Most reactions require existing bonds to break through an endothermic process before new bonds can form exothermically. Different reactions require different amounts of activation energy depending on how easily the bonds in the reactants can be broken.
Potential energy diagram worksheet 1
This document provides instructions for using a potential energy diagram to analyze a chemical reaction. It asks the reader to label different components on the diagram including reactants, products, and activated complex. It also asks the reader to determine values like activation energy and heat of reaction from the diagram and define key concepts such as heat of reaction. Finally, it asks the reader to compare forward and reverse activation energies for an exothermic reaction using the diagram.
IGCSE 11.1.1
Exothermic reactions release energy as heat to the surroundings and can occur spontaneously or explosively. Examples include combustion, respiration, and reactions of metals with acids. Endothermic reactions absorb energy from the surroundings and cause cooling. Reversible reactions are exothermic in one direction as they release heat, but are endothermic in the reverse direction as they absorb the same amount of heat.
exothermic and endothermic reactions ppt
Chemical reactions can be either exothermic or endothermic. Exothermic reactions release energy in the form of heat, causing a temperature increase. Endothermic reactions absorb energy from their surroundings in the form of heat, causing a temperature decrease. Examples of exothermic reactions given are the neutralization of an acid with a base and burning magnesium, while examples of endothermic reactions include the thermal decomposition of limestone and dissolving salts in water. The energy changes in chemical reactions can be explained by considering the energy required to break bonds in reactants versus the energy released when new bonds form in products.
G-10 IG-CHEMICAL ENERGETICS 1 ..pptx
This document discusses chemical energetics and exothermic and endothermic reactions. It defines key terms like exothermic, endothermic, potential energy, and activation energy. Examples of exothermic reactions given are combustion reactions and neutralization which release heat. Burning fuels like methane and hydrogen are also exothermic. Endothermic reactions absorb heat from their surroundings. The document explains that bond breaking requires energy and is endothermic, while bond formation releases energy and is exothermic. Whether a reaction is exothermic or endothermic depends on the strength of bonds broken versus those formed. It concludes with questions for exercises.
Endo and Exo.pptx
Chemical reactions can be classified as either exothermic or endothermic based on whether energy is released or absorbed during the reaction. In an exothermic reaction, more energy is released from bond formation in the products than is required to break bonds in the reactants, resulting in a net release of energy. In an endothermic reaction, more energy is required to break bonds in the reactants than is released in bond formation in the products, resulting in a net absorption of energy. Common examples of exothermic reactions include combustion and yeast reactions, while photosynthesis and dissolving acid in water are typically endothermic.
7.1 Endothermic & Exothermic Reactions - TE.ppt
This document defines and provides examples of endothermic and exothermic chemical reactions. Endothermic reactions absorb energy from their surroundings, such as the thermal decomposition of limestone. Exothermic reactions release energy to their surroundings, like burning magnesium. The document explains that bonds contain energy, and energy is absorbed when bonds break in endothermic reactions, while energy is released when new bonds form in exothermic reactions. Examples of both endothermic and exothermic reactions are given.
Recommended
IGCSE 11.1.13
Activation energy (Ea) is the minimum amount of energy needed to start a chemical reaction by breaking bonds between reactants. Most reactions require existing bonds to break through an endothermic process before new bonds can form exothermically. Different reactions require different amounts of activation energy depending on how easily the bonds in the reactants can be broken.
Potential energy diagram worksheet 1
This document provides instructions for using a potential energy diagram to analyze a chemical reaction. It asks the reader to label different components on the diagram including reactants, products, and activated complex. It also asks the reader to determine values like activation energy and heat of reaction from the diagram and define key concepts such as heat of reaction. Finally, it asks the reader to compare forward and reverse activation energies for an exothermic reaction using the diagram.
IGCSE 11.1.1
Exothermic reactions release energy as heat to the surroundings and can occur spontaneously or explosively. Examples include combustion, respiration, and reactions of metals with acids. Endothermic reactions absorb energy from the surroundings and cause cooling. Reversible reactions are exothermic in one direction as they release heat, but are endothermic in the reverse direction as they absorb the same amount of heat.
exothermic and endothermic reactions ppt
Chemical reactions can be either exothermic or endothermic. Exothermic reactions release energy in the form of heat, causing a temperature increase. Endothermic reactions absorb energy from their surroundings in the form of heat, causing a temperature decrease. Examples of exothermic reactions given are the neutralization of an acid with a base and burning magnesium, while examples of endothermic reactions include the thermal decomposition of limestone and dissolving salts in water. The energy changes in chemical reactions can be explained by considering the energy required to break bonds in reactants versus the energy released when new bonds form in products.
G-10 IG-CHEMICAL ENERGETICS 1 ..pptx
This document discusses chemical energetics and exothermic and endothermic reactions. It defines key terms like exothermic, endothermic, potential energy, and activation energy. Examples of exothermic reactions given are combustion reactions and neutralization which release heat. Burning fuels like methane and hydrogen are also exothermic. Endothermic reactions absorb heat from their surroundings. The document explains that bond breaking requires energy and is endothermic, while bond formation releases energy and is exothermic. Whether a reaction is exothermic or endothermic depends on the strength of bonds broken versus those formed. It concludes with questions for exercises.
Endo and Exo.pptx
Chemical reactions can be classified as either exothermic or endothermic based on whether energy is released or absorbed during the reaction. In an exothermic reaction, more energy is released from bond formation in the products than is required to break bonds in the reactants, resulting in a net release of energy. In an endothermic reaction, more energy is required to break bonds in the reactants than is released in bond formation in the products, resulting in a net absorption of energy. Common examples of exothermic reactions include combustion and yeast reactions, while photosynthesis and dissolving acid in water are typically endothermic.
7.1 Endothermic & Exothermic Reactions - TE.ppt
This document defines and provides examples of endothermic and exothermic chemical reactions. Endothermic reactions absorb energy from their surroundings, such as the thermal decomposition of limestone. Exothermic reactions release energy to their surroundings, like burning magnesium. The document explains that bonds contain energy, and energy is absorbed when bonds break in endothermic reactions, while energy is released when new bonds form in exothermic reactions. Examples of both endothermic and exothermic reactions are given.
5. using electromagnets
Lifts, cars, and other large electrical machines use high currents. Relays use an electromagnet to allow a small current in one circuit to control a large current in another circuit. This allows a small current to control whether a large current flows through another circuit by closing or opening a switch.
2
The circuit for a door bell contains an electromagnet that pulls an armature towards it when the circuit is closed, causing a hammer to strike the bell and produce sound. The movement of the armature then breaks the circuit, allowing the hammer to return to its original position, and this sequence repeats to make the bell ring continuously whenever the door bell button is pressed.
1
An electromagnet on a recycling plant conveyor belt is used to pick up and move metal cans. Electromagnets attract ferromagnetic metals like iron, nickel, and cobalt. Electromagnets have the advantage over permanent magnets in that their magnetic field can be switched on and off by controlling the electric current running through the magnet coils.
5
The document presents results from an experiment investigating electromagnets. It shows that increasing the number of coils or the current in an electromagnet increases the number of drawing pins attracted. Graphs illustrate the direct relationship between the number of coils/current and the number of drawing pins attracted.
4
An experiment was designed to investigate how changing the current and number of coils affects the strength of an electromagnet. The apparatus shown could be used by altering the current running through the coils or changing the number of coils wrapped around the iron core, then measuring the effect on the magnetic force produced. Results could show that increasing either the current or number of coils strengthens the electromagnet.
3
The document discusses two experiments that can investigate factors affecting the strength of an electromagnet other than the presence of an iron core. The first experiment would examine how the number of coils influences the number of drawing pins attracted while keeping current the same. The second would look at how the size of the current affects the number of pins for a constant number of coils.
6. halides
Halides are compounds formed between halogens (F, Cl, Br, I) and metals. They have many industrial, medical, and household applications. Halides can be identified by their reaction with silver nitrate solution to form precipitates of insoluble silver halides. The hydrogen halides are colorless gases with hydrogen fluoride having an unexpectedly high boiling point due to hydrogen bonding between H-F molecules. Larger halide ions are more reactive reducing agents as their outer electrons are farther from the nucleus and more easily donated.
4. g7 and g1
Alkali metals react vigorously with chlorine to form metal chlorides. Lithium reacts with chlorine gas to form lithium chloride according to the chemical equation 2Li (s) + Cl2 (g) → 2LiCl (s). Sodium also reacts with chlorine gas in a similar reaction to form sodium chloride, represented by the equation 2Na (s) + Cl2 (g) → 2NaCl (s).
3. group 7
The document discusses the chemical properties of the halogens (fluorine, chlorine, bromine, iodine, astatine) by examining their reactivity trends down the group in reactions with iron wool and hydrogen. It notes that reactivity decreases down the group, with chlorine causing iron wool to burn brightly, bromine less so, and iodine only slightly. It also examines the halogens' reactions with metals like sodium and nonmetals like hydrogen to form halides, and displacement reactions between halogens and halide ions.
2. group 1
The document discusses the chemical properties of alkali metals. It explains that alkali metals react vigorously with oxygen and water. The reactivity increases down the group as the atoms get larger, shielding the outer electrons from the nucleus and making them easier to lose. Equations for reactions of lithium, sodium, and potassium with oxygen, water, and other substances are provided. Flame tests for group 2 metals are also discussed.
5.displacement
The document discusses the chemical properties of halogens, specifically their oxidizing abilities in displacement reactions. It states that in displacement reactions between halogens and halides, the halogen acts as an oxidizing agent by oxidizing the halide ion and being reduced to form the halide ion itself. The order of increasing oxidizing ability among the halogens is F > Cl > Br > I. An example reaction of chlorine displacing bromine is provided, and it is explained that in the reaction, chlorine is reduced while bromide is oxidized.
1. group 0
The document discusses the properties and discovery of noble gases. It notes that:
- Noble gases (group 0) are very unreactive due to their full outer electron shells. This stability makes them similar across the group.
- Helium was first discovered on the sun in 1868. Argon was discovered in the 1890s after nitrogen from air was found to contain a small quantity of an inert element.
- Ramsay realized argon needed its own group and later discovered it contained traces of other noble gases - neon, krypton, and xenon. Radon was discovered in 1900.
4
The document discusses how changing the pressure affects the equilibrium of reversible gas reactions. It states that if the pressure is increased, the equilibrium will shift in the direction that decreases the number of gas molecules. If the pressure is decreased, the equilibrium will shift in the direction that increases the number of gas molecules. It also provides the example of nitrogen dioxide and dinitrogen tetroxide, explaining that if the pressure of this system is increased, more dinitrogen tetroxide will be produced.
3
Increasing the concentration of a substance in a reversible reaction shifts the equilibrium to decrease the concentration of that substance. For example, adding more water to a reaction producing bismuth oxychloride and hydrochloric acid from bismuth chloride and water would shift the equilibrium right, producing more products. Similarly, adding more chlorine gas to a reaction producing iodine trichloride from chlorine and iodine chloride would make the mixture appear more yellow by shifting the equilibrium to the right.
2
The document discusses exothermic and endothermic reactions. It states that all reactions are exothermic in one direction, releasing heat, and endothermic in the other direction, absorbing heat. It also explains that if temperature is increased, the equilibrium will shift in the endothermic direction to decrease temperature, and if temperature is decreased, the equilibrium will shift in the exothermic direction to increase temperature. It provides the example of nitrogen dioxide and dinitrogen tetroxide, where the forward reaction is exothermic and the backward reaction is endothermic. If temperature is increased in this system, more nitrogen dioxide will be produced.
5
Catalysts speed up chemical reactions by lowering the activation energy of the reaction pathway, without being used up in the process. When added to a reversible reaction, a catalyst increases the rates of both the forward and reverse reactions equally, resulting in no change to the equilibrium position but equilibrium being reached faster. Catalysts are particularly important in industrial applications as they allow reactions to proceed more quickly.
6
Ammonia is produced industrially through the Haber process, which involves reacting nitrogen and hydrogen at high pressures and temperatures. However, the reaction is reversible and does not go to completion. To maximize yield while keeping costs low, manufacturers use a compromise of 450°C and 200 atmospheres, as extremely high pressures require expensive equipment and low temperatures slow the reaction rate. Key costs in ammonia production include raw materials, energy, equipment, wages. Yield is maximized by using a catalyst, removing ammonia as it forms, and recycling unreacted gases.
1
Le Chatelier's principle states that if a system in dynamic equilibrium experiences a change, the equilibrium will shift to counteract the change. Specifically, increasing temperature shifts equilibrium towards the endothermic direction, increasing concentration shifts equilibrium away from that substance, and increasing pressure shifts equilibrium away from the gas producing side.
IB 3.2.2.5
The melting points of elements in group 1 decrease down the group because the atomic radius increases, resulting in weaker metallic bonds and a weaker attraction between the nucleus and delocalized electrons. The melting points given for lithium, sodium, potassium, rubidium, and cesium show this decreasing trend.
IB 3.2.2.4
Electronegativity decreases down a group because the atomic radius increases and shielding increases due to electrons filling new principal energy levels, making the increased nuclear charge less significant. Therefore, iodine has a larger atomic radius and attracts electron density in a covalent bond less strongly than fluorine.
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5. using electromagnets
Lifts, cars, and other large electrical machines use high currents. Relays use an electromagnet to allow a small current in one circuit to control a large current in another circuit. This allows a small current to control whether a large current flows through another circuit by closing or opening a switch.
2
The circuit for a door bell contains an electromagnet that pulls an armature towards it when the circuit is closed, causing a hammer to strike the bell and produce sound. The movement of the armature then breaks the circuit, allowing the hammer to return to its original position, and this sequence repeats to make the bell ring continuously whenever the door bell button is pressed.
1
An electromagnet on a recycling plant conveyor belt is used to pick up and move metal cans. Electromagnets attract ferromagnetic metals like iron, nickel, and cobalt. Electromagnets have the advantage over permanent magnets in that their magnetic field can be switched on and off by controlling the electric current running through the magnet coils.
5
The document presents results from an experiment investigating electromagnets. It shows that increasing the number of coils or the current in an electromagnet increases the number of drawing pins attracted. Graphs illustrate the direct relationship between the number of coils/current and the number of drawing pins attracted.
4
An experiment was designed to investigate how changing the current and number of coils affects the strength of an electromagnet. The apparatus shown could be used by altering the current running through the coils or changing the number of coils wrapped around the iron core, then measuring the effect on the magnetic force produced. Results could show that increasing either the current or number of coils strengthens the electromagnet.
3
The document discusses two experiments that can investigate factors affecting the strength of an electromagnet other than the presence of an iron core. The first experiment would examine how the number of coils influences the number of drawing pins attracted while keeping current the same. The second would look at how the size of the current affects the number of pins for a constant number of coils.
6. halides
Halides are compounds formed between halogens (F, Cl, Br, I) and metals. They have many industrial, medical, and household applications. Halides can be identified by their reaction with silver nitrate solution to form precipitates of insoluble silver halides. The hydrogen halides are colorless gases with hydrogen fluoride having an unexpectedly high boiling point due to hydrogen bonding between H-F molecules. Larger halide ions are more reactive reducing agents as their outer electrons are farther from the nucleus and more easily donated.
4. g7 and g1
Alkali metals react vigorously with chlorine to form metal chlorides. Lithium reacts with chlorine gas to form lithium chloride according to the chemical equation 2Li (s) + Cl2 (g) → 2LiCl (s). Sodium also reacts with chlorine gas in a similar reaction to form sodium chloride, represented by the equation 2Na (s) + Cl2 (g) → 2NaCl (s).
3. group 7
The document discusses the chemical properties of the halogens (fluorine, chlorine, bromine, iodine, astatine) by examining their reactivity trends down the group in reactions with iron wool and hydrogen. It notes that reactivity decreases down the group, with chlorine causing iron wool to burn brightly, bromine less so, and iodine only slightly. It also examines the halogens' reactions with metals like sodium and nonmetals like hydrogen to form halides, and displacement reactions between halogens and halide ions.
2. group 1
The document discusses the chemical properties of alkali metals. It explains that alkali metals react vigorously with oxygen and water. The reactivity increases down the group as the atoms get larger, shielding the outer electrons from the nucleus and making them easier to lose. Equations for reactions of lithium, sodium, and potassium with oxygen, water, and other substances are provided. Flame tests for group 2 metals are also discussed.
5.displacement
The document discusses the chemical properties of halogens, specifically their oxidizing abilities in displacement reactions. It states that in displacement reactions between halogens and halides, the halogen acts as an oxidizing agent by oxidizing the halide ion and being reduced to form the halide ion itself. The order of increasing oxidizing ability among the halogens is F > Cl > Br > I. An example reaction of chlorine displacing bromine is provided, and it is explained that in the reaction, chlorine is reduced while bromide is oxidized.
1. group 0
The document discusses the properties and discovery of noble gases. It notes that:
- Noble gases (group 0) are very unreactive due to their full outer electron shells. This stability makes them similar across the group.
- Helium was first discovered on the sun in 1868. Argon was discovered in the 1890s after nitrogen from air was found to contain a small quantity of an inert element.
- Ramsay realized argon needed its own group and later discovered it contained traces of other noble gases - neon, krypton, and xenon. Radon was discovered in 1900.
4
The document discusses how changing the pressure affects the equilibrium of reversible gas reactions. It states that if the pressure is increased, the equilibrium will shift in the direction that decreases the number of gas molecules. If the pressure is decreased, the equilibrium will shift in the direction that increases the number of gas molecules. It also provides the example of nitrogen dioxide and dinitrogen tetroxide, explaining that if the pressure of this system is increased, more dinitrogen tetroxide will be produced.
3
Increasing the concentration of a substance in a reversible reaction shifts the equilibrium to decrease the concentration of that substance. For example, adding more water to a reaction producing bismuth oxychloride and hydrochloric acid from bismuth chloride and water would shift the equilibrium right, producing more products. Similarly, adding more chlorine gas to a reaction producing iodine trichloride from chlorine and iodine chloride would make the mixture appear more yellow by shifting the equilibrium to the right.
2
The document discusses exothermic and endothermic reactions. It states that all reactions are exothermic in one direction, releasing heat, and endothermic in the other direction, absorbing heat. It also explains that if temperature is increased, the equilibrium will shift in the endothermic direction to decrease temperature, and if temperature is decreased, the equilibrium will shift in the exothermic direction to increase temperature. It provides the example of nitrogen dioxide and dinitrogen tetroxide, where the forward reaction is exothermic and the backward reaction is endothermic. If temperature is increased in this system, more nitrogen dioxide will be produced.
5
Catalysts speed up chemical reactions by lowering the activation energy of the reaction pathway, without being used up in the process. When added to a reversible reaction, a catalyst increases the rates of both the forward and reverse reactions equally, resulting in no change to the equilibrium position but equilibrium being reached faster. Catalysts are particularly important in industrial applications as they allow reactions to proceed more quickly.
6
Ammonia is produced industrially through the Haber process, which involves reacting nitrogen and hydrogen at high pressures and temperatures. However, the reaction is reversible and does not go to completion. To maximize yield while keeping costs low, manufacturers use a compromise of 450°C and 200 atmospheres, as extremely high pressures require expensive equipment and low temperatures slow the reaction rate. Key costs in ammonia production include raw materials, energy, equipment, wages. Yield is maximized by using a catalyst, removing ammonia as it forms, and recycling unreacted gases.
1
Le Chatelier's principle states that if a system in dynamic equilibrium experiences a change, the equilibrium will shift to counteract the change. Specifically, increasing temperature shifts equilibrium towards the endothermic direction, increasing concentration shifts equilibrium away from that substance, and increasing pressure shifts equilibrium away from the gas producing side.
IB 3.2.2.5
The melting points of elements in group 1 decrease down the group because the atomic radius increases, resulting in weaker metallic bonds and a weaker attraction between the nucleus and delocalized electrons. The melting points given for lithium, sodium, potassium, rubidium, and cesium show this decreasing trend.
IB 3.2.2.4
Electronegativity decreases down a group because the atomic radius increases and shielding increases due to electrons filling new principal energy levels, making the increased nuclear charge less significant. Therefore, iodine has a larger atomic radius and attracts electron density in a covalent bond less strongly than fluorine.
More from shaunoff (20)
IGCSE 11.1.10
- 1. Making and breaking chemical bonds Most chemicals will break up (decompose) if they are heated strongly enough. This means that energy is needed to break chemical bonds – an endothermic process. 1 of 35 © Boardworks Ltd 2007
- 2. Making and breaking chemical bonds Most chemicals will break up (decompose) if they are heated strongly enough. This means that energy is needed to break chemical bonds – an endothermic process. 2 of 35 © Boardworks Ltd 2007
- 3. Making and breaking chemical bonds Most chemicals will break up (decompose) if they are heated strongly enough. This means that energy is needed to break chemical bonds – an endothermic process. energy absorbed 3 of 35 © Boardworks Ltd 2007
- 4. Making and breaking chemical bonds Most chemicals will break up (decompose) if they are heated strongly enough. This means that energy is needed to break chemical bonds – an endothermic process. energy absorbed 4 of 35 © Boardworks Ltd 2007
- 5. Making and breaking chemical bonds Most chemicals will break up (decompose) if they are heated strongly enough. This means that energy is needed to break chemical bonds – an endothermic process. energy absorbed Because bond-breaking is endothermic, bond-making must therefore be exothermic. This means that energy is released when chemical bonds are made. 5 of 35 © Boardworks Ltd 2007
- 6. Making and breaking chemical bonds Most chemicals will break up (decompose) if they are heated strongly enough. This means that energy is needed to break chemical bonds – an endothermic process. energy absorbed Because bond-breaking is endothermic, bond-making must therefore be exothermic. This means that energy is released when chemical bonds are made. 6 of 35 © Boardworks Ltd 2007
- 7. Making and breaking chemical bonds Most chemicals will break up (decompose) if they are heated strongly enough. This means that energy is needed to break chemical bonds – an endothermic process. energy absorbed Because bond-breaking is endothermic, bond-making must therefore be exothermic. This means that energy is released when chemical bonds are made. energy released 7 of 35 © Boardworks Ltd 2007
- 8. Making and breaking chemical bonds Most chemicals will break up (decompose) if they are heated strongly enough. This means that energy is needed to break chemical bonds – an endothermic process. energy absorbed Because bond-breaking is endothermic, bond-making must therefore be exothermic. This means that energy is released when chemical bonds are made. energy released 8 of 35 © Boardworks Ltd 2007