visual presentation of a chemical reaction

Posted: 27 February, 2024

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Inspire and energise your classes with these much-loved demonstrations.

These resources have been taken from a vast collection of practical demonstrations for stimulating and motivating students.

If you'd like to buy a copy of Classic chemistry demonstrations visit our online bookshop . If you're a member of the Royal Society of Chemistry, you're entitled to a 35% discount.

‘Dissolving’ polystyrene in acetone

In association with Nuffield Foundation

Investigate what happens to polystyrene when it is placed in propanone (acetone) in this demonstration. Includes kit list and safety instructions.

‘Magic’ writing with colour changing reactions

Reveal invisible messages or pictures drawn with aqueous solutions by spraying them with suitable reagents in this demonstration. Includes kit list and safety instructions.

A controlled explosion using hydrogen and air

Show how a hydrogen–air mixture can gain explosive properties using a plastic drink bottle in this demonstration. Includes kit list and safety instructions.

A hydrogen powered rocket

Try this spectacular demonstration to make a rocket using a plastic drink bottle fuelled by hydrogen and air. Includes kit list and safety instructions.

A simple oscillating reaction

Use this demonstration to illustrate an oscillating reaction as bromate ions oxidise malonic acid to carbon dioxide. Includes kit list and safety instructions.

A solid–solid reaction between lead nitrate and potassium iodide

Use this demonstration with kit list and safety instructions to prove that two solids can react together, making lead iodide from lead nitrate and potassium iodide.

A visible activated complex

A simple demonstration of catalysis also introducing the idea of an activated complex and to allow discussion of the mechanism of catalysis. Includes kit list and safety instructions.

Allotropes of sulfur

Use this practical to explore the changes in the colour and consistency of sulfur as you heat it, melt it and eventually boil it. Includes kit list and safety instructions.

Ammonia fountain demonstration

Try this experiment to make a miniature chemical fountain using only soluble ammonia and atmospheric pressure. Includes kit list and safety instructions.

Ammonium dichromate volcano

Try this demonstration to create a mini volcanic eruption illustrating the decomposition of ammonium dichromate. Includes kit list and safety instructions.

Anodising aluminium

Explore an application of electrolysis in this demonstration by anodising aluminium to improve corrosion resistance. Includes kit list and safety instructions.

Bright sparks

The spontaneous combustion of high surface area iron

Burning milk powder

Gather a Bunsen burner, and some common powdered milk to help students grasp the ideas of surface area and reaction rates. Includes kit list and safety instructions. 

Cannon fire

 Increase the rate of burning with the inclusion of oxygen, in this loud exothermic practical

Catalysing the reaction of sodium thiosulfate and hydrogen peroxide

Illustrate the effect of a catalyst as sodium thiosulfate is oxidised by hydrogen peroxide in this demonstration. Includes kit list and safety instructions.

Catalysts for the thermal decomposition of potassium chlorate

Try this demonstration to investigate the effectiveness of various catalysts for the decomposition of potassium chlorate. Includes kit list and safety instructions.

Chemiluminescence of luminol: a cold light experiment

Show how the energy of a chemical reaction can be given out as light. Includes kit list and safety instructions. 

Colourful reactions using ammonia solution

Turn ammonia solution red, white or blue by adding phenolphthalein, lead nitrate or copper(II) sulfate in this demonstration. Includes kit list and safety instructions.

Combustion of ethanol

Illustrate the large energy changes that take place during the combustion of alcohols with this spectacular demonstration. Includes kit list and safety instructions.

Cracking hydrocarbons in liquid paraffin with a catalyst

Model the industrial process of cracking larger hydrocarbons to produce smaller alkanes in this demonstration or class practical. Includes kit list and safety instructions.

Dehydration of ethanol to form ethene

Use this class practical or demonstration to produce ethene gas as an example of an unsaturated hydrocarbon. Includes kit list and safety instructions.

Determining relative molecular mass by weighing gases

Use this demonstration to determine the relative molecular masses of different gases using the ideal gas equation. Includes kit list and safety instructions.

Determining the relative molecular mass of butane

Use this demonstration to calculate the relative molecular mass of butane using simple apparatus. Includes kit list and safety instructions.

Diffusion of gases: ammonia and hydrogen chloride

A demonstration to show the diffusion of gases, using ammonia solution and hydrochloric acid. Includes kit list and safety instructions.

Dyeing three colours from the same dye bath

Show how dyeing involves chemical interactions between dyes and the molecular nature of different fibres in this demonstration. Includes kit list and safety instructions.

Electrolysis of molten lead(II) bromide

Introduce your students to the study of electrolysis through the production of metallic lead and bromine in this demonstration. Includes kit list and safety instructions.

Endothermic solid–solid reactions

Observe an endothermic reaction between two solids in this demonstration or class experiment. Includes kit list and safety instructions.

Equilibria involving carbon dioxide in aqueous solution

Use this demonstration or class practical to illustrate changes to equilibria in carbonated soda water. Includes kit list and safety instructions.

Estimating the concentration of bleach

Compare the chlorine content and concentration of sodium hypochlorite in different bleaches in this class practical. Includes kit list and safety instructions.

Exploding a tin can using methane

Use this demonstration to illustrate how methane can create an explosive mixture with the oxygen in air. Includes kit list and safety instructions.

Exploding bubbles of hydrogen and oxygen

Create a small explosion in this demonstration by electrolysing water to produce hydrogen and oxygen bubbles. Includes kit list, video and safety instructions.

Extracting iron from breakfast cereal

Try this class practical or demonstration to extract food-grade iron from breakfast cereals using neodymium magnets. Includes kit list and safety instructions.

Flame colours: a demonstration

Explore how different elements rect when exposed to a flame, and discuss how alkali metals, alkaline earth metals, and metal salts change the colour of fire.

Floating and sinking bubbles

Make bubbles of carbon dioxide, hydrogen or methane in this demonstration exploring density, diffusion and solubility. Includes kit list and safety instructions.

Halogen reactions with iron wool

Illustrate an exothermic redox reaction by heating iron wool with chlorine, bromine and iodine with this demonstration. Includes kit list and safety instructions.

Hydrogen peroxide decomposition using different catalysts

Collect a range of catalysts to explore the decomposition of hydrogen peroxide, paying close attention to the varied reaction rates. Includes kit list and safety instructions.

Identifying the products of combustion

Illustrate the presence of water and carbon dioxide in the products of hydrocarbon combustion in this demonstration. Includes kit list and safety instructions.

Indicators and dry ice demonstration

Create bubbles, ‘fog’ and a colour change adding dry ice to alkaline ammonia or sodium hydroxide solution in this demonstration. Includes kit list and safety instructions.

Iodine clock reaction demonstration method

Use this iodine clock reaction demonstration to introduce your students to rates of reaction and kinetics. Includes kit list and safety instructions.

Liquefying chlorine gas

Use this demonstration to produce liquid chlorine and compare it with bromine and iodine in their condensed state. Includes kit list and safety instructions.

Making nylon: the ‘nylon rope trick’

The ‘nylon rope trick’ is a classic of chemistry classrooms, by mixing decanedioyl dichloride and in cyclohexane you can create a solution that will form nylon strings when floated on an aqueous solution of 1,6-diaminohexane. Kit list and safety instructions included.

Making rayon

Use this demonstration to produce rayon fibres in the classroom using cotton wool or filter paper. Includes kit list and safety instructions.

Making silicon and silanes from sand

Create silicon in your classroom using just sand and magnesium. This exothermic practical will show learners the nuances of heat based reactions and how to perform them safely. Kit list and safety instructions included. 

Modelling the greenhouse effect

Use this demonstration to illustrate the greenhouse effect and the role of carbon dioxide as a greenhouse gas. Includes kit list and safety instructions.

Phenol-methanal polymerisiation

Make Bakelite in class and investigate its properties using phenol, methanal and ethanoic acid in this demonstration. Includes kit list and safety instructions.

Reacting zinc and copper(II) oxide

Illustrate competition reactions using the exothermic reaction between copper(II) oxide and zinc in this class demonstration. Includes kit list and safety instructions.

Reactions of chlorine, bromine and iodine with aluminium

Try this demonstration to produce some spectacular exothermic redox reactions by reacting aluminium with halogens. Includes kit list and safety instructions.

Reactivity trends of the alkali metals

Use this experiment to demonstrate the trend in reactivity down group 1 of the Periodic Table, exploring the physical and chemical properties of the alkali metals.

Rechargeable cells: the lead–acid accumulator

Use this practical to demonstrate the chemistry behind rechargeable batteries, using a lead–acid accumulator cell. Includes kit list and safety instructions.

Reduction of copper(II) oxide by hydrogen

Determine the formula of copper(II) oxide by reducing it using hydrogen or methane, in one of three methods available to you in this practical. Includes kit list and safety instructions. 

Sodium ethanoate ‘stalagmite’

Quickly grow your own ‘stalagmite’ from a supersaturated solution of sodium ethanoate in this demonstration. Includes kit list and safety instructions.

Sulfuric acid as a dehydrating agent

Try these two demonstrations to illustrate the difference between dehydration and drying using sulfuric acid. Includes kit list and safety instructions.

The ‘blue bottle’ experiment

In this demonstration, the redox indicator Methylene blue can be oxidised many times by shaking. Includes kit list and safety instructions.

The ‘breathalyser’ reaction | 16–18 years

In association with Nuffield Foundation , By Tim Jolliff and Sandrine Bouchelkia

Try this demonstration to recreate an early ‘breathalyser’ test, passing ethanol vapour through potassium dichromate. Includes kit list and safety instructions

The ‘Old Nassau’ or Halloween clock reaction

Illustrate dramatic colour changes as a result of redox and precipitation reactions in this vivid demonstration. Includes kit list and safety instructions.

The combustion of iron wool

Try this quick teacher demonstration to demonstrate the increase in mass as iron wool is heated in air. Includes kit list and safety instructions.

The cornflour ‘bomb’

Create a small explosion inside a tin can using cornflour in this demonstration, illustrating energy transformation. Includes kit list and safety instructions.

The density of carbon dioxide

Illustrate the higher density of carbon dioxide relative to air by pouring it over a lighted candle in this demonstration. Includes kit list and safety instructions.

The density of ice

Demonstrate to students what happens as ice cubes floating on oil start to melt and the density of the water changes. Includes kit list and safety instructions.

The effect of pressure and temperature on equilibrium | Le Chatelier’s principle

Try this demonstration to explore the effects of pressure and temperature on an equilibrium mixture with your students. Includes kit list and safety instructions.

The equilibrium between two coloured cobalt species

In this demonstration the equilibrium between two different coloured cobalt species is disturbed. Le Chatelier’s principle is used to predict a colour change.

The oxidation states of vanadium

Introduce your students to the idea that different oxidation states of transition metal ions often have different colours, and that electrode potentials can be used to predict the course of the redox reactions. Includes kit list and safety instructions. 

The reaction of aluminium and copper(II) sulfate

Try this practical or demonstration to illustrate the displacement of copper from copper sulfate using aluminium foil, with kit list and safety instructions.

The reaction of ethyne with chlorine

Try this teacher demonstration with your students to illustrate the spontaneous reaction of ethyne and chlorine. Includes kit list and safety instructions.

The reaction of magnesium with steam

Plunge a burning magnesium ribbon into the steam above boiling water and allow the hydrogen that is formed to burn – or collect it over water and test it with a lighted spill.

The silver mirror test with Tollens’ reagent

Try this practical to explore the mirror-making reaction between silver nitrate (Tollens’ reagent) and glucose. Includes kit list, video and safety instructions.

The thermal properties of water

Explore water’s boiling point, specific heat capacity and thermal conductivity in this demonstration. Includes kit list and safety instructions.

The thermite reaction between aluminium and iron(III) oxide

Illustrate a highly exothermic thermite reaction resulting in molten iron in this teacher demonstration. Includes kit list and safety instructions.

Thermal decomposition of nitrates: ‘writing with fire’

Make an invisible message ‘glow’ by applying a lighted splint to filter paper treated with sodium nitrate in this demonstration. Includes kit list and safety instructions.

Turning ‘red wine’ into ‘water’

Use acidified potassium permanganate – or ‘red wine’ – to make ‘water’, ‘milk’ and ‘lemonade’ in this engaging demonstration. Includes kit list and safety instructions.

Turning copper coins into ‘silver’ and ‘gold’

Perform what looks like alchemy with ordinary copper coins in this teacher demonstration. Includes kit list and safety instructions.

Universal indicator ‘rainbow’

Try this demonstration to create a rainbow effect using universal indicator, hydrochloric acid and sodium hydroxide. Includes kit list and safety instructions.

Urea-methanal polymerisation

Explore condensation polymerisation by creating and investigating the properties of a thermosetting polymer in this demonstration. Includes kit list and safety instructions.

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Chemical Reactions: Types of reactions and the laws that govern them

by Anthony Carpi, Ph.D., Adrian Dingle, B.Sc.

Did you know that chemical reactions happen all around us, such as when you light a match, start a car, or even take in a breath of air? But no matter the type of reaction, in every case a new substance is produced and is often accompanied by an energy and/or an observable change.

• The steps from a qualitative science to quantitative one, were crucial in understanding chemistry and chemical reactions more completely.
• When a substance or substances (the reactants), undergo a change that results in the formation of a new substance or substances (the products), then a chemical reaction is said to have taken place.
• Mass and energy are conserved in chemical reactions. Matter is neither created or destroyed, rather it is conserved but rearranged to create new substances. No energy is created or destroyed, it is conserved but often converted to a different form.
• Chemical reactions can be classified into different types depending on their nature. Each type has its own defining characteristics in terms of reactants and products.
• Chemical reactions are often accompanied by observable changes such as energy changes, color changes, the release of gas or the formation of a solid.
• Energy plays a crucial role in chemical reactions. When energy is released into the surroundings the reaction is said to be exothermic; when energy is absorbed from the surroundings the reaction is said to be endothermic

This is an updated version of the module Chemical Reactions (previous version) .

Chemical reactions happen absolutely everywhere. While we sometimes associate chemical reactions with the sterile environment of the test tube and the laboratory - nothing could be further from the truth. In fact, the colossal number of transformations make for a dizzying, almost incomprehensible array of new substances and energy changes that take place in our world every second of every day.

In nature, chemical reactions can be much less controlled than you’ll find in the lab, sometimes far messier, and they generally occur whether you want them to or not! Whether it be a fire raging across a forest (Figure 1), the slow process of iron rusting in the presence of oxygen and water over a period of years, or the delicate way in which fruit ripens on a tree, the process of converting one set of chemical substances (the reactants) to another set of substances (the products) is one known as a chemical reaction .

Figure 1 : A controlled fire in Alberta, Canada, set to create a barrier for future wildfires.

Though chemical reactions have been occurring on Earth since the beginning of time, it wasn’t until the 18th century that the early chemists started to understand them. Processes like fermentation, in which sugars are chemically converted into alcohol , have been known for centuries; however, the chemical basis of the reaction was not understood. What were these transformations and how were they controlled? These questions could only be answered when the transition from alchemy to chemistry as a quantitative and experimental science took place.

• Historical context

Beginning in the early Middle Ages , European and Persian philosophers became fascinated with the way that some substances seemed to “transmute” (or transform) into others. Simple stones, such as those that contained sulfur, seemed to magically burn; and otherwise unimpressive minerals were transformed, like the ore cinnabar becoming an enchanting silvery liquid metal mercury when heated. Alchemists based their approach on Aristotle’s ideas that everything in the world was composed of four fundamental substances - air, earth, fire, and water (Figure 2).

Figure 2 : Aristotle believed that everything in the world was composed of four fundamental substances - air, earth, fire, and water.

As such, they proposed, and spent generations trying to prove, that less expensive metals like copper and mercury could be turned into gold. Despite their misguided approach, many early alchemists performed foundational chemical experiments - transforming one substance into another, and so it is difficult to point to a specific date or event as the birth of the idea of an ordered, quantifiable chemical reaction . However, there are some important moments in history that have helped to make sense of it.

• Lavoisier: Law of Mass Conservation

Antoine Lavoisier was a French nobleman in the 1700s who began to experiment with different chemical reactions . At the time, chemistry still couldn’t be described as being a true, quantitative science. Most of the theories that existed to explain the way that substances changed relied upon Greek philosophy, and there was precious little experimental detail attached to the alchemist’s tinkering.

However, during the second half of the 18th century, Lavoisier performed many quantitative experiments and observed that while substances changed form during a chemical reaction , the mass of the system – or a measure of the total amount of “stuff” present – did not change. In doing so, Lavoisier championed the idea of conservation of mass during transformations (Figure 3). In other words, unlike the alchemists before him who thought that they were creating matter out of nothing, Lavoisier proposed that substances are neither created nor destroyed, but rather change form during reactions . Lavoisier’s ideas were published in the seminal work Traité élémentaire de Chimie in 1789 (Lavoisier, 1789), which is widely hailed as the birth of modern chemistry as a quantitative science.

Figure 3 : Lavoisier's Law of Mass Conservation, which states that substances are neither created nor destroyed, but rather change form during reactions. In this example, the reactants (zinc and two hydrogen chloride molecules) convert into different products (zinc chloride and dihydrogen), but no mass is lost or created.

• Proust: Law of Constant Composition

Joseph Proust was a French actor who followed in Lavoisier’s footsteps. Proust performed dozens of chemical reactions , starting with different amounts of various materials. Over time he observed that no matter how he started a certain chemical reaction , the ratio in which the reactants were consumed was always constant . For example, he worked extensively with copper carbonate and no matter how he changed the ratio of starting reactants, the copper, carbon, and oxygen all reacted together in a constant ratio (Proust, 1804). As a result, in the last few years of the 18th century, Proust formulated the law of constant composition (also referred to as the law of definite proportions, Figure 4).

He realized that any given chemical substance (that we now define as a compound) always consisted of the same ratio by mass of its elemental parts regardless of the method of preparation. This was a huge step forward in modern chemistry since it had been previously believed that the substances formed during chemical reactions were random and disordered.

Figure 4 : An example of Proust's Law of Constant Composition, which states that any compound always consists of the same ratio by mass of its elemental parts, regardless of the method of preparation.

• Dalton: Law of Multiple Proportions

The English chemist John Dalton helped make sense of the laws of conservation of mass and definite proportions in 1803 by proposing that matter was made of atoms of unique substances that could not be created or destroyed (see our module Early Ideas about Matter for more information).

Dalton extended Proust’s ideas by recognizing that it was possible for two elements to form more than one compound , but that whatever the compound was, it would always contain elements combined in whole number ratios (Dalton, 1808). This observation is known as the law of multiple proportions (Figure 5) and with his atomic theory , helped to cement Lavoisier’s observations.

Figure 5 : Dalton's Law of Multiple Proportions, which states that two elements may form more than one compound, but whatever the compound was, it would always contain elements combined in whole number ratios

These advancements, taken together, laid the groundwork for our modern understanding of chemical reactions , chemical equations, and chemical stoichiometry , or the process of expressing the relative quantities of reactants and products in a chemical reaction .

Comprehension Checkpoint

• Types of chemical reactions

There is a staggering array of chemical reactions . Chemical reactions occur constantly within our bodies, within plants and animals, in the air that circulates around us, in the lakes and oceans that we swim in, and even in the soil where we grow crops and build our homes. In fact, there are so many chemical reactions that occur that it would be difficult, if not impossible, to understand them all. However, one method that helps us to understand them is to categorize chemical reactions into a few, general types. While not a perfect system , placing reactions together according to their similarities helps us to identify patterns, which in turn allows predictions to be made about as yet unstudied reactions. In this module, we will consider and provide some context for a few categories of reactions, specifically: synthesis , decomposition, single replacement, double replacement, REDOX (including combustion), and acid-base reactions.

No matter the type of reaction , one universal truth applies to all chemical reactions . For a process to be classified as a chemical reaction, i.e., one where a chemical change takes place, a new substance must be produced. The formation of a new substance is nearly always accompanied by an energy change, and often with some kind of physical or observable change. The physical change can be of different types, such as the formation of bubbles of a gas , a solid precipitate , or a color change. These changes are clues to the existence of a chemical reaction and are important triggers for further research by chemists.

• Synthesis reactions

Prior to Lavoisier’s work, it was poorly understood that there were different gases made up of different elements . Instead, various gases were commonly mischaracterized as types of "air" or air missing parts – for example, terms commonly used were "inflammable air," or "dephlogisticated air." Lavoisier thought differently and was convinced that these were different substances. He conducted experiments where he mixed inflammable air with dephlogisticated air and a spark, and he found that the substances combined to produce water. In response, he renamed inflammable air "hydrogen" from the Greek hydro for "water" and genes for "creator." In so doing, Lavoisier was identifying a synthesis reaction . In general, a synthesis reaction is one in which simpler substances combine to form another more complex one. Hydrogen and oxygen (which Lavoisier also renamed dephlogisticated air) combine in the presence of a spark to form water, summarized by the chemical equation shown below (for more on chemical equations see the section called Anatomy of a chemical equation ), it represents a simple synthesis reaction.

2H 2(g) + O 2(g) → 2H 2 O (l)

• Decomposition reactions

In 1774, the scientist Joseph Priestley turned his curiosity to a mineral called cinnabar – a brick red mineral. When he placed the mineral under sunlight amplified by a powerful magnifying glass, he found that a gas was produced which he described as having an “exalted nature” because a candle burned in the gas brightly (Priestley, 1775). Without realizing it, Priestley had discovered oxygen as a result of a decomposition reaction . Decomposition reactions are often thought of as the opposite of synthesis reactions since they involve a compound being broken down into simpler compounds or even elements . In the case of Priestley’s oxygen, he had broken down mercury (II) oxide (cinnabar) with heat into its individual elements. The reaction can be summarized in the following equation.

• Single replacement reactions

The British chemist and meteorologist John Daniell, invented one of the very first practical batteries in 1836 (Figure 6). In his cell, Daniell utilized a very common single replacement reaction . His early cells were complicated affairs, with ungainly parts and complicated constructs, but by contrast, the chemistry behind them was really quite simple.

Figure 6 : Daniell cell batteries.

In certain chemical reactions , a single constituent can substitute for another one already joined in a chemical compound . The Daniell cell works because zinc can substitute for copper in a solution of copper sulfate, and in so doing exchange electrons that are used in the battery cell. The reaction can be summarized as follows:

This particular single displacement is called a metal displacement since it involves one metal replacing another metal, and many types of batteries are based on metal replacement reactions . However, several other types of single replacement reactions exist, such as when a metal can replace hydrogen from an acid or from water, or a halogen can replace another halogen in certain salt compounds .

• Combustion reactions

The controlled use of fire was a crucial development for early civilization. While it’s difficult to pin down the exact time that humans first tamed the combustion reactions that produce fire, recent research suggests it may have occurred at least a million years ago in a South African cave (Berna et al. 2012).

Chemically, combustion is no more than the reaction of a fuel (wood, oil, gasoline, etc.) with oxygen. For combustion to take place there must be a fuel and oxygen gas . However, these reactions often require activation energy (discussed in more detail in the module Chemical Bonding: The Nature of the Chemical Bond ), which can be provided by a ‘spark’ or source of energy for ignition. Fuel, oxygen, and energy are the three things make up what is known as the fire triangle (Figure 7), and any one of them being absent means that combustion will not take place.

Figure 7 : The fire triangle is made up of three things - fuel, oxygen, and energy.

In the modern world, many of the fuels that are typically burned for energy , are hydrocarbons – substances that contain both hydrogen and carbon (as discussed in more detail in our Carbon Chemistry module). Plants produce hydrocarbons when they grow, and thus make an excellent fuel source, and other hydrocarbons are produced when plants or animals decay over time (such as natural gas , oil, and other substances). When these fuels combust, the hydrogen and carbon within them combine with oxygen to produce two very familiar compounds , water, and carbon dioxide. One simple example is the combustion of natural gas, or methane, CH 4 :

As with the combustion of all fuels, heat and light are products , too, and it is these products that are used to cook our food or to heat our homes.

• Reduction-oxidation reactions

Each of the four types of reaction above are sub-categories of a single type of chemical reaction known as redox reactions. A redox reaction is one where reduction and oxidation take place together, hence the name. The individual processes of oxidation and reduction can be defined in more than one way, but whatever the definition, the two processes are symbiotic, i.e., they must take place together.

In one definition, oxidation is described as the process in which a species loses electrons , and reduction is a process where a species gains electrons. In this way, we can see how the pair must take place together. If a chemical substance is to lose electrons (and therefore be oxidized), then it must have another, interdependent chemical substance that it can give those electrons to. In the process, the second substance (the one gaining electrons) is said to be reduced. Without such an electron acceptor, the original species can never lose the electrons and no oxidation can take place. When the electron acceptor is present, it gets reduced and the redox combination process is complete. Redox reactions of this type can be summarized by a pair of equations – one to show the loss of electrons (the oxidation), and the other to show the gain of electrons (the reduction). Using the example of the Daniell cell above,

The electrons shown being lost by zinc in the first reaction , are the same electrons being accepted by the copper ions in the second. Together, the reactions can be combined to cancel out the electrons on either side of the reactions, into the overall redox reaction:

Other definitions of oxidation and reduction also exist, but in every case, the two halves of the redox reaction remain symbiotic – one loses and the other gains. The loss from one species cannot happen without the other species gaining.

• Double displacement reactions

When soap won’t easily produce a lather in water, the water is said to be ‘hard’. Hard water causes all kinds of problems that go beyond just making it difficult to form a lather. The buildup of compounds in water pipes (known as ‘scale’), can block the flow of water and can cause problems in industrial processes. Textile manufacturing and the beverage industry rely heavily on water. In those situations, the quality of the water can make a difference to the end product , so controlling the water composition is crucial.

Hard water contains magnesium or calcium ions in the form of a dissolved salt such as magnesium chloride or calcium chloride. When soap (sodium stearate) comes into contact with either of those salts, it enters into a double displacement reaction that forms the insoluble precipitate known as ‘soap scum’.

A double displacement reaction (also known as a double replacement reaction) occurs when two ionic substances come together and both substances swap partners. In general:

Where A and C are cations (positively charged ions), and B and D are anions (negatively charged).

In the case of the reaction of soap with calcium chloride, the reaction is:

The solid calcium stearate is what we call soap scum, which is formed by the reaction of the soluble sodium stearate salt (the soap) in a double replacement reaction with calcium chloride.

• Acid-Base reactions

Acid-base reactions happen around, and even inside of us, all the time. From the classic elementary school baking soda volcano to the process of digestion, we encounter acids and bases on a daily basis.

When a hydrogen atom loses its only electron , it forms a positive ion , H + . This hydrogen ion is the essential component of all acids , and indeed one definition of an acid is that of a hydrogen ion donor. Compounds such as the citric acid in lemon juice, the ethanoic acid in vinegar, or a typical laboratory acid like hydrochloric acid, all give their hydrogen ions away in chemical reactions known as acid-base reactions . The chemical opposites of acids are known as bases , and bases can be defined as hydrogen ion acceptors. Whenever an acid donates a hydrogen ion to a base, an acid-base reaction has taken place, for example, when hydrochloric acid donates a hydrogen ion to a base such as sodium hydroxide:

A closer look at this reaction reveals that in water the HCl gives away an H + as shown below:

The resulting species , H 3 O + (the hydronium ion), can, in turn, act as an acid when it comes into contact with any species that can accept a hydrogen ion , such as hydroxide ions from sodium hydroxide:

Combining equations #9a and #9b gives us equation #9c.

Equation #9c can be re-written to show the individual ions that are found in solution , thus:

Removing the spectator ions from the equation above, we get the net ionic equation:

Any chemical reaction that forms water from the reaction between an acid and base as in equation #9e is known as a neutralization reaction.

• Anatomy of a chemical equation

Chemical equations are always linked to chemical reactions since they are the shorthand by which chemical reactions are described. That fact alone makes equations incredibly important, but equations also have a crucial role to play in describing the quantitative aspect of chemistry, something that we formally call stoichiometry .

All chemical reactions take on the same, basic format. The starting substances, or reactants , are listed using their chemical formula to the left-hand side of an arrow, with multiple reactants separated with plus signs. In the case of a reaction between carbon and oxygen:

To the right hand of the arrow one finds the chemical formulas of the new substance or substances (known as the products) that are produced by the chemical reaction . In this case, since carbon dioxide is the result of burning carbon in the presence of oxygen:

Since reactions can result in both physical as well as chemical changes, each substance is given a state symbol written as a subscript to the right of the formula , this describes the physical form of the reactants and products . Common state abbreviations are (s) for solids , (l) for liquids , (g) for gases and (aq) for any aqueous substances, i.e., those dissolved in water.

Finally, in order to ensure that this representation abides by the law of conservation of mass , the equation may need to be balanced by the addition of numbers in front of each species that create equal numbers of atoms of each element on each side of the equation. In the case of the formation of carbon dioxide from carbon and oxygen, there is no need for the addition of such numbers (called the stoichiometric coefficients), since 1 carbon atom and 2 oxygen atoms appear on each side of the equation.

• Energy changes

In nature, chemical reactions are often driven by exchanges in energy . In this respect, reactions are generally separated into two categories – those that release energy and those that absorb energy.

Exothermic reactions are those that release energy to the surroundings (Figure 8, right). Combustion reactions are an obvious example because the energy released by the reaction is converted into the light and heat seen in the immediate surroundings.

By contrast, endothermic reactions are those that absorb energy from the surroundings (Figure 8, left). In this situation, one may have to heat up the reaction or add some other form of energy to the system before seeing the reaction proceed.

Figure 8 : On the left is an endothermic reaction, where energy is absorbed from the surroundings. In contrast, on the right is an exothermic reaction, which releases energy into the surroundings.

In both cases it is important to note that energy is neither created nor destroyed, rather it is transferred from one type of energy to another, for example from chemical energy to that of heat or light . The energy that goes into the formation of chemical bonds is exchanged for other types of energy with the environment around that reaction . A classic example is the photosynthesis reaction, in which plants absorb light energy from the sun in order to create bonds between atoms that make up sugars , which are stored as chemical energy for later use by the plant. The process of respiration is essentially the reverse of photosynthesis, where the bonds in sugar molecules are broken and the released energy is then used by the plant.

• The context of chemical reactions

Chemical reactions happen all around us every day. Whether it is a single replacement reaction in the battery of our flashlight, a synthesis reaction that occurs when iron rusts in the presence of water and oxygen, or an acid-base reaction that happens when we eat – we experience chemical reactions in almost everything we do. Understanding these reactions is not an abstract concept for a chemist in a far off laboratory, rather it is critical to understanding life and the world around us. To truly master chemical reactions, we need to understand the quantitative aspect of these reactions, something referred to as stoichiometry , and a concept we will discuss in another module.

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Types of Chemical Reactions

Intro to Chemical Reactions Chapter 6. How do you know if a chemical reaction occurred? We look for visual signs OR a chemical change -If something bubbles,

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Chemical Reactions: Reactants change chemical and physical properties …. to become new substances made from the same elements; these are called products.

8-2 Types of Chemical Reactions. Combustion: A combustion reaction is when oxygen combines with a hydrocarbon to form water and carbon dioxide. These.

Types of Chemical Reactions We will study 4 types.

There are six main types of reactions

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NaOH + Pb(NO 3 ) 2 Pb(OH) 2 + NaNO 3 ORDER: 1.Balance metals 2.Balance non-metals 3.Balance oxygen 2 nd last 4.Balance hydrogen last Tip for balancing.

Chemical Reactions Chapter 16 Section 4. Synthesis Reactions two or more substances combine to form another substance A + B  AB N 2 (g) +3H 2 (g)  2NH.

Aim: What are the five general types of reactions? Do Now: Write and balance the Chemical Equation 1.Zinc and lead (II) nitrate react to form zinc nitrate.

Types of Chemical Reactions Notes

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Chemistry Solutions

May 2019 | Resource Feature

My Favorite Demonstrations for Teaching Chemical Reactions

By Stephanie Harry

Demos & Labs

One of my favorite units to teach is chemical reactions. Chemistry can be a challenging subject, but in my experience, my students seem to enjoy and understand many of the concepts covered in this unit. Since adopting the flipped learning teaching style, I am continuously looking for interactive reinforcement activities to use during class. In this unit in particular, I feel like I am heading in a good direction.

I begin the chemical reaction unit with a series of demonstrations to teach students how to identify evidence that a chemical reaction has occurred, how to write a word equation to explain a chemical reaction, and how to convert a word equation to a balanced chemical equation.

Before beginning the demonstration, I typically ask my students to provide me with a list of the five indicators that can suggest that a chemical reaction may have occurred. The suggested proofs identified by the students are: change in color, formation of a gas, formation of a precipitate, change in temperature, and production of light.

In my experience, although students can often list these five points, not all students have actually observed them in the chemistry setting. For example, I once had a student who thought adding water to red punch, and then seeing the color fade, was an indication of a chemical reaction. Another student assumed that because food changes temperature when placed in the refrigerator, this was a sign that a chemical reaction had occurred.

To overcome these misconceptions, I explain that when we are speaking of change in color or temperature, we are looking for unexpected changes. However, even in doing this, I had felt that I needed to find a way to provide a visual explanation to support my students’ understanding. This thought led me to create a series of seven chemical reaction demonstrations, based on my experience attending different workshops, collaborating with other teachers, and researching the topic online. These days, I perform these demonstrations in the following order:

• Copper metal granules are added to 1M hydrochloric acid.
• Zinc metal granules are added to 1M hydrochloric acid.
• Aqueous solutions of potassium iodide and lead (II) nitrate are combined.
• Aqueous solutions of potassium thiocyanate and iron (III) nitrate are combined.
• Aqueous solutions of potassium iodide and aluminum nitrate are combined.
• Ammonium dichromate is ignited.
• Magnesium ribbon is ignited.

To transition easily from one demonstration to the next, I set up the materials for the seven demonstrations in the following manner:

• Demonstrations 1 and 2 : I place the 1M hydrochloric acid in a test tube, and each metal sample in a beaker.
• Demonstrations 3, 4, and 5 : I place one of the aqueous solutions in a test tube and the other solution in a dropper bottle. Alternatively, both reactants can be placed in test tubes.
• Demonstration 6 : I place ammonium dichromate in a 1000 mL beaker (or larger) with wire gauze over the opening. Since it will be ignited, I also create a wick using a small piece of paper towel soaked in ethanol. I use tongs to place the wick on top of the ammonium dichromate (like placing a candle into a birthday cake) and then use a burning wooden splint to light the wick.
• Demonstration 7 : I place a precut magnesium ribbon in a beaker.

When I am ready to begin the demonstrations, my students and I put on our personal protective equipment and I outline safety precautions prior to each demonstration. Although my students are not handling the chemicals directly, safety is the priority. One reason for this level of caution is that I plan to move around the classroom with the chemicals (as appropriate) for students to get a closer observation of the results.

I often get very specific about safety instructions. Prior to the first and second demonstrations, I inform students that the reactant, 1M HCl, is a corrosive chemical. It is essential that eyes and skin are protected from contact, and the demonstration should be performed in a fume hood or a well ventilated area. Additionally, prior to Demonstration 6, I explain that it will be performed in a fume hood because ammonium dichromate contains toxic chromium (VI). Finally, before we do Demonstration 7, I remind students to not look directly at the magnesium while it is burning in the reaction.

Figure 2. Summary of safety precautions for each demonstration.

I also use a document camera to display each demonstration, which is connected to a classroom Smart Board, allowing all students to easily see the details of the reactions as they occur. Additionally, this serves as a helpful way to display the written information from the student handout next to the results. During the demonstrations, I compile and record information shared by my students in the data table displayed using the document camera and Smart Board.

Before beginning a reaction, I have my students describe the appearance of the reactant(s) in the data table provided on the student handout (Figure 3). I direct them to record details about color, physical state, texture of the reactants, etc.

Figure 3. Example of activity data table for students to compete.

As I complete each demonstration, I ask the students whether or not a chemical reaction has occurred. To help them determine their answer, I remind them to consider whether the appearance of any of the reactants changed, and record their answers in the appropriate column. I also remind them that chemical equations are written to explain chemical reactions; so, if there is no chemical reaction, then there is no chemical equation. This helps to remind them that just because there is no observable evidence (change in color, formation of a precipitate, etc.), it doesn’t mean that a chemical reaction has not occurred. Some additional notes that may be helpful:

• In this activity, Demonstrations 1 and 5 will not produce a chemical reaction.
• Five of the reactions will produce indicators of a chemical reaction for students to observe.
• During Demonstration 2, students will observe the formation of gas, as bubbles are produced when zinc metal is added to a test tube containing 1M hydrochloric acid. Additionally, they will notice the change in color of the zinc metal.
• Students should note gas formation in several of the demonstrations, including Demonstrations 2 and 6.
• Demonstrations 3, 4, and 6 will each show a distinctive color change.
• Demonstration 3 will produce a precipitate.
• Students will notice light is produced in Demonstrations 6 and 7.

After all seven demonstrations have been completed, I work with students to develop word equations for any demonstrations that resulted in a chemical reaction.

Figure 5. Fill-in-the-blank word equations allow students to indicate the products of the reactions that they observed.

I have found it interesting that after writing the name of the products in Demonstration 3, some students have seen a similarity to Demonstration 4, and were able to correctly predict the products for that reaction. Because these students saw a pattern between the rearrangement of metal and non-metal ions, I introduced them to the term double replacement . This also helped reiterate that all of the elements used in the reactant were present in the product.

In the final component of this activity, students convert the word equations to formula equations. Depending on the level of student understanding, teachers may want to keep the focus on proper formula writing , rather than on actually balancing the equation. My goal at this point in the lesson is to make sure students use their prior knowledge to write the correct formulas for each compound, and place them on the correct side of the chemical equation.

I complete the formula equation for Demonstration 2 with the class and allow the students to complete the other equations on their own. I use this time to move around the classroom and provide individualized help to students, sharing such tips as:

• When completing the formula equation for Demonstration 2, the elemental symbol for zinc is Zn, but the elemental symbol for hydrogen is H 2 .
• Be careful to properly incorporate symbols, states of matter, and yielding in the equation. I find it easier to share this guidance during this activity, since students have visual examples of the physical state of the reactants and products.

These chemical demonstrations have truly been beneficial in helping students identify indicators of chemical reactions, as they offer great visual examples of the five indicators of a chemical reaction. I have found that this lesson helps minimize my students’ misconceptions and gives them a visual connection to prove that a chemical reaction has occurred. Gathering, recording, and displaying information from students to complete the data table, and allowing students to engage in dialogue about the demonstrations, makes the activity both educational for my students, and more manageable for me.

This activity provides teachers with seven chemical demonstrations that can be easily setup and incorporated into a unit on chemical reactions. Chemical demonstrations are a great way to engage students and create interaction. I encourage you to try it with your own students!

Photo credit: (article cover) Kesu01/Bigstockphoto.com

Types of Chemical Reactions Interactive Notes with PowerPoint

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Description

These interactive notes will help students classify and define the major types of chemical reactions through a combination of engaging visual aids and real life applications. The 5 types of chemical reactions explored include synthesis, decomposition, combustion, single replacement and double replacement. The interactive notes align with a PowerPoint slide show and an animated video, both of which are included. 

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With what standard does this resource align?

This resource will support students in moving  towards the mastery of the Next Generation Science Standards.

Middle school: 

NGSS MS-PS1.B: Chemical Reactions: Substances react chemically in characteristic ways. In a chemical process, the atoms that make up the original substances are regrouped into different molecules, and these new substances have different properties from those of the reactants.

High School: 

NGSS HS-PS1-2: Construct and revise an explanation for the outcome of a simple chemical reaction based on the outermost electron states of atoms, trends in the periodic table, and knowledge of the patterns of chemical properties. 

Is there prep involved? 

No prep is required with this all inclusive package. A variety of versions of interactive notes are provided. This allows you to print a version that works best to support your teaching style and your students’ needs.  I have included a PowerPoint slide show and a link to ysci’s animated video that mirrors the interactive notes. In addition, there is a set of practice problems and answer key on identifying the types of chemical reactions. 

How do I use this resource? 

Teacher led note taking: Use the included PowerPoint and interactive notes 

Flipped learning: Students independently watch the video and complete the interactive notes 

In class practice 

Homework 

What’s included? 

1 blank version of the interactive notes 

1 fill-in-the blank version of the interactive notes 

1 version of the interactive notes with guiding questions only  

1 full color completed set of interactive notes 

1 PowerPoint slide show with vibrant images of chemical reactions & real life applications 

1 set of practice problems for identifying the types of chemical reactions with teacher key 

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Instagram: @yscivideo 

YouTube: https://www.youtube.com/yscivideo

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Copyright © 2019 Ysci 

By purchasing this file, you agree to the following terms. All rights reserved by author. This product is to be used by the original downloader only for personal or classroom use only. Copying for more than one teacher, classroom, department, school, or school system is prohibited. Additional licenses can be purchased at a discount for others to use in your department. This entire document, or any parts within, may not be reproduced or displayed for public viewing. You may NOT electronically post this product online including to teacher blogs, classroom websites, school networks or Google Classroom. Failure to comply is a copyright infringement and a violation of the Digital Millennium Copyright Act (DMCA).

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Visualize how the reaction happens in 3D (Use your mouse to zoom and rotate the reaction)

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Examples of Chemical Reactions in Everyday Life

Chemical reactions occur everywhere in the world around you, not just in a chemistry lab. Here are 20 examples of chemical reactions in everyday life and a closer look at what’s happening on a molecular level.

How to Recognize a Chemical Reaction

The first step to recognizing chemical reactions in the world around you is identifying when a reaction is taking place. Chemical reactions cause chemical changes . In other words, substances interact and form new products . Not every change in matter is a chemical reaction . For example, melting ice, tearing a sheet of paper into strips, and dissolving sugar in water are physical changes that don’t change the chemical identity of matter .

Here are some signs of a chemical reaction. If more than one sign is present, it’s likely a reaction has occurred:

• Temperature change
• Color change
• Bubbling or gas production
• Formation of a solid called a precipitate when liquids are mixed

20 Examples of Chemical Reactions in Everyday Life

Here are some broad examples of chemical reactions in daily life:

Photosynthesis

• Aerobic cellular respiration
• Anaerobic respiration (including fermentation )
• Oxidation (including rust)
• Metathesis reactions (such as baking soda and vinegar)
• Electrochemistry (including chemical batteries)
• Soap and detergent reactions
• Acid-base reactions
• Rotting of food
• Electroplating metals
• Disinfecting surfaces and contact lenses
• Leaves changing color with seasons
• Salt keeping ice off roads and helping to freeze ice cream

Examples of Organic Compounds

Some chemicals are inorganic, while those with carbon and hydrogen are organic. Here are examples in everyday life.

A Closer Look at Chemical Reactions in Daily Life

Here is a closer look at some everyday reactions, along with some chemical equations.

You experience combustion reactions when you strike a match, burn a candle, start a campfire, or light a grill. In a combustion reaction, a fuel reacts with oxygen from air to produce water and carbon dioxide. Here is the reaction for the combustion of propane, a fuel used in gas grills and some fireplaces: C 3 H 8  + 5O 2  → 4H 2 O + 3CO 2  + energy 

Plants use a chemical reaction called photosynthesis to convert carbon dioxide and water into food (glucose) and oxygen. It’s a key reaction because it generates oxygen and yields food for plants and animals. The overall chemical reaction for photosynthesis is: 6 CO 2  + 6 H 2 O + light → C​ 6 H 12 O 6  + 6 O 2

Aerobic Cellular Respiration

Animals use the oxygen provided by plants to perform essentially the reverse reaction of photosynthesis to get energy for cells. Aerobic respiration reacts glucose and oxygen to form water and chemical energy in the form of adenosine triphosphate ( ATP ). Here is the overall equation for aerobic cellular respiration: C 6 H 12 O 6  + 6O 2  → 6CO 2  + 6H 2 O + energy (36 ATP)

Anaerobic Cellular Respiration

Organisms also have ways of getting energy without oxygen. Humans use anaerobic respiration during intense or prolonged exercise to get enough energy to muscle cells. Yeast and bacteria use anerobic respiration in the form of fermentation to make everyday products, such as wine, vinegar, yogurt, bread, cheese, and beer. The equation for one form of anerobic respiration is: C 6 H 12 O 6  → 2C 2 H 5 OH + 2CO 2  + energy

Rust, verdigris, and tarnish are all examples of common oxidation reactions. When iron rusts, it changes color and texture to form a flake coating called rust. The reaction also releases heat, but it usually occurs too slowly for this to be noticeable. Here is the chemical equation for the rusting of iron: Fe + O 2  + H 2 O → Fe 2 O 3 . XH 2 O

Electrochemistry

Electrochemical reactions are redox (oxidation and reduction) reactions that convert chemical energy into electrical energy. The type of reaction depends on the battery. Spontaneous reactions occur in galvanic cells, while nonspontaneous reactions take place in electrolytic cells.

Digestion is a complex process that involves thousands of chemical reactions. When you put food in your mouth, water and the enzyme amylase breaks down sugar and other carbohydrates into simpler molecules. Hydrochloric acid and enzymes break down proteins in your stomach. Sodium bicarbonate released into the small intestine neutralizes the acid and protects the digestive tract from dissolving itself.

Soap and Detergent Reactions

Washing your hands with water isn’t a chemical reaction because you’re just mechanically rinsing away grime. But, when you add soap or detergent, chemical reactions occur that emulsify grease and lower surface tension so you can remove oily grime. Even more reactions occur in laundry detergent, which may contain enzymes to break apart proteins and whiteners to prevent clothes from looking dingy.

Just mixing dry ingredients usually doesn’t result in a chemical reaction. But, adding a liquid ingredient often results in a reaction. Cooking with heat also causes reactions. Mixing flour, sugar, and salt is not a chemical reaction. Neither is mixing oil and vinegar. Cooking an egg is a chemical reaction because heat polymerizes proteins in egg white, while the hydrogen and sulfur in the yolk can react to form hydrogen sulfide gas. When you heat sugar, a reaction called caramelization occurs. When you heat meat, it browns due to the Maillard reaction . Baked goods rise due to carbon dioxide bubbles formed by the reaction between baking powder or soda and liquid ingredients.

Acid-Base Reactions

Acid-base reactions occur anytime you mix an acid (e.g., lemon juice, vinegar, muriatic acid, battery acid, carbonic acid from carbonated beverages) with a base (e.g., baking soda, ammonia, lye). A good example of an acid-base reaction is the reaction between baking soda and vinegar to form sodium acetate, water, and carbon dioxide gas: NaHCO 3  + HC 2 H 3 O 2  → NaC 2 H 3 O 2  + H 2 O + CO 2 In general, a reaction between an acid and a base produces a salt and water. For example, if you react muriatic acid (HCl) and lye (NaOH), you get table salt (NaCl) and water (H 2 O): HCl + NaOH → NaCl + H 2 O In this reaction, two clear liquids form another clear liquid, but you can tell a reaction occurs because it releases a lot of heat.

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Creating Models of Chemical Reactions

In this activity, activity overview, template and class instructions, more storyboard that activities.

• This Activity is Part of Many Teacher Guides

Use this lesson plan with your class!

Creating models of chemical reactions is super important to help students understand how things change and how atoms balance. In this activity, students will create a grid that illustrates four chemical reactions . They should be sure to include the reactants, products, and the equation in their finished product. To extend this activity, ask students to add a cell that explains what type of reaction it is (exothermic vs endothermic) and what happens during the process.

• Reaction between sodium and water to produce sodium hydroxide and hydrogen - Na + H 2 O → NaOH + H 2
• Reaction of sulfur trioxide and water to form sulfuric acid - SO 3 + H 2 O → H 2 SO 4

Example Chemical Reactions

• Iron and Sulfur: Exothermic reaction that forms iron sulfide. The word equation for this reaction is iron + sulfur → iron sulfide . The balanced symbol equation is Fe + S → FeS.
• Hydrogen and Oxygen: Reaction that forms water. The word equation for this reaction is hydrogen + oxygen → water . The balanced symbol equation is 2H 2 + O 2 → 2H 2 O
• Methane and Oxygen: Combustion that forms carbon dioxide and water. The balanced symbol equation is CH 4 + 2O 2 → CO 2 + 2H 2 O . This reaction requires students to balance the equation.
• Photosynthesis: Endothermic reaction that produces glucose and oxygen. The word equation for photosynthesis is carbon dioxide + water → glucose + oxygen . The balanced symbol equation is 6CO 2 + 6H 2 O → C 6 H 12 O 6 + 6O 2 . This reaction requires students to balance the equation.

(These instructions are completely customizable. After clicking "Copy Activity", update the instructions on the Edit Tab of the assignment.)

Student Instructions

Model various chemical reactions by creating a T-Chart on Storyboard That. Identify the reactants and products in the reactions and create models of the molecules. Remember: atoms are not created or destroyed in a chemical reaction, so the total number of atoms doesn’t change.

• Click "Start Assignment".
• Reaction between Iron and Sulfur
• Combustion of Hydrogen in Air
• Combustion of Methane in Air
• Photosynthesis
• Write the names of the reactants and products in the title of cell.
• Write the symbols in the description box underneath.
• Add coefficients if necessary to balance the symbol equation.
• Use shapes to model the arrangement of the different atoms for the reactants and products.
• Count the number of each type of atom before and after the reaction to ensure your symbol equation is correctly balanced.

Lesson Plan Reference

Grade Level 6-12

Difficulty Level 3 (Developing to Mastery)

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Using Drawing Technology to Assess Students’ Visualizations of Chemical Reaction Processes

• Published: 26 September 2013
• Volume 23 , pages 355–369, ( 2014 )

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• Hsin-Yi Chang 1 ,
• Chris Quintana 2 &
• Joseph Krajcik 3  

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In this study, we investigated how students used a drawing tool to visualize their ideas of chemical reaction processes. We interviewed 30 students using thinking-aloud and retrospective methods and provided them with a drawing tool. We identified four types of connections the students made as they used the tool: drawing on existing knowledge, incorporating dynamic aspects of chemical processes, linking a visualization to the associated chemical phenomenon, and connecting between the visualization and chemistry concepts. We also compared students who were able to create dynamic visualizations with those who only created static visualizations. The results indicated a relationship between students constructing a dynamic view of chemical reaction processes and their understanding of chemical reactions. This study provides insights into the use of visualizations to support instruction and assessment to facilitate students’ integrated understanding of chemical reactions.

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Acknowledgments

The authors would like to thank the anonymous reviewers for their helpful comments which improved the presentation of the study. We also thank Fang-Chin Yeh for her help with data analysis. The completion of the study was supported by the National Science Council of Taiwan under Grant No. 102-2628-S-017-001-MY3.

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College of Education and College of Natural Science, Michigan State University, East Lansing, MI, USA

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Chang, HY., Quintana, C. & Krajcik, J. Using Drawing Technology to Assess Students’ Visualizations of Chemical Reaction Processes. J Sci Educ Technol 23 , 355–369 (2014). https://doi.org/10.1007/s10956-013-9468-2

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Published : 26 September 2013

Issue Date : June 2014

DOI : https://doi.org/10.1007/s10956-013-9468-2

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