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Ever wondered why a campfire feels warm or why an ice pack cools down your injury? It all comes down to chemistry – specifically, exothermic and endothermic reactions.

These two types of reactions dictate whether heat is released or absorbed.

But what exactly do exothermic and endothermic reactions mean, and why should you care? Let’s break it down in a fun and simple way!

What Are Chemical Reactions?

Let’s get the basics right before diving into exothermic and endothermic reactions.

A chemical reaction occurs when substances (reactants) transform into new substances (products).

Depending on the type of reaction, energy is either absorbed or released. This energy usually appears as heat, but it can also be in the form of light or sound.

Think of it like making a cake. You mix ingredients, put them in the oven, and after some time, you have a delicious treat. Heat (energy) is required to make this transformation happen, just like in some chemical reactions.

Examples of Exothermic and Endothermic Reactions

Exothermic Reactions: The Heat Givers

An exothermic reaction releases heat into the surroundings. It’s like a gift-giver but with energy. The reaction itself produces energy, usually in the form of heat, making the surrounding area warmer.

Examples of Exothermic Reactions

1. Burning a Candle or Wood:

When you burn wood in a fireplace or light a candle on a cake, heat and light are released into the environment. That’s an exothermic reaction at work! The combustion of wood or wax involves breaking chemical bonds in the fuel and combining it with oxygen, producing carbon dioxide, water vapor, and—most noticeably—heat. This is why standing near a campfire feels warm, as the reaction constantly gives off thermal energy.

2. Fireworks Exploding:

Ever noticed how fireworks light up the sky with a bang? That’s a chemical reaction releasing heat, light, and sound all at once. The explosion results from a rapid oxidation reaction, where metal salts and chemical compounds react with oxygen to produce a dazzling display of colors and a burst of energy. The energy released in fireworks not only creates heat but also propels the explosion outward, producing the characteristic loud sound.

3. Condensation of Water Vapor:

When steam cools down and turns into water droplets, it releases heat. This is why a bathroom mirror fogs up after a hot shower. When warm, moisture-laden air comes into contact with a cooler surface, such as a mirror or window, the water vapor loses energy and condenses into liquid droplets, giving off heat in the process. This heat release plays an important role in weather patterns, contributing to cloud formation and even storm development.

4. Respiration (Breathing Process):

The process of breaking down glucose in our body to produce energy is exothermic. That’s how we stay warm! During cellular respiration, glucose reacts with oxygen inside our cells to produce carbon dioxide, water, and energy in the form of ATP (adenosine triphosphate). The breakdown of glucose releases heat, which helps maintain body temperature. This is why exercising makes us feel warmer—our muscles are burning more glucose and releasing more heat as a result.

In an exothermic reaction, the products have less energy than the reactants because some of that energy was released into the environment.

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Endothermic Reactions: The Energy Absorbers

On the flip side, an endothermic reaction absorbs heat from its surroundings. This means the reaction feels cold because it’s taking in energy instead of giving it away.

Examples of Endothermic Reactions

1. Melting Ice:

Ice absorbs heat from its surroundings to melt into water. That’s why your drink gets colder when you add ice cubes. The heat energy is used to break the rigid structure of ice crystals, allowing the molecules to move freely as liquid water. This process continues until all the ice melts, pulling heat from the drink and making it cooler.

2. Boiling Water:

Water needs heat to reach its boiling point. That heat gets absorbed, making this an endothermic process. The added heat increases the energy of the water molecules, making them move faster until they break free from the liquid state and become steam. This is why a pot of boiling water continues absorbing heat even after reaching its boiling point, as the energy is used to turn liquid into gas.

3. Cooking Food:

Whether you’re frying an egg, baking a cake, or roasting vegetables, heat is absorbed by the food to cook it. The heat energy causes chemical changes in proteins, carbohydrates, and fats, transforming raw ingredients into cooked meals. For example, when an egg is fried, the heat absorbed breaks the bonds in the egg proteins, causing them to unfold and solidify into a different texture and color.

4. Photosynthesis (Plants Making Food):

Plants take in sunlight (energy) to convert carbon dioxide and water into glucose and oxygen. Without this endothermic reaction, we wouldn’t have food or oxygen! Chlorophyll in plant cells captures sunlight, and this energy is used to drive a series of chemical reactions that produce glucose, a form of stored energy. This process not only sustains plants but also provides energy for all life on Earth, making it one of the most essential endothermic reactions in nature.

In an endothermic reaction, the products have more energy than the reactants because they absorbed heat during the process.

Key Differences Between Exothermic and Endothermic Reactions

Exothermic and Endothermic Reactions

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Real-Life Applications of Exothermic and Endothermic Reactions

You might be thinking, “Okay, cool… but why does this matter?” Well, exothermic and endothermic reactions impact our everyday lives more than you realize!

Exothermic Reactions in Daily Life:

• Hand Warmers: Those little packets you use in winter contain iron powder that reacts with oxygen, releasing heat to keep your hands toasty.
• Fireplaces & Heaters: Burning fuel releases heat, keeping our homes warm during chilly nights.
• Fireworks & Explosions: These fun displays are just rapid exothermic reactions!

Endothermic Reactions in Daily Life:

• Ice Packs: Instant cold packs use endothermic reactions to absorb heat from your body and reduce swelling.
• Cooking & Baking: Every time you cook, you’re triggering endothermic reactions that make raw ingredients edible.
• Sweating: When we sweat, our body absorbs heat from the environment to cool us down—an endothermic process that keeps us from overheating.

How to Remember the Difference

A quick trick to help you remember exothermic and endothermic reactions is below:

• Exothermic = Exit (Heat exits into the surroundings)
• Endothermic = Enter (Heat enters the reaction)

Easy, right? Now you’ll never mix them up again!

Wrapping It Up

Understanding exothermic and endothermic reactions isn’t just about passing a chemistry test—it’s about seeing how science plays a role in everyday life. From the warmth of a campfire to the cooling effect of an ice pack, exothermic and endothermic reactions are all around us.

So next time you light a candle or sip on a cold drink, take a second to appreciate the chemistry happening right in front of you. Who knew science could be this cool (or hot)?

Got any questions or cool experiments related to these reactions? Drop them in the comments below! Let’s keep the conversation going.

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You may also like... Types of Chemical Reactions in Chemistry

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Master the distinction between ions and radicals with simple definitions, step-by-step comparisons, and real-world examples – only at Selftution.com, the #1 educational website for easy, effective learning!

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Most people worldwide consider ions and radicals the same, but there is a difference between the two. Yes, you read it right; there is a difference between ions and radicals.

In chemistry, ions and radicals are two important concepts that play distinct roles in chemical reactions and processes. Understanding the difference between them can help clarify how substances interact and change.

Ions are charged particles formed when atoms or molecules gain or lose electrons. This charge can be positive or negative, depending on whether electrons are lost or gained.

For example, when sodium loses an electron, it becomes a positively charged ion, while chlorine gains an electron to become a negatively charged ion.

Radicals, conversely, are atoms or molecules with unpaired electrons. These unpaired electrons make radicals highly reactive, as they seek to pair up with other electrons to achieve stability.

Radicals are often involved in chain reactions and play a significant role in processes like combustion and polymerization.

Definition of Ions and Radicals

An ion is an atom of an element or a group of atoms of different elements that behave as a single unit with a positive or negative charge. Whereas, a radical (often called a ‘free radical’) is an atom of an element or a group of atoms of different elements that has at least one unpaired electron. Radicals are neutral, positive, and negative in charge.

In this blog, we’ll explore the fundamental differences between ions and radicals, their unique characteristics, and their roles in chemical reactions.

Difference between ions and radicals

IONS

The molecule of a compound, usually salts in solution states splits into two parts with a positive and negative charge on them. These are ions. The positive ion is called a cation, and the negative ion is an anion. For example, a molecule of sodium carbonate (Na₂CO₃) has two parts: sodium (Na⁺) and carbonate (CO₃^2-). Here, sodium is a positive ion and carbonate is a negative ion.

Ions bear positive or negative charges due to the loss or gain of an electron. An atom or a molecule is neutral as the number of protons in it is equal to that of electrons. So, anytime atoms or molecules lose electrons, it results in the formation of positively charged ions, i.e., cations. Similarly, if atoms or molecules gain extra electrons, it results in negatively charged ions, i.e,. anions. Therefore, the number of protons in ions is never equal to the number of electrons. If protons are more, an ion is positive, and if protons are fewer, then an ion is negative.

Therefore, from the above, the definition of an ion is–

An ion is an atom of an element or a group of atoms of different elements that behave as a single unit with a positive and negative charge on it.

Other than positive and negative, we can also classify ions as –

• Simple and compound ions
• Acidic and basic ions

Simple Ions and Compound Ions:

An ion is considered simple if it consists of a single atom, such as Na⁺, Mg²⁺, or Cl⁻. In contrast, compound ions are made up of more than one atom from different elements. Examples of compound ions include NH₄⁺, SO₄²⁻, and CO₃²⁻. Simple ions are monoatomic, meaning they consist of a single atom, while compound ions can be diatomic or polyatomic, consisting of multiple atoms.

Acidic and Basic Ions:

When salt is in the solution state, it splits into two parts: the positive part is called a basic ion, and the negative part is an acidic ion.

Let’s understand with an example, of a neutralization reaction between potassium hydroxide (a base) and hydrochloric acid. Here, the acid reacts with the base to form a salt (potassium chloride) and water. During the formation of this salt, the base potassium hydroxide contributes to the potassium ion (K⁺). Therefore, (K⁺) is a basic ion. While hydrochloric acid contributes to chloride ions (Cl^–). Therefore, (Cl^–) is an acidic ion.

It is important to note here that some people call “simple and compound ions” “simple and compound radicals”. Similarly, “acidic and basic ions” as “acidic and basic radicals”. There is a reason for that, and we will discuss it after studying radicals.

Examples of positive and negative ions or radicals.

RADICALS

All atoms, except those of inert gases, are chemically unstable. They try to achieve stability by combining with other atoms of the same or different elements, following the octet or duplet rule. However, sometimes even after combining with other atoms, the resulting molecule still has an unpaired electron in its valence shell. Such molecules, with at least one unpaired electron in their outermost shell or valence shell, are called radicals. Due to the presence of this unpaired electron, radicals are highly chemically reactive.

Electrons in atoms usually occur in pairs, with each electron spinning in opposite directions. This opposite spin provides some stability to the atom. However, atoms with an odd number of electrons will always have one unpaired electron. Therefore, these atoms are more unstable than those with an even number of electrons because the unpaired electron lacks a counter-spin. It is important to note that radicals always have an odd number of electrons, which means they will always have at least one unpaired electron.

Mostly, the number of protons and the number of electrons are equal in radicals. Therefore, they are electrically neutral. However, this is not always true. Sometimes, radicals bear positive and negative charges due to the absence or presence of excess electrons. This is the reason why people mistake “positive and negative ions” with “positive and negative radicals”. The radicals with charge are also known as radical ions and are mostly found in organic compounds.

Therefore, from the above, the definition of radical is–

Radical or ‘free radical’ is an atom of an element or a group of atoms of different elements that have at least one unpaired electron. Radicals are neutral, positive, and negative in charge.

Examples of radicals-

Elemental nitrogen and hydroxyl are the best examples of simple free radicals.

Nitrogen:  The atomic number of nitrogen is 7, which means a nitrogen atom has 7 protons and 7 electrons. Because of the odd number of electrons, one electron is always unpaired. Therefore, we consider a nitrogen atom in its free state as a radical. Elemental nitrogen radicals are highly reactive and never exist in the free state in nature.

Hydroxyl: Refer difference between ions and radicals, below.

DIFFERENCE BETWEEN IONS AND RADICALS

To understand the difference between ions and radicals, let’s consider the example of hydroxyl (OH). Hydroxyl ions and hydroxyl radicals are two distinct species that are often mistaken for each other. A hydroxyl ion carries a negative charge, while a hydroxyl radical is neutral.

Hydroxyl radicals are formed as a result of the homolytic fission of water molecules, whereas heterolytic fission results in hydroxyl ions. Before we delve deeper into hydroxyl ions and hydroxyl radicals, let’s first understand homolytic and heterolytic fission.

Homolytic and Heterolytic Fission

‘Fission’ means ‘to split’. During homolytic fission, a molecule of a compound breaks in such a manner that the splitting atoms retain one of the originally bonded electrons. Whereas, in heterolytic fission, one of the atoms moves away with both of the bonded electrons. This results in one set of fragments with a greater number of electrons than the others. The fragment with more electrons bears a negative charge, and the one with fewer electrons bears a positive charge.

Homolytic and heterolytic fission of the water molecule

Now let’s get back to the hydroxyl ion and the hydroxyl radical. During the homolytic fission of a water molecule, it splits into two fragments, hydroxyl (OH) and hydrogen (H). In this fission, the number of protons and the number of electrons in the hydroxyl (OH) fragment is the same, i.e., nine. Therefore, it bears a neutral charge. However, as one unpaired electron exists in this hydroxyl fragment due to an odd number of electrons, it is a hydroxyl radical. Whereas, during heterolytic fission, the water molecule splits in such a way that the hydroxyl fragment (OH) moves away with both of the bonded electrons. This results in more electrons in comparison to protons. It has nine protons and ten electrons. Due to this excess electron, this hydroxyl fragment bears a negative charge and is thus called a hydroxyl ion.

You may also like it …Types of chemical reactions

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Learn the fundamental distinctions between mixtures and compounds through clear definitions, comparison charts, and real-world examples.

Welcome to Selftution.com – making chemistry easy to understand! So, let’s begin.

In chemistry, a mixture and a compound represent two different forms of substances. This post will explore these two forms and examine the key differences between a mixture and a compound.

Additionally, we will discuss why water is classified as a compound, while air is considered a mixture.

Skip to >> Why is water a compound and air a mixture?

Definition of a mixture-

The mixture is an impure substance formed by mixing two or more pure substances (elements and compounds) in any proportion, such that they do not undergo any chemical change and retain their individual property.

To learn more about the mixtures, click here

Definition of Compound –

A compound is a substance formed when two or more elements chemically combine in a fixed ratio.

To learn more about the compounds, click here

DIFFERENCE BETWEEN MIXTURE AND COMPOUND WITH EXAMPLES

Six (6) basic differences between a mixture and a compound are as follows:

01. PURITY

A mixture is an impure substance formed by combining two or more pure substances, whereas a compound is a pure substance.

For example, the air is a mixture and therefore an impure substance, as it contains molecules of different gases. It consists of oxygen (O₂), nitrogen (N₂), carbon dioxide (CO₂), Argon (Ar), water vapor, etc., in varying proportions. In contrast, water (a compound) is a pure substance because it contains only water molecules (H₂O).

Example of a Mixture and a compound

02. PROPERTIES

The components of a mixture do not combine chemically; therefore, they retain their chemical and physical properties. In contrast, a compound is a completely new substance with properties entirely different from those of its components.

For example, a water molecule (H₂O) consists of two hydrogen (H) atoms and one oxygen (O) atom. Water exists as a liquid under normal conditions, whereas hydrogen and oxygen are gases. A mixture of hydrogen and oxygen, when ignited, produces fire, whereas water is commonly used to extinguish fire.

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03. HOMOGENEITY

The mixture may be either homogeneous or heterogeneous, whereas a compound is always homogeneous because it consists of molecules of the same kind.

For example, water has uniform properties throughout, making it homogeneous. Similarly, a mixture of salt and water is homogeneous, whereas a mixture of rice and pulses is heterogeneous.

04. COMPOSITION

In mixtures, the ratio of components can vary. Whereas in the case of a compound, the components are present in a fixed ratio by weight.

For example, we can prepare a mixture of sand and common salt either by mixing 1 gm of salt with 2 gm of sand, or 2 gm of salt with 2 gm of sand. Whereas, we get water, a compound, only when hydrogen and oxygen atoms combine chemically in a fixed ratio of 1:8 by mass.

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05. SEPARATION

We can separate components of a mixture by simple physical methods. Whereas we cannot separate the components of a compound by simply physical means. In compounds, chemical bonds join atoms together. These bonds are very strong and difficult to break. Thus, to get original elements (or atoms) from compounds, we need to apply chemical methods.

For example, we can separate salt from the salt solution simply by evaporation. Whereas, to break the molecule of water into its elements, hydrogen and oxygen, we need to pass an electric current through it.

Interesting home experiment by sciencefix to observe the breaking down of water into its components, hydrogen and oxygen, by passing an electric current through it.

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06. ENERGY CHANGES

The formation of a mixture does not involve any change in energy, whereas the formation of a compound always results in the absorption or release of energy.

For example, during the formation of an aqueous sugar solution, no release or absorption of energy takes place. Whereas, water forms during the combustion of hydrogen gas in the presence of oxygen, which results in the release of a huge amount of energy.

DIFFERENCE BETWEEN MIXTURE AND COMPOUND IN TABULAR FORM

Difference between mixture and compound​

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WHY IS WATER A COMPOUND AND AIR A MIXTURE?

Earlier, water and air were thought to be elements. However, it was later proven that water is a compound, whereas air is a mixture. Below is a list of evidence supporting this fact:

• Water is a compound because hydrogen and oxygen in water are chemically combined in a fixed ratio of 1:8 by mass. In contrast, the main components of air—nitrogen, oxygen, carbon dioxide, and water vapor—are not chemically combined, making air a mixture.
• The chemical composition of water remains constant regardless of its source, whereas the composition of air varies from place to place. For example, during the rainy season, air becomes humid due to increased water vapor. Additionally, impurities such as sulfur dioxide and hydrogen sulfide can alter the composition of air in certain locations.

• Water exhibits properties entirely different from its constituent elements, hydrogen and oxygen. For example, the oxygen in water does not support combustion. In contrast, air is a mixture because its components retain their individual properties. For instance, the oxygen in the air supports combustion.
• Energy changes occur during the formation of a compound, as seen in the formation of water. When you ignite a mixture of hydrogen and oxygen, it releases a significant amount of energy to form water. However, mixing the components of air does not cause any energy change, further proving that air is a mixture.

• The chemical formula of water is fixed as H₂O, whereas the variable composition of air prevents it from being represented by a single chemical formula.

Why is water a compound and air a mixture?

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Struggling to remember metal reactivity order? Learn the complete reactivity series with memory tricks, real-world examples, and practice questions – only at Selftution.com, the best platform for clear, exam-ready learning!

The reactivity series is a systematic arrangement of metals and nonmetals in order of their decreasing chemical activity.

Also known as the activity series, it helps predict how a substance will react with others, particularly in displacement reactions.

This series determines the ability of a metal to replace another in a reaction and indicates how it interacts with oxygen, water (cold, hot, or steam), and acids.

Chemistry revolves around chemical reactions, where substances interact to form new products through transformative processes.

Element reactivity varies based on their position in the reactivity series, a key concept for understanding chemical processes.

This post will explore the significance of the reactivity series, how it influences chemical reactions, and its practical applications in everyday life. Let’s dive in!

METAL REACTIVITY SERIES

Definition of metal reactivity series –

Metal reactivity series is a list in which metals are arranged in the decreasing order of their chemical activity.

Reactivity Series of Metals

The metal activity or reactivity series finds its utility in the study of displacement reactions. It helps us to predict whether a particular metal can displace another metal from a compound or not.

In a displacement reaction, a metal higher up in the reactivity series displaces all other metals in a compound that lies below it. For example, zinc being more active than copper replaces copper sulfate in the solution state to form zinc sulfate and free copper.

The activity of a metal depends upon its capability to lose electrons in the solution state to form positive ions. The more readily metal loses its electrons, the more active it is, and the higher up it is in the reactivity series. Among the most commonly known metals, the most active cesium is at the top and the least active platinum is at the bottom of the reactivity series.

Special Features of the Reactivity or Activity Series of Metals:

1. The ease with which a metal in solution loses electrons and forms a positive ion decreases down the series, i.e., from cesium to platinum.
2. Although a non-metal, hydrogen is included in the reactivity series of metals. It is because, like metals, it too loses an electron and becomes a positively charged (H⁺) ion.
3. Metals at the top of the reactivity series can displace metals that are placed lower from their salt solutions.
4. The activity series facilitates the comparative study of the metals in terms of the degree of their reactivity.
5. Metals at the top of the reactivity series are difficult to obtain from their ores.

The table below explains the reactions of metals with oxygen (air), water, dilute acids, and other salt solutions, at various levels of the reactivity series.

Reactivity of metals with oxygen, water, and acids

NONMETAL REACTIVITY OR ACTIVITY SERIES

Definition of reactivity of nonmetals–

Nonmetal activity series is a list in which nonmetals are arranged in the decreasing order of their reactivity.

Similar to metals, we can also arrange nonmetals in terms of their reactivity. During displacement reactions, a more active nonmetal displaces a less active nonmetal from a compound.

Reactivity Series of Nonmetals

The activity of a nonmetal depends upon its capability to gain electrons in the solution state to form positive ions. The more readily a nonmetal gains electrons, the more active it is, and the higher up it is in the reactivity series.

Frequently Asked Questions

Q1. Which is the most active or reactive metal?

Ans. Cesium and francium are the most reactive metals and are at the top of the reactivity series. However, the quantity of francium produced until now is too small. Therefore, for all practical purposes, we consider cesium as the most reactive metal.

Cesium reacts with water. Video by Periodic Videos

Q2. Which is the least active or reactive metal?

Ans. Platinum is one of the least reactive metals and therefore lies at the bottom of the reactivity series. It has remarkable corrosion resistance, even at high temperatures, and is therefore considered a noble metal. Due to its low reactivity, it found uses in the manufacture of laboratory equipment, electrodes, platinum resistance thermometers, dentistry equipment, and jewelry.

Q3. Which is the most active or reactive nonmetal?

Ans. Fluorine is the most active nonmetal, and it displaces all other nonmetals in salt solutions.

Q4. What determines the reactivity of metals?

Ans. We determine the reactivity of metals based on their capability to lose electrons in the solution state to form positive ions or cations. Therefore, the more readily metal loses its electrons, the more active it is, and the higher up it is in the reactivity series.

Q5. What determines the reactivity of nonmetals?

Ans. We determine the reactivity of nonmetals based on their capability to gain electrons in the solution state to form negatively charged ions or anions. Therefore, nonmetals that gain electrons rapidly are more active than others.

Knowledge of chemistry is endless, so to know more about various types of chemical reactions. Click here

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Learn the key differences between homogeneous and heterogeneous mixtures with simple explanations, real-world examples, and visual aids.

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So, let’s begin.

In chemistry, we group materials around us into pure substances and mixtures.

We further classify mixtures into homogeneous and heterogeneous mixtures, based on how their components are distributed.

These mixtures are all around us, from the air we breathe to the food we eat.

Homogeneous mixtures have a uniform composition, meaning their components are evenly distributed, while heterogeneous mixtures have visibly different parts.

Recognizing these differences helps us better understand the world of materials.

This post will explore the characteristics of homogeneous and heterogeneous mixtures with easy-to-understand examples.

Let’s uncover the science behind these mixtures!

Mixture Definition

A mixture is an impure substance created by combining two or more pure substances (elements or compounds) in any proportion, without undergoing any chemical change, thereby retaining their individual properties.

Most substances we encounter exist in the form of mixtures. The pure substances that combine to form a mixture are referred to as its components or constituents. Mixtures can be created by combining elements, compounds, or both. Depending on the physical state of its components, a mixture can exist in any of the three states of matter: solid, liquid, or gas. Since the components of a mixture do not undergo chemical reactions, they retain their individual properties and can be separated using various physical methods.

Homogeneous and Heterogeneous Mixture Definition and Examples

Examples of Mixture

Some common examples of mixtures in our daily lives include air, milk, fruit juice, medicines, honey, tap water, brass, and bronze.

Air is a mixture of oxygen, nitrogen, and carbon dioxide gases. It also contains water vapor, dust particles, and traces of inert gases.

Mixtures are classified into two types: homogeneous and heterogeneous mixtures. However, before we delve into studying them, let’s first understand some basic characteristics of mixtures.

Characteristics of Mixtures with Examples

In mixtures, the components are loosely combined and retain their individual properties. Based on this, mixtures exhibit the following characteristics:

• No Chemical Change:
  A mixture consists of two or more pure substances that coexist without undergoing any chemical combination. For example, in a mixture of common salt and sand, their particles are not chemically bonded. We can easily separate the salt by dissolving it in water and filtering out the sand.

• Variable Composition:
  The components of mixtures can vary in proportion, making it impossible to represent them with a chemical formula. For instance, you can prepare a mixture of sand and common salt by mixing 1 g of salt with 2 g of sand or 2 g of salt with 2 g of sand.

• No Specific Physical Properties:
  A mixture does not have fixed melting or boiling points. Its properties depend on the proportions of its components. For example, the boiling point of water increases as more common salt is dissolved in it.
• Homogeneity:
  Mixtures can be homogeneous or heterogeneous. For example, a mixture of salt and water is homogeneous, whereas a mixture of salt and sand is heterogeneous.
• Separation:
  Simple physical methods can separate the components of a mixture. For instance, evaporation can recover salt dissolved in water.
• Energy Change:
  Generally, no significant energy change (heat or light) occurs during the formation of mixtures. For example, when sand and salt are mixed, there is no release or absorption of energy.

HOMOGENEOUS AND HETEROGENEOUS MIXTURE

Based on their composition and distribution, we can divide mixtures into two types – homogeneous and heterogeneous mixtures.

Video courtesy eNotes

HOMOGENEOUS MIXTURE

Definition of the homogeneous mixture–

Homogeneous mixture is the one, in which the components are uniformly distributed throughout its volume and cannot be seen separately.

We should not confuse a homogeneous mixture with a compound. In compounds, atoms of various elements combine chemically in a fixed ratio. Whereas, in homogeneous mixtures, components of the mixture stay together in a varying ratio. For more differences between mixture and compound, click here.

Examples of Homogeneous Mixture-

A salt solution in water is an example of a homogeneous mixture, as we cannot see salt particles separately from water. We can prepare a salt solution by mixing one or two tablespoons in 1 liter of water. In both cases, the solution formed will be homogeneous, but the proportions of salt and water are not the same.

Some other examples of homogeneous mixtures are tap water,  milk, air, fruit juice, medicine, etc.

We can form a homogeneous mixture of metals by mixing two or more metals in the molten state. This homogeneous mixture of molten metals, on cooling, forms a solid mixture. We call this solid mixture an alloy. For example, brass is an alloy of copper and zinc. Similarly, bronze is an alloy of copper, zinc, and tin.

HETEROGENEOUS MIXTURE

Definition of the heterogeneous mixture-

A heterogeneous mixture is one in which the components are not uniformly distributed throughout its volume and can be easily distinguished as separate entities.

Examples of Heterogeneous Mixture

Most mixtures found in nature are heterogeneous. For instance, soil is a mixture containing hundreds of elements and compounds, and its composition varies from one location to another. Other examples of heterogeneous mixtures include rocks, a mixture of kerosene and water, and combinations like rice and pulses.

Examples of Homogeneous and Heterogeneous Mixtures

Examples of Homogeneous and Heterogeneous Mixtures

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Selftution.com is the best educational website for students, offering step-by-step guides to make chemistry easy. Learn decomposition, combustion, single & double displacement reactions with real-world examples.

Welcome to Selftution.com – Where complex science becomes simple! So, let’s begin.

Chemical reactions occur all around us, from burning wood to digesting food.

They involve the breaking and forming of chemical bonds, leading to new substances with different properties.

Types of Chemical Reactions

There are different types of chemical reactions, such as hydrolysis, oxidation, synthesis, reduction, combustion, electrolysis, and neutralization.

However, all chemical reactions can be grouped into four main categories:

1. Combination or synthesis reaction
2. Decomposition reaction
3. Displacement reaction, and
4. Double displacement or double decomposition reaction

Types of chemical reactions in chemistry with examples

Additionally, we can classify chemical reactions based on energy changes. Exothermic reactions release energy in the form of heat or light, whereas endothermic reactions absorb energy.

Before we study various types of reactions, it is important to know what a chemical reaction is.

WHAT IS A CHEMICAL REACTION?

A chemical reaction is a process in which two or more substances, when brought into close contact with each other, interact chemically to produce one or more new substances.

Definition of a chemical reaction –

A chemical reaction is a process in which two or more substances called reactants, react with each other to form one or more new substances called products. This may involve the evolution or absorption of energy.

Reactant (A) + Reactant (B) → Product (C) + Product (D)

COMBINATION OR SYNTHESIS CHEMICAL REACTIONS

A chemical reaction in which two or more substances combine to form a single new substance is called a combination or synthesis reaction. For example,

A + B → AB

In the above reaction, a combination of substances A & B (reactants) takes place to give a molecule of a new substance AB (product). It’s important to note that most of the synthesis reactions are exothermic.

There are three types of combination or synthesis chemical reactions:

1. Two elements combine to form a compound.
2. An element and a compound combine to give a new compound.
3. Two or more compounds combine to form a new compound.

Examples of all three types of combination or synthesis chemical reactions:

Types of synthesis or combination chemical reactions with examples

DECOMPOSITION CHEMICAL REACTIONS

A decomposition reaction is one in which a compound splits into simpler substances (elements or compounds) such that the products formed do not recombine to form the original compound. For example,

AB → C + D

Here, the decomposition of the molecule AB takes place to give two or more new substances, C & D. It may occur in the presence of heat or light or by the passage of an electric current. Therefore, decomposition reactions are of three types –

• Thermal decomposition,
• Photo decomposition or photolysis, and
• Electrolysis

Thermal decomposition reactions are those that involve heat. Photo decomposition or photolysis is the process that involves light. Electrolysis is due to the passing of an electric current.

Heat, light, and electricity are different forms of energy that cause the breaking of bonds in the molecules of reactants to yield simpler products. Therefore, decomposition reactions are endothermic.

Similar to combination reactions, based on the types of products formed, decomposition chemical reactions are also of three types:

1. A compound breaks up to form two or more elements.
2. Formation of both elements and compounds due to the breakup of a compound.
3. A compound breaks up to form two or more compounds.

Examples of all types of chemical decomposition reactions:

Types of decomposition chemical reactions with examples

DISPLACEMENT CHEMCIAL REACTIONS

In a displacement reaction, a more active element displaces a less active element from a compound. In such reactions, the active element replaces one element of the reactant molecule. For example,

AB + C → CB + A

Here, C displaces A from AB, since C is chemically more active as compared to A. A displacement reaction can be exothermic or endothermic in nature. We used the reactivity series of metals and nonmetals to identify the chemical activity of an element. In the reactivity series, more active metals are at the top and less active at the bottom.

There are three types of displacement reactions:

1. A more reactive metal displaces a less active metal from its salt solution.
2. A metal more reactive than hydrogen displaces hydrogen gas from an acid.
3. A more reactive non-metal displaces a less active non-metal from its salt solution.

Examples of types of chemical displacement reactions:

Types of displacement chemical reactions with examples

DOUBLE DISPLACEMENT OR DOUBLE DECOMPOSITION CHEMICAL REACTIONS

This is a type of chemical reaction in which two compounds in a solution state exchange their ions or radicals to form new compounds. For example,

AB + CD → CB + AD

Here, AB & CD are the two reactant molecules. They exchange their ions or radicals to form two new molecules, CB and AD. Double displacement or double decomposition reactions are of two types: precipitation reactions and neutralization reactions.

1. Precipitation reaction: A chemical reaction in which two compounds in their aqueous state react to form an insoluble solid (a precipitate) as one of the products is called a precipitation reaction.
2. Neutralization reaction: A chemical reaction in which an acid reacts with a base or an alkali to form a salt and water is called a neutralization reaction.

Examples of types of double displacement or double decomposition chemical reactions for precipitation and neutralization reactions:

Types of double displacement or decomposition chemical reactions with examples

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Learn the exact definition of pure substances in chemistry, key examples (elements & compounds), and how they differ from mixtures.

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In chemistry, the definition of a pure substance differs from how we use the word in everyday life.

Many things we think of as pure, like milk, honey, tap water, and fruit juice, are not pure substances according to chemistry.

We usually call something pure if it has no unwanted materials.

However, in chemistry, these are not pure substances, but mixtures made of different elements and compounds.

For example, milk is a mix of water, fats, proteins, and sugars, while honey is made up of different types of sugars and small amounts of other substances.

To learn more about mixtures, click here.

DEFINITION OF PURE SUBSTANCE IN CHEMISTRY

The definition of a pure substance in chemistry is as follows:

Pure substances consist of only one type of element or compound. They contain identical atoms or molecules throughout and display consistent physical and chemical properties.

Another definition of a pure substance is:

A pure substance is a homogenous material which contains atoms of one kind and has a definite set of physical and chemical properties. However, if a substance consists of two or more differenct kinds of atoms, combined together, then the proportion of these combining atoms by weight must remain constant.

PROPERTIES OF PURE SUBSTANCES

A pure substance exhibits the following properties:

• A pure substance has a fixed melting and boiling point.
• It has its characteristic taste, color, and odor.
• It is always homogeneous.
• We cannot break a pure substance into simpler substances by any physical means.

PURE SUBSTANCE EXAMPLES

According to the definitions above, all elements and compounds qualify as pure substances. Examples of pure substances include:

• Elements: Pure gold, pure silver, pure copper, and pure iron, which consist of only one type of atom.
• Compounds: Pure sugar and pure water, which consist of identical molecules. Although these molecules contain different types of atoms, the ratio of these atoms by weight always remains constant.

Pure Substance in chemistry Definition and examples

IMPURE SUBSTANCES WITH EXAMPLES

Impure substances are those that contain small amounts of other substances mixed with them. These are commonly referred to as mixtures. Many substances we consider “pure” in everyday life are not classified as pure substances in chemistry. For example:

• Tap Water: Clear-looking tap water is not a pure substance because it contains small amounts of dissolved minerals and gases. These dissolved minerals give water its taste; otherwise, pure water is tasteless.
• Fruit Juice: Fruit juice is an impure substance as it contains sugar, minerals, salts, and various organic compounds.

• Milk: Milk is not a pure substance because it consists of different compounds, including fats, carbohydrates, proteins, salts, vitamins, and water. Moreover, the proportions of these compounds vary from one cow to another.
• Honey: Honey is an impure substance as it contains various molecules of substances in addition to sugar.
• Brass: Brass, an alloy made by combining copper and zinc, is not considered a pure substance in chemistry. It contains atoms of two different elements that are not chemically bonded.

FREQUENTLY ASKED QUESTIONS

Question 1. Is air a pure substance?

Ans: Air is not a pure substance; it is a mixture of 21% oxygen, 78% nitrogen, and 0.3% carbon dioxide. The remaining 0.7% consists of inert gases like argon and traces of other gases. This composition varies from place to place due to climatic conditions and the levels of pollutants in the region. The composition of air also changes with altitude. At higher elevations, the percentage of oxygen decreases, which is why mountaineers carry oxygen cylinders. Additionally, air contains dust particles, and their concentration differs depending on the location.

Question 2. Is water a pure substance?

Ans: Water in its pure form, or distilled water, is a pure substance. Pure water is a compound made up of hydrogen and oxygen combined in a 2:1 ratio, respectively. However, clear-looking tap water is not a pure substance in chemistry because it contains small amounts of dissolved minerals and gases. These dissolved minerals give water its taste; otherwise, water would be tasteless.

Question 3. Is sugar a pure substance?

Ans: Yes. Sugar is a pure substance. Sugar in its pure form is a compound with the same kind of molecules. It contains carbon, hydrogen, and oxygen combined in a fixed proportion. Therefore, although it contains different kinds of atoms, the proportion of these combining atoms by weight is always constant.

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Struggling to tell them apart? Selftution.com – the #1 trusted educational website – breaks it down with clear definitions, side-by-side comparisons, and everyday examples to make learning effortless.

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In chemistry, all types of changes can broadly be classified into physical and chemical changes. In this post, we will learn about the difference between physical and chemical changes with examples-

• The characteristics of physical and chemical changes with examples,
• The differences between physical and chemical changes, and
• How to know that a chemical change has taken place?
• Is the burning of a candle a physical or chemical change?

PHYSICAL CHANGE

Physical change definition:

A physical change is a change in which no new substance is formed. The substance may change in shape, size, and physical state and appearance, but not in its chemical composition.

These changes may be temporary (reversible) or permanent (irreversible).

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CHARACTERISTICS OF PHYSICAL CHANGE

• No new substance is formed.
• There is no change in the composition of the substance.
• The properties of the substance remain the same.
• The change is temporary or reversible, but not always.
• There may be changes in size, shape, and state of substance during physical change.
• There may or may not be any exchange of energy during a physical change.

EXAMPLE OF PHYSICAL CHANGES

Some day-to-day examples of physical changes are – the drying of clothes, the ringing of a bell, the melting of ice cream, the melting of butter, the glowing of a bulb, the heating of electric iron, the beating of metals into sheets, the formation of clouds, cutting of the grass, boiling or freezing of water, sublimation of camphor, making of sugar solution, etc.

Examples of physical change.

Let us understand some of them in detail.

Example – 1: Interconversion of states of matter (Reversible physical change)

The interconversion of states of matter is a physical change. You might have observed that ice, when taken out of a refrigerator, melts and changes into water, and when put back in the refrigerator, it freezes into ice again. During the melting or freezing of water, no new substance is formed; only the state of water changes.

Similarly, on boiling, water changes into steam. On cooling, the steam again changes back into water. So again, there is only a change in the state of the water, and no change in the chemical composition of water occurs. 

The picture depicts the interconversion of states of matter with an example of water. Here, water, a form of matter, undergoes interconversion from the solid to the liquid state and from the liquid to the gaseous state on heating. Similarly, on cooling, vapor changes to a liquid and on further cooling to a solid.

On removing the cause of change, the substance returns to its original state. Thus, physical change is a reversible or temporary change.

Example 2: Crushing and making of chalk (Reversible physical change)

To observe a physical change with the help of chalk, crush some pieces of it into dust. Add a little water to the dust and make a thick paste of it. Roll this paste into the shape of a chalk piece, and let it dry. On drying, you will get back the chalk with its original properties.  Therefore, physical change is a reversible change.

Example 3: Sublimation of Ammonium Chloride (Reversible physical change)

Sublimation is the process of changing a solid directly into a gas on heating and vice versa without passing through a liquid state. To demonstrate the same, take some ammonium chloride in a test tube and heat it gently. You will observe that solids change into vapors. When the vapor touches the upper, cooler part of the test tube, it changes back to solid ammonium chloride. This shows that no new substance is formed. Thus, the sublimation of ammonium chloride is a physically reversible change.

However, physical change may not always be reversible. Some examples of irreversible physical change are –

Example – 4: Cutting of a carrot into pieces (Irreversible physical change)

Take a carrot and cut it into pieces. What do you observe? The shape and size of the carrot have changed. Now observe the color of the carrot pieces and taste one of them. There will be no difference in the taste of the whole carrot and its pieces.

Thus, physical change only affects the physical form of a substance, and the chemical composition of the final product remains identical to the original substance. However, in the present case, once the carrot is cut into pieces, we cannot obtain the original carrot. Thus, physical change may not always be reversible.

Example – 5: Cutting of paper into pieces (Irreversible physical change)

Take a piece of paper and cut it into four square pieces. Now, cut each square piece further into four square pieces. Lay them on a table so that the pieces acquire the shape of the original paper. Although the chemical composition of each piece of paper is similar to that of the original paper, you cannot join all the pieces back to make the original paper without cut marks. Thus, cutting paper is an irreversible physical change.

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CHEMICAL CHANGE

Chemical change definition:

A chemical change is a permanent change in which new substance(s) are formed. The new substance(s) shows composition and properties which are completly different from those of the original substance.

CHARACTERISTICS OF CHEMICAL CHANGE

• The change is permanent and irreversible.
• One or more new substances are formed.
• There is an exchange of energy during a chemical change. This means the heat or light or both might be given out or consumed.

EXAMPLES OF CHEMICAL CHANGES

Some day-to-day examples of chemical changes are – the rusting of iron, the ripening of fruits, the bursting of crackers, the souring of milk, the burning of paper or coal, the decomposition of organic matter, etc.

Examples of chemical change.

Let us understand some of them in detail.

Example 1: Burning of paper and coal (Irreversible chemical change)

Take a piece of paper and burn it. It turns into ash. This is a new substance, whose properties are different from those paper. The ash cannot change into paper again. Therefore, the burning of paper is a chemical change.

Coal is primarily made of carbon. On burning, coal reacts with oxygen present in the air to form carbon dioxide gas. We cannot change carbon dioxide back to carbon by any physical method. Therefore, the burning of coal is a chemical change.

Example 2: Rusting of iron (Irreversible chemical change)

Take an iron nail and bring it near a magnet. The magnet attracts the nail. Now leave the nail in tap water for a few days. You will notice that the nail develops a reddish-brown coating all over it. This coating is called rust. Rust is a hydrated oxide of iron, a new substance.  It forms due to a reaction of iron with oxygen dissolved in water.

Again, bring the nail near the magnet. But now the magnet does not attract nails toward itself. This shows that rust is a new substance that is non-magnetic. You cannot get iron back from rust with any physical method. Thus, rusting of iron is an irreversible and permanent chemical change.

Example – 3: Heating of sugar (Irreversible chemical change)

Put some sugar in a pan and heat it. You will notice that sugar first melts and then changes into a reddish-brown solid. On prolonged heating, it turns black and gets charred. Now stop heating and taste it on cooling. You will notice it is not sweet anymore. On heating, the sugar converts to black residue called charcoal, and some gaseous products like carbon dioxide and water.  Charcoal is a form of carbon, an element. You cannot get sugar back from charcoal with any physical method. Therefore, the heating of sugar is an irreversible and permanent chemical change.

HOW TO KNOW THAT A CHEMICAL CHANGE HAS TAKEN PLACE?

A chemical change is usually accompanied by a change in color, the evolution of gas, and the release or absorption of energy.

Color change: Ripening of fruits is a chemical change. The color of raw fruits changes during ripening. Heating sugar and burning paper gives a black substance.

Evolution of a gas: On the reaction of metals like zinc with dilute hydrochloric acid, the evolution of hydrogen gas takes place. Heating of sugar results in the evolution of carbon dioxide gas.

Release or absorption of energy: During a chemical change, energy is released or absorbed in the form of heat, light, or sound. An explosion of a cracker is a chemical change that produces heat, light, and sound energy.

DIFFERENCE BETWEEN PHYSICAL AND CHEMICAL CHANGE

Difference in physical and chemical changes

BURNING OF A CANDLE IS A PHYSICAL OR CHEMICAL CHANGE?

The burning of a candle is an example of a simultaneous physical and chemical change.

When a candle is lit, the wax melts and turns into a liquid state. As the molten wax drops on the floor or table, it solidifies again. Therefore, this is a physical change.

Simultaneously, some of the molten wax rises up the wick, turns into vapor, and burns with the flame to form two new substances: carbon dioxide and water vapor. The amount of wax that gets burned escapes in the form of gases, and the candle becomes smaller and smaller. This is a chemical change.

Thus, the melting of wax is a physical change, and the burning of wax is a chemical change.

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Discover the key differences between physical & chemical properties of metals, explained with clear definitions, comparisons, and real-world applications.

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Some basic physical and chemical properties of metals that differentiate them from non-metals and metalloids are as follows:

Physical Properties of Metals:

• Usually, metals are hard, opaque, possess luster, and conduct heat and electricity.
• They are malleable and ductile, can withstand longitudinal pull, and produce a resonant sound when struck.
• All metals are solid, except mercury and gallium, which are liquids at room temperature.
• Metals usually have high melting and boiling points.

Chemical Properties of metals are:

• Metals possess 1, 2, or 3 valence electrons in the outermost shell of their atoms. Therefore, they lose or donate valence electrons to form positively charged ions (cations).
• They react with cold water, steam, and dilute acids to give hydrogen.
• Chlorides of metals are true salts, and oxides of metals are usually basic.

INTRODUCTION

Before we proceed further with a detailed study of the physical and chemical properties of metals, let’s have a look at some common facts about them:

• There are three types of elements – metals, nonmetals, and metalloids. About 80% of the known elements in the periodic table are metals. To be precise, out of the 118 elements, 94 are metals, 17 are nonmetals, and the remaining 7 are metalloids. However, this number may vary since the nature of some elements is still under research.
• The usage of metals by mankind predates ancient times as far as 3500 B.C. Artifacts made of copper and tin were discovered at the excavation sites of the Indus Valley Civilization. The studies showed the use of iron by Asians in 2000 B.C. and later around 900-500 B.C. in China.

• Occurrence – Metals may occur in free and native states, for example, gold, silver, and platinum. However, most of them occur in a combined state as compounds along with earthly impurities. Metals are mined from the earth, but they are not offered in their pure form for direct use. It gives us in the form of ores. The ore is a mineral from which we can extract metal economically. The process of extracting metal from its ores is called metallurgy.

Ores of some common metals with chemical names and formulas

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PHYSICAL AND CHEMICAL PROPERTIES OF METALS

Although every element has a distinct set of physical and chemical properties, still based on some common features, we can still distinguish metals from nonmetals and metalloids. In this post, we will learn about various physical and chemical properties related to metals.

Physical Properties of Metals

Some common terms used for describing the physical properties of metals are – state, luster,  hardness, malleability, ductility, tensile strength, sonority, conductivity, density, corrosivity, etc.

Physical Properties of Metals

Based on the physical properties, the uses of certain common metals are described below:

The table depicts certain uses of common metals based on their physical properties.

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Chemical Properties of Metals

All metals react with common reagents like oxygen (air), hydrogen, halogens, sulfur, water, acids, etc. However, the extent of the reaction is different in the case of each metal.

Some of the chemical properties of metals are as follows:

• Metals react with oxygen present in the air to form oxides. These oxides dissolve in water to form an alkali.
• They form a range of salts when reacted with halogens (F, Cl, Br, I). For example, NaCl, KBr, etc.
• When reacted with cold water or steam, metals release hydrogen and produce alkali as a by-product.
• Metals react with dilute acids to produce hydrogen gas and a salt solution.

Metals possess 1, 2, or 3 valence electrons in the outermost shell of their atoms. Therefore, they lose or donate valence electrons to form positively charged ions (cations). The chemical properties of metals depend upon their capability to lose electrons in the solution state to form positive ions. The more readily a metal loses its electrons, the more active it is, and the higher up it is in the reactivity series.

Reactivity Series of Metals

Metal reactivity series is a list in which metals are arranged in the decreasing order of their chemical activity.

The metal activity or reactivity series finds its utility in the study of displacement reactions. It helps us to predict whether a particular metal can displace another metal from a compound or not. In a displacement reaction, a metal higher up in the reactivity series displaces all other metals in a compound, which lies below it. For example, zinc being more active than copper replaces it with copper sulfate in a solution state to form zinc sulfate and free copper.

To understand the chemical properties of various metals, refer to the table below, which explains the reactions of metals with oxygen (air), water, dilute acids, and other salt solutions.

Reactivity of metals with oxygen, water, and acids

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TYPES OF METALS

Based on different physical and chemical properties, there are four types of metals: Base metals, Ferrous metals, Noble or Free metals, and Heavy metals.

1. Base Metals – refer to metals that get oxidized easily. For example, iron, copper, nickel, aluminum, lead, zinc, tin, tungsten, molybdenum, tantalum, magnesium, cobalt, bismuth, cadmium, titanium, zirconium, antimony, manganese, beryllium, chromium, germanium, vanadium, gallium, hafnium, indium, niobium, rhenium, and thallium.
2. Ferrous Metals – derived from the Latin word meaning ‘containing iron’. For example, pure iron, carbon steel, stainless steel, cast iron, alloy steel, etc. 
3. Noble or Free Metals – unlike base metals, are resistant to corrosion and, due to their rarity, are considered precious. There are 9 noble or free metals, which are rhenium, ruthenium, rhodium, palladium, silver, osmium, iridium, platinum, and gold.
4. Heavy Metals possess a high density or atomic weight. Some examples of heavy metals are antimony, arsenic, bismuth, cadmium, cerium, chromium, cobalt, copper, gallium, gold, iron, lead, manganese, mercury, nickel, platinum, silver, tellurium, thallium, tin, uranium, vanadium, and zinc.

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Struggling to understand metals, nonmetals, and metalloids? Selftution.com – the #1 trusted educational website – breaks down their differences with clear comparisons, charts, and real-world applications. Master chemistry the easy way!

In the earlier post, Elements and Compounds, we learned that scientists have discovered a total of 118 elements so far. Among them, 94 elements are metals, 17 are nonmetals, and the remaining 7 elements are metalloids.

• Metals are generally shiny, good conductors of electricity, and malleable.
• Nonmetals tend to be dull, poor conductors, and brittle.
• Metalloids have properties that are intermediate between metals and nonmetals.

Understanding the Difference Between Metals, Nonmetals, and Metalloids is essential to grasping how elements behave in nature and their various uses.

Metalloids, for example, exhibit characteristics of both metals and nonmetals, making them unique in their properties.

In this blog, we will explore how these three categories differ in terms of physical and chemical properties, along with their practical applications in daily life.

The above demarcation between metals, nonmetals, and metalloids may vary since the nature of some elements is still under research.

Metals

Metals are hard, opaque, lustrous materials that are good conductors of heat and electricity. They are malleable and ductile, can withstand longitudinal pull, and produce a resonant sound when struck. They form basic oxides on reacting with the air and evolve hydrogen on reaction with water and dilute acids. Click here to know more about the physical and chemical properties of metal.

Nonmetals

Nonmetals are poor conductors of heat and electricity and are non-malleable and non-ductile. They do not exhibit luster and do not produce resonant sounds when struck. Nonmetals form acidic oxides when reacted with oxygen and do not evolve hydrogen when reacted with dilute acids. They are further divided into halogens and noble gases.

• Halogens: We derive the word halogen from  ‘hal‘, which means salt, and ‘gen,’ which means generation. Thus, halogens are salt producers, as they react with metals to form a range of salts. For example, NaCl, KBr, etc.
• Noble gases: These are 6 monoatomic gases, i.e., He, Ne, Ar, Kr, Xe, and Rn. They are chemically inert and almost inactive. Due to its inactive nature, we use argon in light bulbs to prevent the tungsten filament from oxidizing.

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Metalloids

Metalloids are those elements that show properties of both metals and nonmetals. They have physical and chemical properties intermediate between those of metals and nonmetals. Commonly recognized metalloids are Boron (B), Silicon (Si), Germanium (Ge), Arsenic (As), Antimony (Sb), and Tellurium (Te). The less commonly recognized metalloid is Polonium (Po). Some of the uses of metalloids are :

• Boron: used as an insecticide and in fire retardants.
• Silicon: Silicon gel is applied to burn patients to absorb moisture. Silicon, being a semiconductor of electricity, is extensively used for the manufacture of transistors in the electronics industry.
• Arsenic: It finds use for medicinal purposes.
• Antimony: Used as an antiparasitic drug.
• Germanium finds application along with silicon in cell phones as a semiconductor.

The difference in physical properties of Metals, Nonmetals, and Metalloids

The difference in Properties of Metals, Nonmetals, and Metalloids

1. State: Metals and metalloids are generally solid at room temperature, whereas nonmetals exist as liquids, gases, or brittle solids. Metalloids are brittle solids that crumble to powder when struck.
2. Luster: Metals are lustrous, i.e., show brightness, whereas nonmetals are non-lustrous. Metalloids exhibit metallic luster and may look like metals.
3. Nature: Most metals are malleable and ductile, except zinc and mercury. Nonmetals are mostly non-malleable and non-ductile. All metalloids are brittle solids; therefore, they are not used for structural purposes. For example, Silicon is a brittle solid, but it is lustrous and acts as a good conductor of electricity under certain conditions.
4. Tensile Strength: Metals have high tensile strength; therefore, they are extensively used for structural purposes. Whereas, nonmetals and metalloids have poor or low tensile strength. However, as an exception, carbon fibers made out of carbon, a nonmetal, have very high tensile strength.

1. Conduction of Heat/Electricity: Metals are good conductors of heat and electricity, whereas nonmetals are non or poor conductors. Metalloids are semiconductors of electricity and show average transmission of heat under certain temperature conditions.
2. Chemical properties: Metals lose or donate valence electrons and form cations, whereas nonmetals gain or accept electrons to form anions. Thus, metals and nonmetals combine through the ionic bond to form a compound. However, nonmetals also share electrons to form a compound through the covalent bond. Chemically, metalloids behave mostly as nonmetals. However, they can combine with metals to form an alloy.

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