Atoms And Molecules

Hello students, welcome to today's science class. I am so happy to see you all here, ready to learn about one of the most fundamental concepts in chemistry — atoms and molecules. This chapter is extremely important because it forms the foundation for understanding how matter behaves and how different substances are formed. So let's begin our journey into the tiny world of particles that make up everything around us.

Students, have you ever thought about what matter is made of? Since ancient times, philosophers in India and Greece wondered about this very question. Around 500 BC, an Indian philosopher named Maharishi Kanad thought deeply about this. He said that if we keep dividing any piece of matter, we will eventually reach a point where we cannot divide it any further. He called these smallest particles "Parmanu." Another Indian philosopher, Pakudha Katyayama, further elaborated on this idea and said that these particles usually exist in combined forms, giving us various kinds of matter.

At around the same time, Greek philosophers Democritus and Leucippus also thought about this. They said that if we keep dividing matter, we will reach a stage where the particles cannot be divided any further. Democritus called these indivisible particles "atoms," which literally means "cannot be cut" or "indivisible." However, students, all these ideas were based on philosophical thinking. There was no experimental proof available at that time because science was not advanced enough.

It was only by the end of the eighteenth century that scientists began to understand the difference between elements and compounds. They became curious about how and why elements combine with each other, and what happens when they combine. This is when two very important scientists made significant contributions.

Antoine L. Lavoisier, a French chemist, laid the foundation of chemical sciences by establishing two important laws of chemical combination. He conducted many experiments and, along with another scientist Joseph L. Proust, helped establish these laws. Let me explain these laws to you in detail.

The first law is called the Law of Conservation of Mass. Students, let me ask you a question. When a chemical reaction takes place, does the total mass of the substances change? Think about it. If you burn a piece of paper, it seems to disappear, right? But actually, the ash that remains plus the gases produced all together have the same mass as the original paper. This is what the Law of Conservation of Mass states — mass can neither be created nor destroyed in a chemical reaction.

Let me describe an activity that demonstrates this law. Imagine we take two solutions. Let's say we take copper sulphate solution in one container and sodium carbonate solution in another. We prepare a 5% solution of each in water. Now, we take some sodium carbonate solution in a conical flask and put some copper sulphate solution in a small ignition tube. We hang this ignition tube carefully inside the flask, making sure the solutions do not mix. We put a cork on the flask and weigh it carefully. Now, we tilt and swirl the flask so that the two solutions mix together. A chemical reaction takes place, and we can see a precipitate forming. Now, if we weigh the flask again, we will find that the mass remains exactly the same. This is because no mass is lost or gained during the reaction. The cork is put on the flask to ensure that no gases escape, which could otherwise change the mass. So students, always remember — in any chemical reaction, the total mass of reactants equals the total mass of products.

Now, the second law is called the Law of Constant Proportions, also known as the Law of Definite Proportions. Students, this law says that in a chemical compound, the elements are always present in a fixed proportion by mass, no matter where the compound comes from or how it is prepared.

For example, water always has hydrogen and oxygen in the ratio 1:8 by mass. Whether you take water from a river, from rain, or make it in a laboratory, when you decompose 9 grams of water, you will always get 1 gram of hydrogen and 8 grams of oxygen. Similarly, in ammonia, nitrogen and hydrogen are always present in the ratio 14:3 by mass, regardless of the source. This law was stated by Proust as, "In a chemical substance the elements are always present in definite proportions by mass."

Now students, you might be wondering — why do these laws exist? What is the reason behind them? This is exactly what a British chemist named John Dalton thought about. He provided the basic theory about the nature of matter, and his theory explained these laws beautifully.

John Dalton was born in 1766 in England in a poor weaver's family. Can you imagine? He began his career as a teacher at the age of just twelve! Seven years later, he became a school principal. In 1793, he moved to Manchester to teach mathematics, physics, and chemistry at a college. He spent most of his life there teaching and researching. In 1808, he presented his atomic theory, which was a revolutionary turning point in the study of matter.

According to Dalton's atomic theory, all matter — whether an element, a compound, or a mixture — is made up of small particles called atoms. Let me tell you the main postulates of this theory.

First, all matter is made of very tiny particles called atoms, which participate in chemical reactions. Second, atoms are indivisible particles, which cannot be created or destroyed in a chemical reaction. This postulate explains the Law of Conservation of Mass. Third, atoms of a given element are identical in mass and chemical properties. Fourth, atoms of different elements have different masses and chemical properties. Fifth, atoms combine in the ratio of small whole numbers to form compounds. And sixth, the relative number and kinds of atoms are constant in a given compound. This last postulate explains the Law of Definite Proportions.

So students, Dalton's atomic theory was a breakthrough because it gave a scientific basis to the laws of chemical combination that we discussed earlier. You will learn in the next chapter that actually atoms are made up of even smaller particles, but for now, we will consider atoms as the basic building blocks of matter.

Now let's look at some questions based on what we have learned so far.

Question 1: In a reaction, 5.3 grams of sodium carbonate reacted with 6 grams of acetic acid. The products were 2.2 grams of carbon dioxide, 0.9 grams of water, and 8.2 grams of sodium acetate. Show that these observations are in agreement with the law of conservation of mass.

Students, let's check this. The reaction is: sodium carbonate + acetic acid → sodium acetate + carbon dioxide + water.

The total mass of reactants is 5.3 + 6 = 11.3 grams. The total mass of products is 8.2 + 2.2 + 0.9 = 11.3 grams. Since the mass of reactants equals the mass of products, this demonstrates the Law of Conservation of Mass.

Question 2: Hydrogen and oxygen combine in the ratio of 1:8 by mass to form water. What mass of oxygen gas would be required to react completely with 3 grams of hydrogen gas?

Students, since the ratio is 1:8, for every 1 gram of hydrogen, we need 8 grams of oxygen. So for 3 grams of hydrogen, we need 3 × 8 = 24 grams of oxygen.

Question 3: Which postulate of Dalton's atomic theory is the result of the law of conservation of mass?

The answer is the second postulate — atoms are indivisible particles, which cannot be created or destroyed in a chemical reaction.

Question 4: Which postulate of Dalton's atomic theory can explain the law of definite proportions?

This is the sixth postulate — the relative number and kinds of atoms are constant in a given compound.

Great! Now let's move on to the next section — What is an Atom?

Students, have you ever seen a mason building a wall? The wall is made of bricks, and many walls make a room, and many rooms make a building. The building block of the building is a brick. Similarly, think about an ant-hill — it is made of tiny grains of sand. In the same way, atoms are the building blocks of all matter. Everything around you — this table, the air, your body, everything — is made up of atoms.

Now, let's talk about how we represent atoms or elements using symbols. Dalton was the first scientist to use symbols for elements in a specific sense. When he used a symbol, he meant one atom of that element. Later, a scientist named Berzilius suggested that the symbols of elements should be made from one or two letters of the name of the element.

Now students, here is something very important. The first letter of a symbol is always written as a capital letter, and the second letter as a small letter. For example, hydrogen is written as H, not h. Aluminium is written as Al, not AL. Cobalt is written as Co, not CO. Notice the difference — Co means cobalt, but CO would mean carbon and oxygen together! So we must be very careful with capitalization.

Some symbols are derived from the first letter and a letter appearing later in the name. For example, chlorine is Cl, zinc is Zn. And some symbols have been taken from Latin, German, or Greek names. For example, iron is Fe from the Latin name "ferrum," sodium is Na from "natrium," and potassium is K from "kalium." Each element has a unique chemical symbol.

One nanometre is one-billionth of a metre. That is 1 divided by 10^9 metres, or 1 nm equals 10^-9 metres. Equivalently, 1 metre equals 10^9 nanometres.

Let me give you some examples of relative sizes. The radius of a hydrogen atom is about 10^-10 metres. A water molecule is about 10^-9 metres. A haemoglobin molecule is about 10^-8 metres. A grain of sand is about 10^-4 metres. An ant is about 10^-3 metres. An apple is about 10^-1 metres. So you can see how incredibly small atoms are! We cannot see them with our naked eyes. But even though they are so small, they are the basis of everything in our world. Through modern techniques, we can now produce magnified images of surfaces of elements showing atoms.

Now let's discuss atomic mass. Students, one of the most remarkable concepts in Dalton's theory was that each element has a characteristic atomic mass. The theory could explain the law of constant proportions so well that scientists were prompted to measure the atomic mass of atoms.

Since determining the mass of an individual atom was difficult, scientists determined relative atomic masses using the laws of chemical combinations. Let me explain with an example.

Consider carbon monoxide, which is formed by carbon and oxygen. It was observed experimentally that 3 grams of carbon combines with 4 grams of oxygen to form carbon monoxide. In other words, carbon combines with 4/3 times its mass of oxygen.

Scientists initially took 1/16th of the mass of an atom of naturally occurring oxygen as the unit. This was because oxygen reacted with many elements and formed compounds, and this unit gave masses of most elements as whole numbers. However, in 1961, carbon-12 isotope was chosen as the standard reference for measuring atomic masses. One atomic mass unit is equal to exactly 1/12th the mass of one carbon-12 atom. This is called the unified mass, denoted by "u."

Let me give you an easy analogy. Imagine a fruit seller who wants to sell fruits but has no standard weight. He takes a watermelon and says this has a mass of 12 units. He then makes twelve equal pieces of the watermelon and finds the mass of each fruit he is selling relative to one piece of the watermelon. This is similar to how we define atomic mass units.

Here is a table of atomic masses of some elements that you should remember. Hydrogen is 1 u, carbon is 12 u, nitrogen is 14 u, oxygen is 16 u, sodium is 23 u, magnesium is 24 u, sulphur is 32 u, chlorine is 35.5 u, and calcium is 40 u. Notice that chlorine has a decimal atomic mass — this is because chlorine exists in different isotopic forms, and 35.5 is the average atomic mass.

Now, let's answer the in-text questions.

Question 1: Define the atomic mass unit.

Students, the atomic mass unit is defined as a mass unit equal to exactly 1/12th the mass of one carbon-12 atom.

Question 2: Why is it not possible to see an atom with naked eyes?

Because atoms are extremely small in size. Their radius is about 10^-10 metres, which is too small to be seen without powerful microscopes.

Now let's move on to the next section — What is a Molecule?

Students, a molecule is generally a group of two or more atoms that are chemically bonded together, that is, tightly held together by attractive forces. A molecule can be defined as the smallest particle of an element or a compound that is capable of an independent existence and shows all the properties of that substance. Atoms of the same element or of different elements can join together to form molecules.

Now let's discuss molecules of elements. The molecules of an element are constituted by the same type of atoms. Some elements like argon and helium have molecules made up of only one atom — these are called monoatomic molecules. But most non-metals have more than one atom in their molecules. For example, oxygen has two atoms bonded together, so it is called a diatomic molecule, written as O2. If three atoms of oxygen combine, we get ozone, written as O3. The number of atoms constituting a molecule is known as its atomicity.

Let me give you a table of atomicity for some elements. Argon and helium are monoatomic. Oxygen, hydrogen, nitrogen, and chlorine are diatomic. Phosphorus is tetra-atomic, meaning it has four atoms in one molecule. Sulphur is poly-atomic, meaning it has many atoms, usually eight in its most common form, written as S8.

Now, molecules of compounds are different. Atoms of different elements join together in definite proportions to form molecules of compounds. For example, water has hydrogen and oxygen in the ratio 2:1 by number of atoms, which corresponds to 1:8 by mass. Ammonia has nitrogen and hydrogen in the ratio 1:3 by number of atoms, which is 14:3 by mass. Carbon dioxide has carbon and oxygen in the ratio 1:2 by number of atoms, which is 3:8 by mass.

Let me show you how to find the ratio by number of atoms from the ratio by mass, using water as an example. For water, the ratio by mass of hydrogen to oxygen is 1:8. The atomic mass of hydrogen is 1 u, and oxygen is 16 u. So we divide the mass ratio by the atomic mass: for hydrogen, 1 divided by 1 equals 1; for oxygen, 8 divided by 16 equals 1/2. So the simplest ratio is 2:1. That means in one water molecule, there are 2 hydrogen atoms and 1 oxygen atom. This is why the formula of water is H2O.

Now, students, let's talk about ions. Compounds composed of metals and non-metals contain charged species. These charged species are known as ions. An ion may consist of a single charged atom or a group of atoms that have a net charge on them. An ion can be negatively or positively charged. A negatively charged ion is called an anion, and a positively charged ion is called a cation.

For example, sodium chloride, which is common table salt, has positively charged sodium ions, written as Na+, and negatively charged chloride ions, written as Cl-. A group of atoms carrying a charge is known as a polyatomic ion. For example, ammonium is NH4+, hydroxide is OH-, nitrate is NO3-, carbonate is CO3^2-, and sulphate is SO4^2-.

Now, let's discuss valency. Students, the combining power or capacity of an element is known as its valency. Valency can be used to find out how the atoms of an element will combine with atoms of another element to form a chemical compound.

Let me give you an interesting analogy. Think of valency as the number of arms an entity has. Human beings have two arms. An octopus has eight arms. If one octopus wants to catch hold of some people such that all eight arms of the octopus and both arms of all the humans are locked, how many humans can the octopus hold? If we represent the octopus as O and humans as H, we would need four humans to use all eight arms. So the formula would be OH4. The subscript 4 indicates the number of humans held by the octopus.

Similarly, valency tells us how many atoms of one element will combine with atoms of another element. Here is a table of some common ions and their valencies. Sodium is Na+ with valency 1, potassium is K+ with valency 1, magnesium is Mg2+ with valency 2, calcium is Ca2+ with valency 2, aluminium is Al3+ with valency 3, chloride is Cl- with valency 1, oxide is O2- with valency 2, and so on. Some elements can show more than one valency, and in such cases, a Roman numeral is used to show the valency. For example, iron can be Fe2+ or Fe3+, written as iron(II) or iron(III).

Now let's learn about writing chemical formulae. The chemical formula of a compound is a symbolic representation of its composition. Let me tell you the rules for writing chemical formulae.

First, the valencies or charges on the ion must balance. Second, when a compound consists of a metal and a non-metal, the name or symbol of the metal is written first. For example, calcium oxide is CaO, sodium chloride is NaCl, iron sulphide is FeS, and copper oxide is CuO. The non-metals are written on the right. Third, in compounds formed with polyatomic ions, if there is more than one polyatomic ion, we enclose the formula of the ion in brackets and write the number outside. For example, magnesium hydroxide is Mg(OH)2. If there is only one polyatomic ion, brackets are not needed, like in NaOH.

Now let's practice writing some formulae.

Example 1: Hydrogen chloride. Symbol H, Cl. Valency is 1 for both. So the formula is HCl.

Example 2: Hydrogen sulphide. Symbol H, S. Valency is 1 for hydrogen and 2 for sulphur. We criss-cross the valencies, so we get H2S.

Example 3: Carbon tetrachloride. Symbol C, Cl. Valency is 4 for carbon and 1 for chlorine. So we get CCl4.

Example 4: Magnesium chloride. Symbol Mg, Cl. Charge is 2+ for magnesium and 1- for chloride. We criss-cross the charges, so we get MgCl2. This means there are two chloride ions for every magnesium ion.

Example 5: Aluminium oxide. Symbol Al, O. Charge is 3+ for aluminium and 2- for oxygen. We criss-cross: Al2O3.

Example 6: Calcium oxide. Symbol Ca, O. Both have charge 2+. If we criss-cross, we get Ca2O2, but we simplify this to CaO because the numbers cancel out.

Example 7: Sodium nitrate. Symbol Na, NO3. Charge is 1+ for sodium and 1- for nitrate. So we get NaNO3. No brackets are needed because there is only one nitrate ion.

Example 8: Calcium hydroxide. Symbol Ca, OH. Charge is 2+ for calcium and 1- for hydroxide. We criss-cross to get Ca(OH)2. Notice the brackets around OH because there are two hydroxide ions.

Example 9: Sodium carbonate. Symbol Na, CO3. Charge is 1+ for sodium and 2- for carbonate. We criss-cross to get Na2CO3. No brackets are needed here.

Example 10: Ammonium sulphate. Symbol NH4, SO4. Charge is 1+ for ammonium and 2- for sulphate. We criss-cross to get (NH4)2SO4. Notice the brackets around NH4 because there are two ammonium ions.

Now let's answer the in-text questions.

Question 1: Write down the formulae of sodium oxide, aluminium chloride, sodium sulphide, and magnesium hydroxide.

Students, let's work through each one.

For sodium oxide: Na and O. Charge is 1+ for sodium and 2- for oxygen. Criss-cross: Na2O.

For aluminium chloride: Al and Cl. Charge is 3+ for aluminium and 1- for chloride. Criss-cross: AlCl3.

For sodium sulphide: Na and S. Charge is 1+ for sodium and 2- for sulphur. Criss-cross: Na2S.

For magnesium hydroxide: Mg and OH. Charge is 2+ for magnesium and 1- for hydroxide. Criss-cross: Mg(OH)2.

Question 2: Write down the names of compounds represented by the following formulae: Al2(SO4)3, CaCl2, K2SO4, KNO3, CaCO3.

Let's identify each:

Al2(SO4)3 is aluminium sulphate. CaCl2 is calcium chloride. K2SO4 is potassium sulphate. KNO3 is potassium nitrate. CaCO3 is calcium carbonate.

Question 3: What is meant by the term chemical formula?

The chemical formula of a compound is a symbolic representation of its composition, showing the constituent elements and the number of atoms of each combining element.

Question 4: How many atoms are present in an H2S molecule and a PO4^3- ion?

In H2S, there are 2 hydrogen atoms and 1 sulphur atom, so total 3 atoms. In PO4^3-, there is 1 phosphorus atom and 4 oxygen atoms, so total 5 atoms.

Now let's move on to molecular mass. Students, we already discussed atomic mass. Now we can extend this to calculate molecular masses. The molecular mass of a substance is the sum of the atomic masses of all the atoms in a molecule of the substance. It is the relative mass of a molecule expressed in atomic mass units.

Example 3.1(a): Calculate the relative molecular mass of water, H2O.

Atomic mass of hydrogen is 1 u, oxygen is 16 u. Water has 2 hydrogen atoms and 1 oxygen atom. So molecular mass = 2 × 1 + 1 × 16 = 2 + 16 = 18 u.

Example 3.1(b): Calculate the molecular mass of HNO3.

Molecular mass = atomic mass of H + atomic mass of N + 3 × atomic mass of O = 1 + 14 + 48 = 63 u.

Now, students, there is also something called formula unit mass. This is used for ionic compounds whose constituent particles are ions rather than molecules. The formula unit mass is calculated in the same way as molecular mass, but we use the term "formula unit" instead of "molecule."

For example, sodium chloride has formula unit NaCl. Its formula unit mass = atomic mass of Na + atomic mass of Cl = 23 + 35.5 = 58.5 u.

Example 3.2: Calculate the formula unit mass of CaCl2.

Atomic mass of Ca is 40 u. Atomic mass of Cl is 35.5 u. So formula unit mass = 40 + 2 × 35.5 = 40 + 71 = 111 u.

Now let's answer the questions in this section.

Question 1: Calculate the molecular masses of H2, O2, Cl2, CO2, CH4, C2H6, C2H4, NH3, CH3OH.

Let's calculate each one:

H2: 2 × 1 = 2 u O2: 2 × 16 = 32 u Cl2: 2 × 35.5 = 71 u CO2: 12 + 2 × 16 = 12 + 32 = 44 u CH4: 12 + 4 × 1 = 12 + 4 = 16 u C2H6: 2 × 12 + 6 × 1 = 24 + 6 = 30 u C2H4: 2 × 12 + 4 × 1 = 24 + 4 = 28 u NH3: 14 + 3 × 1 = 14 + 3 = 17 u CH3OH: 12 + 3 × 1 + 16 + 1 = 12 + 3 + 16 + 1 = 32 u

Question 2: Calculate the formula unit masses of ZnO, Na2O, K2CO3, given atomic masses of Zn = 65 u, Na = 23 u, K = 39 u, C = 12 u, O = 16 u.

ZnO: 65 + 16 = 81 u Na2O: 2 × 23 + 16 = 46 + 16 = 62 u K2CO3: 2 × 39 + 12 + 3 × 16 = 78 + 12 + 48 = 138 u

Now students, we have covered all the concepts in the chapter. Let's solve the exercises at the end of the chapter.

Exercise 1: A 0.24 g sample of compound of oxygen and boron was found by analysis to contain 0.096 g of boron and 0.144 g of oxygen. Calculate the percentage composition of the compound by weight.

Students, percentage composition means what percentage of the total mass is each element.

Total mass of compound = 0.24 g Mass of boron = 0.096 g Mass of oxygen = 0.144 g

Percentage of boron = (0.096 / 0.24) × 100 = 0.4 × 100 = 40% Percentage of oxygen = (0.144 / 0.24) × 100 = 0.6 × 100 = 60%

So the compound contains 40% boron and 60% oxygen by mass.

Exercise 2: When 3.0 g of carbon is burnt in 8.00 g oxygen, 11.00 g of carbon dioxide is produced. What mass of carbon dioxide will be formed when 3.00 g of carbon is burnt in 50.00 g of oxygen? Which law of chemical combination will govern your answer?

Students, let's first find out how much carbon dioxide is formed when 3 g of carbon reacts with 8 g of oxygen. We get 11 g of carbon dioxide. This tells us that 3 g of carbon reacts with 8 g of oxygen to form 11 g of carbon dioxide.

Now, if we have 3 g of carbon but 50 g of oxygen, what happens? The carbon will completely react with only 8 g of oxygen (because that's all that is needed), and the remaining 42 g of oxygen will be left unused. So the mass of carbon dioxide produced will still be 11 g.

This answer is governed by the Law of Conservation of Mass, because the mass of carbon dioxide produced depends on the amount of carbon that reacts, and excess oxygen does not create more product.

Exercise 3: What are polyatomic ions? Give examples.

Students, polyatomic ions are clusters of atoms that carry a fixed charge on them. They behave as a single unit. Examples include ammonium (NH4+), hydroxide (OH-), nitrate (NO3-), carbonate (CO3^2-), sulphate (SO4^2-), and phosphate (PO4^3-).

Exercise 4: Write the chemical formulae of the following: magnesium chloride, calcium oxide, copper nitrate, aluminium chloride, calcium carbonate.

Let's work through each:

Magnesium chloride: Mg and Cl. Charge is 2+ for magnesium and 1- for chloride. Formula: MgCl2.

Calcium oxide: Ca and O. Both have charge 2+. Formula: CaO.

Copper nitrate: Cu and NO3. Charge is 2+ for copper and 1- for nitrate. Formula: Cu(NO3)2. Note: copper usually has valency 2 in this compound.

Aluminium chloride: Al and Cl. Charge is 3+ for aluminium and 1- for chloride. Formula: AlCl3.

Calcium carbonate: Ca and CO3. Charge is 2+ for calcium and 2- for carbonate. Formula: CaCO3.

Exercise 5: Give the names of the elements present in the following compounds: quick lime, hydrogen bromide, baking powder, potassium sulphate.

Let's identify the elements:

Quick lime is calcium oxide, so it contains calcium and oxygen. Hydrogen bromide is HBr, so it contains hydrogen and bromine. Baking powder is sodium bicarbonate, NaHCO3, so it contains sodium, hydrogen, carbon, and oxygen. Potassium sulphate is K2SO4, so it contains potassium, sulphur, and oxygen.

Exercise 6: Calculate the molar mass of the following substances: ethyne (C2H2), sulphur molecule (S8), phosphorus molecule (P4), hydrochloric acid (HCl), nitric acid (HNO3).

Note: Students, the term "molar mass" here essentially means molecular mass or formula unit mass.

(a) Ethyne, C2H2: 2 × 12 + 2 × 1 = 24 + 2 = 26 u

(b) Sulphur molecule, S8: 8 × 32 = 256 u

(c) Phosphorus molecule, P4: 4 × 31 = 124 u

(d) Hydrochloric acid, HCl: 1 + 35.5 = 36.5 u

(e) Nitric acid, HNO3: 1 + 14 + 3 × 16 = 1 + 14 + 48 = 63 u

Now students, we have covered all the exercises. Let me also explain the group activity about writing formulae. This is a fun activity where you can learn to write chemical formulae by using placards with symbols and valencies. You criss-cross the valencies to get the formula. For example, for sodium sulphate, Na+ and SO4^2-, you need two sodium ions for one sulphate ion, so the formula is Na2SO4. Similarly, for sodium phosphate, Na+ and PO4^3-, you need three sodium ions for one phosphate ion, so the formula is Na3PO4.

Now let me give you a summary of everything we have learned in this chapter.

Students, in this chapter, we learned about atoms and molecules, which are the building blocks of matter. We started with the historical perspective — how ancient Indian philosophers like Maharishi Kanad and Greek philosophers like Democritus thought about the divisibility of matter. We then discussed the two important laws of chemical combination: the Law of Conservation of Mass, which states that mass is neither created nor destroyed in a chemical reaction, and the Law of Constant Proportions, which states that elements in a compound are always present in a fixed proportion by mass.

We learned about Dalton's atomic theory, which provided explanations for these laws. According to Dalton, all matter is made of atoms, atoms are indivisible, atoms of the same element are identical, atoms of different elements are different, atoms combine in simple whole number ratios to form compounds, and compounds have a fixed number of atoms of each element.

We then discussed atoms in detail — their symbols, their extremely small size, and atomic mass. We learned that atomic mass is measured in atomic mass units, where one atomic mass unit is 1/12th the mass of a carbon-12 atom.

We moved on to molecules — the smallest particles of elements or compounds that can exist independently. We learned about monoatomic, diatomic, polyatomic molecules, and the concept of atomicity.

We discussed ions — charged particles that are formed when atoms lose or gain electrons. We learned about cations, anions, and polyatomic ions.

We learned about valency and how to write chemical formulae by criss-crossing the valencies or charges of ions. We practiced writing formulae for various compounds.

We also learned about molecular mass and formula unit mass — how to calculate the mass of molecules and formula units by adding up the atomic masses of constituent atoms.

Finally, we solved all the exercises, which tested our understanding of percentage composition, laws of chemical combination, polyatomic ions, chemical formulae, elements in compounds, and calculation of molecular masses.

This is a very important chapter that forms the foundation for understanding chemistry. Make sure you practice writing chemical formulae and calculating molecular masses regularly. Remember, atoms and molecules are everywhere around us, and understanding them helps us understand the world better.

Thank you for your attention, students. Keep learning, keep exploring, and see you in the next class!