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Welcome to these 9th class Chemistry notes on States of Matter and Phase Changes — Chapter 1 of the Chemistry with Anam course. Before we dive into atoms, bonding, and reactions, it helps to step back and ask a simple question: what is chemistry actually about, and what does matter look like around us? This chapter lays that foundation — covering the branches of chemistry, the different states matter can exist in, and how we classify the substances that make up the world.

1.1  What is Chemistry?

Chemistry is the branch of science that studies the properties, composition, and structure of substances, along with the physical and chemical changes they undergo. Because it's such a huge subject, chemistry is split into several specialised branches — physical, inorganic, organic, analytical, biochemistry, environmental, nuclear, polymer, geochemistry, medicinal, and astrochemistry, to name the major ones. Each branch lets scientists focus deeply on one aspect of how matter behaves, from designing new medicines to studying molecules found in space.

1.2  States of Matter

Matter is anything that has mass and occupies space. In everyday life we mainly encounter three states — solid, liquid, and gas — with plasma often called the fourth state, common in fluorescent tubes, lightning, and stars. What separates these states is how tightly the particles are packed and how strong the forces between them are.

Fig: Particle arrangement in solids, liquids, and gases — packing gets looser and movement gets freer as you move from solid to gas.

Solids have a fixed shape and volume because their particles are locked into position, able only to vibrate in place. Liquids have a fixed volume but take the shape of their container, since their particles stay close but keep moving past one another. Gases have neither fixed shape nor fixed volume — their particles are far apart and move freely, which is why gases are easily compressed while solids and liquids are not. Beyond these three, matter can also exist in intermediate and exotic states, such as liquid crystals, supercritical fluids, and Bose-Einstein condensates.

1.3  Elements, Compounds and Mixtures

All matter can be grouped into three categories:

• Element: the simplest form of matter — a pure substance made of only one kind of atom (e.g. iron, oxygen, gold).
• Compound: a pure substance formed when two or more elements chemically combine in a fixed ratio (e.g. water, carbon dioxide).
• Mixture: formed when substances are physically combined in any ratio, without a new chemical bond forming (e.g. air, rock, milk). Mixtures can be homogeneous (uniform throughout, like salt water) or heterogeneous (uneven, like a sample of rock).

1.4  Allotropic Forms of Substances

Some elements can exist in more than one structural form, called allotropes, each with different physical and chemical properties. Carbon is the classic example — it exists as diamond (an extremely hard, transparent crystal), graphite (soft, layered, and a good conductor), and buckminsterfullerene (a cage-like molecule of 60 carbon atoms). Oxygen also has two allotropes: ordinary oxygen (O₂) and ozone (O₃).

1.5  Solutions, Colloids and Suspensions

When one substance is mixed into another, the result falls into one of three categories, depending on particle size:

Fig: In a true solution particles are invisible and never settle; in a colloid they're small and stay spread out; in a suspension they're larger, visible, and settle over time.

• True solution: solute particles dissolve completely and pass through filter paper leaving no residue (e.g. salt in water).
• Colloidal solution: particles are bigger than in a true solution but still too small to see, and they don't settle over time (e.g. starch solution).
• Suspension: particles are large enough to see, don't dissolve, and settle down if left undisturbed (e.g. chalk in water).

1.6  Saturated Solutions and the Effect of Temperature on Solubility

An unsaturated solution can still dissolve more solute at a given temperature, while a saturated solution has dissolved the maximum amount possible — any extra solute added will simply settle at the bottom. The solubility of a substance is the amount that can dissolve in 100 g of solvent at a particular temperature, and it's different for every solute-solvent pair.

Temperature usually changes solubility: for most solids (like potassium nitrate), solubility rises as temperature increases, which is why crystallisation is often used to purify solids by dissolving them hot and letting pure crystals form as the solution cools. For gases, the opposite is generally true — solubility falls as temperature rises, which is why a fizzy drink goes flat faster when warm.

Key Takeaways — States of Matter and Phase Changes (Chapter 1)

• Chemistry studies the composition, properties, and changes of matter, and is divided into branches like physical, organic, and inorganic chemistry.
• The main states of matter — solid, liquid, gas, and plasma — differ in how closely particles are packed and how freely they move.
• Matter is classified as elements, compounds, or mixtures, and some elements exist as allotropes with different properties.
• Solutions, colloids, and suspensions differ mainly by particle size and behaviour on standing.
• Solubility depends on temperature — most solids dissolve better when hot, most gases dissolve better when cold.

Welcome to these 9th class Chemistry notes on Atomic Structure — Chapter 2 of the Chemistry with Anam course. All matter around us is built from atoms, yet elements can look and behave completely differently from one another. Compare a lump of iron ore to a piece of gold jewellery — iron is dense and dull, while gold is soft, shiny, and easy to shape. These differences all trace back to what's happening inside the atoms themselves. This chapter walks through what atoms are made of, how scientists discovered that, and the key numbers chemists use to describe them — atomic number, mass number, and isotopes.

2.1  Structure of the Atom

For a long time, atoms were believed to be indivisible — the smallest possible pieces of matter. That idea changed once scientists began experimenting with electricity passed through gases at low pressure inside sealed glass tubes.

• Electron: Glowing rays inside these tubes (cathode rays) were shown to be streams of tiny, negatively charged particles — the electron.
• Proton: A second set of experiments revealed positively charged particles travelling the opposite way, later named protons.
• Neutron: Much later, a third, chargeless particle of similar mass to the proton was discovered and named the neutron.
• Nucleus: Rutherford's famous gold-foil experiment then showed that all this positive charge and nearly all the mass of an atom is packed into one tiny central region — the nucleus — while electrons occupy a much larger space around it, arranged in energy levels called shells (K, L, M, N...).

Fig: Bohr model of a Sodium (Na) atom — nucleus at the centre with electrons arranged as K = 2, L = 8, M = 1.

2.2  Atomic Number and Mass Number

Two numbers identify every atom:

• Atomic Number (Z): the number of protons in the nucleus — unique to each element and the basis for its position on the periodic table.
• Mass Number (A): the total number of protons and neutrons in the nucleus. Since electrons are so light, they aren't counted. The number of neutrons can always be found as N = A − Z.

Fig: Working out protons, neutrons and electrons for Carbon-12 using its atomic number and mass number.

2.3  Isotopes and Their Masses

Atoms of the same element always have the same atomic number, but they can have different numbers of neutrons — these variants are called isotopes. Because isotopes have the same number of electrons, they behave almost identically in chemical reactions, even though their masses differ.

Fig: The three isotopes of hydrogen — same proton count, different neutron count (purple = proton, grey = neutron, gold = electron).

Radioactive isotopes: Some isotopes have unstable nuclei and release energy as radiation — this is radioactivity. Radioactive isotopes have real-world uses, including dating ancient organic remains (radiocarbon dating) and helping doctors image organs during diagnosis.

Relative atomic mass: Since most elements exist as a mixture of isotopes, chemists calculate a weighted average mass based on each isotope's natural abundance, measured against carbon-12 as the reference standard.

Key Takeaways — Atomic Structure (Chapter 2)

• Atoms are made of protons, neutrons and electrons — protons and neutrons sit in the nucleus, electrons occupy the surrounding shells.
• Atomic number (Z) = number of protons; Mass number (A) = protons + neutrons.
• Isotopes have the same atomic number but different mass numbers.
• Radioactive isotopes are used in medicine, dating, and energy applications.
• Want to explore this visually? Use the Explore Periodic Table tool on the home page to see any element's electron shells and configuration.