Chapter 4 — Atomic Structure

Early thinking and why experiments mattered: Long ago, people wondered whether matter could be divided forever or whether there was a smallest unit. Some thinkers suggested tiny, indivisible particles and called them atoms. These early ideas were not based on experiments but on reasoning. Later, scientists began to test ideas by observing how substances combined and by using instruments to study electricity and gases. Experimental evidence replaced pure speculation and made the particle idea a scientific model of matter.

Evidence from careful observations: For example, chemical reactions often occur in fixed proportions. When two elements combine the mass ratios are fixed, which suggests discrete particles combining in simple numbers. Gases showed predictable behaviour in pressure and temperature experiments, also pointing to particles moving and colliding. Electricity experiments revealed particles with charge leaving and entering materials. Such experimental facts were crucial in forming a modern view of atoms.

The particle model as a useful tool: The particle model gives a clear way to think about solids, liquids and gases and about mixing, dissolving and chemical change. It explains why a sugar cube dissolves into invisible particles in water, why heating can increase motion of particles and cause expansion, and why reactions sometimes involve rearrangement of atoms rather than creation of new atoms. For Class 7, the particle model is a practical picture: atoms and molecules are tiny units that combine and separate according to simple rules.

Limits and progress: It is important to know that early models were improved over time. A good model must match observations and help make predictions. In later classes you will learn more precise models, but for now concentrate on using the particle idea to explain everyday phenomena and simple chemical changes.

How charged particles were found: Experiments with electric discharge tubes showed streams of tiny particles moving from a negatively charged electrode. These particles were found to be much lighter than atoms and carry a negative charge; they were named electrons. The discovery of electrons showed that atoms were not indivisible but had smaller parts.

Finding positive particles: Other experiments that measured the behaviour of gases and ions revealed a positive particle inside the atom. Studies of hydrogen and of positive rays led scientists to recognise the proton, a particle with positive charge and a mass similar to that of a hydrogen atom. The proton determines the chemical identity of an element because it sets the positive charge in the nucleus.

Why neutrons were proposed: Measurements of atomic mass showed that some atoms were heavier than would be expected from protons alone. To account for this extra mass without extra charge, a neutral particle was proposed and later detected: the neutron. Neutrons sit in the nucleus alongside protons and add mass but no electrical charge.

Relative sizes and charges: For Class 7 it is enough to know that electrons are very light and negatively charged, while protons are heavy (compared to electrons) and positively charged, and neutrons are about as heavy as protons but neutral. The nucleus (protons + neutrons) contains most of the atom's mass, and electrons around it determine chemical behaviour. This picture helps explain why chemical reactions involve electrons while the nucleus stays unchanged in ordinary chemistry.

The nucleus and why it is important: The nucleus is the tiny central part of an atom that holds protons and neutrons. Although it is extremely small compared to the whole atom, it contains almost all the mass. Protons inside the nucleus carry positive charge; neutrons carry no charge. The number of protons gives the element its identity and is written as atomic number.

Electron cloud and shells: Electrons move in regions around the nucleus called shells or energy levels. These shells are often drawn as rings for simplicity. Electrons are negatively charged and much lighter than protons and neutrons. Because electrons are spread over a much larger volume than the nucleus, most of the atom’s volume is empty space. Electrons in different shells have different energies and the inner shells fill before outer shells.

Scale and comparison: If an atom were enlarged to the size of a stadium, the nucleus would be like a small marble at the centre and the electrons would be like tiny moving insects in the stands. This image helps to grasp the idea of mass concentrated in the nucleus and the large empty space around it.

Why structure explains chemistry: The separation of massive nucleus and light electrons explains why chemical reactions mainly involve electrons and why nuclear particles are unchanged in ordinary chemistry. The shell arrangement is the basis for bonding rules: atoms with nearly full or nearly empty outer shells tend to lose, gain or share electrons to achieve stability. For Class 7, practise sketching atoms with labelled nucleus and shells to make this structure clear.

Definitions and notation: The atomic number (Z) is the number of protons in the nucleus and defines the element. The mass number (A) is the total number of protons and neutrons. We often write an atom using nuclear notation: A over Z followed by the element symbol, for example 14/6 C for a carbon atom with mass number 14 and atomic number 6.

How to use these numbers: From Z you immediately know how many protons are present. For a neutral atom, the number of electrons equals Z. To find the number of neutrons, subtract: neutrons = A − Z. This is a simple calculation that tells you the composition of the nucleus. For example, 23/11 Na has 11 protons, 11 electrons (neutral) and 23 − 11 = 12 neutrons.

Why the distinction matters: Atomic number determines chemical properties because it fixes the positive charge in the nucleus and thus the number of electrons surrounding it in a neutral atom. Mass number helps to tell how heavy an atom is and to recognise different isotopes of the same element. Many exam questions ask for these straightforward calculations and for drawing or naming atoms using nuclear notation.

Practical classroom tips: When given a nuclear notation, always list protons, neutrons and electrons separately and explain whether the atom is neutral or an ion. Practice with several examples until you can do the subtraction quickly and without error. Remember that electrons hardly affect mass number, which is why the mass number counts only protons and neutrons.

What electron shells are: Electrons occupy energy levels called shells around the nucleus. In class 7 we label shells as K (1st), L (2nd), M (3rd) and so on, or simply 1, 2, 3. Electrons fill the lowest energy shells first. The arrangement of electrons into shells determines many chemical properties.

How many electrons fit in each shell (class 7 rule): For our level we use the simple capacities: 1st shell (K) up to 2 electrons, 2nd shell (L) up to 8 electrons, 3rd shell (M) up to 8 electrons (we do not need to use the higher number 18 at this stage). Electrons fill shells in order: first K, then L, then M. For example, an atom with 11 electrons will have arrangement 2,8,1.

How to write electron arrangements: Write the number of electrons in each shell separated by commas. This practice quickly shows the number of valence electrons (those in the outermost shell). For instance oxygen (Z = 8) has 2,6 so it has 6 valence electrons; sodium (Z = 11) has 2,8,1 so it has 1 valence electron.

Why this matters for bonding: Atoms with incomplete outer shells tend to gain, lose or share electrons to reach a stable arrangement. Counting electrons in shells is the first step to deciding whether an atom will form an ion or share electrons in covalent bonds. Practise drawing shells and placing dots or small crosses to represent electrons so you become comfortable with the filling order and how valence electrons are identified.

Valence electrons defined: Valence electrons are the electrons in the outermost occupied shell of an atom. These are the electrons that are most easily lost, gained or shared in chemical reactions because they are furthest from the attractive pull of the nucleus.

Stability and the full-shell tendency: Atoms tend to be more stable when their outer shell is full. Noble gases like helium (2), neon (2,8) and argon (2,8,8) have full outer shells and are mostly unreactive. Other elements often change their electron numbers to achieve the same stable configuration, either by losing electrons to empty the outer shell or by gaining electrons to fill it, or by sharing electrons to reach a full shell together.

Predicting behaviour from valence count: Count the outer electrons to predict common behaviour: elements with one valence electron (like sodium) usually lose it to form +1 ions. Those with seven valence electrons (like chlorine) usually gain one to form −1 ions. Atoms with four or five valence electrons often share electrons to form covalent bonds instead of forming ions. This simple rule helps explain why some elements are metals and others are non-metals in terms of electron behaviour.

Examples and classroom practice: Practice by listing valence electrons for common elements and saying whether they will likely lose, gain, or share electrons. Use simple diagrams showing before and after electron counts to explain why sodium becomes Na+ and chlorine becomes Cl−, and why two hydrogen atoms share a pair to make H2. Remember this is a simple model useful for many predictions at this stage.

What an ion is: An ion is an atom (or group of atoms) that has a net electrical charge because it has lost or gained electrons. Protons remain in the nucleus, so changing the number of electrons gives a net positive or negative charge. Ions are common in salts and many solutions.

Forming cations and anions: If an atom loses one or more electrons it becomes a positive ion called a cation. For example a neutral sodium atom (11 protons, 11 electrons) loses one electron and becomes Na+ with 11 protons and 10 electrons. If an atom gains electrons it becomes a negative ion called an anion; chlorine gains one electron to become Cl− with 17 protons and 18 electrons.

Calculating the charge: The charge on an ion equals the difference between the number of protons and the number of electrons. For example, if an atom has 12 protons and 10 electrons the net charge is +2 and the ion is written as 12/?? with 2+ (Mg2+). Always check whether the particle described is neutral or charged by comparing protons and electrons.

Notation and importance: Ion charges are shown as superscripts after the element symbol, for example Na+ or O2−. Ionic charges must be balanced when writing formulas of compounds. Many physical properties, like conductivity in solution, depend on the presence of mobile ions. At Class 7, practise forming simple ions from electron arrangements and writing their symbols with charges.

Basic idea of ionic bonding: Ionic compounds form when atoms transfer electrons so that each achieves a more stable electron arrangement. After the transfer, the positively charged cation and negatively charged anion attract each other by electrostatic force and form a solid ionic lattice. This is typical when metals react with non-metals.

Steps to write a formula: First, identify the ions each element will form using valence electrons. For example sodium (2,8,1) becomes Na+ and chlorine (2,8,7) becomes Cl−. Second, balance the total positive and negative charges so the compound is neutral. For Na+ and Cl− a 1:1 ratio balances charges, giving NaCl. For calcium (Ca2+) and chloride (Cl−), two chlorides are needed for each calcium, giving CaCl2. Always write the formula with the cation first and the anion second.

Worked approach and examples: When you meet a new pair of elements, write their likely ions and their charges. Then ask: how many of each ion are needed so that the sum of positive charges equals the sum of negative charges? For example, aluminium often forms Al3+ and oxide forms O2−. To balance Al3+ and O2− you need two Al3+ (total +6) and three O2− (total −6), giving Al2O3. Practise finding the simplest whole-number ratio that balances charges.

Naming and properties: Name the compound by stating the cation (metal) followed by the anion with an −ide ending for simple non-metals: sodium chloride, calcium chloride, magnesium oxide. Ionic solids are usually hard, brittle and have high melting points; they do not conduct electricity as solids but do when molten or dissolved because ions can move freely. For Class 7, work through a variety of metal + non-metal pairs until balancing becomes routine.

What covalent bonding means: Covalent bonding occurs when two non-metal atoms share one or more pairs of electrons so that each atom counts the shared electrons toward a full outer shell. This sharing forms molecules where atoms are held together by shared electron pairs. Covalent bonds are common between non-metal atoms such as hydrogen, oxygen, nitrogen and carbon.

Simple shared pair rules: In class 7 we focus on single covalent bonds formed by sharing one pair of electrons. The shared pair is often shown as a line between atom symbols (H—H) or as a pair of dots or crosses between symbols in dot-and-cross diagrams. For example, two hydrogen atoms, each with one electron, share one pair to make H2. In water, oxygen shares one electron with each of two hydrogen atoms to form H—O—H.

How to draw and name covalent molecules: Draw the outer electron shells and place shared pairs between atoms. Count electrons to ensure each atom (usually) reaches a full outer shell according to the simple rules used in this class. Names of simple molecular compounds are learned separately and often use prefixes for number of atoms in higher classes, but simple names like water (H2O) and ammonia (NH3) should be memorised.

Properties and differences from ionic compounds: Covalent substances can be gases, liquids or solids with lower melting points than ionic solids. They do not conduct electricity when dissolved because they do not produce free ions. At Class 7, practise drawing dot-and-cross diagrams and simple structural formulas to become familiar with molecule shapes and shared pairs.

Defining isotopes: Isotopes are atoms of the same element that have the same number of protons but different numbers of neutrons. Because protons define the element, isotopes have the same chemical behaviour but different masses. For example carbon-12 and carbon-14 are both carbon because each has six protons, but carbon-14 has two extra neutrons.

Writing isotopes and calculating neutrons: We write isotopes as mass number over atomic number with the element symbol, for example 14/6 C for carbon-14. To find neutrons subtract: neutrons = A − Z. Thus carbon-14 has 14 − 6 = 8 neutrons while carbon-12 has 12 − 6 = 6 neutrons.

Relative atomic mass (simple idea): Natural elements often consist of a mixture of isotopes. Relative atomic mass (often shown on the periodic table) is the weighted average mass of all the isotopes compared to carbon-12, which is defined as exactly 12. Because it is an average of isotopes, relative atomic mass is frequently a decimal number (for example chlorine ~35.5) even though mass numbers of single isotopes are whole numbers.

Uses of isotopes at a basic level: Some isotopes are useful in medicine, industry and archaeology. Radioactive isotopes can be used to trace movements inside the body or to date ancient organic materials (radiocarbon dating with carbon-14). Stable isotopes can help identify sources of substances in environmental studies. For Class 7 the focus is on recognising isotopes, calculating neutrons and understanding the cause of non-integer atomic masses on tables.

Why models are used: Atoms are too small to see directly, so scientists build simple models—pictures and rules—to explain observations and predict behaviour. The nucleus-and-shell model, and dot-and-cross diagrams for bonding, are examples used here to help students visualise particles and bonds.

What models show and what they do not: Models show where particles are typically found (nucleus, shells) and how electrons may be shared or transferred. However, they are simplified: electrons do not orbit like planets and their exact behaviour follows quantum rules learned later. Dot-and-cross diagrams do not show bond angles or real electron motion; they only show electron pairing and transfer in a helpful way.

Using structure to explain properties: The number and arrangement of electrons explain many simple properties. Metals with few valence electrons lose them easily and conduct electricity; ionic compounds made of ions have high melting points and conduct when molten due to mobile charges; covalent molecules may be gases or liquids with lower melting points. Recognising valence electrons and likely ions helps predict formulas and behaviour in simple reactions.

Scientific thinking and progress: Be critical with models: they are tools that change when new evidence appears. As you learn more, your models will become more sophisticated. For Class 7, practise describing what a model successfully explains and name at least one limitation. This habit of questioning and refining ideas is a key part of learning science.

What to review: For successful revision, make a short list of common elements with their symbols and typical electron arrangements: H (1), C (2,4), N (2,5), O (2,6), Na (2,8,1), Mg (2,8,2), Cl (2,8,7), Ca (2,8,8). Practice reading nuclear notation to find protons, neutrons and electrons and practise writing simple ionic and covalent formulas by balancing charges or showing shared pairs.

Practice tasks: Draw shell diagrams for at least five elements, convert nuclear notation like 23/11 Na into protons, neutrons and electrons, and write the formulas and names of ionic compounds such as NaCl, CaCl2, MgO and covalent molecules such as H2O and NH3. Explain why a sodium atom becomes Na+ and why chlorine becomes Cl− using shell diagrams.

Exam skill tips: Many school questions ask for labeled diagrams, calculations of neutrons, drawing of electron arrangements and writing formulas. Show working clearly: write Z and A, do the subtraction A − Z for neutrons, write electron arrangement step by step, and balance charges when making ionic formulas. Use clear labels and neat diagrams to present answers.

Further practice and improvement: After you complete simple examples, try mixed questions under timed conditions to build speed. Check each answer by counting protons and electrons and by ensuring charge balance in ionic compounds. Create a one-page revision sheet with element symbols, typical ions, and electron arrangements; use it for quick daily review. Also practise explaining answers in words and in diagrams—both are commonly required in exams. Regular short practice sessions are more effective than long cramming and will help you remember key facts and methods for tests and future chemistry classes.