Chapter 4 — Atomic Structure and Chemical bonding

Introduction and history
The idea that matter is made of small particles goes back to ancient thinking, but scientific models of the atom developed through experiments and careful reasoning. Early models were attempts to explain measurable phenomena such as gas laws, fixed mass ratios in compounds and electrical observations. As experiments improved, models were revised to accommodate new evidence.

Dalton's atomic theory
John Dalton proposed that matter is composed of indivisible atoms of each element, that atoms of an element are identical in mass and properties, and that compounds form by combination of atoms in fixed ratios. Dalton's ideas explained conservation of mass and definite proportion laws for many chemical reactions and became the first scientific atomic theory used by chemists.

Thomson's discovery of the electron and its model
Cathode ray experiments showed the existence of electrons—tiny negatively charged particles present in atoms. J. J. Thomson proposed that atoms were not indivisible but contained electrons embedded in a diffuse positive medium to keep the atom neutral. This ‘plum pudding’ idea pictured electrons like plums in a pudding of positive charge. The model explained presence of negative particles and neutrality but lacked a concentrated centre that could deflect fast particles strongly.

Rutherford’s gold foil experiment
Rutherford, using alpha particles directed at very thin gold foil, observed that most particles passed through but some were deflected at large angles and a few bounced back. He concluded that atoms have a tiny, dense, positively charged nucleus containing most of the mass, with electrons outside. This nuclear model explained the scattering but raised a new problem: classical physics suggested orbiting electrons should lose energy and spiral into the nucleus, which does not happen.

Bohr’s quantum orbits
Niels Bohr introduced the idea that electrons occupy discrete energy levels or orbits and do not radiate energy while in a fixed orbit. Electrons can move between levels by absorbing or emitting quanta of energy (photons), explaining spectral lines of hydrogen. Bohr’s model bridged classical and quantum ideas and helped calculate energy levels for simple atoms, although it could not explain complex spectra or multi-electron atoms fully.

Progress to modern view
These historical models are important stepping stones. Each model incorporated new experimental facts and improved our understanding. The modern atomic theory uses quantum mechanics to describe electrons as probability clouds (orbitals) rather than fixed orbits. At Class 9 level the progression—Dalton → Thomson → Rutherford → Bohr—teaches how science updates ideas with evidence, and students learn why the nucleus and quantised energy levels are central to chemistry.

Overview of subatomic particles
Atoms are made of three primary subatomic particles: electrons, protons and neutrons. Each type has distinct charge, mass and location within the atom and together they determine chemical and physical properties. Understanding these particles is essential to explain atomic behaviour, bonding and chemical reactions.

Electrons
Electrons are negatively charged particles that exist in regions around the nucleus called shells or energy levels. They are very light: an electron’s mass is about 1/1836 of a proton. Electrons determine chemical reactivity because chemical bonds involve outer (valence) electrons. Electrons can be removed to form positive ions or added to form negative ions, processes that underlie many chemical reactions.

Protons
Protons are positively charged particles found in the nucleus. Each proton has a relative mass near 1 atomic mass unit (u). The number of protons in the nucleus—the atomic number Z—identifies the element and determines many chemical properties because it sets the number of electrons in a neutral atom. For example, an atom with 6 protons is carbon, always.

Neutrons
Neutrons are neutral particles also located in the nucleus and have mass similar to protons (≈1 u). Neutrons add to the mass of the atom and contribute to nuclear stability. Different numbers of neutrons create isotopes of the same element; some isotopes are stable while others are radioactive. Chemically isotopes behave similarly because chemical reactions depend mainly on electrons, not neutrons.

Relative masses and charges
In simple terms, assign relative masses: proton ≈ 1 u, neutron ≈ 1 u, electron ≈ 0 u (very small). Charges: proton = +1, electron = −1, neutron = 0. A neutral atom has equal numbers of protons and electrons, balancing their charges.

How particles were discovered
Electrons were found by cathode ray studies, protons were identified through positive particles in gas discharge and nuclear experiments, and neutrons were detected by observing particles with mass but no charge emitted from certain nuclei. Modern instruments like mass spectrometers and cloud chambers allow measurement of particle masses and behaviour and confirm their roles in atoms.

Chemical implications
Protons set element identity and electron number; electrons determine bonding and chemical reactivity; neutrons influence isotopic mass and nuclear properties. For class-level problems, knowing where each particle is and how to count them is key: protons = Z, neutrons = A − Z, electrons = Z in neutral atoms, or adjusted for ions.

Atomic number (Z)
Atomic number is the number of protons in the nucleus of an atom and is unique for each element. It determines the element’s identity and, for a neutral atom, the number of electrons as well. In chemical notation the atomic number is usually not written explicitly but it is fundamental in the periodic table: elements are arranged by increasing Z.

Mass number (A)
Mass number is the sum of protons and neutrons in the nucleus. It gives the total number of nucleons and is always an integer. Mass number is used to distinguish different isotopes of an element and is included in nuclear notation as A_Z X where X is the element symbol.

Isotopes explained
Isotopes are atoms of the same element (same Z) that differ in neutron number and therefore mass number. For example, carbon-12 (12_6C) and carbon-14 (14_6C) both have Z = 6 but differ in A. Isotopes have nearly identical chemical behaviour because chemistry depends on electrons, but their physical properties (mass, density) and nuclear behaviour (radioactivity) may differ. Isotopes are important in dating (e.g., carbon-14) and medical imaging (radioisotopes).

Stable and radioactive isotopes
Some isotopes are stable; others are unstable and decay by emitting radiation, transforming into other elements or isotopes. Radioactive decay processes (alpha, beta, gamma) change either the number of protons or the energy of the nucleus. Class 9 requires awareness that isotopes exist and that some are radioactive, but not detailed nuclear physics.

Ions and their formation
Ions are charged species formed when atoms gain or lose electrons. A cation is positive (lost electrons), an anion is negative (gained electrons). For example, Na → Na+ + e− and Cl + e− → Cl−. Ionic charges are important to write formulas and to predict compound stoichiometry. For a neutral compound, total positive and negative charges must balance.

Notations and examples
An atom is often written as A_Z X where A is mass number and Z is atomic number. For instance, 35_17Cl indicates chlorine with 17 protons and 18 neutrons. An ion is shown with its charge, e.g., Na+ or SO4^2−. Counting particles: neutrons = A − Z; electrons in an ion = Z − charge (for cation) or Z + charge magnitude (for anion).

Average atomic mass
Natural elements are often mixtures of isotopes; the relative atomic mass shown in the periodic table is a weighted average of isotopic masses. For many calculations at Class 9, mass numbers are used as integers while the concept of average atomic mass explains why table values are not whole numbers.

What are electron shells?
Electrons occupy defined energy levels around the nucleus called shells. These shells are often labelled K, L, M, N or by the principal quantum number n = 1, 2, 3, 4, ... Each shell can hold a limited number of electrons and electrons fill shells in order of increasing energy. The arrangement of electrons among shells explains chemical behavior such as bonding and valency.

Maximum electrons in shells
The capacity of a shell is given by 2n^2 where n is the shell number: K (n=1) holds up to 2 electrons, L (n=2) up to 8, M (n=3) up to 18, and N (n=4) up to 32. For many lighter elements we commonly use 2, 8, 8 as practical distributions for K, L, M when explaining bonding and valence. The detailed orbital filling order (s, p, d) is covered later in higher classes.

Electronic configuration notation
The electronic configuration states how many electrons are in each shell. For example, sodium (atomic number 11) has distribution 2, 8, 1 meaning 2 electrons in K, 8 in L and 1 in M. This notation is simple and useful for elements up to calcium (Z = 20). Another way is orbital notation (1s2 2s2 2p6 ...), but for Class 9 the shell distribution 2, 8, 1 is sufficient and easier to remember for common elements.

Filling order and the Aufbau idea
Electrons occupy lower energy shells first (Aufbau principle). Within a shell, sublevels have specific capacities and energies. For simple class-level problems, follow the rule of filling K first, then L, then M, and so on, until you place all electrons. This explains periodic repetition of properties: when outer shells have similar electron counts, elements display similar chemistry.

Valence electrons and chemical behaviour
The electrons in the outermost shell are valence electrons; they determine how an atom bonds. Atoms seek more stable arrangements often resembling noble gases with full outer shells. Atoms may lose, gain or share electrons to achieve a stable outer configuration. For example, chlorine has valence configuration 7 (2,8,7) and tends to gain one electron to complete its octet and form Cl−.

Writing configurations for common elements
Practice by using atomic numbers: carbon (6) → 2,4; oxygen (8) → 2,6; neon (10) → 2,8; magnesium (12) → 2,8,2; calcium (20) → 2,8,8,2. These help predict valency and bonding patterns. Remember to count electrons carefully and to identify the outermost shell for valence count.

Importance and limitations
Electronic configuration bridges atomic structure and bonding. It helps explain periodic trends and chemical reactivity. At higher levels, quantum mechanics refines the picture with orbitals and spin, but the shell model is a strong foundation for Class 9 chemistry.

Understanding valency
Valency is the combining capacity of an atom, usually equal to the number of electrons an atom will lose, gain or share to obtain a stable outer shell. For many main-group elements this means reaching an octet (eight electrons) or a duet for hydrogen. Valency is a practical concept used to predict how atoms form compounds and what the simplest formula of a compound will be.

Predicting valency from electronic configuration
Valency can be found by inspecting the outer-shell electrons. For example, sodium (2,8,1) has one valence electron and tends to lose it forming Na+ (valency 1). Oxygen (2,6) has six valence electrons and commonly gains two to form O2− (valency 2). Carbon (2,4) has four valence electrons and typically forms four covalent bonds to reach an octet, giving valency 4.

Group trends in the periodic table
Elements in the same group have the same number of valence electrons and therefore similar valencies and chemical behaviour. For instance, group 1 elements all have one valence electron and commonly form +1 ions. Across a period from left to right the number of valence electrons increases by one with each element, changing typical valency and resulting chemical behaviour.

Octet rule and exceptions
Many main-group elements follow the octet rule, striving to have eight electrons in their valence shell. However, there are exceptions: hydrogen follows a duet rule (2 electrons), boron and beryllium may be stable with fewer than eight electrons in some compounds, and heavier elements (from period 3 onward) can sometimes expand their octet using d-orbitals. Class 9 focuses on common octet cases and simple exceptions like hydrogen and boron.

Positive and negative valency
Valency can appear as positive when an atom tends to lose electrons (metals), and negative when it tends to gain electrons (non-metals). For example, sodium shows valency +1 while chlorine shows valency −1. When writing chemical formulas we use absolute values to balance atoms: Na+ and Cl− combine in a 1:1 ratio to form NaCl.

Multiple valencies
Certain elements, especially transition metals and some p-block elements, can exhibit variable valency (e.g., iron can be +2 or +3). This happens because electrons from different shells may be involved in bonding. For Class 9, it is enough to recognise that some elements can show more than one valency and to use the correct valency when writing formulas.

Applications
Valency helps in writing chemical formulas, predicting product compositions in simple reactions, and understanding periodic trends. Practise using valency with common elements to gain fluency in writing formulas and balancing simple chemical equations.

What is ionic bonding?
Ionic bonding is the electrostatic attraction between positively and negatively charged ions that results when atoms transfer electrons. It commonly occurs between metals, which lose electrons, and non-metals, which gain electrons. The transfer leads to the formation of cations and anions which then attract each other to form ionic compounds.

Mechanism of ion formation
Consider sodium and chlorine as a typical example. A sodium atom has configuration 2,8,1 and can become more stable by losing one electron to form Na+. Chlorine has 2,8,7 and can gain one electron to become Cl− and reach an octet. The electron transfer is driven by the tendency to achieve stable electronic configuration and by the favourable energy change when oppositely charged ions form a lattice.

Structure of ionic solids
Ionic compounds form extended three-dimensional crystalline lattices rather than isolated ion pairs. In a lattice each ion is surrounded by several ions of opposite charge in a repeating pattern. The lattice arrangement maximises attractions and minimises repulsions, which explains high structural stability and distinct crystal shapes often observed in ionic salts.

Properties of ionic compounds
Because ions are held strongly in a lattice, ionic solids typically have high melting and boiling points: large energy is required to separate ions. Ionic solids are generally hard and brittle; when a crystal is struck, like-charged ions can be forced next to each other causing repulsion and cleavage. Ionic compounds do not conduct electricity in solid form because ions are fixed, but they conduct when molten or dissolved in water because ions become free to move and carry charge.

Solubility and polarity
Many ionic compounds dissolve in polar solvents like water because solvent molecules stabilise separated ions via ion–dipole interactions. Solubility varies with lattice enthalpy and hydration enthalpy: if hydration energy overcomes lattice energy, the salt dissolves. Non-polar solvents fail to solvate ions effectively, so ionic compounds are usually insoluble in such media.

Writing ionic formulas
To write formulas balance total positive and negative charges. For Mg2+ and Cl−, two chloride ions balance one magnesium yielding MgCl2. Polyatomic ions (e.g., NO3−, SO4^2−) act as single charged units and are balanced similarly; when more than one polyatomic ion is needed, use parentheses in the formula, e.g., Ca(NO3)2.

Limitations and covalent character
Real bonds may show partial covalent character if small highly charged cations polarise large anions. This modifies some properties, but the ionic model remains a useful first approximation for many common salts in Class 9 chemistry.

Definition and basic idea
Covalent bonding forms when two atoms share pairs of electrons so that both achieve a more stable electronic arrangement. Sharing allows atoms to effectively fill their outer shells without full transfer of electrons. Covalent bonds commonly occur between non-metal atoms and lead to discrete molecules or network structures depending on bonding patterns.

Single, double and triple bonds
A single covalent bond is one shared pair of electrons between two atoms (represented A–B or as a pair of dots). Two shared pairs form a double bond (A=B) and three shared pairs form a triple bond (A≡B). Increasing bond order generally increases bond strength and decreases bond length: triple bonds are shortest and strongest of the three types, while single bonds are longest and weakest among them. Bond energy values reflect how much energy is needed to break a bond.

Lewis structures for covalent molecules
Lewis dot structures show valence electrons and help visualise covalent bonding. Shared pairs between atoms appear as lines or paired dots, while lone pairs remain on atoms. Using Lewis structures, you can ensure atoms satisfy the octet (or duet for hydrogen) and predict where multiple bonds are needed to fulfill electron counts. For example, CO2 is drawn as O=C=O to give each atom an octet.

Polar and non-polar covalent bonds
If the atoms involved have different electronegativities, the shared electrons are pulled closer to the more electronegative atom, creating a polar covalent bond with partial charges (δ+ and δ−). When identical atoms bond (e.g., H–H, Cl–Cl), sharing is equal and the bond is non-polar. Polar covalent bonds contribute to molecular polarity and affect properties like solubility and intermolecular interactions.

Molecular properties and states
Covalent substances can be gases, liquids or solids. Simple covalent molecules often have low melting and boiling points because molecules are held by weaker intermolecular forces, not by covalent bonds between molecules. However network covalent solids (e.g., diamond) consist of strong covalent bonds throughout and have very high melting points and hardness.

Bond formation energetics
Forming covalent bonds releases energy and breaking them requires energy. Bond energies can be used qualitatively to explain why some reactions are exothermic: new bonds formed have greater total energy release than energy required to break existing bonds. For Class 9 students, understanding that sharing leads to stable molecules and how polarity affects properties is the key takeaway.

Nature of metallic bonding
Metallic bonding is a model where metal atoms release some of their valence electrons, creating a lattice of positive metal ions immersed in a 'sea' of delocalised electrons. These electrons are not associated with any single atom but move throughout the metal, producing collective bonding that holds the metal structure together. This model explains many characteristic metallic properties.

Delocalised electron sea and conduction
Because electrons are free to move within the metal lattice they can carry electric charge under an applied potential difference. This electron mobility makes metals good conductors of electricity in the solid state. The same delocalised electrons also transport kinetic energy efficiently, giving metals high thermal conductivity.

Malleability and ductility
The ability to deform metals without breaking, called malleability (hammering into sheets) and ductility (drawing into wires), follows from metallic bonding. When layers of metal ions slide past each other, the delocalised electrons continue to hold the lattice together, so bonds are not broken as they would be in an ionic or covalent lattice. This is why metals can be reshaped without shattering.

Lustre and cohesion
Metals are shiny because mobile electrons reflect and re-emit light, creating metallic lustre. The cohesion provided by metallic bonds gives metals strength and explains why they often form dense, closely packed crystal structures. Strength and melting point vary across metals because metallic bonding strength depends on valence electron density and ionic size.

Alloys and modified properties
Alloys are mixtures of metals (sometimes with non-metals) that change the lattice arrangement and electron behaviour to yield altered properties like greater hardness, resistance to corrosion or altered melting points. For example, steel (iron + carbon) is much harder than pure iron because small carbon atoms distort iron’s lattice and inhibit dislocation movement.

Comparison with other bonds
Metallic bonding differs from ionic and covalent bonding: ionic bonds result from full transfer of electrons and electrostatic attraction between ions, while covalent bonds involve sharing between specific atoms. Metallic bonding involves delocalised electrons shared among many atoms. Each bonding type yields different macroscopic properties, and recognising these differences helps explain material uses in daily life and technology.

What are Lewis structures?
Lewis structures, or electron-dot diagrams, use dots to represent valence electrons around the chemical symbol of an atom. Shared pairs forming covalent bonds appear as paired dots or as a line between atomic symbols. Lewis diagrams are a simple visual tool to show how atoms share or transfer electrons to form molecules and ions and to check whether atoms achieve stable electron arrangements.

The octet and duet rules
Most main-group elements tend to achieve an octet of electrons (eight) in their valence shell because noble gas configurations are particularly stable. Hydrogen is stable with two electrons (a duet). These rules explain why atoms gain, lose or share electrons. While the octet rule is a useful guide, remember there are exceptions such as boron compounds and elements beyond the second period that can have expanded octets.

Steps to draw Lewis structures
1. Count total valence electrons from each atom, adjusting for ionic charges.
2. Place the least electronegative atom (except hydrogen) at the centre and connect atoms by single bonds.
3. Distribute remaining electrons to satisfy octets of outer atoms first, placing lone pairs as needed.
4. If the central atom lacks an octet, form multiple bonds by converting lone pairs from outer atoms into shared pairs until the octet is satisfied, bearing in mind exceptions.

Formal charges and resonance
While not required in full detail at Class 9, be aware that sometimes different Lewis structures can be drawn for the same molecule. When electron distribution cannot be represented by a single structure, resonance structures are used to show delocalisation. Formal charges help choose the best Lewis structure: the structure with formal charges closest to zero and negative charges on more electronegative atoms is usually preferred.

Common examples
Water (H2O) is drawn with O having two lone pairs and two single bonds to H: H–O–H. Carbon dioxide (CO2) needs two double bonds to give C an octet: O=C=O. Ammonia (NH3) has three N–H bonds and one lone pair on N. Boron trifluoride (BF3) is a notable exception where boron has only six electrons in many compounds and remains electron-deficient.

Usefulness and limits
Lewis diagrams are widely used to predict molecular shapes, bond counts and reactive sites. They are a simple model and do not show electron probability distributions or provide quantitative bond energies, but they are indispensable for the early study of bonding and reactions in Class 9 chemistry.

Bond polarity and electronegativity
Bond polarity arises from differences in electronegativity—the tendency of an atom to attract shared electrons. When two atoms in a bond have different electronegativities, the shared electron pair is drawn closer to the more electronegative atom, creating a partial negative charge (δ−) there and a partial positive charge (δ+) on the other atom. Examples include H–Cl (polar) and H–H or Cl–Cl (non-polar).

Molecular polarity depends on shape
Molecular polarity is the vector sum of individual bond dipoles and depends crucially on molecular geometry. Even if bonds are polar, a symmetric arrangement can cancel bond dipoles and give a non-polar molecule overall. For instance, CO2 has two polar C=O bonds but is linear, and the equal opposite dipoles cancel, yielding a non-polar molecule. In contrast, H2O has polar O–H bonds and a bent shape so the dipoles do not cancel and the molecule is polar.

How to decide polarity
1. Draw the Lewis structure and determine the geometry of the molecule (using simple VSEPR ideas).
2. Identify polar bonds by differences in electronegativity.
3. Add bond dipoles vectorially: if they cancel, the molecule is non-polar; if they do not, the molecule is polar. For many class-level molecules, simple symmetry arguments are sufficient to decide polarity.

Effects of polarity
Polarity affects many physical properties: polar molecules have stronger intermolecular attractions (dipole–dipole interactions, and hydrogen bonds where applicable) and often higher boiling points than comparable non-polar molecules. Polar substances tend to be soluble in polar solvents like water; non-polar substances dissolve in non-polar solvents. Polarity also affects reactivity and how molecules orient in electric fields.

Special case: hydrogen bonding
When hydrogen is bonded to very electronegative atoms (N, O, F), strong dipole interactions called hydrogen bonds form between molecules. Hydrogen bonding is an extreme case of polarity with large effects on boiling points and other properties, as seen in water’s unusually high boiling point for its molar mass.

Limitations and practical approach
For Class 9 focus on simple molecules and geometries: linear, bent, trigonal planar and tetrahedral. Use electronegativity differences to mark bond polarity and symmetry to test if dipoles cancel. This approach is sufficient to understand common examples and their properties at school level.

Qualitative VSEPR idea
Valence Shell Electron Pair Repulsion (VSEPR) theory explains molecular shapes qualitatively by stating that electron pairs (bonding and lone pairs) around a central atom repel each other and arrange themselves as far apart as possible. This simple idea predicts many common molecular shapes and helps relate geometry to bonding and polarity.

Counting regions of electron density
To predict shape, draw the Lewis structure and count regions of electron density around the central atom — each single, double or triple bond counts as one region, as does each lone pair. The number of regions gives the basic electron-domain geometry: 2 regions → linear; 3 → trigonal planar; 4 → tetrahedral; 5 → trigonal bipyramidal; 6 → octahedral. Then consider how lone pairs compress bond angles, producing shapes such as bent or trigonal pyramidal.

Common shapes and examples
Linear: two bonding regions (e.g., CO2) give 180° bond angles. Trigonal planar: three bonding regions (e.g., BF3) give roughly 120°. Tetrahedral: four bonding regions (e.g., CH4) give about 109.5°. Trigonal pyramidal: three bonding regions + one lone pair (e.g., NH3) gives slightly less than 109.5° due to lone pair repulsion. Bent (angular): two bonding regions + one or two lone pairs (e.g., H2O) produces bond angles smaller than the ideal tetrahedral angle.

Role of lone pairs
Lone pairs occupy more space than bonding pairs because they are localized closer to the central atom and repel bonding pairs more strongly. This decreases bond angles between bonded atoms. For example, the two lone pairs on oxygen in water push the hydrogen atoms closer together to give an angle of about 104.5°, less than the tetrahedral angle.

Multiple bonds and geometry
Multiple bonds (double or triple) act as a single region of electron density for geometry purposes, though they repel slightly more than single bonds and can modify angles a little. For example, the double bond in CO2 counts as one region on carbon but gives a linear arrangement with the two oxygens.

Practical approach for Class 9
Use simple Lewis structures and the count-of-regions method to assign shapes for small molecules. Sketch 3D arrangements using simple models — tetrahedral as a pyramid-like arrangement with one bond drawn behind or as wedges in diagrams. Understanding shapes helps explain polarity and intermolecular forces, linking structure to properties.

Definitions
Bond length is the average distance between the nuclei of two bonded atoms. Bond energy (bond enthalpy) is the energy required to break one mole of bonds in the gas phase, and bond strength refers to how strongly atoms are held together in the bond. These three concepts are closely related: stronger bonds typically have higher bond energies and shorter bond lengths.

Bond order effect
Bond order indicates how many electron pairs are shared between two atoms: single (bond order 1), double (2) or triple (3). As bond order increases, bond length decreases and bond energy increases because more shared electrons produce greater attraction between the bonded nuclei and the shared electrons. For example, a C–C single bond is longer and weaker than a C=C double bond, which is in turn longer and weaker than a C≡C triple bond.

Factors affecting bond length and energy
Atomic sizes influence bond length: bonds between larger atoms are generally longer. Electronegativity differences and bond polarity can alter bond energy since electrostatic contributions change. Resonance and delocalisation can give bonds intermediate character between single and double; these resonance-stabilised bonds have lengths and energies between the single and double values.

Measurement and comparison
Bond lengths are measured by techniques like X-ray diffraction and spectroscopy, giving precise distances in picometres (pm). Bond energies are tabulated as average values because they depend on molecular environment. For Class 9, use comparative reasoning—knowing relative strengths and lengths is more important than exact numbers.

Chemical implications
Bond energies explain why chemical reactions are exothermic or endothermic: breaking bonds requires energy, forming bonds releases energy. A reaction is exothermic if the total energy released in forming new bonds exceeds the energy required to break reactant bonds. Bond strength also influences reactivity: weaker bonds are easier to break and thus more reactive under certain conditions.

Examples and exceptions
O2 has a double bond with significant bond energy; O–O single bonds (as in peroxides) are weaker and easier to break. Iodine molecules have relatively weak I–I bonds and low melting points. Network covalent solids like diamond have extremely strong bonds throughout the structure, leading to very high melting points and hardness, showing how bond types at the microscopic level determine macroscopic properties.

What are intermolecular forces?
Intermolecular forces are attractive interactions between molecules. They are weaker than intramolecular covalent or ionic bonds but crucial in determining physical properties such as boiling point, melting point, viscosity, solubility and state of matter. Understanding these forces explains why substances with similar molecular weights can behave very differently.

London dispersion (van der Waals) forces
All molecules, polar or non-polar, experience London dispersion forces due to momentary fluctuations in electron distribution that create temporary dipoles. These instantaneous dipoles induce dipoles in neighbouring molecules, leading to weak attraction. Dispersion forces increase with molecular size and polarizability (how easily electron clouds distort), so larger molecules have stronger dispersion forces and higher boiling points than smaller molecules of similar type.

Permanent dipole–dipole interactions
Polar molecules possess permanent dipoles because of unequal sharing of electrons in bonds. These dipoles align such that positive ends of one molecule attract negative ends of another, creating dipole–dipole interactions stronger than simple dispersion forces for molecules of comparable size. The strength depends on dipole magnitude and molecular orientation.

Hydrogen bonding
Hydrogen bonding is a particularly strong type of dipole interaction occurring when hydrogen is covalently bonded to a highly electronegative atom (N, O, or F), and this hydrogen interacts with a lone pair on another electronegative atom. Hydrogen bonds are responsible for many unusual properties of water: high boiling point, high surface tension and the fact that ice is less dense than liquid water. Hydrogen bonding also affects biological structures such as DNA base pairing and protein folding.

Consequences for properties and solubility
Strong intermolecular forces lead to higher boiling and melting points; weak forces yield low boiling points and gaseous states at room temperature. Intermolecular forces also determine solubility patterns — polar solvents dissolve polar and ionic substances via favourable interactions, whereas non-polar solvents dissolve non-polar substances. The rule 'like dissolves like' summarises this tendency.

Class 9 focus
Recognise the three main types of intermolecular forces and relate them qualitatively to observed properties. Use simple examples (e.g., CH4 vs H2O, I2 vs NaCl) to reinforce how intermolecular forces shape everyday observations, and understand that stronger intermolecular forces mean higher energies are required to separate molecules.

Electrovalency vs covalency
Electrovalency refers to ionic bonding where electrons are transferred and oppositely charged ions attract. Covalency refers to bonding by sharing electron pairs between atoms. These represent two idealised ends of a bonding spectrum, and many real compounds have bonding character between these extremes. Comparing the two helps predict properties like solubility, melting point and electrical conductivity.

Which type forms between elements?
Large differences in electronegativity, typically between metals and non-metals, favour electrovalent (ionic) bonding because one atom strongly attracts electrons and the other readily loses them. Similar electronegativities, usually between non-metals, favour covalent bonding where electrons are shared. The periodic table gives useful guidance: group 1 or 2 metals with group 16 or 17 non-metals typically form ionic salts, while non-metal–non-metal combinations form covalent molecules.

Physical distinctions
Ionic compounds are usually crystalline, hard, and have high melting/boiling points. They conduct electricity when molten or in solution due to mobile ions but not in solid state. Covalent molecular substances often have lower melting/boiling points, can be gases, liquids or soft solids, and generally do not conduct electricity because no free charge carriers exist. Metallic bonding is distinct, giving metals conductivity in solid state due to delocalised electrons.

Borderline and mixed character
Some compounds show mixed character. Factors like high ionic charge and small ionic sizes increase polarising power and can shift ionic bonds toward covalent character. For instance, AlCl3 shows substantial covalent character due to polarisation of chloride ions by small Al3+. Fajan’s rules (qualitative) help predict such behaviour: small highly charged cations polarise large anions, increasing covalency.

Usefulness of comparative view
Understanding whether a bond is mostly ionic or covalent helps explain and predict substance behaviour. For example, NaCl dissolves readily in water and conducts when dissolved because it is ionic, while sugar dissolves but does not conduct because it is covalent and remains as neutral molecules in solution. Recognising mixed character prevents over-simplification of bonding in real compounds.

Class 9 scope
At this level, focus on clear examples of ionic and covalent bonding, learn the typical properties of each, and use simple rules of electronegativity difference and periodic trends to predict bonding type. Awareness of exceptions such as covalent character in some ionic compounds completes the basic comparative picture.

Why formulas and names matter
Writing correct chemical formulas and names is a basic skill for communicating composition of substances. Formulas tell the kinds and numbers of atoms in a compound; names (systematic or common) allow clear identification. For ionic compounds the formula balances charges; for covalent compounds prefixes or common names indicate numbers of atoms.

Rules for ionic formulas
1. Identify the cation (positive ion) and anion (negative ion) with their charges.
2. Balance the total positive and negative charges so that the compound is neutral.
3. Write the symbols with subscripts showing the number of each ion; omit '1' as a subscript. For example, Mg2+ and Cl− combine: two Cl− balance one Mg2+, giving MgCl2. For polyatomic ions treat the group as a unit and use parentheses if more than one is present, e.g., Ca(NO3)2.

Naming ionic compounds
For binary ionic salts name the cation first then the anion with an -ide ending (e.g., sodium chloride). For cations with variable charges (common in transition metals), Roman numerals indicate charge in the systematic name, e.g., iron(III) chloride for FeCl3. Polyatomic ions have standard names such as sulfate (SO4^2−) and nitrate (NO3−) which should be memorised.

Rules for covalent compound names
Binary covalent compounds (two non-metals) often use prefixes to indicate number of atoms: mono-, di-, tri-, tetra-, etc. The first element retains its name; the second ends with -ide. For instance CO is carbon monoxide, CO2 is carbon dioxide, and SF6 is sulfur hexafluoride. In practice, simple molecules like water (H2O) and ammonia (NH3) are commonly used by their common names rather than systematic names.

Practice strategy
Memorise common ion charges (group 1 → +1, group 2 → +2, group 17 → −1) and common polyatomic ions. Practice converting names to formulas and vice versa by balancing charges and counting atoms. Use parentheses for multiple polyatomic ions and ensure correct subscripts to reflect stoichiometry of neutral compounds.

Examples and pitfalls
Be careful to balance charges rather than valencies directly; for example Al3+ and O2− combine as Al2O3 (two Al3+ give +6 balanced by three O2− giving −6). Also avoid writing charges in the final formula. Correct formula writing is vital for further topics like reaction stoichiometry and molar mass calculations.

Empirical vs molecular formula
An empirical formula gives the simplest whole-number ratio of atoms of each element in a compound, while a molecular formula gives the actual number of atoms of each element in one molecule. For example, glucose has molecular formula C6H12O6 but empirical formula CH2O because the numbers reduce to the smallest whole-number ratio.

How to find empirical formula from composition
When given percentage composition or masses of elements in a sample, a stepwise method finds the empirical formula. Assume a convenient sample mass (often 100 g) so percentages become masses. Convert masses to moles by dividing by atomic masses. Divide all mole values by the smallest mole value to obtain a ratio. If necessary, multiply ratios by a small integer to remove fractional subscripts and obtain whole numbers. This gives the empirical formula.

Obtaining molecular formula
If the molar mass (molecular mass) of the compound is known, divide the molar mass by the empirical formula mass (sum of atomic masses in the empirical formula) to obtain an integer n. Multiply empirical subscripts by n to get the molecular formula. Class 9 may not require detailed determination of molar mass experimentally, but understanding the relation is useful.

Worked example approach
If a compound contains 40% C, 6.7% H and 53.3% O, assume 100 g sample: 40 g C, 6.7 g H, 53.3 g O. Convert to moles: 40/12 = 3.33 mol C, 6.7/1 = 6.7 mol H, 53.3/16 = 3.33 mol O. Divide by smallest (3.33) to get C1H2O1, giving empirical formula CH2O. If molar mass were 180 g mol−1 (as for glucose), empirical mass CH2O is 30, multiplier n = 180/30 = 6, so molecular formula C6H12O6.

Importance and uses
Empirical formulas summarise composition concisely and are used in analytical chemistry, combustion analysis and stoichiometry. Molecular formulas are necessary to understand molecular size and reactions involving moles. Class 9 exercises focus on converting percentages to empirical formulas and recognising the difference between empirical and molecular expressions.

Why ionic compounds may show covalent character
The simple ionic model assumes full transfer of electrons, but in practice some ionic compounds show covalent characteristics. Fajan’s rules provide a qualitative way to predict when an ionic bond will have appreciable covalent character by considering the sizes and charges of ions and their electronic structures.

Main ideas of Fajan's rules
1. Small cations with high positive charge are strongly polarising: they attract electron density from nearby anions, distorting the anion’s electron cloud and causing partial sharing of electrons.
2. Large anions with high negative charge are easily polarised because their outer electrons are farther from the nucleus and less tightly held; this increases covalent character when paired with polarising cations.
3. Cations with a noble gas electronic configuration are less polarising than those with pseudo-noble gas configurations; higher charge and smaller size increase polarisation.
4. High ionic charge and large size difference between ions increase polarisation and thus covalent character.

Practical consequences and examples
Al3+ is small and highly charged and therefore a strong polariser; AlCl3 shows strong covalent character and exists as molecular species in the vapour phase. In contrast, Na+ is larger and singly charged, so NaCl behaves as a typical ionic salt. Among halides, iodide (I−) is large and more polarisable than fluoride (F−), so bonding with a small highly charged cation tends to be more covalent with I− than with F−.

How to use the rules
For a given pair of ions, check cation size and charge and anion size and charge. If cation is small and multiply charged and anion is large with high charge, expect increased covalent character. These rules are qualitative but useful to explain observed deviations from ideal ionic behaviour such as lower melting points or covalent molecular structures.

Limitations
Fajan’s rules are approximate and apply best to simple binary salts. They give no numerical predictions but help bridge ionic and covalent models. At Class 9 level the rules build intuition about why some compounds do not fit neatly into purely ionic or covalent categories.

How bonding determines physical properties
The type of bonding in a substance largely controls its macroscopic properties such as melting point, electrical conductivity and solubility. By examining bonding we can predict whether a material will be a solid, liquid or gas at room temperature and how it will behave in electric fields and different solvents.

Melting and boiling points
Ionic solids typically have high melting and boiling points because strong electrostatic attractions between ions in the lattice require large energy to overcome. Covalent molecular substances usually have low to moderate melting points because they are held together by weaker intermolecular forces; network covalent solids like diamond are exceptions because strong covalent bonds extend through the whole lattice and produce very high melting points.

Electrical conductivity
In ionic solids ions are fixed and cannot move, so they do not conduct electricity in the solid state. When ionic compounds are molten or dissolved in water, ions become mobile and the substance conducts electricity. Metals conduct electricity in the solid state because of delocalised electrons that move freely. Covalent molecular substances generally do not conduct electricity because they lack mobile charged particles, although some covalent network solids (graphite) conduct due to delocalised electrons in layers.

Solubility in solvents
Solubility depends on the nature of interactions between solute and solvent. Polar solvents like water stabilise ions and polar molecules via ion–dipole and dipole–dipole interactions, so ionic and polar covalent substances tend to dissolve in water. Non-polar solvents dissolve non-polar substances by dispersion forces. The adage 'like dissolves like' summarises this behaviour and guides predictions about solubility.

Mechanical properties
Ionic crystals tend to be hard and brittle because displacing layers brings like-charged ions into contact causing cleavage. Metals are malleable and ductile due to the electron sea allowing layers to slide without breaking bonding. Molecular solids are often soft due to weak inter-molecular forces. Thus bonding type explains mechanical differences among materials used in daily life.

Experimental identification
Simple laboratory tests — melting point measurement, conductivity tests in solid and molten states, and solubility trials in water vs organic solvent — provide evidence for bonding type. For example, a white crystalline solid that melts at high temperature, dissolves in water and conducts when molten is likely ionic. These tests connect theory to observation and are commonly used in Class 9 practical work.

Why experiments are useful
Simple experiments let students connect theoretical bonding ideas with real observations. By testing melting points, conductivity and solubility and by building physical models, learners see how ionic, covalent and metallic bonding cause characteristic behaviours. Experiments are essential to develop scientific reasoning and to check predictions made from electronic configurations and bonding models.

Melting point comparisons
Compare the melting points of an ionic salt (e.g., NaCl) and a covalent molecular solid (e.g., iodine). Using small samples and proper apparatus, students will notice that NaCl melts at much higher temperature than iodine. This illustrates strong ionic lattice forces versus weak intermolecular forces in molecular solids. Always follow safety instructions and use small quantities.

Conductivity tests
Set up a simple circuit with a bulb or meter to test conductivity of samples in solid and molten states. Ionic solids like NaCl do not conduct as solids but do conduct when molten or dissolved, showing that mobile ions are required for conductivity. Metals such as copper wire conduct electricity in solid state due to delocalised electrons. Covalent molecular solids like sugar do not conduct in either state, confirming lack of free charge carriers.

Solubility experiments
Test solubility of salt, sugar and oil in water and in a non-polar solvent (teacher-supervised). Salt dissolves in water and not in non-polar solvents; oil dissolves in non-polar solvents but not in water; sugar dissolves in water but does not conduct. These observations demonstrate polarity and the 'like dissolves like' principle as well as effects of intermolecular forces.

Model building and visualization
Construct simple ball-and-stick or paper models for molecules like CH4 (tetrahedral) and H2O (bent) to visualise shapes and bond angles. Draw Lewis structures to plan models and to explain molecular polarity. Physical models help understand 3-D arrangements which are hard to visualise on paper alone.

Recording and explaining results
Encourage students to note observations carefully, propose explanations using bonding models (ionic lattice, covalent molecules, metallic sea of electrons), and relate results to properties like melting point and conductivity. Emphasise safety, small quantities, and teacher supervision. Through these simple experiments students gain a practical appreciation of how microscopic bonding leads to macroscopic behaviour.