BONDING PPT

1. August 2025 ARUSHA SCIENCE SCHOOL CHEMICAL BONDING Elishama Hubi

2. 07 01 02 03 04 05 06 CONTENT Introduction to Chemical Bonding Types of Chemical Bonds Bond Theories Shapes of Molecules and Ions Bond Polarity and Molecular Polarity Intermolecular Forces Energetics of Bonding

3. OVERVIEW ARUSHA SCIENCE Importance of Bonding in Chemistry • A chemical bond is the force of attraction between two or more atoms or ions that holds them together in a stable chemical entity (molecule, ion, or solid lattice). • Bonds form because atoms seek a more stable (lower-energy) electron configuration—often resembling that of a noble gas. • Bonds explain why atoms combine to form substances—they stabilize atoms by achieving electron configurations with lower energy. • Understanding bonding is key to predicting properties such as melting/boiling points, solubility, conductivity, reactivity, and molecular shape. • It forms the foundation for deeper study of molecular interactions, reactivity, and material properties.

4. Role of Valence Electrons in Bonding ARUSHA SCIENCE Meaning Role Role • In ionic bonding, electrons are transferred from one atom to another; in covalent bonding, electrons are shared. • Atoms bond to fill or empty their valence shell—achieving greater stability (e.g., noble gas configuration). • Valence electrons are the electrons in the outermost shell of an atom and are the ones involved in bonding.

5. Duplet Rule (First Shell Elements) • Duplet (or duplet) rule applies to H, He, Li, and Be, since their first energy level can hold a maximum of two electrons. • Exceptions to the Octet Rule • There are three main categories of exceptions: • Incomplete (Electron-Deficient) Octet • Atoms like boron (e.g., in BF₃) have fewer than eight electrons—forming stable compounds despite incomplete octets. • Such electron-deficient species are highly reactive and can accept electrons (e.g., BF₃ accepts a pair from NH₃). • Expanded Octet (Hypervalent Molecules) • Atoms in period 3 and beyond (like P in PCl₅ or S in SF₆) can accommodate more than eight electrons using available d orbitals (sp³d or sp³d² hybridization). • Odd-Electron Species (Free Radicals) • Molecules with odd numbers of electrons (e.g., NO) cannot form complete octets—they contain unpaired electrons and are typically reactive free radicals. Octet Rule & Exceptions The Octet Rule Octet rule: Atoms tend to gain, lose, or share electrons to achieve eight electrons in their outer shell. Commonly applies to elements in the s- and p-blocks. Example: In CO₂, each oxygen and the central carbon achieve stability by sharing electrons to complete their octets. ARUSHA SCIENCE

6. TYPES OF CHEMICAL BONDING ARUSHA SCIENCE

9. 1. IONIC BONDING DefinitionAn ionic bond is formed when one atom transfers one or more electrons to another atom, leading to the formation of positively charged cations and negatively charged anions. These oppositely charged ions are held together by electrostatic forces of attraction. Mechanism of Formation Electron transfer occurs between atoms with a large electronegativity difference . Metals (low electronegativity) lose electrons to form cations. Nonmetals (high electronegativity) gain electrons to form anions. TYPES OF CHEMICAL BONDING Threre are 3 major types of chemical bonding Ionic bonding Covalent bonding Metallic bonding ARUSHA SCIENCE

10. Example: Sodium Chloride (NaCl) ARUSHA SCIENCE Sodium (Na) Electronic configuration = 1s2 2s2 2p6 3s1Na → Na⁺ + e⁻ Ionization Energy (I.E) Chlorine (Cl) Electronic configuration = 1s2 2s2 2p6 3s2 3p5 Cl + e⁻ → Cl⁻ Electron Affinity (E.A) Overall reactionNa⁺ + Cl⁻ → NaCl(s) Lattice Energy (L.E)

11. Example: Sodium Chloride (NaCl) ARUSHA SCIENCE

12. Characteristics of Ionic Cpds ARUSHA SCIENCE • 1. High Melting and Boiling Points • Due to strong electrostatic forces of attraction between oppositely charged ions in the lattice. • Large amounts of energy are required to overcome these forces. • Example: NaCl melts at 801 °C, MgO at ~2852 °C. • 2. Hard and Brittle Nature • Hard because the ions are tightlypacked in a rigid crystalline lattice. • Brittle because a slight displacement of layers causes like charges to come adjacent, resulting in strong repulsion and fracture. • 3. Electrical Conductivity • Solid state: Poor conductor because ions are fixed in position and cannot move freely. • Molten state: Good conductor because ions are free to move and carry electric current. • Aqueous solutions: Good conductor as ions are dissociated and mobile.

13. Characteristics of Ionic Cpds ARUSHA SCIENCE • 4. Solubility • Generally soluble in polar solvents (e.g., water) due to strong ion–dipole interactions. • Insoluble or poorly soluble in non-polar solvents (e.g., hexane). • 5. Crystalline Structure • Exist as giant ionic lattices with regular, repeating arrangements of ions. • Lattice type depends on ion sizes and charges (e.g., rock salt structure for NaCl). • 6. Non-directional Bonding • The electrostatic attraction acts equally in all directions around the ion in the lattice, making the bonding non-directional.

14. 2. COVALENT BONDING DefinitionA covalent bond is formed when two atoms share one or more pairs of electrons to achieve a stable electron configuration (usually the noble gas configuration). The shared electrons are attracted simultaneously by the nuclei of both atoms. Mechanism of Formation Occurs mainly between non-metallic elements with similar or identical electronegativities. Each atom contributes at least one electron to form the shared pair(s). The shared electrons occupy an orbital between the two nuclei, resulting in a region of high electron density. General representation: Single covalent bond: H + H → H–H Double covalent bond: O = O Triple covalent bond: N ≡ N TYPES OF CHEMICAL BONDING Threre are 3 major types of chemical bonding Ionic bonding Covalent bonding Metallic bonding ARUSHA SCIENCE

15. Types of Covalent Bonds ARUSHA SCIENCE Single Bond → one shared pair of electrons (e.g., H₂, Cl₂, CH₄). Double Bond → two shared pairs of electrons (e.g., O₂, CO₂). Triple Bond → three shared pairs of electrons (e.g., N₂, C₂H₂).

16. Types of Covalent Bonds ARUSHA SCIENCE

17. Single Bond ARUSHA SCIENCE Example 1: H₂ moleculeHydrogen atoms each have 1 electron:H (1s¹) + H (1s¹) → H:HOr represented as:H + H → H–H Example 2: Cl₂ moleculeChlorine atoms each have 7 valence electrons:Cl (2,8,7) + Cl (2,8,7) → Cl:ClOr:Cl + Cl → Cl–Cl

18. Single Bond ARUSHA SCIENCE Example 3: CH₄ moleculeCarbon has 4 valence electrons, hydrogen has 1:C (2,4) + 4 × H (1) → H |H–C–H|H (Each C–H bond is a single covalent bond.)

19. Double Bond ARUSHA SCIENCE Example 1: O₂ moleculeOxygen atoms each have 6 valence electrons:O (2,6) + O (2,6) → O::OOr:O = O Example 2: CO₂ moleculeCarbon (4 valence e⁻) bonds with two oxygen atoms (6 valence e⁻ each):O = C = OEach C=O is a double bond (two shared pairs).

20. Triple Bond ARUSHA SCIENCE Example 1: N₂ moleculeNitrogen atoms each have 5 valence electrons:N (2,5) + N (2,5) → N≡N(Three shared pairs between the two nitrogen atoms.) Example 2: C₂H₂ (ethyne) moleculeCarbon atoms (4 valence e⁻ each) bond to each other with a triple bond, and each bonds to one hydrogen atom:H–C≡C–H

21. Characteristics of Covalent Cpds ARUSHA SCIENCE • 1. Low Melting and Boiling Points • Due to weak intermolecular forces (Van der Waals, dipole–dipole, or hydrogen bonds) between molecules. • Exception: Network covalent solids (e.g., diamond, SiO₂) have very high melting points due to strong covalent bonds throughout the structure. • 2. Electrical Conductivity • Generally poor conductors in all states (no free ions or delocalized electrons). • Exceptions: Graphite conducts electricity due to delocalized π-electrons. • 3. Solubility • Non-polar covalent compounds are soluble in non-polar solvents (e.g., benzene). • Polar covalent compounds may be soluble in polar solvents (e.g., sugar in water).

22. Characteristics of Covalent Cpds ARUSHA SCIENCE • 4. Bond Directionality • Covalent bonds are directional — electron density is concentrated between the bonded atoms, giving molecules definite shapes. • 5. Variable Reactivity • Some covalent compounds are highly reactive (e.g., halogens), others are inert (e.g., N₂). • 6. Molecular and Network Structures • Molecular covalent: discrete molecules held by weak intermolecular forces (e.g., H₂O, CO₂). • Network covalent: giant continuous lattice of atoms bonded covalently (e.g., diamond, SiC).

23. 2. METALLIC BONDING • DefinitionMetallic bonding is the type of chemical bonding that occurs between metal atoms, where positively charged metal ions (cations) are surrounded by a “sea” of delocalized valence electrons. These electrons are free to move throughout the lattice, holding the cations together through electrostatic attraction. • Mechanism of Formation • Metals have low ionization energies and tend to lose valence electrons easily. • These electrons become delocalized (free from any specific atom) and move throughout the entire metallic lattice. • The metal atoms become positively charged ions arranged in a regular pattern, embedded in the mobile sea of electrons. • The electrostatic attraction between the positive ions and the electron sea forms the metallic bond. TYPES OF CHEMICAL BONDING Threre are 3 major types of chemical bonding Ionic bonding Covalent bonding Metallic bonding ARUSHA SCIENCE

24. General Representation ARUSHA SCIENCE Metal atoms → Metal cations + Delocalized electronsExample: Na → Na⁺ + e⁻ (delocalized)

25. General Representation ARUSHA SCIENCE

26. Characteristics of Metallic Cpds ARUSHA SCIENCE • 1. High Electrical Conductivity • Free-moving delocalized electrons carry electric current in both solid and molten states. • 2. Thermal Conductivity • Delocalized electrons transfer heat energy rapidly through the lattice. • 3. Malleability and Ductility • Metallic bonds are non-directional; cations can slide over one another without breaking the bond, allowing metals to be hammered into sheets (malleable) or drawn into wires (ductile).

27. Characteristics of Metallic Cpds ARUSHA SCIENCE • 4. Lustre • Free electrons reflect light, giving metals their shiny, metallic appearance. • 5. Variable Melting Points • Transition metals (e.g., tungsten) have very high melting points due to strong metallic bonding. • Alkali metals have lower melting points because of weaker metallic • 6. Non-directional Bonding • The electrostatic attraction acts in all directions between cations and the electron sea, giving metals their unique mechanical properties.

28. Characteristics of Covalent Cpds ARUSHA SCIENCE • Next Part on Types of chemical Bonding • Dative Bonding • Covalent Character of Ionic Compounds (Fajan Rule) on Polarisation • Ionic character of Covalent compunds (Polarity)

29. The Reaction Between Ammonia and Hydrogen Chloride Ammonium ions, NH4+, are formed by the transfer of a hydrogen ion (a proton) from the hydrogen chloride molecule to the lone pair of electrons on the ammonia molecule. COORDINATE OR DATIVE BOND A coordinate bond, also known as a dative covalent bond, is a type of covalent bond where one atom provides both electrons for the shared pair.  Unlike a regular covalent bond where each atom contributes one electron, in a coordinate bond, one atom acts as the sole electron donor while the other accepts them. This type of bonding is crucial in coordination compounds and is often seen in Lewis acid-base reactions. 

30. ARUSHA SCIENCE When the ammonium ion, NH4+, is formed, the fourth hydrogen is attached by a dative covalent bond, because only the hydrogen's nucleus is transferred from the chlorine to the nitrogen. The hydrogen's electron is left behind on the chlorine to form a negative chloride ion. Once the ammonium ion has been formed it is impossible to tell any difference between the dative covalent and the ordinary covalent bonds. Although the electrons are shown differently in the diagram, there is no difference between them in reality.

31. In simple diagrams, a coordinate bond is shown by an arrow. The arrow points from the atom donating the lone pair to the atom accepting it. ARUSHA SCIENCE

32. The structure of Aluminum Chloride Aluminum chloride sublimes (phase transition from solid to gas) at about 180°C. If it simply contained ions, it would have a very high melting and boiling point because of the strong attractions between the positive and negative ions. The implication is that it when it sublimes at this relatively low temperature, it must be covalent. The dots-and-crosses diagram shows only the outer electrons. ARUSHA SCIENCE

33. AlCl3, like BF3, is electron deficient. There is likely to be a similarity, because aluminum and boron are in the same group of the Periodic Table, as are fluorine and chlorine. Measurements of the relative formula mass of aluminum chloride show that its formula in the vapor at the sublimation temperature is not AlCl3, but Al2Cl6. It exists as a dimer (two molecules joined together). The bonding between the two molecules is coordinate, using lone pairs on the chlorine atoms. Each chlorine atom has 3 lone pairs, but only the two important ones are shown in the line diagram.

34. The uninteresting electrons on the chlorines have been faded in color to make the coordinate bonds show up better. There's nothing special about those two particular lone pairs - they just happen to be the ones pointing in the right direction. Energy is released when the two coordinate bonds are formed, and so the dimer is more energetically stable than two separate AlCl3 molecules.

35. Nature of chemical bond http://www.adichemistry.com/general/chemicalbond/fajan/fajans-rules.html • Dative Bonding • Covalent Character of Ionic Compounds (Fajan Rule) on Polarisation • Ionic character of Covalent compunds (Polarity)

36. Upcoming Lesson……… ARUSHA SCIENCE • Energetics of Bonding • Bond length, bond energy, bond strength • Bond energy and bond dissociation enthalpy • Relationship between bond strength and reactivity • Hess’s Law for bond energy calculations • Lattice energy (Born-Haber cycle – qualitative)

37. Bonding Energetics The Relationship between Bond Order and Bond Energy Bond order is a measure of the number of chemical bonds between a pair of atoms.  It indicates the stability and strength of a bond. A higher bond order generally signifies a stronger and more stable bond.  Single bond: Bond order of 1 (e.g., H-H in H2). Double bond: Bond order of 2 (e.g., O=O in O2). Triple bond: Bond order of 3 (e.g., N≡N in N2). Triple bonds between like atoms are shorter than double bonds, and because more energy is required to completely break all three bonds than to completely break two, a triple bond is also stronger than a double bond. Similarly, double bonds between like atoms are stronger and shorter than single bonds. Bonds of the same order between different atoms show a wide range of bond energies The stability of a molecule is a function of the strength of the covalent bonds holding the atoms  together. Bond energy is defined as the energy required to break a particular bond in a molecule in the gas phase.

39. Bond Dissociation Energy (Bond Energy) ARUSHA SCIENCE • Bond Dissociation Energy (also referred to as Bond energy) is the enthalpy change ΔH, heat input) required to break a bond (in 1 mole of a gaseous substance. • What about when we have a compound which is not a diatomic molecule? Consider the dissociation of methane:

40. Bond Dissociation Energy (Bond Energy) ARUSHA SCIENCE There are four equivalent C-H bonds, thus we can that the dissociation energy for a single C-H bond would be: Bond energy is always a positive value - it takes energy to break a covalent bond (conversely energy is released during bond formation)

41. Bond Energy and Enthalpy of Reactions ARUSHA SCIENCE • We can estimate the enthalpy change for a chemical reaction by adding together the average energies of the bonds broken in the reactants and the average energies of the bonds formed in the products and then calculating the difference between the two. • If the bonds formed in the products are stronger than those broken in the reactants, then energy will be released in the reaction ( ΔHrxn <0): • The ≈ sign is used because we are adding together average bond energies; hence this approach does not give exact values for ΔHrxn.

42. Example ARUSHA SCIENCE Let’s consider the reaction of 1 mol of n-heptane (C7H16) with oxygen gas to give carbon dioxide and water. This is one reaction that occurs during the combustion of gasoline: In this reaction, 6 C–C bonds, 16 C–H bonds, and 11 O=O bonds are broken per mole of n-heptane, while 14 C=O bonds (two for each CO2) and 16 O–H bonds (two for each H2O) are formed. The energy changes can be tabulated as follows:

46. Summary Bond order is the number of electron pairs that hold two atoms together. Single bonds have a bond order of one, and multiple bonds with bond orders of two (a double bond) and three (a triple bond) are quite common. In closely related compounds with bonds between the same kinds of atoms, the bond with the highest bond order is both the shortest and the strongest. In bonds with the same bond order between different atoms, trends are observed that, with few exceptions, result in the strongest single bonds being formed between the smallest atoms. Tabulated values of average bond energies can be used to calculate the enthalpy change of many chemical reactions. If the bonds in the products are stronger than those in the reactants, the reaction is exothermic and vice versa. The breakage and formation of bonds is similar to a relationship: you can either get married or divorced and it is more favorable to be married. Energy is always released to make bonds, which is why the enthalpy change for breaking bonds is always positive. Energy is always required to break bonds. Atoms are much happier when they are "married" and release energy because it is easier and more stable to be in a relationship (e.g., to generate octet electronic configurations). The enthalpy change is always negative because the system is releasing energy when forming bond.

47. It is not possible to measure lattice energies directly. However, the lattice energy can be calculated by using a thermochemical cycle. The Born-Haber cycle is an application of Hess’s law that breaks down the formation of an ionic solid into a series of individual steps: ΔH∘f, the standard enthalpy of formation of the compound IE, the ionization energy of the metal EA, the electron affinity of the nonmetal ΔH∘s, the enthalpy of sublimation of the metal ΔH D, the bond dissociation energy of the nonmetal ΔH lattice, the lattice energy of the compound The Born-Haber Cycle Bond energy is defined as the energy required to break a particular bond in a molecule in the gas phase.

48. The Born-Haber Cycle

49. The Born-Haber Cycle

50. Ionic Bonds • Definition: Attraction between oppositely charged ions after electron transfer. • Strength Measurement: Lattice Energy (U) – energy released when gaseous ions form an ionic solid. • Factors Increasing Strength: • Higher ionic charges • Smaller ionic radii • Compact lattice structure • Formula (Coulomb’s Law form):U = (k × Q₁× Q₂) / r • Example Values: NaCl ≈ 787 kJ/mol; MgO ≈ 3795 kJ/mol STRENGTH OF IONIC AND COVALENT BONDS Ionic bond strength is measured by lattice energy, the amount of energy released when gaseous ions come together to form a solid. Stronger ionic bonds occur when the ions have higher charges, smaller sizes, and pack closely in the lattice. Covalent bond strength is measured by bond dissociation energy (bond enthalpy), the energy required to break one mole of covalent bonds in the gaseous state. Stronger covalent bonds have higher bond orders, shorter bond lengths, and suitable differences in electronegativity