Chemistry - Chemical Bonds PDF

Examples: Chemical Bonds 1. Water (H₂O): Noble Gases Oxygen: 6 valence electrons; Noble gases are elements found in Group 18 of...

Examples: Chemical Bonds 1. Water (H₂O): Noble Gases Oxygen: 6 valence electrons; Noble gases are elements found in Group 18 of Hydrogen: 1 valence electron. the periodic table. They have complete valence electron shells, which makes them very stable and largely unreactive under normal conditions. 2. Carbon Dioxide (CO₂): Examples: Carbon: 4 valence electrons; 1. Helium (He) Oxygen: 6 valence electrons 2. Argon (Ar) each. 3. Xenon (Xe) Valency and Lewis Structure Valency 3. Ammonia (NH₃): Combining power of an element determined by Nitrogen: 5 valence electrons; the number of electrons in its outermost shell. Hydrogen: 1 each. It indicates how many electrons an atom will tend to gain, lose, or share when forming a chemical bond. Lewis Structure A diagram that shows the bonding between atoms of a molecule and the lone pairs of Ionic Bonds electrons that may exist. Ionic bonds are formed when one atom donates Key Points: electrons to another, creating ions (charged atoms). 1. Octet Rule: Most atoms strive to have The resulting electrostatic attraction between eight electrons in their valence shell. oppositely charged ions holds the compound 2. Representation: Dots represent together. electrons; lines or pairs of dots represent bonds. Key Formula: Formation of an ionic compound:  Examples: Examples: 1. Water (H₂O): 1. Sodium Chloride (NaCl): Oxygen shares electrons with Sodium (Na) loses one two hydrogens (two single electron to become Na⁺, while bonds). chlorine (Cl) gains one 2. Methane (CH₄): electron to become Cl⁻. Carbon shares electrons with 2. Magnesium Oxide (MgO): four hydrogens (four single Magnesium (Mg) loses two bonds). electrons to become Mg²⁺; 3. Carbon Dioxide (CO₂): oxygen (O) gains two Carbon forms double bonds electrons to become O²⁻. with each of two oxygen 3. Calcium Fluoride (CaF₂): atoms. Calcium (Ca) forms Ca²⁺; 4. Nitrogen (N₂): each fluorine (F) forms F⁻, Two nitrogen atoms form a with two F⁻ balancing one triple bond. Ca²⁺. 5. Oxygen (O₂): Two oxygen atoms form a double bond. Covalent Bonds Forms when two nonmetal atoms share electrons to achieve a full outer shell. Polarity These bonds can be single, double, or triple Electronegativity depending on the number of shared electron A measure of an atom’s ability to attract pairs. electrons towards itself when forming a chemical bond. Key Points: a. The greater the difference in 1. Single Bond: One pair of electrons electronegativity between two atoms, shared. the more polar the bond. 2. Double Bond: Two pairs of electrons Examples: shared. 1. Fluorine: Highest electronegativity 3. Triple Bond: Three pairs of electrons 2. Oxygen: High electronegativity shared. 3. Carbon: Moderate electronegativity Bond Polarity Molecular Polarity Determined by the difference in Determined not only by the polarities of electronegativity between two bonded atoms. individual bonds but also by the molecule’s A large difference creates a polar bond. overall shape. A symmetrical shape may cancel out bond Examples: dipoles, resulting in a nonpolar molecule. 1. Hydrogen Chloride (HCl): Significant difference Examples: between H and Cl → polar 1. Water (H₂O): Bent shape → polar bond. molecules. 2. Water (H₂O): Oxygen is more 2. Carbon Dioxide (CO₂): electronegative than Linear geometry → bond hydrogen, resulting in polar dipoles cancel, resulting in a O-H bonds. nonpolar molecule. 3. Hydrogen Molecule (H₂): 3. Ammonia (NH₃): Both atoms have equal Trigonal pyramidal → polar electronegativity → molecules. nonpolar. 4. Methane (CH₄): Tetrahedral and symmetrical → nonpolar. Molecular Geometry 5. Hydrogen Chloride (HCl): Three-dimensional arrangement of atoms within a molecule, which is determined by the Diatomic molecule → polar repulsions between electron pairs. due to difference in electronegativity. Intermolecular Forces Forces of Attraction Lone Pair formula Forces of attraction can be classified into: 1. Intramolecular Forces: Strong bonds that hold atoms together within a molecule (e.g., ionic, covalent bonds). 2. Intermolecular Forces: Weaker attractions between molecules. Johannes van der Waals Examples: Pioneered the study of intermolecular forces, 1. O₂ in Water: leading to the concept of van der Waals forces. Water molecules (polar) Intermolecular Forces: induce a dipole in nonpolar 1. Dipole-Dipole Interactions oxygen. 2. Hydrogen Bonding 2. CO₂ in a Solvent: 3. London Dispersion Forces Even though CO₂ is nonpolar, its electron cloud can be distorted near ions or polar Ion-Dipole Forces molecules. Occur between an ion and a polar molecule. 3. I₂ in Ethanol: They are particularly significant in solutions Ethanol’s polarity can induce where salts dissolve in polar solvents. temporary dipoles in I₂. Examples: 1. Sodium Ion (Na⁺) in Water: The positive ion interacts with Hydrogen Bonding the negative end of the water A strong type of dipole-dipole interaction that dipole. occurs when hydrogen is bonded to a highly 2. Chloride Ion (Cl⁻) in Water: electronegative atom and is attracted to a lone The negative ion interacts pair on another electronegative atom. with the positive end of the water dipole. Dipole-Dipole Forces 3. Potassium Ion (K⁺) in Ammonia (NH₃): Attractions between polar molecules where the Similar ion-dipole positive end of one molecule is attracted to the interactions occur. negative end of another. Examples: Ion-Induced 1. Water (H₂O): When an ion induces a dipole in a neighboring Extensive hydrogen bonding nonpolar molecule by distorting its electron among water molecules. cloud. 2. Ammonia (NH₃): Hydrogen bonds between Dipole-Induced Forces ammonia molecules. When a polar molecule induces a dipole in a 3. Hydrofluoric Acid (HF): nonpolar molecule. Exhibits hydrogen bonding.  4. Acetone (CH₃COCH₃): The acid-base pair before and after the reaction Exhibits dipole-dipole are known as conjugate pairs. interactions between its polar Examples: groups. 1. Acetic Acid (CH₃COOH) and Acetate Ion (CH₃COO⁻) London Dispersion Forces 2. Hydrofluoric Acid (HF) and Fluoride Ion (F⁻) London dispersion forces are weak, temporary 3. Ammonium Ion (NH₄⁺) and Ammonia attractive forces that result from temporary (NH₃) dipoles induced in atoms or molecules. pH They are present in all molecules but are the The pH of a solution is defined as the negative only intermolecular force in nonpolar logarithm (base 10) of the hydrogen ion molecules. concentration. Key Point: a. Strength increases with larger electron clouds and greater polarizability. Examples: 1. Noble Gases: Argon, Neon, etc. 2. Methane (CH₄): Small, nonpolar molecules. 3. Iodine (I₂): Larger, more polarized molecules. 4. Nonpolar organic compounds: E.g., Hexane (C₆H₁₄). Acids and Bases Conjugate Pairs In acid-base reactions, the acid donates a proton (H⁺) and becomes its conjugate base; the base accepts a proton and becomes its conjugate acid.