Chemistry Notes PDF

Chemistry Notes -JEWEL SHEEN S. DUGENIO WONG, RRT, LPT Matter – Anything that has mass and occupies space. Mass – The amount of matter in an object. Volume – The amount of space an object occupies. Density – Mass per unit volume of a substance. States of Matter – Solid, liquid, gas, and pla...

Chemistry Notes -JEWEL SHEEN S. DUGENIO WONG, RRT, LPT Matter – Anything that has mass and occupies space. Mass – The amount of matter in an object. Volume – The amount of space an object occupies. Density – Mass per unit volume of a substance. States of Matter – Solid, liquid, gas, and plasma. Solid – State of matter with definite shape and volume. Liquid – State of matter with definite volume but no definite shape. Gas – State of matter with no definite shape or volume. Plasma – Ionized gas with high energy and charged particles. Atom – The smallest unit of an element. Molecule – Two or more atoms bonded together. Element – A pure substance made of only one type of atom. Compound – A substance formed from two or more elements chemically bonded. Mixture – A combination of substances not chemically bonded. Physical Property – A characteristic that can be observed without changing the substance's identity. Chemical Property – A characteristic that describes a substance's ability to undergo chemical changes. Melting Point – The temperature at which a solid turns into a liquid. Boiling Point – The temperature at which a liquid turns into a gas. Freezing Point – The temperature at which a liquid turns into a solid. Condensation – The process of gas turning into a liquid. Sublimation – The transition of a substance from solid to gas without passing through the liquid state. Deposition – The transition of a substance from gas to solid without passing through the liquid state. Evaporation – The process of a liquid turning into a gas at the surface. Viscosity – A measure of a liquid's resistance to flow. Buoyancy – The ability of an object to float in a fluid. Pressure – Force exerted per unit area. Thermal Conductivity – The ability of a substance to conduct heat.  Specific Heat – The amount of heat required to change the temperature of a substance. Chemical Change – A change that results in the formation of a new substance. Physical Change – A change that affects the form of a substance but not its chemical composition. Ionic Compounds Definition: Compounds formed by the transfer of electrons between a metal and a non-metal, resulting in positive (cation) and negative (anion) ions held together by ionic bonds. Examples: Sodium chloride (NaCl), Calcium carbonate (CaCO₃). 2. Covalent (Molecular) Compounds Definition: Compounds formed when two or more non-metal atoms share electrons, creating covalent bonds. Examples: Water (H₂O), Carbon dioxide (CO₂). 3. Acids Definition: Substances that release hydrogen ions (H⁺) when dissolved in water, often characterized by a sour taste and the ability to react with bases to form salts. Examples: Hydrochloric acid (HCl), Sulfuric acid (H₂SO₄). 4. Organic Compounds Definition: Compounds primarily composed of carbon atoms bonded with hydrogen, oxygen, nitrogen, or other elements, typically found in living organisms. Examples: Methane (CH₄), Glucose (C₆H₁₂O₆). 5. Complex Compounds (Coordination Compounds) Definition: Compounds consisting of a central metal atom or ion surrounded by non-metal ions or molecules (ligands) coordinated through covalent bonds. Examples: Potassium ferrocyanide [K₄[Fe(CN)₆]], Copper(II) sulfate pentahydrate (CuSO₄·5H₂O). 1. Quantum Mechanics – The branch of physics that describes the behavior of particles at atomic and subatomic scales. 2. Wave-Particle Duality – The concept that particles, such as electrons, exhibit both wave-like and particle-like behavior. 3. Schrödinger Equation – A fundamental equation that describes how the quantum state of a system evolves over time. 4. Wavefunction (Ψ) – A mathematical function describing the quantum state of a particle and its probability distribution. 5. Atomic Orbital – A region in space where there is a high probability of finding an electron. 6. Principal Quantum Number (n) – Specifies the energy level and size of an orbital. 7. Angular Momentum Quantum Number (l) – Defines the shape of an orbital (e.g., s, p, d, f). 8. Magnetic Quantum Number (mₗ) – Specifies the orientation of an orbital in space. 9. Spin Quantum Number (mₛ) – Describes the intrinsic angular momentum (spin) of an electron. 10. Pauli Exclusion Principle – States that no two electrons in an atom can have the same set of quantum numbers. 11. Heisenberg Uncertainty Principle – States that it is impossible to simultaneously know the exact position and momentum of a particle. 12. Electron Configuration – The arrangement of electrons in an atom's orbitals. 13. Aufbau Principle – Electrons fill orbitals starting with the lowest energy level.  14. Hund's Rule – Electrons occupy degenerate orbitals singly before pairing up. 15. Orbital Hybridization – The mixing of atomic orbitals to form new hybrid orbitals. 16. Valence Electrons – Electrons in the outermost shell of an atom that participate in bonding. 17. Energy Levels – Discrete values of energy that electrons in an atom can have. 18. Photon – A quantum of electromagnetic energy emitted or absorbed by electrons during transitions. 19. Emission Spectrum – The spectrum of light emitted by an atom when its electrons drop to lower energy levels. 20. Quantum Tunneling – The phenomenon where particles pass through energy barriers they classically cannot surmount. ~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~ 1. Ion – An atom or molecule with a net electric charge due to the loss or gain of electrons. 2. Cation – A positively charged ion formed by the loss of electrons. 3. Anion – A negatively charged ion formed by the gain of electrons. 4. Electrostatic Force – The force of attraction between oppositely charged ions. 5. Electron Transfer – The movement of electrons from one atom to another during ionic bond formation. 6. Valence Electrons – Electrons in the outermost shell of an atom, involved in bond formation. 7. Octet Rule – Atoms tend to gain, lose, or share electrons to achieve a stable configuration of 8 valence electrons. 8. Ionic Bond – A chemical bond formed through the electrostatic attraction between oppositely charged ions. 9. Lattice Energy – The energy released when ions form a crystalline lattice. 10. Metal – An element that typically loses electrons to form cations in ionic bonding. 11. Nonmetal – An element that typically gains electrons to form anions in ionic bonding. 12. Crystal Lattice – A three-dimensional arrangement of ions in a structured, repeating pattern. 13. Electronegativity – The ability of an atom to attract electrons in a chemical bond. 14. Ionization Energy – The energy required to remove an electron from an atom to form a cation. 15. Electron Affinity – The energy change that occurs when an atom gains an electron to form an anion. 16. Polyatomic Ion – A charged particle composed of two or more covalently bonded atoms. 17. Ionic Compound – A compound formed by ionic bonds, typically between metals and nonmetals. 18. Solubility – The ability of an ionic compound to dissolve in water or another solvent. 19. Conductivity – The ability of ionic compounds to conduct electricity when molten or dissolved in water. 20. Formula Unit – The simplest ratio of ions represented in an ionic compound. ~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~ 1. Molecular Structure – Arrangement of atoms in a molecule. 2. Covalent Bond – Bond formed by the sharing of electrons between atoms. 3. Intermolecular Forces – Weak forces between molecules (e.g., van der Waals, dipole-dipole, hydrogen bonding). 4. Low Melting Point – Characteristic of molecular compounds due to weak intermolecular forces. 5. Low Boiling Point – Result of weak attractions between molecules. 6. Non-conductive – Molecular covalent compounds typically do not conduct electricity in solid or liquid states. 7. Volatility – Tendency of a substance to vaporize; molecular compounds are often volatile. 8. Solubility – Ability to dissolve in a solvent; depends on polarity (polar dissolves in polar, nonpolar in nonpolar). 9. Polarity – Uneven distribution of electron density in a molecule. 10. Nonpolar Molecule – Molecule with even electron distribution and no net dipole moment. 11. Polar Molecule – Molecule with uneven electron distribution and a net dipole moment. 12. Hydrogen Bonding – Strong intermolecular force between molecules with H bonded to N, O, or F. 13. Dipole-Dipole Interaction – Attraction between polar molecules. 14. London Dispersion Forces – Weak intermolecular forces present in all molecules, stronger in larger molecules. 15. Soft Solids – Molecular solids are often soft due to weak intermolecular forces. 16. Low Density – Molecular compounds often have lower densities compared to ionic or metallic compounds. 17. Insulators – Poor conductors of heat and electricity. 18. Molecular Lattice – Arrangement of molecules in a crystalline structure for some molecular solids. 19. Amorphous Solid – Non-crystalline molecular solid with no regular arrangement. 20. Thermal Stability – Limited stability under heat due to weak intermolecular forces. Adjective is a word that describes a noun/pronoun Examples of Nouns: -doctor -cat Examples of Pronoun: -he , she, it. Examples: 1. He is stressed. HE=PRONOUN 2. He is furious to the other man. HE=PRONOUN 3. The girl is excited to travel. GIRL=NOUN