Lecture 3_ Matter and chemical reactions (1)

Unit 1: Matter and Chemical Reactions

Table of Contents

  • Basic particles of matter

  • Chemical bonds: covalent bonds

  • Chemical bonds: hydrogen bonds

  • Chemical bonds: ionic bonds

  • Chemical reactions

1. Basic Particles of Matter

  • Matter: Anything that occupies space and has mass.

  • Element: A substance that cannot be broken down into simpler substances by chemical reactions.

  • Atom: The smallest unit of matter that retains the properties of an element.

  • Molecule: Two or more atoms held together by shared electrons.

  • Compound: A substance consisting of two or more different elements combined in a fixed ratio.

The Bohr Model of Atoms

  • Resembles a solar system with negatively charged electrons orbiting a dense nucleus which contains positively charged protons and neutral neutrons.

  • Quantum mechanics introduced orbital theory expanding the understanding of atomic structure by emphasizing electron clouds.

Atom Structure

  • Nucleus Composition: Protons (1 amu) and neutrons (1 amu) create a dense atomic nucleus.

  • Electrons: Form a negatively charged cloud around the nucleus; they do not have mass.

  • Atomic Number: Number of protons (also equals number of electrons).

  • Atomic Mass: Sum of protons and neutrons; measures the "heaviness" of an atom.

Ions and Isotopes

  • Ions: Atoms that have gained or lost electrons.

    • Example: Sodium (Na+) has lost an electron, Calcium (Ca2+) has lost two.

  • Isotopes: Atoms with an unusual number of neutrons.

    • Example: Carbon-12 (6neutrons), Carbon-13 (7 neutrons), Carbon-14 (8 neutrons).

2. Chemical Bonds

  • Definition: Attraction between atoms that holds them together to form molecules.

  • Types of Chemical Bonds:

    • Covalent Bonds

    • Hydrogen Bonds

    • Ionic Bonds

Valence Electrons

  • Valence Electrons: Electrons in the outermost orbitals of an atom that can be involved in chemical bonding.

  • Covalent Bond: Formed when two atoms share valence electrons.

    • Merging two atomic orbitals into one containing electrons results in a stable covalent bond.

  • Electron Sharing Rule: General pattern of 2, 8, 8, 18 electron shells.

Covalent Bonds

  • Shared electrons between two hydrogen atoms or between hydrogen and carbon demonstrate covalent bonding.

  • Sharing is not always equal; some atoms may "hog" shared electrons.

Electronegativity and Polar Bonds

  • Electronegativity: The ability of an atom to attract electrons towards itself.

    • Example: In water (H2O), electrons are drawn closer to the oxygen atom than to hydrogen atoms due to its higher electronegativity.

  • Polar Covalent Bonds: Occur when electron sharing is unequal, leading to dipoles (partial positive/negative charges).

  • Pure or symmetrical molecules (e.g., O2, H2) exhibit equal sharing of electrons.

3. Chemical Bonds: Hydrogen Bonds

  • Hydrogen atoms (δ+) in polar covalent bonds can attract negatively charged atoms (δ-), forming hydrogen bonds.

  • Examples: Water (H2O) and Ammonia (NH3).

  • Importance of Hydrogen Bonds: Critical for the structure of proteins and the properties of water, making them essential to life.

4. Chemical Bonds: Ionic

  • Ionic Bonds: Formed between metals and nonmetals with substantial differences in electronegativity.

    • Electrons are transferred, leading to the formation of cations (positive) and anions (negative).

    • Example: Sodium chloride (NaCl) forms through ionic bonding.

  • Ionic compounds exist as crystals due to cation-anion interactions.

  • Solubility of Salts: Ionic bonds can be disrupted by hydrogen bonding when ionic compounds are dissolved in water.

5. Chemical Reactions

  • Definition: Rearrangement of atoms to form new compounds with different properties.

    • Evidence of chemical reactions includes heat production, color change, gas production, and precipitate formation.

  • Chemical Equation: Represented as x[A] + y[B] → z[AB], indicating reactants converting to products.

Energy in Reactions

  • Endergonic Reactions: Require energy input (non-spontaneous).

  • Exergonic Reactions: Release energy (spontaneous).

  • Reactions can be coupled to allow energy from one to drive another (e.g., ATP synthesis).

Equilibrium in Reactions

  • Reversible reactions can reach equilibrium, where product concentration stabilizes with reactant concentration.

  • Le Chatelier’s Principle: States that changes in concentration of reactants/products can shift the direction of equilibrium.

Importance of Chemical Reactions

  • Responsible for creating and breaking down biomolecules and macromolecules.

  • Inorganic Compounds: Lack carbon and hydrogen in specific bonds.

  • Organic Compounds: Contain carbon bound to hydrogen, forming essential biomolecules.

  • Type of Reactions:

    • Synthesis Reactions (Anabolic): Smaller molecules combine into larger molecules.

      • Example: Dehydration synthesis, where water is released.

    • Decomposition Reactions (Catabolic): Larger molecules break down into smaller fragments.

      • Example: Hydrolysis, where water is consumed.

Review Summary

  1. Atoms: Smallest unit of matter, composed of protons, neutrons, and electrons.

  2. Matter consists of elements, which have atoms as basic units.

  3. Atoms interact through sharing or transferring electrons to form molecules and compounds.

  4. Ions have atypical electron counts, while isotopes have unusual neutron counts.

  5. Ionic and covalent bonds are strong, with unevenly shared electrons forming weaker hydrogen bonds.