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Chemistry Chapter 2 Part 1: Basic Chemistry Overview

CHAPTER 2 PART 1 — BASIC CHEMISTRY OVERVIEW

  • Objectives covered in this chapter

    • Describe types of energy: kinetic versus potential
    • Explain the difference between an element, an atom, a molecule, and a compound
    • Identify the four main elements of the body and understand atomic structure
    • Explain three different types of chemical bonds
    • Explain the three types of chemical reactions in the body and factors that increase or decrease reaction rates
    • Recognize that lectures include details beyond pure chemistry (applications, biology)

  • Matter and energy (chemistry as the science of matter and energy)

    • Matter: anything that occupies mass and space
    • Energy: the capacity to do work or to put matter into motion
    • Human chemistry is carbon-based and includes organic molecules and biochemicals (hormones, enzymes, neurotransmitters)
    • States of matter: solid (definite shape and volume), liquid (definite volume, variable shape), gas (variable shape and volume)
    • In the body, energy is stored as ATP and constantly cycled between ATP and ADP

  • Forms of energy and energy conversion in the body

    • Four main forms of energy (all convertible into one another):
    • Chemical energy: energy stored in chemical bonds of molecules (e.g., food)
      • When we eat, food is potential energy; chemical energy is released when bonds form/break during metabolism
    • Electrical energy: movement of charged particles (e.g., ions across membranes like Na⁺ and K⁺ across membranes)
    • Mechanical energy: energy directly involved in moving matter (e.g., muscles pulling on bones to move)
      • Chemical energy can be converted into mechanical energy to power movement
    • Radiant (electromagnetic) energy: energy carried by electromagnetic waves; visible light on retina is electromagnetic energy; UV light supports vitamin D synthesis
    • All four forms can be converted from one to another in the body (e.g., chemical → mechanical)

  • Structural organization reminder

    • Levels: atoms → molecules → cells → tissues → organs → organ systems
    • We will discuss atoms and molecules next, then build toward cells and tissues

  • Atoms, elements, molecules, and compounds (key concepts introduced)

    • Elements: basic substances that cannot be simplified further
    • 92 naturally occurring elements; body cannot synthesize all elements; we obtain many from food and environment
    • Atomic number (Z): number of protons in the nucleus; shown in the upper-left corner of element symbols on the periodic table; also equals the number of protons
    • Atomic symbol: one- or two-letter shorthand (e.g., H for hydrogen, He for helium, C for carbon, N for nitrogen, O for oxygen, K for potassium)
    • Four elements that make up about 96% of body mass: H, C, N, O
    • Subatomic particles
    • Protons: positive charge, located in the nucleus
    • Neutrons: neutral charge, located in the nucleus
    • Electrons: negative charge, orbit the nucleus in shells (valence shell dating and electron cloud models)
    • Nucleus and electron shells
    • Nucleus contains protons and neutrons
    • Electrons are outside the nucleus in electron shells (valence shell crucial for bonding)
    • Planetary model vs. orbital (electron cloud) model exist; emphasis on planetary model in this lecture
    • Recognizing the atomic number and symbol
    • Example: Hydrogen (H) has Z = 1; Helium (He) Z = 2; Phosphorus (P) Z = 15; Potassium (K) Z = 19
    • Small but important note on common examination cues
    • Don’t memorize every element’s properties for the exam; focus on recognizing atomic number, symbol, and the concept of protons, neutrons, and electrons

  • The mass number and isotopes

    • Mass number (A): sum of protons and neutrons in the nucleus
    • A = Z + N
    • Examples:
      • Hydrogen: Z = 1, N varies; typical A = 1 (protium), 2 (deuterium), 3 (tritium)
      • Helium: Z = 2, usually N = 2 → A = 4
      • Lithium: Z = 3, common N = 4 → A = 7
      • Carbon: Z = 6, common N = 6 → A = 12
      • Neon: Z = 10, common N = 10 → A = 20
    • Isotopes: structural variants of an element with the same Z (same protons) but different numbers of neutrons (N)
    • If N differs from Z, the atom is an isotope of that element
    • Example: Hydrogen with different neutron counts is still hydrogen but is an isotope of hydrogen
    • Important exam takeaway: mass number involves protons and neutrons only; electrons do not affect the mass number
    • Isotopes recognition: when neutrons differ from protons, you’re looking at an isotope; otherwise it’s the standard element

  • Atoms, molecules, and compounds

    • Atoms form molecules and compounds by bonding
    • Molecule: two or more atoms of the same element bonded together (e.g., H₂, O₂; CH₄ as a molecule of multiple atoms)
    • Compound: two or more different kinds of atoms bonded together (e.g., H₂O, CO₂, C₆H₁₂O₆)
    • Glucose example: C₆H₁₂O₆ is a molecule; when combined with other atoms, it forms compounds
    • Mixtures and their classifications (substances retain chemical identities but mix physically)
    • Solutions: homogeneous mixtures where solute is dissolved in solvent (e.g., salt in water)
    • Colloids: heterogeneous mixtures with large macromolecules or particles that don’t settle quickly (e.g., gelatin, whipped cream, milk, mist)
    • Suspensions: heterogeneous mixtures with large particles that settle over time (e.g., sugar in cold tea, blood sediments)
    • Solvent vs solute
    • Solvent: liquid in which the solute is dissolved (e.g., water)
    • Solute: the substance dissolved in the solvent (e.g., salt, sugar)
    • Homogeneous mixtures vs heterogeneous mixtures
    • Homogeneous: even distribution of solute throughout solvent; you cannot distinguish solute particles visually
    • Heterogeneous: uneven distribution; particles are visible; suspensions are a subtype with visible sedimentation
    • Visual examples and distinctions
    • Jell-O: colloid (solutes not fully dissolved; appears cloudy/opaque)
    • Pumice in air; mist; whipped cream: colloids (small particles but not fully dissolved)
    • Salt in water at room temperature: typically a solution (solute dissolves); in cold tea with sugar, sugar may not fully dissolve leading to a suspension-like appearance until dissolved (temperature and particle size influence outcomes)
    • Summary distinctions
    • Colloid: smaller particles; heterogeneous but does not settle (appears cloudy or opaque)
    • Suspension: larger particles; heterogenous; sediments visible
    • Solution: small particles; homogeneous; particles fully dissolved

  • Chemical bonds: ionic, covalent (polar and nonpolar), and hydrogen bonds

    • Ionic bonds (electrostatic attraction)
    • Complete transfer of one or more electrons from one atom to another
    • Example: Na⁺ (sodium) donates an electron to Cl⁻ (chloride) → NaCl (table salt)
    • Result: cation (positive charge) and anion (negative charge) attracted to each other
    • On the exam you’ll identify anions (negative charge) and cations (positive charge) by their charges, not by the identity of the element
    • Covalent bonds (sharing electrons)
    • Sharing of electrons between atoms to fill outer shells
    • Nonpolar covalent bond: equal sharing of bonding electrons; no partial charges
      • Example: CH₄ (carbon sharing electrons with four hydrogens)
    • Polar covalent bond: unequal sharing; creates partial positive and partial negative charges on atoms
    • Hydrogen bonds
    • A weaker type of bond; intermolecular force between molecules (e.g., water) or intramolecular interactions that influence three-dimensional structures
    • Important for properties like water surface tension and high heat capacity; relates to form and function in biomolecules
    • Key takeaways about bond strengths (summary visual on slide)
    • Ionic bonds are generally strongest, followed by covalent bonds (stronger if nonpolar, often strong), and hydrogen bonds are weaker
    • The octet rule (rule of eight) (electrons in shells)
    • Innermost shell typically holds 2 electrons and is stable when full
    • Outer shells tend to want 8 electrons to be stable (octet rule)
    • Potassium example: outermost shell has only 1 electron, making it highly reactive and seeking to attain stability (either gain or share electrons)
    • Atoms will interact to achieve a full outer shell (eight electrons, or two in the first shell)
    • Inert elements (stability) and reactivity
    • Stable (inert) elements include helium and neon when their valence shell is full (2 in the first shell, 8 in the outer shell)
    • Elements with incomplete outer shells are chemically reactive and will gain, lose, or share electrons to reach stability
    • Examples reinforcing bonding concepts
    • H₂O: hydrogen and oxygen form polar covalent bonds; hydrogen bonds contribute to water’s properties
    • CH₄: methane formed by covalent bonds between carbon and hydrogen
    • NaCl: ionic compound formed by transfer of an electron from sodium to chlorine

  • Chemical reactions and energy considerations

    • Chemical reactions: bonds are formed, rearranged, or broken; represented by chemical equations
    • Reactants → Products
    • Equations include molecular formulas for reactants and products
    • Exergonic vs. endergonic reactions
    • Exergonic: energy released (ΔG < 0); reaction releases more energy than it absorbs
    • Endergonic: energy absorbed (ΔG > 0); requires energy input (e.g., ATP used in cellular processes)
    • Types of chemical reactions
    • Synthesis (anabolic): small reactants combine to form a larger product; involves bond formation
      • Example: A + B → AB
    • Decomposition (catabolic): a large molecule is broken into smaller pieces; bonds broken and energy released or used; often energy input is required to drive bonds breaking
      • Example: AB → A + B
    • Exchange (displacement or replacement): bonds are broken and formed, and components are rearranged
      • Example: AB + CD → AD + CB
    • ATP/ADP as a practical example of energy transfer in reactions
    • ATP (adenosine triphosphate) structure: adenosine + three phosphate groups
    • ADP (adenosine diphosphate) + Pi are produced when one phosphate is removed in an energy-releasing step
    • Notation: ATP ⇄ ADP + Pi (phosphate) with energy transfer driving cellular work
    • ATP hydrolysis typically releases energy used for cellular processes
    • What influences reaction rates
    • Temperature: higher temperature generally increases rate (e.g., faster solubility, faster reaction kinetics)
    • Particle size: smaller particles react faster; larger particle size slows reaction rates
      • Example concept: smaller solutes dissolve faster than large chunks
    • Concentration: higher concentration generally increases rate
    • Catalysts and enzymes: substances that speed up reactions without being consumed; enzymes are biological catalysts
      • Enzymes lower activation energy required for the reaction
      • Enzymes are not consumed during the reaction; they are available for subsequent reactions

  • pH and hydrogen context

    • Hydrogen relates to pH (potential of hydrogen); pH scale is introduced as the measure of hydrogen ion concentration in solutions
    • Hydrogen bonds and pH are linked to biomolecular structure and function; more detailed discussion occurs later in the course

  • Quick recap of key formulas and concepts to remember (LaTeX)

    • Mass number: A = Z + N where Z = atomic number (protons), N = neutrons
    • Atomic number: Z = ext{number of protons}
    • Isotope: variant of an element with different number of neutrons; same Z but different N
    • Common molecule/compound examples:
    • Water: ext{H}_2 ext{O}
    • Methane: ext{CH}_4
    • Glucose: ext{C}6 ext{H}{12} ext{O}_6
    • Sodium chloride (table salt): ext{NaCl}
    • Energy flow in cells (representative): ATP → ADP + Pi; energy release powers work
    • Bond types (simplified strength order): Ionic > Covalent > Hydrogen (note: covalent includes nonpolar and polar cases)
    • Four energy forms to remember: E{ ext{chemical}}, E{ ext{electrical}}, E{ ext{mechanical}}, E{ ext{radiant}}

  • Connections to broader biology

    • The chemistry of atoms, bonds, and reactions underpins all biological macromolecules (carbohydrates, proteins, lipids, nucleic acids)
    • The form of a molecule (shape and bonding) influences its function in the body (form follows function)
    • Understanding energy storage and transfer (ATP/ADP) is foundational for metabolism and physiology

  • Practical implications and examples discussed in lecture

    • Nutrient intake and storage as energy (potential energy stored in fats)
    • The sodium–potassium pump as an example of electrical energy use across membranes
    • Blood sedimentation rate as a clinical indicator connected to colloids/suspensions in blood
    • Real-world applications of mixtures (solutions, colloids, suspensions) in pharmaceuticals and everyday life

  • Final notes for exam prep

    • Be able to distinguish between element, atom, molecule, and compound
    • Recognize and interpret atomic number, mass number, and isotopes from context
    • Identify ionic, polar covalent, nonpolar covalent, and hydrogen bonds from descriptions
    • Differentiate between synthesis, decomposition, and exchange reactions; identify exergonic vs. endergonic processes
    • Understand the octet rule and why stability drives bonding behavior
    • Apply concepts of solutions, colloids, and suspensions with examples (solvent vs solute; particle size effects)
    • Recall ATP/ADP dynamics and the general role of catalysts (enzymes) in speeding reactions without being consumed