Chapter 2 (PART ONE)

2.1 Introduction

• Chemistry is the branch of science that deals with the composition and characteristics of chemicals.

• A knowledge of chemistry is necessary for the understanding of physiology because chemicals participate in body processes.

• Our foods, liquids and medications are composed of chemicals.

• Cells and their organelles are assemblies of molecules, which are composed of chemicals.

2.2 Fundamentals of Chemistry

• Matter is anything that has weight (mass) and takes up space (living and nonliving things).

• All matter is composed of elements.

• Elements are the smallest units of matter with specific chemical properties.

• Living organisms require about 20 elements, of which oxygen, carbon, hydrogen, and nitrogen are most abundant.

• Elements are composed of atoms; atoms of different elements vary in size, weight, and interaction with other atoms.

• Atoms are the smallest unit of an element.

• Attractions between two or more atoms are called chemical bonds.

• Atomic Structure (overview): central nucleus contains protons and neutrons; electrons orbit around the nucleus in shells.

• Protons have a positive charge and are about equal in size to neutrons.

• Neutrons are uncharged.

• Electrons are much smaller than protons and neutrons and have a negative charge.

• An electrically neutral atom has equal numbers of protons and electrons.

• Atoms that gain or lose one or more electrons become ions.

• The periodic table organizes elements and is used to summarize properties and relationships among elements.

• The table is divided into groups such as alkali metals, alkaline earth metals, transition metals, lanthanoids, actinoids, nonmetals, and noble gases (referenced in the diagram).

Periodic Table and shell information (concepts referenced in the figures)

• The first electron shell (closest to the nucleus) holds a maximum of 2 electrons.

• The second and third energy shells each hold a maximum of 8 electrons.

• For elements with atomic numbers up to 18, the outermost shell stability (filled shell) leads to chemical inertness.

• The periodic table shows element symbols, atomic numbers, and typical groupings (e.g., hydrogen H; lithium Li; sodium Na; magnesium Mg; etc.).

• Groups and periods determine common properties such as reactivity and bonding tendencies.

Major Elements of the Body (Table 2.1)

• Major elements (by weight) in the human body:

  • Oxygen, symbol: O —

  • Carbon, C —

  • Hydrogen, H —

  • Nitrogen, N —

  • Calcium, Ca —

  • Phosphorus, P —

  • Potassium, K —

  • Sulfur, S —

  • Chlorine, Cl —

  • Sodium, Na —

  • Magnesium, Mg —

• Total ≈ 99.9 ext{%}

Trace Elements (Table 2.1 continuation)

• Trace elements listed: Chromium Cr, Cobalt Co, Copper Cu, Fluorine F, Iodine I, Iron Fe, Manganese Mn, Zinc Zn.

Atomic Structure and Subatomic Particles

• An atom consists of a central nucleus containing protons and neutrons, and electrons in orbit around the nucleus in shells.

• Protons have a positive charge; neutrons are uncharged; electrons have a negative charge.

• An electrically neutral atom has equal numbers of protons and electrons.

• Atoms that gain or lose electrons become ions (charged particles).

• The nucleus contains protons (p) and neutrons (n); electrons (e) orbit in shells around the nucleus.

• The mass primarily comes from protons and neutrons; electrons have very small mass and are not included in atomic mass calculations.

Atomic Number, Mass Number, and Isotopes

• Atomic number (the number of protons)

• Mass number (protons plus neutrons)

• Electron mass is negligible compared with proton/neutron mass; thus it is not included in the calculation of the atomic or mass numbers.

• Isotopes: Atoms with the same atomic number but different mass numbers are isotopes of the same element.

• Atomic weight (atomic mass) in a sample is the weighted average of the masses of the isotopes present

• Isotopes can be stable; some are radioactive and emit energy or atomic fragments.

• For any element, the number of protons (Z) is the same for all isotopes; the neutron number varies with the isotope.

Determining Subatomic Particles

• Proton number (the number of protons) equals the atomic number:

• Neutron number equals the mass number minus the atomic number: N = A - Z.

• Electron number equals the proton number for neutral atoms: e = Z.

• Example: Sodium (Na) with Z = 11 and A = 23 has

  • Neutron number: N = A - Z = 23 - 11 = 12

  • Electron number: e = Z = 11

Bonding of Atoms (Overview)

• Atoms form chemical bonds by gaining, losing, or sharing electrons.

• Electrons are arranged in shells around the nucleus.

• Shell capacities (up to Z = 18):

  • First shell: maximum 2 electrons

  • Second shell: maximum 8 electrons

  • Third shell: maximum 8 electrons

• Atoms with outermost shells that are filled are stable and chemically inert.

Ionic Bonding

• Atoms with incompletely filled outer shells are reactive as they seek to achieve stability by gaining or losing electrons.

• Ions: atoms that gain or lose electrons become charged ions.

• Oppositely charged ions attract and form ionic (electrovalent) bonds:

  • They tend to form crystal lattices (crystal arrays) rather than discrete molecules.

  • The chemical formula reflects the ratio of ions, not the exact number of ions in a sample.

• Example: Sodium (Na) has 1 electron in its outer shell and tends to lose 1 electron to form a positively charged Na⁺ ion, which can bond with Cl⁻ to form NaCl (table salt).

Covalent Bonding

• Covalent bonds form when atoms share electrons to fill their outer shells and achieve stability.

• Bond types by number of shared electron pairs:

  • Single covalent bond: one shared pair of electrons

  • Double covalent bond: two shared pairs

  • Triple covalent bond: three shared pairs

• Example: H₂ is formed when two hydrogen atoms share one pair of electrons.

Polar Covalent and Hydrogen Bonds

• Carbon and hydrogen can form covalent bonds; many C–H bonds are covalent.

• Polar covalent bonds occur when electrons are not shared equally, causing a dipole moment within the bond.

• Polar molecules: molecules containing polar covalent bonds; overall electrically neutral because they contain equal numbers of protons and electrons, but distribution of charge is uneven.

• Hydrogen bonds: a relatively weak attraction between a slightly positive hydrogen atom (attached to a highly electronegative atom like N or O in one molecule) and a nearby electronegative atom (like N or O) in another polar molecule.

• Significance: Hydrogen bonding contributes to properties such as water's cohesion and many biological interactions, though it is weaker than ionic or covalent bonds.

Connections to Foundational Principles and Real-World Relevance

• The chemical basis of life explains why oxygen, carbon, hydrogen, and nitrogen dominate biological tissues and processes.

• Knowledge of bonding explains why water is a universal solvent and how macromolecules like proteins and nucleic acids form and function.

• Ionic vs covalent bonds influence the structure, properties, and behavior of salts, minerals, and many biomolecules.

• Isotopes and atomic weight underpin dating, tracer techniques, and the behavior of elements in biological systems.

Formulas and Key Equations (LaTeX)

• Neutron number: N = A - Z

• Electron number for neutral atoms: e = Z

• Polar covalent vs nonpolar covalent distinction is described by electronegativity differences (qualitative concept; specific values depend on context).

• Atomic weight as a weighted average of isotopes