Atoms, Ions, and Molecules Lecture Notes

Chapter 02: Atoms, Ions, and Molecules

2.1a Matter, Atoms, Elements, and the Periodic Table

• Matter: Defined as anything that has mass and occupies space.

  • Three forms of matter:

  • Solid (example: bone)

  • Liquid (example: blood)

  • Gas (example: oxygen)

• Atom: The smallest particle that exhibits the chemical properties of an element.

• Elements: There are 92 naturally occurring elements that compose all matter, organized on the Periodic Table of Elements.

Periodic Table of Elements

• Structure: Includes symbols, atomic weights, and atomic numbers for the various elements.

  • Each element characterized by:

  • Atomic number: Number of protons.

  • Element's chemical symbol: Unique identifier for each element.

  • Average atomic mass: Shown under each element symbol.

• Arrangement: Elements are organized by atomic number in rows.

• Example Excerpt from the Periodic Table:

  • Hydrogen (H): Atomic number 1; Average atomic mass 1.008

  • Helium (He): Atomic number 2; Average atomic mass 4.003

  • Lithium (Li): Atomic number 3; Average atomic mass 6.941

  • Transition metals and heavier elements are also included, with additional complexities in their structures.

2.1a Components of an Atom

• Basic Structure of Atoms:

  • Atoms are made up of three types of subatomic particles:

  • Neutrons:

    • Mass of 1 atomic mass unit (amu)

    • No charge

  • Protons:

    • Mass of 1 amu

    • Positive charge (+1)

    • Found in the nucleus

  • Electrons:

    • Mass is approximately 1/1800th of a proton/neutron

    • Negative charge (−1)

    • Orbit the nucleus in regions known as electron orbitals.

2.1a Atoms, Elements, and the Periodic Table 2

• Chemical Symbol: Unique for each element, typically contains one or two letters (e.g., C for carbon).

• Atomic Number: Indicates the number of protons in an atom and is listed above the corresponding element symbol.

• Average Atomic Mass: The combined mass of protons and neutrons, located beneath the symbol.

2.1a Determining Subatomic Particles

• Calculating Subatomic Particles:

  • Proton number = atomic number

  • Neutron number = atomic mass − atomic number

  • Example for Sodium (Na): If sodium has an atomic mass of 23 and an atomic number of 11, then:

    • Neutron number = 23 - 11 = 12

  • Electron number = proton number.

2.1a Atomic Structure Modeling

• Electron Shells:

  • Atoms have shells of electrons surrounding the nucleus. Each shell has assigned energy levels and can hold a limited number of electrons:

  • Innermost shell: Holds up to 2 electrons.

  • Second shell: Holds up to 8 electrons.

  • Electrons fill the closest shells first before moving to outer shells.

2.1b Isotopes

• Definition: Isotopes are atoms of the same element with the same number of protons and electrons but differing numbers of neutrons.

• Chemical Properties: Isotopes exhibit identical chemical properties but differ in atomic mass.

• Example of Carbon Isotopes:

  • Carbon-12: 6 protons, 6 neutrons (most abundant)

  • Carbon-13: 6 protons, 7 neutrons

  • Carbon-14: 6 protons, 8 neutrons

2.1b Radioisotopes

• Characteristics: Radioisotopes have excess neutrons and are chemically unstable; they can emit radiation as they decay into stable isotopes.

• Forms of Radiation:

  • Alpha particles

  • Beta particles

  • Gamma rays

• Half-Life:

  • Physical half-life: Time required for 50% of a radioisotope to become stable.

  • Biological half-life: Time taken for half of radioactive material to be eliminated from the body.

2.1c Chemical Stability and the Octet Rule

• Periodic Table Organization: Organized into columns based on electron arrangement in the valence shell.

• Octet Rule: Atoms tend to lose, gain, or share electrons to achieve a valence shell consisting of eight electrons.

  • Stable and unreactive elements include noble gases.

2.1c Chemical Stability and the Octet Rule 2

• Elements with complete outer shells (8 electrons) are chemically stable.

• Fulfillment of the octet rule leads to chemical reactivity or stability based on electron gain/loss.

2.2 Ions and Ionic Compounds

• Definition: Chemical compounds made from associations of two or more elements joined in a fixed ratio, classified into ionic or molecular compounds.

• Ionic compounds are structures held by ionic bonds in a lattice formation.

2.2a Ions

• Definition of Ions: Ions are atoms that gain or lose electrons, thereby acquiring a positive (cation) or negative (anion) charge.

  • Example Functions:

  • Electrolytes play critical roles in physiological processes, such as sodium ions in sports drinks to replace sweat losses.

2.2a Formation of Cations and Anions

• Cations: Formed by the loss of electrons, resulting in a positive charge.

  • Example: Sodium (Na) donates an electron to reach stability, leading to +1 charge.

• Anions: Formed by the gain of electrons, resulting in a negative charge.

  • Example: Chlorine (Cl) gains an electron to achieve stability, resulting in a -1 charge.

2.2b Ionic Bonds

• Ionic Bonds: Result from the electrostatic attraction between cations and anions.

  • Example: Sodium chloride (NaCl), where Na donates one electron to Cl, forming a rigid crystal lattice structure.

2.3 Covalent Bonding, Molecules, and Molecular Compounds

• Covalent Bonds: Involve shared electrons between atoms of different elements, resulting in molecular compounds (e.g. CO2).

2.3a Chemical Formulas: Molecular and Structural

• Molecular Formula: Indicates the number and type of atoms of each element (e.g. H2O).

• Structural Formula: Shows the arrangement of atoms within a molecule (e.g., O=C=O in CO2).

2.3b Covalent Bonds

• Nature of Covalent Bonds: Involves sharing electrons, commonly formed among biological elements: H, O, N, C.

• Types:

  • Single Bond: One pair of electrons shared (e.g. H2).

  • Double Bond: Two pairs shared (e.g. O2).

  • Triple Bond: Three pairs shared (e.g. N2).

2.3b Covalent Bonds 2 and 3- Carbon Structures

• Carbon Skeleton: Carbon can form chains or rings, allowing for complex biological structures and the flexibility of compounds.

2.3b Nonpolar and Polar Covalent Bonds

• Electronegativity determines electron sharing; equal sharing yields nonpolar covalent bonds, while unequal sharing results in polar covalent bonds.

• Electronegativity Order (most to least): H < C < N < O

2.3c Nonpolar, Polar, and Amphipathic Molecules

• Nonpolar Molecules: Molecules containing only nonpolar covalent bonds.

  • Example: O—O and C—H bonds.

• Polar Molecules: Contain polar covalent bonds, e.g., water (H2O).

• Amphipathic Molecules: Have both polar and nonpolar regions. Example: phospholipids.

2.3d Intermolecular Attractions

• Definition: Weak attractions between molecules crucial for molecular shape and function (e.g. hydrogen bonding).

• Hydrophobic Interactions: Occur between nonpolar molecules in polar substances.

2.4 Molecular Structure and Properties of Water

• Water: Makes up about two-thirds of the human body by weight; it is a polar molecule due to unequal sharing of electrons.

2.4b Properties of Water

• Phases of Water: Three forms based on temperature: gas (vapor), liquid, and solid (ice).

  • Liquid Water Properties:

  • Transports, lubricates, cushions body structures, and helps excrete wastes.

• Cohesion and Adhesion: Water molecules exhibit strong attractive forces to each other, leading to surface tension, while adhesion allows interaction with other molecules.

2.4b High Specific Heat and Heat of Vaporization

• Specific Heat: High amount of energy required to raise water temperature, helping to stabilize body temperature.

• Heat of Vaporization: High energy required to convert water from liquid to gas, aiding in cooling through evaporation.

2.4c Water as the Universal Solvent

• Solvent Properties: Water is known as a universal solvent because it can dissolve many substances, particularly polar molecules and ions.

  • Hydrophilic Substances: Water-loving substances that dissolve easily (e.g., glucose).

  • Hydrophobic Substances: Water-fearing substances that do not dissolve (e.g., oils).

  • Amphipathic Molecules: Partially dissolve with both hydrophilic and hydrophobic regions, as in phospholipids forming cell membranes.