2.1

2.1 Basic Chemistry

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

  • Composed of chemical elements, which are further composed of atoms.

• Primary Elements:

  • In Earth's crust: Silicon (Si), Aluminum (Al), Oxygen (O)

  • In organisms: Carbon (C), Hydrogen (H), Oxygen (O), Nitrogen (N), with Phosphorus (P) and Sulfur (S) also present.

  • These elements are essential building blocks of biological molecules.

2.2 Atomic Structure

• Sub-atomic Particles:

  • Protons: Located in the nucleus, positively charged.

  • Electrons: Orbit the nucleus in shells, negatively charged.

  • Neutrons: Located in the nucleus, have no charge.

• Atomic Number:

  • Defined as the number of protons in the nucleus of an atom.

• Mass Number:

  • Defined by the formula: $\text{mass number} = \text{number of protons} + \text{number of neutrons}$

  • Example: Carbon, represented by the symbol C.

2.3 Isotopes

• Isotopes:

  • Defined as atoms of the same element that have the same number of protons but different numbers of neutrons, resulting in a different mass number.

• Unstable Isotopes:

  • Change over time into stable isotopes, a process that releases energy in the form of rays and subatomic particles.

  • These isotopes are classified as radioactive isotopes.

2.4 Radioactivity

• Uses of Radioactive Isotopes:

  • Low Radiation Levels:

  • Can be used as tracers in medical testing and imaging.

  • Example: In a thyroid scan, a missing area may indicate a tumor that does not absorb radioactive iodine.

  • High Radiation Levels:

  • Radiation can damage cells but is used for sterilizing medical and dental products.

  • Example: Radiation therapy is utilized to kill cancer cells.

  • Potential harmful effects include DNA damage that may lead to cancer.

2.5 Check Your Progress

• Questions:

  1. Discuss the differences in key elements found in Earth’s crust compared to those in living organisms.

  2. Describe how radiation can have both beneficial and harmful effects on human health.

  3. Compare and contrast oxygen-16 and oxygen-18.

2.6 Molecules and Compounds

• Formation of Compounds:

  • Formed when atoms of two or more different elements bond.

• Types of Bonds:

  • Ionic Bonds:

  • Electrons transfer between atoms, leading to the formation of ionic compounds.

  • Covalent Bonds:

  • Electrons are shared between atoms, forming molecules.

  • Example of covalent bonding includes molecules formed by two or more of the same atom (e.g., H2, O2, O3).

2.7 Ionic Bonding

• Ions:

  • Defined as charged particles formed when electrons are transferred from one atom to another.

  • Example: Sodium (Na) donates electrons while Chlorine (Cl) accepts electrons, creating positively and negatively charged ions.

  • Ionic Bonds:

  • Attractive forces between positively and negatively charged ions.

2.8 Covalent Bonding

• Formation of Covalent Bonds:

  • A covalent bond is created when two atoms share electrons.

  • Visual Representation:

  • A single covalent bond may be represented as a line between two atoms.

  • A double covalent bond represents two pairs of shared electrons, and a triple covalent bond represents three pairs.

2.9 Shapes of Molecules

• Three-Dimensional Shapes:

  • Molecules exhibit three-dimensional shapes important for their structural and functional roles.

  • Linear: Molecules with two atoms are linear.

  • Tetrahedral: For instance, methane (CH4) has a tetrahedral shape due to four single covalent bonds.

2.10 Nonpolar and Polar Covalent Bonds

• Nonpolar Covalent Bonds:

  • Occur when electrons are shared equally between atoms.

• Polar Covalent Bonds:

  • Occur when the sharing of electrons is unequal.

  • Example: In water, oxygen has greater electronegativity than hydrogen, resulting in a polar covalent bond.

  • Consequence: The more electronegative atom acquires a partial negative charge, and the other atom obtains a partial positive charge.

2.11 Hydrogen Bonding

• Hydrogen Bonds:

  • Form due to the polarity within water molecules, where hydrogen atoms of one molecule are attracted to oxygen atoms of another molecule.

  • Characteristics:

  • Weaker than ionic or covalent bonds, often represented as dotted lines.

  • Collectively, many hydrogen bonds can confer considerable strength.

  • Importance: Hydrogen bonds play a significant role in the stability of many biological molecules, particularly those with polar covalent bonds involving hydrogen, oxygen, or nitrogen.

2.12 Check Your Progress

• Questions:

  1. Determine whether carbon dioxide (CO2) and nitrogen gas (N2) are classified as molecules, compounds, or both.

  2. Analyze the characteristics that lead hydrogen ions to form polar bonds with a partially positive charge.

2.13 Key Terms

• molecule: A group of atoms bonded together.

• compound: A substance formed when two or more different elements combine chemically.

• ions: Charged particles resulting from the loss or gain of electrons.

• ionic bond: An attraction between oppositely charged ions.

• covalent bond: A bond formed by the sharing of electrons between atoms.

• hydrogen bond: A weak attraction between a slightly positive hydrogen atom and a slightly negative atom (e.g., oxygen or nitrogen).

• electronegativity: A measure of the tendency of an atom to attract a bonding pair of electrons.

2.14 Key Concepts

• Ion Formation: Atoms lose or gain electrons to fill their outer orbital, resulting in ion formation for bonds.

• Ionic Bond: Attracted forces between ions of opposite charges.

• Covalent Bonding: Involves shared electron pairs, resulting in molecule shapes that influence biological functions.

• Polar vs Nonpolar Bonds: Electrons may be shared equally or unequally, defining bond characteristics and creating partial charges.

• Hydrogen Bond Importance: Weaker than ionic and covalent bonds but critical for maintaining cellular structure and function in biological molecules.