Chap2 biology

The Chemical Context of Life

Overview: A Chemical Connection to Biology

Understanding the chemical basis of life is fundamental to biology. Living organisms and their environments are governed by the principles of chemistry and physics. Biology draws upon insights from other sciences to explain the structure and function of life at all levels.

• Life's Hierarchical Organization: Atoms → Molecules → Organelles → Cells → Tissues → Organs → Organisms.

• Emergent Properties: New properties arise at each level of organization that are not present at the previous level.

• Transition from Chemistry to Biology: The boundary between non-living and living matter is crossed as molecules become organized into cells.

Concept 2.1: Matter Consists of Chemical Elements in Pure Form and in Combinations Called Compounds

Elements and Compounds

All matter is composed of elements, which are substances that cannot be broken down into other substances by chemical reactions. Compounds are substances consisting of two or more elements combined in a fixed ratio.

• Element: A pure substance made of only one kind of atom.

• Compound: A substance consisting of two or more elements in a fixed ratio (e.g., NaCl).

• 92 Naturally Occurring Elements: Each element has a unique symbol, often derived from Latin or German names.

• Emergent Properties: Compounds have characteristics different from those of their constituent elements (e.g., Na + Cl → NaCl, edible table salt).

Essential Elements of Life

• About 20-25% of the 92 elements are essential elements required for life.

• Humans need 25 essential elements; plants need 17.

• Major Elements: Oxygen (O), Carbon (C), Hydrogen (H), and Nitrogen (N) make up about 96% of living matter.

• Trace Elements: Required in minute quantities (e.g., iron (Fe), iodine (I)).

• Some trace elements are required only by certain species (e.g., vertebrates require iodine for thyroid function).

Concept 2.2: An Element's Properties Depend on the Structure of Its Atoms

Atomic Structure

Atoms are the smallest units of matter that retain the properties of an element. Each atom consists of a dense nucleus containing protons and neutrons, surrounded by electrons.

• Subatomic Particles:

  • Proton: Positive charge, mass ≈ 1 dalton.

  • Neutron: No charge, mass ≈ 1 dalton.

  • Electron: Negative charge, negligible mass.

• Atomic Number (Z): Number of protons in the nucleus.

• Mass Number (A): Sum of protons and neutrons.

• Atomic Mass: Approximate total mass of an atom, measured in daltons.

Isotopes and Radioactivity

• Isotopes: Atoms of the same element with different numbers of neutrons.

• Stable Isotopes: Do not decay spontaneously.

• Radioactive Isotopes: Decay spontaneously, emitting particles and energy.

• Applications: Used in biological research, medical diagnostics, and dating fossils.

Electron Configuration and Chemical Behavior

The chemical behavior of an atom is determined by the distribution of electrons in electron shells. Electrons have potential energy due to their position relative to the nucleus.

• Electron Shells: Energy levels where electrons are found.

• Valence Electrons: Electrons in the outermost shell; determine chemical reactivity.

• Octet Rule: Atoms tend to fill their valence shell with 8 electrons (except for hydrogen and helium).

• Periodic Table: Arranged by increasing atomic number; elements in the same column have similar valence electron configurations and chemical properties.

Table: Subatomic Particles

Particle	Charge	Mass (daltons)	Location
Proton	+1	1	Nucleus
Neutron	0	1	Nucleus
Electron	-1	~0	Electron cloud

Concept 2.3: The Formation and Function of Molecules Depend on Chemical Bonding Between Atoms

Chemical Bonds

Atoms with incomplete valence shells can interact by sharing or transferring valence electrons, resulting in chemical bonds.

• Covalent Bonds: Sharing of a pair of valence electrons between atoms.

• Molecule: Two or more atoms held together by covalent bonds.

• Single Bond: Sharing of one pair of electrons (e.g., H2).

• Double Bond: Sharing of two pairs of electrons (e.g., O2).

• Nonpolar Covalent Bond: Electrons shared equally (e.g., H2, O2).

• Polar Covalent Bond: Electrons shared unequally due to differences in electronegativity (e.g., H2O).

Ionic Bonds

• Ions: Atoms or molecules with electrical charge due to loss or gain of electrons.

• Cation: Positively charged ion (loss of electron).

• Anion: Negatively charged ion (gain of electron).

• Ionic Bond: Attraction between oppositely charged ions (e.g., Na+ and Cl- form NaCl).

• Salts: Compounds formed by ionic bonds; often found as crystals.

Other Types of Chemical Interactions

• Hydrogen Bonds: Weak bonds between a hydrogen atom covalently bonded to an electronegative atom and another electronegative atom (e.g., between water molecules).

• Van der Waals Interactions: Weak attractions between molecules or parts of molecules that result from transient local partial charges.

Summary Table: Types of Chemical Bonds

Bond Type	Strength	Description	Example
Covalent	Strong	Sharing of electron pairs	H2, O2, H2O
Ionic	Strong (in dry conditions)	Transfer of electrons; attraction between ions	NaCl
Hydrogen	Weak	Attraction between H and electronegative atom	Between water molecules
Van der Waals	Very weak	Transient attractions due to partial charges	Gecko feet adhesion

Key Equations

• Atomic Number: $Z = \text{number of protons}$

• Mass Number: $A = \text{number of protons} + \text{number of neutrons}$

• Atomic Mass (approximate): $\text{Atomic mass} \approx A$

Examples and Applications

• Radioactive Isotopes: Used in PET scans to monitor metabolic processes.

• Trace Elements: Iodine deficiency can cause goiter in humans.

• Hydrogen Bonds: Responsible for water's unique properties, such as high surface tension and specific heat.

Additional info: The notes above expand on the original content by providing definitions, examples, and context for key terms and concepts, as well as organizing the material into a logical, textbook-style structure for effective study.