The Chemical Basis of Life

Matter, Elements, and Compounds

Chemicals make up all physical matter, including human bodies, the bodies of other living organisms, and the physical environment.
Living organisms are composed of matter, which is defined as anything that occupies space and has mass (or weight).
Matter exists in three physical states:
- Solid
- Liquid
- Gas
Matter is composed of chemical elements:
- An element is a pure substance that cannot be broken down into other substances by ordinary chemical means.
- There are $92$ naturally occurring elements in nature, but only a few exist in a pure state.
A compound is a substance consisting of two or more different elements combined in a fixed ratio.
- Compounds are significantly more common in nature than pure elements.
- Sodium chloride, or table salt, is a common compound composed of equal parts of sodium ( $Na$ ) and chlorine ( $Cl$ ).

Chemical Elements in the Human Body

Approximately $25$ elements are essential to human life.
Four elements make up about $96\%$ (specifically $96.3\%$ ) of the total weight of the human body (including water):
- Oxygen ( $O$ ): $65.0\%$
- Carbon ( $C$ ): $18.5\%$
- Hydrogen ( $H$ ): $9.5\%$
- Nitrogen ( $N$ ): $3.3\%$
Seven major elements constitute the remaining $3.7\%$ of human body weight:
- Calcium ( $Ca$ ): $1.5\%$
- Phosphorus ( $P$ ): $1.0\%$
- Potassium ( $K$ ): $0.4\%$
- Sulfur ( $S$ ): $0.3\%$
- Sodium ( $Na$ ): $0.2\%$
- Chlorine ( $Cl$ ): $0.2\%$
- Magnesium ( $Mg$ ): $0.1\%$
Trace elements are essential to life processes but are required only in minute quantities, making up less than $0.01\%$ of human body weight. Trace elements include:
- Boron ( $B$ ), Chromium ( $Cr$ ), Cobalt ( $Co$ ), Copper ( $Cu$ ), Fluorine ( $F$ ), Iodine ( $I$ ), Iron ( $Fe$ ), Manganese ( $Mn$ ), Molybdenum ( $Mo$ ), Selenium ( $Se$ ), Silicon ( $Si$ ), Tin ( $Sn$ ), Vanadium ( $V$ ), and Zinc ( $Zn$ ).

Pathologies associated with trace element imbalances and clinical conditions:
- Goiter: An enlargement of the thyroid gland caused by iodine deficiency in the diet.

*   **Hyperthyroidism / Graves' Disease**: An overactive thyroid condition resulting in symptoms such as exophthalmos (protruding eyes).

Atomic Structure and Subatomic Particles

Each chemical element consists of one specific kind of atom.
An atom is the smallest unit of matter that still retains all the unique physical and chemical properties of an element.
Atoms are composed of three primary subatomic particles:
- Protons: Positively charged particles ( $+$ ).
- Electrons: Negatively charged particles ( $-$ ).
- Neutrons: Electrically neutral particles (no charge).
Spatial organization of subatomic particles within an atom:
- Protons and neutrons are tightly packed together inside the central nucleus.
- Electrons rapidly orbit the nucleus within an electron cloud.
- The electrostatic attraction between the negative charge of the orbiting electrons and the positive charge of the nuclear protons keeps the electrons near the nucleus.
Fundamental atomic properties:
- Atomic Number: The fixed number of protons present in an atom's nucleus, which defines the element.
- Mass Number: The total sum of the number of protons and neutrons in an atom's nucleus.
- Atomic Mass: The total mass of an atom, which is approximately equal to its mass number.
Representative atomic structures:
- Helium ( $He$ ): Atomic number = $2$ , Mass number = $4$ ( $2$ protons, $2$ neutrons, $2$ electrons).

*   **Carbon ( $C$ )**: Atomic number =  $6$ , Mass number =  $12$  ( $6$  protons,  $6$  neutrons,  $6$  electrons).

Isotopes and Biological Applications of Radioactivity

Although all atoms of a specific element share the same atomic number, they can differ in mass number.
Isotopes are atomic variants of the same element that have the identical number of protons but a different number of neutrons.
Isotopes of Carbon:
- Carbon-12 ( $^{12}C$ ): $6$ protons, $6$ neutrons, $6$ electrons (Mass number = $12$ ).
- Carbon-13 ( $^{13}C$ ): $6$ protons, $7$ neutrons, $6$ electrons (Mass number = $13$ ).
- Carbon-14 ( $^{14}C$ ): $6$ protons, $8$ neutrons, $6$ electrons (Mass number = $14$ ).

Radioactive Isotopes:
- Isotopes in which the nucleus is unstable and decays spontaneously, giving off particles and high-energy radiation.
- Applications in medicine and diagnostic imaging:
  - Positron Emission Tomography (PET): Tracks metabolic activity and pinpoints tissue abnormalities or tumors.
  - Early Detection of Alzheimer's Disease: Detects metabolic variations and amyloid accumulation in living brain tissues.
  - Treatment of Hyperthyroidism: Uses targeted radioactive iodine to destroy overactive thyroid tissue.
- Uncontrolled or excessive exposure to radioactive isotopes can cause severe damage to cellular molecules and DNA.

PET scan brain comparison for Alzheimer's disease

Elements are systematically classified in the Periodic Table of the Elements based on atomic number and recurring chemical characteristics.

Electron Configuration and Chemical Bonding

Of the three subatomic particles, only electrons are directly involved in the chemical activity and bonding of an atom.
Electrons reside in distinct energy levels called electron shells surrounding the nucleus:
- The innermost shell accommodates a maximum of $2$ electrons.
- Outer shells accommodate up to $8$ valence electrons in lighter elements.
- The valence shell is the outermost electron shell that governs an atom's reactivity.
- Electrons travel around the electron cloud in specific 3D regions called orbitals.

Electron Distribution Diagrams of First 18 Elements

Atoms with incomplete outer valence shells tend to react with other atoms so that both end up with complete valence shells.
Atoms fulfill their valence shells by sharing, donating, or receiving electrons.
These atomic interactions hold the participating atoms close together, forming stable chemical bonds.

Types of Chemical Bonds: Covalent, Ionic, and Hydrogen

There are three principal types of chemical bonds: covalent bonds, ionic bonds, and hydrogen bonds.
1. Covalent Bonds:
- The strongest category of chemical bond, formed when two atoms share one or more pairs of outer-shell valence electrons.
- Two or more atoms joined together by covalent bonds constitute a molecule.
- Types of covalent bonds based on electron pairs shared:
  - Single Covalent Bond: The sharing of $1$ pair of electrons ( $2$ electrons total). Example: Hydrogen gas ( $H_2$ ).
  - Double Covalent Bond: The sharing of $2$ pairs of electrons ( $4$ electrons total). Example: Oxygen gas ( $O_2$ ).
  - Triple Covalent Bond: The sharing of $3$ pairs of electrons ( $6$ electrons total). Example: Nitrogen gas ( $N_2$ ).
- Molecular example: Methane ( $CH_4$ ) features a central carbon atom sharing single covalent bonds with four surrounding hydrogen atoms.

Alternative Ways to Represent Four Common Molecules

*   **Electronegativity and Bond Polarity**:
    *   **Electronegativity** is an atom's intrinsic pull or attraction for shared electrons in a covalent bond.
    *   **Nonpolar Covalent Bond**: Formed between atoms with identical electronegativity (e.g., diatomic molecules  $H_2$  or  $O_2$ ), resulting in an equal sharing of electrons.
    *   **Polar Covalent Bond**: Formed between atoms with differing electronegativities.
        *   In a water molecule ( $H_2O$ ), oxygen is far more electronegative than hydrogen.
        *   The shared electrons spend more time closer to the oxygen nucleus.
        *   Consequently, the oxygen atom acquires a partial negative charge ( $\delta^-$ ), while the hydrogen atoms carry partial positive charges ( $\delta^+$ ).
        *   Because of these polar covalent bonds, water is classified as a **polar molecule**.

2. Ionic Bonds:
- An ion is an atom or molecule that carries an electrical charge due to the gain or loss of one or more electrons.
  - Cation: A positively charged ion formed when an atom loses electrons.
  - Anion: A negatively charged ion formed when an atom gains electrons.
- An ionic bond is the strong electrostatic attraction that holds two oppositely charged ions together.
- Salt is a synonym for an ionic compound.
- Example: A sodium atom ( $Na$ ) transfers an electron to a chlorine atom ( $Cl$ ), producing a sodium cation ( $Na^+$ ) and a chloride anion ( $Cl^-$ ). Their mutual attraction forms an ionic bond in sodium chloride ( $NaCl$ ).

3. Hydrogen Bonds:
- A weak chemical bond formed between a partially positive hydrogen atom of one polar covalent molecule and a partially negative atom (such as oxygen) of another polar molecule.
- In water, hydrogen bonding occurs between the slightly positive hydrogen atom of one water molecule and the slightly negative oxygen atom of an adjacent water molecule.

Hydrogen Bonding Between Water Molecules

Unique Properties and Functions of Water

Cohesion and Adhesion:
- Cohesion: The tendency of molecules of the same kind to stick to one another. Cohesion is extraordinarily strong in water due to its extensive hydrogen bonding network.
- Adhesion: The attraction between two different kinds of polar molecules.
- Biological importance: Plants rely on cohesion and adhesion to pull water and dissolved nutrients upward from roots to leaves through water-conducting xylem tissue.

Cohesion and Adhesion in Plant Water Transport

Temperature Moderation:
- Because of hydrogen bonding, water possesses a high heat capacity and resists rapid temperature fluctuations.
- Heat: The total thermal energy associated with the random movement of atoms and molecules in a body of matter.
- Temperature: A measure of the average kinetic energy of molecules.
- Thermal mechanics of hydrogen bonds:
  - Heat must be absorbed to break hydrogen bonds.
  - Heat is released when hydrogen bonds form.
Low Density of Ice:
- Ice is less dense than liquid water, allowing ice to float on top of liquid bodies of water.
- When water freezes, each molecule forms stable, spacious hydrogen bonds with four neighboring molecules, creating a crystalline lattice.
- In liquid water, hydrogen bonds constantly break and re-form, allowing molecules to pack closer together.

Universal Solvent of Life:
- Solution: A liquid mixture consisting of a uniform blend of two or more substances.
- Solvent: The dissolving agent of a solution.
- Solute: The substance that is dissolved.
- Aqueous Solution: A solution in which water serves as the solvent.
- Water's versatility as a solvent stems from its molecular polarity. Polar or charged solutes dissolve readily as water molecules cluster around them, forming hydration spheres.
- Example: Table salt ( $NaCl$ ) dissolving in water as polar water molecules separate individual $Na^+$ and $Cl^-$ ions.

Acids, Bases, pH Scale, and Buffers

In aqueous solutions, a small fraction of water molecules break apart into highly reactive ions:
- Hydrogen ions ( $H^+$ )
- Hydroxide ions ( $OH^-$ )
Acids and Bases:
- An acid is a compound that releases $H^+$ ions into a solution, increasing the $H^+$ concentration: $\text{HCl} \rightarrow \text{H}^+ + \text{Cl}^-$
- A base (alkaline substance) is a compound that accepts $H^+$ ions or releases $OH^-$ ions, reducing $H^+$ concentration: $\text{NaOH} \rightarrow \text{Na}^+ + \text{OH}^-$
- Mixing an acid and a base yields water and a salt: $\text{NaOH} + \text{HCl} \rightarrow \text{H}_2\text{O} + \text{NaCl}$
The pH Scale:
- The pH scale measures the relative acidity or basicity of a solution, ranging from $0$ to $14$ .
- Acidic Solutions: pH values from $0$ to less than $7$ (higher $H^+$ concentration relative to $OH^-$ ).
- Neutral Solutions: pH equal to $7.0$ ( $[H^+] = [OH^-]$ ).
- Basic (Alkaline) Solutions: pH values greater than $7$ up to $14$ (higher $OH^-$ concentration relative to $H^+$ ).
- Each integer change on the pH scale represents a tenfold (10-fold) change in $H^+$ concentration.
- Standard reference pH values:
  - Battery acid: $\text{pH } 0$
  - Lemon juice, gastric juice: $\text{pH } 2$
  - Vinegar, cola: $\text{pH } 3$
  - Tomato juice: $\text{pH } 4$
  - Rainwater: $\text{pH } 5.5$
  - Saliva, human urine: $\text{pH } 6$
  - Pure water: $\text{pH } 7.0$
  - Human blood, tears: $\text{pH } 7.4$
  - Seawater: $\text{pH } 8$
  - Milk of magnesia: $\text{pH } 10.5$
  - Household ammonia: $\text{pH } 11.5$
  - Household bleach: $\text{pH } 12.5$
  - Oven cleaner: $\text{pH } 13.5 - 14$

Buffers:
- A buffer is a substance that minimizes changes in pH in biological fluids.
- Buffers operate by accepting excess $H^+$ ions when conditions become acidic and donating $H^+$ ions when conditions become basic.