Introduction to Chemistry in Biology: Atoms, Elements, and Bonding

Hierarchy of Sciences and Life's Foundations

  • The Purity Hierarchy of Sciences:

    • Mathematics is considered the most "pure" science.

    • Physics is applied Mathematics.

    • Chemistry is applied Physics.

    • Biology is applied Chemistry.

    • Psychology is applied Biology.

    • Sociology is applied Psychology.

    • This hierarchy implies that understanding foundational sciences like Chemistry and Physics is crucial for comprehending Biology.

  • Learning Objectives for the Session:

    • Describe the interrelationship between protons, neutrons, and electrons.

    • Understand what isotopes are and their importance.

    • Understand how to "read" the periodic table.

    • Apply the octet rule to predict the number of bonds an atom can form.

  • Biology's Foundation:

    • Biology is fundamentally built upon the principles of Chemistry and Physics.

    • The study of atoms and molecules in living organisms is known as Biochemistry.

    • Examples of critical elements in biology include Hydrogen (H), Carbon (C), Nitrogen (N), Oxygen (O), Phosphorus (P), Sulfur (S), and Calcium (Ca).

    • An atom's basic structure involves a nucleus, electrons, and electron shells.

Chemical Elements Essential for Life

  • Primary Composition of Living Organisms:

    • Living organisms are primarily composed (99\%) of only 7 elements: Carbon (C), Hydrogen (H), Nitrogen (N), Oxygen (O), Phosphorus (P), Sulfur (S), and Calcium (Ca).

  • Most Abundant Elements (approximately 95\% of total mass in living organisms):

    • Oxygen (O): 65\% of human body mass, 25.5\% of all atoms.

    • Carbon (C): 18\% of human body mass, 9.5\% of all atoms.

    • Hydrogen (H): 9\% of human body mass, 63.0\% of all atoms.

    • Nitrogen (N): 3\% of human body mass, 1.4\% of all atoms.

  • Mineral Elements (less than 1\% of total mass):

    • Calcium (Ca)

    • Potassium (K)

    • Chlorine (Cl)

    • Sodium (Na)

    • Magnesium (Mg)

    • Phosphorus (P)

    • Sulfur (S)

  • Trace Elements (less than 0.01\% of total mass):

    • Boron (B)

    • Manganese (Mn)

    • Chromium (Cr)

    • Molybdenum (Mo)

    • Cobalt (Co)

    • Selenium (Se)

    • Copper (Cu)

    • Silicon (Si)

    • Fluorine (F)

    • Tin (Sn)

    • Iodine (I)

    • Vanadium (V)

    • Iron (Fe)

    • Zinc (Zn)

    • Note: Many other trace and mineral elements also have reported functions. For example, Aluminum (Al) is a cofactor for certain chemical reactions in animals but is generally toxic to plants.

Atoms and Atomic Structure

  • Atoms:

    • The smallest functional units of matter that form all chemical substances.

    • Cannot be further broken down by ordinary physical or chemical means.

  • Subatomic Particles: Atoms are composed of 3 types of subatomic particles:

    • Protons (p^+):

      • Positive charge.

      • Located in the nucleus.

    • Neutrons (n^0):

      • Neutral charge.

      • Located in the nucleus.

    • Electrons (e^-):

      • Negatively charged.

      • Orbit the nucleus in electron shells.

      • Crucially determine an atom's chemical properties.

Atomic Number and Mass Number

  • Atomic Number:

    • Defined as the number of protons (p^+) in the nucleus.

    • It is a characteristic identifier for each element (e.g., Hydrogen (H) = 1 proton, Carbon (C) = 6 protons, Oxygen (O) = 8 protons).

    • In an electrically neutral atom, the number of protons (p^+) equals the number of electrons (e^-).

  • Mass Number:

    • Defined as the total number of protons (p^+) plus the number of neutrons (n^0) in the nucleus.

    • Examples: Hydrogen (H) = 1 (often 1 proton, 0 neutrons), Carbon (C) = 12 (often 6 protons + 6 neutrons), Oxygen (O) = 16 (often 8 protons + 8 neutrons).

  • Standard Atomic Notation: An element's symbol is often represented with its mass number as a superscript and its atomic number as a subscript (e.g., $^{12}_{6}C$ for Carbon).

The Periodic Table

  • Organization of Elements: The periodic table systematically arranges all known chemical elements.

  • Information Provided for Each Element:

    • Atomic Number

    • Chemical Symbol

    • Element Name

    • Atomic Mass

  • Reading the Rows (Periods):

    • Elements in the same horizontal row (period) have the same number and types of electron orbitals.

    • Rows also provide an indication of an element's abundance in cells (e.g., High, Low, Trace, Rare/None).

  • Reading the Columns (Groups):

    • Elements in the same vertical column (group) have the same number of electrons in their outermost shell (valence electrons).

Isotopes and Radioisotopes

  • Isotopes:

    • Atoms of the same element that have the same number of protons (p^+), but a different number of neutrons (n^0).

    • This means they have the same atomic number but different mass numbers (e.g., Carbon-12 and Carbon-14).

  • Radioisotopes:

    • Some isotopes are unstable and spontaneously decay, giving off energy in the form of rays (like gamma rays) or subatomic particles.

    • These are widely used in research, medicine, and studies of the history of life.

  • Applications of Radioisotopes:

    • Tracers: Used to reveal the pathway or destination of a substance that has entered a cell, the human body, or an ecosystem (e.g., in PET scans a tracer is used).

    • Radiometric Dating: Used to determine the age of fossils, rocks, and archaeological artifacts (e.g., Carbon-14 dating).

    • Medical Applications:

      • Cobalt-60: Used to irradiate food, making it safer to eat and extending its shelf-life.

      • Iodine-131: Used to study thyroid hormones and visualize the thyroid gland due to its uptake by thyroid cells.

Radiometric Dating and Half-life

  • Half-life:

    • The length of time required for half ( rac{1}{2}) the atoms of a radioactive isotope to change into a more stable product.

    • It is a constant rate, unaffected by external factors like temperature, light, or pressure.

    • All radioactive isotopes have a dependable and characteristic half-life.

  • Example: Carbon-14 Dating:

    • Radioisotope: $^{14}C$

    • Product: $^{14}N$

    • Half-life: 5,730 years

    • Useful dating range: < 50,000 years

Electron Orbitals and Shells

  • Orbitals: Electrons occupy specific regions of space around the nucleus called orbitals.

  • Electron Shells (Bohr Model):

    • In the simplified Bohr model, electron shells are depicted as concentric circles around the nucleus, representing different energy levels.

    • Atoms with more than two electrons have at least two orbitals.

    • Example (Carbon Atom): The nucleus contains 6 protons and 6 neutrons. Electrons occupy shells around it.

    • Energy Levels: Electrons in shells represent different energy levels (e.g., Hydrogen has one electron in its first shell, Carbon has two in its first and four in its second).

The Octet Rule and Chemical Bonding

  • Electron Distribution and Shell Capacity:

    • Each shell contains one or more electron orbitals.

    • The innermost electron shell has 1 orbital and can hold a maximum of 2 electrons.

    • The next two shells can each hold a maximum of 8 electrons.

  • Valence Electrons and Reactivity:

    • Atoms with fewer than 8 electrons in their outermost shell are considered reactive.

    • Electrons located in the outermost shell are called valence electrons; they are available to combine with other atoms.

    • Atoms with oldsymbol{\le 3} valence electrons tend to donate electrons.

    • Atoms with oldsymbol{\ge 5} valence electrons tend to receive electrons.

  • The Octet Rule:

    • States that atoms tend to combine in such a way that they each have eight electrons in their outermost (valence) shell.

    • A full outer shell (typically with 8 electrons, or 2 for the first shell) makes an atom most stable.

In-Class Activity: Bond Calculation (Deductive Reasoning for Reactivity)

  • This exercise helps determine the reactivity and bonding capacity of major elements in biological molecules by examining their atomic numbers.

  • Steps for Carbon (C), Hydrogen (H), Nitrogen (N), and Oxygen (O):

    1. Calculate the number of electrons in the outermost energy shell (valence electrons).

      • Hydrogen (H): Atomic Number 1. 1 electron in the 1^{st} (outermost) shell.

      • Carbon (C): Atomic Number 6. 2 electrons in 1^{st} shell, 4 electrons in 2^{nd} (outermost) shell.

      • Nitrogen (N): Atomic Number 7. 2 electrons in 1^{st} shell, 5 electrons in 2^{nd} (outermost) shell.

      • Oxygen (O): Atomic Number 8. 2 electrons in 1^{st} shell, 6 electrons in 2^{nd} (outermost) shell.

    2. Determine how many electrons each atom needs to acquire via covalent bonding to fill its outer shell (satisfy the octet rule).

      • Hydrogen (H): Needs 1 more electron (to reach 2 in the first shell).

      • Carbon (C): Needs 4 more electrons (to reach 8 in the second shell).

      • Nitrogen (N): Needs 3 more electrons (to reach 8 in the second shell).

      • Oxygen (O): Needs 2 more electrons (to reach 8 in the second shell).

    3. Identify the maximum number of bonds each atom is capable of participating in.

      • Hydrogen (H): 1 bond.

      • Carbon (C): 4 bonds.

      • Nitrogen (N): 3 bonds.

      • Oxygen (O): 2 bonds.

Practice Question Solution

  • Question: You discover an isotope of an element that has 6 electrons in its second and outermost shell, 8 protons, and 6 neutrons. What element is it?

  • Solution Breakdown:

    • The defining characteristic of an element is its number of protons.

    • The element has 8 protons.

    • Looking at the periodic table, the element with an atomic number of 8 is Oxygen (O).

    • The fact that it has 6 electrons in its outermost shell is consistent with Oxygen's electron configuration (2 electrons in the first shell, 6 in the second).

    • The number of neutrons (6) indicates a specific isotope of Oxygen (Oxygen-14, as 8 protons + 6 neutrons = 14 mass number).

  • Correct Answer: C. Oxygen (O)