Biology 1 - Lecture 3: Atomic Structure and Chemical Bonds

Atomic Structure

• Composition of the Atom: An atom consists of a nucleus containing protons and neutrons, and electrons that orbit the nucleus.

• Charge Dynamics: Electrons, which carry a negative charge, are attracted to the nucleus, which carries a positive charge.

• Elemental Identity: The type of atom is determined by the number of protons in the nucleus. Each type of atom has a specific name and symbol, such as:

  • $H$ (Hydrogen)

  • $He$ (Helium)

• Elements: These are substances composed of only one type of atom.

• Atomic Masses:

  • The mass of protons and neutrons is approximately $1\,Da \approx 1.7 \times 10^{-27}\,kg$.

  • The mass of electrons is approximately $0.0005\,Da \approx 9 \times 10^{-31}\,kg$.

  • The total mass of the atom is the sum of the masses of the protons and neutrons; the mass of electrons is considered negligible.

• Atomic Characterization:

  • $Z$ (Atomic Number) = The number of protons.

  • $A$ (Mass Number) = The total number of protons plus neutrons.

Isotopes and Atomic Weight

• Definition: Atoms of the same element can differ from one another in the number of neutrons they contain.

• Example: Carbon isotopes ($C$).

• Stability: Not all isotopes are stable.

• Radioactive Decay: During $\alpha$ or $\beta$ radioactive decay of an atom, one element transforms into an atom of a different element, releasing energy and particles in the process.

• Chemical Behavior: All isotopes of an element react with other atoms in identical ways. In chemistry, only the electrons are considered important.

• Average Atomic Mass (Atomic Weight): This is the weighted average mass of all naturally occurring isotopes of an element.

The Periodic Table

• Organization:

  • Vertical columns (Groups) contain elements with similar chemical properties.

  • Atomic number represents the quantity of protons.

  • Elemental Symbol and Name identify the substance.

  • Molar Mass ($g/mol$) represents the mass of one mole of the substance.

• Elemental Importance in Living Organisms:

  • The Big Six: Six elements highlighted in yellow make up $98\%$ of the mass of most living organisms: Hydrogen ($H$), Carbon ($C$), Nitrogen ($N$), Oxygen ($O$), Phosphorus ($P$), and Sulfur ($S$).

  • Trace Elements: Elements framed in orange are present in small amounts in many organisms.

• Notation: Masses written in parentheses denote unstable elements that decay rapidly to form other elements.

Electron Arrangement in the Atom

• Governing Forces: The attraction of electrons to the nucleus is rooted in the electromagnetic force. Generally, work must be performed to move electrons away from the nucleus, which increases the energy of the system.

• Quantum Mechanics: The spatial movement of electrons is described by quantum mechanics rather than Newtonian mechanics.

• Orbitals: These represent the probability of finding an electron in a specific location.

• Shells and Subshells: Orbitals are organized into shells and subshells. Each shell has a different number of subshells, and each subshell has a specific number of orbitals.

• Energy Levels:

  • The order of orbital energy is not always simple.

  • Example: Carbon ($C$) structure is $1s^2 2s^2 2p^2$.

  • Stable configurations are achieved by placing electrons in order of increasing energy, with a maximum of two electrons per orbital.

• Valence Electrons: The arrangement of electrons in the outermost shell (valence shell) determines the chemical properties of an atom.

The Octet Rule

• Definition: Atoms tend to gain, lose, or share electrons to achieve a configuration of eight valence electrons in their outermost shell, similar to the configuration of noble gases.

• Examples:

  • Carbon ($C$) would like to gain $4$ electrons.

  • Nitrogen ($N$) would like to gain $3$ electrons.

  • Oxygen ($O$) would like to gain $2$ electrons.

Molecules and Chemical Bonds

• Molecules: Formed when two or more atoms create bonds between them. A molecule is the smallest part of a compound that retains its properties.

  • Molecules can be made of identical atoms (e.g., Oxygen $O_2$), dozens of different atoms (e.g., sugar), or millions of atoms (e.g., $DNA$).

• Empirical Formula: Describes the ratio of atoms composing the molecule (e.g., Water: $H_2O$).

• Molecular Mass: The sum of the atomic masses of all atoms composing the molecule.

• Types of Interactions:

  • Covalent Bond: Sharing of electron pairs.

  • Ionic Bond: Attraction of opposite charges after the transfer of electrons.

  • Hydrogen Bond: Sharing of a Hydrogen atom/attraction between partial charges.

  • Hydrophobic Interaction: Interaction of nonpolar substances in the presence of polar substances (especially water).

  • van der Waals Interaction: Interaction of electrons of nonpolar substances.

Covalent Bonds

• Sharing: Two atoms create a partnership of valence electrons (outer shell).

• Bond Orders:

  • Single Covalent Bond: Sharing one pair of electrons ($H-H$, $H_2O$).

  • Double Covalent Bond: Sharing two pairs of electrons ($O=C=O$, $C_2H_4$).

  • Triple Covalent Bond: Sharing three pairs of electrons ($N \equiv N$, $C_2H_2$).

• Depicting Molecules:

  • Lewis Structure: Shows dots representing valence electrons.

  • Structural Formula: Uses lines to represent bonds.

  • Line Diagram (Skeletal Formula): Simplified geometric representation.

  • Condensed Formula: Lists atoms and their counts (e.g., $CH_3CH_2CH_3$).

• 3D Representation:

  • Normal Bond: Bond lies in the plane of the paper.

  • Dashed Bond: Bond extends backwards, away from the viewer.

  • Wedged Bond: Bond protrudes forwards, towards the viewer.

Electronegativity and Polarity

• Electronegativity: The relative measure of an atom's ability to attract (or release) electrons in a covalent bond.

• Non-polar Covalent Bonds: Formed between atoms with close electronegativity values (e.g., Oxygen molecule $O_2$).

• Polar Covalent Bonds: Formed when there is an unequal distribution of electrons. In water ($H_2O$), oxygen is more electronegative than hydrogen, pulling electrons closer to the oxygen atom. This results in:

  • Partial negative charge ($\delta^-$) near oxygen.

  • Partial positive charge ($\delta^+$) near hydrogen.

Ionic Bonds

• Formation: Occurs when there is a large difference in electronegativity between atoms. Instead of sharing, an electron transfers from one atom to another.

• Ions: The atom that loses an electron becomes a positive ion (cation), and the one that gains an electron becomes a negative ion (anion).

• Stability: Electrostatic attraction between the ions stabilizes the molecule.

• Properties: Ionic bonds generally form solids known as salts.

• Solubility (The Universal Solvent): Ionic bonds break in polar solvents like water because water molecules overcome the attraction between the positive and negative ions, creating mobile ions in the solution.

Hydrogen Bonds and Intermolecular Forces

• Hydrogen Bonds: These result from the attraction between positive and negative partial charges across different molecules or different parts of the same molecule. They are much weaker than covalent bonds.

• Hydrophilic vs. Hydrophobic:

  • Hydrophilic: Polar substances that "prefer" the water environment (e.g., salts, sugars).

  • Hydrophobic: Non-polar substances (e.g., oil) that "interfere" with the hydrogen bonds between water molecules, causing non-polar molecules to be attracted to each other in the presence of water.

• van der Waals Interactions: Driven by induced electrical interactions when outer electron clouds of two atoms are very close. They are the weakest of all intermolecular interactions, but a high density of these forces can create very strong overall interactions (e.g., geckos and spiders).

Bond Energy Comparison

• Covalent Bond: $50-110\,kcal/mol$

• Ionic Bond: $3-7\,kcal/mol$

• Hydrogen Bond: $3-7\,kcal/mol$

• Hydrophobic Interaction: $1-2\,kcal/mol$

• van der Waals Interaction: $1\,kcal/mol$

Acids and Bases

• Water Dissociation: A very small portion of water molecules undergo spontaneous dissociation into ions, as the ions are less stable than the full molecule.

  • In pure water: $[OH^-] = [H^+] = 10^{-7}\,M$.

  • Concentration of $H_2O = 55\,M$.

• Definitions:

  • Molar Concentration ($M$): $moles/liter$.

  • Mole: $\approx 6 \times 10^{23}\text{ particles}$.

  • pH (Power of Hydrogen): A quantitative measure of the acidity level of a solution. Calculated as $pH = -\log_{10}([H^+])$.

  • pH of Water: $pH_{water} = -\log_{10}(10^{-7}) = 7$.

• Hydronium Ion: Free protons ($H^+$) do not exist in solution; they bond immediately to water molecules to form hydronium ions ($H_3O^+$).

• Acids: Substances that donate protons ($H^+$) when dissolved in water.

  • Example: Carboxyl group dissociation ($—COOH \rightarrow —COO^- + H^+$).

  • Acidic solutions: [H^+] > [OH^-].

• Bases: Substances that accept protons ($H^+$) when dissolved in water. This leads to the production of hydroxide ions ($OH^-$).

  • Example: Sodium hydroxide ($NaOH \rightarrow Na^+ + OH^-$) or Amino group protonation ($—NH_2 + H^+ \rightarrow —NH_3^+$).

  • Basic solutions: [H^+] < [OH^-].

pH Scale and Common Examples

Below are the pH values for various substances:

• Acidic Mine Water: $-3.6$ to $1.0$

• Car Battery Acid: $0.5$

• Stomach Acid: $2.0$

• Lemon Juice: $2.4$

• Cola: $2.5$

• Vinegar: $2.9$

• Orange Juice: $3.5$

• Beer: $4.5$

• Acid Rain: $5.0$

• Coffee: $5.0$

• Tea: $5.5$

• Milk: $6.5$

• Saliva: $6.5$ to $7.4$

• Distilled Water: $7.0$

• Blood: $7.34$ to $7.45$

• Sea Water: $8.0$

• Hand Soap: $9.0$ to $10.0$

• Ammonia: $11.5$

• Bleach: $12.5$

• Degreasers: $13.5$