Chapter 2: The Chemical Context of Life – General Biology Study Notes

Chemistry: Matter, Elements, & Compounds

Introduction to Matter

All living organisms are composed of matter, which is the foundation of biological systems. Understanding the chemical nature of matter is essential for studying life processes.

• Matter: Anything that takes up space and has mass.

• Elements: Pure substances that cannot be broken down into other substances by chemical reactions.

• All matter is made up of elements.

Example: Water (H2O) is made of the elements hydrogen and oxygen.

Elements in the Human Body

Major and Trace Elements

Elements are classified based on their abundance and role in the human body. Four elements make up the majority of living matter, while others are present in smaller amounts.

Element	Symbol	Percentage of Body Mass (including water)
Oxygen	O	65.0%
Carbon	C	18.5%
Hydrogen	H	9.5%
Nitrogen	N	3.3%
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	Various	<0.01%

• Trace elements are required in minute quantities but are essential for life (e.g., iron, zinc, copper).

• Some elements can be toxic, but certain species may adapt to environments containing toxic elements.

Compounds

Definition and Properties

A compound is a substance consisting of two or more elements combined in a fixed ratio. Compounds have unique properties that are different from those of their constituent elements.

• Example: Sodium (Na, a metal) and chlorine (Cl, a gas) combine to form sodium chloride (NaCl, table salt), which has properties distinct from either element.

Chemistry: Atoms

Structure of Atoms

Each element consists of unique atoms, which are the smallest units of matter that retain the properties of the element.

• Atoms are composed of subatomic particles: protons, neutrons, and electrons.

• Protons (positive charge) and neutrons (neutral) are located in the nucleus.

• Electrons (negative charge) form a cloud around the nucleus.

Atomic Number and Atomic Mass

Definitions and Calculations

The atomic number and atomic mass are fundamental properties of atoms that determine their identity and behavior.

• Atomic number: Number of protons in the nucleus. Also equals the number of electrons in a neutral atom.

• Mass number: Sum of protons and neutrons in the nucleus.

• Atomic mass: Total mass of the atom, approximated by the mass number.

• Proton and neutron mass ≈ 1 dalton (atomic mass unit, amu).

• Electrons are much lighter and are ignored in atomic mass calculations.

Formulas:

• Number of neutrons = mass number - atomic number

$\text{Number of neutrons} = \text{Mass number} - \text{Atomic number}$

Isotopes

Definition and Applications

Atoms of the same element may have different numbers of neutrons, resulting in isotopes. Isotopes have similar chemical properties but different physical properties.

• Stable isotopes do not change over time.

• Radioactive isotopes decay spontaneously, emitting particles and energy.

• Atomic mass of an element is the weighted average of all naturally occurring isotopes.

Example (Carbon):

Isotope	Protons	Neutrons	Stability	Abundance
Carbon-12	6	6	Stable	98.93%
Carbon-13	6	7	Stable	1.07%
Carbon-14	6	8	Radioactive	Trace

$\text{Average atomic mass of carbon} = (98.93\% \times 12\,\text{amu}) + (1.07\% \times 13\,\text{amu}) = 12.0107\,\text{amu}$

• Radioactive tracers are used in medicine and research to track atoms through metabolism and imaging.

Radiometric Dating

Principles and Half-Life

Radiometric dating uses the decay of radioactive isotopes to estimate the age of fossils and rocks. The rate of decay is measured by the half-life, the time required for half of the isotope to decay.

• Parent isotope decays into daughter isotope at a fixed rate.

• Half-life values vary from seconds to billions of years.

• Example: The half-life of Carbon-14 is 5,700 years.

Formula:

$\text{Amount remaining} = \text{Initial amount} \times \left(\frac{1}{2}\right)^{n}$ where $n$ = number of half-lives elapsed

• Scientists measure isotope ratios to calculate elapsed time since formation.

Energy Levels of Electrons

Potential and Kinetic Energy

Electrons possess potential energy due to their position relative to the nucleus. Changes in electron energy levels are fundamental to chemical reactions.

• Kinetic energy: Energy of motion.

• Potential energy: Stored energy due to position or structure.

• Electrons occupy electron shells with characteristic energy levels.

• Electrons can move between shells by absorbing or releasing energy in fixed amounts.

Electron Distribution and Chemical Properties

Valence Electrons and Reactivity

The chemical behavior of an atom is determined by the distribution of electrons, especially those in the outermost shell (valence shell).

• Valence electrons: Electrons in the outermost shell.

• Elements with a full valence shell are chemically inert (unreactive).

• The periodic table shows electron distribution and guides predictions of chemical behavior.

Electron Orbitals

Orbital Shapes and Hybridization

An orbital is a three-dimensional space where an electron is found 90% of the time. The arrangement and hybridization of orbitals determine molecular shapes.

• First shell: 1s orbital (spherical)

• Second shell: 2s (spherical) and three 2p orbitals (dumbbell-shaped)

• Hybridization of orbitals leads to specific molecular geometries (e.g., tetrahedral in methane).

Example: Water (H2O) and methane (CH4) have distinct shapes due to orbital hybridization.

Chemical Bonding Between Atoms

Types of Chemical Bonds

Atoms with incomplete valence shells can share or transfer electrons, forming chemical bonds that hold atoms together in molecules and compounds.

• Covalent bonds: Sharing of electron pairs between atoms.

• Ionic bonds: Transfer of electrons, resulting in attraction between oppositely charged ions.

• Hydrogen bonds: Attraction between a hydrogen atom covalently bonded to an electronegative atom and another electronegative atom.

• Van der Waals interactions: Weak attractions due to transient local charges.

Covalent Bonds

• Single covalent bond: Sharing one pair of electrons.

• Double covalent bond: Sharing two pairs of electrons.

• Valence: Bonding capacity, equal to the number of electrons needed to fill the valence shell.

• Structural formula: Representation of atoms and bonds (e.g., H—H for single bond, O=O for double bond).

Electronegativity and Bond Polarity

• Electronegativity: Atom’s attraction for electrons in a covalent bond.

• Nonpolar covalent bond: Electrons shared equally.

• Polar covalent bond: Electrons shared unequally, creating partial charges.

Example: Water (H2O) is a polar molecule due to unequal sharing of electrons between oxygen and hydrogen.

Ionic Bonds

• Formed when one atom transfers an electron to another, creating ions.

• Cation: Positively charged ion.

• Anion: Negatively charged ion.

• Ionic compounds (salts) are stable when dry but dissociate in water.

Example: Sodium chloride (NaCl) forms from Na+ and Cl- ions.

Hydrogen Bonds

• Form between hydrogen covalently bonded to an electronegative atom (e.g., oxygen, nitrogen) and another electronegative atom.

• Important in water, DNA, and protein structure.

Van der Waals Interactions

• Result from transient uneven electron distribution, creating temporary charges.

• Collectively can be strong, as in gecko toe adhesion.

Chemical Reactions

Making and Breaking Bonds

Chemical reactions involve the making and breaking of chemical bonds, transforming reactants into products.

• Reactants: Starting materials in a reaction.

• Products: Substances formed from a reaction.

• Photosynthesis: Example of a chemical reaction powered by sunlight.

Equation:

$6CO_2 + 6H_2O \xrightarrow{\text{sunlight}} C_6H_{12}O_6 + 6O_2$

• Chemical reactions are reversible; products of the forward reaction can become reactants in the reverse reaction.

• Chemical equilibrium: Forward and reverse reactions occur at the same rate; concentrations of reactants and products remain constant.

Equation:

$3H_2 + N_2 \rightleftharpoons 2NH_3$

Additional info: Some explanations and examples have been expanded for clarity and completeness, including definitions, formulas, and context for biological relevance.