chemistry of life chapter

2.1 Atoms, Ions, and Molecules — Study Notes

1. Matter & Atoms

• Matter = anything that has mass and occupies space.

• All matter is made of atoms, the smallest units that retain the properties of an element.

• Atoms consist of:

  • Protons (+ charge) — determine the element’s identity.

  • Neutrons (neutral) — add mass; influence stability.

  • Electrons (– charge) — determine chemical behavior.

2. Atomic Number & Mass Number

• Atomic number = number of protons.

• Mass number = protons + neutrons.

• Changing neutrons → isotopes (same element, different mass).

3. Ions

• Atoms become ions when they gain or lose electrons.

• Cations = positively charged (lost electrons).

• Anions = negatively charged (gained electrons).

• Ionic behavior is driven by achieving a stable electron configuration (like noble gases).

4. Elements & the Periodic Table

• Elements are organized by atomic number and electron configuration.

• Columns (groups) share similar chemical properties.

• Noble gases are stable → other atoms react to mimic their electron arrangement.

5. Chemical Bonds

Ionic Bonds

• Form between cations + anions.

• Result from electron transfer.

• Create crystal lattices (strong but brittle).

Covalent Bonds

• Form when atoms share electrons.

• Can be:

  • Nonpolar — equal sharing.

  • Polar — unequal sharing → partial charges.

6. Molecules & Compounds

• Molecule = two or more atoms bonded.

• Compound = molecule containing different elements.

• Molecular shape + polarity determine biological function.

7. Water’s Importance

• Water is polar → forms hydrogen bonds.

• Hydrogen bonding gives water:

  • High cohesion & adhesion

  • High specific heat

  • Excellent solvent properties

• Critical for biochemical reactions and temperature regulation.

8. Chemical Reactions

• Involve breaking and forming bonds.

• Reactants → Products.

• Governed by:

  • Concentration

  • Temperature

  • Catalysts (enzymes in biology)

9. pH & Acids/Bases

• pH measures hydrogen ion concentration.

• Acids donate H⁺.

• Bases accept H⁺.

• Biological systems rely on buffers to maintain pH stability.

2.2 Water and Mixtures — Study Notes

1. Water as the Universal Solvent

• Water is the most important inorganic molecule in the body.

• Its polarity allows it to dissolve many substances.

• Polar water molecules surround solutes → forming hydration shells.

• This makes water essential for:

  • Transport

  • Chemical reactions

  • Temperature regulation

  • Lubrication

2. Properties of Water

High Specific Heat

• Water absorbs a lot of heat before its temperature rises.

• Helps maintain stable internal body temperature.

High Heat of Vaporization

• Requires significant energy to turn water into vapor.

• Sweating uses this property to cool the body.

Cohesion & Adhesion

• Cohesion: water molecules stick to each other.

• Adhesion: water sticks to other surfaces.

• Together, they support fluid movement (e.g., blood flow, capillary action).

Surface Tension

• Water molecules at the surface are strongly attracted to each other.

• Important for lung function (alveoli).

3. Mixtures

Mixtures are combinations of substances that are not chemically bonded.

Types of Mixtures

1. Solution

• Small particles dissolved in water.

• Transparent, does not scatter light.

• Particles stay evenly distributed.

• Example: plasma, saline.

2. Colloid

• Medium-sized particles.

• Cloudy, scatters light.

• Particles remain suspended.

• Example: cytosol, proteins in blood.

3. Suspension

• Large particles.

• Cloudy, particles settle if left still.

• Example: blood cells in plasma.

4. Water in Chemical Reactions

Hydrolysis

• Water breaks bonds.

• Used in digestion and breakdown of macromolecules.

Dehydration Synthesis

• Water is removed to form new bonds.

• Used to build proteins, carbohydrates, lipids.

5. Acid–Base Chemistry in Water

• Water can dissociate into H⁺ and OH⁻.

• pH measures the concentration of H⁺.

• Acids increase H⁺.

• Bases decrease H⁺.

• Biological systems rely on buffers to stabilize pH.

6. Water’s Role in the Body

• Transport: blood, lymph, urine.

• Lubrication: serous fluid, synovial fluid.

• Cushioning: cerebrospinal fluid.

• Temperature regulation: sweating, blood flow.

• Chemical reactions: hydrolysis & synthesis.

2.3 Energy and Chemical Reactions — Study Notes

1. What Is Energy?

Energy = the capacity to do work.

Forms of Energy in the Body

• Kinetic energy

  • Energy of motion

  • Examples: muscle contraction, blood flow, ion movement

• Potential energy

  • Stored energy

  • Examples: chemical bonds, concentration gradients

Chemical Energy

• The most important form of potential energy in the body

• Stored in chemical bonds of molecules (glucose, ATP, fatty acids)

2. The First Law of Thermodynamics

• Energy cannot be created or destroyed

• It can only be converted from one form to another

• Example: chemical energy → kinetic energy during muscle contraction

3. The Second Law of Thermodynamics

• Every energy conversion releases heat

• Heat is unusable energy

• This is why the body must constantly produce ATP — energy is always “lost” as heat

4. Chemical Reactions

Chemical reactions occur when chemical bonds are broken, formed, or rearranged.

Types of Chemical Reactions

1. Decomposition (Catabolic)

• Large molecules → smaller molecules

• Releases energy

• Example: digestion, glycogen breakdown

2. Synthesis (Anabolic)

• Smaller molecules → larger molecules

• Requires energy

• Example: protein synthesis, building glycogen

3. Exchange Reactions

• Bonds are broken and formed

• Atoms are swapped between molecules

• Example: ATP production

5. Oxidation–Reduction (Redox) Reactions

• Oxidation = loss of electrons

• Reduction = gain of electrons

• Always occur together

• Critical for:

  • Cellular respiration

  • ATP production

  • Detoxification pathways

6. Reaction Rates

Reaction speed depends on:

Temperature

• Higher temperature → faster reactions

• Lower temperature → slower reactions

Concentration

• More reactants → faster reactions

• Fewer reactants → slower reactions

Presence of Catalysts

• Catalysts speed up reactions without being consumed

• In the body, catalysts = enzymes

7. Enzymes

Enzymes are biological catalysts that:

• Lower activation energy

• Increase reaction speed

• Are specific to substrates

• Are not used up in reactions

Enzyme Function Steps

1. Substrate binds to enzyme’s active site

2. Enzyme-substrate complex forms

3. Reaction occurs

4. Products are released

5. Enzyme is ready to repeat

Factors Affecting Enzyme Activity

• Temperature (too high → denaturation)

• pH (each enzyme has an optimal pH)

• Substrate concentration

8. ATP: The Energy Currency

ATP = adenosine triphosphate

• Stores energy in high‑energy phosphate bonds

• When ATP → ADP + Pi, energy is released

• Used for:

  • Muscle contraction

  • Active transport

  • Synthesis reactions

  • Cellular signaling

9. Activation Energy

• The minimum energy required to start a reaction

• Enzymes lower activation energy → reactions occur faster and at body temperature

2.4 Organic Compounds — Study Notes

1. What Makes a Compound Organic

• Organic compounds contain carbon and are produced by living organisms.

• Carbon can form four covalent bonds, allowing:

  • Long chains

  • Rings

  • Branches

• This flexibility makes carbon the backbone of biological molecules.

2. Functional Groups

Functional groups determine how organic molecules behave.

• Hydroxyl (–OH) → increases solubility; found in sugars & alcohols

• Carboxyl (–COOH) → acidic; donates H⁺

• Amino (–NH₂) → basic; accepts H⁺

• Phosphate (–PO₄³⁻) → energy transfer (ATP), signaling

• Methyl (–CH₃) → nonpolar; gene regulation

3. Four Major Classes of Organic Molecules

1. Carbohydrates

• Made of C, H, O in a 1:2:1 ratio.

• Primary function: energy.

• Types:

  • Monosaccharides → glucose, fructose

  • Disaccharides → sucrose, lactose

  • Polysaccharides → glycogen (human storage), starch (plants)

2. Lipids

• Hydrophobic molecules; not polymers.

• Functions:

  • Long‑term energy storage

  • Insulation

  • Cell membrane structure

• Types:

  • Triglycerides → fats/oils

  • Phospholipids → membrane bilayer

  • Steroids → cholesterol, hormones

3. Proteins

• Built from amino acids linked by peptide bonds.

• Functions:

  • Structure (keratin, collagen)

  • Movement (muscle)

  • Transport (hemoglobin)

  • Enzymes

  • Immunity (antibodies)

• Structure levels:

  • Primary → amino acid sequence

  • Secondary → alpha helices, beta sheets

  • Tertiary → 3D folding

  • Quaternary → multiple chains together

4. Nucleic Acids

• DNA & RNA.

• Store and transmit genetic information.

• Made of nucleotides:

  • Sugar

  • Phosphate

  • Nitrogen base (A, T, C, G, U)

4. Polymers & Monomers

Many organic molecules are polymers built from repeating monomers.

Examples:

• Proteins → amino acids

• Carbohydrates → monosaccharides

• Nucleic acids → nucleotides

Building Polymers

• Dehydration synthesis → removes water to form bonds

Breaking Polymers

• Hydrolysis → adds water to break bonds.

5. Why Carbon Is Essential

Carbon’s bonding ability allows:

• Complex shapes

• Stability

• Diversity of biological molecules

• Ability to form long chains and rings

This is why carbon is the foundation of life’s chemistry.

6. Biological Importance

Organic compounds are essential for:

• Energy production

• Cell structure

• Genetic inheritance

• Hormone signaling

• Enzyme activity

• Membrane formation

• Cell structure

• Genetic inheritance

• Hormone signaling

• Enzyme activity

• Membrane formation