Chapter 2 Notes: Matter, Atoms, and Chemical Reactions

Matter and Energy

Why study general chemistry: its role as the building block for understanding how cells and physiology work. Matter is the basic substance studied in chemistry.
Matter is defined as anything that occupies space and has mass.
Energy is the capacity to do work. Two basic types:
- Kinetic energy
- Potential energy
Three states of matter:
- Solid: holds its shape; fixed volume
- Liquid: shape conforms to the container; fixed volume
- Gas: shape and volume follow the container; no fixed volume or shape

The Building Blocks: Elements and Atoms

Elements: cannot be broken down into simpler substances by chemical means.
Each element has:
- Atomic number (number of protons, p+)
- Atomic symbol
- Mass number (approximately protons + neutrons, p+ + n0)
Atoms are the smallest units of matter that retain element properties and are composed of:
- Protons (positive charge) in the nucleus
- Neutrons (no charge) in the nucleus
- Electrons (negative charge) in the electron cloud surrounding the nucleus
Subatomic structure relationships:
- Nucleus contains protons and neutrons
- Electron cloud surrounds the nucleus
- Overall charge is neutral when the number of protons equals the number of electrons

Common Human Body Elements (Table 2.1 overview)

Major elements (approx. % body mass):
- Oxygen (O): ~65.0%
- Carbon (C): ~18.5%
- Hydrogen (H): ~9.5%
- Nitrogen (N): ~3.2%
- 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%
- Iodine (I): ~0.1%
- Iron (Fe): ~0.1%
Functions (selected examples):
- Oxygen: essential in organic molecules; as an ion (O2−) inherited in metabolism; critical for ATP production in cellular respiration
- Carbon: backbone of all organic molecules (carbohydrates, lipids, proteins, nucleic acids)
- Hydrogen: contributes to organic molecules and water; involved in energy production
- Nitrogen: component of proteins and nucleic acids; essential for genetic material
- Calcium: bone/teeth mineral; Ca2+ involved in muscle contraction, nerve impulse conduction, blood clotting
- Phosphorus: part of calcium phosphate salts in bones/teeth; part of nucleic acids; component of ATP
- Potassium: major intracellular cation; necessary for nerve impulses and muscle contraction
- Sulfur: component of certain proteins
- Sodium: major extracellular cation; water balance; nerve impulses and muscle contraction
- Chlorine: major extracellular anion; helps with water balance
- Magnesium: cofactor in many metabolic reactions
- Iodine: needed for thyroid hormone synthesis
- Iron: component of hemoglobin and some enzymes
Trace elements (less than ~0.01–0.1%):
- Chromium (Cr), Cobalt (Co), Copper (Cu), Fluorine (F), Manganese (Mn), Molybdenum (Mo), Selenium (Se), Silicon (Si), Tin (Sn), Vanadium (V), Zinc (Zn)
- Role: required in minute amounts; often part of enzymes or enzyme activation

Atomic Structure and Electron Arrangement

Atoms are the basic unit that retains element properties.
Subatomic particles:
- Protons (p+): positive charge; located in the nucleus
- Neutrons (n0): no charge; located in the nucleus
- Electrons (e−): negative charge; occupy the electron cloud around the nucleus
Electron arrangement determines chemical reactivity via valence electrons.
Valence shell concept:
- Inert elements have complete valence shells (e.g., Helium with 2e in its only shell; Neon with 8e in its outer shell)
- Reactive elements have incomplete valence shells (e.g., Hydrogen with 1e in its outer shell; Carbon with 4e in its outer shell; Oxygen with 6e in its outer shell; Sodium with 1e in its outer shell)

Bonds and Molecules

Molecules and compounds are formed by chemical bonds and can be made of two or more atoms and/or two or more elements in fixed ratios.
Types of bonds:
- Hydrogen bonds: very weak; easily disrupted by temperature or pH; example is the attraction between a hydrogen atom attached to an electronegative atom (e.g., O, N) and another electronegative atom (e.g., O in H2O)
- Ionic bonds: involve electron transfer; ions attract to form compounds; relatively weak in water
- Examples: Na+ Cl− (table salt), Ca2+ with F− (as CaF2)
- Covalent bonds: strongest; electrons are shared between atoms
- Polar covalent: unequal sharing; partial charges; e.g., H2O
- Nonpolar covalent: equal sharing; e.g., O=O in O2 or C=C in ethene
Covalent bond types by electron sharing:
- Single covalent bonds: share one pair of electrons; e.g., H−H in H2, C−H in methane CH4 (four single bonds)
- Double covalent bonds: share two pairs of electrons; e.g., O=O in O2
- Triple covalent bonds: share three pairs of electrons; e.g., N≡N in N2
Structural formulas use dots/lines to show bonds and lone pairs (e.g., CH4 shows four single covalent bonds from carbon to four hydrogens).

Chemical Reactions and Equations

Chemical reactions involve making or breaking chemical bonds and can rearrange matter:
- Reactants → Products
- Example: Hydrogen gas (H2) and Oxygen gas (O2) react to form water (H2O) (illustrative; actual balanced equation is 2 H2 + O2 → 2 H2O)
Types of reactions:
- Synthesis (Combination): smaller particles bond to form larger, more complex molecules
- Example: Amino acids joined to form a protein
- General form: A + B → AB
- Decomposition: bonds in larger molecules are broken into smaller molecules
- Example: Glycogen broken down to glucose units
- General form: AB → A + B
- Exchange (Displacement): bonds are both made and broken; components are rearranged
- Example: ATP transfers a phosphate group to glucose to form glucose-phosphate
- General form: AB + C → AC + B
Redox (Oxidation-Reduction) reactions:
- Oxidation: loss of electrons and hydrogen
- Reduction: gain of electrons and hydrogen
- Example: Glucose oxidation with oxygen to yield carbon dioxide and water
- Overall representation (simplified):
  ext{C}6 ext{H}{12} ext{O}6 + 6 ext{O}2
  ightarrow 6 ext{CO}2 + 6 ext{H}2 ext{O}

Rate of Chemical Reactions

Factors that influence reaction rate:
- Temperature: higher temperature generally increases rate
- Particle size: smaller particles increase surface area and rate
- Concentration: higher concentration can increase rate
- Catalysts: enzymes in biology; lower activation energy to speed up reactions

Solutions, Solvents, and Ions

Solvent: the dissolving agent (water is the universal solvent in many biological contexts)
Solute: the substance being dissolved
Water and hydration:
- Water molecules surround ions in solution via partial charges: negative regions of water surround cations (e.g., Na+), positive regions surround anions (e.g., Cl−)
Salt in water dissociates into ions: Na+ and Cl− in solution

Acids, Bases, pH, and Buffers

pH scale:
- Acidic: pH 0–6.99
- Neutral: pH 7.00
- Basic (Alkaline): pH 7.01–14
Examples of dissociation:
- HCl → H⁺ + Cl⁻
- NaOH → Na⁺ + OH⁻
Buffers:
- Buffers resist abrupt or large changes in pH
- Example: Carbonic acid/bicarbonate buffer system in blood
- Equilibrium example:
  ext{H}2 ext{CO}3
  ightleftharpoons ext{HCO}_3^- + ext{H}^+
- Forward reaction lowers pH; reverse reaction raises pH
Practical relevance: buffers help maintain homeostasis in bodily fluids, enabling stable enzyme function and metabolism

Quick Reviews and Key Concepts

Quick Review I:
- Terms to know: matter, energy (kinetic and potential), elements, atoms (p, e, n), valence, molecules, compounds, inert vs. reactive elements
- Elements to know (selected): O, C, H, N, Ca, P, K, S, Na, Cl, Mg, I, Fe

- Bonds to know: hydrogen, ionic, covalent (single, double, and triple; polar vs nonpolar)

Quick Review II:
- Terms to know: solvent, solute, pH, acid, base, buffer
- Reactions to know: combination (synthesis), decomposition, exchange, redox
- Major concepts: factors influencing rate of reactions

Practical and Conceptual Connections

Chemistry underlies physiology: the way atoms bond and reactions occur governs metabolism, energy production (ATP), and signaling.
Amino acids linking to form proteins illustrate synthesis reactions important for structure and function.
Breakdown of glycogen to glucose demonstrates decomposition reactions in energy mobilization.
ATP energy transfer to glucose exemplifies exchange reactions in metabolism.
Redox chemistry is central to cellular respiration and energy extraction from nutrients.
pH balance and buffering are essential for enzyme activity and general cellular homeostasis.
Understanding solvent and solute behavior helps explain nutrient transport and ion balance.

Formulas and Key Equations (LaTeX)

Synthesis (Combination):
$A + B \rightarrow AB$
Decomposition:
$AB \rightarrow A + B$
Exchange (Displacement):
$AB + C \rightarrow AC + B$
Redox (oxidation of a fuel, reduction of an oxidant; simplified):
$ext{C}6 ext{H}{12} ext{O}6 + 6 \text{O}2 \rightarrow 6 \text{CO}2 + 6 \text{H}2 ext{O}$
Acid dissociation (examples):
$ext{HCl} \rightarrow ext{H}^+ + ext{Cl}^-$
$ext{NaOH} \rightarrow ext{Na}^+ + ext{OH}^-$
Buffer equilibrium (carbonic acid/bicarbonate):
$ext{H}2 ext{CO}3 \rightleftharpoons ext{HCO}_3^- + ext{H}^+$

Note: The content above mirrors the transcript content, reorganized into structured study notes with clear sections, bullet points, and essential equations included in LaTeX. It covers fundamental definitions, element details, atomic structure, bond types, reaction types, solutions/pH/buffers, and key example reactions relevant to general chemistry for biology and physiology contexts.