Chapter 2: Elements, Atomic Structure, Bonds, Chemical Reactions, and Water

Elements, Compounds, and Mixtures

• Definition of Matter: Matter is defined as anything that takes up space and possesses mass.
• States of Matter: Matter exists in four fundamental physical states:
  • Solid (e.g., diamond)
  • Liquid (e.g., juice)
  • Gas (e.g., clouds)
  • Plasma (e.g., ionized neon gas in neon signs)

• Elements:
  • Pure substances that cannot be broken down into other substances by chemical reactions.
  • There are $118$ known elements, of which $92$ occur naturally in nature.
  • Organized systematically on the Periodic Table of Elements based on atomic mass and atomic number.
  • Every element consists of only one unique type of atom.

• Essential Elements vs. Trace Elements:
  • Essential elements are required in large quantities for an organism to survive, grow, and reproduce.
  • Trace elements are required by organisms in minute quantities.
• Compounds:
  • Chemical substances composed of two or more distinct elements combined in a fixed stoichiometric ratio via a chemical reaction.
  • Common examples include water ($\text{H}_2\text{O}$) and table salt ($\text{NaCl}$).
  • The chemical combination alters the physical and chemical properties of the individual constituent elements (emergent properties).
  • Breaking down a compound into its constituent elements requires a chemical reaction.
• Mixtures:
  • Physical combinations of two or more substances that can be separated without chemical reactions.
  • No new chemical bonds are formed between the mixed components.
  • Individual elements or compounds retain their original physical and chemical properties.
  • Components exist in variable, non-fixed ratios.

• Microscopic Representation of Matter:
  • Elements: Single isolated atoms of one color or homonuclear diatomic molecules composed of identical bound atoms.
  • Compounds: Molecules containing two or more different types of chemically bonded atoms in a uniform structure.
  • Mixtures: Heterogeneous or homogeneous physical blendings of independent atoms, diatomic elements, or compound molecules without chemical bonds between distinct species.

Atomic Structure and Subatomic Particles

• Definition of an Atom: The fundamental unit of matter that retains all physical and chemical properties of an element.
• Subatomic Particles:
  • Atoms are constructed from three principal subatomic particles:
  • Protons: Carry a positive electric charge ($+1$). Reside in the central nucleus.
  • Neutrons: Electrically neutral (charge of $0$). Reside in the central nucleus.
  • Electrons: Carry a negative electric charge ($-1$). Orbit within shells surrounding the central nucleus.
• Mass Comparison of Subatomic Particles:
  • Protons and neutrons possess nearly equal masses (approximately $1\,\text{amu}$ each).
  • Electrons possess a mass roughly $\frac{1}{2000}$ that of a proton or neutron, contributing negligibly to total atomic mass.
• Atomic Spatial Architecture:
  • Atoms consist mostly of empty space.
  • Dense central region: Atomic nucleus (containing protons and neutrons).
  • Surrounding region: Electron cloud/shells (containing orbiting electrons).

• Atomic Number ($Z$):
  • Represents the exact number of protons contained within the nucleus of an atom.
  • Defines the specific identity of an element; this number does not change for a given element.
• Atomic Mass / Mass Number ($A$):
  • Represents the total sum of protons and neutrons within the atomic nucleus ($A = \text{Protons} + \text{Neutrons}$).
  • Can vary among atoms of the same element due to variable neutron counts.
• Formulas for Subatomic Calculations:
  • $\text{Protons} = \text{Atomic Number } (Z)$
  • $\text{Neutrons} = \text{Atomic Mass } (A) - \text{Atomic Number } (Z)$
  • $\text{Electrons} = \text{Atomic Number } (Z)$ (in a neutral, uncharged atom)

• Isotopes:
  • Variant forms of a single element containing identical numbers of protons but different numbers of neutrons.
  • Proton count remains constant, ensuring the atomic number ($Z$) is unchanged.
  • Neutron count varies, thereby altering the atomic mass ($A$).
  • Isotopes of an element exhibit identical chemical behavior in chemical reactions because their electron configurations remain identical.
  • Carbon Isotopes Example:
  • Carbon-12 ($^{12}\text{C}$): $6\text{ protons}, 6\text{ neutrons}$
  • Carbon-13 ($^{13}\text{C}$): $6\text{ protons}, 7\text{ neutrons}$ ($1$ extra neutron)
  • Carbon-14 ($^{14}\text{C}$): $6\text{ protons}, 8\text{ neutrons}$ ($2$ extra neutrons)

• Electron Shells and Energy Levels:
  • Electron shells dictate the potential energy levels of electrons.
  • Electrons closer to the nucleus possess lower potential energy; electrons positioned further away possess higher potential energy.
  • Each discrete shell holds a specific maximum capacity of electrons.
• Valence Shell and Valence Electrons:
  • Valence Shell: The outermost electron shell surrounding the nucleus.
  • Valence Electrons: Electrons residing within the valence shell.
  • The valence shell dictates the chemical reactivity and bonding behavior of the atom.
  • Atoms possessing identical numbers of valence electrons exhibit similar chemical characteristics and reactivity patterns.

Chemical Bonds

• Overview of Chemical Bonding: Chemical bonds are attractive forces holding atoms together to achieve stable, filled valence electron shells.

• Covalent Bonds:
  • Formed when two atoms share one or more pairs of valence electrons.
  • Predominantly form between nonmetal atoms.
  • Classifications by Shared Pairs:
  • Single Covalent Bond: Sharing of $1$ pair of valence electrons (e.g., $\text{Cl}_2$ or $\text{H}_2$).
  • Double Covalent Bond: Sharing of $2$ pairs of valence electrons (e.g., $\text{O}_2$).
  • Triple Covalent Bond: Sharing of $3$ pairs of valence electrons (e.g., $\text{N}_2$).
  • Classifications by Polarity:
  • Nonpolar Covalent Bond: Electrons are shared equally between two atoms due to similar electronegativities.
  • Polar Covalent Bond: Electrons are shared unequally due to differences in electronegativity, creating partial negative ($\delta^-$) and partial positive ($\delta^+$) charges across the molecule (e.g., water, $\text{H}_2\text{O}$).

• Ionic Bonds:
  • Occur when a highly electronegative atom completely strips one or more valence electrons away from a weakly electronegative atom.
  • Electron transfer generates fully charged atoms called ions.
  • Cation: A positively charged ion resulting from the loss of electrons ($\text{Protons} > \text{Electrons}$).
  • Anion: A negatively charged ion resulting from the gain of electrons ($\text{Electrons} > \text{Protons}$).
  • The ionic bond is the electrostatic attraction holding oppositely charged cations and anions together (e.g., solid sodium chloride, $\text{NaCl}$ lattice).

• Hydrogen Bonds:
  • Non-covalent electrostatic attractions between a partially positive hydrogen atom (covalently attached to a highly electronegative atom) and another strongly electronegative atom (typically Nitrogen $\text{N}$, Oxygen $\text{O}$, or Fluorine $\text{F}$).
  • Weak relative to covalent or ionic bonds, but powerful in aggregate.
  • Liquid water molecules are linked via dynamic hydrogen bonds; a single water molecule can form up to $4$ hydrogen bonds simultaneously with surrounding water molecules.
  • Responsible for the unique chemical and physical properties of liquid and solid water.

Chemical Reactions and Dynamic Equilibrium

• Definition of Chemical Reactions: A process in which starting chemical substances (reactants) break and form chemical bonds to yield new substances (products).
• Reaction Representation:$\text{Reactants } (A + B) \rightarrow \text{Products } (AB)$

• Characteristics and Indicators of Chemical Reactions:
  • Starting and ending molecular structures differ chemically.
  • Reactions are frequently irreversible under ambient conditions.
  • Observable physical and chemical indicators include:
  • Color changes
  • Temperature changes (exothermic or endothermic)
  • Odor production
  • Gas production or bubbling
  • Release of sound or light energy
• Chemical Equilibrium:
  • Forward Reaction: Reactants collide to produce products ($A + B \rightarrow C + D$).
  • Reverse/Backward Reaction: Products react to re-form original reactants ($C + D \rightarrow A + B$).
  • Equilibrium Progression Steps:
  1. Reaction starts; high concentration of reactants $\text{A}$ and $\text{B}$, high collision rate, high forward reaction rate, zero products.
  2. Products $\text{C}$ and $\text{D}$ accumulate; reactant collisions decrease, forward rate slows down, reverse reaction begins.
  3. Collisions between products increase while reactant collisions continue to decline.
  4. Dynamic Chemical Equilibrium: Established when the rate of the forward reaction equals the rate of the reverse reaction ($\text{Rate}_{\text{forward}} = \text{Rate}_{\text{reverse}}$).
  • At equilibrium, the relative concentrations of reactants and products remain constant over time.

Properties of Water and Aqueous Solutions

• Polarity of Water: Oxygen holds a higher electronegativity than hydrogen, pulling shared electrons toward itself. This confers a partial negative charge ($\delta^-$) on the oxygen atom and a partial positive charge ($\delta^+$) on each hydrogen atom.
• Cohesion:
  • The attraction of water molecules to other water molecules via hydrogen bonding.
  • Enables bulk transport of water columns against gravity in plants.
• Adhesion:
  • The clinging of water molecules to different polar or charged substances.
  • Works alongside cohesion to facilitate capillary action and movement of water up plant cell walls (xylem vessels).
• Transpiration Process in Plants:
  • Roots absorb water from the surrounding soil.
  • Water moves upward through the vascular plant stems via cohesion and adhesion forces.
  • Water vapor evaporates into the atmosphere through microscopic leaf pores called stomata.

• Surface Tension:
  • A measure of how difficult it is to stretch or break the surface of a liquid.
  • Water exhibits high surface tension due to dense hydrogen bonding networks between surface molecules and a lack of bonding to air molecules above.
• Solute, Solvent, and Solution:
  • Solute: The substance being dissolved (e.g., solid salt, $\text{NaCl}$).
  • Solvent: The dissolving medium capable of surrounding solute particles (e.g., liquid water).
  • Solution: A homogenous liquid mixture of solute dissolved in solvent.
  • Aqueous Solution: Any solution in which water serves as the solvent.
  • Water acts as a versatile solvent due to its molecular polarity, which forms hydration shells around polar molecules and charged ions.

• Hydrophilic vs. Hydrophobic:
  • Hydrophilic: "Water-loving"; affinity for water; includes polar or charged molecules that dissolve readily in water.
  • Hydrophobic: "Water-fearing"; lack of affinity for water; includes nonpolar molecules that do not dissolve in water.

Acid-Base Chemistry and the pH Scale

• Acids:
  • Chemical substances that increase the hydrogen ion ($\text{H}^+$) concentration in an aqueous solution.
  • Maintain a $\text{pH}$ value less than $7$ ($\text{pH} < 7$).
  • Turn litmus paper red.
  • Taste sour, feel corrosive or irritating to tissues.
  • Common examples: stomach acid, lemon juice, coffee, soda.
• Bases:
  • Chemical substances that decrease $\text{H}^+$ concentration (or increase hydroxide ion, $\text{OH}^-$, concentration) in solution.
  • Maintain a $\text{pH}$ value greater than $7$ ($\text{pH} > 7$).
  • Turn litmus paper blue.
  • Taste bitter, feel slippery or soapy to the touch.
  • Common examples: baking soda, ammonia solution, soap, bleach.
• Neutralization Reactions:
  • Occurs when a strong acid reacts with a strong base to form a salt and neutral water:     $\text{Strong Acid} + \text{Strong Base} \rightarrow \text{Salt} + \text{Water}$
  • Specific Chemical Reaction Example:     $\text{HCl} + \text{NaOH} \rightarrow \text{NaCl} + \text{H}_2\text{O}$

• The pH Scale and Mathematical Relations:

  • Measures the relative concentration of hydrogen ions ($[\text{H}^+]$) on a logarithmic scale ranging from $0$ to $14$
  • Each single-unit shift on the $\text{pH}$ scale represents a $10\text{-fold}$ change in $\text{H}^+$ concentration.
  • In any aqueous solution at $25^\circ\text{C}$, the product of ion concentrations is constant:     $[\text{H}^+] \times [\text{OH}^-] = 10^{-14}$
  • Mathematical relationship between $\text{pH}$ and $\text{pOH}$:     $\text{pH} + \text{pOH} = 14$

• Complete Quantitative pH Scale Chart:

  • $\text{pH } 0$: $[\text{H}^+] = 10^0\,\text{M}$, $[\text{OH}^-] = 10^{-14}\,\text{M}$, $\text{pOH} = 14$, $10,000,000 \times$ relative to pure water (Examples: battery acid, strong hydrofluoric acid).
  • $\text{pH } 1$: $[\text{H}^+] = 10^{-1}\,\text{M}$, $[\text{OH}^-] = 10^{-13}\,\text{M}$, $\text{pOH} = 13$, $1,000,000 \times$ relative to pure water (Example: hydrochloric acid secreted by stomach lining).
  • $\text{pH } 2$: $[\text{H}^+] = 10^{-2}\,\text{M}$, $[\text{OH}^-] = 10^{-12}\,\text{M}$, $\text{pOH} = 12$, $100,000 \times$ relative to pure water (Examples: lemon juice, gastric acid, vinegar).
  • $\text{pH } 3$: $[\text{H}^+] = 10^{-3}\,\text{M}$, $[\text{OH}^-] = 10^{-11}\,\text{M}$, $\text{pOH} = 11$, $10,000 \times$ relative to pure water (Examples: grapefruit, orange juice, soda).
  • $\text{pH } 4$: $[\text{H}^+] = 10^{-4}\,\text{M}$, $[\text{OH}^-] = 10^{-10}\,\text{M}$, $\text{pOH} = 10$, $1,000 \times$ relative to pure water (Examples: tomato juice, acid rain).
  • $\text{pH } 5$: $[\text{H}^+] = 10^{-5}\,\text{M}$, $[\text{OH}^-] = 10^{-9}\,\text{M}$, $\text{pOH} = 9$, $100 \times$ relative to pure water (Examples: soft drinking water, black coffee).
  • $\text{pH } 6$: $[\text{H}^+] = 10^{-6}\,\text{M}$, $[\text{OH}^-] = 10^{-8}\,\text{M}$, $\text{pOH} = 8$, $10 \times$ relative to pure water (Examples: urine, saliva).
  • $\text{pH } 7$: $[\text{H}^+] = 10^{-7}\,\text{M}$, $[\text{OH}^-] = 10^{-7}\,\text{M}$, $\text{pOH} = 7$, $1 \times$ relative to pure water (Neutral pure water).
  • $\text{pH } 8$: $[\text{H}^+] = 10^{-8}\,\text{M}$, $[\text{OH}^-] = 10^{-6}\,\text{M}$, $\text{pOH} = 6$, $\frac{1}{10} \times$ relative to pure water (Example: sea water).
  • $\text{pH } 9$: $[\text{H}^+] = 10^{-9}\,\text{M}$, $[\text{OH}^-] = 10^{-5}\,\text{M}$, $\text{pOH} = 5$, $\frac{1}{100} \times$ relative to pure water (Example: baking soda).
  • $\text{pH } 10$: $[\text{H}^+] = 10^{-10}\,\text{M}$, $[\text{OH}^-] = 10^{-4}\,\text{M}$, $\text{pOH} = 4$, $\frac{1}{1,000} \times$ relative to pure water (Examples: Great Salt Lake, milk of magnesia).
  • $\text{pH } 11$: $[\text{H}^+] = 10^{-11}\,\text{M}$, $[\text{OH}^-] = 10^{-3}\,\text{M}$, $\text{pOH} = 3$, $\frac{1}{10,000} \times$ relative to pure water (Example: ammonia solution).
  • $\text{pH } 12$: $[\text{H}^+] = 10^{-12}\,\text{M}$, $[\text{OH}^-] = 10^{-2}\,\text{M}$, $\text{pOH} = 2$, $\frac{1}{100,000} \times$ relative to pure water (Example: soapy water).
  • $\text{pH } 13$: $[\text{H}^+] = 10^{-13}\,\text{M}$, $[\text{OH}^-] = 10^{-1}\,\text{M}$, $\text{pOH} = 1$, $\frac{1}{1,000,000} \times$ relative to pure water (Examples: bleaches, oven cleaner).
  • $\text{pH } 14$: $[\text{H}^+] = 10^{-14}\,\text{M}$, $[\text{OH}^-] = 10^0\,\text{M}$, $\text{pOH} = 0$, $\frac{1}{10,000,000} \times$ relative to pure water (Example: liquid drain cleaner).

• Buffers:
  • Substances that minimize drastic changes in the concentrations of $\text{H}^+$ and $\text{OH}^-$ in a solution.
  • Maintain $\text{pH}$ stability by accepting hydrogen ions when they are in excess and donating hydrogen ions when they are depleted.
• Ocean Acidification:
  • Environmental phenomenon caused by excess atmospheric carbon dioxide ($\text{CO}_2$) dissolving into ocean water.
  • Dissolved $\text{CO}_2$ reacts with water to form carbonic acid, lowering ocean $\text{pH}$ and threatening marine ecosystems.