6.2 Chemistry Fundamentals, Atomic Structure, and Reactions

6.2 Chemistry Fundamentals, Atomic Structure, and Reactions

Core Principle: Chemistry forms approximately 20% to 25% of the CAT-ASVAB General Science subtest. Success requires immediate fluency with subatomic particles, atomic number versus mass number calculations, electron shell configurations, Periodic Table groups and periodic trends, ionic and covalent bonding mechanisms, chemical equation balancing, endothermic and exothermic thermodynamics, and the logarithmic pH scale.

At a pace of roughly 48 seconds per item, candidates must recognize chemical symbols, reaction archetypes, and valence rules on sight — there is time to check an answer, but none to derive one from first principles.

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1. Subatomic Structure & Atomic Anatomy

All physical matter in the universe is composed of atoms, the fundamental building blocks of chemical elements that retain the characteristic properties of that element.

+-----------------------------------------------------------------------------------------+
|                                SUBATOMIC PARTICLE PROPERTIES                            |
+------------+---------------+-------------------+--------------------+-------------------+
| Particle   | Charge        | Mass (amu)        | Mass (kg)          | Location          |
+------------+---------------+-------------------+--------------------+-------------------+
| Proton     | Positive (+1) | ~1.0073 amu (~1)  | 1.673 x 10^-27 kg  | Dense Nucleus     |
| Neutron    | Neutral (0)   | ~1.0087 amu (~1)  | 1.675 x 10^-27 kg  | Dense Nucleus     |
| Electron   | Negative (-1) | ~0.00055 amu (~0) | 9.109 x 10^-31 kg  | Electron Orbitals |
+------------+---------------+-------------------+--------------------+-------------------+

Essential Atomic Metrics

• Atomic Number ($Z$): The number of protons residing in the nucleus. The atomic number uniquely defines the chemical identity of the element (e.g., any atom possessing exactly 6 protons is Carbon; 7 protons is Nitrogen; 8 protons is Oxygen). In an electrically neutral atom: $\text{Number of Protons} = \text{Number of Electrons} = Z$
• Mass Number ($A$): The total integer count of protons + neutrons residing in the nucleus: $A = Z + N \quad \Longrightarrow \quad \text{Neutrons } (N) = A - Z$
• Atomic Mass (Weight): The weighted average mass of all naturally occurring isotopes of an element, measured in atomic mass units (amu or Daltons).
• Isotopes: Atoms of the same element containing an identical number of protons ($Z$) but differing numbers of neutrons ($N$), resulting in different mass numbers.
  • Example: Carbon-12 ($^{12}{\ 6}\text{C}$: 6 protons, 6 neutrons; stable) versus Carbon-14 ($^{14}{\ 6}\text{C}$: 6 protons, 8 neutrons; radioactive isotope used in radiometric carbon dating).
  • Example: Hydrogen isotopes: Protium ($^1\text{H}$: 1p, 0n), Deuterium ($^2\text{H}$: 1p, 1n, used in heavy water reactors), and Tritium ($^3\text{H}$: 1p, 2n, radioactive).

Electron Shells, Energy Levels & The Octet Rule

Electrons occupy quantized three-dimensional probability orbitals organized into principal energy levels ($n = 1, 2, 3, 4\dots$). The maximum electron capacity of each principal shell is determined by the formula:

$\text{Max Electrons in Shell } n = 2n^2$

• Shell $n=1$ ($K$ shell): Holds up to $2(1)^2 = \mathbf{2\text{ electrons}}$
• Shell $n=2$ ($L$ shell): Holds up to $2(2)^2 = \mathbf{8\text{ electrons}}$
• Shell $n=3$ ($M$ shell): Holds up to $2(3)^2 = \mathbf{18\text{ electrons}}$
• Shell $n=4$ ($N$ shell): Holds up to $2(4)^2 = \mathbf{32\text{ electrons}}$

Valence Electrons & The Octet Rule: Valence electrons are the electrons occupying the outermost electron shell. Chemical reactivity is driven by the Octet Rule—atoms tend to gain, lose, or share electrons until their valence shell contains 8 electrons, achieving the stable electronic configuration of a noble gas (Helium is stable with a complete duet of 2 electrons).

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2. The Periodic Table & Periodic Trends

The Periodic Table of Elements organizes all 118 known elements in order of increasing atomic number ($Z$).

+-----------------------------------------------------------------------------------------+
|                                PERIODIC TABLE STRUCTURE                                 |
+------------------------------------+----------------------------------------------------+
| PERIODS (Horizontal Rows, 1 to 7)  | GROUPS / FAMILIES (Vertical Columns, 1 to 18)      |
+------------------------------------+----------------------------------------------------+
| • Indicate number of occupied      | • Elements share identical valence electron counts |
|   electron energy shells           | • Exhibit remarkably similar chemical reactivities |
+------------------------------------+----------------------------------------------------+

Major Chemical Families

1. Group 1: Alkali Metals (Li, Na, K, Rb, Cs, Fr): Contain 1 valence electron ($s^1$). Extremely reactive, soft, silvery metals with low melting points. Readily lose 1 electron to form $+1$ cations ($\text{Na}^+$). React vigorously with water to produce hydrogen gas and basic metal hydroxides.
2. Group 2: Alkaline Earth Metals (Be, Mg, Ca, Sr, Ba, Ra): Contain 2 valence electrons ($s^2$). Highly reactive metals (less reactive than Group 1). Readily lose 2 electrons to form $+2$ cations ($\text{Ca}^{2+}$). Abundant in mineral rock formations.
3. Groups 3–12: Transition Metals (Fe, Cu, Ni, Ag, Au, W, Pt): Hard, dense metallic elements characterized by partially filled $d$-orbitals. Exhibit high tensile strength, high melting points, superior electrical and thermal conductivity, and variable positive oxidation states (e.g., $\text{Fe}^{2+}$ and $\text{Fe}^{3+}$).
4. Group 17: Halogens (F, Cl, Br, I, At): Contain 7 valence electrons ($s^2 p^5$). Highly reactive nonmetals that readily gain 1 electron to form $-1$ halide anions ($\text{Cl}^-$). Form ionic salts when bonded with alkali metals (e.g., $\text{NaCl}$). Exist as diatomic molecules in elemental form ($\text{F}_2, \text{Cl}_2, \text{Br}_2, \text{I}_2$).
5. Group 18: Noble Gases (He, Ne, Ar, Kr, Xe, Rn): Contain a full outer octet of 8 valence electrons (Helium has 2). Chemically inert (unreactive), colorless, odorless monatomic gases under standard conditions.

+-----------------------------------------------------------------------------------------+
|                                  PERIODIC TABLE TRENDS                                  |
+-----------------------------------------------------------------------------------------+
|                                ELECTRONEGATIVITY INCREASES  ----------------------->    |
|                                IONIZATION ENERGY INCREASES  ----------------------->    |
|                              <--- ATOMIC RADIUS INCREASES                               |
|   +---------------------------------------------------------------------------------+   |
|   | Group 1 (Alkali)                                             Group 18 (Nobles) | |  |
|   | Period 1                                                                        | |  |
|   | Period 2                                                     F (Highest EN)     | v  |
|   | Period 3                                                                        |    |
|   | Period 4                                                                        | I  |
|   | Period 5                                                                        | N  |
|   | Period 6                                                                        | C  |
|   | Period 7   Fr (Largest Radius)                                                  | R  |
|   +---------------------------------------------------------------------------------+    |
|                                                                                          |
|   * Atomic Radius INCREASES moving DOWN a Group and LEFT across a Period.                |
|   * Electronegativity & Ionization Energy INCREASE moving UP a Group and RIGHT.         |
+-----------------------------------------------------------------------------------------+

The Three Fundamental Periodic Trends

• Atomic Radius: The distance from the atomic nucleus to the outermost stable electron orbital.
  • Trend: Increases moving down a group (each successive period adds a new principal electron shell). Decreases moving left-to-right across a period (increasing nuclear positive charge pulls the valence electron cloud closer inward). Francium (Fr) has the largest atomic radius.
• Electronegativity: The relative measure of an atom's ability to attract shared electrons within a chemical bond.
  • Trend: Increases moving left-to-right across a period and decreases moving down a group. Fluorine (F) is the most electronegative element ($4.0$ on the Pauling scale), followed by Oxygen ($3.5$) and Nitrogen ($3.0$). Noble gases generally lack defined electronegativity values.
• Ionization Energy: The minimum energy required to remove the most loosely held valence electron from an isolated gaseous atom in its ground state.
  • Trend: Follows electronegativity—increases left-to-right across a period and decreases moving down a group. Helium (He) has the highest first ionization energy.

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3. Chemical Bonding Mechanisms

Chemical bonds are electrostatic attractions holding atoms together to achieve lower potential energy and stable octet configurations:

Hydrogen Bonding (Intermolecular Force)

A special, exceptionally strong intermolecular dipole-dipole attraction occurring when hydrogen is covalently bonded to a highly electronegative, small atom (Nitrogen, Oxygen, or Fluorine).

• Hydrogen bonding gives water ($\text{H}_2\text{O}$) its uniquely high boiling point ($100^\circ\text{C}$), high surface tension, high specific heat capacity, and lower density as a solid than as a liquid (causing ice to float).

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4. Chemical Reactions & Stoichiometry

A chemical reaction involves the breaking of chemical bonds in reactant substances and the formation of new bonds to create chemically distinct products.

The Law of Conservation of Mass

Formulated by Antoine Lavoisier: Matter can neither be created nor destroyed in a chemical reaction. The total mass of the reactants must exactly equal the total mass of the products. Consequently, a chemical equation must have an identical count of each atom on both sides of the reaction arrow.

Balancing Chemical Equations Example:

Unbalanced: $\text{Fe} + \text{O}_2 \longrightarrow \text{Fe}_2\text{O}_3$

• Left: 1 Fe, 2 O | Right: 2 Fe, 3 O
• Multiply $\text{Fe}_2\text{O}_3$ by 2: $\text{Fe} + \text{O}_2 \longrightarrow 2,\text{Fe}_2\text{O}_3$ (4 Fe, 6 O)
• Balance iron and oxygen: $\mathbf{4,\text{Fe} + 3,\text{O}_2 \longrightarrow 2,\text{Fe}_2\text{O}_3}$ (Balanced!)

+-----------------------------------------------------------------------------------------+
|                                FIVE CORE REACTION ARCHETYPES                            |
+-------------------+-----------------------------------+---------------------------------+
| Reaction Type     | General Algebraic Model           | Representative Chemical Example |
+-------------------+-----------------------------------+---------------------------------+
| 1. Synthesis      | A + B  ------->  AB               | 2 H2 + O2  ------->  2 H2O      |
|    (Combination)  |                                   | N2 + 3 H2  ------->  2 NH3      |
+-------------------+-----------------------------------+---------------------------------+
| 2. Decomposition  | AB  ------->  A + B               | 2 H2O2  ------>  2 H2O + O2     |
|                   |                                   | 2 KClO3  ----->  2 KCl + 3 O2   |
+-------------------+-----------------------------------+---------------------------------+
| 3. Single         | A + BC  ------>  AC + B           | Zn + 2 HCl  --->  ZnCl2 + H2    |
|    Displacement   | (Active metal displaces cation)   | Fe + CuSO4  --->  FeSO4 + Cu    |
+-------------------+-----------------------------------+---------------------------------+
| 4. Double         | AB + CD  ----->  AD + CB          | AgNO3 + NaCl -> AgCl(s) + NaNO3 |
|    Displacement   | (Precipitation or Neutralization) | HCl + NaOH  --->  NaCl + H2O    |
+-------------------+-----------------------------------+---------------------------------+
| 5. Combustion     | CxHy + O2  --->  CO2 + H2O + Heat | CH4 + 2 O2  --->  CO2 + 2 H2O   |
|                   | (Hydrocarbon oxidizes rapidly)    | 2 C8H18 + 25 O2 -> 16 CO2+18 H2O|
+-------------------+-----------------------------------+---------------------------------+

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5. Reaction Energetics: Exothermic vs. Endothermic

All chemical reactions involve energy changes governed by the enthalpy of reaction ($\Delta H$):

+-----------------------------------------------------------------------------------------+
|                           REACTION THERMODYNAMICS & ENERGETICS                          |
+------------------------------------+----------------------------------------------------+
| EXOTHERMIC REACTIONS (ΔH < 0)      | ENDOTHERMIC REACTIONS (ΔH > 0)                     |
+------------------------------------+----------------------------------------------------+
| • Releases heat to surroundings    | • Absorbs heat from surroundings                   |
| • Temperature of surroundings rises| • Temperature of surroundings drops (feels cold)   |
| • Products have LOWER potential    | • Products have HIGHER potential                   |
|   energy than reactants            |   energy than reactants                            |
| • Examples: Combustion, dynamite,  | • Examples: Photosynthesis, chemical cold packs,   |
|   hand warmers, acid-base neutral. |   baking soda + vinegar, electrolysis              |
+------------------------------------+----------------------------------------------------+

• Activation Energy ($E_a$): The minimum kinetic energy required by colliding reactant molecules to reach the high-energy transition state and initiate a chemical reaction.
• Catalysts: Substances that accelerate the rate of a chemical reaction by providing an alternative reaction pathway with a lower activation energy ($E_a$). Catalysts are neither consumed nor permanently altered during the reaction (biological catalysts are termed enzymes).

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6. States of Matter, Phase Transitions & Solutions

Matter exists primarily in four classical states: Solid (fixed volume and shape; particles vibrate in fixed lattice), Liquid (fixed volume, variable shape taking container contour; particles slide past one another), Gas (variable volume and shape; particles in rapid, random motion with large intermolecular spacing), and Plasma (high-temperature ionized gas containing free electrons and positive ions; found in stars, lightning, and fusion reactors).

+-----------------------------------------------------------------------------------------+
|                                PHASE TRANSITIONS OF MATTER                              |
+------------------------------------+----------------------------------------------------+
| TRANSITION NAME                    | PHYSICAL STATE CHANGE MECHANICS                    |
+------------------------------------+----------------------------------------------------+
| • Melting (Fusion)                 | Solid  -------->  Liquid  (Endothermic: absorbs Q) |
| • Freezing (Solidification)        | Liquid -------->  Solid   (Exothermic: releases Q) |
| • Vaporization (Boiling/Evap.)     | Liquid -------->  Gas     (Endothermic: absorbs Q) |
| • Condensation                     | Gas    -------->  Liquid  (Exothermic: releases Q) |
| • Sublimation                      | Solid  -------->  Gas     (Direct jump, e.g., Dry  |
|                                    |                           Ice / solid CO2, iodine) |
| • Deposition                       | Gas    -------->  Solid   (Direct jump, e.g., frost|
|                                    |                           forming on cold windows) |
+------------------------------------+----------------------------------------------------+

Solutions, Solutes & Solubility

• Solution: A homogeneous mixture composed of two or more substances.
  • Solute: The dissolved substance present in the lesser quantity (e.g., salt, sugar).
  • Solvent: The dissolving medium present in the greater quantity (water is the "universal solvent").
• Solubility States:
  • Unsaturated: Can dissolve additional solute at the given temperature.
  • Saturated: Contains the maximum equilibrium amount of dissolved solute; excess solute remains as precipitate at the bottom.
  • Supersaturated: Contains more dissolved solute than normal equilibrium capacity, achieved by heating and slowly cooling without agitation.
• Factors Influencing Solubility:
  • Solid in Liquid: Solubility increases with increasing temperature.
  • Gas in Liquid: Solubility decreases with increasing temperature (warm soda goes flat rapidly) and increases with increasing pressure (Henry's Law).

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7. Acid-Base Chemistry & The Logarithmic pH Scale

Acids and bases are defined chemically by two dominant scientific frameworks:

1. Arrhenius Theory: Acids produce hydrogen ions ($\text{H}^+$ / hydronium $\text{H}_3\text{O}^+$) in aqueous solution; Bases produce hydroxide ions ($\text{OH}^-$).
2. Brønsted-Lowry Theory: Acids are proton ($\text{H}^+$) donors; Bases are proton ($\text{H}^+$) acceptors.

+-----------------------------------------------------------------------------------------+
|                                  THE LOGARITHMIC pH SCALE                               |
+-----------------------------------------------------------------------------------------+
|  0 ------- 1 ------- 2 ------- 3 ------- 4 ------- 5 ------- 6 ------- 7                |
|  [<======================= STRONGLY ACIDIC =======================>] [NEUTRAL]          |
|  0: Battery Acid (1 M HCl)         3: Vinegar / Soda                  7: Pure H2O       |
|  1: Stomach Gastric Acid (HCl)     4: Tomato Juice                    (Distilled)       |
|  2: Lemon Juice                    5: Black Coffee / Acid Rain                          |
|                                                                                         |
|  7 ------- 8 ------- 9 ------- 10 ------ 11 ------ 12 ------ 13 ------ 14               |
|  [NEUTRAL] [<======================= STRONGLY BASIC / ALKALINE ======================>] |
|  7.4: Human Blood (Strict)         10: Milk of Magnesia               13: Bleach        |
|  8.0: Seawater                     11: Household Ammonia              14: Lye (1 M NaOH)|
|  9.0: Baking Soda (NaHCO3)         12: Soapy Water                                      |
+-----------------------------------------------------------------------------------------+

Properties of Acids vs. Bases

• The Logarithmic Nature of pH: Because $\text{pH} = -\log_{10}[\text{H}^+]$, each whole unit change on the pH scale represents a 10-fold change in hydrogen ion concentration:
  • A solution with $\text{pH} = 3$ is 10 times more acidic than a solution with $\text{pH} = 4$.
  • A solution with $\text{pH} = 2$ is 100 times ($10^2$) more acidic than $\text{pH} = 4$.
  • A solution with $\text{pH} = 1$ is 1,000 times ($10^3$) more acidic than $\text{pH} = 4$.

Neutralization Reactions

When an acid reacts with an equimolar quantity of a base, the hydrogen ions and hydroxide ions combine to form water and an ionic salt:

$\text{Acid} + \text{Base} \longrightarrow \text{Salt} + \text{Water}$ $\text{HCl (Hydrochloric Acid)} + \text{NaOH (Sodium Hydroxide)} \longrightarrow \text{NaCl (Table Salt)} + \text{H}_2\text{O (Water)}$

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8. Real-World Military & Tactical Applications

1. CBRN Chemical Hazard Identification: Military decontamination units utilize knowledge of pH and neutralization chemistry to neutralize chemical agents (such as neutralizing acidic chlorine or mustard vesicant residues with alkaline hypochlorite solutions).
2. Munitions & Propellant Combustion: Solid rocket boosters and artillery propellants rely on rapid, highly exothermic oxidation-reduction reactions where nitrocellulose combusts to generate immense volumes of expanding hot gases ($\text{CO}_2, \text{H}_2\text{O}, \text{N}_2$) inside gun chambers.
3. Naval Battery & Galvanic Protection: Submarines and surface warships install sacrificial zinc anodes on steel hulls; zinc oxidizes preferentially due to its lower ionization potential, protecting the iron hull from corrosive electrochemical seawater oxidation.