2.4 Chemistry Fundamentals: Atomic Structure, Elements & Chemical Reactions

Chemistry Fundamentals: Atomic Structure, Elements & Chemical Reactions

Chemistry is the central science investigating matter, its properties, how substances combine or separate to form other substances, and how substances interact with energy.

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1. Atomic Structure & Subatomic Particles

All matter is composed of atoms. An atom is the smallest unit of an element that retains the chemical identity of that element.

Subatomic Particle Properties

Fundamental Atomic Metrics

• Atomic Number ($Z$): The number of protons in the nucleus. Defines the identity of an element (e.g., Carbon always has $Z = 6$). In a neutral atom, Protons = Electrons.
• Mass Number ($A$): The total count of protons + neutrons in the nucleus: $A = Z + N$.
• Isotopes: Atoms of the same element (same proton count $Z$) that contain different numbers of neutrons ($N$), resulting in different mass numbers (e.g., Carbon-12 with 6 neutrons vs. Carbon-14 with 8 neutrons).
• Ions: Charged atoms formed when an atom gains or loses electrons.
  • Cation: Positively charged ion formed by losing electrons (e.g., $Na^+$).
  • Anion: Negatively charged ion formed by gaining electrons (e.g., $Cl^-$).

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2. The Periodic Table of Elements

The periodic table arranges chemical elements by increasing atomic number ($Z$) in horizontal rows (Periods) and vertical columns (Groups or Families).

Major Element Groups

• Group 1 (Alkali Metals): Highly reactive metals with 1 valence electron (e.g., $Li, Na, K$). React violently with water.
• Group 2 (Alkaline Earth Metals): Reactive metals with 2 valence electrons (e.g., $Mg, Ca$).
• Groups 3–12 (Transition Metals): Hard, dense metals with variable valence states (e.g., $Fe, Cu, Au, Ag$).
• Group 17 (Halogens): Highly reactive nonmetals with 7 valence electrons (e.g., $F, Cl, Br, I$). Form salts when combined with metals.
• Group 18 (Noble Gases): Inert, unreactive gases with a full outer electron shell (8 valence electrons, octet rule; e.g., $He, Ne, Ar$).

Key Periodic Trends

• Atomic Radius: Increases moving down a group (adding electron shells); decreases moving left-to-right across a period.
• Electronegativity: An atom's ability to attract shared electrons. Increases moving left-to-right across a period and up a group (Fluorine is the most electronegative element).
• Ionization Energy: Energy required to remove an electron from an atom. Increases moving left-to-right across a period and up a group.

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3. Chemical Bonding & Molecular Structure

Atoms form chemical bonds to achieve lower potential energy and stable outer electron configurations (typically 8 valence electrons, fulfilling the Octet Rule).

• Nonpolar Covalent Bond: Equal sharing of electrons between atoms with similar electronegativity (e.g., $O_2, H_2$).
• Polar Covalent Bond: Unequal sharing of electrons due to electronegativity differences, creating partial positive ($\delta+$) and negative ($\delta-$) poles (e.g., $H_2O$).

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4. Chemical Reactions & Balancing Equations

A chemical reaction rearranges atoms to transform reactants into products, breaking and forming chemical bonds.

Law of Conservation of Mass: Mass is neither created nor destroyed during a chemical reaction. The total mass of reactants must equal the total mass of products, requiring chemical equations to be balanced.

Five Primary Types of Chemical Reactions

1. Synthesis (Combination): Two or more reactants combine to form a single product. $A + B \rightarrow AB \quad (e.g., 2H_2 + O_2 \rightarrow 2H_2O)$
2. Decomposition: A single compound breaks down into two or more simpler substances. $AB \rightarrow A + B \quad (e.g., 2H_2O \rightarrow 2H_2 + O_2)$
3. Single Replacement (Displacement): One uncombined element replaces another element in a compound. $A + BC \rightarrow AC + B \quad (e.g., Zn + 2HCl \rightarrow ZnCl_2 + H_2)$
4. Double Replacement: Ions from two reacting compounds swap partners to form two new compounds. $AB + CD \rightarrow AD + CB \quad (e.g., AgNO_3 + NaCl \rightarrow AgCl \downarrow + NaNO_3)$
5. Combustion: A hydrocarbon reacts rapidly with oxygen gas ($O_2$) to release heat, carbon dioxide, and water. $\text{Hydrocarbon} + O_2 \rightarrow CO_2 + H_2O + \text{Heat}$

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5. Acids, Bases, and the pH Scale

Acid-Base Definitions

• Arrhenius Definition: Acids produce hydrogen ions ($H^+$) in aqueous solution; Bases produce hydroxide ions ($OH^-$).
• Brønsted-Lowry Definition: Acids are proton ($H^+$) donors; Bases are proton ($H^+$) acceptors.

The pH Scale

The pH scale measures the logarithmic concentration of hydrogen ions in a solution: $\text{pH} = -\log[H^+]$.

$\text{ACIDIC } (\text{pH } 0\text{--}6.9) \longleftarrow \mathbf{\text{NEUTRAL } (\text{pH } 7.0)} \longrightarrow \text{BASIC / ALKALINE } (\text{pH } 7.1\text{--}14)$

Neutralization Reaction: Combining an acid and a base yields a salt and water: $\text{Acid} + \text{Base} \rightarrow \text{Salt} + \text{Water}$ (e.g., $HCl + NaOH \rightarrow NaCl + H_2O$).

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6. Solutions & Solubility

• Solution: A homogeneous mixture of two or more substances.
• Solute: The substance being dissolved (e.g., salt, sugar).
• Solvent: The dissolving medium (e.g., water, known as the "universal solvent").
• Molarity ($M$): A measure of solution concentration expressed as moles of solute per liter of solution ($M = \frac{\text{moles of solute}}{\text{liters of solution}}$).
• Factors Increasing Dissolution Rate: Increasing temperature (for solid solutes), stirring/agitation, and increasing solute surface area (crushing).