Atomic Structure, Periodic Classification, Chemical Bonding, and Chemical Equations

The Structure of the Atom and Periodic Classification of the Elements

Dalton’s Atomic Theory vs. Modern Atomic Theory

Understanding the evolution of atomic theory is fundamental to chemistry. Dalton's atomic theory laid the groundwork for modern concepts, though it has been refined over time.

• Dalton’s Atomic Theory: Proposed that matter is composed of indivisible atoms, atoms of the same element are identical, and chemical reactions involve rearrangement of atoms.

• Modern Atomic Theory: Recognizes that atoms are divisible into subatomic particles (protons, neutrons, electrons), and atoms of the same element can have different masses (isotopes).

• Example: Isotopes of carbon: 12C and 14C.

Atomic Notation

Atomic notation is a standardized way to represent elements and their isotopes.

• Notation: AZX, where X is the element symbol, A is the mass number, and Z is the atomic number.

• Example: 2311Na represents sodium with mass number 23 and atomic number 11.

Atomic Number, Mass Number, Atomic Mass

These terms describe fundamental properties of atoms.

• Atomic Number (Z): Number of protons in the nucleus; defines the element.

• Mass Number (A): Total number of protons and neutrons in the nucleus.

• Atomic Mass: Weighted average mass of all isotopes of an element (in atomic mass units, amu).

• Formula: $A = Z + N$ (where N = number of neutrons)

The Periodic Table

The periodic table organizes elements by increasing atomic number and recurring chemical properties.

• Groups: Vertical columns; elements in the same group have similar chemical properties.

• Periods: Horizontal rows; properties change progressively across a period.

• Example: Alkali metals (Group 1) are highly reactive metals.

Metals, Non-metals, and Metalloids

Elements are classified based on their physical and chemical properties.

• Metals: Good conductors, malleable, ductile, shiny (e.g., iron, copper).

• Non-metals: Poor conductors, brittle, dull (e.g., oxygen, sulfur).

• Metalloids: Properties intermediate between metals and non-metals (e.g., silicon).

Energy Levels and Subshells

Electrons in atoms occupy specific energy levels (shells) and subshells.

• Principal Quantum Number (n): Indicates the main energy level.

• Subshells: s, p, d, f; each has a specific shape and energy.

• Example: The 2p subshell can hold up to 6 electrons.

Electronic Configuration in Terms of Shells (Quantum Number n)

Electronic configuration describes the arrangement of electrons in an atom.

• Notation: Lists the number of electrons in each subshell (e.g., 1s2 2s2 2p6).

• Aufbau Principle: Electrons fill lower energy orbitals first.

Periodic Trends: Atomic Size, Ionization Energy, Electron Affinity, Electronegativity

Periodic trends describe how certain properties of elements change across the periodic table.

• Atomic Size: Decreases across a period, increases down a group.

• Ionization Energy: Energy required to remove an electron; increases across a period, decreases down a group.

• Electron Affinity: Energy change when an atom gains an electron; generally becomes more negative across a period.

• Electronegativity: Tendency of an atom to attract electrons in a bond; increases across a period, decreases down a group.

Chemical Bonding

Stability of Noble Gases and Other Elements

Noble gases are chemically inert due to their complete valence electron shells.

• Octet Rule: Atoms tend to gain, lose, or share electrons to achieve a full valence shell (usually 8 electrons).

• Example: Neon (Ne) is stable and unreactive.

Valency and Valence Electrons

Valency is the combining capacity of an element, determined by the number of valence electrons.

• Valence Electrons: Electrons in the outermost shell.

• Valency: Number of electrons an atom can gain, lose, or share to achieve stability.

Formulae of Cations/Anions

Cations are positively charged ions; anions are negatively charged ions.

• Cation: Formed by loss of electrons (e.g., Na+).

• Anion: Formed by gain of electrons (e.g., Cl-).

Ionization Energy (Enthalpy of Ionization) and Electron Affinity (Enthalpy of Electron Attachment)

These properties describe the energy changes associated with ion formation.

• Ionization Energy: $\text{A} \rightarrow \text{A}^+ + e^-$

• Electron Affinity: $\text{A} + e^- \rightarrow \text{A}^-$

Ionic and Covalent Bonds

Chemical bonds form when atoms achieve stable electron configurations.

• Ionic Bond: Transfer of electrons from metal to non-metal, forming ions (e.g., NaCl).

• Covalent Bond: Sharing of electron pairs between non-metals (e.g., H2O).

Metallic Bonding

Metallic bonding involves a 'sea' of delocalized electrons shared among metal atoms.

• Properties: High electrical conductivity, malleability, ductility.

Lewis Structures of Bonding Arrangements

Lewis structures represent the arrangement of electrons in molecules.

• Dots: Represent valence electrons.

• Lines: Represent shared electron pairs (bonds).

• Example: Lewis structure of water: H:O:H (with two lone pairs on O).

Chemical Formulae

Chemical formulae indicate the types and numbers of atoms in a compound.

• Empirical Formula: Simplest whole-number ratio of elements.

• Molecular Formula: Actual number of atoms of each element in a molecule.

Chemical Formulae and Chemical Equations

The Mole and Avogadro’s Constant

The mole is the SI unit for amount of substance; Avogadro’s constant defines the number of particles in one mole.

• Avogadro’s Constant: $6.022 \times 10^{23}$ particles/mol.

• Example: 1 mole of H2O contains $6.022 \times 10^{23}$ molecules.

Molar Mass, Atomic Mass

Molar mass is the mass of one mole of a substance; atomic mass is the mass of a single atom.

• Molar Mass: Expressed in g/mol; numerically equal to atomic or molecular mass in amu.

• Formula: $\text{Molar mass} = \frac{\text{Mass of sample (g)}}{\text{Amount of substance (mol)}}$

Empirical and Molecular Formulae

Empirical formula gives the simplest ratio; molecular formula gives the actual number of atoms.

• Example: Glucose: Empirical formula CH2O, molecular formula C6H12O6.

The Law of Conservation of Mass

Mass is neither created nor destroyed in a chemical reaction.

• Implication: Total mass of reactants equals total mass of products.

• Example: $2H_2 + O_2 \rightarrow 2H_2O$ (mass is conserved).

Chemical Equations

Chemical equations represent chemical reactions using symbols and formulas.

• Reactants: Substances consumed.

• Products: Substances formed.

• Balanced Equation: Same number of each atom on both sides.

Stoichiometric Calculations

Stoichiometry involves quantitative relationships in chemical reactions.

• Steps:

  1. Write and balance the chemical equation.

  2. Convert given quantities to moles.

  3. Use mole ratios to find unknown quantities.

  4. Convert moles to desired units (grams, liters, etc.).

• Example: How many grams of H2O are produced from 4 g of H2?