Atomic Structure and the Periodic Table

Atomic Structure:

Atoms are the fundamental units of matter, composed of three main subatomic particles:

1. Protons:

   • Charge: Positively charged (+1).

   • Location: Found in the nucleus.

   • Mass: Approximately 1 atomic mass unit (amu).

   • Role: Determines the element's identity (atomic number).

2. Neutrons:

   • Charge: Neutral (0 charge).

   • Location: Found in the nucleus.

   • Mass: Approximately 1 amu.

   • Role: Contributes to the mass number and stabilizes the nucleus.

3. Electrons:

   • Charge: Negatively charged (-1).

   • Location: Orbit the nucleus in electron shells.

   • Mass: Negligible compared to protons and neutrons (~1/1836 amu).

   • Role: Involved in chemical bonding and reactions.

Key Terms:

• Atomic Number (Z): The number of protons in an atom, unique to each element.

• Mass Number (A): The sum of protons and neutrons in an atom's nucleus.

• Isotopes: Variants of a given element that have the same number of protons but different numbers of neutrons.

Periodic Table:

The periodic table organizes elements based on their atomic number and properties.

Groups:

• Vertical Columns: Elements in the same group have similar chemical properties because they have the same number of valence electrons.

• Group Examples:

  • Group 1: Alkali metals (e.g., Lithium, Sodium)

  • Group 17: Halogens (e.g., Fluorine, Chlorine)

  • Group 18: Noble gasses (e.g., Helium, Neon)

Periods:

• Horizontal Rows: Elements in the same period have increasing atomic numbers from left to right.

• Trends in Periods: Properties of elements change progressively across a period.

Trends:

• Atomic Radius:

  • Increases down a group: Additional electron shells increase the distance between the nucleus and the outermost electrons.

  • Decreases across a period: Increased nuclear charge pulls electrons closer to the nucleus.

• Ionization Energy:

  • Decreases down a group: Outer electrons are further from the nucleus and easier to remove.

  • Increases across a period: Increased nuclear charge makes it harder to remove an electron.

• Electronegativity:

  • Decreases down a group: Atoms with larger radii have a weaker pull on bonding electrons.

  • Increases across a period: Atoms with higher nuclear charge attract bonding electrons more strongly.

Bonding, Structure, and the Properties of Matter

Types of Bonding:

1. Ionic Bonding:

   • Description: Transfer of electrons from metals to nonmetals, forming ions.

   • Example: NaCl (sodium chloride).

2. Covalent Bonding:

   • Description: Sharing of electrons between nonmetals.

   • Example: Di-Hydrogen Mono-oxide(water).

3. Metallic Bonding:

   • Description: Delocalized electrons shared among a lattice of metal cations.

   • Example: Cu (copper).

Structures:

1. Ionic Compounds:

   • Structure: Giant ionic lattices.

   • Properties: High melting and boiling points; conduct electricity when molten or dissolved.

2. Simple Molecular (Covalent):

   • Structure: Small molecules.

   • Properties: Low melting and boiling points; do not conduct electricity.

3. Giant Covalent Structures:

   • Examples: Diamond (each carbon bonded to four others) and graphite (layers of carbon atoms).

   • Properties: High melting and boiling points; graphite conducts electricity due to delocalized electrons.

4. Metallic Structures:

   • Structure: Lattice of metal cations surrounded by a sea of delocalized electrons.

   • Properties: High melting and boiling points; good conductors of electricity and heat.

Properties of Matter:

• States of Matter: 

  • Solid: Fixed shape and finite volume, particles are arranged closely.

  •  Liquid: Variable shape and finite volume, particles are packed but mobile. 

  • Gas: No fixed shape or volume particles can move freely. 

• Changes of State: 

  • Melting:  Solid to Liquid 

  • Freezing: Liquid to Solid

  • Boiling: Liquid to Gas

  • Condensation: Gas to Liquid 

  • Sublimation: Solid to Gas (skips liquid state of matter)

• Intermolecular Forces: are attractive forces that act between molecules or particles in the solid or liquid states. These forces are weaker than bonding forces.

Indicators of Intermolecular Forces: Melting and Boiling Point

• The stronger the intermolecular forces to be broken, the larger the amount of energy needed to break down. Thus, the higher the melting point for solid to liquid transformation. And, the higher boiling point for liquid to gas transformation.

Quantitative Chemistry

Moles:

• Formula: n=m/M​

  • n = moles.

  • m = mass.

  • M = molar mass.

Avogadro's Constant:

• Value: 6.022×1023

Balanced Equations:

• Purpose: Show conservation of mass and moles in chemical reactions.

Concentration:

• Formula: c=n/V

  • c = concentration.

  • n = moles.

  • V = volume.

Yield and Atom Economy:

• Percentage Yield: Percentage Yield=(actual yield/theoretical yield)×100

• Atom Economy: Atom Economy=(mass of desired product /total mass of reactants)×100

Chemical Changes

Types of Reactions:

1. Exothermic: Release energy (e.g., combustion).

2. Endothermic: Absorb energy (e.g., thermal decomposition).

Reaction Rates:

• Factors Affecting Rate: Temperature, concentration, surface area, catalysts.

• Collision Theory: Particles must collide with sufficient energy and correct orientation for a reaction to occur.

Equilibrium:

• Description: Dynamic equilibrium is reached in reversible reactions when the rate of the forward reaction equals the rate of the backward reaction.

Energy Changes

Energy Profiles:

1. Exothermic Reactions:

   • Products have lower energy than reactants.

2. Endothermic Reactions:

   • Products have higher energy than reactants.

Enthalpy Changes (ΔH):

• Exothermic: ΔH

• Endothermic: ΔH

Bond Energy Calculations:

• Energy Change: ΔH=∑Bond Energies of Reactants−∑Bond Energies of Products

Acids and Bases:

1. pH Scale:

   • Measures acidity or alkalinity.

   • pH < 7: Acidic, pH = 7: Neutral, pH > 7: Alkaline.

2. Neutralization:

   • Acid + Base → Salt + Water.

   • Example: HCl+NaOH→NaCl+Water

Electrolysis:

1. Process:

   • Use of electric current to drive a non-spontaneous chemical reaction.

2. Applications:

   • Electroplating, extraction of metals (e.g., aluminum from bauxite).

Redox Reactions:

1. Oxidation: Loss of electrons.

2. Reduction: Gain of electrons.

3. Redox Reactions: Reactions involving both oxidation and reduction processes.

   • Example: 2Mg+O2→2MgO

Rates of Reaction:

1. Measuring Reaction Rates:

   • Change in concentration of reactants or products over time.

2. Factors Affecting Rates:

   • Temperature, concentration, surface area, presence of a catalyst.

Chemical Analysis:

1. Qualitative Analysis: Identifying the presence of specific ions or compounds.

   • Flame tests for metal ions (e.g., sodium yields a yellow flame).

2. Quantitative Analysis: Determining the quantity of a substance present.

   • Titrations to find concentration of acids or bases.