Lesson 2: Atomic Theory & The Mole Concept 

Overview of Atomic Theory

Atomic theory explains that all matter is composed of atoms, the smallest units retaining the chemical properties of an element. Atoms are made up of three fundamental subatomic particles:

• Protons: Positively charged particles located in the nucleus.

• Neutrons: Neutral particles also found in the nucleus.

• Electrons: Negatively charged particles that orbit the nucleus in specific energy levels or orbitals.

Key Concepts

The arrangement of electrons in energy levels determines an element’s chemical properties and reactivity. For instance:

• Elements in Group 1 of the periodic table (e.g., sodium, Na) have one electron in their outermost energy level, making them highly reactive and prone to forming cations by losing this electron.

Example:

A helium atom (He) consists of:

• 2 protons and 2 neutrons in the nucleus.

• 2 electrons orbiting the nucleus in a single energy level.

Isotopes

Isotopes are variations of the same element that differ in the number of neutrons within their nuclei. While their atomic masses differ, isotopes of the same element typically exhibit similar chemical properties due to having the same number of protons and electrons.

Examples:

• Carbon-12: 6 protons and 6 neutrons (stable).

• Carbon-14: 6 protons and 8 neutrons (radioactive, commonly used in radiocarbon dating).

Applications of Isotopes:

• Nuclear Medicine: Isotopes like iodine-131 are used for diagnosing and treating thyroid disorders.

• Archaeology: Carbon-14 helps in dating ancient artifacts.

Ions

Ions are charged particles that form when atoms gain or lose electrons:

• Cations: Positively charged ions formed by losing electrons (e.g., Na⁺).

• Anions: Negatively charged ions formed by gaining electrons (e.g., Cl⁻).

Example:

When sodium (Na) reacts with chlorine (Cl):

• Sodium loses one electron, forming a Na⁺ ion.

• Chlorine gains the electron, forming a Cl⁻ ion.

• The oppositely charged ions combine to form sodium chloride (NaCl).

Chemical Formulas and Names of Compounds

Chemical formulas represent the composition of compounds, showing the type and number of atoms in a molecule.

Examples:

• H₂O: Two hydrogen atoms bonded to one oxygen atom (water).

• CO₂: One carbon atom bonded to two oxygen atoms (carbon dioxide).

• NH₃: One nitrogen atom bonded to three hydrogen atoms (ammonia).

• NaCl: One sodium atom bonded to one chlorine atom (table salt).

• C₆H₁₂O₆: Six carbon atoms, twelve hydrogen atoms, and six oxygen atoms (glucose).

Example Applications:

• Sulfuric Acid (H₂SO₄): The formula reveals that it contains:

  • 2 hydrogen atoms.

  • 1 sulfur atom.

  • 4 oxygen atoms.

• Calcium Phosphate (Ca₃(PO₄)₂): The formula indicates:

  • 3 calcium atoms.

  • 2 phosphate groups.

  • A total of 8 oxygen atoms.

Nomenclature

Chemical nomenclature ensures clarity and uniformity in naming compounds.

Ionic Compounds

• Name the cation first, followed by the anion.

Examples:

• KBr → Potassium bromide.

• MgO → Magnesium oxide.

• Ca(NO₃)₂ → Calcium nitrate.

• Fe₂O₃ → Iron(III) oxide (Roman numerals indicate the oxidation state of the metal).

Covalent Compounds

• Prefixes indicate the number of each atom in the compound:

Prefixes:

• Mono-

• Di-

• Tri-

• Tetra-

• Penta-

• Hexa-

• Hepta-

• Octa-

• Nona-

• Deca-

Examples:

• CO → Carbon monoxide.

• N₂O₅ → Dinitrogen pentoxide.

• PCl₃ → Phosphorus trichloride.

• SF₆ → Sulfur hexafluoride.

Acids

• Acid names depend on the type of anion present:

  • Anion ends in -ide: Add "hydro-" as a prefix and "-ic acid" as a suffix.

    • HCl → Hydrochloric acid.

  • Anion ends in -ate: Replace "-ate" with "-ic acid."

    • H₂SO₄ → Sulfuric acid.

  • Anion ends in -ite: Replace "-ite" with "-ous acid."

    • H₂SO₃ → Sulfurous acid.

Examples:

• HNO₃ → Nitric acid.

• H₂CO₃ → Carbonic acid.

• HClO₂ → Chlorous acid.

• HClO₄ → Perchloric acid.

The Mole Concept

The mole concept is a fundamental principle in chemistry, bridging the gap between the atomic scale and macroscopic measurements. It allows chemists to count and measure substances in a practical and standardized manner.

Avogadro’s Number

Avogadro's number, 6.022×10²³, represents the number of particles (atoms, molecules, or ions) in one mole of a substance. This constant is crucial for performing calculations involving large quantities of particles.

Key Points:

• One mole of atoms: Contains 6.022×10²³ atoms.

• One mole of molecules: Contains 6.022×10²³ molecules.

• One mole of ions: Contains 6.022×10²³ ions.

Examples:

• Oxygen gas (O₂):

  • One mole contains 6.022×10²³ O₂ molecules, equivalent to 1.204×10²⁴ oxygen atoms (since each O₂ molecule has two oxygen atoms).

• Sodium (Na):

  • One mole contains 6.022×10²³ sodium atoms.

• Glucose (C₆H₁₂O₆):

  • Two moles of glucose would contain 1.204×10²⁴ molecules, which is equivalent to 2.90×10²⁵ total atoms (24 atoms per glucose molecule).

Molar Mass

The molar mass is the mass of one mole of a substance, expressed in grams per mole (g/mol). It is numerically equivalent to the relative atomic or molecular mass.

Steps to Calculate Molar Mass:

1. Identify the chemical formula.

2. Use the periodic table to find the atomic masses of the elements.

3. Multiply each atomic mass by the number of atoms of that element in the formula.

4. Sum these values to find the total molar mass.

Examples:

• Sodium chloride (NaCl):

  • Na: 23.00 g/mol, Cl: 35.44 g/mol

  • Molar mass: 23.00 + 35.44 = 58.44 g/mol

• Water (H₂O):

  • H: 1.008 g/mol × 2 = 2.016 g/mol, O: 16.00 g/mol

  • Molar mass: 2.016 + 16.00 = 18.016 g/mol

• Calcium carbonate (CaCO₃):

  • Ca: 40.08 g/mol, C: 12.01 g/mol, O: 16.00 g/mol × 3 = 48.00 g/mol

  • Molar mass: 40.08 + 12.01 + 48.00 = 100.09 g/mol

Applications of the Mole Concept

1. Mass to Moles Conversion

Formula:

Moles=Mass of substance Molar mass\text{Moles} = \frac{\text{Mass of substance}}{\text{Molar mass}}

Example:

For 25 g of NaCl:

Moles=2558.44≈0.428 mol\text{Moles} = \frac{25}{58.44} \approx 0.428 \, \text{mol}

2. Moles to Number of Particles Conversion

Formula:

Particles=Moles×6.022×1023\text{Particles} = \text{Moles} \times 6.022 \times 10^{23}

Example:

For 0.5 mol of water:

Particles=0.5×6.022×1023=3.011×1023 molecules\text{Particles} = 0.5 \times 6.022 \times 10^{23} = 3.011 \times 10^{23} \, \text{molecules}

3. Volume of Gases at STP (Standard Temperature and Pressure)

At STP, one mole of gas occupies 22.4 L.

Example:

For 2 moles of oxygen gas:

Volume=2×22.4=44.8 L\text{Volume} = 2 \times 22.4 = 44.8 \, \text{L}

Percent Composition

Percent composition is the percentage by mass of each element in a compound.

Example Calculation for CO₂:

• Determine molar mass:

  • C: 12.01 g/mol, O: 16 g/mol × 2 = 32 g/mol

  • Total molar mass: 12.01 + 32 = 44.01 g/mol

• Calculate percent composition:

  • Carbon: 12.0144.01×100≈27.29%\frac{12.01}{44.01} \times 100 \approx 27.29\%

  • Oxygen: 3244.01×100≈72.71%\frac{32}{44.01} \times 100 \approx 72.71\%

Chemical Formulas and Reactions

Chemical formulas represent the transformation of reactants into products in a chemical reaction, following the law of conservation of mass.

Example Reactions:

• Formation of Water:

  2H2+O2→2H2O2H_2 + O_2 \rightarrow 2H_2O

  (Two molecules of hydrogen react with one molecule of oxygen to form two molecules of water.)

• Combustion of Methane:

  CH4+2O2→CO2+2H2OCH_4 + 2O_2 \rightarrow CO_2 + 2H_2O

  (Methane reacts with oxygen to produce carbon dioxide and water.)

• Neutralization Reaction:

  HCl+NaOH→NaCl+H2OHCl + NaOH \rightarrow NaCl + H_2O

  (Hydrochloric acid reacts with sodium hydroxide to form sodium chloride and water.)

• Photosynthesis:

  6CO2+6H2O→C6H12O6+6O26CO_2 + 6H_2O \rightarrow C_6H_{12}O_6 + 6O_2

  (Carbon dioxide and water are converted into glucose and oxygen in the presence of light.)

• Decomposition Reaction:

  CaCO3→CaO+CO2CaCO_3 \rightarrow CaO + CO_2

  (Calcium carbonate decomposes into calcium oxide and carbon dioxide.)

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