Jason Cameron

Chapter 1 study guide

7 min read

SCH3U Chapter 1: Atomic Structure and Chemical Nomenclature

1. Chemical Nomenclature

Chemical nomenclature is the system of naming chemical compounds. Understanding this system is crucial for communicating effectively in chemistry.

1.1 Ionic Compounds

Ionic compounds are formed by the transfer of electrons between metals and non-metals.

Univalent Ionic Compounds

  • Composed of ions with only one charge
  • Formula: Metal + Non-metal
  • Naming: [Metal name] [Non-metal root + -ide]
  • Example: NaCl (Sodium chloride)

Multivalent Ionic Compounds

  • Involve metals that can form ions with different charges
  • Formula: Metal with multiple possible charges + Non-metal
  • Naming: [Metal name] [Roman numeral for charge] [Non-metal root + -ide]
  • Example: Fe₂O₃ (Iron(III) oxide)

Polyatomic Ionic Compounds

  • Involve polyatomic ions (groups of atoms that act as a single unit)
  • Formula: Metal/Ammonium + Polyatomic ion
  • Naming: [Metal/Ammonium name] [Polyatomic ion name]
  • Example: NH₄NO₃ (Ammonium nitrate)

1.2 Acids

Acids are compounds that release hydrogen ions (H⁺) in solution.

Binary Acids

  • Contain hydrogen and one other element
  • Formula: H + Non-metal
  • Naming: Hydro- [Non-metal root] -ic acid
  • Example: HCl (Hydrochloric acid)

Oxyacids

  • Contain hydrogen, oxygen, and another element
  • Formula: H + Polyatomic ion containing oxygen
  • Naming: [Polyatomic ion root] -ic/-ous acid
  • Example: H₂SO₄ (Sulfuric acid)

1.3 Acid Salts

Acid salts are formed when only some of the hydrogens in an acid are replaced by a metal.

  • Formula: Metal + Remaining hydrogen + Anion
  • Naming: [Metal name] hydrogen [Anion name]
  • Example: NaHCO₃ (Sodium hydrogen carbonate)

1.4 Radicals

Radicals are atoms or groups of atoms with an unpaired electron.

  • Formula: Group of atoms behaving as a single unit
  • Naming: Based on composition
  • Example: OH⁻ (Hydroxide radical)

1.5 Hydrates

Hydrates are compounds that contain water molecules in their crystal structure.

  • Formula: Ionic compound · nH₂O
  • Naming: [Ionic compound name] [Greek prefix]-hydrate
  • Example: CuSO₄·5H₂O (Copper(II) sulfate pentahydrate)

1.6 Molecular Compounds

Molecular compounds are formed by the sharing of electrons between non-metals.

  • Formula: Non-metal + Non-metal
  • Naming: [Greek prefix] [First element root] [Greek prefix] [Second element root + -ide]
  • Example: CO₂ (Carbon dioxide)

1.7 Deriving Chemical Formulas

To derive chemical formulas:

  1. Determine the charges of the ions involved
  2. Use the criss-cross method to balance charges
  3. Simplify the formula if possible

Example: Aluminum oxide

  • Al³⁺ and O²⁻
  • (Al₃)(O₂)
  • Simplified to Al₂O₃

2. Atomic Structure

Understanding the structure of atoms is fundamental to chemistry.

2.1 Basic Structure of the Atom

  • Protons: Positively charged particles in the nucleus
  • Neutrons: Neutral particles in the nucleus
  • Electrons: Negatively charged particles orbiting the nucleus
  • Nucleus: Central part of the atom containing protons and neutrons
  • Electron shells: Energy levels where electrons are found

Electron Configuration:

  • First shell (K): 2 electrons
  • Second shell (L): 8 electrons
  • Third shell (M): 18 electrons
  • Fourth shell (N): 32 electrons

2.2 Development of Atomic Theory

  1. John Dalton (1766-1844)

    • Proposed that all matter is composed of indivisible particles called atoms

  2. J.J. Thomson (1856-1940)

    • Discovered the electron
    • Proposed the “plum pudding” model of the atom

  3. Ernest Rutherford (1871-1937)

    • Discovered the nucleus through his gold foil experiment
    • Proposed the nuclear model of the atom

  4. Niels Bohr (1885-1962)

    • Proposed that electrons orbit the nucleus in fixed energy levels
    • Developed the Bohr model of the atom

  5. Erwin Schrödinger (1887-1961) and Werner Heisenberg (1901-1976)

    • Developed quantum mechanical model of the atom

2.3 Atomic Number and Mass Number

  • Atomic Number (Z): Number of protons in an atom’s nucleus
  • Mass Number (A): Total number of protons and neutrons in an atom’s nucleus

Relationship: A = Z + N (where N is the number of neutrons)

2.4 Atomic Diagrams

Bohr Diagram

  • Shows electrons in circular orbits around the nucleus
  • Each orbit represents an energy level

Bohr-Rutherford Diagram

  • Similar to Bohr diagram, but includes protons and neutrons in the nucleus

Lewis Diagram

  • Represents valence electrons as dots around the element symbol

2.5 Emission Spectra

  • Caused by electrons dropping from higher to lower energy levels
  • Each element has a unique emission spectrum due to its unique electron configuration
  • Used to identify elements in unknown samples

3. Ions and the Octet Rule

3.1 The Octet Rule

  • Atoms tend to gain, lose, or share electrons to achieve a stable electron configuration of 8 valence electrons (like noble gases)
  • This rule explains many aspects of chemical bonding

3.2 Formation of Ions

  • Process: Atoms gain or lose electrons to achieve a stable electron configuration
  • Metals: Form cations (positive ions) by losing electrons
  • Non-metals: Form anions (negative ions) by gaining electrons

Determining Ion Charge

  1. Look at the group number on the periodic table
  2. For main group elements:
    • Metals: lose all valence electrons
    • Non-metals: gain electrons to complete octet

Multivalent Ions

  • Ions of the same element with different charges
  • Example: Iron (Fe²⁺ and Fe³⁺)

Polyatomic Ions

  • Ions composed of multiple atoms that behave as a single unit
  • Example: Ammonium (NH₄⁺)

4. Isotopes and Atomic Mass

4.1 Isotopes

  • Atoms of the same element with the same number of protons but different numbers of neutrons
  • Have the same atomic number but different mass numbers

4.2 Isotopic Abundance

  • The relative amount of each isotope of an element found in nature
  • Expressed as a percentage or fraction

4.3 Isotope Notation

Standard format: ₍ᵧ₎ˣElement

  • x = mass number (A)
  • y = atomic number (Z)
    Example: ₍₁₁₎²³Na (Sodium-23)

4.4 Average Atomic Mass Calculations

Average Atomic Mass = Σ (Isotope Mass × Relative Abundance)

Example:
For chlorine with two isotopes:

  • Cl-35 (75.77%) and Cl-37 (24.23%)
    Average Atomic Mass = (34.97 × 0.7577) + (36.97 × 0.2423) = 35.45 u

Calculating Isotopic Abundance

Given the average atomic mass and individual isotope masses, set up an equation and solve for the unknown abundance.

5. The Periodic Table and Periodic Law

5.1 Groups and Periods

  • Groups: Vertical columns (similar chemical properties)
  • Periods: Horizontal rows (increasing atomic number)

5.2 Chemical Families and Locations

  1. Alkali Metals (Group 1)
  2. Alkaline Earth Metals (Group 2)
  3. Halogens (Group 17)
  4. Noble Gases (Group 18)
  5. Transition Metals (Groups 3-12)
  6. Lanthanides and Actinides (f-block elements)

5.3 Element Categories

  • Metals: Left and center of the periodic table
    • Good conductors of heat and electricity
    • Malleable and ductile
    • Usually solid at room temperature (except mercury)
  • Non-metals: Right side of the periodic table
    • Poor conductors of heat and electricity
    • Brittle in solid form
    • Can be solid, liquid, or gas at room temperature
  • Metalloids: Along the stair-step line between metals and non-metals
    • Have properties of both metals and non-metals
  • Most reactive metals: Group 1 (Alkali Metals)
  • Most reactive non-metals: Group 17 (Halogens)
  • Least reactive: Group 18 (Noble Gases)

Reactivity is related to electron configuration:

  • Metals: easier to lose electrons = more reactive
  • Non-metals: easier to gain electrons = more reactive
  • Noble gases: full outer shell = least reactive

6.1 Atomic Radius

  • Definition: Half the distance between the nuclei of two adjacent atoms of the same element in a molecule
  • Trend:
    • Decreases from left to right across a period
    • Increases from top to bottom down a group
  • Reason:
    • Across a period: Increased nuclear charge pulls electrons closer
    • Down a group: New electron shells increase size

6.2 Ionization Energy

  • Definition: Energy required to remove the most loosely bound electron from an atom in its gaseous state
  • Trend:
    • Increases from left to right across a period
    • Decreases from top to bottom down a group
  • Reason:
    • Across a period: Higher nuclear charge and smaller atomic radius make it harder to remove electrons
    • Down a group: Greater distance from nucleus makes it easier to remove electrons

6.3 Electron Affinity

  • Definition: Energy change when a neutral atom in its gaseous state accepts an electron
  • Trend:
    • Generally increases from left to right across a period
    • Decreases from top to bottom down a group
  • Reason: Similar to ionization energy, related to nuclear charge and atomic size

6.4 Electronegativity

  • Definition: Ability of an atom to attract shared electrons in a chemical bond
  • Trend:
    • Increases from left to right across a period
    • Decreases from top to bottom down a group
  • Reason:
    • Across a period: Higher nuclear charge and smaller atomic radius increase electron-attracting power
    • Down a group: Greater distance from nucleus decreases electron-attracting power

Cheat sheet

Rows - Period

> 1.7 is ionic

Compound is non polar if:
Contains Diatomic element (HofBrinCL)
Nobel gasses

Valence Shell Electron Pair Repulsion
LP-LP > LP- BP > BP-BP

  • Refer to central atom as “A”
  • Attached atoms are labeled “X”
  • Lone pairs on central atom are labeled “E”

Intramolecular forces: Bonds/forces WITHIN the molecule
Intermolecular forces: Non-Bonding forces BETWEEN molecules

Graph View