URP: Molecules, Compounds & Nomenclature- chapter 3 - Detailed Study Notes
URP: Molecules, Compounds & Nomenclature - Detailed Study Notes
1. Introduction to Nomenclature
- Making questions based on subtitles aids in note-taking. Variants can foster diverse learning techniques among students.
- Closing statements or questions can solidify understanding.
- Reference: Read Section 3-2, pages 59-60 (not covered in class).
- Two bond types:
- Covalent Bonds:
- Definition: Involve the sharing of electrons between atoms.
- Ionic Bonds:
- Definition: Involve the transfer of electrons from one atom to another.
3. Molecular Compounds
- Definition: Compounds containing covalent bonds between non-metal atoms.
- Characteristics:
- Composed of discrete units known as molecules.
- Represented by chemical formulas that indicate:
- Elements present using chemical symbols.
- Relative number of atoms of each element using subscripts.
- Example: Glucose
- Chemical formula: C6H{12}O_6
- Elements present: Carbon (C), Hydrogen (H), Oxygen (O).
- Atom count per molecule: 6C, 12H, 6O.
4. Representing Compounds
- Types of Chemical Formulas:
- Empirical Formula:
- Simplest representation of a formula.
- Indicates atoms present and their relative numbers, with subscripts in the simplest whole number ratio.
- Example: P2O5 is empirical for P4O{10}, P6O{15}, and itself.
- Example: CH2O is empirical for CH2O (formaldehyde), C2H4O2 (acetic acid), C6H{12}O6 (glucose).
- Molecular Formula:
- Shows actual number of atoms present for each element in a molecule.
- Examples:
- CH_2O for formaldehyde (the empirical formula).
- C6H{12}O_6 for glucose (molecular is 6 times the empirical).
- Does not indicate how atoms are connected.
- Structural Formula:
- Displays the order of bonding and types of bonds.
- Example: For ethylene, C2H4 structural representation can be shown with connectivity.
- Definition: Simplified representation written in one line.
- Example: CH3COOH or CH3CO_2H
4.3 Molecular Models
- Ball-and-stick Model:
- Represents atoms as balls and bonds as sticks, showcasing bond lengths and geometries.
- Space-filling Model:
- Shows that atoms in a molecule occupy space, implying actual atomic contact; models are built to scale.
5. Ionic Compounds
- Definition: Compounds featuring ionic bonds (often between metals and non-metals).
- Process:
- Metal donates electrons to form cations.
- Non-metal accepts electrons to form anions.
- Example: Calcium Chloride (CaCl2)
- Calcium donates two electrons to form Ca^{2+}, chlorine accepts one to form Cl^–.
- To maintain neutrality, one Ca^{2+} must pair with two Cl^– ions.
- Crystal structure: Each Ca^{2+} ion is surrounded by six Cl^– ions.
6. Naming Compounds
- Organic Compounds:
- Formed by carbon and hydrogen or carbon with oxygen, nitrogen, etc. (details in CHEM 101).
- Inorganic Compounds:
- Comprise elements excluding carbon and hydrogen.
6.1 Binary Compounds
- Definition: Compounds formed between two elements.
- Binary Ionic Compounds (Type I):
- Metal ion forms one type of cation (Groups 1, 2, Al³⁺, etc.).
- Ionic compounds must maintain electrical neutrality.
- How to Name:
- Order: Cation name, then anion name.
- Example: Na^+ is named Sodium, Cl^– is named Chloride leading to Sodium Chloride (NaCl).
- Example Formulations:
- Sodium bromide (NaBr), Calcium bromide (CaBr₂).
- Note: Transition metals form multiple cations requiring distinction.
- Systematic naming requires Roman numerals to denote cation charge (e.g., Iron (II) for Fe^{2+}).
7. Charges of Ions
- Example Table (Ions with Charges):
- Alkali metals: +1 (e.g., Li^+, Na^+, K^+)
- Alkaline earth metals: +2 (e.g., Be^{2+}, Mg^{2+})
- Transition metals vary (e.g., Fe^{2+}, Cu^{2+}).
8. Naming Binary Ionic Compounds (Type II)
- Definition: Metal ion forms multiple cations (e.g., Fe^{2+}, Fe^{3+}).
- Naming Rules:
- Unmodified metal name followed by cation charge in Roman numerals (e.g., Fe^{3+} = Iron(III)).
9. Polyatomic Ions
- Definition: Ions made up of more than one atom. Examples include:
- Ammonium ion: NH_4^+
- Carbonate ion: CO_3^{2–}
- Notable Polyatomic Ions with Charges:
- Hypochlorite: ClO^{–}
- Acetate: CH_3COO^{–}
10. Naming Polyatomic Ions and Oxyanions
- Naming convention for oxoacids:
- Anions without oxygen change from -ate to -ic.
- Example: Sulfate (SO₄²⁻) to Sulfuric acid (H₂SO₄).
11. Hydrates
- Definition: Formula units with water molecules incorporated in their solid structures.
- Example:
- CoCl2•6H2O (Cobalt (II) chloride hexahydrate).
- Properties:
- Losing water makes compounds anhydrous.
12. Binary Covalent Compounds (Type III)
- Formed from two non-metals.
- Naming Rules:
- Use prefixes to denote the number of atoms present (e.g., mono, di, tri).
- Prefix 'mono' is omitted for the first element.
13. Acid Nomenclature
- Binary Acids:
- Form: Hydro + (Non-metal root) + -ic + acid.
- Example: HCl = Hydrochloric acid.
- Oxoacids:
- Formed from oxyanions and yield H⁺ in solutions.
- Example:
- H2SO4 is sulfuric acid from sulfate.
14. Molecular Mass and Calculations for Compounds
- Molecular Mass: Sum of atomic masses of atoms in a formula or molecular unit.
- Molar Mass: Mass of one mole of entities (atoms, ions, molecules).
- Formulas used to determine mass percent of elements in compounds based on their chemical formulas and molar masses.
- Establishing formulas from experimentally determined percent compositions to find molecular and empirical formulas.
17. Combustion Analysis
- Technique to find amounts of C and H in combustible organic compounds by burning and measuring resultant products, which can convert mass data into empirical formula estimates.
18. Multiconcept Examples
- Example Problems: Provided for practice in applying nomenclature, mass calculations, and empirical formula determination in diverse scenarios (details in examples).
19. Learning Objectives
- Students should learn to:
- Understand and apply nomenclature rules, bond types, and representation of compounds.
- Calculate molecular masses and percent compositions.
- Determine empirical and molecular formulas based on analytical data.