CHE130 L6

Important Reminders

  • Review Practice Problems from today

  • Read sections 3.2-3.3

  • Homework Assignments:

    • HW 2 due Monday, September 22

    • HW 3 due Thursday, September 25 (not yet posted)

    • Exam 1 scheduled for Wednesday, October 1

Element Data Table

Selected Elements with Atomic Masses

  • H (Hydrogen): 1.008

  • Li (Lithium): 6.941

  • Be (Beryllium): 9.012182

  • Na (Sodium): 22.989770

  • Mg (Magnesium): 24.3050

  • He (Helium): 4.002602

  • Al (Aluminum): 26.981538

  • Si (Silicon): 28.0855

  • P (Phosphorus): 30.973761

  • S (Sulfur): 32.066

  • K (Potassium): 39.0987

  • Ca (Calcium): 40.078

  • Sc (Scandium): 44.95591

  • Ti (Titanium): 47.867

  • V (Vanadium): 50.9415

  • Cr (Chromium): 51.9961

  • Mn (Manganese): 54.938044

  • Fe (Iron): 55.845

  • Co (Cobalt): 58.933194

  • Ni (Nickel): 58.933194

  • Cu (Copper): 63.546

  • Zn (Zinc): 65.38

  • Additional elements include Rb, Sr, Y, Zr, Nb, Mo, etc.

Examples of Ionic Compounds Formation

  • From Calcium and Chlorine:

    • Formation of CaCl_2 (calcium chloride)

  • From Ammonium and Chlorate:

    • Formation of NH4ClO3 (ammonium chlorate)

  • From Aluminum and Sulfate:

    • Formation of Al2(SO4)_3

Example Compounds

  • Ionic compounds: Ca^{2+} and Cl^- form CaCl_2

  • Aluminum carbide (AlC_3):

  • Calcium sulfate (CaSO_4)

  • Mercury compounds:

    • Hg_2O (mercury(I) oxide)

    • HgO (mercury(II) oxide)

Writing Formulas of Ionic Compounds from Names

  • Iron (III) Sulfide: Fe2S3

  • Magnesium Sulfate Heptahydrate: MgSO4 \cdot 7H2O

  • Gallium (III) Nitride: GaN

  • Iron (III) Chloride Dihydrate: FeCl3 \cdot 2H2O

  • Sodium Phosphate: Na3PO4

Naming Binary Molecular Compounds

  • Naming Rules:

    • Names use a different set of rules compared to ionic compounds.

    • Explicit identification of atom ratios required.

  • Examples:

    • SO2 (sulfur dioxide), SO3 (sulfur trioxide)

    • Prefixes denote number of atoms (e.g. mono-, di-, tri-, etc.).

    • Negative suffix ‘-ide’ used for nonmetallic elements in the compound.

    • When one atom of the first element is present, 'mono-' prefix typically is omitted.

Acids

  • General Definition: Compounds that release hydrogen ions (H^+) when dissolved in water.

  • Aqueous State: Denoted with (aq)

  • Naming Binary Acids:

    1. Change “hydrogen” to hydro–.

    2. Modify other nonmetallic element’s name by adding the suffix “–ic”.

    3. Add the word “acid”.

  • Examples:

    • HF(g) → Hydrogen fluoride; HF(aq) → Hydrofluoric acid

    • HCl(g) → Hydrogen chloride; HCl(aq) → Hydrochloric acid

  • Naming Oxyacids: Include hydrogen, oxygen, and at least one other element.

    • Modify according to root names of related anions and add “acid”.

Formula Mass and the Mole Concept

  • Definition of Formula Mass: The sum of average atomic masses of all atoms in a substance’s formula.

  • Ionic vs. Covalent Compounds:

    • Ionic compounds consist of cations and anions; molecular mass not applicable to ionic compounds.

    • Not referred to as molecules.

The Mole

  • Definition: An amount unit representing the number of entities (atoms, molecules, or ions) in a substance.

  • Defined as the quantity containing the same number of discrete entities as atoms in 12 g of carbon-12:
    N_A = 6.02214179 \times 10^{23}

  • Molar Mass: The mass in grams of one mole of a substance; same numerically as atomic/formula mass in amu.

Percent Composition

  • Definition: The percentage by mass of each element in a compound.

  • Example: For a compound with 10.0 g sample containing 2.5 g H and 7.5 g C, the calculations yield the percent composition of each element.

Empirical and Molecular Formulas

  • Empirical Formula: Shows the simplest whole-number ratio of elements in a compound.

  • Determining Formulas:

    1. Measure masses of elements, derive moles, divide by the smallest molar amount to yield the empirical formula.

  • Example: If a compound is made from 1.71 g C and 0.287 g H, the method described will determine the empirical formula.

Problems for Practice

  • Calculate percent compositions, empirical formulas, and handle complex questions like chlorate ions in a specific mass of compound.

  • Determine identity and molar mass for compounds with given percent compositions and other constraints.

Conclusion of Lecture 6

- Review all definitions, calculations of molar masses, molecular and empirical formulas, and practice multiple problem sets provided throughout the notes.

These notes are formatted to provide a thorough understanding of Chemistry lecture content, building from foundational topics such as atomic mass and the mole to more complex ideas such as compound naming and empirical formula determination.