Chapter_3_Lecture_Notes_-_CHEM_1113_Broering

Page 1: Introduction to Chemical Bonds

• Chemical Bonds: Attraction forces between atoms which hold them together in a compound.

  • Formed due to attractive forces between protons and electrons.

• Types of Chemical Bonds:

  1. Ionic Bonds: Formed between a metal and a nonmetal.

  2. Covalent Bonds: Formed between two nonmetals.

  3. Metallic Bonds: Formed between two metals.

• Examples:

  • Atoms of Argon (Ar)

  • Molecules of Bromine (Br2)

  • Formula units of Sodium Chloride (NaCl)

Page 2: Ionic Compounds

• Ionic Bonds:

  • Occurs between a metal (cation) and a nonmetal (anion).

  • Involves transfer of electron(s) from the metal to the nonmetal.

  • Compounds formed through ionic bonds are termed ionic compounds.

  • Cation and anion are held together by electrostatic (Coulombic) attractions.

  • In a crystal lattice structure, this attraction is referred to as lattice energy.

• Empirical Formulas:

  • Represent the smallest ratio of atoms in a compound.

  • They may or may not indicate how many of each type of atom are present.

  • Example: Empirical formula for NaCl is NaCl.

Page 3: Molecular Compounds

• Covalent Bonds:

  • Formed by the sharing of one or more pairs of electrons between two atoms.

  • The bond is a balance between attractive and repulsive forces:

    • Attractive forces: protons of atom A attract electrons of atom B.

    • Repulsive forces: electron-electron and proton-proton repulsions.

• Molecular Compounds: Formed from covalently bonded atoms.

• Chemical Formulas: Example of Ethanol: C2H6O (2 Carbon, 6 Hydrogen, and 1 Oxygen).

Page 4: Metallic Bonds

• Metallic Bonds:

  • Occur in a lattice of metal atoms surrounded by a sea of shared electrons.

• Electron Mobility: Metals are excellent conductors of electricity due to free-flowing electrons.

Page 5: Classification of Elements and Compounds

• Elements:

  • Atomic Elements: Exist as single atoms (e.g., Noble gases).

  • Molecular Elements: Exist as diatomic or larger molecules (e.g., H2, O2, P4)

    • Allotropes: Different molecular forms of the same element (e.g., O2 vs O3 for oxygen).

Page 6: Classification of Compounds

• Compounds:

  • Molecular Compounds: Formed by nonmetal atoms through covalent bonding (e.g. CCl4).

  • Ionic Compounds: Composed of cations and anions (e.g. NaCl).

    • May have simple or polyatomic ions such as NO3- (nitrate) or NH4+ (ammonium).

Page 7: Formulas for Ionic Compounds

• Chemical Formulas of Ionic Compounds:

  • Neutral compounds with a net charge of zero.

  • Written as empirical formulas; charges can be balanced using the criss-cross method.

Page 8: Criss-Cross Method for Formulas

• Examples:

  • Aluminum Oxide: Al2O3

  • Barium Fluoride: BaF2

• Steps: Determine the charge on cation and anion, adjust subscripts to balance overall charge.

Page 9: Naming Binary Ionic Compounds

• Rules for Naming:

  1. Name the cation (metal).

  2. Name the anion (nonmetal) by changing the ending to -ide.

• Example Names:

  • MgCl2: magnesium chloride

  • KBr: potassium bromide

Page 10: More Examples of Naming Compounds

• Naming Examples:

  • Lithium and Bromine: LiBr (Lithium bromide)

  • Magnesium and Oxygen: MgO (Magnesium oxide)

Page 11: Transition Metal Ionic Compounds

• Most transition metals can form multiple cations. The name includes a Roman numeral to reflect its charge (e.g., Copper(I) oxide vs Copper(II) oxide).

Page 12: Naming Compounds with Polyatomic Ions

• Polyatomic Ions: Charged groups of multiple atoms bonded by covalent bonds.

• Naming: Name the cation and add the name of the polyatomic ion without changing its ending.

Page 13: Common Polyatomic Ions

• List of Common Polyatomic Ions:

  • Ammonium (NH4+)

  • Hydroxide (OH-)

  • Nitrate (NO3-)

  • Sulfate (SO4 2-) ... (further ions listed).

Page 14: Naming Acids

• Naming rules for binary acids:

  1. Add prefix hydro- to the second element's name.

  2. Replace the second element's ending with -ic acid.

Page 15: Naming Oxyacids

• Change -ate to -ic and -ite to -ous when naming acids derived from polyatomic ions.

• For prefixes like Per- or Hypo-, these carry over to the acid's name.

Page 16: Concept Test

• Determine the Name: For a formula such as H2SO4, know how to name the resulting compound based on its composition and structure.

Page 17: More Naming Practice

• Given Compounds: Practice naming and type of bonds present (ionic or covalent).

Page 18: Summary of Naming/Practice

• Understanding naming conventions for acids and ionic compounds helps in identifying elements and formulas.

Page 19: The Mole: Introduction

• Mole Definition: 1 mole = 6.022 × 10^23 particles.

• Conversion factors relate microscopic measurements to macroscopic scales.

Page 20: Relationship between Mass and Moles

• Chemist's Dozen (Mole): Allows relating mass of substances to number of particles.

• Example showing relationships between grams and moles.

Page 21: Molar Mass

• Molar Mass: The mass in grams of 1 mole of a substance.

• Example: Molar mass of chlorine (Cl) is approximately 35.45 g/mol.

Page 22: Mass-Mole Conversion

• Formula for conversions between mass and moles:

  • Mass (g) to moles: g × (1 mol/molar mass).

Page 23: Concept Test on Moles

• Problems relevant to finding moles of specific scenarios.

Page 24: Moles of Atoms in Compounds

• Relates the number of moles from chemical formulas to moles of components within those formulas.

Page 25: Practice Problems on Mole Calculation

• Solving practice problems of moles using given data.

Page 26: Mole Ratios in Compounds

• Using ratios in compounds for calculating moles different atoms in the compound.

Page 27: Empirical and Molecular Formulas

• Comparison and calculation of empirical formulas from molecular data.

Page 28: Calculation Method for Empirical/Molecular Forms

• Steps for determining chemical formulas using experimental data.

Page 29: Percent Composition Calculation

• Percent Composition: Calculation methodology for percent of an element in a compound based on mass.

Page 30: Practice Problems on Percent Composition

• Example calculations for better understanding of empirical formulas.

Page 31: Empirical vs Molecular Formulas

• Discuss the concepts of empirical and molecular formulas and when they are applicable.

Page 32: Naming and Distinguishing Compounds

• Methods for distinguishing various forms of chemical compounds and their adherent nomenclature practices.

Page 33: Awareness on Empirical Formula Determination

• Consider empirical formulas from mass percentages and result analysis.

Page 34: Review on Empirical Relationships

• Recap knowledge on how empirical data guides molecular formula representations.

Page 35: Empirical Formula Practice Problem

• A summation of practices for finding empirical formulas from chemical compositions.

Page 36: Concept Test on Molecular Formula Assessment

• Testing understanding of molecular formula based on percent compositions.

Page 37: Determining Empirical Formulas through Mass Relationships

• Detailed example of determining empirical formulas based on compound combustion analysis.

Page 38: Model for Finding Molecular Formulas

• Working through examples of determining empirical formula types based on mass data.

Page 39: Application of Combustion Analysis Techniques

• Combustion analysis methodologies for determining empirical formulas in various organic compounds.

Chemical Bonds: Forces holding atoms in compounds, created by attractions between protons and electrons.

Types of Bonds:

• Ionic Bonds: Between metals (cation) and nonmetals (anion), involving electron transfer. Form ionic compounds (e.g., NaCl).

• Covalent Bonds: Sharing of electrons between nonmetals, forming molecular compounds (e.g., C2H6O).

• Metallic Bonds: In metal lattices with shared electrons, allowing conductivity.

Classification:

• Elements:

  • Atomic (single atoms) and Molecular (diatomic or larger).

  • Allotropes represent different molecular forms.

• Compounds: Molecular (covalent bonding, e.g., CCl4) and Ionic (cations/anions, e.g., NaCl), including polyatomic ions.

Naming Compounds:

• Binary Ionic: Name metal, change nonmetal ending to -ide (e.g., MgCl2 to magnesium chloride).

• Polyatomic ions retain names (e.g., Cu(NO3)2 as copper nitrate).

• Acids: Binary acid naming uses hydro- and -ic; oxyacids change -ate to -ic and -ite to -ous.

The Mole: 1 mole = 6.022 × 10^23 particles; relates mass to numbers of particles via molar mass.

Empirical and Molecular Formulas: Understand and determine based on mass relationships and percent composition.

Lattice energy is the energy released when gaseous ions form an ionic solid. It measures the strength of the electrostatic attractions between cations and anions. Higher lattice energy indicates greater stability of the ionic compound.