Regents review chemistry

🧲 Bonding & Intermolecular Forces (IMF) Study Guide

🔗 Types of Chemical Bonds

Ionic Bonds

• Formed between metal + nonmetal

• Electrons are transferred

• High melting/boiling points

• Conduct electricity when dissolved (electrolytes)

• Example: NaCl

Covalent Bonds

• Formed between nonmetals

• Electrons are shared

• Polar: Unequal sharing (ΔEN = 0.4–1.7)

• Nonpolar: Equal sharing (ΔEN < 0.4)

• Example: H₂O (polar), Cl₂ (nonpolar)

Metallic Bonds

• “Sea of mobile electrons” among metal atoms

• Conduct electricity in solid and liquid phases

• Example: Cu, Fe

📊 Electronegativity & Bond Type

• Use Table S (Periodic Table) for electronegativity values

• ΔEN ≥ 1.7 → Ionic

• ΔEN < 1.7 → Covalent

🧪 Molecular Polarity

• A molecule is polar if:
  

  • It has polar bonds

  • It is asymmetrical in shape

• Nonpolar molecules = symmetrical

  • Examples: CH₄, CO₂

💧 Intermolecular Forces (IMF)

Hydrogen Bonding

• Between H and N, O, or F

• Strongest IMF

• Causes high boiling points (e.g., H₂O, NH₃)

Dipole-Dipole

• Between polar molecules

• Medium strength

London Dispersion Forces (LDF)

• Between all molecules

• Only IMF for nonpolar molecules

• Strength increases with molecular size

• Weakest IMF

🌡 Physical Properties & IMF

• Stronger IMF → Higher boiling/melting point

• Ionic compounds > Covalent compounds in melting point

🧊 Solubility Rules

• Like dissolves like

  • Polar dissolves polar

  • Nonpolar dissolves nonpolar

⚡ Electrical Conductivity

• Ionic Compounds:

  • Conduct in aqueous or molten form

  • Do not conduct in solid form

• Covalent Compounds:

  • Usually nonconductors

🧬 Network Solids

• Strong covalent bonds in large lattice

• Very high melting points

• Do not conduct electricity

• Examples: Diamond (C), SiO₂

🧠 Regents Tips

• Use Table S: Electronegativity

• Use Table F: Solubility of ionic compounds

• Draw Lewis dot diagrams to show bonds

• Check molecular shape for polarity (linear, bent, tetrahedral)

🧪 Regents Chemistry Study Guide: Formulas, Moles, and Equations

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Key Concepts

• Mole: A unit representing 6.02 \times 10^{23} particles (Avogadro’s number).

• Molar Mass (g/mol): The mass of 1 mole of a substance; calculated by summing atomic masses from the periodic table.

• Formula Mass: Sum of atomic masses of all atoms in a formula (used for ionic compounds).

• Molecular Mass: Same as formula mass but used for molecular compounds.

• Percent Composition:
  \text{Percent by mass} = \frac{\text{Part mass}}{\text{Total mass}} \times 100

• Empirical Formula: The simplest whole-number ratio of atoms in a compound.

• Molecular Formula: Actual formula; may be a multiple of the empirical formula.

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Key Skills

1. Calculate Formula/Molecular Mass

   • Add up the atomic masses of all atoms in a compound.

2. Determine Moles from Mass
   \text{Moles} = \frac{\text{Given mass (g)}}{\text{Molar mass (g/mol)}}

3. Determine Mass from Moles
   \text{Mass (g)} = \text{Moles} \times \text{Molar Mass}

4. Calculate Percent Composition

   • Find total molar mass.

   • Divide each element’s contribution by the total and multiply by 100.

5. Determine Empirical Formula
   

   • Convert percent to grams (assume 100 g sample).

   • Convert grams to moles.

   • Divide all by the smallest number of moles.

   • Multiply to get whole numbers if necessary.

6. Determine Molecular Formula
   \text{Molecular formula} = (\text{Empirical formula})_n
   

   • Where n = \frac{\text{Molar mass}}{\text{Empirical formula mass}}

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Practice Question Types (with tips)

• Molar mass of compounds

  • Use periodic table values; include units.

• Grams ↔ Moles conversions

  • Always check if you are given grams and need moles, or vice versa.

• Percent composition

  • Double-check each element’s contribution.

• Empirical formula

  • Round carefully and check if ratios make sense.

• Molecular formula

  • Always confirm molar mass is given or can be derived.