Chemical Bonds and Melting Points

Diamond Melting Point

• Diamonds melt at 7000°C.
• This is due to the incredibly strong covalent bonds within their structure.

Covalent Bonds in Diamonds

• Diamonds consist of carbon atoms bonded covalently in a network structure.
• Covalent bonds are exceptionally strong.
• Melting or boiling covalent molecules (e.g., sugar, water) involves breaking intermolecular bonds (weak bonds between molecules), not covalent bonds.

Intermolecular vs. Intramolecular Bonds

• Intramolecular: Occur within a molecule (e.g., covalent bonds).
• Intermolecular: Occur between molecules (e.g., weak forces like London dispersion).

Ionic Bonds

• Ionic bonds form between positively and negatively charged ions.
• Sodium chloride (NaCl) melts around 800°C, which is stronger than intermolecular forces but weaker than covalent bonds in diamond.

Metallic Bonds

• Metallic bonds hold metals together via a "sea of electrons" surrounding cations (positive metal ions).
• Strong metallic bonds result in high melting points for metals.

Melting Points and Bond Strength

• Stronger bonds correlate with higher melting points (diamonds, some ionic compounds, metals).
• Weaker intermolecular forces result in lower melting points (sugar, water, various molecular compounds).

Network Covalent Solids

• Diamond, graphite, and silicon dioxide (sand) are examples of network covalent solids.
• Carbon atoms in diamond form a 3D network.
• Graphite's carbon atoms are arranged in layered sheets.
• Sand's SiO₂ molecules are strongly covalently bonded in a network.

Solubility and Solutions

• Ionic and molecular compounds dissolve in solutions.
• Network covalent and metallic solids generally do not.
• Substances dissolve in solutions when there's attraction between solute and solvent.

Polarity of Molecules

• Nonpolar molecules: Atoms are identical or evenly distributed around a central atom (held together by London Dispersion forces).
• Polar molecules: Asymmetrical, causing partial charges (stronger attraction than nonpolar).
• Superpolar Molecules (Hydrogen Bonding): Polar molecules with hydrogen attached to highly electronegative elements (oxygen, nitrogen, fluorine) exhibit strong hydrogen bonding, almost as strong as ionic bonds (e.g., water).

Bond Strength Order

• Ionic bonds are the strongest.
• Hydrogen bonds (in superpolar molecules) are next.
• Polar molecules exhibit weaker attractions.
• Nonpolar molecules have the weakest forces.

Hydrocarbons

• Hydrocarbons (carbon and hydrogen) exist as gases, liquids, or solids at room temperature depending on their size.
• Larger hydrocarbons possess stronger London Dispersion Forces due to greater electron cloud size.

London Dispersion Forces

• Temporary positive and negative charge shifts in nonpolar molecules create weak attractive forces (London Dispersion Forces).

Solubility and Attraction

• Substances dissolve if there's attraction between solute and solvent.
• Polar solutes dissolve in polar solvents; nonpolar solutes dissolve in nonpolar solvents; ionic compounds dissolve in polar solvents.

Wax in Hexane vs. Water

• Wax dissolves in hexane (nonpolar solvent).
• Wax does not dissolve in water (polar solvent) due to lack of attraction between nonpolar wax and polar water.

Ionic Solutes in Water

• Water dissolves ionic compounds by aligning with the ions: oxygen (negative) towards positive ions, and hydrogen (positive) towards negative ions.
• This strong attraction leads to dissolution.

Polar Solutes in Water

• Water dissolves polar molecules similarly, using its partial charges to break apart the solute.

Nonpolar Solutes in Water

• No dissolution as no attraction exists between nonpolar molecules and polar water molecules.

Concentration (Molarity)

• Molarity = Moles of solute / Liters of solution
• Convert grams of solute to moles using molar mass before calculation.

Saturation and Solubility

• Saturated solution: Solution with maximum solute at a given temperature.

• Solubility Chart: Historical method to find solubility values for various substances at different temperatures by looking at a chart.

• Supersaturated solution: Contains more solute than the saturated amount—unstable.

• Unsaturated solution: Less solute than the saturated amount, more solute can be dissolved.

• Temperature generally increases solubility of most ionic compounds (more solute dissolves).

• Solubility of some ionic compounds (e.g., NaCl) is relatively unaffected by temperature.

Calculation of Molarity, Moles, and Liters

• Molarity (M): moles of solute/ liters of solution

• To find any of the unknowns (moles, liters, or molarity) you need two of the other variables.

• Convert grams to moles using molar mass.