CHEM11 Unit 9 Solutions: Conductivity

Electrolyte (4)

• A solution that conducts electricity well (has ions)

• Ionic compounds are made when atoms transfer electrons to achieve a full outer shell, creating cations and anions, held together by strong electrostatic forces (they are made of ions)

• When breaking apart these compounds, ions are produced

• A solution will have high electrical conductivity if it has a large number of ions floating around

  • The higher the concentration of ions, the higher the conductivity (i.e. better electrolyte)

Nonelectrolyte

A solution that does not conduct electricity (does not have ions)

• Covalent bonds do not involve ions, but the sharing of electrons, which means there are no charged particles (ions) created during their formation

  • Except certain acids/bases that are covalent

Conductivity of Pure H2O

• Water by itself will dissociate into very small amounts of H+ and OH- ions.

  • Self-ionization at a very small scale (extremely small concentration)

• Therefore pure water has no conductivity

Examples of Electrolytes (4 main groups)

• Acids beginning with H

  • HCl, HBr, HI, HF

• Bases with OH

  • LiOH, NaOH, KOH

• Acetic acids

  • CH3COOH, C6H5COOH, CHOOH

• Ammonia

  • NH3

Conductivity of Ionic Solids

Solids = non-electrolyte

• No conductivity

• Metals have a sea of electrons but are not electrolytes

• The ions in a solid are locked in a rigid lattice structure, not able to move freely, meaning no flow of charge to conduct electricity

Conductivity of Ionic Liquids

Liquids = electrolyte

• Ions are free to move

• These free-moving ions can carry charge, making melted ionic compounds good conductors

Conductivity of Ionic Gases

Gases = nonelectrolyte

• Rarely exist in a gaseous phase because of their high melting and boiling points

Conductivity of Aqueous Ionic Compounds

Aqueous = electrolyte

• Many ionic compounds dissolve in water, dissociating into their ions

• These ions move freely in the solution, allowing the flow of charge and good conductivity

• Strong electrolytes = fully dissociate

• Weak electrolytes = partially dissociate

Conductivity of Covalent Compounds

Non-electrolytes

• Poor conductors due to their lack of free-moving ions/electrons

Concentrations of Solutions Affects Conductivity

When the concentration increases, the conductivity increases because more ions per unit of volume

• 0.01M NaCl = low conductivity

• 0.5M NaCl = medium conductivity

• 2.0M NaCl = higher conductivity

Formula (molecular equation)

• Normal balanced equation with states

Complete ionic equation

• Breaks up any aqueous species into their ions

• Includes all ions

Net Ionic Equation

• Cancels out species that appear on both sides

Spectator ions

• Ions that do not take part in the reaction (species that are cancelled out)

Solubility

The maximum amount of solute that can be dissolve in a solvent (usually in g/L)

• Solubility changes with temperature

Saturated/Saturation

The point at which no more solute can dissolve in a solvent

• The solubility of NaCl is around 360g/L at 25 degrees. Therefore, you can dissolve around 360 grams of NaCl in water before it becomes saturated (the water won’t dissolve anymore NaCl)

Solution

A homogenous mixture (particles that are evenly mixed)

Why does solubility vary with temperature?

• The Kinetic Molecular Theory → a higher temperature = higher kinetic energy (faster moving particles)

• When no more solute will dissolve, the solution is saturated (reached saturation point)

Miscibility

Miscible = liquids can mix with each other

Immiscible = liquids cannot mix with each other

• Polar liquids will be miscible with another polar liquid

  • Water and alcohol will be mixed in any proportions

• Non-polar liquids will be miscible with another non-polar liquid

  • Salad oil and motor oil will not mix

• Non-polar liquid will not be miscible with a polar liquid (water and oil do not mix)

  • “Like dissolves like”

What happens when an ionic solid is placed in water?

• Ionic solid will break apart/dissociate into their ions

  • The partially positive end of a water molecule (H) will be attracted to the negative ion in an ionic solid

  • The partially negative end of a water molecule will be attracted to the positive ion of an ionic solid

  • These attractions disrupt the electrostatic forces in the lattice

Hydrated

• When surrounded by attached water molecules, the ion is said to be hydrated

Solvation

• The process where the ionic solid is broken up in water

  • Process of dissolving

• Solvation can also happen to polar molecules instead of ions in water

• Solvation does not occur between polar and non-polar substances

• If both solvent and solute are non-polar, solvation may occur through London Dispersion Forces

1. Non-polar solvents can only dissolve/mix with non-polar compounds

2. Polar solvents can dissolve/mix with polar or ionic compounds

Steps for making a solution from solid compounds:

1. Using a balance, obtain the required mas of solute in a beaker

2. DIssolve the solid in distilled water

3. Quantitatively transfer the solute and rinse all equipment into a clean volumetric flask

4. Add distilled water up to the calibration line on the flask. An eye dropper may be used.

5. Stopper the volumetric flask and mix the contents thoroughly.

Standardized Solution

A solution whose concentration is known

Equivalence Point

The point in the titration where moles H+ = moles OH-

Endpoint

The point in the titration where the indicator changes colour