SCH3U Grade 11 Chemistry Solutions and Solubility Test

SCH3U1 Grade 11 Chemistry Unit 3 Solutions and Solubility Unit Test Study Notes

Effect of Temperature on Solubility
– every unique pure substance has its own unique solubility based on the types of bond present
– units used to describe/measure solubility is: mass of solute/100mL of solvent

Solubility of Solid
– trend: solubility of solids increase as temperature increases
– energy is required to break apart bonds of solids when dissolved in water
– as temperature increases, there is more energy to break these bonds

SATP/STP: standard ambient temperature and pressure

S = soluble
SS = slightly soluble
I = insoluble

Solubility of Gases
– trend: solubility of gases decrease as temperature increases
– with more energy, the gas particles escape the solution

Solubility of Liquids
– trend: solubility of liquids is not affected by the temperature
– solute: liquid with less amount
– solvent: liquid with greater amount

Water
– universal solvent
– small size, highly polar nature, and the capacity to form hydrogen bonds makes water very successful at dissolving solutes
– water has a permanent dipole
– the negative end is attracted to the positive end, causing a special type of attraction called hydrogen bonding

Hydrogen Bonding
– any substance containing hydrogen and oxygen/fluorine/nitrogen
– doesn’t create an actual bond, uses strong intermolecular forces to create a force of attraction
– hydrogen bonded compounds are likely to dissolve in water

Properties
– water is held together by covalent bonds which re stronger than hydrogen bonds
– however, the hydrogen bond is stronger than regular dipole-dipole attractions
– this results in higher boiling points because more energy is required to break apart these bonds
– hydrogen bonds also result in higher surface tension

Volume/Volume Concentration (V/V) %
– volume of solute(mL)/volume of solution(mL) X 100%

Mass/Volume Concentration (m/V) %
– mass of solute(g)/mass of solution(mL) X 100%

Mass/Mass concentration (m/m) %
– mass of solute(g)/mass of solution(g) X 100%

Very Low Concentrations
– parts per million (ppm) = mass of solute/mass of solution X 106
– parts per billion (ppb) = mass of solute/mass of solution X 109

Molarity
– molar [ ]
– moles of solute/1L of solution

Dilutions
– reducing concentration of a solute by adding additional solution to the mixture
– standard/stock solution: one where the [ ] is known
– c1V1 = c2V2
– c1 is the initial [ ]
– c2 is the final [ ]
– V1 is the initial volume
– V2 is the final volume

Double Displacement Reactions
– 2 possible outcomes
– Compounds remain as ions and no reaction occurs (NR)
– New compounds created that consist of 2 of the following: solid precipitate, gas, or water

Net Ionic Equations
– an ionic compound dissociates in water and is broken up into its constituent ions
– the above occurs before a double displacement reaction happens
– net ionic equation only contains the new product and the constituents that produce this compound
– spectator ions are any ions not involved in the creation of the new product

Acids
– sour taste
– no texture
– conducts electricity in an aqueous solution
– pH less than 7
– turns litmus paper red
– phenolphthalein is colourless
– Acid + Metal  H2(g)
– Acid + Carbonate  CO2(g) + H2O(l)

Bases
– bitter taste
– slippery texture
– conducts electricity in an aqueous solution
– pH greater than 7
– turns litmus paper blue
– turns phenolphthalein pink

Arrhenius Theory
– an acid is any substance that will ionize in water to produce H ions
– a base is a substance that will dissociate in water to produce OH ions
– H ions cannot exist alone, and thus exist attached to H2O, creating hydronium: H3O
– Only valid for reactions in water

Bronsted-Lowry Theory
– acids are substances that have an H ion removed
– conjugate base is paired with the acid, and becomes the new base
– bases are substances that have an H ion added
– conjugate acid is paired with the base and becomes the new acid

Strong Acids
– will completely dissociate

Weak Acids
– will only have some of the solution dissociated
– indicated in a chemical equation by a double arrow

Monoprotic Acid
– can only give up 1 H ion

Diprotic Acid
– can give up 2 H ions (H2SO4, H2CO3)

Triprotic Acid
– can give up 3 H ions (H3PO4)

pH and pOH
– pH = -log[H or H3O]
– [H or H3O] = 10-pH
– pOH = -log[OH]
– [OH] = 10-pOH
– pH + pOH = 14

Neutralization
– Acid + Base  Salt + Water
– Titrations are done to determine the number of moles when the number of moles of H and OH are equal
– Equivalence point: the point when titration is complete (H = OH)
– End point: a sudden change occurs during a titration
– Equivalence is theoretical and determined by calculations
– End is experimental and determined by indicators