SCH3U Grade 11 Chemistry Solutions and Solubility Test 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 gasses 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 – Water is very effective at dissolving solutes due to its small size, highly polar nature, and ability to form hydrogen bonds.
– 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 – Covalent bonds, which are more powerful than hydrogen bonds, hold water together. – 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 colorless
– 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