Solutions

1. Solutions

2. Types of Solutions • Solid in Solid • Copper in silver (sterling silver) • Zinc in copper (brass) • Solid in Liquid • Salt in water (ocean water) • Solid in Gas • Microscopic particulates in air • Mothball particles in air

3. Types of Solutions • Liquid in Solid • Mercury in silver amalgams (tooth fillings) • Liquid in Liquid • Ethylene glycol in water (engine antifreeze) • Methanol in water (gas line antifreeze) • Liquid in Gas • Water vapour in air

4. Types of Solutions • Gas in Solid • Hydrogen in palladium (purification of hydrogen) • Poisonous gases adsorbed in carbon (charcoal filter) • Gas in Liquid • Carbon dioxide in beverages (carbonated beverages) • Oxygen in water (supporting aquatic life) • Gas in Gas • Oxygen in nitrogen (air)

5. Electronegativity and Water • Water is made up of three atoms – two hydrogen atoms and an oxygen atom • The oxygen atom forms a covalent bond with each of the hydrogen atoms • The oxygen atom has a higher electronegativity than the hydrogen atoms • Electronegativity is the ability of an atom of an element to attract electrons when the atom is in a compound. • In general, electronegativity values decrease from top to bottom within a group and increase from left to right across a period.

6. Electronegativity Values for Selected Elements Oxygen has a higher electronegativityvalue than hydrogen, therefore when they form a covalent bond, oxygen attracts the electron pair more strongly.

7. Water – A Polar Molecule • A polar bond is a covalent bond in which there is a separation of charge between one end and the other - in other words in which one end is slightly positive and the other slightly negative. • Because the oxygen atom in water is more electronegative than the hydrogen atoms, the electron pair spend more time near the oxygen atom giving it a slightly negative charge. As a result, the hydrogen atoms are slightly positive.

8. Polarity • An electrostatic attraction between the partial positive charge near the hydrogen atoms and the partial negative charge near the oxygen results in the formation of a hydrogen bond. • This bonding between water molecules results in strong intermolecular forces which leads to a high boiling point in water. • The ability of ions and other molecules to dissolve in water is due to polarity. For example, sodium chloride is shown in its crystalline form and dissolved in water.

9. The Solution Process • In the previous example, as individual solute ions (Na+ and Cl-) break away from the crystal, the negatively and positively charged ions become surrounded by solvent (water) molecules and the ionic crystals dissolve. • Polar solvents, such as water dissolve ionic compounds and other polar compounds. • Nonpolar solvents, such as gasoline dissolve nonpolar compounds. • Water is polar and oil is nonpolar – do not mix

10. Factors Affecting Solution Formation • The compositions of the solvent and the solute determine whether a substance will dissolve. • The following three factors determine how quickly a substance will dissolve: • Stirring (agitation)- fresh solvent is continually brought into contact with the surface of the solute. • Temperature – higher temp = higher kinetic energy = increased collisions between solvent and solute • Surface Area – the more surface of the solute that is exposed, the faster the rate of dissolving

11. Dissociation of Ionic Compounds in Solution • When an ionic compound dissolves in solution, it dissociates into individual ions. Eg) NaCl Na+ + Cl- CaCl2  Ca+ + Cl- + Cl- • Each individual ion becomes surrounded by water molecules.

12. Covalent Compounds in Solution • Covalently bonded compounds do not dissociate in solution. • Covalent compounds will only dissolve in water if the covalent bonds are polar (like dissolves like) • When a covalent compound dissolves, water molecules surround each individual molecule as opposed to each individual ion. • Electricity can be conducted in an ionic solution (due to the presence of charged particles in solution)but not a covalent one.

13. Solubility • Solubility – is the amount of solute that dissolves in a given quantity of a solvent at a specified temperature and pressure to produce a saturated solution. • Saturated Solution – Contains the maximum amount of solute for a given quantity of solvent at a constant temperature and pressure. • Unsaturated Solution – contains less solute that a saturated solution at a given temp. And pressure. • Supersaturated Solution – contains more solute than it can theoretically hold at a given temperature.

14. Concentrations of Solutions • The concentration of a solution is a measure of the amount of solute that is dissolved in a given quantity of solvent. • A dilute solution is one that contains a small amount of solute. • A concentrated solution contains a large amount of solute.

15. Molarity • The most important unit of concentration is Molarity (M). Molarity is the number of moles of solute dissolved in one litre of solution. • Molarity (M) = moles of solute Litres of solution • The volume involved is the total volume of the resulting solution, not the volume of the solvent alone. • When the symbol M is accompanied by a numerical value, it is read as “molar”.

16. Making Dilutions • Diluting a solution reduces the number of moles of solute per unit volume, but the total number of moles of solute in solution does not change. • Because the number of moles of solute remains unchanged upon dilution, the moles of solute can be represented as moles of solute = M1x V1 = M2 x V2 • M1 andV1 are the molarity and volume of the initial solution, and M2 and V2 are the molarity and volume of the diluted solution.

17. Dilution Example • How many millilitres of aqueous 2 M MgSO4 solution must be diluted with water to prepare 100 mL of aqueous 0.4 M MgSO4? Note: Volumes can be in litres or millilitres as long as the same units are used for both V1 and V2.

18. Percent Solutions • The concentration of a solution can also be described by the percent of a solute in the solvent. • Percent solutions can be expressed in three ways: • % volume/volume (% v/v) – the ratio of the volume of the solute to the volume of the solution. Used when both the solute and solvent are liquids. • % mass/mass (% m/m) – the ratio of the mass of the solute to the mass of the solution. This is also referred to as % weight/weight (% w/w). Used when the solute is a solid. • % mass/volume (% m/v) – the ratio of the mass of the solute (g)to the volume of the solution (mL). This is also referred to as % weight/volume (% w/v). Not a true percentage because of different units in the numerator and denominator.

19. Calculating Percent Solutions • % volume/volume = volume of solute x 100% volume of solution Eg) What is the percent by volume of ethanol (C2H6O) in the final solution when 85 mL of ethanol is diluted to a volume of 250 mL with water?

20. Calculating Percent Solutions • % mass/mass = mass of solute x 100% mass of solution Eg) You want to make 2000g of a solution of glucose and water that has a 2.8% concentration of glucose. How much glucose and water would you need to make this solution?

21. Calculating Percent Solutions • % mass/volume = mass of solute x 100% volume of solution Eg) Calculate the % m/v of a 200 mL solution containing 6 g of NaCl.

22. Other Measures of Concentration • Parts per million (ppm), parts per billion (ppb) and parts per trillion (ppt) are other measures of concentration. • They are used to denote very low concentrations. • Considered to be dimensionless quantities because no particular unit of measurement is used. • One part per trillion (1 ppt) is a proportion equivalent to one-twentieth of a drop of water diluted into a two-meter-deep, Olympic-size swimming pool.