Ch. 8 - Solutions, Acids & Bases. Solution = a homogeneous mixture of 2 or more substances

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1 Ch. 8 - Solutions, Acids & Bases Solution = a homogeneous mixture of 2 or more substances Solute substance whose particles are dissolved in a solution Solvent substance in which the solute dissolves in 1

2 Solutes/solvents can be liquids, solids &/or gases Air = gases dissolved in gases Kool-aid = solids dissolved in water Solid-Solid solutions particularly of metals are sometimes called ALLOYS 2

3 3 ways things can dissolve in water: 1. Dissociation of ionic compounds = an ionic compound is placed in a solution; the solution exerts a greater pull on the ions than the attraction that exists between the ions; ions are pulled apart or dissociated ; is a physical change When an ionic compound dissolves in water, the charged ends of water molecules surround the oppositely charged ions. 3

4 2. Dispersion of molecular compounds covalent polar molecules are attracted to polar solvent molecules; when enough water (solvent) molecules surround the crystal structure, the water s pull due to its polarity is greater than the pull of the covalent crystal structure and the molecules are pulled away from the crystal and become evenly dispersed throughout the solvent; is a physical change 4

5 3. Ionization of molecular compounds process in which neutral molecules gain or lose electrons; is a chemical change because 2 new ionic substances are formed. H 2 O + HCl H 3 O +1 + Cl -1 5

6 Physical properties of solutions are often different than those of the solute/solvent alone. For example: 1. Conductivity = is greater for an ionic solution than for the solid ionic compound 2. Freezing point = mixing two substances such as salt and water, will lower the freezing point of water because the formation of solvent ice crystals is disrupted by the solute The presence of solute particles affects how a solvent freezes. A Pure water freezes in a hexagonal pattern. B In water salted with MgCl 2, the dissociated Mg 2+ and Cl ions disrupt the formation of ice crystals. 3. Boiling point = can be changed by combining solutes and solvents 6

7 During the formation of a solution, energy is either released (exothermic) or absorbed (endothermic). Heat of solution = difference between energy required to break bonds between solute particles and the bonds between solvent particles vs. the energy formed when solutes/solvent particles are attracted to each other EX: Cold packs consist of 2 bags - one with ammonium nitrate and the other with water. The energy required to combine water and the ammonium nitrate is greater than that given off when the ammonium nitrate is broken apart (a positive heat of solution). Heat is pulled from the water and result is a cold pack! Hot packs use substances with a negative heat of solution (breaking of the solute bonds creates more energy than is required to dissolve in water so heat is given off.) 7

8 8

9 Factors that affect the rate of dissolving do so by affecting the collision rate of solute & solvent particles. They include: 1. Surface area increase surface area = increase in dissolving rate 2. Stirring increased stirring = increased rate of collision and therefore, increased dissolving rate 3. Temperature increased temperature = more molecular motion = increased rate of collision between solute and solvent particles = increased dissolving rate 9

10 8.2 - Solubility & Concentration Solubility = the maximum amount of a solute that dissolves in a given amount of solvent (usually expressed as grams of solute per 100 grams of solvent at a certain temperature) 10

11 Solution solubility is classified as: 1. Saturated solution = contains as much solute as the solvent can hold at a given temperature; if you add more solute, it won t dissolve 2. Unsaturated solution = contains less than the maximum amount of solute than can be dissolved; you can add more solute & it will dissolve 3. Supersaturated solution = contains more solute than it should be able to hold; is a very unstable solution; created by heating a solution, dissolving more solute in it & then slowly cooling it. In the photo sequence, a single crystal of sodium acetate, NaC 2 H 3 O 2, is added to a supersaturated solution of sodium acetate in water. The excess solute rapidly crystallizes out of the solution. 11

12 Factors that can affect solubility: 1. Polarity of the solvent = likes dissolve likes ; polar solvents dissolve polar solutes better than if they were of different polarities 2. Temperature = for solids being dissolved in a liquid, solubility will increase as the temperature is increased; gases are the opposite - increasing temperature decreases the solubility of a gas in a liquid 3. Pressure = Increasing pressure on a gas in a liquid solution increases its solubility 12

13 Concentration of a solution = amount of solute dissolved in a specified amount of solution 13

14 SOLUBILITY CURVE FOR VARIOUS SALTS 14

15 Solubility for Gases 15

16 Section Acids & Bases Acid = a compound that produced hydronium ions (H 3 O + ) when dissolved in water; is considered a proton donor HCl + H 2 O --> H 3 O + + Cl - Properties of Acids: 1. Sour taste Ex: citric acid in food 2. Reactive with metals Ex: foil over spaghetti sauce 3. Produce color changes in indicators Ex: Litmus paper (an indicator changes color when exposed to acids/bases) --> blue paper turns red in acid Common acids: Acetic acid (CH 3 COOH) = vinegar carbonic acid (H 2 CO 3 )= carbonated beverages hydrochloric acid (HCl)= stomach acid sulfuric acid (H 2 SO 4 ) = car batteries 16

17 Uses of acids 17

18 Base = compound that produces hydroxide ions (OH - ) when dissolved in water; considered a proton acceptor NaOH --> Na + + OH - Properties of Bases: 1. Bitter taste Ex: Cacao beans contain the base theobromine, which makes chocolate bitter 2. Slippery feel Ex: Soap, cleaning products 3. Color changes in indicator Ex: Bases turn red litmus paper blue; phenolphthalein turns magenta in the presence of a base Ex: Hydrangeas color depends on the acidity of the soil (purple if acidic, pink if soil is basic) 18

19 Common Bases: Aluminum hydroxide (Al(OH) 3 ) = deodorant calcium hydroxide (Ca(OH) 2 ) = concrete Magnesium hydroxide (Mg(OH) 2 ) = antacid Sodium hydroxide (NaOH) = drain cleaner 19

20 Sodium bicarbonate, or baking soda, is often added to swimming pools to regulate the acidity of the water. 20

21 21

22 Neutralization = reaction between an acid and a base where the negative ions in an acid combine with the positive ions in a base to produce an ionic compound called a salt and water. NaOH + HCl --> NaCl + H 2 O Common salts: Sodium chloride (NaCl) Sodium carbonate (Na 2 CO 3 ) Potassium chloride (KCl) Magnesium chloride (MgCl 2 ) 22

23 ph scale --> 0-14; measure of the concentration of hydronium ions in a solution; equals -log (H 3 O + concentration) Acids --> ph of less than 7 Neutral solution --> ph = 7 Bases --> ph greater than 7 The lower the ph value, the greater the H ion concentration. (high ph = low H concentration) 23

24 Strong acids = when placed in water, they ionize almost completely Ex: HCl, H 2 SO 4, HNO 3 Strong bases = when placed in water, they completely dissociate Ex: NaOH, Ca(OH) 2, KOH Weak acids/weak bases ionize or dissociate only slightly in water. Buffer solutions = prepared by mixing weak acid or base with its salt; buffers are solutions that are resistant to large changes in ph 24

25 Buffer solutions = prepared by mixing weak acid or base with its salt; buffers are solutions that are resistant to large changes in ph 25

26 26

27 Electrolyte --> a substance that ionizes or dissociates into ions when dissolved in water (salts, strong acids and strong bases are all examples of electrolytes); electrolytic solutions are good conductors of electricity Batteries & fuel cells use electrolytic solutions 27

28 Attachments cold packs.pdf Acids vs. Bases Comparison.notebook

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