Prentice Hall Physical Science

Prentice Hall Physical Science Chapter 8 Notes Ms. S. Phelps

8.1 Formation of Solutions - for a solution to form, one substance must dissolve in another - remember that a solution is a homogeneous mixture A. Dissolving - a solution has two parts 1. a solute which is the substance that is dissolved 2. a solvent which is what the solute dissolves in - solutes and solvents can be any state of matter - the solutions you are most familiar with have water as a solvent - substances can dissolve in water in three ways – dissociation, dispersion, and ionization - DISSOCIATION – polar water molecule pull the ions in an ionic compound apart (fig. 3, p. 229) - DISPERSION – when a solute breaks into small pieces that spread throughout the water evenly (fig. 4, p. 230) - IONIZATION – occurs when the solute and the solvent lose or gain electrons and it involves a chemical change B. Properties of Liquid Solutions - the conductivity, freezing point, and boiling point are the physical properties of a solution that differ from those of the solute and the solvent - CONDUCTIVITY – when ions are dissolved they can move freely and therefore conduct electricity - FREEZING POINT – the presence of ions in water interfere with its freezing process and cause the freezing point to lower (the solution will freeze at a lower temperature) - BOILING POINT – solutions also have higher boiling points than the solvent alone C. Heat of Solution - when solutions are formed, energy is either released or absorbed (hot packs and cold packs (p. 233) D. Factors Affecting Rates of Dissolving - surface area, stirring, and temperature affect how quickly a solution forms - SURFACE AREA – the greater the surface area of the solute the quicker the solution forms be cause there are more places for collisions to take place EX: crushed solutes dissolve faster (think powdered drinks) - STIRRING – increases that rate at which a solution forms because the solute is moving around and colliding more - TEMPERATURE – increasing temperature increases the rate at which a solution forms because the particles move faster and collide more often

1 | P a g e Prentice Hall Physical Science Chapter 8 Notes Ms. S. Phelps

8.2 Solubility and Concentration A. Solubility - solubility is the maximum amount of solute that dissolves in a given amount solvent at a constant temperature - saturated solutions contain as much solute as a solvent will hold at a given temperature; if you add more solute it will not dissolve - unsaturated solutions have less than they can hold; if you add more it will dissolve - supersaturated solutions contain MORE solute than they can normally hold at a given temperature; if you add more solute, even one crystal, all of the extra solute will come out of solution B. Factors Affecting Solubility - the polarity of the solvent, temperature, and pressure affect solubility (how much will dissolve) - POLARITY OF SOLVENT – “likes dissolve likes” polar solvents dissolve polar solutes and nonpolar solvents dissolve nonpolar solutes EX: oil and water - TEMPERATURE – increasing temperature usually increases solubility, but it decreases the solubility of gases - PRESSURE – increasing pressure increases the solubility of a gas (carbonated drinks) C. Concentration - concentration is the amount of solute dissolved in a specific amount of solvent - high concentration means there is a lot of solute in the solvent and low concentration means there is a small amount of solute in a solvent - percent by volume = volume of solute/volume of solution *100% - percent by mass = mass of solute/mass of solution * 100%

8.3 -8.4 Acids and Bases and Strength of Acid sand Ba 1. General Definitions: Acid: a substance which when added to water produces hydrogen ions [H+]. Base: a substance which when added to water produces hydroxide ions [OH-]. 2. Properties: Acids:

 react with zinc, magnesium, or aluminum and form hydrogen (H2(g)) 2-  react with compounds containing CO3 and form carbon dioxide and water  turn litmus red  taste sour (lemons contain citric acid, for example) DO NOT TASTE ACIDS IN THE LABORATORY!! Bases:  feel soapy or slippery  turn litmus blue  they react with most cations to precipitate hydroxides  taste bitter (ever get soap in your mouth?) DO NOT TASTE BASES IN THE LABORATORY!! + - 3. Water dissociation: H2O(l) → H (aq) + OH (aq) + - equilibrium constant, KW = [H ][OH ] / [H2O] Note: water is not involved in the equilibrium expression because it is a pure liquid, also, the amount of water not dissociated is so large compared to that dissociated that we consider it a constant + - -14 Value for Kw = [H ][OH ] = 1.0 x 10 + - -14 Note: The reverse reaction, H (aq) + OH (aq) → H2O(l) is not equal to 1 x 10 [H+] for pure water = 1 x 10-7 [OH-] for pure water = 1 x 10-7 Definitions of acidic, basic, and neutral solutions based on [H+]  acidic: if [H+] is greater than 1 x 10-7 M

2 | P a g e  basic: if [H+] is less than1 x 10-7 M  neutral: if [H+] if equal to 1 x 10-7 M Example 1: What is the [H+] of a sample of lake water with [OH-] of 4.0 x 10-9 M? Is the lake acidic, basic, or neutral? Solution: [H+] = 1 x 10-14 / 4 x 10-9 = 2.5 x 10-6 M Therefore the lake is slightly acidic Remember: the smaller the negative exponent, the larger the number is. Therefore:  acid solutions should have exponents of [H+] from 0 to -6.  basic solutions will have exponents of [H+] from -8 on. Example 2: What is the [H+] of human saliva if its [OH-] is 4 x 10-8 M? Is human saliva acidic, basic, or neutral? Solution: [H+] = 1.0 x 10-14 / 4 x 10-8 = 2.5 x 10-7 M The saliva is pretty neutral. 4. pH relationship between [H+] and pH + pH = -log10[H ] Definition of acidic, basic, and neutral solutions based on pH acidic: if pH is less than 7 basic: if pH is greater than 7 neutral: if pH is equal to 7 The [H+] can be calculated from the pH by taking the antilog of the negative pH Example 3: calculate the [OH-] of a solution of baking soda with a pH of 8.5. Solution: First calculate the [H+] if pH is 8.5, then the antilog of -8.5 is 3.2 x 10-9. Thus the [H+] is 3.2 x 10-9 M Next calculate the [OH-] 1.0 x 10-14 / 3.2 x 10-9 = 3.1 x 10-6 M Example 4: Calculate the pH of a solution of household ammonia whose [OH-] is 7.93 x 10-3 M. Solution: This time you first calculate the [H+] from the [OH-] 7.93 x 10-3 M OH- = 1.26 x 10-12 M H+ Then find the pH -log[1.26 x 10-12] = 11.9 Now you try a few by yourself. You can then check your answers using the Java applet that follows, but remember, you won't learn how to do them if you don't try by yourself first. Practice #1. What is the pH of a solution of NaOH that has a [OH-] of 3.5 x 10-3 M? Practice #2. The H+ of vinegar that has a pH of 3.2 is what? Practice #3. What is the pH of a 0.001 M HCl solution?

How can pH be determined experimentally? By using pH paper or a pH meter 5. Strength of Acids and Bases: Acids 1. Strong Acids:  completely dissociate in water, forming H+ and an anion. + 1- example: HN03 dissociates completely in water to form H and N03 . The reaction is + 1- HNO3(aq) → H (aq) + N03 (aq) + - A 0.01 M solution of nitric acid contains 0.01 M of H and 0.01 M N03 ions and almost no HN03 molecules. The pH of the solution would be 2.0.  There are only 6 strong acids: You must learn them. The remainder of the acids therefore are considered weak acids.

1. HCl

2. H2SO4

3. HNO3

4. HClO4

5. HBr

6. HI + + -  Note: when a strong acid dissociates only one H ion is removed. H2S04 dissociates giving H and HS04 ions. + 1- H2SO4 → H + HSO4 + 1- A 0.01 M solution of sulfuric acid would contain 0.01 M H and 0.01 M HSO4 (bisulfate or hydrogen sulfate ion). 2. Weak acids: 3 | P a g e  a weak acid only partially dissociates in water to give H+ and the anion for example, HF dissociates in water to give H+ and F-. It is a weak acid. with a dissociation equation that is + - HF(aq) ↔ H (aq) + F (aq)  Note the use of the double arrow with the weak acid. That is because an equilibrium exists between the dissociated ions and the undissociated molecule. In the case of a strong acid dissociating, only one arrow ( → ) is required since the reaction goes virtually to completion.  An equilibrium expression can be written for this system: + - Ka = [ H ][F ] / [HF]  Which are the weak acids? Anything that dissociates in water to produce H+ and is not one of the 6 strong acids. 1. Molecules containing an ionizable proton. (If the formula starts with H then it is a prime candidate for being an acid.) Also: organic acids have at least one carboxyl group, -COOH, with the H being ionizable. 1- + 2- 2. Anions that contain an ionizable proton. ( HSO4 → H + SO4 ) 3. Cations: (transition metal cations and heavy metal cations with high charge) + + also NH4 dissociates into NH3 + H Bases 1. Strong Bases:  They dissociate 100% into the cation and OH- (hydroxide ion). + - example: NaOH(aq) → Na (aq) + OH (aq) a. 0.010 M NaOH solution will contain 0.010 M OH- ions (as well as 0.010 M Na+ ions) and have a pH of 12.  Which are the strong bases? The hydroxides of Groups I and II.  Note: the hydroxides of Group II metals produce 2 mol of OH- ions for every mole of base that dissociates. These hydroxides are not very soluble, but what amount that does dissolve completely dissociates into ions. 2+ - exampIe: Ba(OH)2(aq) → Ba (aq) + 2OH (aq) a. 0.000100 M Ba(OH)2 solution will be 0.000200 M in OH- ions (as well as 0.00100 M in Ba2+ ions) and will have a pH of 10.3. 2. Weak Bases: What compounds are considered to be weak bases? 1. Most weak bases are anions of weak acids. 2. Weak bases do not furnish OH- ions by dissociation. They react with water to furnish the OH- ions. Note that like weak acids, this reaction is shown to be at equilibrium, unlike the dissociation of a strong base which is shown to go to completion. 3. When a weak base reacts with water the OH- comes from the water and the remaining H+ attaches itsef to the weak base, giving a weak acid as one of the products. You may think of it as a two-step reaction similar to the hydrolysis of water by cations to give acid solutions. examples: + - NH3(aq) + H2O(aq) → NH4 (aq) + OH (aq) + - methylamine: CH3NH2(aq) + H20(l) → CH3NH3 (aq) + OH (aq) - - acetate ion: C2H3O2 (aq) + H2O(aq) → HC2H302(aq) + OH (aq) - General reaction: weak base(aq) + H2O(aq) → weak acid(aq) + OH (aq) Since the reaction does not go to completion relatively few OH- ions are formed. Acid-Base Properties of Salt Solutions: definition of a salt:  an ionic compound made of a cation and an anion, other than hydroxide.  the product besides water of a neutralization reaction determining acidity or basicity of a salt solution: 1. split the salt into cation and anion 2. add OH- to the cation a. if you obtain a strong base. the cation is neutral b. if you get a weak base, the cation is acidic 3. Add H+ to the anion a. if you obtain a strong acid, the anion is neutral b. if you obtain a weak acid. the anion is basic 4. Salt solutions are neutral if both ions are neutral 5. Salt solutions are acidic if one ion is neutral and the other is acidic 6. Salt solutions are basic is one of the ions is basic and the other is neutral. 7. The acidity or basicity of a salt made of one acidic ion and one basic ion cannot be determined without further information. Examples: determine if the following solutions are acidic, basic, or neutral Click on each one to find out the answer.

4 | P a g e KC2H3O2 NaHPO4

Cu(NO3)2 LiHS

KClO4 NH4Cl 6. Acid-Base Reactions:  Strong acid + strong base: HCl + NaOH → NaCl + H2O + - net ionic reaction: H + OH → H2O Strong acid + weak base:

example: write the net ionic equation for the reaction between hydrochloric acid, HCl, and aqueous ammonia, NH3. What is the pH of the resulting solution? Strong base + weak acid:

example: write the net ionic equation for the reaction between citric acid (H3C6H507) and sodium hydroxide. What is the pH of the resulting solution? 7. Titrations 1. Nomenclature: these are terms that are used when talking about titrating one substance with another. You need to learn these definitions well enough to explain them to someone else.  titration  titrant  indicator  equivalence point  end point  titration cuve 2. Strong acid-strong base titration example: titration curve pH at equivalence point species present appropriate indicators 3. Strong acid-weak base titration example titration curve pH at end point species present appropriate indicators 4. Weak acid-strong base titrations

5 | P a g e example: titration curve for the titration of vinegar with NaOH pH at end point- approximately 8.5 species present- H2O and NaC2H3O2 appropiate indicator-phenolphthalein

Note: no matterwhat type of titration you do, at the equivalence (end) point the number of moles of H+ is equivalent to the number of moles of OH-. This applies whether you have weak or strong acids and/or bases. + Problems: l. Citric acid (C6H807) contains a mole of ionizable H /mole of citric acid. A sample containing citric acid has a mass of 1.286 g. The sample is dissolved in 100.0 mL of water. The solution is titrated with 0.0150 M of NaOH. If 14.93 mL of the base are required to neutralize the acid. then what is the mass percent of citric acid in the sample? 2. A sample of solid calcium hydroxide is mixed with water at 30 oC and allowed to stand. A 100.0 mL sample of the solution is titrated with 59.4 mL of a 0.400 M solution of hydrobromic acid. a. What is the concentration of the calcium hydroxide solution? o b. What is the solubility of the calcium hydroxide in water at 30 C? Express your answer in grams of Ca(OH)2 / 100 mL water? 8. Three models of acids: l. Arrhenius Model Basis for the model--action in water  acid definition: produces H in water solution  base definition: produces OH1- in water solution 2. Bronsted-Lowry Model Basis for the model-- proton transfer  acid definition: donates a proton ( H )  base definition: accepts a proton  conjugate acid definition: the acid becomes the conjugate base after it donates the proton because it can now accept it back.  conjugate base definition: the base becomes the conjugate acid after it accepts the proton because it can now donate it back. 3. Lewis Model Basis for model--electron pair transfer  acid definition: accepts a pair of electrons  base definition: donates a pair of electrons

6 | P a g e