Part 2 Solutions, Acids & Bases - Lesson 1

Part 2 Solutions, Acids & Bases - Lesson 1
C7 & C9 – Properties of Acids and Bases

Properties of Acids and Bases
Acids and Bases can both be defined empirically and theoretically. You are expected to be able to provide any of the following definitions and use them to answer questions.

Empirical Definitions
Acid – a substance which dissolves in water to produce a solution that: Tastes sour Turns blue litmus red Conducts electricity Reacts with active metals to produce H2(g) Neutralizes Bases Base – a substance which dissolves in water to produce a solution that: Tastes bitter; feels slippery Turns red litmus blue Neutralizes acids

Theoretical Definitions
a) Original Arrhenius: Acid – a substance that forms an acidic solution by ionizing in water to produce free hydrogen ions (H+(aq)) in solution Example: HCl (aq)  H+ (aq) + Cl- (aq) Base – a substance that forms a basic solution by dissociation in water to produce free hydroxide ions (OH-(aq)) in solution Example: NaOH(aq)  Na+(aq) + OH–(aq)

Theoretical Definitions
b) Updated Definition: Acid – a species that forms an acidic solution by reacting with water to produce hydronium ions (H3O+(aq)) Example: HCl (aq) + H2O(aq) H3O+(aq) + Cl- (aq) Base – a species that forms a basic solution by reacting with water to produce hydroxide ions (OH-(aq)) Example: NH3 (aq) + H2O(aq) NH4+(aq) + OH–(aq) The hydronium ion (hydrated proton) – was discovered by Paul Giguère at the Université Laval in 1957.

The Hydronium Ion Theoretical chemists thought it was unlikely that a hydrogen ion, which is a tiny proton with a very high charge-to-size ratio, could exist on its own in aqueous solution. Instead they thought it would bond strongly to polar water molecules. Paul Giguère provided empirical evidence for this type of bonding when he discovered hydrated protons (H3O+(aq)); commonly called hydronium ions The modern view of acids and bases is that the hydronium ions (H3O+(aq)); are responsible for acidic properties and hydroxide ions (OH–(aq)) are responsible for basic properties

Learning Tips: The terms “acid” and “base” refer to chemical substances, whereas the terms “acidic” and “basic” refer to the properties of a solution Chemists have known about some acids for hundreds of years. Because these acids are commonly known by familiar names, their IUPAC names are not often used. You should know both forms of naming for acids. Do you remember how to name acids?

Naming Acids Systematic IUPAC Traditional aqueous hydrogen chloride
HCl (aq) aqueous hydrogen chlorate HClO3(aq) aqueous hydrogen chlorite HClO2(aq) Traditional hydrogen ______ide = hydro______ic acid ex. hydrogen chloride = hydrochloric acid hydrogen ______ate = __________ic acid ex. hydrogen chlorate = chloric acid hydrogen ______ite = __________ous acid ex. hydrogen chlorite = chlorous acid Remember: acid contains sulfur = you add a ur H2SO4(aq) sulfuric acid acid contains phosphorus = add an or H3PO4(aq) phorphoric acid

Practice – Naming Acids
IUPAC HClO4(aq) HClO(aq) HNO2(aq) HNO3(aq) H2SO4(aq) H2SO3(aq) H3PO4(aq) Traditional aqueous hydrogen perchlorate perchloric acid aqueous hydrogen hypochlorite hypochlorous acid aqueous hydrogen nitrite nitrous acid aqueous hydrogen nitrate nitric acid aqueous hydrogen sulfate sulfuric acid aqueous hydrogen sulfite sulfurous acid aqueous hydrogen phosphate phosphoric acid

Homework Acid Naming WS WB pg 49 Q 11 WB pg 60 Q 1-5

C8 - pH and pOH Calculations
Lesson 2 C8 - pH and pOH Calculations

Pure Water Pure water actually self ionizes (called “auto-ionization”), so it contains H+(aq) and OH-(aq) ions, but their concentrations are so low that a conductivity test is negative. In a sample of pure water, about two out of every billion molecular collisions are successful in forming hydronium and hydroxide ions 2H2O(l)  H3O+(aq) and OH-(aq) In pure water at SATP, the hydronium ion concentration is very low; about 1 x 10-7 mol/L This concentration is often negligible and will show no conductivity unless very sensitive equipment is used (pg. 238 – Figure 1)

Pure Water Adding acid to water adds H+(aq) ions causing the H+(aq) concentration to increase, thus it makes the solution conductive Adding base to water adds OH-(aq) ions causing the OH-(aq) concentration to increase, thus it makes the solution conductive Aqueous solutions exhibit a wide range of hydronium ion concentrations – from more than 10 mol/L for concentrated HCl(aq) to less than mol/L for concentrated NaOH(aq) This range is called pH; meaning “power of hydrogen” “the negative of the base ten exponent for the hydronium ion concentration” [H3O +(aq)] = pH

pH – power of hydrogen [H3O +(aq)] = 10 –pH
This range is called pH; meaning “power of hydrogen” “The negative of the base ten exponent for the hydronium ion concentration” [H3O +(aq)] = 10 –pH 1 x 101 mol/L 1 x 10-7 mol/L 1 x mol/L Acidic solution Neutral Basic solution pH = -1 pH = 15 pH = 7

pH – power of hydrogen [H3O +(aq)] = 10 -pH
The pH scale is used to communicate a broad range of hydronium ion concentrations. Most common acids and bases have pH values between 0 and 14

pH changes Changes in pH can be deceptive. Adding vinegar to pure water might change the pH from 7 to 4. While this change of 3 pH units may not appear significant, the change in hydronium ion concentration is 103 or 1000 times larger

Practice Try pg. 239 #1-3

pOH ?? pH + pOH = 14

pH Calculations pH = -log [H3O +(aq)]
Do you think solutions always have a pH that is an integer or simply a power of 10? No, scientists often need pH measurements to one or more decimal places So our definition of [H3O +(aq)] =10 –pH must be improved so we can convert numbers like 6.7 x 10-8 mol/L to a pH Our new definition: pH = -log [H3O +(aq)] pH = -log [ 6.7 x 10-8] * the units are dropped because a log has no units pH = - ( ) pH = – but how many sig digs can it have?

pH Calculations Sig digs for pH:
“The number of digits following the decimal point in the pH value is equal to the number of sig digs in the hydronium ion concentration.” [H3O +(aq)] = 6.7 x (two sig digs) pH = (two sig digs)

pH Calculations So from [H3O +(aq)] to pH we use:
pH = -log [H3O +(aq)] pH = -log (4.5 x 10-10) pH = 9.35 (two sig digs) But to go from pH to [H3O +(aq)] we can still use: [H3O +(aq)] =10 –pH [H3O +(aq)] = [H3O +(aq)] = 4.5 x mol/L Since pH has no units, the definition of pH includes the requirement that concentration be in mol/L; you will need to add the units to your answer.

Using your calculator:
Go to pg. 241 and read the two Learning Tips Numbers in scientific notation are best entered using the exponent key (EE) – because the calculator treats the entry as one value. The 10x key is not recommended because you may obtain the incorrect result in some situations Try it: Turn [H3O +(aq)] = 4.7 x mol/L into a pH value Calculator: (-) log 4 . 7 2nd , (-) 1 1 enter

Using your calculator:
A solution has a pH of Calculate its hydronium ion concentration. [H3O +(aq)] =10 –pH [H3O +(aq)] =10 –5.3 [H3O +(aq)] = x 10-6 Sig digs? (pH = 5.3 – only 1 sig dig) = 5 x mol/L Try it with your calculator: Calculator: 2nd log (-) 5 . 3 enter

pOH and Hydroxide ion Concentration
Although pH is used more commonly, in some applications it is more practical to describe hydroxide ion concentration. The definition of pOH follows the same format as pH Example: Calculate the hydroxide ion concentration of water with a pOH of 6.3. pOH = -log [OH -(aq)] [OH -(aq)] =10 –pOH [OH -(aq)] =10 –pOH [OH -(aq)] =10 – 6.3 = 5 x 10-7 mol/L

Summary pH = -log [H3O+(aq)] [H3O+(aq)] =10 –pH pOH = -log [OH -(aq)]
[OH -(aq)] =10 –pOH The number of digits following the decimal point in a pH or pOH value is equal to the number of significant digits in the corresponding hydronium or hydroxide concentration. For both pH and pOH, an inverse relationship exist between the ion concentration and the pH or pOH. The greater the hydronium ion concentration, the lower the pH is.

Example What is the hydronium concentration of a solution with a pH of 3.35?

Example What is the [OH-] of a solution with a pH of 8.605

Example 5.00g of HCl is dissolved in 250mL of water. Find the [H3O+], [OH-], pH and pOH

Practice WB pg 61 Q 6-9 pH and pOH WS

C7 & C8 - Acid – Base Indicators
Lesson 3 C7 & C8 - Acid – Base Indicators

Acid – Base Indicators Substances that change colour when the acidity of the solution changes are known as acid-base indicators A very common indicator used is litmus, which is obtained from lichen Litmus paper is prepared by soaking absorbent paper with litmus solution and then drying it.

Acid – Base Indicators Acid-base indicators are unique chemicals because they can exist in two forms, each with a distinctly different colour The form of the chemical depends on the acidity of the solution. They usually have very complicated formulas, so simple abbreviations are used: Lt – litmus; Bb – bromothymol blue; In – indicator (generic) The two forms of any indicator depend on whether a particular hydrogen atom is present in the indicator’s molecule. In general, the lower pH form is designated HIn(aq) In general, the higher pH form is designated In-(aq)

pH range of colour change
Acid – Base Indicators Other acid-base indicators: Because the chemical structure of each indicator is different, the pH at which the indicator changes from the HIn(aq) form to the In-(aq) form is different for each indicator. (See inside back cover of textbook) Common name Color of HIn(aq) pH range of colour change Color of In-(aq) Bromothymol blue Yellow Blue Phenolphthalein Colourless Pink

A Close Look at Litmus Paper
The colour changes of litmus are a little more complicated than what you have learned previously. There is a “fuzzy” region around the neutral point (pH=7) where the colour is not easily distinguished. Between , litmus is in the process of changing colour and you will see mixtures of red and blue. These intermediate colours are somewhat useful, with the exception of bromothymol blue, where the green is a distinct colour in the middle of the range Common name Color of HIn(aq) pH range of colour change Color of In-(aq) Litmus Red Blue

Uses of Acid-Base Indicators
Acid base indicators have two primary uses: Indicators are commonly used to mark the end of a titration (you will learn more about this in Chapter 8) Indicators are used to estimate the pH of a solution by using a number of different indicators. In this method, acid-base indicators are used to replace the more expensive pH meter, although indicators are not as accurate. Example: Solution After addition to samples of the solution … pH range implied by each indicator pH of solution A ● methyl red was red ● thymol blue turned yellow ● methyl orange turned red pH must be < 4.8 pH > 2.8 but < 8.0 pH must be < 3.2 Combining these results, the pH must be between 2.8 and 3.2

Practice Lab Exercise 6.B pg. 247
Complete the Analysis Portion of the lab below Solution After addition to samples of the solution … pH range implied by each indicator pH of solution A ● methyl violet was blue ● methyl orange was yellow ● methyl red was red ● phenolphthalein was colourless B ● indigo carmine was blue ● phenol red was yellow ● bromocresol green was blue ● methyl red was yellow C ● thymol blue was yellow ● bromocresol green was yellow ● methyl orange was orange

Rainbow Connection

Practice WB pg 62 Q 10-11

C9 - Explaining Acids and Bases
Lesson 4 C9 - Explaining Acids and Bases

Do you remember? Arrhenius Definition of Acids:
A substance that forms an acidic solution by dissolving in water to produce free hydrogen ions (H+(aq)) in solution Example: HCl (aq)  H+ (aq) + Cl- (aq) Arrhenius Definition of Bases: A substance that forms a basic solution by dissolving in water to produce free hydroxide ions (OH-(aq)) in solution Example: NaOH(aq)  Na+(aq) + OH–(aq) What was the problem??

Arrhenius Arrhenius did not know that a hydrogen ion is better described as a hydronium ion; a hydrogen ion bonded to a water molecule Easy to predict HCl(aq): HCl (aq)  H+ (aq) + Cl- (aq) ACID But what about: NaHCO3(aq) or CO2(g) ACID or BASE??

Do you remember? Modified Definition of an Acid:
A species that forms an acidic solution by reacting with water to produce hydronium ions (H3O+(aq)) Example: HCl (aq) + H2O(aq) H3O+(aq) + Cl- (aq) Example: NaHSO4(aq)  Na+(aq) + HSO4- (aq) Dissociation... HSO4- (aq) + H2O(aq) H3O+(aq) + SO42- (aq) Then reacts with water

Do you remember? Modified Definition of a Base:
A species that forms an basic solution by reacting with water to produce hydroxide ions (OH-(aq)) Example: Ca(OH)2 (aq)  Ca2+(aq) + 2 OH- (aq) There is no need to show the reaction with water because hydroxide ions are already present in Ca(OH)2 so a simple dissociation produces the hydroxide ions Example: NH3 (aq) + H2O(aq) NH4+(aq) + OH–(aq) The modified definition (reaction with water) is necessary here to explain how hydroxide ions form in solution. Example: Na2CO3 (aq)  2 Na+(aq) + CO32-(aq) Dissociation... CO32- (aq) + H2O(aq) HCO3-(aq) + OH- (aq) then reacts with water

Summary Write the chemical formulas for the reactants: molecule or polyatomic ion water. • Ignore any Group 1 and 2 cations, and Group 7 anions in the compound formula. Evidence from many compounds shows that these ions do not produce acidic or basic solutions. • If the substance is a non-metal oxide (e.g. CO2(g), SO2(g)), use two moles of water for every mole of the substance in the reactants. Note the evidence provided. If the final solution is acidic, write hydronium ions as the first product. If the final solution is basic, then hydroxide ions are the first product. Complete the other product by determining the combination of atoms and charge required to balance the chemical equation. • The other product should be a recognizable chemical formula—usually a polyatomic ion on your polyatomic ion chart (see the inside back cover).

Practice In a test of the modified Arrhenius theory, a student tested the pH of a solution made by dissolving solid sodium cyanide in water, and found it to have a pH greater than 7. Can the modified Arrhenius theory explain this evidence? Provide your reasoning. NaCN(s)  soluble ionic compound – therefore dissociates: NaCN (s)  Na+(aq) + CN-(aq) then CN- reacts with water CN-(aq) + H2O(l)  OH-(aq) + HCN(aq) The modified Arrhenius theory can explain the basic character of a sodium cyanide solution because it is possible to write a balanced chemical equation with valid products, including the hydroxide ion.

Practice Carbon dioxide is a major air pollutant from the combustion of fossil fuels. Suggest a possible chemical reaction that explains the acidity of a carbon dioxide solution. CO2(g) + 2 H2O(aq)  H3O+(aq) + HCO3-(aq) The modified Arrhenius theory can explain the acidic character of a carbon dioxide solution because it is possible to write a balanced chemical equation with valid products, including the hydronium ion.

More Practice.... HI(aq) – explain acidic properties
HI(aq) + H2O(l)  H3O+(aq) + I- (aq) NaCH3COO(aq) – explain basic properties NaCH3COO(aq)  Na+ (aq) + CH3COO-(aq) Simple Dissociation CH3COO-(aq) + H2O(l)  CH3COOH (aq) + OH-(aq) HOCl(aq) – explain acidic properties HOCl + H2O(l)  H3O+(aq) OCl- (aq)

Try Three More... H3PO4(aq) – explain acidic properties
H3PO4(aq) + H2O(l)  H3O+(aq) + H2PO4- (aq) Na2SO4 (aq) – explain basic properties Na2SO4 (aq)  2 Na+ (aq) + SO42-(aq) Simple Dissociation SO42-(aq) + H2O(l)  HSO4- (aq) + OH-(aq) Sr(OH)2(aq) – explain basic properties Sr(OH)2(aq)  Sr2+(aq) OH- (aq)

Will I ask you to predict if something is an acid or base?
No, for example: Na2HPO4(s) → 2 Na+(aq) HPO42-(aq) Will the solution be acidic, basic, or neutral? We can write valid equations to predict that either hydronium ions or hydroxide ions will form when hydrogen phosphate ions react with water: HPO42-(aq) + H2O(l) → H3O+(aq) PO43-(aq) Or HPO42-(aq) + H2O(l) → OH-(aq) H2PO4-(aq) Nothing that you have studied so far in this textbook enables you to predict which is correct. If you know that the solution turns blue litmus red, then you can select one of the equations to explain the evidence.

Neutralization Reactions
Can now be defined as the reaction between hydronium and hydroxide ions to produce water. Example: HCl(aq) + NaOH(aq)  NaCl(aq) + H2O(l) If we mix an acid and a base, we must be mixing solutions of hydronium and hydroxide ions, so the modified Arrhenius theory predicts: H3O+(aq) + OH-(aq)  2H2O(l)

Summary Acids are substances that react with water to produce hydronium ions Most bases are substances that react with water to produce hydroxide ions Neutralization can be explained as the reaction between hydronium ions and hydroxide ions to produce water.

Practice pg. 253 #1,2,4,5 WB pg 62 Q 12,

C10 – Strengths of Acids and Bases
Lesson 5 C10 – Strengths of Acids and Bases

What do you remember? ACIDS BASES Tastes sour Turn blue litmus red
Fill in as many empirical properties for acids and bases as you can ... ACIDS BASES Tastes sour Turn blue litmus red pH less than 7 Neutralize bases React w/ active metals to produce H2(g) React w/ carbonates to produce carbon dioxide Taste bitter Feel slippery Turn red litmus blue pH greater than 7 Neutralize acids

Do all acids and bases have the same properties and to the same degree?
Two different acids of the same concentration were tested with a pH meter. HCl(aq) 1.0 mol/L = pH = 1 CH3COOH 1.0 mol/L = pH = 4 ??? How does this happpen ???

Strong and Weak Acids The concept of strong and weak acids was developed to describe and explain the differences in properties of acids. An acid can be described as a weak acid if its characteristic properties (under the same conditions) are less than those of a common strong acid There are relatively few strong acids: HClO4(aq), HI(aq), HBr(aq), HCl(aq) H2SO4(aq), HNO3(aq)

The Difference: Using the Modified Arrhenius Theory
Strong Acids: have high conductivity, high rate of reaction w/ metals and carbonates and a relatively low pH These empirical properties suggest many ions are present (lots of H3O+ ions present); which is consistent with the idea that strong acids react completely (>99%) with water to form hydronium ions HCl(aq) + H2O(l)  H3O+(aq) + Cl-(aq) >99%

The Difference: Using the Modified Arrhenius Theory
Weak Acids: have low conductivity, a lower rate of reaction w/ active metals and carbonates and a relatively higher pH These empirical properties suggest fewer hydronium ions are present Based on this evidence, a weak acid reacts incompletely (<50%) with water to form relatively few hydronium ions CH3COOH(aq) + H2O(l)  H3O+(aq) + CH3COO-(aq) <50%

WEAK ACIDS Because of this difference, weak acids are much safer to handle – you can even eat and drink many of them. Tannic Acid in Tea Stearic Acid in Animal Fat Lactic acid in many dairy products and builds up in muscles Citrus acid in fruits Carbonic acid in drinks

Practice – Strong and Weak Acids
Pg. 255 #1-5

Strong and Weak Bases Strong bases – have a high electrical conductivity, fast reaction rate and a very high pH (>>7) Weak bases – have a low electrical conductivity (if molecular bases), slower reaction rate and a pH closer to, but greater than, 7 How do we explain the difference?

NH3 (aq) + H2O(aq) OH–(aq) + NH4+(aq)
Strong and Weak Bases Strong Bases – all soluble ionic hydroxides that dissociate completely (>99%) to release hydroxide ions NaOH(s)  Na+(aq) + OH-(aq) Weak Bases – an ionic or molecular substance that reacts partially (<50%) with water to produce relatively few hydroxide ions NH3 (aq) + H2O(aq) OH–(aq) + NH4+(aq) <50%

Practice Explain the weak base properties of baking soda.
NaHCO3 (aq)  Na+(aq) + HCO3-(aq) HCO3- (aq) + H2O(aq) H2CO3(aq) + OH- (aq) Explain why sodium acetate has a pH of 8 when dissolved in water. NaCH3COO(aq)  Na+ (aq) + CH3COO-(aq) CH3COO-(aq) + H2O(l)  CH3COOH (aq) + OH-(aq) <50% <50%

Practice Pg. 257 #6, 7, 9

Summary Strong Acids Weak Acids Strong Bases Weak Bases
Empirical properties (need same concentration & temperature) Very low pH Med to low pH Very high pH Med to high pH High conductivity Low conductivity Low conductivity* Fast reaction rate Slow reaction rate Modified Arrhenius Theory Completely react with water to form H3O+(aq) ions Partially react with water to form H3O+(aq) ions Completely react with water to form OH-(aq) ions Partially react with water to form OH-(aq) ions * Applies only to weak bases that are molecular

Practice WB pg 49 Q 11

C10 – Polyprotic Acids and Bases
Lesson 6 C10 – Polyprotic Acids and Bases

Polyprotic Acids Some acids have only one acidic hydrogen atom in their compound formula (HA); these are called monoprotic acids i.e. HI(aq), HCl(aq), HBr(aq), HF(aq) HCN(aq) Some acids contain more than one acidic hydrogen and can react more than once with water (HxA); these are called polyprotic acids i.e. H2SO4(aq), H3PO4(aq) Example: H3PO4(aq) + H2O(l)  H3O+(aq) + H2PO4-(aq) H2PO4-(aq) + H2O(l)  H3O+(aq) + HPO42-(aq) HPO42-(aq) + H2O(l)  H3O+(aq) + PO43-(aq) <50% Each of the resulting ions can react with an additional water to produce more hydronium ions, but these reactions are even less complete <1% ~o%

Polyprotic Acids With all of these possible reactions you might think that phosphoric acid is very acidic, but only the first reaction is significant. Also, the first reaction is only 50% complete, so it is actually a weak acid and it has a pH noticeable higher than a strong acid at the same concentration. Based on this evidence: In general, polyprotic acids are weak acids whose reaction with water decreases with each successive step. Acid Concentration pH HCl(aq) 0.1 mol/L 1.0 H3PO4(aq) 1.7 NOTE: H2SO4(aq) is a notable exception to this rule. It is a strong acid because it’s first reaction with water is essentially complete. However the second reaction is much less than 50% complete.

Polyprotic Bases The same concept applies to bases. Some bases, like CH3COO-(aq), are monoprotic bases, meaning they can react with water only once to produce hydroxide ions CH3COO-(aq) + H2O(l)  OH-(aq) + CH3COOH (aq) Others, like CO32- (aq) are polyprotic bases because they can react more than once with water, but the reactions are much less than 50% CO32-(aq) + H2O(l)  OH-(aq) + HCO3-(aq) HCO3-(aq) + H2O(l)  OH-(aq) + H2CO3(aq) <50% <1%

Polyprotic Bases Bases like Ba(OH)2 are not polyprotic bases. This is because there is no possibility of successive reactions with water. Barium hydroxide dissociates to produce hydroxide ions, it does not matter that there are two moles of hydroxide ions for every one mole of Ba(OH)2 Ba(OH)2 (aq)  Ba2+(aq) + 2OH- (aq)

Polyprotic Bases In general, polyprotic bases are weak bases, whose reaction with water decreases with each successive step. Evidence shows that a carbonate ion solution is less basic than a strong base at the same concentration Base Concentration pH NaOH(aq) 0.1 mol/L 13.0 Na2CO3(aq) 11.4 Practice pg. 259 #1-8

Unit B Review We just finished Unit B 
Things to work on to prepare for Unit 2 Exam WB pg Topic Review Questions

Part 2 Solutions, Acids & Bases - Lesson 1

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