Acids and Bases: an Introduction

Acids Sour Taste Electrolytes Identified by the presence of H+ ions (Arrhenius). Proton (H+) donor (Bronsted-Lowry). Common Acids: vinegar, citric acid, sulfuric acid, hydrochloric acid. There are many different ways acids can be defined. For our purposes, acids have an H and produce H ions and bases have an OH and produce hydroxide ions. Acids—donates hydrogen ions Vinegar (acetic acid), hydrochloric acid (HCl)—present in our stomach/stomach acid that helps digest food, citric acid present in fruits such as organges and lemons, sulfuric acid (H2SO4) present in car batteries Write down compounds: HCl, H2SO4, HNO3 NEVER add water to acid, ALWAYS ACID TO WATER (A—W), wear eye protection when working with them.

ACIDS An acid  H+ in water HNO3, nitric acid HNO3  H+ + NO3- Strong acids are strong electrolytes and completely dissociate in water. HNO3, nitric acid HNO3

ACIDS An acid  H+ in water HCl hydrochloric HBr hydrobromic The following are examples of strong acids. HCl hydrochloric HBr hydrobromic HI hydroiodic HNO3 nitric H2SO4 sulfuric HNO3

HCl (aq)  H3O+ (aq) + Cl-(aq) The Hydronium Ion Free H+ is highly reactive and does not actually exist in nature. It reacts with water to form the hydronium ion, H3O+. HCl (aq)  H3O+ (aq) + Cl-(aq) HCl H 2 O 3 + Cl - hydronium ion

Weak Acids WEAK ACIDS are weak electrolytes and do not completely dissociate in water. CH3CO2H acetic acid H2CO3 carbonic acid H3PO4 phosphoric acid Dissociation of a Weak Electrolyte

Bases Bitter taste Slippery Like acids, they are electrolytes Identified by the presence of OH- ions (Arrhenius). Proton (H+) acceptor (Bronsted-Lowry). Proton acceptor—receives hydrogen ions Compounds: NaOH (lye), Mg(OH)2—anyone know what this is? Milk of Magnesia, baking soda, ammonia (NH3) Bases can be used as antacids—acid reflux.

BASES Base  OH- in water NaOH(aq)  Na+(aq) + OH-(aq) 04m08an1 NaOH is a strong base Strong bases are strong electrolytes and soluble in water

Strength of Acids and Bases Based on the concentration of H+ or OH- ions in a solution. Strong Acids/Bases: completely dissociate into ions in a solution. Weak Acids/Bases: do NOT completely dissociate into ions in a solution. Think about strong vs. weak coffee or tea. More stuff in concentrated solution—therefore stronger. Strong: completely break up into ions, write on board: HCl --- H+ + Cl- , we don’t have to worry about factoring water in. Weak: acids/bases not completely broken into ions, must write reaction in reverse as some ions will reform. Write acetic acid (weak acid) on board– CH3CO2H + H2O --- H3O+ + CH3CO2- (we can write H as hydronium ion H3O) **Give List**

Strong Acid Example: HCl

Weak Acid Example: CH3CO2H

“BIG 6”---Strong Acids (Know them!!) HClO4 HI HCl HNO3 HBr H2SO4

Strong Bases (Know them!!) Group I metal hydroxides (NaOH, KOH, etc.) Soluble/Slightly soluble Group II metal hydroxides ( Ca(OH)2, Sr(OH)2, Ba(OH)2 )

Conjugate Acids/Bases Acids and bases are related to each other through the addition/loss of hydrogen ions Conjugate acid-base pairs Acids produce conjugate bases Bases produce conjugate acids

Conjugate Examples HA + H2O  H3O+ + A- HNO3 + NH3  NH4+ + NO3-

Conjugate Acid/Base Strength Stronger the acid, the weaker the conjugate base Stronger the base, the weaker the conjugate acid Weak acids/bases have strong conjugate bases/acids

Acid Nomenclature Binary acids– containing only 2 elements (one is hydrogen) 1) Prefix “hydro—” with binary acid 2) root name for second element 3) End the name with “IC acid” Oxyacids– acids containing hydrogen, oxygen, and a nonmetal 1) use the given polyatomic ion name from anion 2) add “IC acid”

Naming Acids Binary acids contain only two different elements. Formula Name HF hydroflouoric HCl hydrochloric HBr hydrobromic HI hydriodic

Naming Acids Oxyacids contain a polyatomic ion containing oxygen. Formula Name Polyatomic Anion CH3COOH Acetic Acid CH3COO- H2CO3 Carbonic acid CO32- HNO3 Nitric Acid NO3- H2SO4 Sulfuric Acid SO42- H3PO4 Phosphoric Acid PO43-

Example 1: HBr Write name Hydrobromic acid

Now you try, HCl Write name. Hydrochloric acid

Example 2: Hydrofluoric acid Write chemical formula HF

Now you try, hydriodic acid. HI

Example 3: H2SO4 Write the name. Sulfuric acid -name these acids based on the anion, look at your reference packet when you notice that it is a polyatomic ion. Sulfate—drop –ate, and add –ic acid ending.

Example 4: HClO2 Write the name. -ous ending due to one less oxygen Chlorous acid.

Try: HClO3 HClO HClO4 Chloric acid, 2) hypochlorous acid, 3) perchloric acid

Base Nomenclature Name of cation, name of anion/hydroxide Ex. NaOH Sodium hydroxide.

Brainstorm… Briefly brainstorm how we measure the amount of hydrogen ions in a solution? How can we quickly measure the acidity or basicity of a solution?

pH How do we measure the amount of “stuff” in a solution? **concentration** How do we determine a solution’s concentration? **molarity Same way with acids and bases except we look at the hydrogen and hydroxide ions to determine how concentrated (ie acidic or basic) a solution is. How can we measure the acidity or basicity of a solution? How would this measurement relate to the hydrogen and hydroxide ion concentrations in solution?

Let’s take a look at water 2H2O  H3O+ + OH- [H3O+][OH-]= 1x10-14M2 [H3O+]= 1x10-7M [OH-]= 1x10-7M Kw= 1x10-14M2 When water is neutral (neither acidic or basic), hydrogen and hydroxide ions dissociate and each have a concentration of 1x10-7 mol/L or M. This value is obtained at 25 degrees Celcius and equilibrium. If we combine these concentrations together, the product = 1x10-14M = Kw (water dissociation equilibrium constant) Ions dissociate in water is small compared to when acid or base is added SO we will be looking at concentrations of the acid or base added. We are only going to be dealing with strong acids and bases

Example 1 An acid is added to water and gives a hydroxide ion concentration [OH-] of 1.0x10-12M. What is the hydrogen ion concentration [H+] ? What happens to the hydrogen ion concentration when we add acid to water? What about hydroxide ions with base? Example 1: [OH-] = 1.0x10-12M *Write equation, 1.0x10-14M2 = [H+][OH-], plug in what we know and solve. 1.0x10-14M = [H+] (1.0x10-12M), [H+] = 0.01M or 1x10-2M **So this value tells us the concentration of a solution, how acidic it is? Is this value easy to understand? NO, so the pH scale was developed to easily determine how acidic or basic a solution is. how many of you have used a pH meter or pH indicators for a pool, hot tub, etc.? Why do you need to know the pH, what do you do? We have certain pH levels that must be maintained in our body for us to function normally (blood pH, stomach acid pH, etc.). *may mention buffers as a way of ensuring pH levels remain stable*

What is pH? Pouvoir hydrogene: “hydrogen power” pH = measure of [H3O+]------- Acidity [H3O+] expressed in powers of 10 Ex. 10-14 to 10-1

pH Scale Range from 0-14. NEUTRAL, pH=7. (pure water) BASE, pH > 7. (ocean water, milk of magnesia, baking soda) ACID, pH < 7. (stomach acid/HCl, vinegar, soft drinks) Logarithmic scale= base 10 system, relationship between a solution’s acidity/basicity and the concentration of H+ ions in solution Ex. [H+] = 1x10-7 is a pH of 7 (take the negative log), neutral. [H+] = 1x10-10 is a pH of 10, basic, more or less hydrogen ions in solution? LESS * draw out number [H+] = 1x10-2 is a pH of 2, acidic, more or less hydrogen ions in solution? MORE *draw out number Blood pH is 7.3-7.5, mostly neutral. Hydrogen ion concentration is inversely proportional to the scale (pH goes down/increase H ions, pH goes up/decrease H ions)—try with OH ions too. *have students write it down.

pH Scale Draw pH scale. Look at common things on scale.

How do we measure the pH of a solution? Acid-base indicators (ex. litmus paper) pH meter Ask students if they have measured pH in a pool, etc. What did they use? Once we get the pH of a solution, how can we find the concentration of H ions in a solution? Use the log relationship between pH and H ions Types of indicators: methyl red, phenolphtalein, methyl orange (indicators have a certain pH range for a color change)