Unit 8: Acids, Bases, and Salts

Unit 8: Acids, Bases, and Salts

Topic 1: Properties of Acids and Bases

i. PROPERTIES OF ACIDS A) Acids have a _______________taste (but NEVER TASTE unknown substances) Examples: Acetic acid (vinegar), citric acid (orange juice), lactic acid (yogurt) SOUR

i. PROPERTIES OF ACIDS B) Acids have a __________________ 0 7 14
General formula = HX (Where X = ANION such as Cl-) pH LESS THAN 7

i. PROPERTIES OF ACIDS C.) Aqueous acids solutions are ELECTROLYTES. That means they CONDUCT ELECTRICITY. The strength of an electrolyte depends on the NUMBER OF IONS. Strong acid = strong electrolyte = good conductor Weak acid = weak electrolyte = poor conductor

i. PROPERTIES OF ACIDS D.) Acids cause acid-base INDICATORS to change colors. Table M lists several common indicators. Example: Litmus paper turns PINK in an acidic environment

i. PROPERTIES OF ACIDS E) Acids react with certain METALS to produce HYDROGEN GAS (H2 (g) ) with a SINGLE REPLACEMENT reaction. The metals ABOVE hydrogen in Table J react with acid to produce H2 (g) Example: 2 HCl(aq) + 2 Zn(s) → 2 ZnCl (aq) + H2(g)

i. PROPERTIES OF ACIDS F)Acids react with BASES to form a SALT AND WATER. These types of reactions are called NEUTRALIZATION reactions (DOUBLE REPLACEMENT reactions). Example: HCl(aq) + NaOH(aq)  NaCl + HOH Acid + Base Salt + Water

i. PROPERTIES OF ACIDS G) Examples of common acids (Table K)
(H2CO3)- Carbonic acid in soda (HC2H3O2)- Acetic acid in vinegar (HCl)- Hydrochloric acid in stomach acid (H2SO4)- Sulfuric acid in car batteries (HNO3) - Nitric acid in explosives (H3PO4)- Phosphoric acid in flavorings

iI. PROPERTIES OF bases Bases taste BITTER and feel SLIPPERY (but NEVER TASTE unknown substances) Examples: antacids, soaps, ammonia-based cleaning products

iI. PROPERTIES OF bases B. Bases have pH VALUES > 7 (fall on HIGHER end of pH SCALE) General formula = XOH (Where X = CATION such as Na+)

iI. PROPERTIES OF bases C.) Aqueous solutions bases solutions are ELECTROLYTES. That means they CONDUCT ELECTRICITY_. The strength of an electrolyte depends on the NUMBER OF IONS. Strong base = strong electrolyte = good conductor Weak base = weak electrolyte = poor conductor

iI. PROPERTIES OF bases D. Bases cause acid-base INDICATORS to change color. Table M lists several common indicators. Example: LITMUS paper turns BLUE in an alkaline environment

iI. PROPERTIES OF bases E. Bases react with ACIDS to form a SALT AND WATER. These types of reactions are called NEUTRALIZATION reactions (DOUBLE REPLACEMENT reactions). Example: 2 HBr(aq) + Ca(OH)2(aq) → CaBr2 + 2 H2O Acid + Base  Salt + Water

iI. PROPERTIES OF bases F. Examples of common bases (TABLE L)
NaOH- sodium hydroxide (LYE) soaps, drain cleaner Mg (OH)2 - magnesium hydroxide-antacids Al(OH)3-aluminum hydroxide-antacids, deodorants NH4OH-ammonium hydroxide- “ammonia”

III. PROPERTIES OF SALTS
Salts are NEUTRAL ionic substances that have positive ions other than HYDROGEN (H+) and negative ions other than HYDROXIDE (OH-)

III. PROPERTIES OF SALTS
Salts are formed from NEUTRALIZATION REACTIONS and are NEUTRAL acid + base → SALT + water Examples: LiBr, KI, CaCl2, NaNO3

Topic 2: Naming Acids and Bases

IV. Naming acids and bases
Naming Binary Acids (hydrogen with nonmetal): - Name the hydrogen in the acid hydro- - Name the anion and change ending of the name to –ic acid Examples: 1. HCl (aq) - _ ____________________________ 2. HI (aq) - _____________________________ 3. HBr(aq) - ____________________________ 4. H2S(aq) - _____________________________ Hydrochloric acid Hydroiodic acid Hydrobromic acid Hydrosulfuric acid

Naming Ternary Acids (hydrogen with Polyatomic ions) Go to Table ______ and name the PAI (polyatomic ion) DO NOT use the prefix hydro – If the PAI ends in - ate, it changes to –ic acid If the PAI ends in – ite, it changes to –ous acid Examples: 1) H2CO3 (aq) - _ ____________________________ 2) H3PO4 (aq) - _____________________________ 3) H2SO4 (aq) - _____________________________ 4) H2SO3(aq) - _____________________________ E Carbonic acid Phosphoric acid Sulfuric acid Sulfurous acid

Naming Bases: - Bases are all ____________________, therefore, they are named just like ________________ Name 1st element Use roman numeral if the metal has more than one charge listed 2nd half of the name will typically be _______________ Examples: 1. NaOH (aq) - ____________________ IONIC COMPOUNDS IONIC COMPOUNDS HYDROXIDE Sodium hydroxide

Naming Bases (continued): 2. Ca(OH)2 (aq) - _____________________________ 3. Fe(OH)3 (aq) - _____________________________ 4. Co(OH)2 (aq) - _____________________________ Calcium hydroxide Iron (III) hydroxide Cobalt (II) hydroxide

Topic 3: Acid/Base Theories

V. What is an acid (Arrhenius acid)?
H2O To be an acid, it must be dissolved in __________ When acids are dissolved in ______________ , they ______________ and form ______________________ IMPORTANT: ACIDS ARE SUBSTANES THAT CONTAIN… See Table ___ in Reference Tables for ________________ H2O BREAK APART FREELY MOVING IONS H+ IONS THAT IONIZE WHEN DISSOLVED IN WATER COMMON ACIDS K

V. What is an acid (Arrhenius acid)?
Example: When an acid is dissolved in water, the H+ leaves the acid and bonds to the water molecule to form a ________________________ HCl (g) + H2O (l)  H3O + (aq) + Cl - (aq) The acid HCl contains 1 H+ which combines with 1 H2O to form 1 H3O+ HYDRONIUM ION (H3O+) Dissociation of acids in water animation

VI. What is a Base? (Arrhenius Base)
IMPORTANT: BASES ARE SUBSTANES THAT CONTAIN… See Table _________ in Reference Tables for ____________________ Example: When a base is dissolved in water, the OH - 1 leaves the base NaOH (s) ______________________ OH- (Hydroxide) ions dissolved in aqueous solution. L COMMON BASES Na + 1 (aq) + OH - 1 (aq)

VII. Arrhenius acids and Bases
A. Review Arrhenius Acids and Bases Acids: ___________ ions are the only____________ ions in water Bases: ___________ ions are the only ___________ ions in water Bronsted-Lowry (Alternate Theory)Acids and Bases Acids: are proton ___________________________________ Bases: are proton ___________________________________ H + (aq) POSITIVE OH --- (aq) NEGATIVE DONORS, GIVE AWAY H+ (aq) to another substance ACCEPTORS, RECEIVE H+ (aq) from another substance

VIII. Bronsted-Lowry (Alternate Theory)Acids and Bases
Acids: are proton ___________________________________ Bases: are proton ___________________________________ DONORS, GIVE AWAY H+ (aq) to another substance ACCEPTORS, RECEIVE H+ (aq) from another substance

IX. Conjugate Acid-Base Pairs
Consists of two substances related by the loss or gain of a single hydrogen ion The ionization of HCl: HCl (aq)  H+1 (aq) + Cl-1 (aq) Animation and Animation 2 Example: HCl (aq) + HOH (l)  Cl -1 (aq) + H3O+1 (aq) One acid and one base per side IX. Conjugate Acid-Base Pairs Conjugate base Conjugate acid Acid Base One that gains H+ is the conjugate acid on opposite side One that loses H+ is the conjugate base on opposite side

Example: NH3 (aq) + HOH (l)  NH4+1 (aq) + OH-1 (aq) HOH (l) + HOH (l)  H3O+ (aq) + OH-1 (aq) HC2H3O2 (aq)+ H2O (l)  H3O + (aq) + C2H3O2 – (aq) Conjugate acid Conjugate base Base Acid Conjugate acid Conjugate base Base Acid Conjugate acid Conjugate base Acid Base

Examples: Based on the acid-base theory (Bronsted-Lowry Acid and Base), identify the acid and base for the following: Remember: B/L acid: LOSE H+ in the reaction B/L base: GAIN H+ in the reaction NH3 + H2SO4  NH4 + + HSO4 – HBr + NH2 -  NH3 + Br - NH3 + HCl  NH4 + + Cl - IX. Conjugate Acid-Base Pairs B A CA CB A B CA CB B A CA CB

X. Weak vs. strong acids Acids always produce H3O+ (hydronium) ions when dissolved in water. The more H+ that dissociate when in water, the stronger the acid will be. Examples of Strong Acids – ____________________ Example of Weak Acid – ______________ H Cl, H NO3, H Br, H I CH3CH2COOH (HC2H5O2) (organic acids) WEAK VS. STRONG ACID ANIMATION Weak vs. Strong Acid Animation #2

XI. Strong vs. weak bases Like acids, the more ions that dissociate in water, the stronger the base is. Bases formed with group 1 and 2 metals are typically stronger. Examples of Strong Bases – __________________ Example of Weak Bases – ___________________ LiOH, NaOH, Mg(OH)2 Al(OH)3, NH4OH

Topic 4: Indicators

XII. How to know if a solution is acidic or basic
Indicator – Acid-Base Indicators (Found on Table ___________) Acid-Base Indicators and narrowing down pH using multiple indicators (mixture of indicators gives great range of colors, pH paper) Remember: An acid has a pH _________________ and a base has a pH _________________ a substance that changes color if an acid or base is added M less than 7 above 7

Tips Using Table M: If the indicator changes TO the listed color, pH is ___________ the pH range. If the indicator does NOT change, pH is __________ the pH range. The pH range does NOT indicate that the solution has pH within that range above below

Tips on using Table M Tip 1- Notice "Common Acid Base Indicators," acid is on the left and base is on the right. Low pH numbers are on the left and high pH numbers are on the right. They match up. Tip 2 - Look at Methyl Orange, red is on the left, 3.2 is on the left. To the left (or below) a pH of 3.2 the color is red. To the right (above) a pH of 4.4 the color is yellow. Tip 3 - Between the 2 numbers will never be asked on a test. This is where the indicator changes color.

Example: A solution has a pH of 5.5. What color changes will happen to following indicators. 1) Methyl orange __________ 3) Bromscresol green____________ 2) Bromthylmol blue ________ 4) Thymol blue ______________ yellow blue yellow yellow

Example: A solution yields the following results when tested with various indicators: Methyl Orange = yellow Phenolphthalein = clear Bromcresol Green = blue Thymol Blue = yellow - Can the pH be: a) 2.8 b) c) d) 4.8  pH above 4.4  pH below 8.0  pH above 5.4  pH below 8 XXXX XXXX XXXX

Topic 5: pH Scale

XIII. The Power of Hydronium (What ph really means)
The pH scale is easiest method to use for comparing the strengths of acids and bases. Water breaks up very slightly to form hydrogen ions and hydroxide ions: H2O  H +1 + OH – 1 In neutral water, the concentration of H + = 1.0 x M, so the pH = _____ 7

A pH of 0-7 is __________ When the concentration of H+ is... A pH of 7 is ____________ M, pH = ____, , pH =___ A pH of 7-14 is _____________ , pH = _____ acidic 1 2 neutral basic 3

The pH scale is logarithmic. Each change of a single pH unit signifies a tenfold change in the concentration of the hydrogen ion. Therefore, the[H+] is ten times greater in a solution of pH of 5 as in a solution of pH of 6. Examples: A solution with pH of 3 is 10 times more acidic than solution with a pH of ____ B) A solution with pH of 3 is ___________ times more acidic than solution with a pH of 6. C) A solution with a pH of 2.0 has a hydronium ion concentration 10,000 times greater than a solution with a pH of ___________ 4 (for every power of 10, pH goes UP 1) 1,000 6

The pH scale is logarithmic. Each change of a single pH unit signifies a tenfold change in the concentration of the hydrogen ion. Therefore, the[H+] is ten times greater in a solution of pH of 5 as in a solution of pH of 6. Examples (continued): D) A sample of base has pH of 9. What would the pH of a base 1,000 more basic be? E) A sample of acid has a pH of 3. What would be the pH of an acid 100 times more acidic be? 12 (for every power of 10, the pH goes UP 1 becomes more basic) 1 (for every power of 10, the pH goes DOWN 1 becomes more acidic)

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Think about this… If you add a strong acid to a base, what would happen to the pH of the mixed solution? pH would decrease

Topic 6: Neutralization

XIV. Acid and base neutralization
Neutralization: When an acid and base are mixed, the ____________________ and the ______________________ combine to form _______________. The anion of the acid and the cation of the base come together to form a _________________. Neutralization of Acid and Base Video Clip Double Replacement Animation Rainbow Connection Demo (Titration with pH indicators) Salt – __________________________________________ H+ of the acid OH --- of the base Water (H2O) salt An ionic compound (metal with nonmetal/PAI) that can be formed by acid-base neutralization

Examples: 1) HCl (aq) + NaOH (aq)  _____ _____ ______ ___________ 2) H2SO4(aq) + 2 __________  K2SO4 (aq) + 2 _________ _______ _______ ___________ ___________ 3) 2 HNO3(aq) + __________  Ca(NO3)2 (aq) + 2 __________ _______ _______ _______ _________ DOUBLE REPLACEMENT REACTIONS!!! HOH (l) + NaCl(aq) salt acid base water HOH(l) KOH (aq) acid salt base water Ca(OH)2 (aq) HOH (l) acid base salt water

Example: Write the equation for the neutralization reaction between dilute nitric acid (HNO3) and potassium hydroxide (KOH). (DON’T FORGET TO BALANCE) 1 1 1 1 ___ HNO3( aq) + ___ KOH (aq)  ___ H(OH)(l) + ___ K(NO3)(aq) Potassium nitrate water

Example: Write the equation for the neutralization reaction between dilute Mg(OH)2 and HCl. (DON’T FORGET TO BALANCE) 1 2 2 1 ___ Mg(OH)2 ( aq) + ___ HCl (aq)  ___ H(OH)(l) + ___ MgCl2(aq)

XV. Titration (video demo)
Definition: process of adding a known amount of solution of known concentration to determine the concentration of another solution moles of H+1 (aq) = moles of OH-1 (aq) Molarity of Acid * Volume of Acid = Molarity of base * Volume of Base Titration Equation (see Table T): Titration Animation Titration Animation #2 One mole of H + neutralizes one mole of OH -- MA(VA) = MB(VB) OR (#H+)(MA)(VA) = (#OH---)(MB)(VB)

XV. Titration Example: What is the concentration of hydrochloric acid solution if 50.0 mL of a M KOH are needed to neutralize 20.0 mL of the HCL solution of unknown concentration? (#H+)(MA)(VA) = (#OH---)(MB)(VB) #H+= 1 MA = ? MA = M VA = 20.0 mL #OH--- = 1 MB = M VB = 50.0 mL

XV. Titration Example: What is the concentration for sulfuric acid solution if 50 mL of a M KOH are needed to neutralize 20 mL of the sulfuric solution of unknown concentration? H2SO4 (#H+)(MA)(VA) = (#OH---)(MB)(VB) #H+= 2 MA = ? MA = M VA = 20.0 mL #OH--- = 1 MB = M VB = 50.0 mL

XV. Titration Example: How many mL of 2.0 M H2SO4 are required to neutralize 30.0 mL of M NaOH? (#H+)(MA)(VA) = (#OH---)(MB)(VB) #H+= 2 MA = 2.0 M VA = 7.5 mL VA = ? #OH--- = 1 MB = 1.0 M VB = 30.0 mL

XV. Titration Example: How many mL of 0.10 M Ca(OH)2 are required to neutralize 25.0 mL of M HNO3? (#H+)(MA)(VA) = (#OH---)(MB)(VB) #H+= 1 MA = 0.50 M VB = 62.5 mL VA = 25.0 mL #OH--- = 2 MB = 0.10 M VB = ??

Unit 8: Acids, Bases, and Salts

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Unit 8: Acids, Bases, and Salts

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