Acids and Bases
Acids and Bases in Medicine, Health, and Cosmetics
- Aluminum hydroxide, magnesium hydroxide, calcium carbonate: antacids -Magnesium hydroxide: antiperspirant/deodorant - Hydrochloric acid: stabilizer for many medications - Carbonic acid: regulator of acidity in blood.
A net reaction involves 2 pairs:
- An acid-conjugate base pair - A base-conjugate acid pair
examples of salts
- Antacids - Agricultural lime added to acidic soil
examples of Acids in Food
- Benzoic acid: preservative - Carbonic acid: carbonated drinks - Acetic acid: Vinegar - Tartaric acid: soft drinks, tartness to food - Citric acid: tartness, citrus fruits, preservative
Tools used in titration
- Buret - Erlenmeyer flask - volumetric pipet
Lewis concept
- Covers more situations than Bronsted-Lowry or Arrhenius - Considers reactions that do not involve H⁺ transfer ***Useful when describing reactions that produce new covalent bonds through electron transfers
Base buffers of interest
- HCN and conj base CN⁻ w/ pKa = 9.2 - NH₄ and conj base NH₃ w/ pKa = 9.2 - The bicarbonate buffer system is important for maintaining blood pH.
Common pH values:
- Household bleach: pH ≈ 13 - Ammonia-based window cleaner: pH ≈ 11 - Blood: pH ≈ 7.4 - Pure water: pH =7 - Vinegar: pH ≈ 2
Examples of Acids and Bases in Cleaning
- Hydrochloric acid (HCl): household cleaners and swimming pool maintenance - Sodium and potassium hydroxides: manufacture of soaps, detergents, and cleaners; drain cleaner - Sodium hypochlorite: bleach, disinfectant - Ammonia: window cleaner, household cleaner
Examples of Acids and Bases in Manufacturing
- Hydrochloric acid: used in electroplating and leather processing - Nitric acid: used in the production of fertilizers and explosives; used to etch metals and purify gold - Ammonia: used in the production of fertilizers and the manufacture of nitric acid; used to neutralize acid in petroleum and in the production of latex
Examples of acids reacting with metal
- Metal etching - Acid rain damage (like on vehicles)
Bronsted-Lowry concept
- Useful in ALL cases described for Arrhenius concept. - Useful when compounds do not fit Arrhenius definitions, particularly bases other than hydroxide ion -Useful for situations NOT involving H2O. *** Useful when describing reactions involving H+ transfers
Arrhenius concept
- Useful when discussing aqueous solutions **Useful when compounds fit the Arrhenius definitions of acid or base.
Strong acid combined with weak base (example: HCl + NH3 → NH4Cl):
- Water not formed as product; no hydroxide ions - Solution pH < 7: ammonium salt dissociates to form some H⁺
Sugar: Does it form ions in water? Conductivity? Electrolyte or non-electrolyte?
- does not conduct electricity - nonelectrolyte
buret
- narrow glass calibrated in 0.1 mL increments. • has a stopcock to control the flow of liquid out the bottom. - is used to measure the amount of base added during titration.
Physical and Chemical Properties of Bases
- taste bitter - are slippery (like soaps!!) - Turn red litmus paper blue - turn colorless phenolphthalein pink - will neutralize acids - may react with some metal salts. - may react with carbon dioxide
Physical and chemical properties of acids
- taste sour - are neutralized by bases (to form salts) 💥 - turn blue litmus paper red. - turn methyl orange solution red - react with carbonates and bicarbonates. - react with metals.
Salts can be acidic, basic, or neutral.
...
A solution of HCl has [H+] = 2.40 × 10-4 M and [OH-] = 4.17 × 10-11 M. What is the pH of this solution? What is its pOH?
1). Write the appropriate equation and identify the variables. 2). Substitute the given values and solve. pH = -log[H⁺] pH= - log[2.40 × 10-4] pH = -log[0.00024] = -(-3.62) = 3.62 pH + pOH = 14 3.62 + pOH = 14 pOH = 10.38 = 10.4
2 methods for measuring pH change
1. chemical indicator 2. pH meter
H2co3
ACID: carbonic acid weak acid can sometimes be a conjugate acid.
Is HNO3 an acid or base?
ACID: nitric acid
Reactions of Acids with Bicarbonate and Carbonate
Acid reacts with carbonate or bicarbonate compound to form CO₂,H₂O, and a salt.
What is formed when acid reacts with metal?
Acid reacts with metal to form a salt and hydrogen GASSS. (JUST h)
HSO4-
Acid. Proton donator (hint: Bronsted-Lowry acid)
Which of the 3 concepts could be used to define KOH as a base?
All of them! (Hydroxide is present)
Examples of bases
Ammonia (NH₃) Sodium hydroxide (NaOH) Potassium hydroxide (KOH) 💥carbonate (CaCO3) **Bases OFTEN have the hydroxide compound (HO) , like Potassium/Sodium hydroxide, but NOT always!! :)
Reactions of Bases with Metal Salts
Base reacts with soluble metal salt to form INsoluble metal hydroxide.
Which of the 3 concepts could be used to define H2O as an acid?
Bronsted-Lowry and Lewis
Which combo will make a buffer?***
CH3COOH and CH3COONa
carbonate formula
CO₃²⁻
Examples of Reactions of Bases with Metal Salts
CuSO₄(aq) + 2NaOH (aq)→Cu(OH)₂(s) + Na₂SO₄ (aq) ... (1 more ex chose not to show)
Weak acids and bases partially dissociate.
Examples: HCN, HF, NH3 HCN⇌H⁺ + CN⁻
Strong acids and bases dissociate completely in solution.
Examples: HCl, HNO3, NaOH
What happens when there are too many bases in human blood?
H ions from bicarbonate ions react w/ the bases.
during alkalosis, which chemical is removed from the body by the kidneys?
HCO₃⁻
Which acid is used in electroplating and leather processing
Hydrochloric acid
Examples of acids
Hydrochloric acid (HCl) in stomach Acetic acid (HCH₃COO) (A dilute solution is vinegar.) Sulfuric Acid (H₂SO₄) Oxalic acid (H₂C₂O₄)
Mg(OH)₂ + 2HNO₃→ Mg(NO₃)₂ + 2H2O What is the Net Ionic equation?
H⁺ + OH⁻→H₂O https://www.youtube.com/watch?v=Jia6f1HBFgA
self-ionization of water
H₂O + H₂O ⇌ H₃O⁺ + OH⁻
Water reacts with more water to produce hydronium and hydroxide
H₂O + H₂O↔H₃O⁺ + OH⁻ In aqueous solution at 25°C: [H₃O⁺][OH⁻]=10⁻¹⁴ For pure water at 25°C: [H₃O⁺] = [OH⁻] = 10⁻⁷ M x 10⁻⁷ = 10⁻¹⁴ [H⁺]
hydronium ion (formula?)
H₃O⁺
Hydronium ion
H₃O⁺ (aq) ** The brackets (that SHOULD be) on the outside of the lewis structure mean a concentration of hydronium ions in solution
equivalence point
In titrations it's the point at which the # of moles of one reactant has been added in stoichiometric quantity to react completely with the moles of the other reactant. - Moles of base added = moles of acid in flask - pH in flask changes dramatically
salt
Ionic compound formed from combining the acid or the base (or the neutralization of acid with a base) made of the acid's anion and base's cation When an acid reacts with a base, a salt and water are formed. Neutralization reaction: ACID + BASE = SALT + WATER
NH3 (ammonia)
Lewis BASE
hydroxide ion (formula?)
OH⁻
Alkalosis
Occurs if blood pH is *above 7.45*
Which of the 3 concepts apply to aqueous solutions?
Only the Arrhenius concept
Which of the 3 concepts could be used to define SO2 as a base?
Only the lewis concept. There's no hydroxide so it can't be Arrhenius Doesn't seem to be a H⁺ transfer, which might be why it's not Bronsted-Lowry
Weak acid combined with strong base (example: CH3COOH + NaOH → Na+ CH3COO− + H2O):
Solution pH > 7; acetic acid not completely ionized
Buffers in biology
The pH in the body is regulated by a buffer in BLOOD based on the following reaction: H⁺(aq) HCO₃⁻ ↔H₂CO₃ (aq) ↔H₂O(l) + CO₂(g) ** Healthy pH range is about 7.35 to 7.45. Exercise increases blood flow [CO₂] and [H+ ] A buffer system shifts equilibrium to compensate for changes:
What is true about solutions of salt and water?
They DO conduct electricity
Important Formula for H3O+ and OH- in pH and pOH
[OH⁻] = (1.0 x 10⁻¹⁴)/[H₃O⁺]
pH indicator
a chemical compound that changes color in response to changes in solution pH - compound that changes color at a specific pH range. - pH meter that reads solution pH directly.
Base
a chemical substance that decreases H+ concentration in aqueous solution or that increases OH⁻ concentration in solution.
pH
a measure of the hydronium ion concentration in an aqueous solution; pH = -log [H⁺] - Acidic: 1-6 -Neutral: pH=7 - Basic(alkaline): 8-14
pOH
a measure of the hydroxide ion concentration in an aqueous solution; pOH = -log[OH− ] pOH is the opposite of pH AND they are inversely related - As pH increases, pOH decreases - Acids have low pH and high pOH. - Bases have high pH and low pOH. While pH is calculated based on the concentration of Hydrogen ions, [H+], pOH is calculated based on the concentration of Hydroxide ions, [OH-], present in the solution
titration
a solution of UNKNOWN concentration is reacted with a STANDARD solution of known concentration; the goal of titration is to use the known concentration and reaction stoichiometries to find the unknown concentration 💥Method for determining the concentration of a solution
standardized base solution
a solution of known concentration
Lewis Acid
a substance that ACCEPTS (-) electrons to form a covalent bond Is usually the more positive reactant.
lewis base
a substance that DONATES ELECTRONS to form a covalent bond Is usually the more negative reactant
Acid
a substance that INCREASES the concentration of H+ (H3O+ ) ions in solution H⁺: hydrogen ion H₃O⁺: hydronium ion
Arrhenius acid
a substance that, when dissolved in water, increases the concentration of H₃O⁺ (or H⁺) (remember bases often contain hydroxide or OH right together, while an acid does NOT contain OH).
Arrhenius base
a substance that, when dissolved in water, increases the concentration of OH⁻
buffer
an aqueous solution containing a weak acid and its conjugate base OR a weak base and its conjugate acid. This solution resists changes in pH when small quantities of an acid or base are added. ** It resists change because of an equilibrium. EX: H₃O⁺(aq) HCO₃⁻ ↔H₂CO₃ (aq) H₂O(l)
What is the buffer system that is found in the body?
bicarbonate buffer system
What is produced when a strong acid reacts w/ the bicarbonate buffer system in the human body? ***
carbonic acid
salt characteristics
compound formed from a reaction of a acid and a base (it's NEUTRAL) ***Usually an electrolyte because it ionizes in water. - most salts are ionic and soluble in water pH: 7
conjugate acid-base pair
consists of two ions or molecules related by the loss or gain or one hydrogen ion
What is salt's conductivity?
do NOT have low conductivity
salt solution
form ions in the solution - Conductivity - Electrolytes: form ions in AQ solution and have high conductivity. - Nonelectrolytes: do not form ions in AQ solution and do not conduct electricity.
How do you chose between Bronsted-Lowry, Arrhenius, or both?
if there is H2O in the product side, it's probable that it's both Bronsted-Lowry AND Arrhenius. However, if no H2O is present, it's probably ONLY Bronsted-Lowry. Arrhenius never appears on its own!
Spectator ions
ions that do not participate directly in the reaction
BH3
lewis acid
Which pair of substances would most likely result in the production of a solid when reacting with an acid?
metal and carbonate
NaCl
neither a base nor an acid; it is a compound of a base and acid. pH= 7
A solution of KOH has a pH of 12.75 and a pOH of 1.25. Find the concentration of H3O+ and OH- ions in the solution.
pH = -log[H3O⁺] 12.75 = -log[H3O⁺] -12.75 = log ₁₀[H3O⁺] 10^(-12.75) = [H3O⁺] =1.78e¹³ pOH = −log[OH⁻] 1.25 = -log₁₀ [OH⁻] -1.25 = log₁₀ [OH⁻] 10^(-1.25) = [OH⁻] 0.056 = [OH⁻] 5.6 x 10^⁻² = [OH⁻]
pH and pOH equations
pH = -log[H₃O⁺] pOH = −log[OH⁻] [H₃O⁺] = 10^(⁻ph) [OH⁻] = 10^(-pOH)
Henderson-Hasselbalch equation
pH = pKa + log [A-]/[HA] Applies to weak acids and bases pKa = negative logarithm of Ka Ka = acid disassociation constant [HA] = concentration of an acid [A-] = concentration of conjugate base 💥pH of desired solution should be close to pKa of acid form of buffer 💥The optimal ratio of [A-] to [HA] for a buffer is 1:1
antacid
prevents or corrects acidity, especially in the stomach. ** It's like an ANTI-acid
Bronsted-Lowry base
proton acceptor. **often negative or neutral. NH₃ + H⁺ →NH₄⁺ Ammonia (NH3) is the bronsted-lowry base here.
Bronsted-Lowry acid
proton donator HCL → Cl⁻ + H⁺
net ionic equation
shows ONLY the ions involved in the reaction. The spectator ions are removed.
total ionic equation
shows all the ions involved in the reaction.
Conjugate acid (it's like the resulting product of a base reactant)
the acid formed as a result of the ADDITION of proton to a Bronsted-Lowry base.
Conjugate base (it's like the resulting product of acid reactant)
the base formed as a result of the REMOVAL of a proton from a Bronsted-Lowry acid.
Acidity
the degree to which a substance is acidic - Lower pH indicates greater acidity
Alkalinity
the degree to which a substance is basic - Higher pH indicates greater alkalinity
Hydrolysis
the disassociation of salts in water - examples are KCI, NaNO₃
Cola contains phosphoric acid and carbonic acid. Which is the most likely reason that cola also contains a large amount of sugar or other sweetener?
to counteract the sourness of the acids! Interesting..
What is the MOST important reason for using a base instead of water to rinse out the burette?
to prevent the dilution of the base, which may affect results
volumetric pipet
used to measure the acid volume before the titration begins
Erlenmeyer flask
vessel for the acid-base reaction
NH4+
weak acid also the conjugate acid of ammonia Bronsted-Lowry acid
Why do buffer systems either use a weak acid or a weak base?
weak acids/bases are partially dissociate
A base is added to an acid during a titration. A pH meter is used to monitor the reaction's progress. When should the addition of base stop?
when adding a few drops of base causes a large increase in pH
How is a buffer prepared?
• A buffer is made by mixing a weak acid/base with its conjugate. • Adding a strong ACID shifts equilibrium to the left. A strong BASE shifts it to the right • Adding a base would not necessarily remove the acidity, and would add many hydroxide ions.
phosphate buffer
• In mammals, cellular pH ranges from 6.9 to 7.4 - A phosphate buffer is effective in maintaining pH in this range.
Step 3 of performing a strong acid-base Titration
• Position the Erlenmeyer flask under the buret. • Slowly add base to acid while swirling the flask to keep contents thoroughly mixed. • Observe any changes as base is added. • Stop additions when the pH increases suddenly • Record the final buret volume. - Locate bottom of meniscus - Record volume to the nearest 0.01 mL
Acid Buffer Systems
• Weak acids only ionize to a small extent: HA ↔H+ + A- .
If [H+ ] = 3.87 × 10^(−11) M, what is [OH- ]?
• Write the appropriate equation pH = −log[H⁺] - Substitute the given values and solve. pH = -log [H⁺] = -log [ 3.87 x 10^-11] = -(-10.4) = 10.4 pOH + pH = 14 pOH + 10.4 = 14 pOH = 3.6
chemical indicators
• measure pH INdirectlu. • change color only in specific pH ranges. • are of limited use. • are pretty cheap
pH meter:
• measures pH directly. • can be used in any acid-base titration. • requires a source of electricity. • requires calibration •expensive.
Step 2 of performing a strong acid-base Titration
∙ Prepare acid by: - clean and rinse volumetric pipet - filling volumetric pipet with acid solution. • transferring acid solution to the flask. ∙ Prepare to measure pH change by: - adding a few drops of chemical indicator or, if using a pH meter, by calibrating the pH meter and then inserting the pH electrode into the acid solution.
Step 1 of performing a strong acid-base Titration
∙ Prepare buret: - clean with soap and water - rinsing w/ water - rinsing with ***standardized base solution (a solution of known concentration) - filling with standardized base solution. ∙ Record the buret volume by: - locating the bottom of the meniscus - recording the volume to the nearest 0.01 mL
How to Carry out Titration Calculations
∙Gather the known information: Includes - concentration of base - initial buret volume of base - final buret volume of base - volume of acid ∙ Write the balanced chemical equation for the reaction ∙ Subtract initial volume from final buret volume to find the volume of base needed to reach the equivalence point ∙ Set up a series of conversions using dimensional analysis to solve for the moles of base used to reach the equivalence point ∙ Use stoichiometric relationships given by the balanced chemical equation to convert moles of base added to moles of acid neutralized ∙ Convert acid volume to liters. Then, divide the moles of acid by the volume of acid to find molarity of the acid
Reactions of Bases with Carbon Dioxide
💥Base reacts with carbon dioxide to form CARBONATE (CO) 2LiOH (aq) + CO2 (g) → Li₂CO₃ (aq) + H2O(l)