Acids and Bases

pKa and pKb

-logKa=pKa. Ka=10^(-pKa). -logKb=pKb. Kb=10^(-pKb).

pH of acid

0-6. Closer to 0=strong acid, closer to 7=weak acid.

Monoprotic acid

1 hydrogen ion produced. Ex: HCl, HNO3, CH3COOH

Creating buffer solutions

1. Weak acid + salt solution for A1-. Ions need to be able to find each other and rejoin quickly. It's in water, so to help, you need to make sure you have enough of the A1- in there. See notes for example. 2. Weak base + salt (not using fluorine) solution for base to make sure there's enough B1+ in there to convert back. See notes for example. These ways are the best way...here is the less preferable way: 3. Weak acid and strong base. Weak acid will have higher concentration and volume while strong base has lower concentration and volume. Harder to do, because of you put in too much, it neutralizes everything that's there. (Like 1.0 M and 500 mL of weak acid with 0.1 M and 100mL of strong base.)

pH + pOH =

14.0. (No sit figs)

Brønsted-Lowry theory

1923. Includes non-water based solutions. State: acid is a proton donor, base is a proton acceptor. (Proton=hydrogen ion.) Conjugate acid is formed when the base accepts the proton. Conjugate base is what is left after acid donates proton.

Lewis Theory

1923. States: acid accepts a pair of electrons, base donates pair of electrons. The electrons don't actually move—formation of a coordinate bond. Electrons are donated by setting up coordinate bond. Electrons are accepted by attaching to coordinate bond. No conjugate acids or conjugate bases in this theory.

Diprotic acid

2 hydrogen ions produced. Ex: H2SO4

Triprotic acid

3 hydrogen ions produced. Ex: H3PO4.

pH of neutral solution

7

pH of alkaline solutions

8-14. Closer to 8=weak base, closer to 14=strong base.

Alkali

A base that is soluble in water and produces hydroxide ions.

In reactions of both strong and weak acids, water acts as...

A base, accepting a proton.

An indicator acts like

A buffer. Stays combined in acid to keep buffering going. In base, wants to release ions and get OH1- with H1+.

pH scale

A simple and effective way of representing the concentration of hydrogen ions in a solution. These concentrations are often very low, so we use a logarithmic scale to increase understanding. pH = -log [H¹+(aq)]; [H¹+] = 10^-pH. pH means potential of the hydronium ion/of hydrogen.

Buffer

A solution that resists changes in pH upon the addition of small amounts of a strong base or strong acid, or upon the dilution of the buffer through the addition of water. May be composed of a weak acid and it's conjugate base of a weak base and its conjugate acid. Can also be prepared from partial neutralization of a weak acid with a strong base. Weak acid is present in excess, producing a conjugate acid-base pair with the salt formed. HA ➡️⬅️ H1+ + A1-. We normally change the concentration of hydrogen ions. Increase, shift to produce more reactants. If exposed to a strong acid, will shift to make more reactants because extra hydrogen ions must be compensated for. If add base, hydroxide ions will form water with hydrogen ions and more product will be produced. Only a narrow range that this is workable in though. Buffer solutions must be weak acids or weak bases. Which one depends on what ions we need.

amphiprotic species

A substance that has the ability to act either as a Bronsted Lowry acid or a Bronsted Lowry base depending on the reaction in which it is taking part. One example is water. Polyprotic species are often involved in reactions in which they behave amphiprotically. Amino acids also act as amphiprotic species. All 2-amino acids contain a weakly acidic carboxyl group and a weakly basis amino group. In the ionized form, the compound acts as an acid in the presence of a strong base and as a base in the presence of a strong acid. Subset of amphoteric—but specifically Brønsted-Lowry.

A change of one pH unit is equal to...

A tenfold change in hydrogen ion concentration (since pH scale is a logarithmic scale to base 10).

Occam's Razor and relation to pH scale

A theory should remain as simple as is possible while maintaining a high capacity for gaining understanding. pH scale does a good job of this, clearly distinguishing between acids, neutral solutions, and bases/alkaline solutions.

Titration

A volumetric analysis technique that involves a reaction between a substance of unknown concentration with a standardized solution (the titrant). The titrant is delivered in small increments from a burette into the solution being analyzed. Progress of reaction can be monitored using several techniques. Data loggers combined with a pH probe can be used to collect data that can be plotted to produce a pH curve. A pH curve enables analysis of characteristic features of the titration. An acid-base indicator undergoes a color change as the titration approaches and reaches the equivalence point.

Acid and base conductivity

All acids and bases dissociate to some degree in water and create ions. Conductivity of an aqueous solution depends on concentration of ions present. Can be measured in a simple experiment with a power pack and graphite electrodes connected to an ammeter. Voltage applied must be identical for each solution. Strong acids and bases are strong electrolytes, so they display higher conductivity than weak acids and bases (higher ammeter reading).

Monitoring the rate of reactions of acids with metals, metal carbonates, and metal hydrogencarbonates

All produce a gas, so the rate of gas evolution is monitored. Can be monitored through loss of mass. A series of such experiments can distinguish a strong acid from a weak acid.

Nitrogen base

All weak, like ammonia NH3. Amines, amides, nitriles. (One of these I think has some sort of carbon attachment and NH2.) They produce OH1- in one step.

How does the enthalpy change of neutralization change for a strong acid versus a weak acid?

Almost identical. The neutralization reaction removes the ionized species from dissociation reaction, driving the reaction to completion. A strong acid or base is completely dissociated in solution, so only enthalpy consideration in this neutralization reaction is the exothermic formation of water from hydrogen ions and hydroxide ions. For a neutralization reaction with a weak acid or base: they exist mainly in their undissociated forms in aqueous solution, and the ionization of a weak acid or base is mildly endothermic. Therefore, the enthalpy change of neutralization will be slightly less exothermic for a strong base-weak acid neutralization reaction than for a strong base-strong acid neutralization reaction. The weaker the acid, the more endothermic the dissociation reaction, the lower the enthalpy change of neutralization.

Strong acid

An effective proton donor that is assumed to completely dissociate in water; Ex: HCl, H2SO4, HNO3. Reactions go to completion. The conjugate base of a strong acid is a very weak base—low proton affinity.

Titration

Analytical testing process that can be useful for acids and bases but also with redox. Solution of unknown concentration (?M). M leads to mol (c=n/V), from which we can get pH. Can be either an acid or a base. 1. Check litmus to see if acid or base. 2. If unknown is acid, use known strong base solution. Unknown may be weak or strong so choose a strong base like KOH or NaOH, 1.0 M to make math easy. (If base, use strong acid.) After adding one or two drops of a liquid indicator to a set volume of the unknown, start dropping in known with a burette. Burette tells #mL added. Gradually turns it from one to the other. When color changes briefly but returns to normal—that is the transition period/interval. If you have a pH meter, you record the values. Add drops until it changes all the way fully. Don't add another drop once the color has completely changed. Must constantly be stirred. See notes for equations and pH curves.

Arrhenius theory

Before he published, they'd found many solutions with similar properties consistently. He states: acid is a substance that dissolves in water to give H1+ (H3O1+). Base is a substance that dissolves in water to give OH1-. Theory still works today, mainly because most acids and bases are water solutions. Meets criteria of acids and bases that most people come into contact with (water-based, in biological things. BUT not all acids and bases dissolve in water, so need the expand the theory.

conjugate acid-base pair

Brønsted-Lowry theory. The conjugate acid and conjugate base differ from each other by a single proton. When you make the substance a reactant in the reaction, it will either lose or gain a proton, which indicates whether it's a conjugate acid or a conjugate base.

List the weak bases you need to know.

CaCO3, NH3.

Organic acids

Carboxyl group. Also sometimes called carboxylic acids. IUPAC naming: butanoic acid (4 carbons, single bonds, one carboxyl group). Butanedioic (4 carbons, single bonds, two carboxyl groups). Ethanoic acid=acetic acid.

Selection of an indicator

Choice depends on relative strengths of acid and base and therefore the pH of the equivalence point. All common indicators can be used in strong acid strong base titrations due to very steep rise near equivalence point, covering pH range of most acid base indicators. Otherwise...indicator must change color within transition interval. Strong acid strong base: phenol red: pKa=7.9: pH range 6.8-8.4: yellow in acid, red in base. Strong acid weak base: methyl orange: pKa=3.7: pH range 3.1-4.4: red in acid, yellow in base. Weak acid strong base: phenolphthalein: pKa=9.6: pH range 8.3-10.0: colorless in acid, pink in base.

endpoint of titration

Color change on indicator.

Strong base

Completely dissociates in water; Ex: all group 1 hydroxides (all also dissolve in water), like NaOH and KOH. A metal hydroxide doesn't act as a Brønsted-Lowry base because it can't accept a proton, but in solution, the OH1- acts as a Brønsted-Lowry base, accepting a proton: OH1- (aq) + H3O1+ (aq) —> 2 H2O (l).

Strong acids and strong bases are assumed to...so concentration of a strong monoprotic acid...

Completely ionize/dissociate in aqueous solutions...so concentration of a strong monoprotic acid will equal the concentration of hydrogen ions.

Salt

Compound composed of an anion and a cation.

How does acid deposition affect the environment?

Deforestation; mineral leaching from soils leading to elevated acid levels in lakes and rivers; uptake of toxic minerals from soil by plants; reduction in pH of lake and river systems; increased uptake of toxic metals by shellfish and other marine life, which can affect fishing industry and ultimately people's health; corrosive effects on marble, limestone, and metal buildings, bridges, and vehicles.

Weak acid

Dissociates only partially in water. Poor proton donor. Dissociation of weak acid is a reversible reaction that reaches equilibrium. At equilibrium, only a small proportion of the acid molecules have dissociated. Conjugate base of a weak acid has a higher affinity for a proton (stronger) than does conjugate base of a strong acid.

Ionization constant=

Dissociation constant. When it doesn't completely ionize, we don't know how much of it did. So, we use these calculations to figure it out. See notes for equations and strong vs. weak.

Explain pKa

Doesn't exactly equal pH but helps us understand pH. pKa<2, strong acid. >2 and <7, weak acid. >7 and <10, weak base. >10, strong base.

Buffer capacity

Effectiveness of buffer to resist changes in pH. Depends on molar concentration of the acid and conjugate base: higher concentration, increases effectiveness at resisting changes in pH.

Standard enthalpy change of neutralization

Energy change associated with the formation of 1 mol of water from the rxn between a string acid and a string base under standard conditions. Negative value, as neutralization is exothermic. -#kJ/mol.

Post-combustion methods

Focus on several complementary technologies to remove SO2, nitrogen oxides, heavy metals, and dioxins from the combustion gases before they are released into the atmosphere.

Base dissociation constant

For: B (aq) + H2O (l) —> BH1+ (aq) + OH1- (aq). Kb = ([BH1+][OH1-])/[B].

Acid dissociation constant

For: HA (aq) + H2O (l) —> A1- (aq) + H3O1+ (aq): Ka = ([A1-][H3O1+])/[HA]. For WEAK ACID.

In aqueous solution a proton can either be represented as...

H1+ or H3O1+ (hydrogen or hydronium)

List the weak acids you need to know.

H3PO4 (phosphoric), CH3COOH (acetic/ethanoic acid).

List the strong acids you need to know.

HCl (hydrochloric), HNO3 (nitric), H2SO4 (sulfuric).

Équivalence point

Halfway through transition interval.

Calcination

Heating of materials to very high temperatures in air in order to bring about their thermal decomposition (lime stone), the removal of water from a hydrated compound (bauxite), or the removal of a volatile material from minerals and ores.

List the strong acids

Hydrochloric acid (HCl), sulfuric acid (H2SO4), nitric acid (HNO3)

Hydrogen ion vs. hydronium ion

Hydrogen ion is not viable as a structure in water, but we sometimes use it because it makes balancing easier. In reality, H1+ attaches immediately via a coordinate bond to H2O.

Auto-ionization/self ionization of water

H₂O (l) ➡️⬅️ H¹+ (aq) + OH¹− (aq). Kc = ([hydrogen ion][hydroxide ion])/([water]). As H2O is constant, Kw=[hydrogen ion][hydroxide ion]=1.0 x 10^-14 at 298 K. This expression is the ion product constant for water. In pure water, [hydrogen ion] = [hydroxide ion] = √(1.0 x 10^-14) = 1.0 x 10^-7. Above 60°F, Kw changes significantly enough for us to change it (but still neutral).

Starting point of a pH curve

Indicates the relative strength of acid or base

Salt hydrolysis

Ionic salts from neutralization dissociate completely in water, and pH of resulting solution is dependent on reaction of the salt with water. Hydrolysis is the ionization of water that results from reaction with an ionic salt. Salts of strong and weak acids and bases react in different ways with water and dictate the type and degree of the salt hydrolysis that results and its effect on the solution's pH.

Hydroxide and carbonate bases are - solutes, so...

Ionic, so dissociation. Just putting more space between the ions so they can do things.

Temperature dependence of Kw

Ionization of H2O is endothermic. Rise in temp will favor forward reaction and increase [H3O1+] and [OH1-]. Increase Kw and decrease pH. pH decreases but SOLUTION REMAINS NEUTRAL because there is an equal amount of OH1-. It's just that pH doesn't register this.

Relationship between Ka and Kb

KaKb=Kw. Stronger acid, larger Ka, weaker conjugate base, smaller Kb of conjugate base. Stronger base, larger Kb, weaker conjugate acid, smaller Ka of conjugate acid.

Electrophiles

Lewis acids. An electron-deficient species that can accept a lone pair from a nucleophile. Loves negative because it's more positive. Positive looking for negative, like metal in ligand.

Nucleophiles

Lewis bases. Loves positive because it has a lone pair. Negative looking for positive.

High conductivity means

More ions so lower or higher pH (more towards extremes of scale).

To determine equivalence point

Must complete titration first and then find it on the graph.

List the strong bases you need to know.

NaOH, KOH.

For a neutralization reaction, Gibbs free energy will be...

Negative

What bases don't go to completion in reaction with acids?

Nitrogen bases

Does the number of hydrogen ions produced predict the strength of an acid? If so, how so?

No

Does the Lewis theory falsify the Brønsted-Lowry theory?

No, it just expands it/extends our understanding. It is more general/inclusive than the other definitions.

Oxyacids

Nonmetal, hydrogen, and oxygen. Named for the nonmetal other than H and O.

Hydronium paper

Not really an indicator—color tells us the pH.

Weak base

Partially ionizes or dissociates, like NH3.

List the weak acids

Phosphoric acid (H3PO4), acetic/ethanoic acid (CH3COOH)

Ionization

Polar covalent + H2O —> ions in solution. Strong has one-sided arrow/complete ionization (no leftover acid), weak has two-sided arrow/partial ionization. With weak, usually a very small number of ions are produced—doesn't even go halfway. Monoprotic acids are ionized in one step, diprotic in 2, triprotic in 3, etc. The second and third steps can never go to 100% completion, but for a strong acid like H2SO4, it's pretty close, so we just kind of say completion. There still, though, has to be some solvent left. So...For weak, every step has partial ionization/double-sides arrow. For strong, first step always has single-sided arrow/complete ionization and rest have double-sides arrow/partial ionization. Bases follow same rules.

All acids are - solutes

Polar covalent—no ions. Must ionize.

SB + WAsalt

Polyatomic ➡️⬅️ H1+ + OH1-. Basic

pOH

Potential of hydroxide. pOH=-log[OH¹-]. 0-7, base (closer to 0 means stronger). 7-14, acid (closer to 14 means stronger). 7 is neutral. Not often used, don't really have pOH meters.

Brønsted-Lowry acid

Proton donor

Some bases don't...and do what instead?

React directly with aqueous acids...instead, they dissolve in water to create an alkaline solution, which neutralizes the acid. CaO2 does this. Many other bases are insoluble in water, too.

Acid + hydroxide base —>

Salt + water. The salt produced in this neutralization reaction is made of a cation from the base and an anion from the acid.

Point of inflection of pH curve

Sharp rise in pH at the equivalence point

Electrolyte

Solution containing ions.

Amphoteric

Species that behave both as an acid and a base.

Half-equivalence point

Stage of the titration at which half of the amount of weak acid has been neutralized and pKa=pH.

Strength vs. concentration

Strength is #ions in solution. Only applies when ions are in the solution- ionization or dissociation has occurred. Weak=small # ions=partial ionization or dissociation. Strong=large #ions=complete ionization or dissociation. Concentration is amount of solute in solution. Dilute=small amount solute went in. Concentrated=large amount solute went in. Can be applied to any solution. No matter how concentrated a weak acid is, it will always be weak.

What is the preferable way to titrate?

Strong acid, strong base. We rarely do weak and weak, because it gives sketchy calculations.

Talk about Ka and pKa for strong vs. weak acids

Strong: high Ka, low pKa. Weak: low Ka, high pKa.

Arrhenius acid

Substance that ionizes in water to produce hydrogen ions.

Properties of bases

Taste bitter, pH greater than 7.0, litmus is blue, phenolphthalein is pink, methyl orange is yellow, corrosive, form hydroxide ions in solution, are electrolytes, slippery feeling.

Properties of acids

Taste sour, pH less than 7.0, litmus is red, phenolphthalein is colorless, methyl orange is red, corrosive, form hydronium ions in solution, are electrolytes because they contain ions, can conduct electricity (usually don't use this), no distinctive feeling.

Pre-combustion methods to reduce sulfur emissions

Techniques used on fuels before their combustion. Mineral beneficiation (physical cleaning) involves crushing coal then floating it to decrease amount of sulfur and other impurities. These techniques can be very useful.

neutralization reaction

The combination of an acid and a base is well known as a neutralization reaction involving the combination of the hydrogen ions and hydroxide ions to produce water.

What is responsible for the physical and chemical properties of acids?

The hydronium ion. More H3O1+, more acidic. Registered with the pH scale. Characteristic ion.

Acid rain

The most prevalent form of acid deposition. Caused by increased industrialization and economic development in many parts of the world that have led to rapidly increasing emissions of nitrogen and sulfur oxides. These oxides are they cause. Rainwater is naturally acidic already due to presence of dissolved CO2, which forms H2CO3 (typically pH 5.6). Acid rain has a pH of less than 5.6. Caused mainly by SO2, NO, and NO2. These are products of natural occurrences (like volcanic eruptions and decomposition of vegetation) and man-made primary pollutants from combustion of fossil fuels containing high levels of sulfur impurities. Formation principally of HNO3 and H2SO4. Major global environmental problem.

Lewis theory and coordinate bonding

The presence of at least 1 lone pair of electrons allows a substance to act as a Lewis base, because it can use those electrons to form a coordinate bond. Transition elements have a partially occupied d subshell, so they can form complex ions with ligands that possess a lone pair of electrons. In this case, the metal atom or ion is acting as the Lewis acid and the ligand is the Lewis base. H2O, NH3, CN1-, Cl1-, and OH1- can act as Lewis Bases in forming complexes.

Binary acids

Two elements like HCl, HBr. Nonmetal and hydrogen. Named for their nonmetal.

Indicator

Typically a weak acid or base that displays a different color in acidic or alkaline solution. Many indicators in aqueous solutions behave as weak acids: Ka=([H1+][In1-])/[HIn]. Midpoint of the color change is observed when [HIn]=[In1-]. At this point, [H1+] = Ka and pH=pKa. The color change for most indicators takes place over a range of pH = pKa ±1. Color of indicator depends on pH of solution: if acidic, indicator exists as HIn. If basic, indicator exists as In1-. Indicator can also be a weak base, in which case ion would give base color and non-ion would give acid color. Most are weak acids: HIn ➡️⬅️ H1+ + In1-. Some are weak bases: BOH ➡️⬅️ B1+ + OH1-. When it goes to ions, it goes to a different color solution.

How do buffer solutions work?

Weak acid and conjugate base (or vice versa) are mixed in equimolar concentrations. Equilibrium position changes in accordance with Le Châtelier's: 1. Adding small amounts of a strong acid will increase hydronium conc, favoring reverse reaction and using extra hydronium to maintain pH. In contrast, small amounts of added strong base react with the hydronium to form water; forward reaction is favored, replenishing hydronium ions and maintaining pH.

All carbonates are

Weak and hard to dissolve.

Talk about Kb and pKb for strong vs. weak bases.

Weak: low Kb, high pKb. Strong: high Kb, low pKb.

How does a hydronium ion form?

When water forms a coordinate bond with a proton.

triprotic acid

an acid able to donate three protons per molecule

monoprotic acid

an acid that can donate only one proton (hydrogen ion) per molecule

diprotic acid

an acid that can donate two protons per molecule

Lewis acid

electron pair acceptor

Lewis base

electron pair donor

Carbonate base

m#(CO3)#. #s depend on oxidation numbers.

Hydroxide base

m(OH)#. # depends on the metal.

SA + WBsalt

metal(H2O)n ➡️⬅️ metal(OH)H2On-1 + H1+. Acidic

Strong acid + weak base —>

pH < 7

Strong acid + strong base —>

pH = 7

Weak acid + strong base —>

pH > 7

Equivalence point for weak base strong acid

pH<7.

Equivalence point for strong acid strong base

pH=7.0

Equivalence point for weak acid strong base

pH>7

Equivalence point for weak base weak acid

pH≈7.

Half equivalence point for weak base strong acid

pKb=pOH

Brønsted-Lowry base

proton acceptor

Acid + metal carbonate/metal hydrogencarbonate —>

salt + carbon dioxide + water.

Acid + metal —>

salt + hydrogen gas. Give off hydrogen gas at different rates according to the reactivity of the metal and the strength and concentration of the acid. Salt formed depends on the acid from which it was produced.

Acid deposition

the process by which acid-forming pollutants are deposited on the Earth's surface