Chem Chapter 19
lewis acids and bases
- Gilbert Lewis did work on bonding that led to a new concept of acids and bases - according to Lewis, an acid accepts a pair of electrons and a base donates a pair of electrons - this definition is more general than those offered by Arrhenius or by Bronsted and Lowry - the lewis definitions include all the Bronsted-Lowry acids and bases - ex. the reaction of H+ + OH- - the hydrogen ion donates itself to the hydroxide ion - the hydroxide ion can bond to the hydrogen ion because it has an unshared pair of electrons. the hydrogen ion accepts the pair of electrons - ex. reaction when ammonia dissolves in water
ion product constant for water
- Kw - the product of the concentrations of the hydrogen ions and the hydroxide ions 1.0 x 10⁻¹⁴
how buffers work
- a buffer solution is better able to resist drastic changes in pH than is pure water - the reason is fairly simple - a buffer solution contains one component that can react with hydrogen ions (a hydrogen-ion acceptor) and another component that can react with hydroxide ions (a hydrogen-ion donor) - these components act as reservoirs of neutralizing power that can be tapped when either hydrogen ions or hydroxide ions are added to the solution - the ethanoic acid-ethanoate ion buffer can be used to show this - when the acid is added to the buffer, the ions act as a hydrogen-ion "sponge" - as the ethanoate ions react with the hydrogen ions, they form ethanoic acid - the weak acid does not ionize extensively in water, so the change in pH is very slight - when hydroxide ions are added to the buffer, the ethanoic acid and the hydroxide ions react to produce water and the ethanoate ion - it is not a strong enough base to accept hydrogen ions from water to a great extent - therefore, the reverse reaction is minimal and the change in pH is very slight
buffer
- a solution in which the pH remains fairly constant when small amounts of acid or base are added - a buffer is a solution of a weak acid and one of its salts or a solution of a weak base and one of its salts
conjugate acids and bases
- all gases become less soluble in water as the temperature rises - thus, when the temperature of an aqueous solution of ammonia is increased, ammonia gas is released - this release acts as a stress on the system - in response to this stress, NH4+ reacts with OH- to form more NH3 and H20 - in the reverse reaction, ammonium ions donate hydrogen ions to hydroxide ions - thus, NH4+ acts as a bronsted-lowry acid, and OH- acs as a bronsted-lowry base - in essence, the reversible reaction of ammonia and water has two acids and two bases - in the equation, the products of the forward reaction are distinguished from the reactants by the use of the adjective conjugate - comes from Latin word meaning "join together"
acid-base indicators
- an indicator is often used for initial pH measurements and for samples with small volumes - an indicator is an acid or a base that dissociates in a known pH range - indicators work because their acid form and base form have different colors in solution - HIn - the acid form of the indicator (HIn) is dominant at low pH and high [H+] - the base form of the indicator (In-) is dominant at high pH and high [OH-] - the change from dominating acid form to dominating base form occurs within a narrow range of about two pH units - within this range, the color of the solution is a mixture of the colors of the acid and the base forms - if you know the pH range over which this color change occurs, you can make a rough estimate of the pH of a solution - at all pH values above this range, you would only see the color of the vase form - for a more precise estimate of the solution's pH, you could repeat the test with indicators that have different pH ranges for their color change - indicators have certain properties that limit their usefulness - the pH values of indicators are usually given for 25°C - at other temperatures, an indicator may change color at a different pH - if the solution being tested is not colorless, the color of the indicator may be misleading - dissolved salts in a solution may affect an indicator's dissocation - using indicator strips can help overcome these problems - an indciator strip is a piece of paper or plastic that has been soaked in an indicator, and then dried - the paper is dipped into an unknown solution - the color that results is compared with a color chart to measure the pH - some indicator paper has absorbed mulltiple indicators - the colors that result will give you a wide range of pH values
arrhenius base
- bitter taste - the taste is a dangerous property to test - the slippery feel of soap - will cause an indicator to change color - also form aqueous solutions that are strong or weak electrolytes
the capacity of a buffer
- buffer solutions have their limits - as acid is added to an ethanoate buffer, eventually no more ethanoate ions will be present to accept the hydrogen ions - at that point, the buffer can no longer control the pH - the ethanoate buffer also becomes ineffective when too much base is added - in that case, no more ethanoic acid molecules are present to donate hydrogen ions - adding too much acid or base will exceed the buffer capacity of a solution - two common buffer systems help maintain optimal human blood pH
conjugate acid-base pair
- conjugate acids are always paired with a base, and conjugate bases are always paired with an acid - the pair consists of two ions or molecules related by the loss or gain of one hydrogen ion - the ammonia molecule and the ammonium ion are a conjugate acid-base pair - the water molecule and the hydroxide ion are also a conjugate acid-base pair - the dissociation of hydrogen chloride in water provides another example of conjugate acids and bases
strong base
- dissociates completely into metal ions and hydroxide ions in aqueous solution - some strong bases, such as calcium hydroxide and magnesium hydroxide, are not very soluble in water - the small amounts of these bases that dissolve in water dissociate completely
arrhenius acid
- distinct properties - acids give foods a tart or sour taste (lemons) - aqueous solutions of acids are strong or weak electrolytes (can conduct electricity) - electrolyte in a car battery - cause certain chemical dyes (indicators) to change color - many metals react with aqueous solutions of acids to produce hydrogen gas
calculating pH from [H+]
- expressing [H+] in scientific notation can make it easier to calculate pH - when the coefficient is 1, the pH of the solution equals the exponent with the signed changed from a minus to a plus - when the coefficient is more than 1, you need a calculator
the pH oconcept
- expressing hydrogen-ion concentration in molarity is not practical - a more widely used system for expressing [H+] is the pH scale, proposed in 19009 by danish scientist Soren Sorensen - ranges from 0 to 14
concentration of ions in aqueous solution
- for any aqueous solutions, the product of the hydrogen-ion concentration and the hydroxide-ion concentration equals 1.0x10⁻¹⁴ - [H+] x [OH-] = 1.0 x 10⁻¹⁴ - this equation is true for al dilute aqueous solutions at 25°C - when susbtances are added to water, the concentrations of H+ and OH- may change - however, the product of [H+] and [OH-] does not change
calculating [H+] from pH
- if the pH is an integer, it is easy to find the value - however, most pH values are not whole numbers - you need to calculate using the antilog (10ⁿ) function to get an accurate value for the hydrogen-ion concentration
calculating pH from [OH-]
- if you know the [OH-], you can find its pH - recall that the ion-product constant for water defines the relationship between [H+] and [OH-] - therefore, you can use the ion-product constant for water to determine [H+] for a known [OH-], then calculate the pH
strong acid
- in general, they are completely ionized in aqueous solution - ex. hydrochloric acid, sulfuric acid
measuring pH
- in many situations, knowing the pH is useful - a custodian might need to maintain the correct acid-base balance in a swimming pool - a gardener may want to know if a certain plant will thrive in a yard - a doctor might need to diagnose a medical condition - either acid-base indicators or pH meters can be used to measure pH
acid dissociation constant
- in strong acids, the [H3O+] is high in aqueous solution - weak acids remain largely undissociated - for dilute aqueous solutions, the concentration of water is a constant - this constant can be combined with Keq to give an acid dissociation constant - ratio of the concentration of the dissociated form of an acid to the concentration of the undissociated form - the dissociated form includes both H3O+ and the anion - the Ka reflects the fraction of an acid that is ionized - for this reason, dissociation constants are sometimes called ionization constants - if the degree of dissociation or ionization of the acid in a solution is small, the value of the dissociation constant will be small - weak acids will have small Ka values - the stronger an acid is, the larger its Ka value will be - some of the acids have more than one dissociation constant because they have more than one ionizable hydrogen - ex. diprotic - the Ka decreases from the first ionization to the second
weak acid
- ionizes only slightly in aqueous solution - ex. ethanoic acid
hydrogen-containing compounds not acids
- methane (CH4) - the four hydrogen atoms in methane are attached to the central carbon by weakly polar C-H bonds - thus, methane has no ionizable hydrogens and is not an acid - ethanoic acid, which is commonly called acetic acid, is an example of a molecule that contains both hydrogens that do not ionize and a hydrogen that does ionize - although its molecules contain four hydrogens, ethanoic acid is a monoprotic acid - the three hydrogens attached to a carbon atom are in weakly polar bonds - they do not ionize - only the hydrogen bonded to the highly electronegative oxygen can be ionized - for complex acids, you need to look at the structural formula to recognize which hydrogens can be ionized
pH
- negative logarithm of the hydrogen-ion concentration - may be represented mathematically using the following equation: pH = -log[H+]
salts
- not only table salt, NaCl - ionic compounds consisting of an anion from an acid and a cation from a base - a reaction between an acid and base will go to completion when the solutions contain equal numbers of hydrogen ions and hydroxide ions - the balanced equation provides the correct ratio of acid to base
[H+] of a solution is less than 1x10⁻⁷
- pH is greater than 7.0 - basic
[H+] of a solution is greater than 1x10⁻⁷
- pH is less than 7.0 - acidic
titration
- process of adding a measured amount of a solution of known concentration to a solution of unknown concentration 1. a measured volume of an acid solution of unknown concentration is added to a flask 2. several drops of an indicator are added to the solution while the flask is gently swirled 3. measured volumes of a base of known concentration are mixed into the acid until the indicator just barely changes color
arrhenius
- proposed a new way of defining and thinking about acids and bases - acids are hydrogen-containing compounds that ionize to yield hydrogen ions in aqueous solution - bases are compounds that ionize to yield hydroxide ions in aqueous solution
weak base
- reacts with water to form the conjugate acid of the base and hydroxide ions - for a weak base, the amount of dissociation is relatively small - ammonia is an example - while equilibrium is established, only about 1 percent of the ammonia is present as NH4+ - this ion is the conjugate acid of NH2 - the concentrations of NH4+ and OH- are low and equal - the equilibrium-constant expression for the dissociation of ammonia in water is as follows Keq=[NH4+][OH-]/[NH3][H20] - the concentration of water is constant in dilute solutions - this constant can be combined with the Keq for ammonia to give a Kb for ammonia Keq*[H2O] = Kb =[NH4+][OH-]/[NH3]
salt hydrolysis reaction
- recall that a salt is one of the products of a neutralization reaction - a salt consists of an anion from an acid and a cation from a base - the solutions of many salts are neutral - salts that form neutral solutions include sodium chloride and potassium sulfate - some salts form acidic or basic solutions - the pH at the equivaence point for the weak acid-strong base titration is basic - for a strong acid-strong base titration, the pH at the equivalence point is neutral - the difference in pH exists because hydrolysis occurs with some salts in solution - sodium ethanoate is the salt of a weak acid and a strong base - in solution, the salt is completely ionized - the ethanoate ion is a bronsted-lowry base, which means it is a hydrogen-ion acceptor - ir reacts with water to form ethanoic acid and hydroxide ions - at equilibrium, the reactants are facored
arrhenius bases
- sodium hydroxide is an ionic solid - it dissociates into sodium ions and hydroxide ions in aqueous solution - extremely caustic - this property is the reason that sodium hydroxide is a major component of products that are used to clean clogged drains - potassium hydroxide is another ionic solid - it dissociates to produce potassium ions and hydroxide ions in aqueous solution - sodium and potassium are group 1A elements - alkali metals - react violently with water - the products of these reactions are aqueous solutions of a hydroxide and hydrogen gas - very soluble in water - concentrated solutions of these compounds is easy - the solutions would have the typically bitter taste and slippery feel of a base, but the solutions are extremely caustic to the skin - calcium hydroxide and magnesium hydroxide are compounds of group 2A metals - these compounds are not very soluble in water - their solutions are always very dilute, even when saturated - a saturated solution of calcium hydroxide has only 0.165g/100g water - magnesium hydroxide is even less soluble than calcium hydroxide with only 0.0009g/100g water
concentration vs. strenght
- sometimes people confuse the concepts - the words concentrated and dilute indicate how much of an acid or base is dissolved in solution - these terms refer to the number of moles of the acid or base in a given volume - the words strong and weak refer to the extent of ionization or dissociation of an acid or base
buffers intro
- suppose you add 10 mL of 0.10M sodium hydroxide to 1L of pure water - the pH wil increase about 4 pH units - from 7.0 to 11.0 - this change is a relatively large increase in pH - now consider a solution containing 0.20M each of ethanoic acid and sodium ethanoate - this solution has a pH of 4.76 - if you add 10mL of 0.10M sodium hydroxide to 1L of this solution, the pH increases 0.01 units - if 10mL of acid had been added instead of the base, the amount of change in pH would have also been small
acid-base reactions
- suppose you mix a solution of a strong acid with a strong base - in general, acids and bases react to produce a salt and water - the complete reaction of a strong acid and a strong base produces a neutral solution
salt hydrolysis
- the cations or anions of a dissociated salt remove hydrogen ions from, or donate hydrogen ions to, water - salts that produce acidic solutions have positive ions that release hydrogen ions to water - salts that produce basic solutions have negative ions that attract hydrogen ions from water
conjugate acid
- the ion or molecule formed when a base gains a hydrogen ion - ex. NH4+ is the conjugate acid of the base NH3
conjugate base
- the ion or molecule that remains after an acid loses a hydrogen ion - ex. OH- is the conjugate base of the acid H20
ion-product constant for water
- the ionization of water is a reversible reaction, so Le Châtelier's principle applies - adding either hydrogen ions or hydroxide ions to an aqueous solution is a stress to the system - in response, the equilibrium will shift toward the formation of water - the concentration of the other ion will decrease - in any aqueous solution when [H+] increases, [OH-] decreases. vice versa
equivalence point
- the point at which neutralization occurs - that indicator that is chosen for a titration must change color at or near the pH of the equivalence point
end point
- the point at which the indicator changes color
base dissociation constant
- the ratio of the concentration of the conjugate acid times the concentration of the hydroxide ion to the concentration of the base - the general form of the expression for the base dissociation constant is Kb=[conjugate acid][OH-]/[base] - you can use this equation to calculate the Kb of a weak base - you need to know the initial concentration of the base and the concentration of hydroxide ions at equilibrium - if you know the pH, you can calculate [H+] and the corresponding [OH-] - the magnitude of Kb indicates the ability of a weak base to compete with the very strong base OH- for hydrogen ions - bases such as ammonia are weak realtive to the hydroxide ion, the Kb for such a base is usually small - the smaller the value of Kb, the weaker the base
self-ionization of water
- the reaction in which water molecules produce ions - can be written as a simple dissociation - in water or in an aqueous solution, hydrogen ions are always joined to water molecules as hydronium ions - yet chemists may still refer to these ions as hydrogen ions or even protons - the self-ionization of water occurs to a very small extent
standard solution
- the solution of known concentration - you can use a similar procedure to find the concentration of a base using a standard acid
Bronsted-Lowry acid
- the two scientists were working independently - each chemist proposed the same definition of acids and bases - according to the bronsted-lowry theory, an acid is a hydrogen-ion donor and a base is a hydrogen-ion acceptor - this theory includes all the acids and bases that Arrhenius defined - it also includes some compounds that Arrhenius did not classify as bases - you can use this to understand why ammonia is a base - ammonia gas is very soluble in water. when ammonia dissolves in water, hydrogen ions are transferred from water to ammonia to form ammonium ions and hydroxide ions - ammonia is a bronsted-lowry base because it accepts hydrogen ions - water is a bronsted-lowry acid because it donates hydrogen ions
calculating dissociation constants
- to calculate it of a weak acid, you need to know the initial molar concentration of the acid and the [H+] of the solution at equilibrium - you can use these data to find the equilibrium concentrations of the acid and the ions - these values are then substituted into the expression for Ka - in general, you can find the Ka of an acid in water by substituting the equilibrium concentrations of the acid [HA] from the dissociation of the acid [A-] and the hydrogen ion [H+] below Ka=[H+][A-]/[HA]
arrhenius acids
- vary in number of hydrogens they contain that can form hydrogen ions - nitric acid has one ionizable hydrogen, so nitric acid is classified as a monoprotic acid - prefix+protic - protic=proton - not all compounds that contain hydrogen are acids - also, some hydrogens in an acid may not form hydrogen ions - only a hydrogen that is bonded to a very electronegative element can be released as an ion - recall that such bonds are highly polar - when a compound that contains such bonds dissolves in water, it releases hydrogen ions - in an aquesous solution, hydrogen ions are not present, they are joined to water molecules as hydronium ions
amphoteric substances
- water appears in both the list of acids and the list of bases - sometimes water accepts a hydrogen ion, sometimes it donates one - how water behaves depends on the other reactant - a substance that can act as either an acid or a base is said to be amphoteric - water is this - in the reaction with hydrochloric acid, water accepts a proton and is therefore a base - in the reaction with ammonia, water donates a proton and is therefore an acid
hydrogen ions from water
- water molecules are highly polar and are in constant motion, even at room temperature - on occasion, the collisions between water molecules are energetic enough for a reaction to occur - when this happens, a hydrogen ion is transferred from one water molecule to another - a water molecule that gains a hydrogen ion becomes a hydronium ion - a water molecule that loses a hydrogen ion become a hydroxide ion
basic solution
- when sodium hydroxide dissolves in water, it forms hydroxide ions in solution - in such a solution, they hydrogen-ion concentration is less than the hydroxide-ion concentration - remember, the hydrogen ions are present from the self-ionization of water - a basic solution is one in which [H+] is less than [OH-] - basic solutions are also known as alkaline solutions
acidic solution
- when some substances dissolve in water, they release hydrogen ions - in hydrochloric acid, the hydrogen-ion concentration is greater than the hydroxide-ion concentration - a solution in which [H+] is greater than [OH-] is an acidic solution - in acidic solutions the [H+] is greater than 1x10⁻⁷
pH meters
- your chemistry laboratory probably has a pH - a pH meter is used to make rapid, continuous measurements of pH - the measurements are typically accurate to within 0.01pH unit of the true pH - if the pH meter is connected to a computer or chart recorder, the user will have a record of the pH changes - a pH meter can be easier to use than liquid indicators or indicator strips - hospitals use pH meters to find small but meaningful changes in the pH of blood and other body fluids - sewage, industrial wastes, and soil pH are also easily monitored with a pH meter - the color and cloudiness of the solution do not affect the accuracy
concentration of pure water at 25°C
1x10⁻⁷M number of H+ ions and OH- ions are equal in pure water
acid-base definitions
Acid: arrhenius - H+ producer bronsted-lowry - H+ donor lewis - electron-pair acceptor Base: arrhenius - OH- producer bronsted-lowry - H+ acceptor lewis - electron-pair donor
hydrogen ions
H+
hydronium ion
H3O+ the ion that forms when a water molecule gains a hydrogen ion
equilibrium constant vs. acid dissociation constant
Keq = [H3O+][CH3COO-]/[CH3COOH][H2O] Keq*[H2O] = Ka = [H3O+][CH3COO-]/[CH3COOH]
hydroxide ions
OH-
ionizable
a hydrogen atom that can for a hydrogen ion
hydrolysis
a hydrogen ion is split off a water molecule -lysis = to separate or loosen strong acid + strong base → neutral solution strong acid + weak base → acidic solution weak acid + strong base → basic solution
lewis acid
a substance that can accept a pair of eelctrons to form a covalent bond
lewis base
a substance that can donate a pair of electrons to form a covalent bond
neutralization reaction
acid + base → salt + water
strong and weak acids and bases
acids and bases are classified as either strong or weak based on the degree to which they ionize in water
neutral solution
any aqueous solution in which [H⁺] and [OH⁻] are equal
caustic substance
can burn or eat away materials with which it comes in contact
neutralization
occurs when the number of moles of hydrogen ions is equal to the number of moles of hydroxide ions
buffer capacity
the amount of acid or base that can be added to a buffer solution before a significant change in pH occurs
equivalent
two things that are equal
