Chem4A

Addition of base to a carbonic acid/bicarbonate buffer system

1. Added alkali perturbs equilibrium H2CO3 + OH- HCO3- + H2O 2. Some carbonic acid reacts with excess alkali, restoring the equilibrium Thus, most of the excess OH- is neutralised. In the presence of carbonic acid and bicarbonate, pH will increase less following addition of alkali than if acid had been added to pure water (buffering)

Bases

A base is any compound, X-, which reacts with H+ (in water, H3O+) acting as a proton acceptor. H+ +X-H—Xor,inwater: H3O+ + X- H—X + H2O

The pH scale

A convenient and widely used way of reporting small concentrations of hydrogen ions (H+) without resorting to scientific notation. In pure water, [H+] = 10-7 M. Equivalently, we report this as: pH = 7 Thus, pH is equal to the negative logarithm (to base 10) of the hydrogen ion concentration, or where [H+] is ALWAYS given in units of mol/l (M, molar) pH = -log10 [H+] pH below 7 means excess H+ = increased acidity Increasing concentration of H+ = adding more acid

Acid-Base Regulation

A pH buffer works chemically to minimise changes in the pH of a solution - The body uses pH buffers to prevent sudden changes in acidity

ionisation

Any process in which atoms become charged. When an atom has gained or lost one or more electrons. If an atom or molecule gains an electron, it is negatively charged; losing an electron makes it positive.

Minerals in living organisms

Not abundant (5%), yet important: -Bone: calcium hydroxyapatite: Ca, P, O, H -Various salts dissolved in water. Salts are composed of metals and electronegative non-metals • Na: Group I - 1 outer electron, readily lost: Na+->Na+ + 1 e- • Cl: Group VII - 7 outer electron, readily gains one more: Cl + 1 e- -> Cl- • Thus, Na and Cl react (violently!) to form NaCl = Na+Cl- • Other salts form similarly, with balancing of + and - charge, e.g., Ca -> Ca2+ + 2e- forms CaCl2, or CaO, for example

A medical example

Old-style antacids • Indigestion can be caused by too much HCl in the stomach • Indigestion tablets contain bases such as: - magnesium hydroxide: Mg(OH)2 - calcium carbonate: CaCO3 - aluminium hydroxide Al(OH)3

When hydrochloric acid (gas) dissolves in water, what happens?

The dissociated H+ ion (proton) attaches to a water molecule

dissociation

The temporary or reversible process in which a molecule or ion is broken down into smaller molecules or ions -The separation of ions that occurs when an ionic compound dissolves -The separating of a molecule into simpler molecules, atoms, radicals, or ions the disconnection or separation of something from something else or the state of being disconnected.

Dissociation constants

a specific type of equilibrium constant that measures the tendency of a larger object to dissociate reversibly into smaller components. Describes the strength of the bonds in question. This indicates the binding strength between the chemical components of a complex product. It is the reverse of the rate of association between the two reactants which would form this complex product. The equation is AB ⇋ A+B.

ph scale

measurement system used to indicate the concentration of hydrogen ions (H+) in solution; ranges from 0 to 14. O is most acidic + and 14 is most basic + or alkaline.

• pH

pH is a measure of the concentration of hydrogen ions (H+) in a solution. pH = -log [H+] • pH scale range: 1-14: Below 7: acid pH, e.g., pH 1 = 1 M H+ Above 7: basic pH, e.g., pH 14 = 1 M OH- pH = 7: neutral pH (E. g: pure water at 25 C) • Excess H+ and OH- in the two compounds combine to form water • The remaining components form a salt Acid + Base Salt + H2O acid base salt acid base salt

base strength

related to the equilibrium position of a reaction, and is therefore a thermodynamic property. strong: dissociates completely strong electrolyte. Willingness or ability to accept a proton

pKa

the ionization constant of a chemical compound. The pH at which the dug will exist as 50% ionized and 50% non-ionized. All drugs are salts of a weak acid or base. The pH at which the concentration of the salt equals the concentration of the acid. The quantities in square brackets symbolize the molar concentrations pKa=-log Ka

pOH

the negative of the common logarithm of the hydroxide ion concentration of a solution. The negative logarithm of the hydroxide ion concentration of a solution; a solution with a pOH above 7.0 is acidic, a solution with a pOH below 7.0 is basic, and a solution with a pOH of 7.0 is neutral. pOH = -log[OH-]

Acid/base

• An acid is a substance that produces H+ when dissolved in water (an acid donates H+) •A base is a substance that produces OH- ions when dissolved in water (a base accepts H+) Neutralisation • Reaction of an acid with a base/alkali is called a neutralisation reaction

Buffers

• Contain both a weak acid and its conjugate base, which are in equilibrium • The equilibrium is described by Ka or, equivalently, by pKa • A buffer solution with a pH at or near the pKa value for the buffer acid/base pair will have equal concentrations of the two components • Around this pH value, the solution resists pH changes from additions of excess acid or base

Base

• Forward reaction: when an acid, HA, dissociates to release a H+, it also givesA-, abase. • Reverse reaction: the base accepts a H+ and yields an acid • An acid and the base resulting from its dissociation are called a conjugate acid and base, respectively. Conjugate acid-base pair (Again, a chemical equilibrium). Acids and Bases are never alone

What is the relationship between acidity and alkalinity?

• Pure water self-dissociates: H2O H+ + OH- Keq = 1.8 × 10-16 M • Added acids react with OH- to make more water, until the concentrations once more satisfy equilibrium conditions: H2O H+ + OH- • Similarly, adding alkali will deplete H+. • In pure water (pH = 7), the equilibrium constant for self-dissociation dictates that: [H+] × [OH-] = 10-7 M × 10-7 M = 10-14 M2. • In dilute acid or base solution, this relationship still applies (because most of the solution is still water), but either acid or base is now in excess, and the other is consumed to restore equilibrium conditions: Excess acid: [H+] × [OH-] = 10-14 M2 Excess base: [H+] × [OH-] = 10-14 M2 pH above 7 means less H+ and correspondingly more OH- = increased alkalinity. Increasing concentration of OH- = adding more alkali. More H+ and less OH- less H+ and More OH-

We can define pOH in analogy to pH

• pH = - log10 [H+] • pOH = - log10 [OH-] (concentrations in M) - Less often used than pH, because pH and pOH are so closely related that only one scale is needed. • In dilute aqueous solution, [H+] × [OH-] = 10-14 M2. Equivalently: log [H+] + log [OH-] = -14 or: pH + pOH = 14

Importance of acids in Biology - examples

Lab: Changes in acidity in solutions of molecules, e.g., culture media for cells: - Microorganisms can change shape, mutate, become pathogenic, - Extremes of acidity can kill cells - Enzymes and other proteins can change or lose activity. Natural environment: - Too much acid in rain water can be toxic to plants - Soil acidity is one factor determining which plants grow - Ocean acidification is related to rising atmospheric CO2 levels, dissolves corals Health and food industries; Acidity must be regulated to ensure that: - Products have consistent well defined properties - Health problems for consumers are avoided - Regulatory requirements are met

salts

-A class of ionic compounds that can be formed during the reaction of an acid and a base. Acids and bases react with each other to form salts. -Compounds made of a metal and nonmetal that are formed when acids and bases react. -Ionic compounds (contain anions and cations), when neutral these have no acid/base characteristics (do not contain H+ or OH-), when it is an acid salt it contains cation, anion, and H. An intermediate between an acid and a neutral salt. - Salts are compounds made of a metal and nonmetal that are formed when acids and bases react. -Sodium chloride, sodium nitrate, and calcium sulfate are all salts.

Acid

-NOTHING TO DO WITH CONCENTRATION, how well a solution dissolves/ionizes/dissociates in water. -Strong donates protons easily & weak clings to their protons. -Classified by degree of ionization: strong acids lose their H+ ions easily and completely. Increasing polarity and decreasing bond energy. -Any compound that increases the number of hydronium ions when dissolved in water. • A strong acid completely dissociates into ions in water • Leads to a high concentration of H+, equal to the concentration of acid that was added. • A weak acid will only partially dissociate into ions in water • The resulting increase in H+ ions is much smaller than the concentration of the weak acid. E.g., Acetic (a.k.a. ethanoic) acid. • At equilibrium, only ≈ 1 in 250 acetic acid molecules dissociates.

Addition of acid to a carbonic acid/bicarbonate buffer system

1. Added acid perturbs equilibrium 2. Some bicarbonate reacts with excess acid, restoring the equilibrium Thus, most of the excess H+ is neutralised. In the presence of carbonic acid and bicarbonate, pH will drop less following addition of acid than if acid had been added to pure water (buffering)

Buffers:

A substance that consists of acid and base forms in a solution and that minimizes changes in pH when extraneous acids or bases are added to the solution. Used to regulate pH, so pH homeostasis can be maintained. Substances that function to prevent radical changes in pH. Act as reservoirs for H+. Donates H+ when H+ concentration falls and accepts when rises.Buffers resist changes in pH.

Base

A substance that decreases the hydrogen ion concentration in a solution. Any ionic compound that increases the number of hydroxide ions when dissolved in water. pH above 7, releases OH- ions when dissolved in water. Similarly, base strength describes the readiness of an alkali to release OH- ions. • A strong base completely dissociates into ions in water • Leads to a high concentration of OH-, equal to the concentration of total alkali that was added. E.g., Sodium hydroxide • A weak base will only partially dissociate into ions in water • The resulting increase in OH- ions is much smaller than the total concentration of base that was added. E.g., Ammonia.

Acids

An acid is any compound which dissociates in water to release H+ and something else, X-: H—XH+ + X- or, in water: H—X + H2OH3O+ + X- Examples: • Hydrochloric acid (X = Cl-) in stomach acid • Acetic acid (CH3-COOHCH3-COO- + H+) in vinegar

Weak acid

An acid that is only slightly ionized in aqueous solution. Weak Acids and pKa The strength of an acid, HA, can be determined by the equilibrium constant for its dissociation reaction, also called its acid dissociation constant, Ka • When Ka << 1, [HA] > [H][A-] at equilibrium, and HA is not significantly dissociated. Thus, HA is a weak acid when Ka < 1. •The lower the value of Ka. -Ka =[H+] [A-] [HA], the weaker the acid. • The value of Ka can also be represented as pKa. The larger the pKa, the weaker the acid.

ions

Charged atoms/ Charged particles/Atoms with a charge. Atom or group of atoms(polyatomic) with a positive or negative electrical charge due to a loss or gain of electrons. Ex: Na+, Cl- , NO3- electrically charged particles found both inside and outside a neuron; negative ions are found inside the cell membrane in a polarized neuron Electrically charged particles derived from the salt or other chemical dissolved in the water.

Bases

Compounds that reduce the concentration of hydrogen ions in a solution. Solutions that have more OH- than H+ ions. Proton acceptors

Conjugate acid-base pair

Consists of two substances related by the loss or gain of a single hydrogen ion. The conjugate base of an acid is formed when the acid donates a proton. In the equation, OH- is the conjugate base to the acid H2O, because H2O donates a hydrogen ion to form OH-, the conjugate base.

Buffers

Definition: An aqueous solution of a weak acid, which resists changes in pH upon addition of small amounts of acids or bases. Example: Conjugate acid-base pair H2O + When both conjugate acid and base are at equal concentrations, small additions of acid or alkali will be absorbed by the appropriate reaction partners, so that pH changes remain small. The solution is said to be "buffered" against pH change.

Ionic bonds are very stable in solid crystals but can dissociate in water

Dissolved ions are selectively enriched in, or excluded from, cells. ION | Extracellular | Cystolic Na+ | 145 |5-15 K+ | 5 | 140 Mg2+ |1-2 |0.5 Ca2+ | 1-2 |0.0001 Cl- |110 |5-15*

What is a chemical equilibrium?

In general, equilibrium reactions convert reactants (A, B, ...) into products (X, Y, ...), but the reaction also proceeds in the opposite direction, thus: A + B (+ ...) X + Y (+ ...) When reactants are mixed, products will be formed at a defined rate. The concentrations of reactants and products will change over time: reactants A, B, etc. are depleted, whereas products X, Y, etc., accumulate. As product molecules accumulate, some will convert back to reactants. Eventually, the rates of forward and reverse reactions will become equal, and the concentrations will no longer change. This is when when the reaction is said to have reached equilibrium.

When the reaction has reached equilibrium, there is no net change in the concentrations of any of the reactants or products.

It was found that, regardless of the starting concentrations of any of the reactants or products, once equilibrium is reached, the following quantity, called the equilibrium constant, Keq, remains fixed: where [ ]eq indicates molar (M, mol/l) concentrations at equilibrium. Note: Keq is a constant under a particular set of physical conditions, but can change when those conditions change. Keq = [X]eq [Y] eq ... = [products]eq [A]eq [B]eq ... [reactants]eq Keq allows us to quantify how much product can be obtained from any reaction AT EQUILIBRIUM (and under specified conditions) A+ B

Water can spontaneously dissociate

One molecule of water donates a H+ to another, and thus serves as an acid. In the reverse reaction, OH- accepts a proton, and thus serves as a base. In a narrower sense of the word, "base" (or "alkali") is used to designate a compound containing, specifically, hydroxide ions. -Self-dissociation of water is an example of an equilibrium reaction, which spontaneously proceeds in both directions (hence the double arrow) -Self-dissociation of pure water is extremely rapid, but happens to only a very small proportion of the water molecules present (hence the uneven lengths of the two arrows). The ions combine rapidly again to re-form water.

Acids and bases

Strong acids and strong bases dissociate completely into ions. Weak acids and weak bases dissociate only partially into ions. Acid: hydrogen containing compounds that release H+ ions when dissolved in water e.g. HC.

acid strength

Strong donates protons easily & weak clings to their protons. Classified by degree of ionization: strong acids lose their H+ ions easily and completely