CAPÍTULO VII ASPECTOS DIVERSOS
3. Ayuda mutua:
Equal numbers of molecules of different gases occupy the same volume in isothermal and isobaric conditions. At 0oC and 100 kPa, 22.71 L per mole and 24.79 L/mol for the other.
• Gather and process information to write the ionic equations to represent the ionisation of acids
• Define acids as proton donors and describe the ionisation of acids in water
Acids react with water in solution to form a solution containing hydronium ions and its conjugate base.
• Identify acids including acetic acid, citric acid (2-hydroxypropane-1,2,3-tricarboxylic acid), hydrochloric acid and sulfuric acid
o Acetic acid (CH3COOH) aka ethanoic acid – present in vinegar
o Citric acid (C6H8O7) – occurs in citrus fruit, also widely used as a
food additive for flavour or as a preservative
o Hydrochloric acid (HCl) – produced by stomach lining glands to break down food molecules, also made commercially to clean metals, brickwork, neutralising bases etc.
o Sulfuric acid (H2SO4) – synthetic acid manufactured to make fertilisers, synthetic fibres
etc.
Phosphoric acid (H3PO4) weak
Polyprotic acids have more than one ionisable hydrogen per formula unit. For others see book
• Describe the use of the pH scale in comparing acids and bases
• Identify pH as –log10[H+] and explain that a change in pH of 1 means a tenfold change in
[H+]
pH scale is a scale of measurement for hydrogen ion concentration. pH = –log10[H3O+]
This obeys the significant figure rule Water self ionises:
H2O + H2O H3O+ + OH-
Kw = [H3O+][OH-] = 1.00 x 10-14at 298K
• Describe acids and their solutions with the appropriate terms weak, strong, concentrated and dilute
• Describe the difference between a strong and weak acid in terms of an equilibrium between intact molecules and its ions
Note: In exams, define concentration and strength if used in question.
A strong acid is one in which all acid present in solution has ionised to hydrogen ions (no degrees of strength), no equilibrium is formed. A weak acid is one in which only some of acid molecules present in solution have ionised to form hydrogen ions, forming an equilibrium between intact molecules and ions. The fraction of molecules ionised is called the degree of ionisation
(concentration of H+/concentration of acid originally).
• Plan and perform a first-hand investigation to measure the pH of identical concentrations of strong and weak acids
Prac – Relative strength of acids
Aim: To compare the relative strengths of different acids using a variety of methods Method:
1). 50mL of 0.1M hydrochloric, acetic, oxalic, citric and sulfuric acid were prepared in labelled 250 mL beakers
2). A pH meter was used in each beaker and the reading recorded
3). A pH strip was placed into each beaker for a short period and its colour compared with a chart 4). A few drops of universal indicator were dropped into each beaker and the colour compared to a colour chart
Results: #####
The strongest was sulfuric acid. It is a strong acid and is diprotic, meaning that the concentration of H3O+ ions is twice the concentration of
hydrochloric acid. HCl is strong but monoprotic, meaning concentration is identical to HCl concentration. Oxalic acid is diprotic and citric is triprotic,
but both are weak and do not completely ionise. The tendency for their conjugate bases to re-bond with hydrogen ions limits the concentration of H3O+ in solution. Acetic is the weakest, being
monoprotic and have a low degree of ionisation. Oxalic – C2H2O4
Citric – C6H8O7
Acetic/ethanoic – CH3COOH
Safety:
HCl - corrosive, vapour can burn mouth, throat and eyes
Oxalic acid – corrosive to tissue, corrosive to respiratory tract if inhaled Wear safety glasses, goggles, use lower concentrations and smaller amounts
• Compare the relative strengths of equal concentrations of citric, acetic and hydrochloric acids and explain in terms of the degree of ionisation of their molecules
An acid is stronger than another if it has a higher degree of ionisation.
Citric is a triprotic acid, acetic is monoprotic. The degrees of ionisation are: HCl – 0.010/0.01 = 1
Citric acid – 2.74 x 10-3/0.01 = 0.274
Acetic acid – 4.17 x 10-4 / 0.01 = 0.0417
HCl is the strongest acid and has the highest degree of ionisation.
• Use available evidence to model the molecule nature of acids and simulate the ionisation of strong and weak acids
• Gather and process information to explain the use of acids as food additives Acids are added to food to:
o Improve taste – e.g. carbonic acid in soft drinks, acetic acid in vinegar
o Preserve food – increases acidity to point where bacteria can no longer survive e.g.
coating freshly cut fruit with citric acid
o Increase nutritional value – e.g. adding ascorbic acid (Vitamin C, an antioxidant)
• Identify examples of naturally occurring acids and bases and their chemical composition Acids:
o Ascorbic acid C6H8O6 – occurs widely in fruit and vegetables, essential to health
o Citric acid – found in citrus fruits
o Lactic acid CH3CH(OH)CO2H – produced by anaerobic respiration in cells, found in
muscle tissue and milk
o HCl – stomach to break down food Bases:
o Ammonia NH3 – result of decomposition of proteins, or anaerobic decay of organic
o Carbonates – e.g. calcium carbonate (limestone), magnesium carbonate
o Metallic oxides – e.g. Iron (III) oxide, copper oxide and titanium oxide, found in minerals. Metals are extracted from them.
• Calculate pH of strong acids given appropriate hydrogen ion concentrations • Outline the historical development of ideas about acids including those of:
o Lavoisier o Davy o Arrhenius o Lavoisier (1780)
• acids were substances that contained oxygen
• Disproved since some oxygen-containing compounds such as metallic oxides were basic, and distinctly acidic substances such as hydrochloric acid contained no oxygen • Wrong but stimulated research
o Davy (1815)
• suggested that acids were substances that contained replaceable hydrogen. • Bases were substances that reacted with acids to form salt and water.
• These definitions worked well for most of that century, but the definition made no attempt to interpret the properties, only classify the substances
o Arrhenius (1884)
• Interpreted acidic properties in terms of ionisation to form H+, and weak/strong in
terms of degrees of ionisation
• the conductivity of acid solutions and their reaction with many metals to form hydrogen gas evidenced that acidic solutions contained hydrogen ions
• acids were substances that ionised in solution to produce hydrogen ions • A base is a substance that in solution produced hydroxide ions
• He defined strong acids as those that ionised completely and weak as those that partially ionised. General equations:
HA(aq) H+(aq) + A-(aq)
XOH(aq) X+(aq) + OH-(aq)
Weaknesses were:
Does not take into account role of solvent in ionisation of acid (ionisation
results from reaction of acid with solvent)
Acid-base reactions can occur in solvents where there is no ionisation Not all acidic/basic substances (e.g. metallic oxides) ionised to produce
hydrogen/hydroxide ions
• Outline the Bronsted-Lowry theory of acids and bases
An acid is a proton donor, a base is a proton acceptor. Gives the broadest definition of acid/base theory (it means that acids must have hydrogen). This definition :
o Does not restrict bases to those which ionise to produce hydroxide ions, such as in the case of metal oxides and ammonia
o Explains how neutralisation reactions don’t require dissolution of ions into aqueous
solution e.g. NH3 + HCl in benzene (direct proton transfer)
o Exchange of proton relies on relative properties of both substances involved, accounting for the role of the solvent
o Shows that hydrolysis of salts to change pH were acid or base reactions
• Trace developments in understanding and describing acid/base reactions ?
• Describe the relationship between an acid and its conjugate base and a base and its conjugate acid
When an acid loses a proton, the resulting ion is called a conjugate base. If the acid is not strong, this conjugate base can re-take a proton to reform the acid, resulting in an equilibrium reaction. Vice versa for bases
• Identify a range of salts which form acidic, basic, or neutral solutions and explain their nature
A salt is an ionic compound containing a cation not H+ and an anion not O2- or OH-. In aqueous
solution, salts completely dissociate into ions.
Type Acidic Basic Neutral
Salt Ammonium nitrate
(NH4NO3) Sodium hydrogen sulfate (NaHSO4) Anything containing Al3+, Fe3+, HSO 4- or H2PO4- Sodium acetate (NaCH3COO) Potassium nitrite (KNO2) Sodium carbonate (Na2CO3) Anything containing F-, S2- etc. Sodium chloride Potassium nitrate (KNO3) Sodium sulfate (Na2SO4)
The pH of a salt solution depends on the nature of its ions, many cations/anions serve as acids or bases. Some generalisations:
o Neutral salts have anions which are the conjugate base of strong acids, and cations the
conjugate acid of strong bases, since their reaction to accept/give protons is negligible o Basic anions react with water to form hydroxide ions in solution. Reaction is equilibrium,
occurs to small extent since conjugate acid is stronger than water, and conjugate base is stronger than basic anion
o Acid anions contain hydrogen atoms to react with water to form hydronium ions, derived from polyprotic acids. The anion resulting from hydrolysis of polyprotic acids is
amphiprotic, and whether it is acidic or basic depends on the tendency for one hydrolysis reaction (proton donation or proton accept) to occur over the other
Some examples: (Basic anions) S2–(aq) + H
2O(l) ↔ HS–(aq) + OH–(aq)
F–(aq) + H
2O(l) ↔ HF(aq) + OH–(aq)
(Acidic cations)
[Fe(H2O)6]3+(aq) + H2O(l) ↔ [Fe(OH)(H2O)5]2+(aq) + H3O+(aq)
• Perform an investigation to identify the pH of a range of salt solutions
Self explanatory. You could use NaCl + KOH for neutral, sodium bicarbonate (NaHCO3) and
Risk analysis:
Hazard Risk Control
Ammonium chloride Released vapour causes coughing, shortness of breath
Wear safety goggles Use small amounts to minimise vapour • Identify conjugate acid/base pairs
Acid Base Conjugate base Conjugate acid
HCl Cl- H2SO4 HSO4-,(and SO42-?) HNO3 NO3- NH4+ NH3 OH- H 2O CN- HCN CO32- HCO3- CH3COO- CH3OOH
• Identify amphiprotic substances and construct equations to describe their behaviour in acidic and basic solutions
Note: Hydrolysis is when a substance reacts with water
Amphiprotic substances can act as both a proton donor and proton acceptor. They react to both accept protons and donate protons. Their behaviour changes whether in aqueous solution or alkaline/acid solution.
e.g. HCO3- (hydrogen carbonate)
In aqueous solution:
HCO3-(aq) + H2O(l) H3O+(aq) + CO32-(aq)
HCO3-(aq) + H2O(l) H2CO3(aq) + OH-(aq)
In basic/acidic solution:
HCO3-(aq or s) + OH- H2O(l) + CO32-(aq)
HCO3-(aq or s) + H+(aq) H2CO3(aq)
Reactions go to completion since products cannot perform the reverse reaction (???). This also applies to HSO3- (hydrogen sulfite) and HSO4-. Water is amphiprotic.
• Identify neutralisation as a proton transfer reaction which is exothermic
Neutralisation reactions are proton transfer reactions, and involve the reaction between an acid and a base. They are exothermic and thus have a negative enthalpy change. The net ionic reaction in Arrhenius theory is:
OH-
(aq) + H+(aq) H2O(l)
Acids and bases not fitting in Arrhenius theory do not necessarily produce salt and water during neutralisation reactions e.g. neutralisation of ammonia. In LB theory, an acid and base react to form conjugate base and conjugate acid. The acid gives a proton to the base. Reactions between strong acids and bases form very weak conjugate acids/bases and go virtually to completion since the back reaction has almost no tendency to occur. Otherwise, reactions are equilibria.
• Describe the correct technique for conducting titrations and preparation of standard solutions
Volumetric analysis is a form of chemical analysis where the concentration of a substance is determined.
Determining the composition of a solution require titration against another solution of known concentration, called the standard solution. The substance dissolved is a primary standard. Equipment:
A primary standard:
o Must be obtainable in very pure form and have known formula
o Should not alter weight unintentionally during preparation/titration e.g. absorbing moisture from air
o Have a reasonably high formula mass to minimise weighting errors
o Purified by drying in oven and cooling in dessicator to eliminate moisture and prevent its absorption
e.g. oxalic acid, sodium carbonate Use of equipment:
o Pipette – solution to be used is first drawn in above mark, then solution let out until meniscus at mark, solution let out through gravity with tip against wall of container o Burette – first, rinse with portion of solution to be dispensed, overfilled then excess
allowed to run out Preparation:
o Accurately measure mass of primary standard e.g. electronic beam balance o Rinse a volumetric flask and beaker with distilled water
o Pour the primary standard into beaker and dissolve with distilled water, less than intended volume of final solution
o Pour into volumetric flask, and repeat a few more times o Use a pipette to add final few drops to complete solution Titration curves:
Strong acid strong base:
Equivalence point at pH 7, steep curve. Indicator used should have colour endpoint near equivalence point. Using indicator changing during equivalence point is inaccurate, too difficult to tell exact colour shade needed
Strong base/weak acid e.g. NaOH, acetic acid
Equivalence point in basic range since salt formed is basic, the anion is a conjugate base of weak acid and thus a weak base. Equilibrium reaction occurs, weak base reacts to reform conjugate acid, thus decreasing acidity since less H3O+
Strong acid/weak base
Similar to above with equivalence point acid. Special case when CO2 formed during reaction e.g.
HCl + Na2CO3, CO2 forms carbonic acid.
Weak/weak
Not good since gradient around equivalence point is quite shallow, big volume difference between indicator endpoints and equivalence points, needs to fit indicator very well or use one which changes during equivalence, hard to distinguish.
• Perform a first-hand investigation and solve problems using titrations and including the preparation of standard solutions, and use available evidence to quantitatively and qualitatively describe the reaction between selected acids and bases
Prac – Titration
Aim: To standardise HCl and NaOH solutions using titration Method:
Preparing standard:
1). A clean 250mL beaker was placed on an electronic balance, zeroed, and had 2.650g of Na2CO3
added using a spatula
2). Approx 100 mL of distilled water was added to the beaker, and solution stirred using stirring rod 3). A 250mL volumetric flask was rinsed with distilled water*
4). The Na2CO3 solution was poured into the volumetric flask, and step 2 repeated
5). The stirring rod/beaker were thoroughly washed using a wash bottle, the runoff dripping into the vol. flask
6). Using a 25mL pipette, the volumetric flask was filled to the 250mL mark 7). A stopper was placed on the flask and contents swirled to mix
8). The pipette was rinsed with the unknown HCl*
9). 10 mL of unknown concentration HCl was poured using a pipette into a 50mL beaker washed with distilled water*, and a few drops of methyl orange added
10). A burette was washed and filled with the Na2CO3 standard solution
11). The standard solution was quickly drained into the beaker to find an approximate end point 12). Steps 8-10 were repeated 3 times but more accurately
*to ensure no cross-contamination
See book for calculations
• Qualitatively describe the effect of buffers with reference to a specific example in a natural system
Buffer solutions resist changes in pH. It contains comparable amounts of a weak acid/base and its conjugate base/acid. Take for example an acetic acid (CH3COOH) and sodium acetate
(NaCH3COO) system (acidic buffer).
Addition of sodium acetate would increase the concentration of CH3COO – ions on the right. The
equilibrium shifts to the left, but due to the unchanged concentration of H3O+ ions, it stays enough
to the right for dissociation to cause a net increase in CH3COO – ions. This net increase enhances
buffering capacity. When hydronium ions are added to solution, the equation will shift to the left according to Le Chatelier’s principle to reduce the concentration of H3O+ ions. When hydroxide
ions are added, the CH3COOH will react to form water and CH3COO-, reducing OH- concentration.
In both cases, pH change is reduced. Buffer in natural system
Carbonic acid / bicarbonate ion buffer system in mammalian blood: H2CO3(aq) + H2O(l) HCO3-(aq) + H3O+(aq)
Maintains the blood at around 7.4 for optimum function, too high/low can result in death. • Analyse … to assess the use of neutralisation reactions as a safety measure or to
minimise damage in accidents or chemical spills
o Many acids and bases are corrosive, can damage materials if spilt Neutralisation reactions can:
o Reduce or nullify corrosive properties of spill, minimising damage
o Utilise common, cheap, safely handled/stored materials and produces relatively harmless
products e.g. sodium bicarbonate NaHCO3
Sodium bicarbonate is amphiprotic, so it can be used for both acidic and basic spills:
) ( 2 ) ( 2 3 ) ( 3 ) (aq HCO aq CO aq H Ol OH− + − → − + ) ( ) ( 2 ) ( 2 ) ( 3 ) (aq NaHCO aq CO g H Ol Na aq H+ + → + + +
Vinegar – commonly used in cooking, contains acetic acid:
) ( 2 ) ( 3 ) ( 3 ) (aq CH COOH aq CH COO aq H Ol OH− + → − +
o Degree of reaction can be controlled by using different amounts of neutralising substance, excessive amounts are wasteful and some
Overall, it is a useful, convenient and safe technique if used appropriately
• Describe the differences between the alkanol and alkanoic acid functional groups in carbon compounds
• Explain the difference in melting point and boiling point caused by straight-chained alkanoic acid and straight-chained primary alkanol structures
Strong intermolecular forces – higher BP/MP compared to similar mole mass (roughly similar dispersion forces)
-OH
carboxylic acid group
Forms One hydrogen bond between molecules
Two polar bonds between molecules (C- O polar and C – O – H hydrogen bond),
higher boil/melt point None ionised, neutral Small no. ionised, acidic
• Identify the IUPAC nomenclature for describing the esters produced by reactions of straight-chained alkanoic acids from CI to C8 and straight-chained primary alkanols from C1 to C8
Alkyl (e.g. methyl) alkanoate (e.g. formate, acetate, propanoate etc.)
• Identify esterification as the reaction between an acid and an alkanol and describe, using equations, examples of esterification
Esters – carboxylic acids combined with alcohols, equilibrium reaction Standard equation:
o H2O molecule released (H of alkanol OH of acid)
o OR’ (R’ dummy variable) of alkanol C of acid
e.g. 3 ( ) 3 ( ) 2 4 3 3(l) 2 (l) SO H l l CH OH CH COOCH H O COOH CH + ← → +
(acetic acid + methanol)
) ( 2 ) ( 3 2 ) ( 2 3 ) ( 2 4 l l SO H l l CH CH OH HCOOCH CH H O HCOOH + ← → +
(formic acid + ethanol)
• Describe the purpose of using acid in esterification for catalysis
o The acid acts as a dehydration agent, removing water from the reaction (Le chatelier
argument), thus increasing yield
o Acts as a catalyst, lowering activation energy to speed reaction o Only small amounts of acid required
o Most common is sulfuric, others include tosic, scandium (III) triflate • Explain the need for refluxing during esterification
Reflux – the backflow of reactants into the reaction vessel o Reactants and ester products can be volatile
o Reaction is carried out at high temperatures to speed reaction, causing evaporation of alcohol and ester and thus loss of reactants/products
o Vapour is also dangerous as it is flammable and toxic
o To avoid loss and prevent diffusion, a cooled condenser is placed over the reaction vessel, covering it
o The vapour condenses here and runs back into the reaction vessel, which Increases yield and saves on resources
Allows the reaction to be carried at higher temperatures (faster) Prevents flammable gases from escaping
• Outlines some examples of the occurrence, production and uses of esters
• Identify and describe the use of esters as flavours and perfumes in processed foods and cosmetics
Esters occur naturally, and are identified as fragrances and flavours in fruit and flowers e.g. orange