Module 6 Acid Base Reactions PDF

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Detailed notes on module 6, HSC Chemistry...

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Module 6: Acids and Bases

Properties of Acids and Bases IQ1: What is an acid and what is a base?

 Investigate the correct IUPAC nomenclature and properties of common inorganic acids and bases Acids are compounds which form hydrogen ions (H+) in aqueous solution. The hydrogen ion attaches to a water molecule and forms hydronium ions (H3O+). Acids are substances that dissociate and ionise in water, forming hydrogen ions. When reacted with bases, they will undergo neutralisation reactions and react with some metals to form hydrogen gas. E.g. HNO3 and CH3COOH, C6H8O7 (citric acid) Properties of Acids o o o o o o

Sour taste Conduct electricity in solution (mostly water soluble) Turn blue litmus red Corrosive React with bases to neutralise them They have a pH below 7

Bases are compounds which form hydroxide ions (OH-) in aqueous solution. Bases dissociate in water to produce hydroxide ions. E.g. NaOH, KOH. Alkalis are a term used to describe water-soluble bases. When a solution has a pH greater than 7.0, it is said to be alkaline (basic). Properties of Bases o o o o o o

Bitter taste Soapy feel in aqueous solution Conduct electricity in solution (not all bases are soluble) Turn red litmus blue Corrosive React with acids to neutralise them

Three main types are: metal hydroxides, metal oxides and metal carbonates/hydrogen carbonate with many exceptions i.e. ammonia (NH3). Oxyanions: anions containing one or more oxygen atoms bonded to another element such as sulphate and carbonate. The central atom exists in a more reduced state, bearing the suffix ‘-ite’. E.g. oxidation state of N is +5 in nitrate (NO3-) while it is +3 in nitrite (NO2-). For these acids, the ‘-ite’ is dropped and replaced with ‘-ous’ Amphoteric Compounds that can act as both an acid and a base. Common amphoteric oxides are those of beryllium, aluminium, zinc, tin and lead.

 Conduct an investigation to demonstrate the preparation and use of indicators as illustrators of the characteristics and properties of acids and bases and their reversible reactions Indicators: substances which change colour based on the pH of the environment. They are able to determine whether a substance is an acid or a base. Indicators provide a qualitative indication of the concentration of hydrogen ions in solution.

Indicators can be derived from natural sources (e.g. litmus from lichens) whilst others are synthetic chemicals (e.g. phenolphthalein). One of the most used indicators is universal indicator, a mixture of different indicators and providing colour changes from pH levels of 1-11>.

How Indicators Work Most indicators are weak acids or bases. HIn(aq) and In-(aq) are different colours.

Through the principles of Le Chatelier’s, when an acid is added to the solution, the addition of H + disturbs the equilibrium where it will shift left in order to minimise the disturbance. Thus, will turn the colour of HIn(aq). In the case of an addition of a base, the [In-] is increased as [H+] reacts with the base. To ensure an accurate measurement of pH, most labs consist of multiple indicators as they cover larger spectrums of the pH range.

 Predict the products of acid reactions and write balanced equations to represent:  Acids and bases Reactions of acids and bases are referred to as neutralisation reactions. Acid + base  salt + water Net ionic equation of these reactions is: H+(aq) + OH-(aq)  H2O(l). All of these reactions are exothermic. A salt is an ionic compound formed when the hydrogen is replaced with the metal ion (or ammonium). The solutions are said to be neutralised when the concentrations of the hydronium and hydroxide ions within the mixture are equal. The presence of the salt can be confirmed by evaporating the water.

 Acids and carbonates Acid + carbonate  salt + carbon dioxide + water During this reaction, bubbles form due to the presence of carbon dioxide. Hydrogen carbonates also undergo similar reactions. Acid + hydrogen carbonate  salt + carbon dioxide + water To confirm the production of carbon dioxide, the limewater teste can be used.

 Acids and metals Dilute acid + metal  salt + hydrogen gas Hydrogen ion (H+) is produced by the acid in aqueous solution, reacting with the metal. Net ionic equation: H+(aq) + Metal(s)  metal ion(aq) + H2(g). Reaction is responsible for acid rain. When testing for the presence of hydrogen gas, a ‘pop’ test can be used. The metal is “eaten away” and dissolved into liquid forming a soluble ionic compound

 Investigate applications of neutralisation reactions in everyday life and industrial processes Neutralisation reactions are present in society due to the importance to be able to change the pH of a substance. Everyday They can be used to treat water that is too basic or acidic such as balancing the pH of a river, drinking water (must be a pH of 7) or pools (maintained close to 7.4) and aquarium water. Antacid: our stomach produces hydrochloric acid to aid the digestion of food and kill any harmful micro-organisms however, when too much hydrochloric is produced the stomach is upset. Antacids contain bases such as magnesium or aluminium hydroxide which neutralises the excess acid. It could also consist of magnesium oxide, magnesium carbonate, calcium carbonate and sodium hydrogen carbonate. Industry Agriculture: the neutralisation of soil is crucial as the ability of plants to take up nutrients from the soil is influenced by how acidic or alkaline the soil is. When determining the pH level, white insoluble

barium sulfate (BaSO4) is sprinkled to provide a neutral, white background which makes colour changes clear unlike the dark colour of soil. Fertilisers – generally made from the neutralisation of ammonia and sulfuric acid or nitric acid.

 Conduct a practical investigation to measure the enthalpy of neutralisation From the law of conservation of energy: energy is needed to break bonds and is released when new bonds are formed. All neutralisation reactions are exothermic and follow the same net ionic equation. Standard enthalpy of neutralisation (ΔHneut) is the enthalpy change when solutions of an acid and alkali react together under standard conditions (25°C and 100kPa) to produce exactly one mole of water. Enthalpy of neutralisation is the enthalpy change when an acid and an alkali react to produce 1 mole of water. ΔHneut = q/n Literature value of enthalpy is -57.62 Kj/mol, results usually different values due to heat loss to surroundings and uneven distribution of heat.

Accuracy: Use a more accurate thermometer (digital), use more accurate scales, use a larger ΔT. Other neutralisation reactions provide similar enthalpy of neutralisation values as the net ionic equation is the same, sodium and potassium cations and the chloride and nitrate anions are spectator ions. Provide net ionic. Therefore, the bonds broken and formed are the same, resulting in the same enthalpy of neutralisation. If a solid base was used instead of an aqueous solution, as the solid needs to dissolve before it reacts which is an exothermic process, the values measured of the enthalpy of neutralisation will be more negative.

 Explore the change in definitions and models of an acid and a base over time to explain the limitations of each model, including: Lavoisier (1780): after experimenting on oxides of non-metals, he concluded that acids were substances that contained oxygen atoms. However, it was found that there were many acids without oxygen. Davy (1815): proposed acids contained replaceable hydrogen atoms. However, he had no explanation to how the molecules interacted and why some substances that didn’t contain hydrogen weren’t acidic e.g. methane, water, or sodium hydrogen carbonate. Liebig (1838): stated that active metals would displace hydrogen from acids.

 Arrhenius’ theory In 1884, Arrhenius proposed that an acid was a substance that ionised in solution to produce hydrogen ions and that a base was a substance that in solution produced hydroxide ions. He used HA to represent an acid: HA(aq)  H+(aq) + A-(aq). It was successful in accounting the behaviour of acid-

base reactions and could explain strong and weak acids as being due to complete or partial ionisation. Limitations o o

o o o

Doesn’t account for acid-base behaviour not in water solution, restricted to aqueous solution Couldn’t explain why some ions that didn’t contain hydrogen, hydroxide or oxide showed acid-base behaviour in water solution, e.g. ammonia, the basic nature of ammonia and why it didn’t produced hydroxide ions in solution Couldn’t explain the neutralisation of acids without hydroxide ions or neutralisation by gases Could not explain why neutralisation of some acids with bases produced a salt that was NOT neutral Could not explain why metallic oxides (Na2O) and carbonates (Na2CO3) are basic

 Brønsted-Lowry theory In 1923, Johannes Brønsted and Thomas Lowry proposed substances behaves as an acid when it donates one or more protons (H+) to a base and a substance behaves as a base when it accepts one or more protons from an acid. In summary: acids are proton donators and bases are proton acceptors. Acid-base reactions involve proton transfer from acid to base. Self-ionisation of water, acid-base reactions in non-aqueous solution and the acidic/neutral/basic nature of salts. Central to this theory is the concept of conjugate species. As hydronium ions and a salt is produced, if the reverse reaction were to occur, the hydronium ion acts as the conjugate acid and the salt acts as the conjugate base. Limitations o o

For a substance to be identified as an acid or base, a H+ must be transferred in the reaction. Doesn’t explain the acid-base behaviour of substances in non-aqueous solutions where a proton is not involved. These include of many non-metal oxides e.g. CO2¸ SO2, SO3 and substances like AlCl3 and BF3 which produce acidic solutions and can react with bases.

Lewis’ Definition Gilbert Lewis proposed a definition that was not limited or restricted by the chemical environment, allowing an acid or base to be identified even if no solvent is present. He proposed: an acid is an electron pair acceptor and a base is an electron pair donor. All Brønsted-Lowry acids and bases are Lewis acids and bases but the reverse is not true.

Using Brønsted-Lowry Theory IQ2: What is the role of water in solutions of acids and bases?

 Conduct a practical investigation to measure the pH of a range of acids and bases A more precise measure of the degree of acidity and alkalinity of solution is found through the pH scale, standing for hydrogen power – based on the concentration of hydrogen ions in solution.

pH = -log10[H3O+] or [H3O+] = 10-pH There are many ways of measuring pH such as indicators, pH meters and pH probes that give a digital readout of the pH of a solution. The pOH scale is a measure of basicity. pOH = -log10[OH-] or [OH-] = 10-pOH. At 25°C, the pOH and pH of a solution have the following relationship: pH + pOH = 14

When writing pH, 3.14 is two significant figures as the ‘3’ is a place holder. A pH probe or pH meter is used to measure pH. These are advantageous over indicators as they provide quantitative data however, are more expensive, must be calculated, take more time and are fragile.

 Calculate pH, pOH, hydrogen ion concentration ([H+]) and hydroxide ion concentration ([OH-]) for a range of solutions pH = -log[H+], [H3O+] = 10-pH pOH = -log[OH-] At 25°C: o o o

Neutral solution: both pH and pOH = 7 Acidic solution: pH < 7 and pOH > 7 Basic solution: pH > 7 and pOH < 7

pH of Strong and Weak Acids pH represents the relative strength and concentration of different acidic solutions. o o o

The higher [H+] is, the lower pH is As strong acids have complete ionisation, this means there is higher [H +], causing a lower pH. As weaker acids do not have complete ionisation, there is a lower [H +], causing a higher pH. The higher the acid concentration, the lower the pH

Calculating Ka from pH The pH of a monoprotic strong acid can be calculated using [H +] = [HA] and pH = -log10[H+]. However, as weak acids exist in equilibrium, the extent of dissociation will change. These require the acid dissociation constant Ka.

 Conduct an investigation to demonstrate the use of pH to indicate the differences between the strength of acids and bases Acid concentration does not relate to its pH as some acids donate a proton more readily than others. The Brønsted-Lowry theory defines the strength of an acid as its ability to donate protons to a base and the strength of a base is measured by its ability to accept protons from an acid.

The reaction of acids with water is referred to as ionisation reactions as ions are formed. When an acid ionises, it produces hydronium ions in aqueous solutions. E.g. HCl(aq) + H2O(l)  H3O+(aq) + Cl-(aq), or HCl(aq)  H+(aq) + Cl-(aq). Acid strength can be measured through determining how much of the acid molecules ionise to produce a hydrogen ion. Strong Acids Strong acids are strong electrolytes that ionise completely when bubbled through water. Reactions use ➔ to signify the dissociation reaction is complete. Strong acids readily donate protons. Therefore, solutions of strong acids contain ions with virtually no unreacted acid molecules remaining. E.g. HCl, H2SO4 and HNO3. Stronger acids have larger Ka values and their conjugate bases have inverse strength to their acids. There are 5 common strong acids: HCl, HBr, HI, H2SO4 and HNO3. Weak Acids Weak acids are weak electrolytes that ionise partially and use reversible reactions in the partial dissociation of a weak acid. E.g. acetic acid (CH3COOH) and carbonic acid (H2CO3). CH3COOH(aq) + H2O(l) ⇌ CH3COO-(aq) + H3O+(aq). When a base dissolves in water, it forms separate ions, these reactions are referred to as dissociation reactions. Bases usually dissociate to produce hydroxide ions in aqueous solution. E.g. NaOH(s)  Na+ K2O(s) + H2O(l)  2K+(aq) + 2OH-(aq). Strength of bases depend on the number of molecules (aq) + OH (aq) or ionic compounds that dissociate in aqueous solution, determined by how much the species have dissociated. Strong Bases Strong bases completely dissociates in water, releasing both anions and cations. The anions react completely with the water, accepting a proton to form hydroxide ions. Strong bases include Group 1 and 2 metal hydroxides. E.g. NaOH, KOH. However it is more correct to state them as an ionic compound that is a source of the strong base OH-. Weak Bases

Weak bases are which only a small proportion of the species produce hydroxide ions in aqueous solution. The reactions use reversible signs. E.g. ammonia (NH3), methylamine (CH3NH2) Do not use ‘strong’ and ‘weak’ to define ‘concentrated’ or ‘diluted’. Polyprotic Acids Monoprotic acids: acids that give up one proton per molecule when they ionise. pH of monoprotic acids:  

A strong monoprotic acid will ionise completely, so the [H3O+] will be equal to the concentration of the acid A weak monoprotic acid will only ionise to a small extent, producing very few [H 3O+]

Polyprotic acids: acids that can give up more than one proton, it is the ability to donate more than one proton not how readily the protons ionise i.e. diprotic – acid can donate two protons. pH of polyprotic acids: 



H2SO4, a strong acid ionises in two steps, a large number of the molecules ionise in the first equation and in the next ionisation of the weak acid, HSO4-, not many of the molecule will react with water and ionise. The total concentration of hydronium ions is the number produced by the first ionisation plus the number produced by the second ionisation. Therefore, strong diprotic acids have a hydronium ion concentration greater than a monoprotic acid of the same concentration and a lower pH.

 Write ionic equations to represent the dissociation of acids and bases in water, conjugate acid/base pairs in solution and amphiprotic nature of some salts, for example: Strength of Conjugate Acid-Base Pairs In all Bronsted-Lowry acid-base reactions, there are two conjugate acid-base pairs. Ex: HF(aq) + OH-(aq) ⇌ H2O(l) + F-(aq) HF is the conjugate acid of F- and are a conjugate acid-base pair, OH- is the conjugate base of H2O and is another conjugate acid-base pair. The stronger an acid is, the weaker its conjugate base is. Similarly, the stronger a base is, the weaker its conjugate acid is. This is as stronger acids have a stronger tendency to give up H+ ions, thus its conjugate base has a weaker tendency to accept a proton. A weaker acid/base will have a relatively strong conjugate base/acid. Acid + Base ⇌ Conjugate Base + Conjugate Acid Amphiprotic Nature An amphiprotic substance is on that can act as both a proton donor and an acceptor (acid and base), depending on which substance it is introduced to. E.g. H2O, HCO3-, H2PO4-. Amphiprotic substances are Bronsted-Lowry acids and bases.

Non-Neutral Salt Solutions Many ions display acidic or basic properties. This means that if the ionic compound containing an acidic or basic ion is dissolved, the resulting solution will not be neutral (ion produces H3O+ or OHions). To determine: o o o

Write the dissociation equation of the compound into its two ions Determine if either ion is acidic or basic by examining their conjugates Write an equation showing the acidic or basic ion reacting with water to form H+/H3O+ if acidic or OH- if basic.

Self-Ionisation of Water Water’s amphiprotic nature allows it to react with itself to form hydronium and hydroxide ions.

It is an equilibrium reaction as water is considered an extremely weak acid/base as a few water molecules ionise and donate a proton to another molecule. The forward reaction occurs to a small extent resulting in a small equilibrium constant. The self-ionisation constant (KW) Is used to represent the ionic product constant for water. KW = [OH-][H3O+] = 1.0 X 10-14 The [H3O+] and [OH-] are very low where in pure water at 25° (298K) the product of their molar concentrations is always 1.0 X 10-14, meaning individual they are 1.0 X 10-7M. In aqueous solution if either [OH-] or [H3O+] increases, the other must decrease proportionally.

Metal Oxides o o

Non-metals from the right hand side of the periodic table tend to form acidic oxides Metals from the left hand side of the periodic table tend to form basic oxides

 Sodium hydrogen carbonate  Potassium dihydrogen phosphate

 Construct models and/or animations to communicate the differences between strong, weak, concentrated and dilute acids and bases As the [OH-] and [H3O+] is 1.0 X 10-7M, the pH of pure water is, -log 10[H3O+] = 7.00, which is why a pH of 7 is considered neutral. o o o o

A concentrated solution is one in which the total concentration is high. A dilute solution is one in which the total concentration is low. Strong acids are acids in which almost all of the acid molecules ionise to produce a hydrogen ion. Weak acids are acids in which only some of the molecules ionise

Weak acids and bases present equilibrium reactions whereas strong acids and bases produce only forward reactions. In acidic substances, H3O+ ions are formed by the reaction of the acid with water as well as the selfionisation of water. This means the [H3O+] will be more than 10-7 M but as KW must remain constant, the [OH-] in the acidic solution must be less than 10-7 M. The opposite is true for basic solutions. At 25°: o o o

Pure water and neutral solution: [H3O+] = [OH-] = 10-7 so pH = 7 Acidic solution: [H3O+] > 10-7M and [OH-] < 10-7 , 10-7M so pH 10-7M so pH >7

Acid dissociation constant (Ka): represents the ionisation of a weak acid.

+¿ ¿ H3O ¿ −¿ A¿ ¿ ¿ K a=¿ As weak acids ionise to a small extent, the Ka values are very small. For strong acids, ionisation is almost 100% so the Ka values are very...