The Acidic Environment

Classify common substance as acidic, basic of neutral

• Alkali: Bases that dissolve in water

• All of the substances we use are acidic, basic or neutral. This classification is usually based on the concentration of hydrogen ions produced when the substance is dissolved in water.

•  Acids:

• Common acids:

•  Hydrochloric acid (HCl), sulfuric acid (H2S04) and nitric acid (HN03). vinegar, lemon juice, aspirin Bases:

•  Strong bases completely ionize in aqueous solutions to give OH - and a cation Weak acids do not completely ionize in solution but exist in equilibrium. Common bases:

•  Na0H Ba(0H)2

• K20

Identify that indicators such as Litmus, phenolphthalein, methyl orange and bromothymol blue can be used to determine the acidic or basic nature of a material over a range ,is identified by

change in indicator colour.

Indicators: Substances that change colour depending on the concentration of hydronium ( H O + ) ions in a solution to determine whether a solution is acidic or basic.

Examples of indicators include litmus, phenolphthalein, methyl orange and Bromothymol blue.

Identify and describe some everyday uses of indicators including the testing

of soil acidity/basicity

• Everyday uses of indicators include:

• The monitoring of the pH-level of fish-tank environments

• Some marine animals are very sensitive to pH levels of water Sea water is about pH 8.5

• Testing the acidity or alkalinity of soils

• because some plants need an acidic soil–E.g. azaleas and camellias

• While others need an alkaline soil–e.g. most annual flowers and vegetables

• Agricultural Lime, calcium carbonate increases the basicity while Manure increases acidity.

• Testing home swimming pools

•  

• These need to be approximately neutral (close to pH: 7.4), though adding chemicals to sanitise the water can change its acid-alkali balance. If it is too acidic, it is dangerous to humans, if it is too basic, then pests grow

Perform a first hand investigation to prepare and test a natural indicator

Aim: to prepare and test a natural indicator

Method: Thinly slice some red cabbage leaves and place them in a beaker

Add water into the beaker and boil the mixture while carefully stirring with a rod

Allow the mixture to cool down and then filter out the cabbage leaves from the solution. Use the solution to test the pH of chemicals

Colour of indicator: Neutral: Blue/purple Acidic: pink to red Basic: Green to yellow

Identify oxides of non-metals which acts as a acids an describe the

conditions under which they act as acids

Acidic Oxides: One, which reacts with water to form acids, or one, which reacts with bases to form salts.Examples-SO2, SO3,CO2

Basic Oxides: One which reacts with acids to form salts, but does not react with alkali solutionsExamples- CuO, Fe2O3

Neutral Oxides: Examples- CO, NO, N2O

Analyse the position of these non-metals in the periodic table and outline the relationship between position of elements in the periodic

table and acidity/basicity of oxides.

Group 1 oxides are strongly basic. The basicity increases down the group

Group 7 oxides are strongly acidic. The acidity decreases down the group.

Generally, the acidity of oxides increases across the periodic table and decreases down the periodic table. In group 3, the trend down the group is that it starts from highly acidic, then to amphoteric, then basic.

Most transition metal oxides are basic although some are amphoteric.

Define Le Chatelier's principle

Chemical Equilibrium: When the forward rate of reaction equals the reverse rate and the concentrations of all the species in the system are constant.

Due to the equilibrium, the reactants are forming products at the same rate at which the products are being converted back to reactant, therefore the concentration of species are constant.

Le Chatelier's Principle: If a system at equilibrium is disturbed (i.e. Changes to equilibrium are made e.g. Adding more reactants), then the system adjusts itself so as to minimize the disturbance, re-establishing equilibrium.

Identify factors which can affect the equilibrium in a reversible reaction

Factors which affect the equilibrium in a reversible reaction:

Change in concentrationChange in temperature- If the temperature is lowered, the exothermic reaction is favoured .If the temperature is increased, the endothermic reaction is favoured.Change in total gas pressureAn increase in pressure leads to decrease in volume leading to increase in concentrationA decrease in pressure leads to an increase in volume leading to decrease in concentration.  By increasing the total gas pressure on the equilibrium, the equilibrium shifts to counteract the increase in pressure

Describe the solubility carbon dioxide in water under various conditions as an

equilibrium process and explain in terms of Le Chatelier's principle

Carbon dioxide gas is soluble in water. The following equation describes its equilibrium:

CO2 (g) CO2 (aq) Carbon dioxide gas dissolves in water.

CO2 (aq) +H 2O(l) H+(aq)+HCO3- (aq)

When there is more carbon dioxide, this reaction increases concentration of hydrogen ions in water (Le Chatelier's principle); therefore the water becomes more acidic and tastes sharp

Example- Soda water contains carbon dioxide in water

-Identify natural and industrial source of sulpher dioxide and oxides of nitrogen.

-Descriptive using equations, examples of chemical reactions which release sulpher

dioxide and chemical reactions which release oxides on nitrogen.

Natural sources of Sulphur dioxide ( SO2 ):

Geothermal hot springs

Emissions from volcanoes.

Bushfires

Natural decay of vegetation on land, wetlands and oceans

Industrial sources of sulphur dioxide: 

The burning of fossil fuels which contain sulphur, such as coal and oil ( industrial emissions) S (s) + O2 (g) = SO2 (g)

When metals are extracted from sulfide ores.

E.g. 2ZnS(s)+3O2 (g) = 2ZnO(s) + 2SO2 (g)

0ther examples of industries which emit sulphur dioxide include:

Petroleum refineries Cement manufacturing Metal processing facilities

0xides of nitrogen: Dinitrogen monoxide (N20), nitric oxide (N0), nitrogen dioxide (N02)

Natural sources of oxides of nitrogen ( NO, NO2 , N2O ):

Lightning O2 (g)+N2 (g = 2NO(g)

Analyse information from secondary sources to summaries the industrial

origins of sulpher dioxide and oxides of nitrogen and evaluate reasons for

concerned concern about their release into the environment

 

• Reasons for concern about release of SO2 and NOx

•  These oxides combine with water to form an acidic solution in the atmosphere, creating acid rain.  

•  Quick summary of effects–Kills marine life–Kills plants and forests–Erodes the build environment

• Increase in oxides have led to increased production of photochemical smog, especially in large cities which brings forth respiratory problems for humans and animals, i.e. breathing difficulties, tissue damage,

Assess the evidence which indicates increase in atmospheric concentration

of oxides of sulpher and nitrogen

• Evidence of the oxides of sulphur and nitrogen being formed:

• 0xides of sulphur and nitrogen are mainly found in the atmosphere as a result to a number of natural process such as:  

• Volcanic action, which produces sulphur dioxide

• Natural decay of vegetation on land and sea ( SO2 ) Lightning which produces nitric oxide

• Due to the industrial revolution and increased human population, there has been an increase of these oxides. Manmade sources of the pollutants include

Calculate volumes of gases given masses of some substances in

reactions, and calculate masses of substances given gaseous volumes, in reactions involving gases at 0°C and

100 kPa or 25°C and 100 kPa

• At 100kPa (1 bar):

-STP (standard temperature and pressure):

273K (0oC), molar volume of gases is 22.71 L/mol

-SLC (standard laboratory conditions) or RTP (room temp and pressure):

-298K (25oC), molar volume of gas is 24.79 L/mol

At 101.3kPa (1atm):

-STP:

273K (0oC), molar volume of gases is 22.41 L/mol

-SLC or RTP:

298K (25oC), molar volume of gas is 24.47 L/mol  

• In a reaction involving gases only, moles ratio can be considered as volume ratios

Explanations formation and effects of acidic rain

What is Acid Rain-

• Acid rain is rain with a higher concentration of hydrogen ions ( H +) than normal, which causes the rain to be more acidic.

•  This rise of acidity is caused by high emission of SO2 and NOx , which is released into the atmosphere by the burning of fossil fuels.

•  The major cause of acid rain is the burning of fossil fuels from power plants, which produces electricity, and automobiles.

Formation of acid rain:

•  Pollutants such as smoke and fumes released from burning fossil fuels rise up into the atmosphere and combine with moisture (water) to form acid rain

• The main chemicals in these pollutants are SO2 and nitrogen oxides ( NOx ) The sulfur dioxides and nitrogen oxides react with water in the atmosphere to produce mild acidic solutions of sulfurous acid and nitric acid.

• Rainwater, snow, fog and other forms of precipitation containing these mild acidic solutions fall into the earth as acid rain

• Effects of acid rain:

• Destruction of plants and forests. The change of pH levels in soil due to the acid levels in acid rain creates difficulty for plants to absorb sufficient minerals and kills important micro-organisms which help sustain life. Also leaves of trees are damaged and lost, due acidic rain

Define acids as proton donor and describe the ionization of acids in

water

• When acids dissolve in water, it can donate a proton, ie a Hydrogen ion ( H +) to a water molecule.

• The proton ion in water solution may be represented as H+(aq) or H3O+(aq ) (hydronium)

•  

• For example, the ionization of HCl

Identify acids including acetic(ethanoic), citric (2-

hydroxypropane -1,2,3- tricarboxylic), hydrochloric and sulfuric acid

• Acetic/Ethanoic acid ( CH3COOH ):

• Acetic acid is a natural acid and is a component in vinegar

•  Citric Acid ( C6 H8O7 ):

•  Citric acid is also a naturally occurring acid, found in almost all living things 0ccurs in citrus fruits such as lemons and oranges

•  Hydrochloric acid ( HCl )

•  Another natural acid (it is present in the stomach for digestion)

•  Sulfuric acid ( H 2 SO4 )

•  Non-natural acid, produced for many industrial processes

• Major uses include: fertilizers, manufacture of chemicals, electrolyte in batteries

Describe the use of pH scale in comparing acids and bases

The pH scale is a measure of the concentration of hydrogen ions in solution. Low pH means a high level of proton (hydrogen ions) concentration and thus very acidic and a high pH means a low level of proton concentration, so it is basic. "

Describe acids and their solutions with the appropriate use of the terms

strong, weak, concentrated and diluted

• Strong acids: ones that disassociate completely or almost completely in water, forming H

• Weak acids: those that partially ionize in water

•  Concentrated acids: are those that have a large amount of solute in a given amount of solution. It has a high number of moles per litre

•  Dilute acids: those that have low amount of solute ionized in a given amount of solution.

• Important: The strength of an acid does not depend on the concentration. Strength depends on the type of acid, eg HCl can never be called a weak acid

Identify pH as - log10 (H+) and explain that a change in pH of 1 means a ten-

fold change in (H+)

• pH is a measure of the concentration of hydrogen ion, hence: pH=-log10 [H + ]

• By definition, [H +][OH -] =10^-14 ie. pH +pOH =14

• [H3O+]=10^-pH . So if pH = 7, then the concentration of H3O+= 10^-7 mol/L (neutral) 

• A neutral substance is one where [H+]= [0H-] = 1x10-7molL-1

• A change in pH of 1 means a ten-fold change in [H+]. This means that a solution of pH 1 has 10 times the concentration of [H+] in pH 2 and 100 times the concentration in pH 3. The pH scale is a logarithmic scale,

Compare the relative strengths of equal concentrations of citric, acetic and hydrochloric acids and explain in terms of the degree of ionization of

their molecules

• Citric and acetic acid are both very weak acids. Acetic acid is a weaker acid with a pH of around 2.9 while citric acid has a pH of around 2.1  

• Hydrochloric acid is a much stronger acid, around pH 1. This means that HCl has a very high degree of ionization, so the forward reaction in:

HCl(aq) +H 2O(l) =H 3O(aq)+Cl(aq) is much more favoured, producing more ions. Acetic acid on the other hand has a low degree of ionization, and mainly exists in water as acetic

acid molecules rather than ions.

Describe the difference between a strong and a weak acid in terms of an

equilibrium between the intact molecule and its ions"

• Strong acids disassociate completely or almost completely in water, releasing H ions in aqueous solution. Since the production of ions is much, much more favoured, therefore the equilibrium lies well to the right. For example:

HCl(aq) +H 2O(l) = H3O(aq)+Cl-(aq)

• Weak acids only ionize partially in water, resulting in an equilibrium between its intact molecules and ions

Outline the historical development of ideas about acids including those of -

Lavoisier - Davy - Arrhenius

Lavoisier:

• Proposed that an acid must contain oxygen

• Non metal oxides reacted with water to form acidic solution Davy:

•  Redefined an acid as a substance containing hydrogen (rather than oxygen), which could be partially or totally replaced by metals when it reacted together, to produce a salt. He realized this by discovering that HCl had acidic properties.

Arrhenius:

Suggested that acids ionized in solution to produce hydrogen ions as the only positive ion in solution

Bases produced hydroxide ions as the only negative ion in solution Completely ionized acids were strong acids, weak if partially ionized Arrhenius proposed these ideas only when the acid was in an aqueous state

 

Outline the Bronsted-Lowry theory of acids and bases

• Bronsted-Lowry acid: Any molecule or ion that acts as a proton donor (hydrogen ions)

• Bronsted-Lowry base: Any molecule of ion that can act as a proton acceptor

• The Bronsted-Lowry theory of acids and bases states that an acid is a substance which in solution donates a proton, while a base is a substance that accepts a proton

• HCl (g) +H 2O(l) =H 3O +Cl-

• HCl (acid)is the proton donor, since it gives a H+ ion Water (base) is the proton acceptor, accepting the H+ ion

•  Monoprotic acids: Acids that ionize and lose one protons Diprotic acids: Acids that ionize and lose two protons

Describe the relationship between an acid and its conjugate base and a base

and its conjugate acid

• Acids donate a proton to form its conjugate base, while bases accept a proton to form its conjugate acid.

• Acid +Base=Conjugate base (ofacid )+Conjugate acid (of base)

eg. HCl (aq) +H O(l) = Cl- +H3 O+

• Cl - is the conjugate base of HCl (acid)

• H3 O + is the conjugate acid of H O (base)

Identify a range of salts which form acidic, basic or neutral

solutions and explain their acidic neutral or basic nature

Identify conjugate acid/base pairs

Acid +Base=Conjugate base (of acid )+Conjugate acid (of base)

Strength Acid Conjugate Base

Strong H2SO4HCl

HSO −4

Cl −

Weak CH 3COOH CH3COO-−

H 2CO3 HCO3 −

   NH 4+ NH

   HSO4− SO4 (2−)

4 SO4

Very Weak H 2O OH−

Identify amphiprotic substances and construct equations to describe their

behaviour in acidic and basic solutionsAmphiprotic : substance that can act either as a proton donor (acid) or a proton acceptor (base)

Examples include: H 2O

Some More points described in detail in the Notes “The Acidic Environment” are listed

Below

• Identify neutralization as a proton transfer reaction which is exothermic

• Describe the correct technique for conducting titrations and preparation of standard solutions"

• Qualitatively describe the effect of buffers with reference to a specific example in a natural system"

• "Analyse information from secondary sources to assess the use of neutralization reactions as a safety measure or to minimize damage in accidents or chemical spills

• "Describe the differences between the alkanols and alkanoic acid functional groups in carbon compounds

• Identify the IUPAC nomenclature for describing the esters produced by reactions of straight chained alkanoic acids from C1 to CB and straight- chained primary alkanols from C1 to CB

• Explain the difference in melting point and boiling point cause by straight-chained alkanoic acid and straight-chained primary alkanols structures“

• Identify Esterification as the reaction between an acid and an alkanols and describe using equations, examples of esterification

• Describe the purpose of using acid in esterification for catalysis

• Explain the need for refluxing during esterification

• Identify data, plan, select equipment and perform a first hand investigation to prepare an ester using reflux

• Outline some examples of the occurrence, production and uses of ester

• Process information from secondary sources to identify and describe the uses of esters as flavours and perfumes in processed foods and cosmetics

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Reference and acknowledgment

- Compiled Notes of Isaac Seunglee Suh