Chemistry Notes
HSC Core Topic 2
The Acidic Environment
1. Indicators were first identified with the observation that the colour of some flowers depends on soil composition
Recall:
An acid is a substance which in solution produces hydrogen ions, H+ or more strictly H3O+, sometimes called the hydronium ion.
A base is a substance which either contains the oxide O2- or hydroxide ion OH- or which in solution produces the hydroxide ion.
A soluble base is called an alkali.
Common acids are hydrochloric acid (HCl), sulfuric acid (H2SO4) and nitric acid (HNO3).
Common bases are sodium hydroxide (NaOH), barium hydroxide (Ba(OH)2), potassium oxide (K2O), magnesium oxide (MgO), iron III oxide (Fe2O3), copper hydroxide (Cu(OH)2), and ammonia (NH3). Of these seven substances only the first three and ammonia are alkalis (soluble in water).
Common properties of all acids are:
· Acids have a sour taste
· Acids sting or burn the skin
· In solution, acids conduct electricity
· Acids turn blue litmus (a vegetable dye) red
Common properties of alkalis are:
· Alkalis have a soapy feel
· Alkalis have a bitter taste
· In solution, alkalis are good conductors of electricity
· Alkalies turn red litmus blue
· Classify common substances as acidic, basic or neutral
All of the substances we use are acidic, basic or neutral, eg. vinegar, lemon juice, aspirin (acid); cleaning ammonia, sodium carbonate, oven cleaners (basic).
· 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, and that the range is identified by change in indicator colour
· Identify data and choose resources to gather information about the colour changes of a range of indicators
An indicator is a substance (usually a vegetable dye) which in solution changes colour depending on whether the solution is acidic or alkaline.
Simple indicators can show the nature of substances over a range through colour change
Indicator / Highly acidic / Slightly acidic / Neutral / Slightly alkaline / Highly alkalineMethyl orange / red / Yellow
Bromothymol blue / Yellow / blue
Litmus / Red / blue
Phenolphthalein / colourless / red
· Identify and describe some everyday uses of indicators including the testing of soil acidity/basicity
Indicators provide a cheap and convenient way of determining the acidity or alkalinity of substances. Some everyday uses of indicators are:
· Testing the acidity or alkalinity of soils (because some plants need an acidic soil – azaleas and camellias – while others need an alkaline soil – most annual flowers and vegetables)
· Testing home swimming pools (these need to be approximately neutral, though adding chemicals to sanitise the water can change its acid-alkali balance)
· Monitoring wastes from photographic processing (discharges to the sewerage system must be nearly neutral: photographic solutions are often highly alkaline).
· Perform a first-hand investigation to prepare and test a natural indicator
Natural indicators such as the above can be created by taking a relevant plant, crushing it into a powder if needed, and adding water/methanol, etc to extract the dye. (To make indicator paper, soak the paper in the indicator.)
2. While we usually think of the air around us as neutral, the atmosphere naturally contains acidic oxides of carbon, nitrogen and sulfur. The concentrations of these acidic oxides have been increasing since the Industrial Revolution
· Identify oxides of non-metals which act as acids and describe the conditions under which they act as acids
Oxides are a class of compound that often displays acidic or basic properties.
An acidic oxide is one which either:
· Reacts with water to form an acid or
· Reacts with bases to form salts (or does both).
Common acidic oxides are carbon dioxide, CO2, and diphosphorus pentoxide, P2O5, because
CO2(g) + H2O(l) -à H2CO3(aq) (carbonic acid)
Or CO2(g) + 2NaOH(aq) -à H2O(l) + 2Na+(aq) + CO32-(aq) (sodium carbonate)
P2O5(s) + 3H2O(l) -à 2H3PO4(aq) (phosphoric acid)
Or P2O5(s) + 6NaOH(aq) -à 3H2O(l) + 6Na+(aq) + 2PO43-(aq) (sodium phosphate)
A basic oxide is one that:
· Reacts with acids to form salts but
· Does not react with alkali solutions (such as NaOH or KOH).
Common basic oxides are copper oxide, CuO, and iron III oxide, Fe2O3, because
CuO(s) + H2SO4(aq) -à CuSO4(aq) + H2O(l)
Fe2O3 + 6HNO3 -à 2Fe(NO3)3(aq) + 3H2O(l)
· Analyse the position of these non-metals in the Periodic Table and generalise about the relationship between position of elements in the Periodic Table and acidity/basicity of oxides
The acidity of the oxide of an element increases from left to right on the periodic table.
· Define Le Chatelier’s principle
Chemical equilibrium is where a reaction does not go to completion, since the products are causing the back reaction at the rate of the forward reaction occurring. The relative amount of the products and reactants at equilibrium can be affected by shifting the equilibrium in order to favour one side.
Le Chatelier’s principle states that if an equilibrium is disturbed, the system adjusts itself to minimise the disturbance.
· Identify factors which can affect the equilibrium in a reversible reaction
Factors that can affect the equilibrium in a reversible reaction are changes in:
· Concentrations of species involved in the reaction
· Total pressure on a reaction involving gases
· Temperature
· Describe the solubility of carbon dioxide in water under various conditions as an equilibrium process and relate this to Le Chatelier’s principle
Solubility of CO2 in water:
CO2(g) + H2O(l) H2CO3(aq) H<0
Disturbance / Effect on equilibrium (right means inc. solubility)Increased pressure / Favours reaction with less moles of gas in product, reducing pressure, therefore shifts right, CO2 dissolves more.
Increase in concentration of CO2 in system / System aims to reduce concentration of CO2 and thus it reacts, ie, shifts equilibrium right
Increase in temperature / Favours endothermic reaction, ie back reaction, ie shifts left, and CO2 is less soluble
· Identify natural and industrial sources of sulfur dioxide and oxides of nitrogen
Natural and industrial sources of SO2 and nitrogen oxides
· Sulfur dioxide (SO2)
o Geothermal hot springs, volcanoes
o Processing or burning fossil fuels (S(compounds) + O2(g) --> SO2(g))
o Extracting metals from sulfide ores
(eg 2ZnS(g) + 3O2(g) --> 3ZnO(s) + 2SO2(g))
· Oxides of nitrogen: nitrous oxide (N2O), nitric oxide (NO), nitrogen dioxide (NO2)
o Lightning (O2(g) + N2(g) --> 2NO(g))
o Same in high temps of combustion chambers (cars, power stations)
o Further reaction of 2NO(g) + O2(g) --> 2NO2(g)
o N2O formed by bacteria (from nitrogenous fertiliser, causes greenhouse effect)
o Mixture of NO and NO2 referred to as NOx
· Describe, using equations, examples of chemical reactions which release sulfur dioxide and chemical reactions which release oxides of nitrogen
These pollutant oxides can become dissolved in water droplets, within large volumes of air.
SO2(g) + H2O(l) --> H2SO3(aq)
2H2SO3(aq) + O2(g) --> 2H2SO4(aq) (catalysed)
2NO2(g) + H2O(l) --> HNO2(aq) + HNO3(aq)
2HNO2(aq) + O2(g) --> 2HNO3(aq) (catalysed)
§ Assess the evidence which indicates increases in atmospheric concentration of oxides of sulfur and nitrogen
§ Choose resources, gather and analyse information from secondary sources to summarise the industrial origins of the above gases and evaluate reasons for concern about their release into the environment
There has been an increase in industrial areas since the Industrial Revolution, in SO2 and NOx pollution. This led to regulations to control emissions from factories, power stations and motor cars. The annual average concentration of SO2 and NO2 in most large cities around the world is 0.01 ppm for each gas. This is about 10 times the value for clean air, though a concentration of 0.01 ppm for either gas is not harmful. Globally, because SO2 and NO2 are washed out of the atmosphere by rain, there appears not have been any significant build-up of their concentrations over the last century or so (unlike CO2 which there has been a 30% increase and N2O increased as well).
Sulfur dioxide irritates the respiratory system and causes breathing difficulties at concentrations as low at 1 ppm. People suffering from asthma and emphysema are particularly susceptible. The effects of sulfur dioxide are magnified if particulates are present also. Nitrogen dioxide irritates the respiratory tract and causes breathing discomfort at concentrations above about 3 to 5 ppm and at higher concentrations does extensive tissue damage. Concentrations above 3 ppm have rarely been reached even in heavily polluted cities. The main problem with NO2 is that it leads to the formation of ozone in what is called photochemical smog. This form of air pollution in which sunlight acts upon nitrogen dioxide in the presence of hydrocarbons and oxygen to form ozone and other pollutants called peroxyaclynitrates (PANS). Ozone has harmful effects at concentrations as low as 0.1 ppm. Hence releasing these oxides into the atmosphere at high concentrations can be harmful.
§ 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 101.3kPa or 25°C and 101.3kPa
At 0oC and 1atm pressure, molar volume of gases is 22.41 L.
At 25oC and 1atm pressure, molar volume of gas is 24.47 L.
See CCHSC for examples
· Explain the formation and effects of acid rain
High SO2 and NOx emission cause acidic rain (as seen above): acid rain is rain with a higher H+ concentration than normal, > 10-5 molL-1. Regular rain unaffected by these pollutants contains some acidic carbonic acid (from CO2) and usually contains 10-6 – 10-5 molL-1 H+.
Acid rain causes:
· Increasing lake acidity (affecting fish populations)
· Damage to pine forests
· Erosion of marble and limestone buildings/decorations (these Contain carbonates which react with acids)
· Process and present information from secondary sources to describe the properties of sulfur dioxide and the oxides of nitrogen
Properties of SO2 and NOx
SO2 / N2O / NO / NO2Colourless
Pungent odour
Soluble / Colourless
Sweet smell
Insoluble / Colourless
No smell
Insoluble / Reddish-brown
Choking odour
Soluble in water
· food preservative
· bleaching
· fumigant / · anaesthetic
· propellant / · synthesizing nitric oxide / · nitric acid
· fertilisers
· explosives
· Breathing difficulties at 1ppm / · Breathing difficulties at 5ppm
· Leads to forming O3
3. Acids occur in many foods, drinks and even within our stomachs
· Define acids as proton donors and describe the ionisation of acids in water
A previous (Arrhenius) definition was that acids form H+ in water (ionisation reaction). The H+ joins with water to make the hydronium ion (H3O+).
According to the Lowry-Brönsted definition (1923), acids are proton donors and bases are proton acceptors.
· Identify acids such as acetic acid (ethanoic acid), citric acid (2-hydroxypropane-1-2-3-tricarboxylic acid), vitamin C and hydrochloric acid as naturally occurring acids, and acids such as sulfuric acid and hydrobromic acid as manufactured acids
Naturally occurring acids / Manufactured acidsAcetic acid (ethanoic acid) – vinegar
Citric acid (2-hydroxypropane-1,2,3-tricarboxylic acid, C6H8O7)
Vitamin C (ascorbic acid, C6H8O6)
Hydrochloric acid (in our stomachs) / sulfuric acid
hydrobromic acid
nitric acid
phosphoric acid
· Describe the use of the pH scale in comparing the concentrations of acids and alkalis
The pH scale is a useful scale for comparing the concentrations of acids and alkalis, since the concentrations of the ions themselves involved range over a very large linear scale. It is more convenient to use a smaller, closer scale, and thus pH is used.
· Describe acids and their solutions with the appropriate use of the terms strong, weak, concentrated and dilute
A strong acid is one where all its molecules ionise to form H+.
A weak acid is one where only some of its molecules ionise to form H+.
A concentrated solution is where the total concentration of the acid is high.
A dilute solution is where the total concentration of the acid is low.
· Identify pH as –log10[H+] and explain that a change in pH of 1 means a ten-fold change in [H+]
pH = -log [H3O+] A change by 1 in pH is a change tenfold in [H+].
Note: the number of decimal places for pH should equal the number of significant figures for [H3O+]
Since pH = -log10[H+] [H+] = 10-pH
pOH = -log10[OH–] pH + pOH = 14 pH = 14 + log10[OH–]
A change of 1 in pH relates to a tenfold change in [H+]
A neutral substance is one where [H+] = [OH–], at 25oC, [H+] = 1x10–7molL-1
· Compare the relative strengths of equal concentrations of citric, acetic and hydrochloric acids and relate this to the degree of ionisation of their molecules
We say that HA is stronger than HB for weak acids, if:
HA(aq) H+(aq) + A–(aq lies further to the right than HB(aq) H+(aq) + B–(aq)
Thus the degree of ionisation of HA is greater than HB
For equal concentrations of the following acids, different pHs were found, and it was deduced that HCl is strong, citric acid is weak and acetic acid is weaker.
· Describe the difference between a strong and a weak acid in terms of an equilibrium between the intact molecule and its ions
For strong acids, there is no equilibrium in the ionisation reaction (the back reaction does not occur), while for weak acids there is.
· Gather and process information from secondary sources to write ionic equations to represent the ionisation of acids
Hydrochloric acid is a strong acid. This means that when we dissolve hydrogen chloride gas in aqueous solution all the HCl molecules react with water to form hydrogen ions:
HCl(g) + H2O(l) -à H3O+(aq) + Cl-(aq)
For a strong acid the ionisation reaction with water goes to completion.
On the other hand acetic acid is a weak acid, because when pure covalent acetic acid is dissolved in aqueous solution, only some of the acetic acid molecules actually react with water to form hydrogen ions:
CH3COOH(aq) + H2O(l) ß-> H3O+(aq) + CH3COO-(aq)
This reaction is an equilibrium one: it does not go to completion.
· Gather and process information from secondary sources to explain the use of acidic oxides such as sulfur dioxide as food additives