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Aberami Purushothamar2009

Module 2 Acidic Environment1. Indicators were identified with the observation that the

colour of some flowers depends on soil composition.

Classify common substances as acidic, basic or neutral.

Acidic Basic Neutral

Vinegar

Lemon juice

Aspirin

Vitamin C

Cloudy ammonia

Washing soda

Antacid tablets

Oven cleaners

Water

Salt

Sugar

Identify that indicators such as litmus, phenolphthalein, methyl orange andbromothymol blue can be used to determine the acidic or basic nature of a materialover a range, and that the range is identified by change in indicator colour.

An indicator is a substance that changes colour in presence of an acid or a base.

Indicator Colour in acid Colour in baseLitmus Red (below pH = 5) Blue (above pH = 7.5)Phenolphthalein Colourless (below pH =

8)Red (above pH = 10)

Bromothymol blue Yellow (below pH = 6) Blue (above pH = 7.5)Methyl orange Red (below pH = 3) Yellow (above pH = 4.5)

Identify and describe some everyday uses of indicators including the testing of soil

acidity/basicity.

Testing Soil

Neutral white powder (e.g. Barium sulfate) to see colour change

Then universal indicator is added

Others

Checking the water in swimming pools using litmus paper.

Testing aquarium water.

Identify data and choose resources to gather information about the colour changesof a range of indicators.

Indicator pH Range Colour RangeIndicator pH range Colour changeMethyl orange 3.1 4.4 Red yellow

Methyl red 4.4 6.0 Pink yellowBromothymol blue 6.2 7.6 Yellow blue

Litmus 6.2 7.4 Red bluePhenol red 6.8 8.4 Yellow red

Phenolphthalein 8.3 10.0 Colourless red

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2. While we usually think of the air around us as neutral, theatmosphere naturally contains acidic oxides of carbon,nitrogen and sulfur. The concentrations of these acidicoxides 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.

React with water to form an acid.

React with bases to form a salt.

They are oxides of non-metals i.e. NO2, CO2.

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/basicityof oxides.

Metals (left side of Table) form basic oxides.

Non-metals (right side of Table) form acidic oxides.

Define Le Chateliers principal.

If a system at equilibrium is disturbed, then the system adjusts itself tominimise the disturbance.

The forward or reverse reaction will occur at a faster rate until equilibrium isre-established.

Identify factors which can affect the equilibrium in a reversible reaction.

Temperature.

Concentration.

Pressure (with gases)

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

equilibrium process and explain in terms of Le Chateliers principal.

CO2 dissolves in water producing carbonic acid (exothermic)

CO2(g) + H2O (l) H2CO3 Increase in pressure moves right.

Adding excess CO2 without changing volume moves right.

Adding excess CO2 without changing pressure no change to equilibrium.

Adding NaOH neutralise moves right.

Adding H+ - moves left.

Increasing temperature favours reactants - moves left.

Identify natural and industrial sources of sulfur dioxide and oxides of nitrogen.

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Sulfur Dioxide

Natural sources geothermal hot springs/volcanoes ( of it).

Industrial process processing and burning fossil fuels.

Nitrogen Oxide

Natural sources lightening and certain bacteria.

Industrial process combustion (both power plants and cars).

Describe, using equations, examples of chemical reactions which release sulfur

dioxide and chemical reactions which release oxides of nitrogen.

Sulfur Dioxide

Smelting metals.

4FeS2(s) + 11O2(g) 2Fe2O3(s) + 8SO2 (g)

Oxides of Nitrogen

Lightening and in combustion of car engines.

N2 (g) + O2(g) 2NO (g) then2NO (g) + O2(g) 2NO2 (g)

Assess the evidence which indicates increases in atmospheric concentration of

oxides of sulfur and nitrogen.

SO2 and NO2 are water-soluble and therefore washed out of the atmosphereby rain

there is little significant build-up of their concentrations over the last 100years

However, it is difficult to be sure because there is a lack of data for periods

before the 1950s It has only been in the last few decades that we can measure concentrations

of these gases below 0.1 ppm with sufficient accuracy.

Calculate volumes f gases given masses of some substances in reactions, and

calculate masses of substances given gaseous volumes, in reactions involving gasesat 0oC and 100kPa or 25oC and 100kPa.

The Mole Method

Step One. Write a balanced equation

Mg (s) + 2HCl (aq) MgCl (aq) + H2 (g)

Step Two. Moles of known substance n=m/M

n= 6.5/24.3 = 0.267 mol Step Three. Determine the moles ration of unknown to known1:1

Step Four. Find the number of moles of the unknown by multiplying with theknown by mole ratio

1 x 1 = 1

Step Five. Times the number of moles by the constant needed.= 0.267 x 24.82

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= 6.63 L

Explain the formation and effects if acid rain.

Acid rain has a higher hydrogen ion concentration.

Acidic oxides dissolve in the water droplets.

How acid rain is formed:

SO2(g) + H2O (l) H2SO3 H2SO3 + O2 H2SO4 (sulfuric acid)

If the quantity of acid rain is greater than the capacity of an environment toneutralise it, then the following can occur:

Soil pH can drop, making it difficult to for plants to absorb sufficient mineral cations

Soil chemistry can change, leading to death of important micro-organisms, andrelease ofnormally insoluble aluminium and mercury into soil water

Protective waxes can be lost from leaves, causing extensive leaf damage

Buildings constructed from carbonates, such as concrete, mortar, limestone andmarble, canbe gradually dissolved away

Aquatic animals can die as water pH drops below 5

Smog and acid rain can combine to form killer fog

Stop acid rain

The only way to stop it is stop emitting SO2 and NO2.

Now legal limits on the amount factories can emit.

Analyse information from secondary sources to summarise the industrial origins ofsulfur dioxide and oxides of nitrogen and evaluate reasons for concern about theirrelease into the environment.

Sulfur Dioxides

After the industrial revolution in the 1800s, there was a great increase inemissions of SO2 to the atmosphere surrounding industrial cities, mainly fromburning coal and extracting metals.

Air quality significantly deteriorated significantly until the 1950s,

Smelters convert minerals into metals i.e. zinc:

2ZnS (s) + 3O2 (g) 2ZnO (s) + 2SO2 (g)

Oxides of Nitrogen

Serious pollution from NOx did not appear until the 20th century as electricitygeneration and motorcar usage expanded dramatically.

With the high temperatures of coal-fired power stations led to the reactionbetween nitrogen and oxygen.

N2 (g) + O2 (g) 2NO (g) Reacts slowly with more oxygen.

2NO (g) + O (g) 2NO2 (g)

Concern

NOx is a potent greenhouse gas.

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SO2 irritates respiratory system.

2SO2 (g) + 4H2O (l) 2H2SO4 (aq)

3. Acids occur in many foods drinks and even within ourstomachs.

Define acids as proton donors and describe the ionisation of acids in water.

Proton Donor

An acid is a proton donor.

Ionise when it contacts water to release hydrogen ions.

Each hydrogen ion (H+) consists of one proton.

HCl (aq) H+ (aq) + Cl- (aq)

Ionisation of Acids

H+ released do not exist alone.

They attach themselves to a water molecule forming a hydronium ion.

HCl + H2O H3O+ + Cl-

Identify acids including acetic (ethanoic), citric (2-hydroxypropane-1,2,3-

tricarboxylic), hydrochloric and sulfuric acid.

Acetic (ethanoic) Acid

CH3COOH

Found in vinegar and wine.

Citric Acid (2-hydroxypropane-1,2,3-tricarboxylic)

CH2COOHCOHCOOHCH2COOH

Found in citric fruits i.e. oranges.

Hydrochloric Acid

HCl

Found in stomach acid.

Sulfuric Acid

H2SO4 Acid rain.

Describe the use of the pH scale in comparing acids and bases.

The pH scale is used to compare the concentration of hydrogen ions [H+] in solutions of acids andbases.

pH = -log[H+]

pOH = -log[OH-]

pH + pOH = 14A pH of 7 is neutral, where the concentration of H+ and OH- is equalpH < 7; acidic

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pH > 7; basic

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

weak, concentrated and dilute.

A strong acid fully ionises into H+ ions and its conjugate base in solution and is a goodconductor.

e.g. HCl H+ + Cl-

A weak acid does not fully ionise in solution and is a poor conductor.e.g. CH3COOH H+ + CH3COOA

concentrated acid has a relatively high number of acid molecules per volume ofsolution.

e.g. 4M HCl

A dilute acid has a relatively low number of acid molecules per volume of solutione.g. 0.01M HCl

Identify pH as log10 [H+] and explain that a change in pH of 1 means a ten-foldchange in [H+].

p means -log10

pH = -log10[H+]

If a substance has a molarity of 1M, it has a pH of 1

If a substance has a molarity of 0.1M, it has a pH of 2

a change in pH of 1, is a ten-fold change in concentration of hydrogen ions

Compare the relative strengths of equal concentrations of citric, acetic and

hydrochloric acids and explain in terms of the degree of ionisation of theirmolecules.

HCl is a strong acid, fully ionises in water.

Acetic acid (ethanoic) is weak, partially ionises in water. Citric acid partially (least).

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

equilibrium between the intact molecule and its ions.

In a strong acid, such as hydrochloric acid, an equilibrium is formed duringionisation:

HCl(aq) + H2O(l) H3O+(aq) + Cl(aq)

In the equilibrium of the strong acid, the equation completely lies on the rightside (near 100%

ionisation). The molecule of the strong acid completely ionises.

In a weak acid, such as acetic acid, an equilibrium is formed duringionisation:

CH3COOH(aq) + H2O(l) H3O+(aq) + CH3COO(aq)

In the equilibrium of the weak acid, the equation lies mostly on the left(partial ionisation). The molecule of the weak acid is in solution with few of itsions.

Use available evidence to model the molecular nature of acids and simulate the

ionisation of strong and weak acids.

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Gather and process information from secondary sources to explain the use if acids

as food additives.

Jams

Citric acid often added to jams to give sharp taste.

The acidity helps prevent the growth of microbes so they dont decomposefood

The lowing of pH makes enzyme reactions slow down, slowing the spoiling offood.

Identify data, gather and process information from secondary sources to identify

examples of naturally occurring acids and bases and their chemical composition.

Citric acid naturally occurs in citrus fruit.

Hydrochloric acid is found in stomachs where it aids digestion.

Process information from secondary sources to calculate pH of strong acids given

appropriate hydrogen ion concentrations.

4. Because of the prevalence and importance of acids,they have been used and studied for hundreds of years.Over time, the definitions of acid and base have beenrefined.

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

Lavoiser - 1780

Found that oxides of non-metals produced acidic solutions.

Defined acids as a substance containing oxygen.

Davy - 1815

Showed HCl contained no oxygen, thus disapproved Lavoisiers theory.

Proposed that acids contained replaceable hydrogen atoms.

Proposed that bases reacted with acid to form salt and water

Arrhenius - 1884

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Acids ionise in aqueous solutions to produce hydrogen ions.

Strong acids ionise 100%, weak acids only partially.

A base is a substance that in solution produced hydrogen ions

Outline the Brnsted-lowry theory of acids and bases.

Acid-base Reaction

acid is a proton donor while base is a proton acceptor

HCl(g) + NH3(g) NH4Cl(s)

A proton is transferred from the hydrogen chloride to the ammonia.

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

Conjugate Pairs

An acid donates a proton to form what is called its conjugate basee.g. HCl + H2O H3O+ + Cl-

Similarly, a base accepts a proton to form what is called its conjugate acide.g. NH3 + H2O NH4+ + OH-

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

their acidic, neutral or basic nature.

Neutral Salts

Are formed when a strong acid and a strong base react.

NaCl(aq) Na+(aq) + Cl-(aq)

H2O(l) H+(aq) + OH-(aq)

The ions form a strong base and a strong acid.

The concentration of hydrogen ions equals the concentration of hydroxideions, neutral.

Acidic Salts

Are formed by a strong acid reacting with a weak base.

NH4Cl(aq) NH4+(aq) + Cl-(aq)

H2O(l) H+(aq) + OH-(aq)

The ions form a weak base and a strong acid.

This leaves an unbalance supply of hydrogen ions, thus the acidic solution.

Basic Salts

Are formed by a weak acid reacting with a strong base.

When sodium acetate dissolves in water Na+, CH3COO-, H+ and OH-

The ions form a weak acid and a strong base.

This leaves an unbalanced supply of hydroxide ions.

Identify conjugate acid/base pairs.

Whenever an acid and a base react, they form their conjugates.HCl + H2O H3O+ + Cl-

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HCl acid1H2O base1H3O+ - conjugate acid of base1Cl- - conjugate base of acid 1

Identify amphiprotic substances and construct equations to describe their behaviour

in acidic and basic solutions.

Can act as a base or an acid (accept or donate protons).

Water is acting as an acid, giving up a proton.

H2O(l) + NH3(g) NH4+(aq) + OH-(aq)

Water is acting as a base, accepting a proton.

H2O(l) + HCl(g) H3O+(aq) + Cl-(aq)

Identify neutralisation as a proton transfer reaction which is exothermic.

Is a proton transfer reaction that consists of an acid reacting with a base.

Is exothermic. Ionic equation: H+(aq) + Cl-(aq) + Na+(aq) + OH-(aq) H2O(l) + Na

+(aq) + Cl-

(aq)

Net ionic equation: H+(aq) + OH-(aq) H2O(l)

Is the reaction between hydrogen and hydroxide ions

Describe the correct technique for conducting titrations and preparation of standard

solutions.

Methods of Calculation1. Write the equation

2. Find the molar relationship.3. Calculate the number of moles of the known substance.4. Use the relationship to determine the moles of the unknown.5. Calculate the molarity(c) of the unknown.6. Write down answer including units.

Standard solution

Accurately known concentration.

To be suitable must be water soluble, accurately known formula, high purityand be stable in air i.e. lose water.

Preparation

Accurately weigh a calculated amount of solid. Dissolve it in water, transferring the entire dissolved solid to a volumetric

flask, adding water to prepare a fixed volume of solution.

Concentration calculated in mol L-1

Other

A standard solution can be reacted with a solution of unknown concentrationusing titration technique.

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One reactant in solution is slowly added to another reactant in solution untilan end point is reached.

The end point of the titration is usually indicated by a change in colour of asmall amount of indicator solution added to the mixture of reactants.

For an acid-base titration an indicator is selected that changes colour at thepH of the salt solution formed at the point of neutralisation.

This is known as the equivalence point.

Calculations are carried out to three significant figures.

Qualitatively describe the effect of buffers with reference to a specific example in a

natural system.

Buffer Control

Controls the level of acidity or basicity in a solution.

If an acid or base is added to a buffer solution there is barely a change in pH.

Contain equal concentrations of a weak acid and its conjugate base.

Blood is a buffer solution, containing carbonic acid.

H2CO3(aq) H+(aq) + HCO-(aq) If more CO2 is dissolved more H ions will form decreasing pH, equilibrium

moves left minimising change.

Gather and process information from secondary sources to trace developments in

understanding and describing acid/base reactions

Originally an acid was perceived to be a substance which had a sour taste,and which reacted with certain metals (Zn and Fe).

Around 1780, Antoine Lavoisier proposed that acids were substances whichcontained oxygen.

He justified this by citing non-metal oxides such as sulfur dioxide (SO2) wereacidic.

However, his theory was disproved by the existence of acids such ashydrochloric acid which do not contain oxygen.

In 1815, Humphry Davy suggested that acids contained replaceablehydrogen. That is, hydrogen that could easily be displaced by reaction withmetals. For example, HCl reacts with zinc to produce zinc chloride andhydrogen gas.

This theory was accepted for most of the 19th century,

In 1884, Svante Arrhenius proposed that an acid was a substance thatproduced hydrogen ion in water.

He also defined bases as substances which produced hydroxide ions in water.However there are some flaws in Arrhenius definition. The ionisation of anacid is something that happens in a solution, not in isolation.

He does appropriately address the role of the solvent in ionisation.

In 1923, two scientists Bronsted and Lowry independently devised a newdefinition. They defined acids as proton donors and bases as protonacceptors. This is the current definition.

Analyse information from secondary sources to assess the use of neutralisation as a

safety measure or to minimise damage in accidents or chemical spills.

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Neutralisation reactions are used for safety purposes in laboratories andindustry since many acids and alkalis are very corrosive.

It is important to neutralise spills of these substances quickly.

Sewage authorities put strict limits on the pH of factory and laboratoryeffluent discharge.

Sodium carbonate (Na2CO3) is widely used to neutralise acid spills since it is astable solid, which is easily stored, cheaply available, and if too much is used,it is less danger than other bases.

5. Esterification is a naturally occurring process which can beperformed in the laboratory.

Describe the differences between the alkanol and alkanoic acid functional groups in

carbon compounds.Alkanols

Are alcohols derived from alkanes.

Contain the function group OH.

Alkanoic Acids

Are carboxylic acids derived from alkanes.

Contain the carboxylic acid functional group COOH.

Identify the IUPAC nomenclature for describing the esters produced by reactions ofstraight-chained alkanoic acids from C1 to C8 and straight-chained primary alkanolsfrom C1 to C8

The alkanol loses its ending -anol, and thyl is added. The alkanoic acidbecomes an -oate

e.g. methanol + ethanoic acid becomes methyl ethanoate

Alkanol Alkanoic acidC1 Methyl MethanoateC2 Ethyl EthanoateC3 Propyl PropanoateC4 Butyl ButanoateC5 Pentyl PentanoateC6 Hexyl HexanoateC7 Hepyl HeptanoateC8 Octyl Octanoate

Explain the difference in melting point and boiling point caused by straight-chained

alkanoic acid and straight-chained primary alkanol structures.

The acids have three polar bonds compared to the two in the alcohols.

Thus the molecules are held together more strongly in the acid and theboiling point of the acid is greater.

Alcohols have a higher melting and boiling points than their parent alkanes(due to dipole-dipole attractions and hydrogen bonds).

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Identify esterification as the reaction between an acid and an alkanol and describe,

using equations, examples if esterification.

Esterification

Is the reaction between an alkanoic acid and an alkanol.

Esters are prepared by the condensation of an alkanol and an alkanoic acid.

They are equilibrium reactions.

Alkanol + alkanoic acid ester + water.

OH from acid and H from the alkanol functional group combine.

Describe the purpose of using acid in esterification for catalysis.

Esterification catalysed by adding a small amount of concentrated sulfuricacid (dehydrates so produces more water).

This condensation reaction moves the equilibrium to the right increasing theyield of ester by removing water molecules.

Acts as a catalyst too (it has two roles!).

Explain the need for refluxing during esterification

Endothermic reaction requires heat to shift eqm to left and produce moreyield.

When reaction is heated, volatile chemicals could escape and cause problemsi.e. an explosion.

Refluxing makes the procedure safer.

A condenser which is cooled by water is placed on the top of the reactionvessel.

Any volatile components pass into the condenser, causing the gas to form a

liquid (not highly flammable).

Outline some examples of the occurrence, production and uses of esters.

Esters occur naturally in the form of flavourings and scents.

The odours and flavours of fruits cause by the esters.

Some esters are used in the industry as solvents as they are able to dissolvemany polar and non-polar substances.

Manufactured esters are used for flavouring food, perfume and colouringcosmetics.

Ethyl acetate is nail polish remover.

1258017237_2009_Chemistry_Notes

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