Module 2- The Acidic Environment
October 4, 2017 | Author: rbtlch1n | Category: N/A
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HSC Chemistry Summary Module 2- The Acidic Environment
Acid
Anion
Typical salt
nitrious
nitirc
sulfurous
sulfuric
hydroiodic
hydrobromic
hydrochloric
H2CO3
HNO2
HNO3
H2SO3
H2SO4
HI
HBr
HCl
PO43-
CO32-
NO2-
NO3-
SO32-
SO42-
I-
Br-
Cl-
Phosphate
Carbonate
Nitrite
Nitrate
Sulfite
Sulfate
Iodide
Bromide
Chloride
Mg(HCOO)2
Na3PO4
CaCO3
NaNO2
PB(NO3)2
Na2SO3
K2SO4
AgI
KBr
NaCl
CaF2
carbonic
H3PO4
HCOO-_ Formate (methanoate) CH3COO- acetate
KCN
Fluoride
phosphoric
HCOOH
cyanide
ZnS
F-
Formic (Methanoic)
CH3COOH
CN-
sulfide
HF
acetic
HCN
S2-
Hydrofluoric
hydrocyanic
H2S
Ag(CH3COO)
Hydrogen sulfide
1
Robert Lee Chin
HSC Chemistry Summary Module 2- The Acidic Environment
The Acidic Environment: 1. Indicators x
Classify common substances as acidic, basic or neutral
Acids Acids are substances capable of providing hydrogen ions (H+) for chemical reactions. Free ions are only available in solutions where the proton is stabilised by a solvent molecule. In an aqueous solution it exists as the hydronium ion: H H 2 O o H 3O Properties of acids: -Sour taste -Sting or burn skin -Conduct electricity as an aqueous solution -Turn blue litmus red -React with active metals to produce a salt and hydrogen gas e.g. Na (s) H 2 SO 4(aq) o H 2(g) NaSO 4(aq) -React with many carbonates to produce salt, water and CO2 e.g. (NH 4 ) 2 CO 3(s) 2HCl (aq) o H 2 O (l) CO 2(g) 2NH 4 Cl -React with bases, neutralising to form water and a salt e.g. HNO 3(aq) NaOH (aq) o H 2 O (l) NaNO 3 Common Substance
Name of acid
Lemon juice
Citric/2-hydroxylpropane1,2,3-tricarboxylic acid
Cream of tartar
Vinegar Fizzy drink
Aspirin
Car battery acid Vitamin-C tablets Yogurt Wine, bananas
Tartaric
Chemical Formula KHC4H4O6 C6H8O7
Acetic/ethanoic
C2H4O2
Carbonic
sulphuric
ascorbic lactic
H2CO3 C9H8O4 H2SO4 C6H8O6 C3H6O3
Tartaric
C4H6O6
Acetylsalicylic
Uses Whipping eggwhites Flavour for food Preservative, flavouring food Fizzy taste Pain relief medicine Car battery Dietary supplement Detergents, Biopolymer precursor Food flavouring
Bases Bases are substances that react with acids to form salts or form the hydroxide ion (OH-) in solution. A soluble base is called an alkali. Metal oxides act as bases when in solution e.g.: CaO (s) H 2 O (l) o Ca 2 (aq) 2OH Properties of alkalis: -bitter taste -Soapy, slippery feel -Conduct electricity as an aqueous solution Robert Lee Chin
2
HSC Chemistry Summary Module 2- The Acidic Environment -Turn red litmus blue Common Substance Ammonia
Name of base
Chemical Formula
Uses
Ammonia
NH3
Hand soap Detergent Antacid
Magnesium/Aluminium hydroxide Sodium hydrogen carbonate Sodium hydroxide
Mg(OH)2 Al(OH)3 NaHCO3
Cleaning agent, insect stings Cleaning Agent Cleaning agent Relieve Indigestion
Bicarbonate of soda Lye water
NaOH
Used in baking Additive in some foods cleaning agent
Neutral substances Neutral substances are neither acidic nor basic. Examples are pure water, pure alcohol and sugar. The salts formed in neutralisation acids are neutral as are some oxides.
x
Common Substance Pure water Table salt sugar
Name of substance
Chemical Formula
Uses
dihydrogen oxide Sodium chloride Sucrose
H2O
NaCl C12H22O11
Pure alcohol
Ethanol
C2H5OH
Essential for life Food additive, preservative Food ingredient and preservative Cleaning agent, preservative
Perform a first-hand investigation to prepare and test a natural indicator
Experiment: Extracting and using a natural indicator Aim: To prepare an indicator solution from red cabbage and test the resulting indicator on a range of substances Equipment: 2-3 large red cabbage leaves, shredded 500 mL beaker 250 mL beaker Tripod, gauze mat and Bunsen burner Test tubes & and test tube rack Universal indicator (optional)
Robert Lee Chin
Approx, 2mL solution of each: 0.1 mol L-1 NaOH 0.1 molL-1 HCl white vinegar household ammonia lemon juice lemonade bicarbonate of soda washing powder antacid tablet (grind into powder) salt water 3
HSC Chemistry Summary Module 2- The Acidic Environment Method: 1/
Place shredded cabbage leaves in 500 Ml beaker and just cover with distilled water (about 200 mL). Slowly boil the cabbage leaves until the water turns a dark reddish-purple and the leaves lose most of their colour.
2/
Allow to cool and pour the liquid off into a clean 250Ml beaker. This is the red cabbage indicator.
3/
Place 2 mL of NaOH and HCl into separate test tubes. Add a few drops of red cabbage indicator until a definite colour is observed. Record the colour of the indicator
4/
Repeat step 3 with the other substances and record results. Classify the substances as acidic, neutral or basic.
5/
Optional: Test each of the solutions with universal indicator to check your classification.
Results: Substance NaOH(aq) HCl(aq) white vinegar Household ammonia Lemon juice Lemonade Bicarbonate of soda Antacid Salt water
Red cabbage indicator colour yellow Red Pink Dark green
Acidic/basic/neutral neutral Acidic Acidic Basic
Universal indicator colour Purple Red Red Blue-green
Red Purple-magenta Blue-green
Acidic Acidic Basic
Red Red Blue-green
Cloudy purple Purple
Slightly basic Neutral
Lime green Dark green
As a generalisation, the red cabbage indicator turned acidic substances red and basic substances blue. Neutral substances stayed the same colour as the red cabbage indicator (purple)
Robert Lee Chin
4
HSC Chemistry Summary Module 2- The Acidic Environment x
Identify data and choose resources to gather information about the colour changes of a range of indicators
x
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 Indicator Litmus
Phenolphthalein Bromothymol Blue Methyl Orange Universal Indicator
Highly acidic red
Slightly acidic red
colourless yellow
colourless yellow
red red
yellow Orangeyellow
Neutral Reddishblue colourless green yellow Green
Slightly basic blue
Highly basic blue
pink blue
red blue
yellow Blue
yellow Purple
Memory assist: Acids turn blue litmus red (Blue in acid goes Red- BAR)
x
Solve problems by applying information about the colour changes of indicators to classify common substances as acidic, neutral or basic
Investigation: Testing the acidity of household substances Aim: To determine the acidity/basicity of some household substances using some indicators Equipment: Small test tubes and test tube rack Beaker Dropper bottles containing: -Phenolphthalein -Litmus -Methyl orange -Universal indicator -Bromothymol Blue
Substances to be tested: -distilled water -drain cleaner -ammonia -vinegar -lemonade -baking soda -shampoo -conditioner -egg white -antacid -lemon juice
Method: Robert Lee Chin
5
HSC Chemistry Summary Module 2- The Acidic Environment 1/
Each substance will be tested using 5 different indicators. Pour about 20 mL of each of the substances into separate test tubes. For drain cleaner, dissolve about a teaspoonful into 200 mL of distilled water, and then pour into the test tubes.
2/
Add one drop of a different indicator to each one of the substances. Mix thoroughly and record the observed colour. Repeat for each substance.
Results: Substance
Phenolphthalein
Litmus
Methyl orange
Distilled water Drain cleaner Ammonia Vinegar Lemonade Baking Soda
Clear Pink Magenta Clear Clear Magenta
Purple Violet Purple Pink R ed Blue
Shampoo
White
Conditioner
White
Pinkpurple Violet
Egg white
Magenta-pink
Purple
Orange
Antacid Lemon juice
Pink Clear
Blue R ed
Yellow R ed
Universal indicator
Orange Yellow Yellow Turquoise Orange Blue-grey R ed R ed R ed R ed Yellow Bluegreen Yellow Pink Yellow
Lime green Bluegreen R ed R ed
Bromothymol Blue
Blue-green Light blue Blue Yellow Yellow Blue Yellow Yellow Blue Blue Yellow
As a generalisation (based on the results): -Phenolphthalein reacts ONLY with basic substances, turning them pink, then red for higher pH -Blue Litmus turns stronger bases purple and strong acids red -Methyl orange turns bases yellow and acids red -Universal indicator turns bases blue and acids red -Bromothymol remains blue in bases and yellow in acids x
Identify and describe some everyday uses of indicators including the testing of soil acidity/basicity
Water Testing pH levels in swimming pools need to regularly tested and maintained between 7.27.8. Above this will encourage growth of bacteria, mould and algae. Above 7.8 and below 7.2 will cause irritation to skin and eyes. A pool pH kit is used to measure the pH level. If it is too low, bicarbonate of soda is added. If too high, chlorine bleach powder.
Robert Lee Chin
6
HSC Chemistry Summary Module 2- The Acidic Environment Fish in aquariums are sensitive to the pH. Too acidic or alkaline water will kill certain fish. Testing of soil pH Many plants can only tolerate a certain pH range in the soil. For example, carnivorous plants prefer acidic soils while beetroot thrives in slightly alkaline soil. To test the pH, a white unreactive powder is first mixed with the soil to absorb moisture before adding universal indicator. Effluent Testing pH can be used to assess the levels of certain types of industrial pollution. Indicators are used to monitor the pH of waste water and natural waterways.
2. Acids in our Environment x
Identify oxides of non-metals which act as acids and describe the conditions under which they act as acids
We can distinguish whether an oxide is an acid or a base by observing its effects on an indicator or seeing if it reacts with an acid or base In general, oxides of metals act as bases; they turn litmus red. They react with water to form an alkaline solution: Metal oxide water o alkaline solution E.g. MgO (s) H 2 O o Mg(OH) 2(aq) Basic oxides react with acids to form water and a salt: Metal oxide acid o water salt E.g. MgO (s) 2HCl o H 2 O (l) Mg(Cl) 2(aq) Oxides of non-metals act as acids; they turn litmus blue. They react with water to form acids: Non - metal oxide water o acid E.g. SO 2 (s) H 2 O o H 2 SO 3 (aq) Acidic oxides react with bases to form water and a salt: Non - metal oxide acid o water salt E.g. SO 2 (aq) 2NaOH o H 2 O (l) Na 2 SO 3 (aq) Basic oxides do not react with alkali solutions
Robert Lee Chin
7
HSC Chemistry Summary Module 2- The Acidic Environment x
Analyse the position of these non-metals in the Periodic table and outline the relationship between position of elements in the periodic table of elements and acidity/basicity of oxides.
In general, the oxides of elements on the LHS (metals) form basic oxides and the oxides of elements on the RHS (non-metals) form acidic oxides. The noble gases do not form oxides. H Be
Zn Zr
= amphoteric oxides
= basic oxides
= acidic oxides
= neutral oxides
x
B
C
N
O
F
Al
Si
P
S
Cl
Ge
As
Se
Br
Sn
Sb
Te
I
Pb
Bi
Po
At
Define Le Chatelier’s principle
Revision of Equilibrium: Many reactions are reversible reactions i.e. forwards and reverse reactions occur at the same time. In an undisturbed, closed system, these reactions will eventually reach a state of equilibrium Features of a system at equilibrium: 1) It is a closed system- no energy or matter leaves or enters 2) Macroscopic properties (e.g. colour, temperature, state, pressure) remain constant 3) Concentrations of products and reacts remain constant 4) Rate of → reaction = rate of ← reaction 5) Microscopic changes DO occur 6) There will ALWAYS be some product & reactant Le Chatelier’s principle applies to systems already in equilibrium that then undergo some change.
Robert Lee Chin
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HSC Chemistry Summary Module 2- The Acidic Environment Le Chatelier’s principle states that if a system in equilibrium is disturbed/changed, then the system adjusts itself to minimise this change These changes are: -concentration of products + reactants -temperature (different effects for endo- and exothermic reactions) -pressure & volume (only if gases are involved) Equilibrium and Indicators: Indicators can be written as HIn, where ‘H’ is the hydrogen atom and ‘In” is the indicator molecule. Indicator reactions are reversible reactions. The equilibrium situation is: HIn H+ + InColour 1 Colour 2 If an alkali is added, the forward reaction is favoured, so more product is formed and colour 2 appears. If an acid is added, the reverse reaction is favoured, so more reactants form and colour 1 appears. Other reactions e.g. combustion reactions; reactions between acids and metals are not reversible- they go to completion. x
Identify factors which can affect the equilibrium in a reversible reaction
By changing concentration, pressure or temperature of reactants and products, we can affect the equilibrium point. Concentration: Increasing concentration of reactants will drive the reaction forward, while increasing the products will drive the reaction in the reverse direction. For example, in A+B C+D, increasing reactants will drive the reaction forward, producing more products, thus reducing the concentration of A and B and maintaining equilibrium. Temperature:
Reaction
Exothermic: A+B C+D + heat Endothermic: A+B + heat C+D
Effect on equilibrium if temperature increases Shifts left- favours reactants Shifts right- favours products
Pressure: (For reactions involving gases) If pressure is increased, the equilibrium will favour the side with the lower amount of substances because this will reduce the number of particles per volume.
Robert Lee Chin
9
HSC Chemistry Summary Module 2- The Acidic Environment x
Describe the solubility of carbon dioxide in water under various conditions as an equilibrium process and explain in terms of Le Chatelier’s principle.
Carbon dioxide comes from volcanic gases, burning of organic matter and respiration of plants and animals. It exists in sea and other natural waters and forms 0.030.04%/V of the atmosphere. The concentration of CO2 in the atmosphere will continue to increase due to more animals, more machines & factories, fewer rainforests and increase in temperature (this means CO2 is les soluble in water, so more is released). When CO2 dissolves in water, an equilibrium forms: CO2(g) + H2O(l) H2CO3(aq) (carbonic acid) The solubility of carbon dioxide in water can be explained in terms of De Chatelier’s principle. Changing the concentration, pressure, temperature or adding chemicals that react with products or reactants alters the equilibrium. Concentration: If the concentration of CO2 is increased, the equilibrium will shift to the right to use up the extra carbon dioxide (and if CO2 concentration is decreased, it will shift to the left to produce more carbon dioxide). If more H2CO3 is added, the equilibrium will shift to the left to use up the extra carbonic acid (and if H2CO3 is removed, it will shift to the right to make more H2CO3). Pressure: Increasing the pressure of the carbon dioxide means will force the equilibrium to use up more CO2 so there are fewer particles. The equilibrium will move to the right, so more carbonic acid will be formed and the solution will become more acidic. Temperature: The reaction is exothermic, so can be written as: CO2(g) + H2O(l) H2CO3(aq) + heat Increasing the temperature will cause the equilibrium to shift to the left to use up the added heat. This is why a warm can of fizzy drink is less fizzy than a cold can- less CO2 can be dissolved Adding reactive chemicals: As carbonic acid forms, it ionises, so equilibrium occurs: H2CO3(aq) 2H+(aq) + CO32-(aq) So the equation can be written as: CO2(g) + H2O(l) 2H+(aq) + CO32-(aq) If we add OH- ions, they will react with the H+ ions, removing them from solution. The equilibrium will shift to the right to make more H+ ions (more CO2 dissolves). The concentration of CO2 in the atmosphere will continue to increase due to more animals, more machines & factories, fewer rainforests and increase in temperature (this means CO2 is les soluble in water, so more is released). Robert Lee Chin
10
HSC Chemistry Summary Module 2- The Acidic Environment x
Identify natural and industrial sources of sulfur dioxide and oxides of nitrogen
Sulfur Dioxide Sulfur dioxide is a colourless, toxic, gas with a pungent odour. It irritates the eyes, damages the respiratory tract and can cause asthma. Industrial sources account for over 75% of all emissions, in particular, combustion of fossil fuels Natural Sources 1. Burning organic matter (bushfires)
1.
2. Decay of organic matter
2.
3. Volcanic and hot spring emissions
3. 4. 5.
Industrial Sources Combustion of fossil fuels (esp. Power plants, vehicles) Smelting of sulphide ores into metal (Pb, Zn, Cu) Manufacture of sulphuric acid, paper, food processing, sewage treatment Petroleum refineries Burning garbage
Oxides of nitrogen There are 3 oxides of interest, all of which cause damage to the respiratory system, increasing the risk of respiratory infections and asthma. Nitrogen dioxide, NO2 Names
Nitrogen (IV) oxide
Colour pH
-Red brown -Acidic -Poisonous
Dinitrogen monoxide, N2O -Nitrogen (I) oxide -Nitrous oxide (aka ‘laughing gas’) -Colourless -neutral
Nitrogen monoxide, NO -Nitrogen (II) oxide -Nitric oxide Colourless -Not acidic, but reacts with oxygen, forming acidic NO2
In the atmosphere, these oxides are oxidised to nitric acid, nitrates and nitrites which settle or are washed away by rain. Strong sunlight causes oxides of nitrogen to react with hydrocarbons, forming photochemical smog. This became a major problem during the industrial revolution in the mid 20th century. Nitrogen oxide Nitrogen dioxide, NO2 Dinitrogen monoxide, (nitrous oxide) N2O Nitrogen monoxide (nitric oxide), NO
Robert Lee Chin
Natural sources -Action of sunlight on nitrogen monoxide and oxygen -Produced by soil bacteria -Produced by soil bacteria -Lightning
Industrial sources -Combustion of fossil fuels in vehicles and power stations -Fuel for racing cars -Sedative/analgesic (‘laughing gas’) -Burning organic matter -Combustion of fossil fuels in vehicles and power stations 11
HSC Chemistry Summary Module 2- The Acidic Environment x
Describe, using equations, examples of chemical reactions which release sulfur dioxide and chemical reactions which release oxides of nitrogen
Reactions releasing sulfur dioxide Manufacture of Iron(II) sulfate is prepared from tanks of dilute sulphuric acid used to clean iron sheets before plating/galvanising. Iron(II) sulfate heated above 300°C decomposes into Iron(III) oxide, sulfur dioxide and sulfur trioxide: 00q C 2FeSO 4(s) 3 o Fe 2 O 3(s) SO 2(g) SO 3(g) Iron sulphide (Pyrite/Fool’s Gold) is a source of sulfur dioxide when roasted in air. The other product is Iron(III) oxide: eat 4FeS 2(s) 7O 2(g) h o 2Fe 2 O 3(s) 4SO 2(g) Oxidation of hydrogen sulphide during the decay of organic matter produces sulfur dioxide and water: xidation H 2S( g ) o o SO 2 ( g ) H 2 O ( l ) Smelting of metal ores (copper, lead, zinc) releases sulfur dioxide. E.g. smelting zinc sulphide releases sulfur dioxide and zinc oxide: eat ZnS(s) O 2(g) h o SO 2(g) ZnO (s) In the laboratory, sulfur dioxide is prepared by heating copper with sulphuric acid. Copper sulfate and water are by-products: eat Cu (s) 2H 2 SO 4(aq) h o SO 2(g) CuSO 4(aq) 2H 2 O (l) Sulfur dioxide is also produced when a sulphite e.g. sodium sulphite (Na2SO3) is treated with dilute acid: Na 2 SO 3(s) 2H (aq) o SO 2(g) H 2 O (l) 2Na (aq) Reactions releasing nitrogen oxides High temperatures (e.g. combustion engines, lightning), nitrogen and oxygen combine to form nitric oxide: eat N 2(g) O 2(g) h o 2NO (g) This nitric oxide can slowly react with oxygen to form nitrogen dioxide: 2NO 2(g) O 2(g) o 2 NO 2(g) Industrially, nitric oxide is prepared by catalytic oxidation of ammonia, producing water as a by-product: xidation 4 NH 3( g ) 5O 2 ( g ) o o 6H 2 O ( l ) 4 NO 2 ( g ) In the laboratory, nitric oxide is produced using copper and nitric acid, producing water and copper(II) nitrate as by-products: 2Cu (s ) 4HNO 3( aq ) o 2 NO 2 ( g ) 2H 2 O ( l ) Cu ( NO 3 ) 2 ( aq )
Robert Lee Chin
12
HSC Chemistry Summary Module 2- The Acidic Environment In the laboratory, nitrogen dioxide is prepared by heating lead(II) nitrate crystals. The by-products are lead oxide and oxygen: eat Pb( NO 3 ) 2 (s ) h o 2 NO 2 ( g ) PbO (s ) O 2 ( g ) When heated, nitrogen dioxide forms nitric oxide and oxygen: eat 2 NO 2 ( g ) h o 2 NO ( g ) O 2 ( g ) Heating ammonium nitrate produces nitrous oxide and water: eat NH 4 NO 3 (s ) h o N 2 O ( g ) 2H 2 O ( l ) x
Assess the evidence which indicates increases in atmospheric concentration of oxides of sulfur and nitrogen
Ice Core Samples Ice core samples show Dinitrogen monoxide (N2O) levels have increased by 10%. Damage Damage to buildings, forests and aquatic organisms provides the most obvious evidence for increasing levels of sulfur and nitrogen oxides. Human health is affected in the form of respiratory diseases. Difficulties obtaining evidence Unlike carbon dioxide, sulfur and nitrogen dioxide are highly water soluble, so the validity of atmospheric measurements is questionable. These oxides also occur in much smaller concentrations than carbon dioxide (≈380ppm). The instruments used to measure these changes have only been available since the 1970’s. x
Explain the formation and effects of acid rain
Acid rain is acidic because it contains dissolved acidic oxides i.e. carbon, nitrogen and sulfur dioxides). Acidic oxides are released by several pathways. Natural sources are volcanoes & geysers; decaying vegetation. Industrial sources include the combustion of fossil fuels in industry and vehicles. In the atmosphere, they dissolve to form weak acids. These acidic particles can precipitate as rain, hail, snow of fog. For example, nitrogen dioxide forms nitric and nitrous acid: 2NO 2(g) H2O (l) o HNO 3(aq) HNO 2(aq)
Robert Lee Chin
13
HSC Chemistry Summary Module 2- The Acidic Environment Formation of acid rain: H2S
O2
SO2
O2
Acid Rain SO3
H2SO4 H2SO3
Volcanoes & geysers
Factories and Industries
Vehicles
Decaying vegetation
The effects of acid Rain Acid rain causes defoliation, stunted growth and decreased ability of plants to withstand frost. It can also leech into the soil and cause chemical reactions that affect plants Sulphuric acid ionises in water and removes plant nutrients. Sulfate ions leech out calcium and magnesium ions, which are essential for healthy soil. Normally insoluble compounds such as aluminium sulfate dissolve in acidic water, releasing toxic aluminium ions into the soil. Acid rain lowers pH in lakes and streams, killing aquatic life such as fish and crayfish. Aluminium ions cause reproductive problems and clog fish gills. Sulfate ions reduce visibility, especially in major cities in the US. Acid rain corrodes metals, stone structures, and paint. It is especially harmful to calcium carbonate in concrete, limestone and marble. Inhalation of sulfate ions has contributed to chronic respiratory diseases (lung cancer, bronchitis, asthma) in humans.
Robert Lee Chin
14
HSC Chemistry Summary Module 2- The Acidic Environment x
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 100kPa or 25°C and 100kPa
Use the following relationships:
number of moles number of moles
number of moles
number of moles
mass molar mass volume molar volume concentration (molL1 ) u volume (L) number of particles 6.022 u 10 23
Steps: 1. Change the mass/volume of given substances to moles 2. Write a balanced equation for the reaction to find mole ratios 3. Change the moles you have calculated back into the units the question asks. Examples: 1) a) Phosphorus trioxide, P2O3, slowly reacts with water forming phosphorus acid, H3PO3. Write a balanced equation for this reaction 2P2 O3( s ) 6 H 2 O(l ) o 4 H 3 PO3( aq ) b) When phosphorus acid is heated, it decomposes into phosphoric acid, H3PO4 and phosphine, PH3. Write a balanced equation for this reaction. eat 4H 3 PO 3(aq) h o 3H 3 PO 4(aq) PH 3(g) c) 7.10 L of phosphine gas was collected at 25°C, 100kPa. Show that the mass of phosphine gas is 9.72 g. 7.10 0.2864... moles 24.79 The mass of phosphine gas is the number of moles times the molar mass of phosphine: 0.2864... moles u [30.97 3(1.008)] 9.73607...g , which is approximately equal to 9.72 g.
The number of moles of phosphine gas is:
d) What mass of pure, solid phosphorus trioxide was involved in this reaction? 1 mole of phosphine is formed by 4 moles of phosphoric acid. The number of moles § 7.10 · of phosphoric acid is: 4¨ ¸ 1.13736...moles © 24.79 ¹
Robert Lee Chin
15
HSC Chemistry Summary Module 2- The Acidic Environment 4 moles of phosphoric acid is produced by 2 moles of phosphorus trioxide i.e. molar ratio of phosphorus trioxide to phosphoric acid is 1:2. The number of moles of § 7.10 · phosphorus trioxide is: 4¨ ¸ u 0.5 0.56868...moles © 24.79 ¹ The mass of phosphorus trioxide is: 0.56868...moles u [30.97 3(16)] 44.90885... # 44.91g (2 d.p.) x
Identify data, plan and perform a first-hand investigation to decarbonate soft drink and gather data to measure the mass changes involved and calculate the volume of gas released at 25°C and 100 kPa.
Investigation: Carbon dioxide in carbonated mineral water Background: As heat is applied to a soft drink, the amount of dissolved and reacted carbon dioxide decreases, and thus more and more escapes as a gas i.e. increased heat causes CO2(g) solubility to decrease. Equipment: -300 mL bottle of mineral water with flat base -500 mL beaker, gauze mat, tripod, Bunsen burner, digital balance -marble (calcium carbonate) heating chips Method: 1/ Weight the capped bottle of mineral water 2/
Pour 200mL of water in the beaker
3/
Uncap the soft drink, being careful not to spill any drink. Reserve cap. Mineral water Beaker with water & marble chips
Tripod & gauze Bunsen burner
4/
Rest the soft drink in the ‘water bath’. Gently heat the water and let boil for 35 mins.
5/
Take off heat and let bottle dry completely. Reweight combined soft drink and cap.
Robert Lee Chin
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HSC Chemistry Summary Module 2- The Acidic Environment Results: Total mass before heating = 520 g Total mass after heating = 510 g Mass difference = 520-520 = 10g Calculations Moles CO 2(g) escaping
Volume CO 2(g)
mass difference molar mass CO 2
moles CO 2(g) u molar volume
5.6340909... # 5.63 L (3 s.f.) x
10 [12 2(16)]
5 moles 22
5 u 24.79 22
Analyse information from secondary sources to summarise the industrial origins of sulfur dioxide and oxides of nitrogen and evaluate reasons for concern about their release into the environment
The main industrial sources of sulfur dioxide and oxides of nitrogen are: -combustion of fossil fuels in vehicles and power stations -burning of organic matter -smelting of metal sulphides -production of sulfuric acid, paper, food production, car fuels -petroleum refining Reasons for concern When these acidic oxides are released into the atmosphere (air pollution) by industries, they dissolve in the water to form acidic rain. The acidic particles produced can fall as gaseous or solid precipitates i.e. rain, hail, snow, dew, fog. Areas most affected are the USA, Canada and North-Western Europe. Australia has had less of a problem due to small population, isolation from other counties, coastal winds and low sulfur content in or fossil fuels. Acid rain causes damage to plant-life, aquatic life and ecosystems, man-made structures, causes respiratory diseases in humans and reduced visibility in major cities.
Robert Lee Chin
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HSC Chemistry Summary Module 2- The Acidic Environment
3. Acids occur in many foods drinks, and even within our stomachs x
Define acids as proton donors and describe the ionisation of acids in water
Acids are substances that release hydrogen ions (protons) when dissolved in water. Acids can be defined as proton donors. For example, sulphuric acid ionises to give hydrogen ions and sulphate ions: H 2SO 4(aq) o 2H (aq) SO4 2 (aq) This reaction is known as an ionisation reaction and is exothermic The H+ ions do not exist alone. The attach themselves to water molecules to form a hydronium ion, H3O+. So the ionisation of sulphuric acid can be written more correctly: H 2SO 4(aq) H2O (l) o 2H 3O (aq) SO4 2 (aq) A base can be defined as a proton acceptor. For example, when potassium hydroxide ionises in water it forms hydroxide ions and potassium ions: NaOH (aq) o Na (aq) OH (aq) The hydroxide ions can accept H+ ions to form water: OH (aq) H (aq) o H 2 O (l) Some acids, such as acetic acid (CH3COOH) are weak acids. Acetic acid has 4 hydrogen ions, but only one actually ionises in water (strong covalent bonds prevent the other hydrogen atoms from ionising): CH3COOH(aq) H+(aq) + CH3COOHx
Identify acids including acetic acid (ethanoic), citric (2-hydroxypropane1,2,3-tricarboxylic), hydrochloric acid and sulphuric acid
Name Acetic (ethanoic) acid
Formula CH3COOH
Citric acid (2-hydroxypropane1,2,3-tricarboxylic)
C6H8O7
Robert Lee Chin
Sources -Vinegar (4% solution) -Bacterial action -citrus fruits -Antioxidant additive -produced fermentation of
Other Info. -Pungent, colourless -Used to make acetates -Colourless, crystalline solid -Found in blood and urine -Added to jam 18
HSC Chemistry Summary Module 2- The Acidic Environment
Hydrochloric acid
HCl
sugar by Aspergillus fungus -Stomach acid
Sulfuric acid
H2SO4
-acid rain
Lactic acid
C3H6O3
-stiff muscles -yogurt, whey
Methanoic (formic) acid Ascorbic acid
H2CO2
-ants
C6H8O8
-fruits and vegetables -blood (metabolically active compound)
x
-antioxidant (food additive) Uses: -Industry -Cleaning brickwork Uses: -explosives -fertilisers -petroleum refining Uses: -lacquers and inks -cosmetics -pharmaceuticals Uses: -Rubber processing Uses: - antioxidant (food additive) -blood cell formation, tissue growth and healing
Describe the use of the pH scale in comparing acids and bases
The pH scale is used to compare the concentration of hydrogen ions in acid and base solutions. The following table relates pH to the concentration of hydrogen ions, hydroxide ions and example of common substances for given pH. pH
[H+]
[OH-]
0 1 2 3 4 5 6 7 8 9 10 11 12 13 14
100 = 1 10-1 10-2 10-3 10-4 10-5 10-6 10-7 10-8 10-9 10-10 10-11 10-12 10-13 10-14
10-14 10-13 10-12 10-11 10-10 10-9 10-8 10-7 10-6 10-5 10-4 10-3 10-2 10-1 100
Example substance 1 M HCl 0.1 M HCl Stomach acid Lemon juice Beer Acid rain Urine Pure water Sea water Toothpaste Detergent Ammonia Drain cleaner 0.1 M NaOH 1 M NaOH
For aqueous solutions the product of the concentration of hydrogen ions and hydroxide ions is the same, regardless of whether the solution is an acid, base or mixture.
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HSC Chemistry Summary Module 2- The Acidic Environment [H+] x [OH-] = ionisation constant, Kw = 10-14 at 25°C For acidic solutions, [H+] greater than 10-7 molL-1 and pH less than7 For basic solutions, [OH-] less than 10-7 molL-1 and pH greater than 7
x
Identify pH as -log10 [H+] and explain that a change in pH of 1 means a tenfold change in [H+]
Because pH is a logarithmic scale, a change in pH of 1 indicates the hydrogen ion concentration has changed by a factor of 10. Mathematically, the pH of a solution is given by: pH log10 [H ], where [H ] is the concentration of hygrogen ions in moles per litre To find the pH using a calculator: 1/ Tap the minus key 2/ Type in the [H+] (e.g. 2.0 x 10-5) 3/ Tap the log key and press enter x
Process information from secondary sources to calculate pH of strong acids given appropriate hydrogen ion concentrations
In strong acids, all hydrogen is assumed to ionise. The concentration of hydrogen ions will depend on the number of hydrogen ions an acid can donate i.e. monoprotic acids release one H+ per molecule, diprotic releases two and triprotic releases three. Examples: 1) Calculate the pH of a sulfuric acid solution of molarity: 0.001 molL-1
H 2SO 4 ( aq ) o 2H ( aq ) SO 4
2
( aq )
From the above balanced equation, 1 mole of sulfuric acid produces 2 H+. Therefore, [H+] = 2(0.001) = 2.0 x 10-3 mol L-1. pH
log[H ] log[2.0 u 10 3 ] 2.69897... # 2.70 (2 s.f.)
2) Calculate the pH of 0.02molL-1 acetic acid if 3% ionises in water.
CH 3 COOH ( aq ) o H ( aq ) CH 3 COO ( aq ) [H + ] = 0.03 x 0.002 = 6.0 x 10 -5 mol L-1 pH
log[H ] log[6.0 x 10 -5 ] 4.2218... # 4.2 (2 s.f .)
3) Calculate the pH of the solutions produced by: a) Dissolving 2 g of NaOH and making volume to 2L Robert Lee Chin
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HSC Chemistry Summary Module 2- The Acidic Environment
mass 2 0.05moles molar mass (23 16 1) moles 0.05 2.5 u 10 2 molL1 volume 2
moles NaOH molarity
NaOH ( aq ) o Na ( aq ) OH ( aq ) Sodium hydroxide is a strong base, so will ionise completely. [OH - ] 2.5 u 10 2 molL1 [H ] u [OH - ] 10 -14 [H ]
pH x
10 -14 [OH - ]
10 -14 2.5 u 10 2
4.0 u 10 13
log[H ] log[4.0 u 10 13 ] 12.39794... # 12.4 (3 s.f .) Solve problems and perform a first-hand investigation to use pH meters/probes and indicators to distinguish between acidic, basic and neutral chemicals
Investigation: Using pH meter to distinguish between acidic, basic and neutral chemicals Aim: To use pH meters to determine the pH of certain chemicals Equipment: -pH meter -buffer solution -0.1 M of the following solutions: HCl, NaCl, NaOH -test tubes and test tube rack Method: 1/ Calibrate the pH meter using the buffer solution. 2/ Place 25mL of each solution into separate test tubes 3/ Record the pH of the each of the solutions by placing the tip of the probe into the solution. Rinse tip of probe with distilled water in between substances. Results:
Substance HCl NaCl NaOH
Robert Lee Chin
pH 1 7 13
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HSC Chemistry Summary Module 2- The Acidic Environment Conclusion: -HCl has a low pH and is therefore acidic. -NaCl has a medium pH, and is therefore neutral -NaOH has a high pH, and is therefore basic. x
Plan and perform a first-hand investigation to measure the pH of identical concentrations of strong and weak acids
Investigation: measuring the pH of identical concentrations of strong and weak acids Aim: To determine the pH of identical concentrations of weak and strong acids. Equipment: -pH meter -buffer solution -0.1 M of the following solutions: CH3COOH, HCl, H2SO4 Method: 1/ Calibrate the pH meter using buffer solution. 2/ Place 25mL of each solution into separate test tubes 3/ Record the pH of the each of the solutions by placing the tip of the probe into the solution. Rinse tip of probe with distilled water in between substances. Results: Substance CH3COOH HCl H2SO4 x
pH 3 1 0.7
Degree of ionisation 1% 100% 100%
Gather and process information from secondary sources to write ionic equations to represent the ionisation of acids
For strong acids that ionise 100% e.g. HCl, H2SO4, the ionic equation can be written with one arrow from left to right. For example, the ionic equation for a strong monoprotic acid, hydrochloric acid: HCl ( aq ) o H ( aq ) CL ( aq ) Ionic equation for a strong diprotic acid, sulphuric acid: 2 H 2 SO 4 ( aq ) o 2H ( aq ) SO 4 ( aq ) For weak acids, the equation will be written with reversible arrows to indicate that the equilibrium point has a significant amount of both reactants and products. For example, the ionisation of carbonic acid: H2CO3(aq) 2H+(aq) + CO32-(aq)
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HSC Chemistry Summary Module 2- The Acidic Environment For organic acids such as acetic acid and citric acid, the H+ from the =COOH group ionises. For example, the ionisation of acetic acid: CH3COOH (aq) CH3COO-(aq)+ H+(aq) x
Describe the difference between a strong and a weak acid in terms of an equilibrium between the intact molecule and its ions
A strong acid is one where nearly 100% of the molecules ionise in an aqueous solution. For example, hydrochloric acid: HCl ( aq ) H 2 O ( l ) o H 3 O ( aq ) Cl ( aq ) A weak acid is one that does not fully ionise. For example, when hydrogen cyanide is placed in water, less than 1% ionises and an equilibrium situation is set up HCN(aq) + H2O(l) H3O++CNAdding more water increases the degree of ionisation, but the concentration will not increases because there is more solution. x
Use available evidence to model the molecular nature of acids and simulate the ionisation of strong and weak acids _ _
+ _ _ _ _
_ _
+ +
_ _
+ +
_ _
Acid molecule
+
Strong acids e.g. HCl, H2SO4 disassociate almost entirely in water to form positive hydrogen ions (protons) and anions _ _ _ + _ +
_ _
+ + +
_ _
_ _ _ _
Weak acids do not dissociate entirely in water. Most of the acid molecules remain in the solution.
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HSC Chemistry Summary Module 2- The Acidic Environment x
Describe acids and their solutions with the appropriate use of the terms strong, weak, concentrated and dilute
A concentrated solution is one in which the total concentration of solute species is high. A 10molL-1 solution would be called concentrated. A dilute solution is one in which the total concentration of solute species is low. A strong acid is one in which all the acid present in solution has ionised to form hydrogen ions. There are few neutral acid molecules left. A weak acid is one in which only some of the acid molecules present in the solution have ionised to form hydrogen ions. Weak acids for equilibrium reactions with water x
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
If we compare different acids of equal concentrations, the pH will depend on the number of H+ ions that ionise in solution. For strong acids, the acid will ionise 100% e.g. HCl (aq) o H Cl For weak acids such as acetic and citric acid, only 1% ionises e.g. CH3COOH CH3COO- + H+ x
Gather and process information from secondary sources to explain the use of acids as food additives
Microorganisms such as clostridium botulism produce toxins in food, which can cause severe food poisoning. Acids are used to preserve foods because many microorganisms including yeasts, moulds and bacteria, are pH sensitive and are killed when exposed to acidic conditions. Some acids act as antioxidants by retarding the oxidation of certain chemicals in food e.g. vitamin C. The addition of acids extends the shelf life of many processed food products including dairy, baked goods, cured meats, fruits and vegetables. In some cases, acids also give a unique flavour to some foods e.g. picked vegetables, sweet and sour sauces Common acids used as preservatives include acetic acid in vinegar. Vinegar is often used to pickle vegetables for canning. Other foods may be fermented to produce acids by bacteria or fungi. For example, the fermentation of milk to yogurt converts lactose to lactic acid. Sulfur dioxide is the only acidic oxide is commonly used as a food preservative. It is added to foods such as dried fruit and preserved deli meats because it maintains the appearance of the food and helps prevent rotting. Other acids used as food preservatives include phosphoric acid, citric acid, propanoic acid, benzoic acid, sodium nitrate and sorbic acid. Robert Lee Chin
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HSC Chemistry Summary Module 2- The Acidic Environment x
Identify data gather and process information from secondary sources to identify examples of naturally occurring acids and bases and their chemical composition Name/s
Composition
Acid/base
pH in Natural source/s natural form 1. 0 Gastric juices
Hydrochloric acid Acetic acid
HCl
a c id
CH3COOH
a c id
2.0-3.0
Salicylic acid Tartaric acid
C6H4(OH)COOH C4H6O6
a c id a c id
caffeine Calcium carbonate Sodium hydroxide Ammonia
C8H10N4O2 CaCO3
base base
2.5 2.5 (2.5% solution) 8.0 9.0
NaOH
base
13
NH3(aq)
base
11.5
Vinegar, fruits and vegetables Plants e.g. rhubarb Wine fermentation Coffee, tea Chalk, marine shells, eggshells Burnt ashes, lye water All living organisms
4. Definition of acids and bases x Outline the historical development of acids including those of: -Lavoisier -Davy -Arrhenius Antoine Lavoisier (1743-94) was a French chemist who demonstrated that combustion reactions involved oxygen. Experimentation led him to believe that acids were composed of two substances, one of them being oxygen. He believed oxygen was present in all acids and as responsible for acidity Humphry Davy (1778-94) was an English chemist who was famous for electrolysis experiments. In 1810, he decomposed hydrochloric acid and found it was composed of hydrogen and chlorine and did not contain oxygen. He observed that metals could displace hydrogen from acids to form salts. He concluded that all acids contain hydrogen. In 1884, Swedish chemist Svante Arrhenius (1859-1927) proposed definitions for acids and bases. He suggested that acids were neutral substances that produce hydrogen ions as the only poitive ion in an aqueous solution and that bases are substances that produce hydroxide ions as the only negative ion in an aqueous solution. His theory was limited because it applied only to aqueous solutions, only accounted for substances containing hydrogen or hydroxide ions and could not explain amphoteric substances such as zinc oxide
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HSC Chemistry Summary Module 2- The Acidic Environment x
Gather and process information from secondary sources to trace developments in understanding and describing acid/base reactions Scientist/s
x
Lavoisier
Acid definition n/ A
Base definition n/A
Davy
n/ A
n/A
Arrhenius
Acid ionises in water to form protons and anion
Base ionises in water to form hydroxide ion and cation
BrönstedLowry
Proton donators
Proton acceptors
Notes Oxygen is present in all acids and is responsible for the acidity Acids contain hydrogen. They do not have to contain oxygen -Applies only to aqueous solutions -Only accounts for substances containing hydrogen or hydroxide ions -Cannot explain amphoteric substances Acids must contain hydrogen Each acid has a conjugate base
Outline the Brönsted-Lowry theory of acids and bases
An acid-base reaction is one in which a proton is transferred from an acid to a base. An acid is defined as a proton acceptor while a base is a substance that accepts a proton from an acid. The Brönsted-Lowry theory is advantageous because it is able to explain: - non-aqueous reactions -why some salts can act as acids or bases -why some substances are amphoteric. It forms the basis for the qualitative treatment of acid-base equilibriums and pH calculations x x
Describe the relationship between an acid and its conjugate base and a base and its conjugate acid Identify conjugate base pairs
Every acid has its conjugate base a substance with exactly one less proton. An acid transfers a proton to its conjugate base in an acid-base reaction. Together, this acid and base form a conjugate pair. The products of an acid-base reaction are another acid and base, so there are always two conjugate acid-base pairs in each reaction. For example: CH 3COOH ( aq ) H 2 O ( l ) o CH 3COO H 3O acid 1
base 2
Robert Lee Chin
base 1
acid 2
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HSC Chemistry Summary Module 2- The Acidic Environment CH3COOH is an acid and CH3COO- is its conjugate base. Together they form a conjugate pair. Similarly, H3O+ is an acid, and H2O is its conjugate base. A strong acid has a weak conjugate base and a strong base has a weak conjugate acid: Strongest Acids Acid Base Weakest Bases HCl C lH2SO4 HSO4 HNO3 NO3H3O+ H2O HSO4 SO42H3PO4 H2PO4CH3COOH CH3COOH2CO3 HCO3H2S HSNH4+ NH3 H2O OHWeakest Acid
x
HSOH-
S2O2-
Strongest Acid
Identify amphiprotic substances and construct equations to describe their behaviour in acidic and basic solutions
Amphoteric= a substance that can act as both an acid and a base e.g. zinc and aluminium oxide Amphiprotic (as defined by the Brönsted-Lowry theory) = an amphoteric substance that can donate or accept protons i.e. it can act as a conjugate acid and a conjugate base. Amphiprotic substances include water (H2O), ammonia (NH3), hydrogen carbonate ion (HCO3-) & phosphane (PH3). H+
Water acting as a proton donator: H2O(l) + NH3(g)
NH4+(aq) + OH-(aq)
H+
Water acting as a proton acceptor: H2O(l) + HCl(aq) x
H3O+(aq) +Cl-(aq)
Identify neutralisation as a proton transfer reaction which is exothermic
In a neutralisation reaction, hydrogen ions and hydroxide ions form water. Neutralisation reactions usually occur between a strong acid and a strong base. Robert Lee Chin
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HSC Chemistry Summary Module 2- The Acidic Environment For example, the reaction between hydrochloric acid and sodium hydroxide: NaOH ( aq ) HCl ( aq ) o H 2 O ( l ) NaCl ( aq ) base
acid
water
salt
The net ionic equation shows that it is a proton transfer reaction: OH ( aq ) H ( aq ) o H 2 O ( l ) Almost all neutralisation reactions are exothermic, releasing about 57 kJ of heat per mole of water formed i.e. ΔH=-57kJ. x
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
Experiment: Preparation of a standard solution of hydrochloric acid Aim: To prepare a primary standard solution of sodium carbonate and use it to determine the concentration of a hydrochloric acid solution. Equipment: -Small beaker -Pipette (20mL) -250 mL volumetric flask -Wash bottle with distilled water -Approx 0.1 molL-1 HCl (100mL) -Stirring rod
-Electronic scale-Burette -3 x 250mL conical flasks -Burette clamp and retort stand -Pipette filler -Approx 2.0g dried Na2CO3 -Suitable indicator (methyl red)
Safety: Wear safety glasses. Hydrochloric acid is corrosive, so avoid contact with skin. If contact occurs, wash well with soap and water. Do NOT pipette by mouth: use pipette filler. Many indicators are poisonous and should be handled with care Method: A Preparing the primary standard 1/ accurately weigh 2.0 g of sodium carbonate in a small beaker 2/ Add a small amount of distilled water to beaker and stir ti dissolve sodium carbonate. Use a wash bottle with distilled water to wash out all the sodium carbonate solution into the funnel. 3/ Rinse the beaker and stirring rod with small amounts of distilled water and transfer the wash water into the flask 4/ Add distilled water to the volumetric flask until it is about two-thirds full. Fit the stopper and shake to dissolve all the sodium carbonate. When all is dissolved, top up the flask until the bottom of the meniscus is level with the mark.
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HSC Chemistry Summary Module 2- The Acidic Environment B Standardising the hydrochloric acid/ solution 5/ Rinse the burette with distilled water and then with HCl, discarding the rinsings. 6/ Set up the burette with burette clamp and fill with HCl. Record the starting volume. 7/ Rinse the pipette with sodium carbonate solution and place in a clean 250 mL beaker. This exact volume is known as the aliquot. Add 2-3 drops of indicator. 8/ Place the flask under the burette and run HCl into the flask, swirling continuously until colour changes. This is the end point. Record end volume. 9/ Calculate the volume of HCl used and record 10/ Refill the burette and repeat steps 5-9 at least twice more until three precise results are obtained. C Standardising the 10% vinegar solution 11/ Measure 25mL vinegar using measuring flask 12/ Place into clean conical flask and fill to 250mL solution 13/ Rinse the burette with distilled water and then with 10% vinegar, discarding the rinsings. 14/ Set up the burette with burette clamp and fill with 10% vinegar. 15/ Rinse the pipette with sodium carbonate solution and place in a clean 250 mL beaker. This exact volume is known as the aliquot. Add 2-3 drops of indicator. 16/ Place the flask under the burette and run 10% vinegar into the flask, swirling continuously until colour changes. This is the end point. 17/ Calculate the volume of 10% vinegar used and record 18/ Refill the burette and repeat steps 5-9 at least twice more until three precise results are obtained. Results: Mass Na2CO3 used = 2.0g Volume Na2CO3 used (L) Initial burette reading (L) Final burette reading (L) Volume HCl used (L) Volume 10% vinegar used
Trial 1 0.02
Trial 2 0.02
Trial 3 0.02
Trial 4 0.02
0.031
0.05
0.03
0.03
0.007
0.03
0.01
0.011
0.024
0.02
0.02
0.019
0.091
0.074
-
-
Calculations: Molarity of Na2CO3 solution: mass moles Na 2 CO 3 molar mass
Volume Na 2 CO 3 Molarity Na CO Robert Lee C2hin 3
250mL moles volume
2.0 2(23) 12 3(16) 0.25L 1
53 0.25
0.0754molL1
1 moles 53 29
HSC Chemistry Summary Module 2- The Acidic Environment
Moles Na2CO3 used in titration: moles concentration u volume 0.0754... u 0.02
0.0015094... # 0.0015 moles (3 s.f .)
Average volume of acids used: 0.024 0.02 0.02 0.019 Average volume HCl used 0.02075L 4 0.091 0.074 Average volume 10% CH 3 COOH used 0.0825L 2
Na 2 CO 3( aq ) 2HCl ( aq ) o CO 2 ( g ) H 2 O ( l ) 2 NaCl ( aq ) 1 mole of sodium carbonate neutralises 2 moles of hydrochloric acid.? 0.0015...moles sodium carbonate will neutralise 2(0.0015...) moles hydrochloric acid. Na 2 CO 3( aq ) 2CH 3 COOH ( aq ) o CO 2 ( g ) H 2 O ( l ) 2 NaCH 3 COO ( aq ) 1 mole of sodium carbonate neutralises 2 moles of acetic acid.? 0.0015...moles sodium carbonate will neutralise 2(0.0015...) moles acetic acid. Molarity of acid solutions:
2(0.0015...) 0.14548761... # 0.015 molL1 (3s.f .) 0.02075 moles 2(0.0015...) Molarity of 10% CH 3 COOH solution 0.003659... # 0.004 molL1 (3s.f .) volume 0.0825 Molarity of 100% CH 3 COOH solution 0.376 molL-1 (3 s.f .) Molarity of HCl solution
x
moles volume
Describe the correct technique for conducting titrations and preparation of standard solutions
A titration is a method used to experimentally determine the molarity of a solution. It is a volumetric analysis technique. A solution of a known concentration called the standard solution is added to a solution of unknown concentration until the neutralisation reaction is complete.
Standard Solutions Robert Lee Chin
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HSC Chemistry Summary Module 2- The Acidic Environment There are two types of standard solutions: Primary standard and secondary standard. Standard solutions are also known as titrants. A primary standard is a solution that has been made by dissolving an accurately measured mass of solute in a small amount of solvent and made to the required volume in a volumetric flask. A secondary standard is a solution whose concentration has been found by titrating against a primary standard. For a chemical to be suitable to prepare as a primary standard solution it must: *Be a water soluble solid *Be obtainable in pure form *Have an accurately known formula *Be stable in air To prepare the standard solution: 1) Accurately weigh a calculated amount of solid 2) Dissolve it in water 3) Transferring ALL of the dissolved solid to a volumetric flask 4) Adding water to prepare a known volume of solution The reaction is complete at the equivalence or end point. This is when the molar ratio of H+ to OH ions is equal i.e. Moles acid u cv acid Moles base u cv base . The solution changes colour at the equivalence point. Selecting the Indicator The equivalence point is not always at pH=7. The salts formed by combining different strong and weak acids have acidic or basic properties. Thus, an indicator must be chosen that changes colour near the equivalence point Indicator litmus Bromothymol blue Methyl orange Phenophalein
Acid red yellow Re d clear
Colour in:
Base blue blue
pH change 6.0-8.0 6.2-7.6
Yellow pink
3.1-4.4 8.3-10.0
Strong acid and strong base:
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HSC Chemistry Summary Module 2- The Acidic Environment
14
Equivalence point
pH 7
Volume acid in base Strong acid and weak base: 14
pH 7 Equivalence point
Volume acid in base Weak acid and strong base:
14 Equivalence point pH 7
Volume acid in base
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HSC Chemistry Summary Module 2- The Acidic Environment Titration Equipment The main equipment includes: -pipette and burette to measure volume of reactants -Flask to mix reactants
20mL pipette
The pipette measures a fixed volume of solution to provide a fixed number of moles of one reactant. Before using, it must be rinsed with distilled water, then with the solution to be used. Rinsing with the solution removes any water which would alter the volume and hence, number of moles of the solution being drawn. The solution should be drawn so that the bottom of the meniscus is in line with the etched line. The volume measured by the pipette is called an aliquot. The flask should be rinsed only with distilled water. It does not matter if it is wet, as this will not alter the number of moles of solution used (this has already been accurately measured by the pipette).
Burette
The burette allows the exalt volume of the reactant required to reach the equivalence point. Like the pipette, it must first be rinsed with distilled water, then with the solution to be used. The volume delivered by the burette is called a titre. Titration Procedure 1. Ensure all equipment is cleaned and rinsed with correct liquid 2. Add one solution to the burette 3. Use pipette to measure volume of other solution 4. Transfer this to a conical flask 5. Add a few drops of the suitable indicator 6. Perform a rough titration to find endpoint. 7. Repeat carefully until at least three readings within 0.1 mL of each other are obtained 8. Perform calculations.
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HSC Chemistry Summary Module 2- The Acidic Environment x
Choose equipment and perform a first-hand investigation to identify the pH of a range of salt solutions
Investigation: Determining the pH of salt solutions Aim: To select equipment and perform an experiment to determine the pH of various salt solutions Equipment: Test tubes and test tube rack Universal indicator Demineralised water 5mL of a 0.1 M solution of the following salts: -Ammonium Chloride (NH4Cl) -sodium carbonate (Na2CO3) -sodium hydrogen sulfate (NaHSO4) -potassium nitrate (KNO3) -ammonium acetate (CH3COONH4) Safety: Wear safety glasses Method: 1/ Place 5mL of each salt solution and demineralised water into separate test tubes 2/ Use a few drops of universal indicator to determine the pH Results: Salt Ammonium chloride Sodium carbonate Sodium Hydrogen sulfate Potassium nitrate Ammonium acetate Demineralised water
NH 4 Cl ( aq ) o NH 4
( aq )
Formula NH4Cl
Experimental pH 6.0
Acid/neutral/base Weakly acidic
Na2CO3 NaHSO4
11.0 3.0
Moderately acidic Strongly acidic
KNO3 NH4CH3COO H2O
7.5 7.0 7.0
neutral Neutral neutral
Cl ( aq )
weak acid
weak base
Na 2 CO 3(sq ) o 2 Na ( aq ) Co 3 neutral
2
( aq )
strong base
NaHSO 4 ( aq ) o Na (aq)
HSO 4
( aq )
neutral strong acid
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HSC Chemistry Summary Module 2- The Acidic Environment x
Qualitatively describe the effect of buffers with reference to a specific example in a natural system.
A buffer is a chemical that controls the pH of a solution. Buffer solutions are usually a mixture of a weak acid and the salt of that acid or a weak base and the salt of that base e.g. hydrogen carbonate ion (HCO3-) and carbonate ion (CO32-). In a buffer, equilibrium is established between the weak acid and its conjugate base. There are two reactions involved. For example, the buffer system involving hydrogen carbonate ions and carbonate ions: If an acid is added to the buffer, hydrogen ions are removed: H ( aq ) HCO 3 ( aq ) o H 2 CO 3 ( aq ) If a base is added to a buffer, hydroxide ions are removed: 2 OH ( aq ) HCO 3 ( aq ) o H 2 O ( l ) CO 3 ( aq ) In living organisms, blood is a buffered solution containing carbonic acid and sodium bicarbonate: CO2(g) + H2O(l) H2CO3(aq) H+(aq) + HCO3-(aq) The more CO2 that dissolves, the more H+ will form. The equilibrium shifts to the left to resist this change. If the pH is increasing, more carbonic acid will dissolve and the equilibrium will shift to the right to minimise the change. Carbonic acid is a weak acid, so a change in hydrogen concentration will not affect the pH much. x
Analyse information from secondary sources to assess the use of neutralisation reactions as a safety measure or to minimise damage in accidents or chemical spills
Neutralisation reactions are commonly utilised for safety in laboratories where acids and bases are used. When selecting the appropriate neutralisation reagent, the following factors need to be considered: -speed of neutralisation -need for reagent that will not be harmful in excess -safe to handle and store -cost -ability to use for both acids and bases i.e. amphiprotic Common neutralising reagents include the hydrogen carbonate ion found in sodium hydrogen carbonate. When the carbonate ion is used for acid spills, it combines with a hydrogen ion, forming water and carbon dioxide: H ( aq ) HCO 3 ( aq ) o H 2 O ( l ) CO 2 ( g )
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HSC Chemistry Summary Module 2- The Acidic Environment When it is used for base spills, it combines with the hydroxide ion, forming water and the carbonate ion: 2 OH ( aq ) HCO 3 ( aq ) o H 2 O ( l ) CO 3
x
Identify a range of salts which form acidic, basic or neutral solutions and explain their acidic, neutral or basic nature
A salt formed by a strong acid and a strong base is neutral e.g. NaCl: Sodium chloride forms by reacting sodium hydroxide with hydrochloric acid. When sodium chloride dissolves in water, the sodium chloride forms Na+ and Cl-. The water forms H+ and OH-: NaCl(aq) H2O(l)
Na+(aq) + Cl-(aq) H+(aq) + OH-(aq) H+ Cl-
ClOH-
Na+ Cl
Na+
-
OH H
H+
+
-
Na+
OH-
Sodium ions are attracted to hydroxide ions, forming NaOH, a strong base that completely ionises. Chlorine ions are attracted to hydrogen ions, forming HCl, a strong acid which ionises completely. Thus, the concentration of hydrogen ions equals the concentration of hydroxide ions and the solution is neutral. A salt formed by a strong acid and a weak base is acidic e.g. NH4Cl: Ammonium chloride forms by reacting hydrochloric acid and ammonium hydroxide. When ammonium chloride dissolves in water, it forms NH4+ and Cl-. The water forms H+ and OH-. NH4Cl(aq) NH4+(aq) + Cl-(aq) + H2O(l) H (aq) + OH-(aq)
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HSC Chemistry Summary Module 2- The Acidic Environment Cl-
H+
Cl-
NH4+
H+
NH4OH
Cl
NH4+
-
OH-
H+
NH4OH OH-
Ammonium ions are attracted to hydroxide ions, forming NH4OH, a weak base. Chlorine ions are attracted to hydrogen ions, forming hydrochloric acid, a strong acid. Thus, the concentration of hydrogen ions is greater than the concentration of hydroxide ions and the solution is acidic. A salt formed by a weak acid and a strong base is basic e.g. CH3COONa: Sodium acetate forms by reacting acetic acid and sodium hydroxide. When sodium acetate dissolves in water, the sodium acetate forms Na+ and CH3COO-. The water forms H+ and OH-. CH3COONa (aq) Na+(aq) + CH3COO-(aq) + H2O(l) H (aq) + OH-(aq)
CH3COOH Na+
OHH+ Na
+
OH-
CH3COOOH-
OHNa+ Na+
CH3COOH
Sodium ions are attracted to hydroxide ions, forming NaOH, a strong base. Acetate ions are attracted to hydrogen ions, forming acetic acid, a weak base. Thus, the concentration of hydroxide ions is greater than the concentration of hydrogen ions and the solution is basic.
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HSC Chemistry Summary Module 2- The Acidic Environment
4. Esterification x
Describe the differences between the alkanol and alkanoic functional groups in carbon compounds
The alkanol functional group is the hydroxyl group, –OH. The alkanoic acid functional group is the carboxyl group, –COOH. The carboxyl group makes alkanoic acids polar molecules. Alkanoic acids form hydrogen bonds so are water soluble. The general formula for alkanoic acids is RCOOH, where R represents the akyl chain with the formula CnH2n+1. For example, pentanoic acid is C4H9COOH. Remember that one C from the akyl group is already included on the COOH functional group. Thus, for pentanoic acid (5 carbons), the number of carbon atoms in the R group is 4. Structural formula for Pentanoic acid: CH3
CH2
CH2
O CH2
C
-COOH functional group
OH x
Identify the IUPAC nomenclature for describing the esters produced by reactions of straight chain alkanoic acids from C1 to C8 and straight chain primary alkanols from C1 to C8
Esters are named with the alkanol functional group first, replacing the suffix ‘anol’ with “yl”. When writing the chemical formula, the acid comes first, followed by the –COO– functional group, then the alkanol group. The suffix ‘oic acid’ is replaced by ‘anoate’ For example: onc. H 2SO 4 C 4 H 9 OH (l) C 4 H 9 COOH c o C 4 H 9 COOC 4 H 9 ( l ) H 2 O ( l ) onc. H 2SO 4 pentoic acid c o butyl pentonoate H 2 O
Butanol alkanol Carbon atoms 1 2 3 4 5 6 7 8
alkanoic acid First Part Methyl Ethyl Propyl Butyl Pentyl hexyl Heptyl octyl
Robert Lee Chin
ester
water
Second Part Methanoate Ethanoate propanoate butanoate pentanoate Hexanoate Heptanoate octanoate 38
HSC Chemistry Summary Module 2- The Acidic Environment x
Explain the differences in melting point and boiling point caused by straight chain alkanoic acid and straight chain primary alkanol structures
Alkanoic acids
Alkanols Alkanes, alkenes & alkynes
Boiling poi nt
olleacr ualned do not form hydrogen bonds. They Alkanes, alkenes anCdaarblkoynnaetsoamres npoern-mpo only have weak dispersion forces and thus low boiling points. The high melting and boiling points of alkanols are due to the hydrogen bonding of the O in one molecule, and the H from the -OH group in a nearby molecule. They also have 1 centre of polarity, forming dipole bonding. Akyl chain Akyl chain
O
H
H
O
Hydrogen bond
In alkanoic acids each carboxyl group is able to form two strong hydrogen bonds. This is because they have two O groups and plenty of hydrogen groups. They have 2 centres of polarity and dipole bonding. This gives alkanoic acids an even higher boiling point than their corresponding alkanol. O Akyl chain
C
C OH
x
HO Akyl chain
O Hydrogen
Identify esterification as the reaction bbeotw ndesen an acid and an alkanol and describe, using equations, examples of esterification
Esters are produced in a condensation reaction between an alkanol and an alkanoic acid called esterification. This is a reversible reaction that forms an equilibrium situation. Robert Lee Chin
39
HSC Chemistry Summary Module 2- The Acidic Environment A molecule of water is condensed out during the reaction. Use of tracers indicates that the OH comes from the alkanol and the O from the acid. For example: O CH3
C
OH
+
ethanoic acid
x
O
Conc. H2SO4
HO
CH2
→
CH3
C
O
CH3
+
H2 O
Heat
methanol
methyl ethanoate
water
Describe the purpose of using acid in esterification for catalysts
Esterification is a slow process that does not reach completion at room temperature because it forms an equilibrium situation. Concentrated sulphuric acid is used as a catalyst. Also, the acid is hydroscopic meaning it absorbs water, shifting the equilibrium to the right and producing more ester. x
Explain the need for refluxing during esterification
The reflux system consists of a reflux condenser fixed onto a reaction flask. The reaction flask is heated to speed up the reaction. The reflux condenser prevents the loss of volatile reactants (i.e. alcohol) or products during heating. It is open at the top to avoid the dangerous build-up of pressure. x
Outline some examples of the occurrence, production and use of esters
Esters occur widely in living things, esp. fruits and flowers. They are responsible for many aromas and flavours in foods. The aroma and flavour of fruits and flowers is from a complex mixture of esters and other compounds but esters are responsible for the main aroma. Fats and oils are the long chain “fatty acid” esters of the triple-alcohol molecule, glycerol. These form the long-term energy storage sites in plants and animals. Name Ethyl ethanoate Ethyl butanoate Pentyl ethanoate Octyl ethanoate
Structure CH3COOC2H5 CH3COOC4H9 C4H9COOC2H5 C7H15COOC2H5
Use Nail polish remover Pineapple Banana Orange
Esters are produced via the reflux of an alcohol, alkanoic acid and a catalyst on an industrial scale. Robert Lee Chin
40
HSC Chemistry Summary Module 2- The Acidic Environment Uses of esters include artificial flavours for drinks and processed foods, industrial solvent in the plastic industry and in cosmetics such as shampoos and lipstick. x
Identify data, plan, select equipment and perform a first-hand investigation to prepare an ester using reflux.
Experiment: Preparation of an ester Aim: To prepare an ester using reflux Equipment: The following alkanols: -methanol -Ethanol -Concentrated sulfuric acid -1.0 molL-1 Na2CO3 solution -retort stand and clamps -conical flask -Funnel -separating funnel -boiling chips (not marble) -condenser with rubber tubing -Bunsen burner -clay triangle -tripod
The following alkanoic acids: -butanoic acid -Glacial acetic acid -Salicylic acid
Safety: Wear safety glasses at all times. Sulfuric acid is corrosive. Clean up spills immediately and wash affected area with large quantities of water. Organic chemicals are flammable. Do not allow liquids or vapours to come into contact with sparks or flames and avoid inhaling vapours.
Method: 1/ Add a few boiling chips to the funnel. Place 8mL of one alkanol, 24mL of one of the alkanoic acids and 1Ml concentrated sulfuric acid into a flask using funnel.
Robert Lee Chin
41
HSC Chemistry Summary Module 2- The Acidic Environment 2/
Set up equipment as shown below:
water out Condenser
Retort stand and clamps
water in Flask with reaction mixture and boiling
chips
Bunsen, tripod clay & clay triangle
3/
Connect tubing to tap and condenser and turn on water so a uniform flow is achieved.
4/
Heat the mixtures over a steady Bunsen burner for 30 minutes (do not let mixture boil too vigoursly) and allow to cool for 5 minutes. Turn off water.
5/
Carefully remove the flask and pour contents into a separating funnel containing 15mL water. Stopper the funnel and shake. Allow layers to separate, drain off and discard aqueous layer.
6/
Add 15mL sodium carbonate solution. This will neutralise the acid and prevent the reaction from going backwards. Shake and drain the lower layer. The ester should be in the separating funnel.
7/
Carefully smell the ester and describe the smell.
Results:
Alkanol ethanol
Alkanoic acid Acetic acid
Ester Ethyl acetate
ethanol methanol
Butanoic acid Salicyclic acid
Ethyl butanoate Methyl salicycoanate
Robert Lee Chin
Aroma Nail polish remover; acetone Banana, fruity Oil of wintergreen
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