Chemistry depth study PDF

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Yr 11 chemistry depth study...

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Chemistry Depth Study Pre-excursion research: Module 5: Inquiry question: Describe and analyse the processes involved in the dissolution of ionic compounds in water To understand the processes involved in the dissolution of ionic compounds in water, we need to first understand the molecular structure of ionic substances. Here, the charged particles are arranged in a crystal lattice structure alternating between positive (cations) and negative (anions) ions electrostatically attracted to each other. Due to this, there are no covalent bonds between each of the individual atoms and they are only held together by electrostatic forces. Also, it is important to understand that the water molecule is polar; that is one side (dipole) - the side with the large electronegative oxygen atom is, on average, more negative than the other dipole - the one with the smaller, less electrostatic hydrogen atoms. This is as the oxygen atom is more electrostatic than the hydrogen atoms and thus the electrons being shared spend a greater proportion of their time near this dipole, making it electrostatic. The electrons therefore spend less time near the hydrogen dipole, making it positive. This property of water (polarity) makes it a good solvent. When an ionic substance is submerged in water, the water molecules form their own electrostatic bonds with the individual atoms of the crystal lattice, called an ion-dipole bond. Water can attach to both positive and negative ions as it has a positive and a negative dipole. The ion in question is then pulled away from the crystal lattice and then attracts other water molecules, binding them to it (the number of water molecules required depends on the charge exhibited by the target ion); when this is complete the ion is known as a aquo-cat/anion. Inquiry Question: predict the formation of a precipitate given the standard reference value for K sp The Ksp value refers to the solubility constant for a certain ionic compound at standard temperature (25 degrees Celsius). This is provided when required on the datasheet.

Since dissolving a substance is an equilibrium, equilibrium calculations can also be used to determine the position of the equilibrium and thus predict the favoured direction in which the reaction will occur. Since the equilibrium in question refers to the dissolution of a substance in water, assuming that the forward reaction is the dissolution, a forwards favoured reaction will mean that more substance will dissolve in water. Since it is assumed that the substance in question is already fully dissolved, there will be no change. On the other hand, a backwards favoured reaction will mean the exact opposite, with the dissolved substance becoming a solid again, or precipitation. From the above, if the current position of the equilibria is calculated and compared with a known critical value (Ksp), then the movement of the reaction can easily predicted and thus the formation of a precipitate; for the below assume that Qsp is the calculated current equilibrium constant x+ x− AB (s) ↔ A(aq) + B (aq) If: -

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Qsp < Ksp then the current equilibrium is below the critical point and will attempt to move towards it in accordance with LCP; the forward reaction is favoured and thus no precipitate is formed. In other words, more of the substance can be dissolved and therefore the solution is NOT saturated. Qsp = Ksp then the current equilibrium is at the critical point and is thus at a dynamic equilibrium. The rates of reaction of the forwards and backwards reactions are exactly equal to each other and thus no macroscopic/observable change will occur. The solution is saturated as it cannot hold any more solute. Qsp > Ksp then the current position of the equilibrium is above the critical point and in accordance to LCP will attempt to decrease the concentration of the ions and will thus favour the backwards reaction. This will lead to the formation of the solid compound, which is by definition a precipitate. - Note: In some cases, if the solution is heated to a higher temperature then it is made saturated and cooled with ABSOLUTELY no disturbance, then the solution will not precipitate. However, this is known as a supersaturated solution and is highly unstable. Here, if any disturbance is to occur (even the disturbance caused by dust is enough), then a precipitate will form.

Since for an ionic compound to dissolve in water both ions must be soluble, the solubility of ionic compounds can be decreased even further by the addition of another ionic compound which contains one of the ions in question; the shifts in equilibrium can be exploited to reduce the solubility of a certain ion. One application of this is in intestinal CAT scans, where the normally radio transparent gut tissue is made opaque through the ingestion of a solution of barium sulfate, which acts as a radio-contrastive agent. However, barium sulfate is extremely toxic in its aqueous form as it can be absorbed into the bloodstream. To mitigate this, the drink must contain a suspension of barium sulfate particles, with a minimal amount of aqueous ions. To achieve this, a less toxic substance containing a sulfate ion is added to maintain the slurry suspension of solid barium sulfate so the patient is not poisoned.

Module 6: Inquiry Question: investigate applications of neutralisation reactions in everyday life and industrial processes Common uses: - Treating acidic soil - Plants cannot grow well in acidic soil - For example lime fertilisers, powdered lime (CaO) or limestone (CaCO3 ) are added to neutralise - Treating basic soil - Compost of vegetables and leaves decompose to liberate carbon dioxide gas - Preventing tooth decay - Decaying food particles produces acid which causes tooth decay - Alkaline toothpaste and brushing neutralises the acid - Baking powder helps bake cakes (I know this seems like a very cheap use but then again, it is a very everyday life use!) - Baking powder contains sodium hydrogen carbonate and a weak acid - The sodium hydrogen carbonate and water react (similar to acid acting on carbonate) - Forms carbon dioxide which helps the cake rise - Shampoo and Conditioner - Shampoos are mild alkaline, causing the small scales on hair to become unmanageable - Hair conditioners are mild acid which neutralises the alkali and cause scales to close up (and makes hair super smooth) Real life industrial applications: - Treating effluents (liquid waste and sewerage) with bases before being released into river streams - Note that liquid wastes contain sulphuric acid which is highly corrosive - Neutralising acidic gases (for example carbon dioxide and sulphur dioxide) that are waste released from power stations, helps minimise pollution - Use of ammonia solution (NH4 OH, basic solution of pH 11-12) to neutralise lactic acid produced by bacteria in latex - Useful in the rubber industry, to prevent the coagulation of latex Use in medicine involving anti-acids: - Aluminium Hydroxide (Al(OH)3) treats heartburn, stomach problems, acid indigestion and reduces phosphate levels - Magnesium Hydroxide (Mg(OH)2) neutralises the excess acid in the stomach - Mild acids which are alkaline in nature (for example vinegar) to treat wasp stings - Bases which are acidic in nature (for example baking soda) to treat bee stings and ant bites

Inquiry Question: Write ionic equations to represent the dissociation of acids and bases in water, conjugate acid/base pairs in solution and amphiprotic nature of some salts, for example: sodium hydrogen carbonate & potassium dihydrogen phosphate Definitions (to begin off with): - Amphiprotic means capable of acting both as a proton donor and proton acceptor - An amphiprotic substance is capable of acting as both a Bronsted-Lowry acid or Bronsted-Lowry base - Bronsted-Lowry acid/base - An acid is a substance that donates one or more protons or hydrogen ions (H+ ) - A base is a substance that accepts one or more protons - Bronsted-Lowry acid/bases just mean that bases become the mirror companion of acids (as one donates and the other receives it) - Conjugate acid/base pairs - Consider H A + B ⇌ HB + + A− (where HA denotes an acid and B is base molecule) - In forward reaction, HA is proton donator, losing H+ whereas B is proton acceptor - Therefore reactant acid produces a base while reactant base produces an acid, all on the product sides Sodium Hydrogen Carbonate is an amphiprotic salt, hence can act as both a Bronsted-Lowry acid or base. - Bronsted-Lowry acid (proton donor) - FULL IONIC: Na+ + HCO3 (s) + H2O (l) ⇌ H3O+ (aq) + CO32- (aq) + Na+ (aq) - When dissolved in water, it reacts and donates a proton (H+ ) - Na+ is a spectator ion so can be cancelled - Conjugate pair of HCO3 → CO3 2- and H2O → H3O+ - Bronsted-Lowry Base (proton acceptor)  + Na+ (aq)  ⇌ 2H+ (aq) + CO32- (aq) + OH- (aq) - FULL IONIC: Na+ + HCO3 (s) + H2 O (l) - When dissolved in water, it reacts with water and accepts a proton - Once again, Na+ is a spectator ion - Conjugate pair of HCO3 → CO3 2- and H2O → OHPotassium dihydrogen phosphate is another amphiprotic substance. - Bronsted-Lowry acid (proton donor) - 2- - FULL IONIC: K+ (aq) + H2PO + H3O+ (aq) + K+ (aq)  4 (aq) + H2O (l) ⇌ HPO4 (aq) - When dissolved in water, it donates a proton - K+ is a spectator ion - Bronsted-Lowry base (proton acceptor) - FULL IONIC: K+ (aq) +H2PO4 - (aq) + H2O (l) ⇌ 3H+ (aq) + PO4 3- (aq) + OH- (aq) + K+ (aq) - When dissolved in water, it accepts a proton - K+ is a spectator ion

Dissociation of strong acid (HCl in this case) in water forms conjugate acid and bases - As HCl acts as Bronsted-Lowry acid, the water acts as the base (proton acceptor)

Dissociation of weak acid (CH3COOH for example) in water

Dissociation of strong base (NaOH) in water - Once again forms conjugate pairs and spectator ions are irrelevant in net equation

Dissociation of weak base (NaF) in water

Inquiry Question: Describe the importance of buffers in natural systems In biological systems, the pH of the environment is extremely important; for instance, enzymes (biological catalysts) are most efficient at a very specific pH level, with the enzyme being denatured in environments which are either too basic or too acidic. In our bodies, the blood is kept at a constant pH level of 7.4; whilst this sounds relatively simple, in practice it is very difficult as many of the biological processes produce hydrogen ions in the form of a hydronium molecule, whilst others produce bicarbonate ions. By definition, this increases/decreases the pH of the blood as these processes are independent of each other. To maintain the target pH level, the blood contains a variety of buffers: A buffer is a chemical solution containing a weak acid and its conjugate base; this allows it to maintain a constant pH, even with the addition of bases or acids. This is due to the acid and its conjugate base being held in a dynamic equilibrium. [Assume that the acid is on the left whilst the base in on the right] When an acid is added, the concentration of H + ions increase; in accordance to LCP, the system opposes that change and favours the forward reaction in an attempt to counteract the change. Hence, the concentration of the weak acid on the left is decreased as it is converted to its conjugate base. Remembering the fact that the conjugate base contributes to the pH of the system, an increase in the concentration of the base counteracts the change induced by the addition of the acid, hence maintaining or buffering the pH of the solution. And vice versa for the addition of bases. In the human body, the main buffer for the blood is the presence of a bicarbonate ion ( H 2 CO 3 ) and carbonic acid ( HCO 3− ) system, with the equilibrium reaction of: + − H 2 CO 3 + H 2 O ↔ HCO 3 + H 3 O When any acidic substance enters the bloodstream, the bicarbonate ions neutralize the hydronium ions forming carbonic acid and water. Carbonic acid is already a component of the buffering system of blood. Thus hydronium ions are removed, preventing the pH of blood from becoming acidic as shown below: + − HCO 3 + H 3 O → H 2 CO 3 + H 2O On the other hand, when a basic substance enters the bloodstream, carbonic acid reacts with the hydroxide ions producing bicarbonate ions and water. Bicarbonate ions are already a component of the buffer. In this manner, the hydroxide ions are removed from the blood, preventing the pH of blood from becoming basic; also shown below: − H 2 CO 3 + OH − → H CO 3 + H 2O In addition to this, there are also many other buffer systems in the body including (but not limited to): the phosphate buffer found inside cells and the haemoglobin buffer, which can either bind with protons or oxygen, regulating the acidity of the blood. However, this one is more complex as it involves a protein.

Module 7: Inquiry Question: Examine the environmental, economic and sociocultural implications of obtaining and using hydrocarbons from the Earth To begin, hydrocarbons are compounds that consist of only hydrogen and carbon atoms. They come in two types: - Acyclic hydrocarbons are hydrocarbons of linear structure. They are mainly derived from petroleum and natural gas - Cyclic hydrocarbons are hydrocarbons that have a cyclic/circular structure (they have phenyl ring, C6H5 which has three alternating C=C bonds) Sourcing hydrocarbons: - Petroleum and natural gas can be found at different levels under the ground - Formed when dead aquatic organisms are covered in sand, clay and mud sediments at the bottom of the sea - Due to the lack of oxygen in that environment, with high heat and high pressure conditions, the dead matter is converted into petroleum and natural gas - Petroleum is less dense than water and hence will situate on top of the water - The natural gas is less dense than the both and hence will be above both but become trapped underneath the impervious rocks - Hence, natural gas can be sourced from these impermeable/impervious rocks located under the Earth - Coal - Formed by the decomposition of dead plants and trees which are covered over by soil and mud - Trapped dead matter are in an environment lacking oxygen but has high heat and pressure, decomposing to peat which is compressed by being trapped which forms coal Uses: - Petroleum consists of hydrocarbons from C1 to C30 - Used in motor vehicles as fuel (octane) - Paraffin wax used in candles and machine lubricants (C20 and higher) - Natural Gas consists of light hydrocarbons (for example methane, ethane, propane, butane) - Coal (consists of rich carbon compounds) - Can be burnt in an anaerobic environment (no oxygen in environment) to produce methane

Environment: Only a small portion of hydrocarbons are toxic to the environment, not all. Also note that sourcing fuels can also damage the environment (for sources below the earth, drilling could damage the environment) - Obtaining petroleum and natural gas involves drilling the earth (as they are trapped underneath) - To drill underneath, drilling machines use lubricants that are dispersed and can pollute the ocean - Lubricant traces contain barium ions (as they drill into the sea, leaving these traces behind) which are toxic, interfering with enzyme activities and can kill organisms - Machine lubricants also leave potassium ions - Dead plants means less source of oxygen which is necessary for aquatic fish and bacteria to decompose dead matter - A high concentration of potassium can result in algae growing uncontrollably (eutrophication)→ algae can cover water surface hence less sunlight, necessary for plant survival - Leaving waste hydrocarbons are also toxic to aquatic animals - Increase in ion concentrations can disrupt osmotic balance in ecosystem → deforms aquatic fauna and flora - To locate good sources of hydrocarbon deposits, sound waves are sent throughout the sea - This noise pollution can disorient and disrupt living organisms → results in imbalance in the ecosystem and result in death of organisms - Once obtaining hydrocarbons they must be transported to the refinery through ships - Oil spillage can occur via collisions with rocks or other ships - Oil spilling also pollutes the waters and is toxic to aquatic life - Hydrocarbon in atmosphere - They can enter the human body and cause respiratory problems like irritation Economic: The many environmental problems and possibilities for problems and disasters to occur emphasises the preservation of aquatic ecosystems. Hence, to preserve this, a lot of money must be used and also people must note that aquatic organisms make up a large amount of economic revenue (sashimi, sushi, fish and chips, seaweed).

For example, in the Philippines, aquatic organisms make up over 550$ USD. So, any spillage or harm done to the water environment will definitely take a harsh reduction in economic benefits (also social implications like poverty and reducing the country’s reputation for safe fish food and money). Sociocultural:

Workers that help drill will be involved with these hydrocarbon pollution and lubricant waste, all of which holds the chance of inhaling the toxic hydrocarbon. If problems occur to the aquatic environment, supply of fish will be scarce and price will not be affordable. Presence of toxic hydrocarbons also can damage the safety of our drinking water due to sea pollution. Treating this polluted water would incur more cost to consumers and to the industry.

Module 8: Inquiry Question: conduct qualitative investigations or data processing – using flame tests,precipitation and complexation reactions as appropriate – to test for the presence in aqueous solution of the following ions: – cations: barium (Ba2+), calcium (Ca2+), magnesium (Mg2+), lead(II) (Pb2+), silver ion (Ag+ ), copper(II) (Cu2+), iron(II) (Fe2+), iron(III) (Fe3+) – anions: chloride (Cl– ), bromide (Br– ), iodide (I – ), hydroxide (OH– ), acetate (CH3COO– ), carbonate (CO3 2– ), sulfate (SO4 2– ), phosphate (PO4 3– ) Flame tests take advantage of the fact that when exposed to sufficient heat energy as provided by the fire, certain electrons in certain cations will absorb the heat, moving to a higher energy level in the process. However, this higher energy orbital is unstable and the electron will invariably decay to a lower energy orbital, releasing the originally input energy as electromagnetic radiation in the form of a photon. Remembering that the wavelength of electromagnetic radiation corresponds to its energy level, it can be deduced that energy sources of the same energy will release a certain wavelength of light. Applying this to the flame test, if the solution being tested is pure, then the energy differential will be consistent and thus so will the wavelength of light being emitted. Since colour is simply a certain wavelength of electromagnetic radiation with a wavelength that is visible, we can work backwards and determine the analyte by comparing the emitted colour with a table of known colours. (We hope that the wavelength corresponding to the energy differential happens to lie in the narrow visible light spectrum). Expected colour values for required cations; Cation

Expected Colour

Barium 2+

(Take your pick) - yellow-green, apple-green, or lime-green

Calcium 2+

orange-red

Magnesium 2+ (Continued - Magnesium 2+

(none), but for burning Mg metal intense white

(Continued) - Cation

Expected colour

Lead(II) 2+

Blue/white - very beautiful actually. Faint, almost transparent flame.

Silver 1+

??? i cant seem to find colour for silver

Copper(II) 2+

Copper green

Iron(II) 2+

Gold, when very hot such as an electric arc, bright blue, or green turning to orange-brown. Bunsen burner is not sufficient to heat it up to such a temperature thus it is invisible/ colourless when tested in bhhs lab

Iron(III) 3+

Orange-brown (brown? Very bright)

Precipitation reactions are a displa...