Chapter 8/ 18 Flashcards
Acids and bases
Bronsted-Lowry acid
Donates a proton (is a proton donor)
- all contain hydrogen as they must be able to donate a hydrogen ion to another species
Bronsted- Lowry base
Accepts a proton (is a proton acceptor)
- all must have a lone pair of electrons to form a bond w/ the hydrogen ion donated to them by B-L acid
Proton
Refers to a hydrogen ion (H+) which exists as the hydronium ion (H3O+) in aqueous solution
Amphiprotic
Species that are able to act as either a Bronsted-Lowry acid or a Bronsted-Lowry base depending on what it is reacting with eg. water
Classification of acids
Can be monoprotic, diprotic or triprotic depending on no. of hydrogen ions they release into solution
- Hydrochloric acid and nitric acid = monoprotic acids
- release only one hydrogen ion in solution - Sulphuric acid = diprotic acid
- releases two hydrogen ions in solution
NB/ ethanoic acid is monoprotic despite having 4 hydrogen atoms
- only one of these is released when ethanoic acids dissociates in solution
Alkali
Bases that are soluble in water
Reaction at equilibrium
Reversible reactions
- when a reaction can occur in both directions w/ forward reaction occurring at same rate as reverse reaction = reaction is at equilibrium
Conjugate acid-base pair
A pair that differs by a proton (H+)
Conjugate acid of a base: add one proton (H+)
- have one more positive charge
Conjugate base of an acid: remove one proton (H+)
- have one less positive charge
Charge on the species must have changed as protons have a positive charge
Amphiprotic species
- a species that can act as both a Bronsted- Lowry acid and a Bronsted-Lowry base depending on what it’s reacting with
- amphiprotic refers to Bronsted-Lowry theory ONLY
Amphoteric species
Refers to substances that can act as either acids or bases
- use isn’t limited to Bronsted-Lowry
Acid reactions with metals
acid + metal –> salt + hydrogen gas
- unreactive metals (those below hydrogen in the activity series) don’t react w/ dilute acids
Test for hydrogen gas
- ignite a small volume of gas in an inverted boiling tube
- flammable gas burns w/ a distinctive sound, known as a ‘squeaky pop’
What are some bases?
Metal oxides
Metal hydroxides
Aqueous ammonia
NB/ alkalis are soluble bases
Neutralisation reactions
Metal oxide + acid –> salt + water
Metal hydroxide + acid –> salt + water
Acid reactions with metal carbonates/ hydrogencarbonates
Metal carbonate + acid –> salt + water + carbon dioxide
Metal hydrogencarbonate + acid –> salt + water + carbon dioxide
Test for carbon dioxide
- tested for by bubbling the gas through limewater [an aqueous solution of calcium hydroxide Ca(OH)2 ]
- if CO2 is present, limewater turns a ‘milky’ colour as a solid precipitate of calcium carbonate is formed
Neutralisation reactions
An acid reacts w/ a base (or an alkali) to produce a salt and water
- reaction between an acid and a base is exothermic
- heat is released, enthalpy change is negative
Equation for a neutralisation reaction
Acid + base (or alkali) –> salt + water
NB/ salt produced in the reaction depends on the parent acid and base that react
Antacid tablets
Ease symptoms of excess stomach acid that causes indigestion or heartburn
- active ingredients of most antacid tablets: metal carbonates or hydrogen carbonates or insoluble metal hydroxides
- react with excess stomach acid in neutralisation reactions to relieve symptoms of heartburn
Acid-base titration
Determines unknown conc. of an acidic or a base solution using a solution of known conc. (standard solution)
- an acid-base indicator is used to determine point where acid neutralises the base or vice versa
- indicators are chosen due to their bright and easily identifiable colour changes
Method for an acid-base titration
- burette is filled w/ alkali of known conc. (titrant)
- titrant is added to acid of unknown conc. (analyte) until end-point is reached
- end point is signified w/ use of an indicator
Thermometric titrations
- can be carried out using a simple calorimeter
- heat is released when an acid and base react together- exothermic reaction
- enthalpy change that occurs = enthalpy change of neutralisation
- it is the enthalpy change when one mole of water is formed in the reaction of an acid and a base
Strong acid + strong base reaction
When any strong acid + strong base react together, enthalpy change = -57 kJ/mol
- because net reaction is the same
- one mole of water is being formed from one mole of H+ ions and one mole of OH- ions
Define pH
pH of a solution = a measure of the conc. of hydrogen ions [H+] in a solution
pH = -log[H+(aq)]
pH scale
- an inverse scale
- higher the conc. of H+ ions in solution, lower pH value
- lower the conc. of H+ ions, higher the pH value
- scale is logarithmic to base 10 (change of 1 pH unit corresponds to 10x change in Hydrogen ion conc.)
pH and conc. of hydroxide ions
- higher conc. of OH- ions = higher pH value
- lower conc. of OH- = lower pH value
Acids and alkaline solutions
Water
- [H+] = [OH-]
- pH = 7
- neutral solution
Acid solution:
- [H+] > [OH-]
- pH lower than 7
- litmus turns red
Basic solution:
- [OH-] > [H+]
- pH higher than 7
- litmus turns blue
Calculating conc. of H+ ions in a solution from its pH value
pH = -log [H+(aq)] [H+] = 10^-pH
pOH of a solution
pOH = -log[OH-]
Acid-base indicator
A weak acid or a weak base in which the dissociated and undissociated forms have different colours
- depending on conc. of H+/OH- ions in solution, one colour or the other is shown
Litmus paper
- a common indicator (can’t be used to determine a pH value for a solution)
- a qualitative indicator
- made from a natural source, lichens
- in acidic solutions, litmus appears red
- in alkaline solutions, litmus appears blue
Methyl orange
Acid: red
Alkali: yellow
Phenolphthalein
Acid: colourless
Alkali: pink/purple
Universal indicator
A mixture of indicators that produces different colours in solutions of different pH
- pH of aqueous solutions can be measured using universal indicator as either a solution or in paper form
Indicators
May be used to measure pH of a solution
- makes use of the fact that their colour changes w/ pH
pH meter
Produces a more accurate method of measuring pH, provided they’re correctly calibrated
- electrode of pH meter is placed in solution to be tested
- a voltage is generated that’s converted into a pH meter reading displayed on the screen
- pH meter is calibrated using buffers of known pH, usually of pH 4.0, 7.0 and 10.0
Water as a conductor
Pure water is a v. poor conductor of electricity
- means it must have a low conc. of mobile ions responsible for electrical conductivity of solutions
- water molecules do dissociate (or ionise) to a v. small extent
- conc. of water is constant
Ionic product constant of water
Kw = [H+][OH-]
- Kw is temp. dependent
- so, at a certain temp., product of [H+] and [OH-] is a constant
- so, as [H+] of a solution increases, [OH-] decreases
- neutral solution has a [H+] = [OH-]
Strong acids and bases
For a strong acid or base, [H+] or [OH-] is equal to its conc.
- because strong acids and bases are completely ionised in solution
Strength of an acid or base
Acids and bases can be classified as strong or weak
- strength of an acid/ base refers to its degree of dissociation in aqueous solutions
- strong acids and bases completely dissociate in solution
- weak acids and bases only partially dissociate in solution
Dissociate
Means to break apart an ionic compound into its constituent ions
Concentrated solutions
Have a high concentration, measured in mol/dm^3
Strong acids
When a strong acid dissolves in water, it completely dissociates into ions
- all molecules of acid react w/ water to produce hydronium (H3O+) or hydrogen (H+) ions
HA (g or l) –> H+ (aq) + A- (aq)
Can also be written involving formation of hydronium ion (H3O+)
HA (g or l) + H2O (l) –> H3O+ (aq) + A- (aq)
- position of equilibrium lies to the right, means that strong acid completely dissociates into its constituent ions
- solution of a strong acid contains virtually no undissociated HA molecules
- conc. of H3O+ ions in solution = initial conc. of strong acid
NB/ for a strong acid, conc. of H+ ions = initial conc. of strong acid
Organic acids
- weak acids
- when water is added, only a small % of its molecules dissociate to form H3O+ ions
- an equilibrium is established, where majority of acid molecules don’t dissociate into ions
- position of equilibrium lies to left
Dissociation of a weak acid
HA (aq) H+(aq) + A-(aq)
Can also be written as:
HA(aq) +H2O (l) H3O+(aq) + A-(aq)
Strong base
One which completely dissociates into ions in aqueous solution
- strong bases include metal hydroxides of group 1
General equation:
MOH(aq) –> M+(aq) + OH-(aq)
- position of equilibrium for dissociation of a strong base lies to the right
Conc. of OH- ions = initial conc. of strong base
Weak bases
All bases, except for hydroxides of group 1 and 2, are weak
- made up of molecules that react w/ water to release hydroxide ions
General equation for their dissociation:
B(aq) + H2O BH+(aq) + OH-(aq)
- at equilibrium, most of the weak base molecules have not dissociated into ions
- position of equilibrium lies to the left
Weak acids and bases
- for weak acids and bases, equilibrium conc. of H3O+ ions or OH- ions are much lower than for strong acids and bases
- these differences in conc. can be used to distinguish between solution of weak and strong acids and bases
Equimolar solutions
Solutions that have equal concentrations
- Weak and strong acids and bases of the same concentration
- distinguished due to their different properties eg. pH, electrical conductivity, and in case of acids, by their reactions w/ active metals
Acids, bases and conductivity
Weak acids and bases have a lower conc. of mobile ions in solution- as they only partially ionise in solution
- hence, are a poor conductor of electricity
Strong acids and bases have a higher conc. of mobile ions in solution- completely ionise in solution
- good conductors of electricity
NB/ when comparing electrical conductivity of weak and strong acids and bases, solutions must be of equal conc.
Acids, bases and reactivity
Bases do not react with metals
Acids and conjugate bases
Strong acids - weak conjugate bases
- a stronger acid will ionise more
- position of equilibrium lies to the right
- conjugate base is reactively weak
Weak acids- strong conjugate bases
- conjugate base is relatively strong
- equilibrium lies to the left
Acid deposition
Refers to all types of acid precipitation
- can be classified as dry deposition eg. acidic gases and solid particles or wet deposition e.g. acid rain
Acidity of rainwater
Natural rainwater is acidic- pH of 5.6
- acidity is due to presence of carbon dioxide in the atmosphere
- carbon dioxide reacts w/ water to form carbonic acid
Acid rain
- has pH of less than 5.0
- formed when acidic gases eg. sulphur dioxide and oxides of nitrogen dissolve in water in atmosphere to produce sulphuric acid and nitric acid
Formation of nitrogen monoxide
- During lightning storms by reaction of nitrogen and oxygen (2 predominant atmospheric gases)
- By reactions that take place in internal combustion engines
- nitrogen reacts directly w/ oxygen at high temp. to produce nitrogen monoxide (a free radical- has an unpaired electron in valence shell of nitrogen atom)
- takes place in both petrol and diesel engines; level of pollution is greater from diesel engines as operating temp. is higher
- use of a catalytic converter in vehicle exhaust fumes reduces level of these emissions
Nitrogen oxide undergoes further oxidation in atmosphere to produce nitrogen dioxide (also a free radical)
- nitrogen dioxide reacts w/ water to form nitric and nitrous acid
Sulphur dioxide
- occurs naturally in the atmosphere as it’s released during volcanic eruptions
- also produced when coal and oil that contain sulphur are burned in power stations
- sulphur is a contaminating presence in coal and fuel oil used in power stations
- sulphur (present in coal or fuel oil) is oxidised during combustion process and forms sulphur dioxide
Photochemical oxidation
- sulphur dioxide can also undergo photochemical oxidation in the atmosphere
- process takes place in water droplets where sulphur dioxide is dissolved, catalysed by particulates eg. soot and fine metallic particles
- reaction is complex, and isn’t a direct reaction w/ atmospheric oxygen
- interactions w/ hydroxyl radicals, ozone or hydrogen peroxide are involved
- sulphur dioxide and sulphur trioxide react w/ water in the atmosphere to form sulphurous and sulphuric acid
Dry deposition
- acidic particles and gases fall to ground via dust and smoke in absence of precipitation
- this deposition can be washed into streams, lakes or rivers, causing harm to biological systems
Environmental impact of acid deposition
- can have serious consequences for living organisms eg. fish and trees
- can enter aquatic ecosystems either directly eg. rain or indirectly as run-off
- pH below 5.5, fish can’t survive, and prevents fish eggs from hatching
- aluminium ions leached from soil by acid rain, enter rivers and streams and have harmful effects on fish gills
- fish exposed to aluminium ions secrete excess mucous around gills, preventing oxygen uptake, leading to death by asphyxiation
- loss of some species can cause knock-on effect through food chain- adversely affecting other organisms
- biodiversity within a lake or river can be reduced by acid deposition as they become dominated by acid-tolerant species
pH of soil
- determines which species of plants grown in a certain location
- aluminium is present in soil at high pH values as insoluble aluminium hydroxide
- when pH falls, aluminium ions are released into soil where it can damage roots of plants, preventing uptake of water
- acid deposition also has potential to leech essential plant nutrients from soil
Acid deposition and effect on buildings
Statues and other architectural structures are often made from limestone or marble- both forms of calcium carbonate
- calcium carbonate can react w/ sulphuric acid forming calcium sulphate, carbon dioxide and water
- structures of iron, steel or aluminium can also be damaged by acid deposition
- rate of rusting, being an electrolytic process, is increased by greater conc. of ions in acidic rainwater
Acid deposition and effects on human health
- acid rainwater and moisture in air can adversely affect mucous membranes and lungs
- causes irritation and possibly exacerbating symptoms for people w/ asthma and other respiratory conditions
Post-combustion methods
A method of reducing amount of acid deposition
- involves removing sulphur oxides from exhaust gases once they have been formed by reacting w/ a base
Pre-combustion methods
A method of reducing amount of acid deposition
- involves removing sulphur before coal is combusted
Hydrodesulphurisation
A pre-combustion method by which sulphur is removed from refined petroleum products eg. gasoline, jet fuel, kerosene and diesel fuel before combustion
- takes place at high temp. and pressures in presence of a catalyst
- sulphur is removed from product in form of hydrogen sulphide
Flue gas desulphurisation
- a post-combustion method used in power stations
- levels of sulphur dioxide emissions in flue gases can be reduced by passing them into a flue gas desulphurisation tower
- in this tower, gases are passed through a sprayed aqueous suspension of calcium carbonate and calcium oxide
- product is calcium sulphite, which is further oxidised to produce calcium sulphate
Rectification of acid deposition in soils and lakes
Soils and lakes whose pH has been significantly lowered by acid deposition can be treated w/ calcium hydroxide or limestone
- neutralises acidity and raises pH
Redox reaction
A reaction that involves both oxidation and reduction
- The two processes always take place together
- if one reactant is oxidised during a reaction, another must be reduced
Oxidation
- gain of oxygen
- loss of hydrogen
- loss of electrons
- results in an increase in oxidation state
Reduction
- loss of oxygen
- gain of hydrogen
- gain of electrons
- results in a decrease in oxidation state
Oxidising agent
A species that oxidises other species
- readily accept electrons
- reduced in the course of the reaction
NB/ an oxidising agent is reduced
Reducing agent
A species that reduces other species
- readily donate electrons
- oxidised in the course of the reaction
NB/ a reducing agent is oxidised
What are oxidation states?
Oxidation states that are assigned to atoms in a compound give a measure of control that an atom has over electrons in the compound
- concept of oxidation states assumes that atoms lose or gain electrons according to their electronegativity values, regardless of whether they actually form ions or not
- if an atom loses control over electrons- oxidised
- if an atom gains control over electrons- reduced
Written with a sign (+/-) and a numerical value
+ = atom has lost electron control
- = atom has gained electron control
Numerical value= takes into account no. of electrons that atom has lost or gained control over
Define oxidation state
A number which, together with its sign, indicates the gain or loss of electron control of an atom during a reaction
- oxidation NUMBER is a numerical part of that value
- it is used as a Roman numeral in the chemical formula of a compound
Rules for determining oxidation states
- Free elements (O2, Cl2, N2) have oxidation state = 0
- Sum of oxidation states of all atoms in a compound = net charge on compound
- Alkali metals (Li, Na, K, Rb and Cs) in compounds = +1
- Fluorine in compounds is always -1
- Alkaline earth metals (Be, Mg, Ca, Sr, Ba and Ra) and Zn = +2
- Hydrogen in compounds = +1, except in certain metal hydrides = -1
- Oxygen in compounds = -2, except in peroxides = -1
- Halogens in compounds = -1
- Charge on a metal ion = same as its oxidation state
- Sum of oxidation states in a polyatomic ion must add up to charge on ion
Exceptions to oxidation state
Some atoms can show variable oxidation states
- includes transition elements and non-metal elements
Disproportionation
Occurs when same species is oxidised and reduced simultaneously during a reaction to form 2 different products
- eg. catalytic decomposition of hydrogen peroxide
Lewis acid-base reactions
the transfer of an electron pair, with the subsequent formation of a coordinate bond
Bronsted- Lowry theory of acids and bases
acids and bases are defined in terms of the transfer of protons (H+ ions)
Lewis theory of acids
According to the Lewis theory, an acid is an electron pair acceptor and a base is an electron pair donor
Lewis acid
an electron pair acceptor
- electrophiles
Lewis base
an electron pair donor
- nucleophiles
Bonding in complex ions according to Lewis theory
Complex ions are formed when a ligand uses a lone pair of electrons to form a coordinate covalent bond w/ a central metal ion
- in this complex ion, the ligands act as Lewis bases, and the central metal ion acts as a Lewis acid
Dimer
a compound formed by the reaction of 2 similar monomers
