chemistry chapter 4

pg 129 - acids and bases

• the ph scale is used to indicate wether a substance is acidic or alkaline

• you can measure ph using an indicator ( a dye that changes colour depending on wether the substance is above or below a certain Ph, e.g universal indicator)

• red = acid (0)

• purple = alkaline (14)

• green = neutral (7)

• you can also measure Ph with a Ph probe, which is attached to a Ph meter and meausres Ph electronically: place in substance and it will display a number, telling you exactly what Ph the substance is, making it more reliable than an indicator

• acids and bases neutralise eachother

• an acid is a substance that creates an aqueous solution with a Ph of less than 7, they form H+ ions (protons) in water

• a base is a substance that has a Ph greater than 7

• an alkaline is a substance that when dissolved in water (aqueous) will have a greater Ph than 7, it produces OH- ions in water

• acid + base = salt + water - this reaction is called neutrailisation

• can also be written as: H+ + OH- = H20 (water)

• when acids and bases react their product will have a neutral Ph (water has a Ph of around 7)

pg 130 - strong and weak acids

• acids will ionise in aqueous solutions

• to ionise means particles dissociate and release charged particles ( in this case protons)

• strong acids will ionise fully in aqueous solutions (water) meaning all particles will dissociate and release H+ ions

• weak acids will don’t fully ionise in aqueous solutions, meaning only a small amount of particles will dissociate and release H+ ions

• examples of strong acids: sulfuric, hydrochloric, nitric

• examples of weak acids: ethanoic, citric, carbonic

• ionisation of weak acids is a reversible reaction, as there is equilibrium between the dissociated and non dissociated particles

• reactions of acids involve H+ ions reacting with other substances, the higher the concentration of H+ ions, the faster the reaction will be, meaning that stronger acids will react faster than weaker acids even if their concentrations are the same

• the pH of a substance measures the concentration of H+ ions in the solution

• when the pH scale decreases by 1 (gets more acidic), the H+ ions increase by a factor of 10 (because the more acidic the solution is, the more H+ ions it will produce)

• e.g a solution that has a pH of 4 will have 10x more H+ ions than a solution with a pH of 5

• for a decrease of 2, the H+ ions will increase by 100x (and so on…)

• acid strength = amount of acid molecules ionising in water

• concentration is different to strength, concentration is how much acid there is in a volume of water

• pH will decrease with increasing acid concentration regardless of wether it’s strong or weak (the lower pH, the more acidic it is)

pg - 131 reactions of acids

• metal oxides and hydroxides are bases (a substance that has a pH greater than 7)

• even bases that aren’t soluble in water will take part in neutralisation reactions

• all metal oxides and hydroxides will react with acids to produce water and salt

• the salt produced depends on the type of acid and base that’s reacted- e.g hydrochloric acid + copper oxide will produce copper chloride and water

• metal carbonates are also bases and will produce carbon dioxide when reacted with an acid, as well as a salt and water

• for example, hydrochloric acid + sodium carbonate - sodium chloride + water + carbon dioxide

• practical to make soluble salts using an insoluble base:

• pick an insoluble base, such as an insoluble metal oxide, hydroxide or carbonate

• warm the dilute acid with a bunsen burner

• add the insoluble base a small bit at a time until it stops reacting/ when the base is in excess

• stir and if the excess solid sinks to the bottom, the acid has been neutralised

• filter out the excess solid to get the salt

• to get pure crystals, use crystalisation

pg 132 - the reactivity series

• used to compare the reactivity of substances against eachother

• the higher a substance is in the list the more reactive it is e.g potassium

• when a metals reacts with water or acid, they form positive ions, the higher up the list the substances are, the more positive ions they form/ they loose electrons easily

• some metals react with acid to create a salt and hydrogen gas

• the speed of the reaction is indicated by the rate in which the metals produces bubbles when reacting with the acid (this is hydrogen gas it’s giving off)

• the more reactive a metal is the faster the reaction will go

• some metals higher up the reactivity will react explosively e.g potassium

• less reactive metals react less violently

• you can measure reactivity with temp change, if you use the same mass and surface area of the metal each time, the more reactive the metal, the more temp change there is

• metals also react with water - metal + water - metal hydroxide + hydrogen

• reactive metals (potassium, magnesium, sodium) will react with water and less reactive metals won’t (zinc, copper, iron)

pg 133 - separating metals from metal oxides

• metals react with oxygen in a process called oxidation

• when metals gain oxygen, they become metal oxides, the metal oxides are often the ores that the metals need to be extracted from if we’re going to use them

• a reaction that separates a metal from its oxide is called a reduction reaction

• to extract a metal from a metal oxide, we use reduction, which is loss of oxygen (e.g copper oxide is reduced to just copper)

• reduction = loss of oxygen

• oxidation = gain of oxygen

• we can also separate metals from their oxides using reduction by carbon and electrolyisis

• electrolysis is very expensive and only used for metals that are more reactive (higher up in the reactivity series) than carbon, as then they can’t be reduced by carbon

• reduction by carbon works only with metals lower down than it in the reactivity series

• this is because carbon can only take oxygen from metals that are less reactive than itself

• in this reaction, the metal oxide is reduced by carbon, meaning the metal looses oxygen (reduction) and carbon gains oxygen (oxidation)

• reaction example : sodium oxide + carbon - sodium + carbon dioxide

• some metals are so unreactive they’re found in the earth as just metal (gold)

pg 134 - redox reactions

• OILRIG - oxidation is loss of electrons, reduction is gain of electrons

• oxidation/loss of electrons usually looks like this in an ionic equation - Mg - Mg2+ + 2e-, magnesium has gone from a mg ion to an mg atom, meaning it has lost electrons because it has lost its charge

• reduction/gain of electrons usually looks like this in an ionic equation - S2- - S + 2e- sulfur has gained 2 electrons, making it S2-, as when an atom gains electrons, they become negative charged, whereas when they loose them they become positive charged, and metal atoms are positive charged

• reduction and oxidation happen at the same times, hence the name REDOX

• in a redox reaction, one elements will reduce the other whilst the other oxidises the other

• displacement reactions are redox reactions

• other redox reactions include metals reacting with water and metals reacting with acid

• in displacement reactions, more reactive metals with displace less reactive metals from their compounds according to the reactivity series

• the metal atom is always the oxidised element (as it looses electrons)

• the metal ion is always the reduced elements (as it gains electrons)

• in ionic equations, there are spectator ions, these are the ions that remain unchanged during the reaction (they don’t get oxidised or reduced)

• when given a reaction like Mg + ZnCl2 - MgCl2 + Zn write out all of the ions separatly

• in this equation, it would look like: Mg + Zn+2 + 2Cl- - Mg+2 + 2Cl- + Zn

• chlorine is the spectator ion, as it doesn’t change so can be crossed out, leaving the ionic equation: Mg + Zn2+ - Mg2+ + Zn

• an ionic equation focusses only on the elements that get oxidised or reduced

• in this equation, the reduced element would be magnesium, as it has gained 2 electrons to become a charged ion Mg2+

• the oxidised element would be zinc, as it has lost electrons to be come a zinc atom

pg 135 - electrolysis

• electrolysis splits things up using electricity

• it is very expensive and requires lots of energy

• it can be used to extract metals from their oxides when they can’t be reduced by carbon due to them being more reactive

• since electrolysis requires electricity to split things up, the ionic compounds have to be molten, as when they’re solid they’re in fixed positions and can’t carry charge

• molten ionic liquids are always broken up into their elements, lead bromide is split up into lead and bromine

• you can tell which element is going to be attracted to the anode or cathode by using the periodic table, if the element is non metal, it has negative charge so will be attracted to the positive anode, as opposite charges are attracted to eachother

• anode = positive

• cathode = negative

• positive metal ions are reduced at the cathode

• negative non metal ions are oxidised at the anode

• REDCAT , ANOX - REDuction at CAThode, ANode OXidation

• adding cryolite will bring down the melting point - this is used in electrolysis of aluminium oxide

• hydrogen will be produced at the cathode if the metal is more reactive than hydrogen

• electrolytes are made of inert metals like platinum

pg 136 - electrolysis of aqueous solutions

• water (H2O) is also split during electrolysis

• it’s split into H+ ions and OH- ions (also called hydrogen and hydroxide)

• the H+ ions will be attracted to the negative cathode and the OH- ions will be attracted to the positive anode

• as well as the ionic compound in the solution, for example: copper sulfate

• this means that aqueous solution will contain, copper (Cu2+), sulfur (SO2-), hydrogen (H+) and hydroxide (OH-)

• H+ and Cu2+ will be attracted to the cathode and SO2- and OH- will be attracted to the anode

• at the cathode (-) : if H+ ions and metal ions are present then hydrogen will be produced, if the metal ion involved is less reactive than hydrogen then a solid, pure layer of the metal will be produced

• at the anode (+) : if halide ions are present then they will be produced at the cathode, if not then oxygen will be produced

• you can use gas tests to see what is produced

• oxygen : relights a glowing splint

• hydrogen : makes a squeaky pop

• chlorine : bleaches damp litmus paper white