C5 Chemical Changes Flashcards
Acid
When dissolved in water, its solution has a PH value less than 7. Acids are proton(H+ ion) donors, forms H+ ions in water. They ionise in aqueous solution.
Alkali
A base which dissolves in
water to form a solution with a Ph greater than 7. Alkalis form OH- ions in water. They are soluble hydroxides
Base
A base is a substance with a PH greater than 7.
Displacement reaction
A reaction in which a more reactive element takes the place of a less reactive element in one of its compounds or in solution. For example, when adding copper metal to a colourless silver nitrate solution. The more reactive metal will form ions(e.g. copper ions ) and the less reactive metal will end up being a regular solid(e.g. silver will be shown as grey crystals).
Equilibrium
The point in a reversible reaction at which the forward and backward rates of reaction are the same. This is because, as the reactants react, their concentrations fall, so the forward reaction will slow down. But as more products are made and the concentrations rise, the backward reaction speeds up. After a while the forward reaction goes at exactly the same rate as the backward one, but there’s no oferall effect. The amount of substances present in the reacting mixture remain constant as well as the concentrations of reactants and productd not changing.
Neutral
A solution with a PH value of 7 that is neither acid nor alkaline/something that carries no electrical charge
Neutralisation
The chemical reaction of an acid with a base in which a salt and water are formed.
Ore
Rock which contains enough material to make it economically worthwhile to extract the metal
Oxidation
A reaction where oxygen is added to a substance/electrons are lost from a substance,
Reduction
A reaction in which oxygen is removed from a metal oxide or electrons are gained.
e.g. 2CuO + C: 2Cu + CO2
Strong acids
These acids completely ionise in an aqueous
solution. The concentration of H+ ions is higher than in weak acids, so the rate of reaction will be faster and strong acids will be more reactive than weak acids of the same concentration.
Salt
A compound formed when some/all of the hydrogen in an acid is replaced by metal or ammonium ions
Weak acids
Acids that do not ionise completely in aqueous solutions. The ionisation of a weak acid is a reversible reaction. As the molecules of the weak acid split up to form H+ ions and negative ions, the ions recombine to form the original molecules again. A position of equilibrium is reached where both whole molecules(the majority) and their ions(the minority) are present.
Examples of acids(name 3)
Stomach acid
Vinegar
Lemon juice
Acid rain
Examples of neutral substancws
Pure water
Examples of alkalis
Washing-up liquid
Soap powder
Bleach
Caustic soda(drain cleaner)
What is an indicator?
An indicator is a dye that changes colour depending on whether it’s above or below a certain PH.
Wide range indicators contain a mixture of dyes that means they gradually change colour over the broad range of PH, useful for estimating the PH of a solution.
What is a Ph probe?
A pH probe attached to a pH meter is used to measure pH electronically. The probe’s placed in the solution you’re measuring and the Ph is given on a digital display as a numerical value, so is more accurate than an indicator
Neutralisation reaction(between acids and bases)
Acid + base: salt + water
Examples of strong acids
Sulfuric acid, hydrochloric acid and nitric acid
Examples of weak acids
Ethanoic, citric and carbonic acids
When only can equilibrium be reached
If the reversible reaction occurs in a closed system
What factors affect the position of equilibrium
The temperature - heating a reaction moves the equilibrium to the right(concentration of products greater than concentration of reactants) and cooling it moves it to the left(the concentration of the reactants is greater than the concentration of the products)
The pressure(only in equilibrium involving gases).
The concentration of the reactants and products
How does the PH of an acid or alkali relate to the concentration of H+ ions in the solution
For every decrease of 1 on the PH scale, the concentration of H+ ions increases by a factor of 10.
Factor H+ ion concentration changes by = 10-x
Concentration
How much acid there is in a certain volume of liquid. A dilute solution is not very concentrated
Acid strength
The proportion of the acid molecules that ionises in water
Why does acid concentration affect PH?
PH decreases with increasing acid concentration
Neutralisation reaction(acid + metal oxide)
Acid + metal oxide: salt + water
Neutralisation reaction(acid + metal hydroxide)
Acid + metal hydroxide: salt + water
Acid and metal carbonate reaction
Acid + metal carbonate: Salt + water + carbon dioxide
Practical(making soluble salts using an insoluble base)
1)Use an acid and insoluble base.
2)Gently warm the dilute acid using a Bunsen burner, then turn it off.
3) Add the insoluble base to the acid a bit at a time until no more reacts. You know when the acid has been neutralised, because, after stiring, the excess solid will sink to the bottom of the flask.
4)Filter out the excess solid to get the salt solution.
5)To get pure, solid crystals of the salt, gently heat the saturated solution using a water bath or electric heater to evaporate some water(to make it more concentrated), then stop heating and leave the solution to cool. Leave the solution at room temperature.Crystals of the salt should form which can be filtered out of the solution and dried(crystallisation).
Reactivity series of metals(with water + acid)
From most reactive to least reactive:
Very reactive metals:
Potassium
Sodium
Lithium
Calcium
Fairly reactive metals:
Magnesium
Aluminium
Carbon
Zinc
Iron
Tin
Lead
Not very reactive metals:
Hydrogen
Copper
Silver
Gold(so unreactive it’s in the Earth as a metal itself)
Metals’ reactions with water
Potassium, sodium, lithium, calcium: they fizz, giving off hydrogen gas, leaving behind
an alkaline solution of metal hydroxide
Magnesium, aluminium, zinc and iron: React slowly with water
Tin + lead: slight reaction with steam
Copper, silver + gold: no reaction
Metals’ reactions with dilute acid
Potassium, sodium, lithium, calcium: they explode with dilute acid
Magnesium, aluminium, zinc and iron: They fizz, giving off hydrogen gas and forming a salt
Tin + lead: React slowly with warm acid
Copper, silver + gold: no reaction
Metals’ reaction with acid(equation)
Acid + metal: salt + hydrogen
Reaction of metals with water
Metal + water: Metal hydroxide + hydrogen
e.g. Ca(s) + 2H2O(l): Ca(OH)2(aq) + H2(g)
Reduction with carbon(equation)
Metal oxide + carbon:(heat): Metal + carbon dioxide
Metals below carbon in the reactivity series can be extracted by reduction using carbon, but metals higher than carbon in the reactivity series are extracted using electrolysis.
Redox reactions
When oxidation(loss of electrons) and reduction(gain of electrons) occur at the same time.
e.g.
Iron atoms are oxidised to Fe2+ ions when they react with dilute acid: Fe + 2H+: Fe 2+ + H2
Iron atoms lose electrons and are oxidised by hydrogen ions: Fe - 2e-: Fe2+
Hydrogen ions gain electrons and are reduced by the iron atoms: 2H+ + 2e-: H2
Example of a redox reactio
Displacement reactions:
The metal ion always gains electrons and is reduced.
The metal atom always loses electrons and is oxidised.
e.g. Reaction of magnesium with zinc chloride:
Mg(s)+ ZnCl2(aq): MgCl2(aq) + Zn(s)
Ionic equation:
Mg(s) + Zn2+(aq)+ 2Cl-(aq): Mg2+(aq) + 2Cl-(aq) + Zn(s)
The chloride ions don’t change in the reaction, they’re spectator ions. Ionic equations here just focus on the substances oxidised or reduced.
Electrolysis experiments, testing gaseous products
Chlorine bleaches damp litmus paper, turning it white.
Hydrogen makes a squeaky pop with a lighted splint.
Oxygen will relight a glowing splint.
Test for hydrogen(reactivity test)
Goes out with a squeaky pop
Lit splint
Test for most Group 1 elements(alkali metals).
They react with non-metals to form ionic compounds, white solids that dissolve in water to form colourless solutions
Potassium reaction with water
The hydrogen produced burns violently with a lilac flame, quickly melts to form a ball, disappears rapidly, often with a small explosion
Sodium reaction with water
Fizzes rapidly, melts to form a ball,quickly becomes smaller until it disappears
Lithium reaction with water
Fizzes steadily, slowly becomes smaller until it disappears
Nitrate ion formula
-
NO
3
Sulfate ion formula
2-
SO
4
Carbonate ion formula
2-
CO
3
Hydrogencarbonate formula
-
HCO
3
Hydroxide formula
OH -
Ammonium formula
+
NH
4
Lead formula
2+
Pb
Zinc ionic formula
2+
Zn
Silver ionic formula
+
Ag
Acetate ion formula(acetic acid)
CH3COO-
Iron ion formula
Fe 3+
Whether it is worth extracting a particular metal depends on:
How easy it is to extract it from its ore.
How much metal the ore contains
The changing demands for a particular metal
How are ores extracted?
Ores are mined from the ground.
Some need to be concentrated before the metal is extracted and ourified. For example, copper ores are ground up into a powder. They are then mixed with water and a chemical making the copper compound repel water. Air is then bubbled through the mixture and the copper compound floats on top as a froth. The rocky bits sink and the concentrated copper compound is scraped off the top.
Chemical reduction
The removal of oxygen from a compound
Reduction of oxide by hydrogen equation
Metal oxide + hydrogen :(heat) Tungen + water(as steam)
Why are alkali metals never added to acid?
If the metal is very reactive, the reaction with acid is too violent to be carried out safely and alkali metals are very reactive.
Sulfuric acid formula
H2SO4
Nitric acid formula
HNO3
Reaction between magnesium and sulfuric acid
Chemical reaction:
Mg(s)+ H2SO4: MgSO4(aq)+ H2(g)
Ionic equation:
Mg(s) + 2H+(aq): Mg2+(aq)+ H2(g)
Hydrogen ions will be displaced from the solution by magnesium as magnesium is more reactive than hydrogen.
Mg(s): Mg2+(aq) + 2e-(A magnesium atom loses the two electrons from its outer shell(oxidation) and gives them to two hydrogen ions from the acidic solution, 2H+(aq), forming two H atoms). The hydrogen ions gain electrons(reduction).They hond to each other(sharing a pair of electrons in a covalent bond):
2H+(aq)+ 2e-: H2(g)
What is the charge on salts
No overall charge
What is the formula of ion sulfate?
The formulae of the ions are Fe3+ and SO4 2-.
LCM of 3 and 2=6
+6 charges from 2Fe3+ ions to cancel out -6 charges from 3 SO4 2- ions.
Formula of ion sulfate:
Fe2(SO4)3
Ammonium
NH4
How to collect pure, dry samples of a salt from neutralisation
1) Carry out the experiment with the indicator to see how much acid reacts with the alkali.
2)Run that volume of acid into the solution of alkali again, but this time without the indicator.
3)Crystallise and dry the crystals of the salt from the reaction mixture
Acid + alkali:
Metal/ammonium salts + water
Nitrate
NO3
Carbonate
CO3
What is the best way to dilute a solution?
By adding water
What is a problem of producing carbon dioxide gas in an experiment?
The carbon dioxide gas may escape
Why are alloys harder than oure metals?
Layers of atoms,in an alloy are distorted
Why do ionic compounds have a high melting point?
Giant structure
Strong electrostatic forces between oppositely charged ions
Experiment(to compare the reactivity of different metals)
1) Measure the same mass of each metal you are comparing.
2) Measure a set volume of a salt solution(e.g. silver nitrate solution) into a beaker.
3) Place one metal into the salt solution.
4)Using a thermometer, measure the temperature of the reaction mixture.
5) Then place the same mass of the next metal into the reaction mixture.
6)The greater the temperature change of the solution, the more reactive the metal is.
Experiment(to compare the reactivity of different metals with acids)
1) Add the metals to the dilute hydrochloric acid.
2) Measure temperature change/compare rate of bubbling
Properties of most transition metals
They form ions with different charges.
They have high melting points.
What is an accurate way of measuring the amounts of acid and alkali used to react with one another in an experiment?
A pipette
How to complete titration(measure amount of acid that reacts with an alkali)
1) Add potassium hydroxide solution to the conical flask.
2)Add a few drops of indicator.
3)Add the sulfuric acid from the burette until the colour of the indicator changes.
4) Read the volume of the burette
Chemical equation for reduction of
oxide to obtain tungsten by hydrogen
WO3 + 3H2(g): W(S)+3H2O(g)
Tungsten oxide + hydrogen: tungsten + water(steam)
Balanced symbol equation for reaction between zinc and hydrochloric acid
Zn(s)+2HCl(aq): ZnCl2(aq)+H2(g)
Are metal carbonates soluble in water?
No
Ionic equation for reaction between hydrogen and hydroxide ions in neutralisation
2H+(aq)+2e-: H2(g)
Ionic equation for displacement
Magnesium + copper sulfate:magnesium sulfate + copper
Mg(s)+Cu2+(aq): Mg2+(aq)+Cu(s)
Ionic equation for reaction between magnesium and sulfuric acid
Reaction: Mg(s)+ H2SO4(aq): MgSO4(aq)+ H2(g)
Ionic equation:
Mg(s)+ 2H+(aq): Mg2+(aq)+ H2(g)
Half equation - magnesium ato:s change into positive magnesium ions:
Mg(s): Mg2+(aq)+ 2e-
Half equation- magnesium atom loses two electrons from its outer shell to two hydrogen ions from the acidic solution, forming two H atoms which bond to make a molecule of hydrogen gas:
2H+(aq)+ 2e-: H2(g)
Observations in the reactions of two different metals with a certain acid?
If one metal’s more reactive than the other, when it reacts with a certain acid, there will be more bubbles produced
How to make soluble salts using an insoluble base
1) Pick the right acid and insoluble base.
2) Gently warm the dilute acid using a Bunsen burner, then turn off the Bunsen burner.
3) Add the insoluble base to the acid a bit at a time, until no more reacts(the base is in excess). You’ll onow when all the acid has been neutralised, because after stirring, the excess solid will sink to the bottom of the flask.
4) Then filter out the excess solid to get the salt solution.
5) To get pure, solid crystals of the salt, gently heat the solution using a water bath or a electric heater to evaporate some of the water and stop heating it, leaving the solution to cool. Crystals of the salt should form, which can be filtered out of the solution and then dried
