Chemistry End of Year Revision

Redox Reactions

A redox reaction is a reaction in which electrons move from one element to another, i.e. one element is reduced and the other is oxidized. It is the shortened form of reduction-oxidation reactions.

• one element is oxidized and the other is reduced simultaneously

• when one element is oxidized, the other is reduced. this is because when electrons are lost, they must be gained somewhere

Oxidation is the loss of electrons from an element in its free state or compound.

Oxidation can be defined as the:

• gaining of oxygen

• loss of hydrogen

• loss of electrons

• gain in oxidation number

Reduction is the gain of electrons from an element in its free state or compound.

Reduction can be defined as:

• the loss of oxygen

• the gain of hydrogen

• the gain of electrons

• the loss in oxidation number

The movement of metals from their atoms to their ions is oxidation.

The movement of metals from their ions to their atoms is reduction.

The movement of non-metals from their atoms to their ions is reduction.

The movement of non-metals from their ions to their atoms is reduction.

Even during partial transfers (e.g. ammonia that has a reversible reaction) one element is more electronegative than the other (nitrogen) i.e. it brings electrons towards it. This is the element that is reduced. The more electropositive one (hydrogen) loses these electrons and is oxidized.

Oxidation Number

The oxidation number of an element is how many electrons the element has gained, lost, or is sharing after a reaction or chemical bonding. Oxidation and Reduction can be defined in terms of oxidation number as well.

Oxidation, in terms of oxidation number, is the gain in the oxidation number of an element.

Reduction, in terms of oxidation number, is the decrease in the oxidation number of an element.

Do NOT get the oxidation number and charge number confused! For the oxidation number, the sign comes first (e.g. -2) while in charge number it would be a 2- charge.

For example, in the reaction of Na + Cl → NaCl, sodium goes from an oxidation state of 0 to an oxidation state of +1, which means it has been oxidized. Cl goes from an oxidation state of 0 to an oxidation state of -1, which means it has been reduced.

The oxidation number of elements can be calculated, but there are some rules that need to be followed.

1. The oxidation number of an element in its free state is zero

2. The oxidation number of an ion in a polyatomic ion is just the charge of the ion

3. The sum of oxidation numbers of a polyatomic ion is equal to the charge on the ion

4. The sum of oxidation numbers in a neutral compound is zero.

5. Hydrogen is always a +1 charge, except in hydrides. (e.g. MgH)

6. Oxygen is always a -2 charge, except in peroxides. (e.g. H2O2)

7. All elements in polyatomic ions and molecules are subject to change except hydrogen and oxygen. Sometimes the oxidation number of a compound is written in its name. (e.g. potassium dichromate(VI))

Some elements are named based on their electron number.

For example, sodium, chlorine and oxygen form four compounds named sodium chlorate, with differing amounts of oxygen in them.

NaClO - Sodium chlorate(I)

NaClO2 - Sodium chlorate (II)

NaClO3 - Sodium chlorate (III)

NaClO4 - Sodium chlorate (IV)

The oxidation number of the chlorine is different each time, which is why it is the one being changed.

When a polyatomic ion has a roman numeral in front of a particular atom, look at the name of the atom it is in front of and that will tell you its oxidation number.

• potassium permanganate (VII) - this means that the Mn in the KMnO4 has a 7+ charge

In Terms of Electron Numbers

Oxidation and reduction can also be defined in terms of electron numbers as stated above.

We can write half equations to simplify what happens during a redox reaction.

For example:

Mg (s) + H2SO4 → MgSO4 (aq) + H2O (l)

We know that Mg goes from an oxidation state of 0 (as it is in its free state) to an oxidation state of 2 (as calculated in the ion).

This means that it has been oxidized, i.e. it loses electrons. To represent this in a half equation, we write:

Mg - 2e- → Mg2+

But we can’t have subtraction signs in this, so we add 2e- to both sides.

Mg → Mg2+ + 2e-

Oxidizing and Reducing Agents

Oxidizing agents are elements or substances that cause other substances to oxidize. They themselves become reduced.

Reducing agents are elements or substances that cause other substances to be reduced. They themselves are oxidized.

There is a lot of important oxidizing and reducing agents to learn:

Oxidizing Agents

Other common oxidizing agents: oxygen, chlorine

• NOTE THAT POTASSIUM DICHROMATE (VI) AND MANGANATE (VII) NEED TO BE DISSOLVED IN DILUTE ACID FIRST! THAT’S WHY THEY ARE ACIDIFIED

Reducing Agents

Other common reducing agents: reactive metals, hydrogen, carbon, carbon monoxide

It is important to note that whenever a substance containing iodine is oxidized, the strong colour of brown/black may mask any colour changes.

• e.g. KMno4 + KI → Mn2+ I2

• here the potassium manganate (VII) turns colourless, but that is masked because the iodine is freed and it turns brown

Substances that can act as both acids and bases

H2O2/Hydrogen Peroxide

Hydrogen peroxide is originally an oxidizing agent and can act as an oxidizing agent (when it reacts with KI) or a reducing agent (when it reacts with stronger oxidizing agents like KMnO4).

SO2/Sulphur Dioxide

Sulphur dioxide is originally a reducing agent and can act as an oxidizing agent (when it reacts with stronger reducing agents like H2S) or as a reducing agent (when it reacts with K2Cr2O7-)

NaNO3/Sodium Nitrate

Sodium nitrate can oxidize the iodide ions to iodine in aqueous potassium iodide but it reduces the Mn7+ ion to Mn2+ in the KMnO4 solution.

How to test for oxidizing agents and reducing agents

To test if something is an oxidizing agent, react it with a popular reducing agent, one being potassium iodide, which shows a clear colourless to brown colour.

To test if something is a reducing agent, react it with a popular oxidizing agent, one being acidified potassium manganate(VII) which goes from purple to colourless.

Examples of Redox Reactions in real life

Sodium Chlorate: more commonly known as bleach, it oxidizes dyes back to their colourless state while being reduced itself.

Rusting: When iron and its alloys come into contact with moisture in the air, an oxide layer (from the O in H2O) forms on the metal as it forms hydrated iron molecules

Browning of food: When food is exposed to the surface, enzymes called phenolases also come into contact with the air, and this causes them to oxidize chemicals in the food into melanins, which are brown and give the food the brown colour.

Preservation of food: sodium sulphite and sulphur dioxide are used as food preservatives to prevent food from browning as they prevent the oxidation of chemicals in the cells into melanin.

Acids and Bases

Acids

Acids are known as proton donors. This is because they give up their H+ ions into solutions. H+ ions are called protons because they just have a singular proton.

Acids can also be defined as substances that react with bases to produce salt and water.

An antacid is a substance taken by people to neutralize stomach acid. It is a base.

All acids have hydrogen ions, but not all hydrogen compounds are acids.

Properties of acids

• Sour

• Corrosive

• pH less than 7

• Turn blue litmus paper red

• Acids only exhibit these properties when dissolved in water! it is the freeing of the H+ ion in the acid that causes these reactions. IT TAKES WATER TO MAKE AN ACID!

• KNOW THESE FOR EXAM!

• pH is the potential of hydrogen dissociation

Acids and their reactions

1. Acids react with metals to produce salt and hydrogen

   • Zn + HCl → ZnCl + H2

2. Acids react with bases to produce salt and water.

   • MgO + H2SO4 → MgSO4 + H2O

3. Acids react with (metal) hydrogen carbonates to produce carbon dioxide, salt, and water

   • MgHCO3 + HCl → CO2 + H2O + MgCl

• these are also the ways to produce salts!

• KNOW THESE FOR THE EXAM!

It is worth noting that ammonium (NH4+) salts react with acids to produce ammonia, a salt, and water.

It is also worth noting that acids give off heat when dissolved in water as new bonds are being formed

Zn (s) + CuSO4 (aq) → ZnSO4 (aq) + Cu (s)

• in this reaction, the overall ionic equation is Zn + Cu2+ → Zn2+ + Cu

• The zinc causes the copper ion to be reduced to its atomic form

• The copper ion causes the zinc to be oxidized to its ionic form

• Visually, copper sulphate is blue, and when it reacts with zinc it forms ZnSO4, a clear solution, with bits of copper metal as a precipitate as copper is displaced

Basicity of Acids

The basicity of an acid is essentially how many hydrogen ions per molecule of the acid are given off in the water.

Monobasic acids give off one hydrogen ion per molecule.

• HCl, CH3COOH

Dibasic acids give off two hydrogen ions per molecule.

• H2SO4

Tribasic acids give off three hydrogen ions per molecule.

• H3PO4

Examples of Acids in Real life

Ascorbic acid (vitamin c)

Citric acid (lime, lemon) (organic, and so its weak

Lactic acid (muscles)

Methanoic/formic acid (ants/bee sting)

Ethanoic/acetic acid (vinegar)

Acids in all three states

Acids may be:

• Solid: citric acid, ascorbic acid

• Liquid: nitric acid, phosphoric acid

• Gas: HCl

Types of Acids

MINERAL ACIDS (STRONG ACIDS)

Mineral acids were originally obtained from minerals.

• HCl

• HNO3

• H2SO4

• H3PO4

ORGANIC ACIDS (WEAK ACIDS)

Organic acids are obtained from plant and animal materials.

• CH3COOH - Ethanoic acid (vinegar)

• Tartaric acid (grapes)

• Citric Acid

• Lactic Acid

Something to note is that these weak acids almost always contain Carbon while the strong acids don’t contain any carbon but instead have hydrogen.

Bases

Bases are proton acceptors. They usually take the place of the hydrogen ions in an acid. They react with an acid to produce a salt and water.

A base is typically anything that reacts with an acid, it doesn’t have to be a metal oxide or hydroxide.

• sodium carbonate is a popular base

Properties of Bases

• Bitter

• Soapy

• Slippery (due to conversion of skin oils to soap)

• pH greater than 7

• turn red litmus paper blue

• React with ammonium (an acid) salts to produce ammonia (a base)

• KNOW THESE FOR EXAM!

• Just because something is a base doesn’t mean it can’t damage your skin, strong bases can definitely damage your skin

Alkalis

Alkalis are essentially soluble bases. They are substances that release hydroxide ions when in water.

• e.g. NaOH

Types of Oxides

Acidic Oxides

Acidic oxides are oxides of non-metals that react with bases to produce salt and water.

• carbon dioxide, sulphur dioxide, nitrogen dioxide, silicon dioxide

ACIDIC OXIDES VS ANHYDROUS OXIDES

Acidic oxides need to be dissolved in water to exhibit acidic properties.

• e.g. CO2, SO2

Anhydrous oxides exhibit acidic properties without containing any water.

• e.g. H2SO4, HCl

Basic Oxides

Basic oxides are oxides of metals that react with acids to produce salt and water.

• lithium oxide, calcium oxide, lead oxide, magnesium oxide, copper oxide (BLACK)

Amphoteric oxides

Amphoteric oxides are oxides that can act as either basic or acidic oxides, depending on what they are reacting with.

• aluminium oxide, zinc oxide, lead oxide

• aluminium oxide reacting with hydrochloric acid (acting as a base) vs aluminium oxide reacting with NaCl (acting as an acid)

Neutral oxides

Neutral oxides are unreactive and don’t produce anything.

• carbon monoxide, nitrogen monoxide, dinitrogen monoxide.

ACID ANHYDRIDES

Acid anhydrides are non-metallic oxides that produce an acid when dissolved in water.

Indicators

• focus on phenolphthalein!

Strength of Acids and Bases

An acid or base’s strength is defined by how well they ionize in water. Strong acids and bases will always ionize completely in water, and they will always conduct electricity well when dissolved in water.

• Examples of strong acids: H2SO4, HCl, HNO3, H3PO3

• Examples of strong bases: NaOH, LiOH, KOH

• Example of a dissociation: HNO3 + Water → H+(aq) + NO3-(aq)

• We say this reaction is COMPLETE as it is IRREVERSIBLE

• Sulphuric acid is so strong it destroys the litmus paper it reacts with so it seems like there’s no reaction, but instead it’s so corrosive and dehydrating that it dries out the dyes on the paper

Weak acids and bases don’t ionize completely (partially) in water, and the reaction is sometimes reversible (e.g. ammonium) and their acids do not conduct electricity well.

• Example of weak acids: CH3COOH, HCOOH, NH4

• Example of partial dissociation: CH3COOH + Water ⇌ CH3COO- + H+

• We say this reaction is PARTIAL as it is REVERSIBLE

• When these acids react with bases the reaction is a lot slower, i.e., moderate, in comparison to strong acids

Properties of Strong Acids and Bases

• ionize completely

• conduct electricity

• strong electrolytes

• High conc of h+ ions

The greater the oxygen ions, the greater the strength of the acid.

• H2SO4 > H2SO3

The fewer hydrogen ions a base has, the stronger it is.

• O > OH > H2O > H3O

Non-Oxidizing Acids and Oxidizing Acids

Non-oxidizing acids are acids whose anions are weaker oxidising agents than hydrogen, i.e they do not readily accept electrons.

• e.g. most organic acids, cold HCl, H2SO4 and H3PO4

Oxidizing acids are acids whose anions are stronger oxidizing agents than hydrogen, i.e., they readily accept electrons

• e.g. concentrated H2SO4, HNO3,

Uses of Acids & Bases

• Antacids- neutralize stomach acids and prevents indigestion

• Carbon dioxide - Fire extinguisher for metal fires

• Baking powder

Important Precipitation Reactions

Silver + Carbonate/Iodine - Yellow Precipitate

Lead + Iodine - Yellow Precipitate

Lead + Carbonate - White Precipitate

Lead/Silver + Sulphide - Black Precipitate

Lead/Silver + Chloride - White Precipitate

Group 2 Metals + Carbonate - White Precipitate

Salts

Salts are formed when metal or ammonium ions replace the hydrogen ions in an acid.

There are two types of salts: normal and acidic salts.

Acidic sats are salts that are formed when only some of the hydrogen ions in an acid are replaced, and thus can only be formed by dibasic and tribasic acids (acids with more than one hydrogen to replace).

Normal salts are salts that are formed when all of the hydrogen ions in an acid are replaced. Normal salts can be formed by all types of acids (mono, di and tribasic acids)

How to Prepare a Soluble Salt by Titration

1. Obtain a fixed volume of NaOH solution (25cm3)

2. Add an indicator to the NaOH solution.

3. Put the acid that will create the salt (whether it be a chloride, sulphate, etc. salt) in the burette

4. Titrate until you find the point of neutralization

5. Do the experiment again without the indicator

6. Take the result and heat up the solution with a Bunsen burner until it reaches the point of saturation, then let it cool and the salt crystals will form.

7. Filter the solution to get the salt crystals as residue.

8. Wash the salt crystals in cool distilled water and dry them between filter paper.

Soluble acid + soluble base = soluble salt (Titration)

• Ca(NO3)2 + HCl = CaCl + HNO3 (pretty strong acid here)

• Ba(OH) + HCl = BaCl + H2O

Metal + Metal = Salt

• Fe + S = FeS

• this is direct combination, works for all different combinations of metals

Insoluble base + acid = soluble salt

• Zn + HCl = ZnCl + H

• PbO + HNO3 = Pb(NO3)2 + H2O

Soluble acid (or base) + Soluble base = insoluble salt (Precipitation)

• Pb(NO3)2 + KI = KNO3 + PbI (s)

• AgNO3 + CaCO3 = Ca(NO3)2 + Ag2CO3

• When preparing, mix solutions, filter them, then dry them

Acidity and Alkalinity of Salts

• Strong acid + Strong base = neutral salt (NaCl)

• Strong acid + weak base = acidic salt (NH4Cl)

• Weak acid + strong base = basic salt (CH3COONa)

You can tell if a solution will make water more acidic, basic or neutral based on the fact that they will either receive (base) or donate (acid) hydrogen ions into the solution.

Water of Crystallisation

Water molecules may become part of the crystal structure of salts when forming a salt from crystallization from an aqueous solution.

• this is called water of crystallisation

• also known as hydrated salts

• Evaporating these salts leaves behind a powdery (non-crystalline) form of the salt

• Salts without water of crystallization are called anhydrous salts.

Dangers of Salts

NaNO3/NaNO2/Sodium Benzoate - carcinogenic- may cause cancer, may increase asthma and hyperactivity in children

NaCl - hypertension

Usage of Salts

CaCO3 - cement

NaCl/NaNO2 - to preserve food

Ammonia Salts - smelling

Epsom salts (magnesium sulphate) - for relaxation purposes, improves plant growth

NaClO - in bleach, it removes dye stains

CaSO4 - plaster of paris, for making casts

Moles

A mole is a fundamental quantity that represents 6.02 x 10^23 of a thing, usually used to represent particles of a substance.

A mole can also be defined as how many particles are in 12g of carbon-12.

The relative atomic mass (which is the same as the molar mass) of an element is the average mass of all the existing isotopes of the element compared to 1/12th of a carbon-12 atom.

• e.g. chlorine has a relative atomic mass of 35.5 because the mass of its two isotopes Cl-35 (34.96g) and Cl-37 (36.96g) give it an average mass of 35.5

• ONE DIFFERENCE BETWEEN MOLAR MASS AND AMU IS THAT AMU HAS NO UNITS! THE MOLAR MASS UNIT IS G/MOL

The relative atomic mass is not the same as the mass number of an element, as it is the AVERAGE of multiple isotopes’ mass numbers.

• It’s only the same when the element has no isotopes.

The molar mass of an element is the grams per 1 mole of an element.

• 1 mole of atoms = molar mass of the element

Metal or noble gases are made of atoms.

Covalent/molecular compounds are made of molecules made up of atoms

Ionic compounds are made of formula units made up of ions

PLEASE NOTE:

If the molecule is diatomic, then it contains 1.2 x 10^24 particles (6.02 x 10^23 times 2) since it has two atoms.

• However, if we count by MOLECULES (covalent bond between atoms) then it has 6.02 x 10^23 MOLECULES

Please note the difference between atoms and molecules when it comes to diatomic compounds.

When calculating the molar mass of a compound, please check the subscript to see if you need to multiply the molar mass by anything!

• e.g. Cl2 would be 35.5 x 2 = 71

Mass concentration is the mass of solute per 1dm3 of solution.

• Mass concentration = mass/volume

• g/dm3

Molar concentration is the moles of solute per 1dm3 of solution.

• Molar concentration = moles/volume

• mol/dm3

A standard solution is a solution whose volume is accurately known and is made by dissolving a known volume of solute in a solvent to produce a specific volume of a solution.

• you either know the molar concentration or the mass concentration

A standard solution is made in a volumetric flask, which is glassware with an elongated neck and calibrated to hold a set volume.

• Essentially the amount of solute you want per unit volume is weighed out using a scale

• Water is added but not to the mark yet

• The solute is added to the water

• The water is then filled to the mark on the volumetric flask

• The volumetric flask is shaken upside down

Molar volume is the volume of one mole of gas at either standard pressure & temperature or room temperature & pressure.

Avogadro’s Law states that the volume of different gases at constant temperature and pressure have the same number of particles or moles.

• this means they have the same number of moles too

Two types of volume

RTP

• 24dm3

• 25*C

• 101.3kPA

STP

• 22.4dm3

• 0*C

• 101.3kPA

LIMITING REAGENTS

The limiting reagent is the reagent that gives the lower number of moles of product.

This can be calculated using the molar ratio.

• e.g. CaCO3 + 2HCl → CaCl2 + CO2 + H2O

• Comparing CaCO3 to CO2, and then 2HCl to CO2, HCl is the limiting reagent, because it produces less moles due to it’s molar ratio.

• Let’s say CaCO3 has 0.20 moles. Then CO2 would also have 0.20 moles

• Let’s say HCl has 0.274 moles. Then CO2 would have 0.137 moles and therefore be less than the CaCO3 so it would be the limiting reagent.

Empirical and Molecular Formula

The empirical formula of a compound is the simplest form of the compound that simply shows the ratio of each element in the compound.

The molecular formula is the actual compound.

When given the question, the main thing you’re going to do is find the moles of everything there.

But to do that, you need to find out what elements will be in the ratio.

e.g. 1.50 g sample of hydrocarbon undergoes complete combustion to produce 4.40 g of CO2 and 2.70 g of H2O. What is the empirical formula of this compound?

We know that because the question said “hydrocarbon” then the elements involved are carbon, hydrogen and oxygen.

To find out the mass of each, compare the mass of what the compound usually would be, i.e. its molar mass, to the mass in the question.

For example, carbon dioxide (CO2) usually has 44 grams per mole. This means that for every mole of carbon dioxide, there’s 12g of carbon. In the question, however, there is only 4.40g per mole of carbon dioxide.

So essentially:

44 grams of CO2 = 12g of carbon dioxide

4.40 grams of CO2 = x of carbon dioxide

Cross-multiply and find your answer.

Do the same for hydrogen, and oxygen since you know what carbon and hydrogen are just subtract them from the total.

To find the empirical formula you need the mass, molar mass, moles, and finally mole ratio.

The mole ratio is just the moles of element/moles of the smallest element.

You may also be asked to find the molecular formula.

To do this, you need the relative molecular formula and divide it by the mass of the empirical formula. When you get this value, multiply it by the empirical formula.

MF = MFM/EFM

Mole factor = molar mass of molecular formula (which is just the molar mass of the compound)/mass of empirical formula

then

MF x Emp F = Mol F

IMPORTANT: IF YOU ARE NOT GIVEN A MASS OF THE COMPOUND, ASSUME YOU HAVE BEEN GIVEN 100G!!

[e.g On analysis, a sample of a sugar was found to contain 40.0% carbon, 6.7% hydrogen and 53.3% oxygen. The molar mass of the compound is 108 gmol-1 . Determine the molecular mass of the compound. (Assume that mass of the sample given is 100g)]

In this question, we have the molar mass, which is the grams per mole of a substance. But we actually work with the mass of the compound in this case. Grams and grams per mole are different. So we assume the mass to be 100g, then 40% would just be 40g, 6.7% would be 6.7 grams and so on.

When they give us the molar mass of the compound that’s the MOLECULAR formula MASS or weight. (g/mol)

Titration

Titration is a technique in which a solution with a known concentration is used to find the concentration of another solution.

The solution of known concentration is called the titrant and it’s usually in the burette.

The solution of unknown concentration is called the analyte and it’s usually in the conical flask.

• also called titrand

Titration is usually a neutralization reaction, where an acid and a base react to form salt and water.

An indicator is used in titration experiments to signify the end of the reaction. This is because indicators usually have different colours in different types (basic or acidic) of solutions.

Some popular indicators:

• phenolphthalein

• screened methyl orange

• methyl orange

Aliquot - A smaller sample of a solution taken from a stock solution

Titre - the concentration of a solution determined by titration. the volume of solution x it takes to neutralize solution y.

Equivalence point - the point at which a neutralization reaction truly ends, i.e. just enough base was added to just enough acid. the H+ ion concentration matches the OH+ concentration

Endpoint - the point at which an acid-base indicator changes colour

We titrate when we want to:

• determine the concentration of a substance

• create a sample of soluble salt

• standardize a solution

pipette: a standard tube for transferring small amounts of water

burette: a graduated cylinder that stores a set amount of water, and has a stop-cock that slowly lets water out. used for delivering KNOWN volumes of liquid

conical flask: flat bottom, used to manipulate solutions

retort stand: supports other glassware with a clamp

When we do titration experiments, we usually select the results within 0.1cm3 of each other, and we don’t use rough titrations!

Steps to carry out titrations

1. add a known volume of a substance (let’s say acid) with an unknown concentration to a conical flask using a pipette

2. add a few drops of indicator

3. Put the known concentration of base into the burette

4. record starting volume on the burette (make it zero)

5. slowly add the solution in the burette to the one in the conical flask while swirling

6. stop when there is a change in the colour of the indicator. the endpoint is reached

7. record the final volume

8. FInal volume-inital volume = volume used in the experiment

9. Do more than one trial and average the volume used in the experiment (titre)

Things in Particular to know for the exam

1. Know the particular reactions an acid has with metals, bases, metal carbonates

2. Oxygen relights a flame

3. Hydrogen makes a pop sound

4. Test for carbon dioxide, it’s acidic, so it turns blue litmus red (or limewater)

5. Acid anhydrides are compounds that when you react them with water they form an acid (important!)

6. Bases include metal oxides, hydroxides and AMMONIA GAS (BECAUSE IT ACCEPTS A HYDROGEN ION, and it can donate two electrons)

7. KNOW ALLLLL THE OXIDES SUPER IMPORTANT

8. Amphoteric oxides examples: Aluminium oxide, lead oxide

9. Molar masses are given!!

10. For the titration questions, know how to find average volume, and unknown solutions, similar to the question in the lab. you find the average volume of the TITRANT!

11. Oxidizing and reducing agents will come throughout the paper

12. Empirical and Molecular formula definitely coming