O/L Chemistry Formulae and Reactions
Introduction
In Chemistry, often for most people the most difficult part is the Mathematics of Chemistry. And O/L students who are studying for Edexcel or Cambridge Chemistry exams wonder what the most common formulae and reactions are to occur in the exam that require this integrated mathematical understanding instead of plain memorisation. So, what could be the formulae and reactions you need to know?
To summarise
,
each formula and reaction provided in this article has a broad reach not only within Chemistry but spans to Physics and Biology. They are some of the most basics of chemistry formulae and reactions that are required in Edexcel and Cambridge O/L Chemistry and are also carried to A/L and applied more but are less revisited at a conceptual level. Understanding it now could make it much easier later in A/L and any science related field of study.
Formulae
1. Equation for Density
Density (g/cm3) = m (g )/ V (cm3); Where m is mass and V is volume
Mass refers to the amount of a substance, and volume represents how much space that substance can take up.
Gases expand to fill the entire space.
Liquids expand to only take the shape of the space below it.
And solids do not expand nor do they fill in space and take up finite volume.
Density refers to the amount of substance (mass) in a specific space (volume); this can also relate to concentration of elements (any solids, liquids and gases).
Ex: There are two blocks, one block weighs 10 g and a volume of 10 cm3
Meaning the density of this is 1 g/cm3
Another also weights 10 g but has a volume of 5 cm3
Meaning the density of this is 2 g/cm3
Twice as dense as the first block but is smaller (compact) and weights the same.
2. Number of Moles
Number of Moles = Mass (g) / Mass of 1 mole (g)
n = 1.71 moles (3 s.f)
3. Concentration
Concentration (mol/dm3)= moles (mol) / volume(dm3)
Ex: 100 dm3 of a solution has 5 grams of NaCl
What is its concentration?
Concentration =
= 0.05 g/dm3
4. Percentage Yield
Then 2.8/3.21 yields 87.2% of copper oxide obtained
5. Atom economy Percentage
% atom economy = (Moles of desired products)/(Moles of all products) × 100
Atom economy = (32 x 100)/(32 + 58.5 +18) = 29.5%
6. Relative atomic mass
This is an average of all available atomic masses of each isotope, weighed by how common each isotope is available in nature. This is the reason why certain elements on the periodic table have decimal places for their atomic masses.
7. Heat energy
8. Enthalpy
= -377.25 J transferred to the surroundings
Reactions
In many fields of study, more often than not symbols are used to identify and reduce the use of full form words when producing answers.
Some of the most common symbols used in Chemistry are in the table below.
| Symbol | Meaning |
|---|---|
| → | Yields or produces |
| (s) | Solid |
| (l) | Liquid |
| (g) | Gas |
| (aq) | Aqueous |
| ⇌ | Reversible reaction |
| Δ (Delta) | Difference between Similar Parameters. Or. Heat supplied to the system |
| Pt | Catalyst used (eg: Platinum) |
Alkene
Cn H2n
Alkenes are reactive, as they have a C=C double bond making a weak spot for other elements to break into and attach (react), making them unsaturated hydrocarbons.
An example of the use of Alkenes is Polythene.
Alkane
Cn H2n+2
Alkanes are fairly unreactive, due to strong C-C and C-H bonds (single bonds). Every available bond is taken up by hydrogen making them saturated, making no weak spot for other chemicals to displace easily.
An example of the use of Alkanes are petrol and diesel.
Alcohol
Cn H2n+1OH
Take an Alkane and replace one H with an OH group, and an Alcohol is formed, changing how alcohols behave with something like water ( -OH can hydrogen bond with water due to the same energy that is required to break them, is used to make them). They burn cleanly and can be made by fermenting sugars with yeast. The naming always ends in ‘-ol’ making a reminder that there is an -OH group.
An example of the use of Alcohols is ethanol, which is found in drinks and sanitizers and are used for biofuels.
Carboxylic Acid
Cn H2n+1COOH
A hydrocarbon chain ending in a -COOH (Carboxyl) group
A Carboxylic acid is what you get when you oxidize an Alcohol, the -OH group gains an extra oxygen to become -COOH, which is acidic as it can break off as a H+ ion. Meaning they react the same was HCl (inorganic acid).
An example of the use of Carboxylic Acid is vinegar, made from ethanoic acid.
Esters
Carboxylic Acid + Alcohol ⇌ Ester + H2O
The -OH end of the carboxylic acid and the -OH group of the alcohol meet in the middle: the alcohol loses an H, and the Carboxylic acid loses the OH, forming water.
Note: It is crucial that the “the alcohol loses an H, and the Carboxylic acid loses the OH” is understood as this is how the process actually takes place and not the other way around.
Even though the other way still gives you an Ester + water, it is wrong.
Leaving the Alcohol and Acid to merge to form an Ester.
No oxidation state changes occur, making it not a redox reaction.
The double head arrow means it’s a reversible reaction, and it also means that the reaction does not complete, it reaches an equilibrium and requires a small amount of a catalyst to speed up the process.
An example of the use of Esters is the pleasant fruity smell you get from perfumes.
Combustion
Hydrocarbon + O2 → CO2 + H2O
Every hydrocarbon is built from Carbon and Hydrogen. When you burn it in plenty of oxygen (complete combustion), every Carbon atom bonds with an Oxygen atom creating CO2 and every Hydrogen atom bonds with an oxygen atom creating H2O. If there is not enough oxygen to bond, it will produce CO (Carbon monoxide) and Carbon, more commonly known as soot found in vehicle exhausts.
Cracking
Long chain Alkane → Shorter alkane + Alkene
Necessary when breaking down a low used molecule into highly useful molecules.
Ex: C10H22 → C8H18 + C2H4 Acids, Bases & salt
Redox
A reaction that involves both reduction and oxidation is called a redox reaction.
Reduction = Gain of electrons
Oxidation = Loss of electrons
This is easily remembered with the Mnemonic OIL RIG
(O)xidation (i)s (L)oss of electrons, (R)eduction (i)s (G)ain
Metal + Acid
1. Metal + Acid Salt + H2
Mg + 2HCl → MgCl2 +H2
Mg has valency of +2, Cl has valency of -1.
Cross valencies: Mg2+ needs 2 Cl– ions MgCl2.
Two H+ ions each gain 1 electron H2 gas
Mg loses 2 electrons: Mg → Mg2+ + 2e- (Oxidised)
Two H+ ions gain electrons: 2H+ + 2e- → H2
Cl doesn’t change making it a spectator ion balancing the Mg2+
Alkali metal + water
2. Alkali metal + water Alkali metal hydroxide + hydrogen
2Na + 2H2O → 2NaOH + H2
Na is neutral it loses 1 electron (oxidised – Loss of electrons)
Na → Na+ + e-
Na has valency +1
Two of the H+ from water each gain one electron (reduced – Gained electrons)
2H+ + 2e- → H2
OH has valency -1
Cross valencies cancel → NaOH
Not a Redox reaction
Acid + Base
3. Acid + Base → Salt + H20
H3PO4 + NaOH → NaH2PO4 + H2O
Na+ from NaOH replaces one H+ in H3PO4 Na has valency 1, so it displaces one H(+1) to form NaH2PO4. The H+ and OH– combine to form H2O
The base’s OH– accepts the H+ (proton) from the acid. No electrons are transferred.
Metal Oxide + Acid
4. Metal Oxide + Acid → Salt + H20
CuO + H2SO4 → CuSO4 + H2O
Cu has a valency of +2
SO4 has valency of -2.
Cross valencies cancel → CuSO4 (1:1 ratio)
Since Cu was already in its +2 state before the reaction, no electron transfer occurs.
H has valency +1
O has valency -2
Making 2H+ + O2- → H2O
Metal Hydroxide + Acid
5. Metal Hydroxide + Acid → Salt + H20
Zn(OH)2 + H2SO4 → ZnSO4 + 2H2O
Zn has valency +2
SO4 has valency -2
Cross valencies Cancel → ZnSO4 (1:1)
Each OH– ion neutralizes one H+
Yielding 2H2O
Carbonate + Acid
6. Carbonate + Acid → Salt + CO2 + H20
Na2CO3 + 2HCl → 2NaCl + H2O + CO2
Na has valency +1
Cl has valency -1
Cross valencies cancel → NaCl (1:1)
But 2Na present → 2NaCl
CO32- reacts with 2H+: CO32- + 2H+ Yielding H2CO3 (Carbonic acid)
Which immediately decomposes to provide H2O + CO2
Metal Carbonate + Acid
7. Metal Carbonate + Acid → Salt + H20 +CO2
Li2CO3 + 2HCl → 2LiCl +H2O +CO2
Li has valency +1
Cl has valency -1
Cross valencies cancel → LiCl (1:1)
But 2Li present → 2LiCl
CO32- reacts combined with 2H+. Yields H2CO3
Which immediately decomposes to provide H2O + CO2
Types of reactions
Redox reactions
Displacement reactions (Halogens)
Cl₂ + 2KBr → 2KCl + Br₂
Cl₂ + 2KI → 2KCl + I₂
More reactive Halogens displace less reactive ones
Displacement reactions for metals
Fe + CuSO₄ → FeSO₄ + Cu
Zn + CuSO₄ → ZnSO₄ + Cu
| Type | General Equation | Example |
|---|---|---|
| Combustion | CxHy + O2 → CO2 + H2O | CH4 + 2O2 → CO2 + 2H2O |
| Synthesis | A + X →AX | 2H2 + O2 → 2H2O |
| Decomposition | AX → A + X | 2H2O → 2H2 + O2 |
| Neutralisation | AX + BY → BX +AY | H2SO4 + CuO → CuSO4 + H2O |
| Single-Displacement | A + BX →AX + B | Zn + CuSO4 → ZnSO4 + Cu |
| Double-Displacement | AX + BY →AY + BX | H2SO4 + CuO → CuSO4 + H2O |
| Reactivity of Metals | Reactivity of non-metals (Halogens) |
|---|---|
| K | F2 |
| Ca | Cl2 |
| Na | Br2 |
| Mg | I2 |
| Al | |
| Zn | |
| Fe | |
| Pb | |
| H | |
| Cu | |
| Ag | |
| Au |
The reactivity decreases as you go down the table. Meaning each element displaces the element below it.
Metal Reactivity:
The higher a metal sits, the more easily its atoms give away electrons to form positive ions, that is the meaning of ‘reactive’. It can rip away those electrons from the ions of the lower metals. Forming, for example, potassium ions and kicking out the free neutral metal.
Reactivity of Halogens:
Halogens react by grabbing an extra electron to become a negative ion. The higher the halogen is, the stronger its pulling force is and the reactivity decreases as you go down the group because outer shells get further from the nucleus and are shielded by inner electrons.
