Chemical Reactions and Equations-Notes
Chemistry - Notes
Chemical Rection
A chemical reaction can be confirmed by observing one or more of the following changes:
- Change in state
- Change in colour
- Evolution of a gas
- Change in temperature
CHEMICAL EQUATIONS
A chemical equation is a simple way to show what happens in a chemical reaction using symbols and formulas. It tells us which substances react together (called reactants) and what new substances are formed (called products). In a chemical equation:- Reactants: are written on the left side
- Products: are written on the right side
- An arrow \(\rightarrow\): connects them, showing the direction of change
- Numbers: show how many atoms or molecules are involved
Chemical equations help us understand reactions quickly and clearly, showing exactly what changes during the process without using long descriptions.
Writing a chemical equation
When a magnesium ribbon is burnt in oxygen, it gets converted to magnesium oxide. This description of a chemical reaction in a sentence form is quite long. It can be written in a shorter form. The simplest way to do this is to write it in the form of a word-equation.The word-equation for the above reaction would be
\[\underset{\quad Reactants}{\text{Magnesium} + \text{Oxygen}} \rightarrow \underset{Product}{\text{Magnesium oxide}}\] A chemical equation represents a chemical reaction. Formulae of magnesium, oxygen and magnesium oxide, the above word-equation can be written as \[Mg+O_2 \rightarrow MgO\]
Balanced Equation
A balanced chemical equation has the same number of atoms on the reactant and product sides for each element involved in the reaction.This is a condition that the equation must meet in order to be consistent with the law of conservation of mass and constant proportions.
The word equation for the reaction of zinc with sulphuric acid may be represented as
\(\text{Zinc } + \text{ Sulphuric acid } \rightarrow \text{ Zinc sulphate } + \text{ Hydrogen}\)
The above word-equation may be represented by the following chemical equation
\[Zn + H_2SO_4 → ZnSO_4 + H_2\]
Steps to Balance a Chemical Equation
- Start by writing the correct formulae for reactants and products.\[Fe+H_2O\rightarrow Fe_3O_4+H_2\]
- List Number of Atoms for Each Element
Create a table or list to count the number of each atom present on both sides.Element Number of atoms in reactants (LHS) Number of atoms in Products (RHS) Fe 1 3 H 2 2 O 1 4 - Begin with the Most Complex Molecule
Balance atoms in the compound containing the greatest number of different elements (often metals or polyatomics)
Using these criteria, we selevt \(Fe_3O_4\) and the elemet oxugen in itAtoms of Iron In reactants In Products Initial 1 in (\(Fe\)) 3 in (\(Fe_3O_4\)) To Balance \(1 \times 3\) \(3\) - Balance Atoms One by One
Balance the other elements (except hydrogen and oxygen) next. - Balance Hydrogen and Oxygen Atoms at the End
After all other elements have been balanced, balance hydrogen and oxygen atoms.Atoms of Oxygen In reactants In Products Initial 1 in (\(O\)) 4 in (\(Fe_3O_4\)) To Balance \(1 \times 4\) \(4\)
Equation now becomes
\[Fe + 4H_2O\rightarrow Fe_3O4 + H_2\]
- Make Sure that All the Atoms Are Balanced on Both Sides
Recount the number of atoms of each element on both sides to verify the equation is balanced. - Write the Final Equation with Smallest Whole Number Coefficients
The fully balanced equation should use the smallest whole number coefficients
Balance Equation\[3Fe + 4H_2O\rightarrow Fe_3O4 + 4H_2O\] -
Writing Symbols of Physical States
To make a chemical equation more informative, the physical states of the reactants and products are mentioned along with their chemical formulae.
The gaseous, liquid, aqueous and solid states of reactants and products are represented by the notations (g), (l), (aq) and (s), respectively. The word aqueous (aq) is written if the reactant or product is present as a solution in water. \[3Fe(s) + 4H_2O(g)\rightarrow Fe_3O4(s) + 4H_2O(g)\]
| Atoms of Hydrogen | In reactants | In Products |
|---|---|---|
| Initial | 8 in (\(4H_2O\)) | 2 in (\(H2\)) |
| To Balance | \(8\) | \(2\times 4\) |
Example:
Balace the given equations
- \(CO(g) + H_2(g)\rightarrow CH_3OH(l)\)
Balanced Equation: \[CO(g) + 2H_2(g)\rightarrow CH_3OH(l)\]Elemet In Reactant (LHS) In Product (RHS) Carbon 1 (in \(CO\)) 1 (in \(CH_3OH\)) Oxygen 1 (in \(CO\)) 1 (in \(CH_3OH\)) Hydrogen 2 in \(H_2\)
to balance: \(2\times 2\)4 (in \(CH_3OH\)) - \(HNO_3 +Ca(OH)_2 → Ca(NO_3)_2 + H_2O\)
Balanced Equation is:\[2HNO_3 + Ca(OH)_2 \rightarrow Ca(NO_3)_2 + 2H_2O\]Element In Reactant (LHS) In Product (RHS) Calcium 1 (in \(Ca(OH)_2\)) 1 (in \(Ca(NO_3)_2\)) Nitrogen 1 (in \(HNO_3\))
To balance: \(1 \times 2\)2 (in \(Ca(NO_3)_2\)) Partially Balanced Equation:
\[2HNO_3 + Ca(OH)_2 \rightarrow Ca(NO_3)_2 + H_2O\]Oxygen 8 (6 in \(2HNO_3\) and 2 in \(Ca(OH)_2\)) 7 (6 in \(Ca(NO_3)_2\) and 1 in \(H_2O\))
To balance O: \(2 \times H_2O\)Balanced Equation:
\[2HNO_3 + Ca(OH)_2 \rightarrow Ca(NO_3)_2 + 2H_2O\]Hydrogen 4 (2 in \(2HNO_3\)) 4 (4 in \(2H_2O\)) - \(NaCl + AgNO_3 → AgCl + NaNO_3\)
Equation is already in balanced state: \[NaCl + AgNO_3 → AgCl + NaNO_3\]Element In Reactant (LHS) In Product (RHS) Sodium 1 (in \(NaCl\)) 1 (in \(NaNO_3\)) Silver 1 (in \(AgNO_3\)) 1 (in \(AgCl\)) Nitrogen 1 (in \(AgNO_3\)) 1 (1 in \(NaNO_3\)) Oxygen 3 (in \(AgNO_3\)) 3 (in \(NaNO_3\)) - \[BaCl_2 +H_2SO_4 \rightarrow BaSO_4 +HCl\]
Balanced Equation:\[BaCl_2 +H_2SO_4 \rightarrow BaSO_4 +2HCl\]Element In Reactant (LHS) In Product (RHS) Barium 1 (in \(BaCl_2\)) 1 in (\(in BaSO_4\)) Sulpher 1 (\(in H_2SO_4\)) 1 (in \(BaSO_4\)) Chlorine 2 (in \(BaCl_2\)) 1 (in \(HCl\))
to balance:
\(1\times 2 HCl\)Partially Balanced Equation:\[BaCl_2 +H_2SO_4 \rightarrow BaSO_4 +2HCl\] Oxygen 4 (in \(H_2SO_4\)) 4 (in \(BaSO_4\)) Hydrogen 2 (in \(H_2SO_4\)) 2 (in \(2HCl\))
TYPES OF CHEMICAL REACTIONS
There are five main types of chemical reactions, each describing how reactants change into products:- Combination reaction: Two or more reactants combine to form one product.
- Decomposition reaction: A single reactant breaks down into more than one product.
- Displacement reaction: An atom or set of atoms is displaced by another atom in a molecule.
- Double displacement reaction: Two compounds exchange ions or elements.
- Oxidation and Reduction reaction: A reaction that involves a transfer of electrons between two species.
Combination Reaction
Calcium oxide reacts vigorously with water to produce slaked lime (calcium hydroxide), releasing a large
amount
of heat.
\(\underset{\text{Quick Lime}}{CaO(s)} + H_2O (l) \rightarrow \underset{\text{Slaked
Lime}}{Ca(OH)_2(aq)} +
Heat\)
- Burning of coal \[C(s) + O_2(g) \rightarrow CO_2(g)\]
- Formation of water from \(H_2(g)\) and \(O_2(g)\)
\[2H_2(g) + O_2(g) \rightarrow 2H_2O(l)\]
Exothermic Reaction
An exothermic reaction is a chemical reaction that releases energy in the form of heat or light. In
contrast, an endothermic reaction absorbs energy from its surroundings.
Examples of exothermic reactions
- Combustion: A substance reacts rapidly with oxygen to release heat and energy. For example, the combustion of wood releases heat and light.
- Neutralization: Neutralization An acid and a base react to
produce
salt and water.
For example,
when sodium hydroxide \((NaOH)\) reacts with hydrochloric acid \((HCl)\), it produces salt and water. \[NaOH + HCl \rightarrow NaCl + H_2O + Heat\] - Nuclear fission: A heavy nucleus decomposes into smaller nuclei when bombarded with neutrons.
- Respiration Energy: is released when carbon dioxide in food
is
broken down into glucose, which then reacts with oxygen in cells.
\(\underset{Glucose}{C_6H_{12}O_6(aq)} + 6O_2(aq) \rightarrow 6CO_2(aq) + 6H_2O(l) + energy\)
- Calcium oxide and water: When calcium oxide combines with water to form calcium hydroxide, heat is evolved. \[CaO + H_2O \rightarrow Ca(OH)_2 + Heat\]
Endothermic Reaction
An endothermic reaction is a chemical or physical change that absorbs heat from its surroundings. The
word
"endothermic" comes from the Greek words endon (inside) and therm (heat).
Examples of endothermic reactions:
- Melting ice cubes
- Evaporating liquid water
- Photosynthesis, and
- Dissolving ammonium nitrate in water.
Decomposition Reaction
When single reactant breaks down to give simpler products. This is a decomposition reaction.
\[FeSO_4.7H_2O\overset{\Delta}{\longrightarrow}Fe_2O_3+SO_2+SO_3\] Ferrous sulphate crystals \((FeSO_2.7H_2O)\) lose water when heated and the colour of the crystals changes. It then decomposes to ferric oxide \((Fe_2O_3)\), sulphur dioxide \((SO_2)\) and sulphur trioxide \((SO_3)\). Ferric oxide is a solid, while \(SO_2\) and \(SO_3\) are gases.
Decomposition of calcium carbonate to calcium oxide and carbon dioxide on heating is an important decomposition reaction used in various industries. \[\mathrm{CaCO_3}\xrightarrow{\Delta}\mathrm{CaO + CO_2}\] \[ \underset{\text{Lead Nitrate}}{2\mathrm{Pb(NO_3)_2}} \xrightarrow{\Delta} \underset{\text{Lead Oxide}}{2\mathrm{PbO}} + 4\mathrm{NO}_2 + \mathrm{O}_2 \] \[2\mathrm{AgCl}\xrightarrow{\text{Sun Light}}2\mathrm{Ag} + \mathrm{Cl_2}\] \[2\mathrm{AgBr}\xrightarrow{\text{Sun Light}}2\mathrm{Ag} + \mathrm{Br}_2\]Displacement Reaction
A displacement reaction is a chemical reaction where a more reactive element replaces a less reactive
element in
a compound. This reaction can occur in solution or in molten form.
Reactivity Series:
rhyme to learn:
Please Stop Calling Me A Cute Zebra. I Like Her Calling Me Smart Goat Please.
\[
\color{red}\bigg\Downarrow\text{Decreasing Order of Reactivity}~\bigg\Downarrow
\]
\[
\begin{array}{cc}
\left\{
\begin{array}{l}
\textcolor{red}{\mathbf{P}} - \text{Potassium} \\
\textcolor{red}{\mathbf{S}} - \text{Sodium} \\
\textcolor{red}{\mathbf{C}} - \text{Calcium} \\
\textcolor{red}{\mathbf{M}} - \text{Magnesium} \\
\textcolor{red}{\mathbf{A}} - \text{Aluminium} \\
\textcolor{red}{\mathbf{C}} - \text{Carbon (Non-Metal)} \\
\textcolor{red}{\mathbf{Z}} - \text{Zinc} \\
\textcolor{red}{\mathbf{I}} - \text{Iron} \\
\textcolor{red}{\mathbf{L}} - \text{Lead} \\
\textcolor{red}{\mathbf{H}} - \text{Hydrogen (Non-Metal)} \\
\textcolor{red}{\mathbf{C}} - \text{Copper} \\
\textcolor{red}{\mathbf{M}} - \text{Mercury} \\
\textcolor{red}{\mathbf{S}} - \text{Silver} \\
\textcolor{red}{\mathbf{G}} - \text{Gold} \\
\textcolor{red}{\mathbf{P}} - \text{Platinum}
\end{array}
\right.
\end{array}
\]
Examples of Displacement Reaction
\[\begin{array}{r|l}\mathrm{Fe}\\\uparrow
&\mathrm{Fe} +\mathrm{CuSO_4}\rightarrow\mathrm{FeSO_4 + Cu}
\\\mathrm{Cu}
\end{array}\]Fe is more reactive than Cu, therefore it will
displace Cu
\[\begin{array}{r|l}\mathrm{Zn}\\\uparrow
&\mathrm{Zn + CuSO_4}\rightarrow\mathrm{ZnSO_4+Cu}
\\\mathrm{Cu}
\end{array}\]
Zn is more reactive than Cu, therefore it will
displace Cu
\[\begin{array}{r|l}\mathrm{Pb}\\\uparrow
&\mathrm{Pb+CuCl_2}\rightarrow\mathrm{PbCl_2 + Cu}
\\\mathrm{Cu}
\end{array}\]
Pb is more reactive than Cu, therefore it will
displace Cu
Double Displacement Reaction
A double displacement reaction is a chemical reaction that occurs when two ionic compounds exchange ions
to form
two new compounds. It's also known as a double replacement or metathesis reaction.
How it happens
- The reactants dissolve in water and break apart into ions
- The positive ions (cations) of one compound switch places with the negative ions (anions) of the other compound
- The new compounds form
\[\mathrm{Na^+SO_4^{2-}+Ba^{2+}Cl^-_2} \rightarrow \mathrm{\underset{Precipitate}{{Ba}^{2+}{SO_4}^{2-} \downarrow }+Na^+_2Cl^-}\] white precipitate of \(\mathrm{BaSO_4}\) is formed by the reaction of \[ \mathrm{SO_4}^{2-}\text{ and } \mathrm{Ba}^{2+} \]The other product formed is sodium chloride which remains in the solution. Such reactions in which there is an exchange of ions between the reactants are called double displacement reactions.
Precipitation reaction: Any reaction that produces a precipitate can be called a precipitation reaction.
Redox Reaction
A redox reaction is a chemical reaction where electrons are transferred between two substances, or species. It is also known as an oxidation-reduction reaction. What happens in a redox reaction?
- The substance that loses electrons is oxidized.
- Loss of Electron … Oxidation
- Loss of Hydrogen … Oxidation
- Gain of Oxygen … Oxidation
- The substance that gains electrons is reduced.
- Gain of Electron … Reduction
- Gain of Hydrogen … Reduction
- Loss of Oxygen … Reduction
- • The oxidation number of an atom, molecule, or ion changes.
- Copper oxide and hydrogen react to form copper and water. In this reaction, hydrogen is oxidized to water while copper oxide is reduced to copper. \[ \mathrm{Cu^{2+}O_2^{2-} + H_2^0\rightarrow Cu^0 + {H_2}^{2+}O^{2-}}\\ \mathrm{Cu^{2+}+2e^-\rightarrow Cu^0}\\\Rightarrow\text{ Gain of Electron - Reduction}\\ \mathrm{H^0-1e^{-}\rightarrow H^+}\\\Rightarrow \text{Loss of Electron - Oxidation} \] During this reaction, the copper(II) oxide is losing oxygen and is being reduced. The hydrogen is gaining oxygen and is being oxidised.
- \[ \mathrm{Zn^{2+}O_2^{2-} + C^0\rightarrow Zn^0 + {C}^{2+}O^{2-}}\\ \mathrm{Zn^{2+}+2e^-\rightarrow Zn^0}\\\Rightarrow\text{ Gain of Electron - Reduction}\\ \mathrm{C^0-2e^{-}\rightarrow C^{2+}}\\\Rightarrow \text{Loss of Electron - Oxidation} \] During this reaction, the \(\mathrm{ZnO}\) is losing oxygen and is being reduced. The Carbon is gaining oxygen and is being oxidised.
- \[ \mathrm{MnO_2 + 4HCl\rightarrow MnCl_2+ H_2O+ Cl_2}\\ \] Loss of Hydrogen by HCl to form \(\mathrm{Cl_2}\Rightarrow \text{ Oxidation}\\\) Loss of Oxygen by \(\mathrm{MnO_2 }\text{ to form }\mathrm{MnCl_2}\Rightarrow \text{ Reduction}\)
- Magnesium metal and oxygen react to form magnesium oxide. In this reaction, magnesium is oxidized
Oxidising Agent
An oxidizing agent is a substance that causes other substances to oxidize by gaining electrons from them. It is also known as an oxidizer or oxidant.
How does an oxidizing agent work?
- In a redox chemical reaction, an oxidizing agent gains electrons from a reducing agent.
- The oxidizing agent's oxidation state decreases, while the reducing agent's oxidation state increases.
- The oxidizing agent is reduced, while the reducing agent is oxidized.
- Oxygen \(\mathrm{(O_2)}\)
- Hydrogen peroxide \(\mathrm{(H_2O_2)}\)
- Halogens, such as chlorine, fluorine, iodine, and bromine
- Halogens have higher electronegativities than other groups of elements.
- This means they have a greater ability to acquire electrons.
- This means they can easily draw electrons to their nuclei.
- Fluorine is the most potent elemental oxidizing agent because of its strong electronegativity.
Reducing Agent
A reducing agent is a substance that donates electrons to another substance in a chemical reaction. This process is called reduction, and the reducing agent is also known as an electron donor.
How does a reducing agent work?- In a redox reaction, one reactant transfers electrons to the other.
- The reducing agent loses electrons and becomes oxidized.
- The other reactant gains electrons and is reduced.
- The process releases heat.
Examples of reducing agents: earth metals, formic acid, and sulfite compounds.
Corrosion
Corrosion is the process by which materials, like metals and non-metals, deteriorate over time. It can occur due to chemical and/or electrochemical reactions with the environment.
- Corrosion is a natural process that causes metals to break down into less desirable substances.
- It can also be defined as the process by which something deteriorates due to oxidation, a chemical reaction with the air.
- Corrosion can occur on materials such as ceramics or polymers.
How does corrosion happen?
- Corrosion occurs when metals react with air moisture, acids, and other gases in the atmosphere.
- This reaction causes damage or disintegration of metal as it interfaces with the environment.
Prevention of Corrosion: Corrosion is the deterioration of metals due to chemical reactions—often with oxygen and moisture in the environment. There are several proven methods to prevent or minimize corrosion, used in homes, industry, and infrastructure projects.
- Painting: Applying paint creates a barrier on metal surfaces, keeping air and moisture away, thus preventing rust on gates, railings, and bridges.
- Oiling and Greasing: Covering iron and steel tools or machinery parts with oil or grease is recommended; this protective layer stops moisture and oxygen from reaching the metal’s surface.
- Galvanisation: Iron or steel objects are coated with a thin layer of zinc. Zinc is more reactive, so it corrodes first, saving the underlying iron from rusting—even if the surface is scratched.
- Alloying: Making alloys by mixing metals (e.g., iron, nickel, and chromium to make stainless steel) creates materials that do not rust or corrode easily. Stainless steel utensils, pipes, and fixtures are common examples.
- Electroplating and Chrome Plating: Iron and other metals can be coated with non-corrosive metals such as tin, chromium, or other protective materials through processes like electroplating, further preventing corrosion.
- Everyday Examples
- Rust-free stainless steel is used in kitchen utensils and water tanks.
- Galvanized iron pipes are common in plumbing.
- Painted gates and fences last longer and stay rust-free.
- Oiling bicycle chains or machine parts prevents rust.
Rancidity
Rancidity is the process by which food containing fats and oils gets spoiled. This spoilage happens due to a chemical reaction called oxidation—when fats and oils react with oxygen from the air. As a result, the food develops an unpleasant smell and taste, making it unfit for consumption.
How Does Rancidity Occur?
- When foods containing fats and oils are exposed to air for a long time, their molecules react with oxygen (this is called oxidation).
- Oxidation leads to the formation of new compounds. Many of these have a foul smell and bad taste.
- Heat, light, and moisture can speed up rancidity.
- If oil or ghee is kept open for several days, it starts smelling bad and tastes unpleasant. This is because of rancidity.
Examples of Rancidity
- Potato chips, namkeen, or snacks taste and smell bad if not stored in airtight packets for several days.
- Butter or ghee left open develops a foul smell.
- Cooking oil stored for long may taste and smell bad when used.
Prevention of Rancidity
There are several simple methods to prevent or slow down rancidity:
- Storing food in airtight containers: Limits the oxygen in contact with food, which slows the oxidation.
- Refrigeration: Keeping food in the fridge slows down the chemical reactions that cause rancidity.
- Adding antioxidants: Chemicals like BHA (Butylated Hydroxyanisole) and BHT (Butylated Hydroxytoluene) are sometimes added to foods to prevent oxidation.
- Packaging with nitrogen gas: Chips and snacks are often packed in nitrogen gas instead of air to keep them fresh and prevent rancidity.
- Keeping food away from light and heat: Heat and sunlight can speed up oxidation, so storing food in a cool, dark place helps.
