TYPES OF CHEMICAL REACTIONS at a glance

  Chemical reactions are reactions in which elements or/and compounds combine chemically to form new substances.

There are different types of chemical reactions, they include 

1.   Combinations reactions 

2.   Decomposition reactions 

3.   Displacement reaction

4.   Double decomposition reaction

5.   Thermal Dissociation reaction 

6.    Reversible reaction

7.  Catalytic reaction

🔗 Combination Reaction (Short Note)

A combination reaction is a chemical reaction in which two or more substances combine to form one single product.

It is also called a synthesis reaction.

General Form

A + B → AB

Examples

1. Formation of magnesium oxide
   
   2Mg_(s) + O_2(g) →2MgO_(s)
   
   

2. Formation of water
   
   2H_2(g) + O_2(g) →2H₂O_(l)
   
   

3. Formation of ammonia  
   
   N_2(g) + 3H_2(g) →2NH_3(g)
   
   

4. Formation of calcium oxide
   
   CaO_(s) + CO_2(g)→ CaCO_3(s)

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🧠 Important Tip

In a combination reaction, many reactants give one product.

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🧪 Decomposition Reaction

A decomposition reaction is a chemical reaction in which one compound breaks down into two or more simpler substances when heat, electricity, or light is applied.

General Form

AB → A + B

Types and Examples

1. Thermal Decomposition (by heat)

CaCO_3(s) {heat} CaO_(s) + CO_2(g)

(Calcium carbonate breaks into calcium oxide and carbon dioxide.)

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2. Electrolytic Decomposition (by electricity)

2H₂O_(l) {electricity} 2H_2(g) + O_2(g)

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3. Photochemical Decomposition (by light)

2AgCl_(s) {sunlight} 2Ag_(s) + Cl_2(g)

🧠 Important Tip

If one compound splits into two or more products, it is a decomposition reaction.

🔁 Displacement Reactions 

A displacement reaction is a chemical reaction in which a more reactive element replaces a less reactive element from its compound.

It usually occurs between a metal and a salt solution.

General Equation

A + BC → AC + B
(Where A is more reactive than B)

Examples

1. Zinc and copper (II) sulphate
   Zn_(s) + CuSO_4(aq) → ZnSO_4(aq) + Cu_(s)

2. Zinc displaces copper because zinc is more reactive.

3. Iron and copper (II) sulphate
   
   Fe_(s) + CuSO_4(aq) → FeSO_4(aq) + Cu_(s)
   
   

4. Copper and silver nitrate

5. Cu_(s) + 2AgNO_3(aq) → Cu(NO₃)_2(aq) + 2Ag_(s)
   
   

Important Points

• Only a more reactive metal can displace a less reactive metal.

• The reaction depends on the reactivity series.

🧠 Important Tip

If a metal is higher in the reactivity series, it will displace a metal below it from solution.

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🔄 Double Decomposition Reaction

A double decomposition reaction (also called double displacement or metathesis reaction) is a chemical reaction in which two compounds exchange their ions to form two new compounds.

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General Form

AB + CD → AD + CB

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Examples

1. Reaction between sodium chloride and silver nitrate
   
   NaCl_(aq) + AgNO_3(aq) → AgCl_(s) + NaNO_3(aq)
   
   

2. Reaction between barium chloride and sodium sulphate
   
   BaCl2(aq) + Na₂SO_4(aq) →BaSO_4(s) + 2NaCl_(aq)

3. Reaction between hydrochloric acid and sodium hydroxide
   
   HCl_(aq) + NaOH_(aq) → NaCl_(aq) + H₂O_(l)

(This is also a neutralization reaction.)

Important Points

• The reaction usually occurs in aqueous solution.

• One of the products is often a precipitate, gas, or water.

🧠 Important Tip

If two compounds exchange ions to form new compounds, it is a double decomposition reaction.

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🔥 Thermal Dissociation (Short Note)

Thermal dissociation is a process in which a compound splits into simpler substances when heated, and the reaction is reversible.

When the temperature is lowered, the products can recombine to form the original compound.

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General Form

AB →  A + B

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Examples

1. Ammonium chloride
   
   NH₄Cl_(s) \xrightleftharpoons{heat} NH_3(g) + HCl_(g)
   
   

2. Dinitrogen tetroxide
   [
   N₂O_4(g) \xrightleftharpoons{heat} 2NO_2(g)
   ]

3. Calcium carbonate
   [
   CaCO_3(s) \xrightleftharpoons{heat} CaO_(s) + CO_2(g)
   ]

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🧠 Important Tip

If a substance breaks on heating and reforms on cooling, it shows thermal dissociation.

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🔄 Reversible Reaction

A reversible reaction is a chemical reaction in which the products can react together to reform the original reactants.

It occurs in both forward and backward directions at the same time.

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Symbol

[
A + B \rightleftharpoons C + D
]

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Examples

1. Formation of ammonium chloride
   
   NH_3(g) + HCl_(g) → NH₄Cl_(s)

2. Dinitrogen tetroxide and nitrogen dioxide

3. N₂O_4(g) ⇌ 2NO_2(g) 

4. Haber process
   
   N_2(g) + 3H_2(g)⇌ 2NH_3(g)
    

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🧠 Important Tip

If a reaction can go both forward and backward, it is a reversible reaction.

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⚡ Catalytic Reaction

A catalytic reaction is a chemical reaction in which a substance called a catalyst increases the rate of the reaction without being used up or changed permanently.

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Catalyst

A catalyst is a substance that speeds up a chemical reaction but remains unchanged at the end of the reaction.

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Examples

1. Decomposition of hydrogen peroxide
   
   

2.   2H₂O_2(aq) ---{MnO₂}--->2H₂O_(l) + O_2(g)
   
   

3. Haber process (manufacture of ammonia)
   
   N_2(g) + 3H_2(g) ---{Fe}---> 2NH_3(g)
   
   

4. Contact process (manufacture of tetraoxosulphate VI acid)
   
   2SO_2(g) + O_2(g) ---{V₂O₅}---> 2SO_3(g)
   
   

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🧠 Important Tip

A catalyst alters the rate of a reaction but is not used up in the process.

⚡ IONIC THEORY & ELECTROLYSIS — AT A GLANCE /summary

 

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🔹 Ionic Theory

Ionic compounds dissociate into charged particles (ions) when dissolved in water or melted.

This process is called ionization.

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🔹 Electrolysis

The decomposition of a compound by passing electricity through its molten form or solution.

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🔹 Electrolytes

Substances that conduct electricity in molten or aqueous state and decompose.

Type	Description	Examples
Strong	Fully ionize	NaCl, acids, alkalis
Weak	Partially ionize	CH₃COOH, NH₃
Non	Do not ionize	Sugar, alcohol, oil

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🔹 Important Terms

Term	Meaning
Anode	Positive electrode (oxidation)
Cathode	Negative electrode (reduction)
Cation	Positive ion → cathode
Anion	Negative ion → anode
Electrolytic cell	Container with electrodes & electrolyte

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🔹 Factors That Control Discharge of Ions

1. Position in electrochemical series

2. Concentration of ions

3. Nature of electrodes

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🔹 Products of Electrolysis

Electrolyte	Cathode	Anode
Acidified water	Hydrogen (H₂)	Oxygen (O₂)
Brine (NaCl)	Hydrogen (H₂)	Chlorine (Cl₂)
CuSO₄ solution	Copper (Cu)	Oxygen (O₂)

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🔹 Uses of Electrolysis

• Extraction of metals (Na, Al, Mg)

• Purification of copper

• Electroplating

• Production of H₂, Cl₂, NaOH

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🧠 Important Quick Tip

The ion that is discharged depends on electrochemical series, concentration and electrode used.

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🧪 Carbon and Its Allotropes – Summary

 

Carbon is found in Group IV, Period II of the periodic table. Its electronic configuration is 1s² 2s² 2p². It occurs naturally in different forms called allotropes.

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🔹 Allotropy

Allotropy is the ability of an element to exist in two or more different forms in the same physical state.

• Crystalline allotropes: Diamond, Graphite, Fullerenes
• Amorphous forms: Coal, Charcoal, Coke, Soot, Lampblack

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💎 Diamond

Diamond is a pure crystalline form of carbon with a strong tetrahedral structure.

• Hardest natural substance
• High melting point
• Does not conduct electricity
• Transparent and shiny

Uses: cutting tools, drilling, jewelry, precision instruments.

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✏️ Graphite

Graphite has flat layers of carbon atoms with free electrons.

• Soft and slippery
• Good conductor of electricity
• High melting point

Uses: pencil lead, lubricant, electrodes, crucibles.

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⚽ Fullerenes

Fullerenes (e.g. C₆₀) are spherical carbon molecules called buckyballs. They are used in medicine, electronics and materials science.

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🖤 Amorphous Carbon

• Charcoal – absorbs gases and colours
• Carbon black & lampblack – used in tyres, inks and polish
• Coal – used mainly as fuel

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🪨 Types of Coal

• Peat – about 60% carbon
• Lignite – about 67% carbon
• Bituminous – about 88% carbon
• Anthracite – about 94% carbon (hardest and purest)

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🔥 Destructive Distillation of Coal

Heating coal to a high temperature in the absence of air produces:

• Coke
• Coal gas
• Coal tar
• Ammoniacal liquor

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🔥 Fuel Gases

• Producer gas – CO + N₂
• Water gas –      CO + H₂
• Synthetic gas – CO + H₂

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🧪 Chemical Properties of Carbon

• Burns in oxygen to form CO₂ or CO
• Combines with elements like sulphur and hydrogen
• Acts as a reducing agent in metal extraction
• Is oxidized by strong acids to form CO₂

🧪 Kinetic Theory of Gases at a glance

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🔹 Meaning

The kinetic theory of gases states that gases are made up of tiny particles (molecules or atoms) which are in constant random motion. The behavior of gases is explained based on the motion and collisions of these particles.

🔹 Main Ideas of the Theory

• Gas particles are very small.

• They are far apart from one another.

• They move freely and rapidly in all directions They collide with one another and with the walls of their container.

🔹 Assumptions (Postulates) of the Kinetic Theory of Gases

1. Gases molecules move randomly in straight lines colliding with one another r and with the walls of the container 

2. Collisions between gas molecules and the walls of the container are perfectly elastic (no energy is lost)

3. The volume of the gas molecules is negligible compared to the volume of the container.

4. There are no forces of attraction between gas molecules.

5. The average kinetic energy of gas molecules is directly proportional to the absolute temperature.

6. Gas pressure is caused by the continuous collision of gas molecules with the walls of the container.

🔹 Explanation of Gas Properties Using Kinetic Theory

1. Pressure

Gas pressure is due to the collisions of gas molecules with the walls of the container.

2. Volume

Gases occupy the entire volume of their container because the molecules move freely.

3. Temperature

When temperature increases, gas molecules gain kinetic energy and move faster.

4. Diffusion

Gases mix easily because their particles move freely and randomly.

🔹 Limitations of the Kinetic Theory

• It assumes gas molecules have no volume.

• It ignores forces of attraction between molecules.

• It does not apply well to real gases at high pressure or low temperature.

🧠 Important Tip

Increase in temperature leads to increase in the kinetic energy of gas molecules. 

The study of the relationship between the variables above as regards the behaviour of gases studied by scientists like Boyle's, Charles, Avogadro's, Gay Lussacs, Grahams and Dalton. Each relationship is discussed in separate posts

 

🧪 Laboratory Safety Rules and Guidelines at a glance

🔹 Meaning

Laboratory safety rules are instructions that guide students and scientists on how to work safely in the chemistry laboratory to prevent accidents and injuries.

🔹 General Safety Rules

1. Always wear lab coat, goggles and gloves

2. Do not eat, drink or chew anything in the laboratory

3. Read instructions before starting any experiment

4. Handle chemicals carefully

5. Do not taste or smell chemicals directly

6. Keep the laboratory clean and tidy

7. Work only when the teacher is present

🔹 Handling Chemicals

• Read the label before using any chemical

• Do not mix chemicals unless instructed

• Use small quantities of chemicals

• Do not return unused chemicals to bottles

🔹 Handling Apparatus

• Do not use broken or cracked glassware

• Handle hot objects with tongs

• Turn off gas and electrical appliances after use

• Keep flammable materials away from fire

🔹 In Case of Accident

• Report immediately to the teacher

• Wash spilled chemicals with plenty of water

• Use fire extinguishers or sand for fire

• Do not panic

🔹 Importance of Laboratory Safety

• Prevents accidents

• Protects life and property

• Ensures smooth experiments

• Maintains a good learning environment

🧠 Important Tip

Most laboratory accidents occur due to carelessness and failure to follow safety rules.

🧪 Objective Questions

1. Which of the following is a laboratory safety rule?
A. Eating in the lab
B. Wearing a lab coat
C. Running in the lab
D. Playing with chemicals

2. The main reason for wearing goggles in the laboratory is to
A. look smart
B. protect the eyes
C. increase vision
D. avoid reading

3. Chemicals should never be tasted because they are
A. expensive
B.  poisonous
C. colourless
D.  hot

4. Broken glassware should be
A. used carefully
B. thrown on the floor
C. reported to the teacher
D. ignored

5. Which of the following should be done before starting an experiment?
A. Eat food
B. Read instructions
C. Run around
D. Touch chemicals

6. In case of chemical spill on the skin, one should
A. wipe with cloth
B. wash with plenty of water
C. ignore it
D. cover it

7. Which of the following is NOT allowed in the laboratory?
A. Wearing gloves
B. Drinking water
C. Using tongs
D. Wearing goggles

8. Flammable substances should be kept
A. near fire
B. in open flames
C. away from fire
D. on the floor

9. When heating substances, you should use
A. hands
B. tongs
C. books
D. paper

10. The safest behavior in the laboratory is to
A. follow safety rules
B. rush experiments
C. play with chemicals
D. ignore instructions

✍️ Theory Questions

1. What are laboratory safety rules?

2. State five laboratory safety rules.

3. Give four reasons why safety rules are important in the chemistry laboratory.

4. What should be done when chemicals spill on the skin?

5. Mention four ways to prevent accidents in the laboratory.

🧪 Hazards, Causes and Prevention in the Chemistry Laboratory

🔹 Meaning of Laboratory Hazards

Laboratory hazards are dangerous situations or substances in the chemistry lab that can cause injury, illness, fire, or damage if not handled properly.

⚠️ Common Laboratory Hazards

1. Chemical hazards – toxic, corrosive or flammable chemicals

2. Fire hazards – Bunsen burners, alcohol, gas leaks

3. Glassware hazards – broken test tubes, beakers

4. Electrical hazards – faulty wires, wet hands

5. Biological hazards – harmful microorganisms

🔥 Causes of Laboratory Accidents

1. Carelessness or playing in the lab

2. Not wearing protective clothing

3. Wrong handling of chemicals

4. Spilling chemicals

5. Using broken or damaged equipment

6. Poor ventilation

7. Not following instructions

🛡 Prevention of Laboratory Accidents

1. Always wear lab coat, goggles and gloves

2. Read labels on chemical bottles carefully

3. Do not eat or drink in the lab

4. Handle glassware with care

5. Keep flammable substances away from fire

6. Wash hands after experiments

7. Report spills and accidents immediately

8. Follow the teacher’s instructions

🧠 Important note

Most laboratory accidents occur due to carelessness and improper handling of chemicals 

🧪 Objective Questions

1. A laboratory hazard is
A. a useful chemical
B. a dangerous condition in the laboratory
C. laboratory equipment
D. a laboratory rule

2. Which of the following is a chemical hazard?
A. Broken glass
B. Acid
C. Water
D. Paper

3. Wearing goggles in the laboratory is to
A. look smart
B. prevent eye injury
C. make experiments faster
D. increase concentration

4. Which of the following can cause fire in the laboratory?
A. Sand
B. Spirit lamp
C. Salt
D. Water

5. Spilling chemicals on the skin should be treated by
A. wiping with cloth
B. washing with plenty of water
C. covering with paper
D. ignoring it

6. Which of the following is NOT a laboratory hazard?
A. Broken beaker
B. Toxic gas
C. Notebook
D. Open flame

7. Eating in the laboratory is dangerous because
A. food is expensive
B. chemicals may enter the body
C. it causes noise
D. it wastes time

8. Fire in the laboratory can be caused by
A. acids
B. water
C. flammable liquids
D. glass

9. Which safety equipment protects the hands?
A. Goggles
B. Gloves
C. Lab coat
D. Mask

10. The best way to prevent laboratory accidents is to
A. rush experiments
B. follow safety rules
C. ignore instructions
D. avoid chemicals

✍️ Theory Questions (WAEC / NECO)

1. Define laboratory hazards.

2. List three types of laboratory hazards.

3. State four causes of laboratory accidents.

4. Mention five safety precautions in a chemistry laboratory.

5. Explain why eating in the laboratory is dangerous.

 

Equillibrium at a glance revision

 🧪 Chemical Equilibrium – At a Glance

🔹 Meaning

Chemical equilibrium is the state in a reversible reaction when the rate of the forward reaction equals the rate of the backward reaction.

        

                                   🔹 Key Features

• The reaction is dynamic (still going on).

• Concentrations of reactants and products remain constant.

• It occurs only in a closed system.

🔹 Reversible Reaction

A reversible reaction is one that can go both forward and backward.

Example:

N₂ + 3H₂ ⇌ 3NH₃

🔹 Le Chatelier’s Principle

When a system is in equilibrium and it is disturbed by an external constraint the equilibrium will adjust itself so as to oppose the disturbance in order to restore equilibrium.

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🔹 Factors Affecting Equilibrium

Change	Effect
Increase in concentration of reactants	Shifts equilibrium to the right
Increase in concentration of products	Shifts equilibrium to the left
Increase in pressure (gases)	Favors the side with fewer gas molecules
Increase in temperature Decrease in temperature	Favors the endothermic reaction Favors the exothermic reaction
Catalyst	Does not change equilibrium position

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🔹 Role of a Catalyst

A catalyst alters both forward and backward reactions but does not change the equilibrium position.

🔹 Important Tip

At equilibrium, reactions do not stop — only the rates become equal.

Kinetic Theory of Matter Revision

📌 Kinetic Theory of Matter – At a Glance

Meaning:
The kinetic theory of matter states that all matter is made up of tiny particles which are in constant random motion.

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🔍 Main Ideas

• Matter is made up of tiny particles.
• These particles are always moving.
• There are spaces between the particles.
• There are forces of attraction between particles.
• Particles possess kinetic energy.

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📊 Particles in Different States of Matter

State	Arrangement	Movement
Solid	Closely packed	Vibrate in fixed positions
Liquid	Close but not fixed	Slide past one another
Gas	Far apart	Move freely and rapidly

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🔥 Effect of Heating

• Particles gain more kinetic energy.
• They move faster.
• The substance expands.

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🧪 Evidence that Particles Are in Motion

• Diffusion
• Brownian motion
• Osmosis
• Evaporation
• Expansion on heating

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🎯 Gas Pressure

Gas pressure is caused by continuous collision of gas particles with the walls of the container.

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📝 Important Tip

Increase in temperature → increase in kinetic energy → particles move faster. 

Kinetic Theory of Matter note for students

 According to the kinetic theory the particles that make up matter that is, atoms, molecules and ions are constantly in motion and hence possess kinetic energy. The particles in a given substance do not possess equal amount of energy, therefore, we use the term average kinetic energy of a substance. A change in temperature will cause a change in the average kinetic energy of a system or substances.  Increasing the temperature will lead to an increase in the kinetic energy of a substance and hence it can lead to a change in the state. Matter can exist in any one of three states.

States Of Mater 

1. Solid State-: The particles of mater in the solid state are held together by strong forces of cohesion such that the particles can only vibrate about a fixed point. A solid therefore possess only vibrational motion

 

Properties of a solid

a. It has definite or fixed shape
b. It has definite or fixed volume and
c. It cannot be compressed

                                                                    The solid state

2 The liquid State- The particles of a liquid are held by weaker forces of cohesion than those in solids. As a result, the particles in liquids can vibrate as well as translate (flow / move). Thus, the particles of a liquid possess both vibrational and translational energy.

Properties of a liquid

a. Have no definite shape but takes the shape of its container
b. Have a definite or fixed volume and
c. It cannot be compressed.

                                                                  The liquid state

3. The Gaseous State: - The forces of cohesion in gaseous molecules/particles are negligible as a result gaseous particles occupies their entire volume. 

Properties of a gas

a. A gas has no definite or fixed shape
b. No definite or fixed volume and
c. It can be compressed.

                                                                The Gaseous state

States of Matter and Particle Motion

• Solid:
  Particles are closely packed and vibrate about fixed positions. They have low kinetic energy.

• Liquid:
  Particles are close but can slide past one another. They have moderate kinetic energy.

• Gas:
  Particles are far apart and move freely at high speed. They have high kinetic energy.

Evidence/Phenomenon to show that the particles of mater are in constant motion

1. DIFFUSION: this is the movement of solute particles from a region of higher concentration to a region of lower concentration.
diffusion can occur in solids, in liquids as well as in gases. It is fastest in gases and slowest in solids

Example of Diffusion

*** When you open a bottle of perfume in one corner of a room, after a short time people in other parts of the room can smell it. (diffusion in gases)

*** If you drop a crystal of potassium permanganate into water, the purple colour slowly spreads through the water without stirring.

That spreading is diffusion (diffusion in liquids)

***When a piece of copper is placed in contact with a piece of zinc and the two metals are heated for a long time, atoms of copper slowly move into the zinc and atoms of zinc move into the copper. After some time, an alloy (brass) is formed (diffusion in solids).

2. BROWNIAN MOTION: - This is the irregular or zigzag movement of small particles in a liquid or gas due to constant collisions with the molecules of the liquid or gas

Example of Brownian Motion

When smoke particles are seen in a beam of sunlight in a dark room, they move about randomly. That movement is also Brownian motion.

3. OSMOSIS: - This is the movement of water molecules from a region of higher concentration to a region of lower concentration through a semi-permeable membrane

Example of Osmosis

If a peeled potato is cut into strips and placed in salt water, after some time the potato becomes soft and shrinks.

This happens because water moves out of the potato cells (from a region of higher water concentration inside the potato to a region of lower water concentration in the salt solution) through a semi-permeable membrane.
This movement of water is called osmosis.

4. Evaporation

When water is left in an open container, it slowly changes into vapour even without boiling.
This happens because some water particles are moving fast enough to escape from the liquid into the air.

5. Expansion when heated When a solid, liquid or gas is heated, it expands.

This is because its particles move faster and spread farther apart.

Example:
A heated metal rod becomes longer.

6.  Sublimation

Substances like camphor or naphthalene disappear slowly when left in the open.
Their particles move directly from solid to gas because they are in constant motion.

7. Gas pressure

Air inside a balloon push against the walls of the balloon.
This is due to continuous movement and collision of gas particles.

   

 Change of State

A change of state is the physical process by which a substance changes from one state of matter to another without any change in its chemical composition. These changes occur as a result of gain or loss of heat energy, which affects the kinetic energy of the particles.

According to the kinetic theory of matter, all matter is made up of tiny particles that are in constant motion. The speed of these particles determines the state of matter.

Types of Change of State

1. Melting (Solid → Liquid)

Melting occurs when a solid is heated and changes into a liquid. Heat energy supplied increases the kinetic energy of the particles, causing them to vibrate more rapidly until they overcome the forces holding them together.

Example: Ice melting into water.

2. Freezing (Liquid → Solid)

Freezing is the change of a liquid into a solid when heat is removed. The particles lose kinetic energy and become fixed in position.

Example: Water freezing to form ice.

3. Boiling / Vaporization (Liquid → Gas)

Boiling occurs when a liquid changes into a gas at a fixed temperature called the boiling point. At this point, particles gain enough kinetic energy to escape from the liquid.

Example: Water changing to steam at 100°C.

  4. Condensation (Gas → Liquid)

Condensation occurs when a gas loses heat and changes into a liquid. The particles lose kinetic energy and move closer together.

Example: Water droplets forming on a cold surface.

5. Sublimation (Solid → Gas)

Sublimation is the direct change of a solid into a gas without passing through the liquid state.

Examples: Iodine, naphthalene, and dry ice.

6. Deposition (Gas → Solid)

Deposition is the direct change of a gas into a solid without passing through the liquid state.

Example: Frost formation.

Role of Heat Energy

• Heating: Increases particle kinetic energy → change to a higher energy state.

• Cooling: Decreases particle kinetic energy → change to a lower energy state.

Important Exam Points 

• Change of state is a physical change.

• No new substance is formed.

• Temperature remains constant during change of state until the process is complete.

• Explained using kinetic energy of particles.

OBJECTIVE QUESTIONS

1. The kinetic theory of matter states that matter is made up of

A. ions
B. molecules
C. tiny particles
D. compounds

2. According to the kinetic theory, particles of matter are always
A. at rest
B. vibrating only
C. in constant motion
D. fixed in position

3. Which of the following best explains diffusion?
A. Attraction between particles
B. Movement of particles from high to low concentration
C. Chemical reaction
D. Expansion of solids

4. Brownian motion is caused by
A. gravity
B. heat
C. collision of molecules
D. evaporation

5. Which state of matter has particles that are far apart and move freely?
A. Solid
B. Liquid
C. Gas
D. Plasma

6. The force of attraction between particles is strongest in
A. gases
B. liquids
C. solids
D. vapour

7. When a solid is heated, its particles
A. stop moving
B. move faster
C. move closer
D. disappear

8. Which of the following shows that gas particles are in motion?
A. Crystallization
B. Diffusion of gas
C. Freezing
D. Condensation

9. The random movement of smoke particles in air is called
A. diffusion
B. evaporation
C. Brownian motion
D. osmosis

10. The kinetic energy of particles increases when
A. temperature decreases
B. temperature increases
C. pressure decreases
D. volume decreases

11. In which state of matter do particles vibrate about fixed positions?
A. Gas
B. Liquid
C. Solid
D. Vapour

12. Osmosis occurs because of
A. random motion of particles
B. chemical reaction
C. evaporation
D. heating

13. Which of the following best describes particles in a liquid?
A. Fixed and tightly packed
B. Far apart and free
C. Close together and able to move
D. Completely motionless

14. The spreading of perfume in a room is due to
A. osmosis
B. diffusion
C. evaporation
D. freezing

15. The kinetic theory explains that gas pressure is due to
A. weight of gas
B. collisions of particles with container walls
C. chemical reactions
D. gravity

16. Which of these is evidence that particles of matter are in motion?
A. Expansion when heated
B. Rusting
C. Burning
D. Melting

17. When temperature increases, the average kinetic energy of particles
A. decreases
B. remains constant
C. increases
D. becomes zero

18. Particles in a gas have
A. very strong forces of attraction
B. weak forces of attraction
C. no energy
D. fixed positions

19. Diffusion occurs fastest in
A. solids
B. liquids
C. gases
D. crystals

20. The kinetic theory of matter is used to explain
A. chemical reactions
B. structure of atoms
C. behaviour of solids, liquids and gases
D. electricity

Theory Questions 

A crystal of potassium permanganate is dropped into a beaker of water. After some time, the purple colour spreads throughout the water even without stirring.

(a) Name the process responsible for this.
(b) Explain why the colour spread

 region of lower concentration through a semi permeable membrane.

Boyles Law

 BOYLE’S LAW

Boyles states that the volume of a given mass of gas is inversely proportional to the pressure provided the temperature remains constant.

 This means that:

• When pressure increases, volume decreases

• When pressure decreases, volume increases

Mathematically,

       V α 1/P

       V = k/P

       PV = k

Hence,         P₁V₁ = P₂V₂

Boyle’s law can be represented graphically as shown below.

1. Pressure vs Volume (inverse curve):

 


2. Pressure vs 1/Volume (straight line):

  

The graph shows that if the pressure is doubled, the volume is reduced to half its former value and if it is halved, the volume is doubled.

 

EXPLANATION OF BOYLE’S LAW USING THE KINETIC THEORY

If a gas is compressed into a smaller space (when the volume of fixed mass of gas is decreased) the molecules of the gas will collide with each other more rapidly ( i.e the gas particles hit the walls of the container more often). This gives rise to an increase in pressure. However, If the volume is increased, the particles have more space to move, so the pressure decreases. 

Examples of Boyle’s Law

• When you push the plunger of a syringe, the air inside is compressed and the pressure increases.

• A bicycle pump works because reducing the volume of air increases its pressure.

Conclusion

Boyle’s Law shows the relationship between the pressure and volume of a gas at constant temperature.

example of calculations on Boyles law 

1. 200cm3 of a gas has a pressure of 510mmHg. What will be its volume if pressure in increased to 780mmHg, assuming there is no change in temperature?

    Solution:

    V₁ = 200cm³, P₁ = 510mmHg,       P₂ = 780mmHg V₂ =?

    Using the expression for Boyle’s law:

       P₁V₁ = P₂V₂

   V₂ = P₁V₁  =  510mmHg x 200cm³ = 130.769 = 131 cm³
                P₂              780mmHg

EASYKEMISTRY

CHEMISTRY TEST – BOYLE’S LAW
Time: 30 minutes

Name: __________________________ Class: __________ Date: __________

Choose the correct option from A–D

1. Boyle’s law states that for a fixed mass of gas at constant temperature,
A. pressure is directly proportional to volume
B. pressure is inversely proportional to volume
C. pressure is equal to volume
D. pressure is proportional to temperature

2. Which of the following is kept constant in Boyle’s law?

A. Pressure
B. Volume
C. Temperature
D. Mass and volume

3. If the volume of a gas is reduced to half at constant temperature, the pressure will

A. remain the same
B. be doubled
C. be halved
D. become zero

4. The mathematical expression for Boyle’s law is

A. V = kP
B. PV = k
C. P + V = k
D. P = k 
    V

5. A gas has a volume of 20 cm³ at a pressure of 2 atm. What will be its volume at 4 atm?

A. 5 cm³
B. 10 cm³
C. 20 cm³
D. 40 cm³

6. According to Boyle’s law, when pressure decreases, the volume of a gas

A. decreases
B. increases
C. remains constant
D. becomes zero

7. Which of the following devices works based on Boyle’s law?

A. Thermometer
B. Barometer
C. Syringe
D. Voltmeter

8. A graph of pressure against volume for a fixed mass of gas at constant temperature is

A. a straight line
B. a curve
C. a horizontal line
D. a vertical line

9. A gas occupies 40 cm³ at 1 atm. What will be its pressure if the volume is reduced to 10 cm³?

A. 2 atm
B. 3 atm
C. 4 atm
D. 5 atm

10. Boyle’s law is valid only when the

A. pressure is constant
B. volume is constant
C. temperature is constant
D. gas is solid

11.   A gas occupies 100 cm³ at 2 atm. What will be its volume at 1 atm, temperature remaining constant?

A. 50 cm³
B. 100 cm³
C. 150 cm³
D. 200 cm³

12.   Which of the following graphs best represents Boyle’s law?

A. Pressure vs Temperature
B. Volume vs Temperature
C. Pressure vs Volume
D. Mass vs Volume

13.  If the pressure of a gas is increased four times, its volume will become
A. four times
B. half
C. one quarter
D. double

14.  A gas occupies 60 cm³ at 3 atm. What will be its volume at 6 atm?
A. 10 cm³
B. 20 cm³
C. 30 cm³
D. 40 cm³

15.  Boyle’s law is useful in explaining the operation of
A. a thermometer
B. a hot-air balloon
C. a bicycle pump
D. a barometer

16.  A gas has a pressure of 4 atm and a volume of 50 cm³. What is the value of PV?
A. 50
B. 100
C. 150
D. 200

17.  A gas occupies 80 cm³ at 5 atm. What will be its volume at 10 atm?
A. 40 cm³
B. 60 cm³
C. 80 cm³
D. 160 cm³

18.  Which of the following statements is correct?
A. Pressure decreases when volume decreases
B. Pressure increases when volume decreases
C. Pressure is constant when volume changes
D. Pressure does not depend on volume

19.  Boyle’s law does not apply when
A. temperature is constant
B. pressure is constant
C. temperature changes
D. volume changes

20.  A gas at 1 atm occupies 500 cm³. What pressure will it have if its volume becomes 250 cm³?
A. 0.5 atm
B. 1 atm
C. 2 atm
D. 4 atm

THEORY QUESTIONS

1. State Boyle’s law. Explain the conditions under which the law is valid.

2. Describe an experiment to verify Boyle’s law. Include a labeled diagram of the apparatus used.

3. Define pressure and volume as used in Boyle’s law and state their SI units.

4. A fixed mass of gas occupies a volume of 40 cm³ at a pressure of 100 kPa.
   Calculate the new volume when the pressure is increased to 200 kPa, assuming temperature remains constant.

5. Explain why Boyle’s law does not hold for real gases at very high pressure.

6. State the mathematical expression of Boyle’s law and explain the meaning of each symbol used.

7. Sketch and explain the graph of pressure against volume for a gas obeying Boyle’s law.

8. Sketch and explain the graph of pressure against the reciprocal of volume (1/V) for Boyle’s law.

9. A gas initially at pressure ( P1 ) and volume ( V1 ) changes to pressure ( P2 ) and volume ( V2 ).
   Derive the Boyle’s law equation relating these quantities.

10. Mention two practical applications of Boyle’s law and explain any one of them.

11.  Explain Boyle's law using the kinetic theory

 

GENERAL GAS EQUATION at a glance

 GENERAL GAS EQUATION

The General Gas Equation is a formula that shows the relationship between the pressure, volume, and temperature of a gas.

It is a combination of both Boyle’s and Charles law.

It is written as:

PV = K

T

P1V1  = P2V2 

   T             T

Explanation

The general gas equation combines the three gas laws:

• Boyle’s Law (pressure and volume)

• Charles’ Law (volume and temperature)

• Pressure Law (pressure and temperature)

It helps us calculate any one of the gas properties if the others are known.

Example

If the pressure, volume, and temperature of a gas are known, the number of moles can be calculated using:

n = PV
      RT

THEORY QUESTIONS 

1. What is the volume at s.t.p of a fixed mass of a gas that occupies 700cm³ at 25^oC and 0.84 x 10⁵ Nm^-2pressure?

    Solution:

   T₁ = 273K, P₁ = 1.01 x 10⁵Nm^-2, T₂ = 25^oC = (25 + 273) = 298K, P₂ = 0.84 x 10⁵Nm^-2,

    V₂ = 700cm³, V₁ =?

   Using the general gas equation

P₁V₁ = P₂V₂
   T₁        T₂

     V₁ = P₂V₂T₁ = 0.84 x 10⁵Nm^-2 x 700cm³ x 273K = 533.337 =533cm³
              P₁T₂          1.01 x 10⁵Nm^-2 x 298K

     = 0.84 x 10⁵Nm^-2 x 700cm³ x 273K = 533.337 
                       1.01 x 10⁵Nm^-2 x 298K

                                           =533cm³
              

Conclusion

The general gas equation is very useful in solving problems involving gases in chemistry.

OBJECTIVE QUESTIONS 

1. The general gas equation is expressed as
   A. ( PV = RT )
   B. ( PV = nRT )
   C. ( V = nRT )
   D. ( P = nRT )

2. In the general gas equation ( PV = nRT ), the symbol R represents
   A. rate constant
   B. gas density
   C. universal gas constant
   D. pressure constant

3. Which of the following quantities must be in Kelvin when using the gas equation?
   A. Pressure
   B. Volume
   C. Temperature
   D. Amount of gas

4. The SI unit of pressure used in the general gas equation is
   A. atmosphere
   B. mmHg
   C. pascal
   D. bar

5. The general gas equation combines which gas laws?
   A. Boyle’s and Charles’ laws only
   B. Boyle’s, Charles’ and Graham’s laws
   C. Boyle’s, Charles’ and Avogadro’s laws
   D. Dalton’s and Avogadro’s laws

6. The value of the universal gas constant R is approximately
   A. 0.082 J mol⁻¹ K⁻¹
   B. 8.31 J mol⁻¹ K⁻¹
   C. 82.06 J mol⁻¹ K⁻¹
   D. 1.00 J mol⁻¹ K⁻¹

7. Which of the following is the unit of R when pressure is in pascals?
   A. L atm mol⁻¹ K⁻¹
   B. J mol⁻¹ K⁻¹
   C. cm³ atm mol⁻¹ K⁻¹
   D. Nm⁻² mol⁻¹

8. If the temperature of a gas increases while pressure is constant, the volume will
   A. decrease
   B. remain constant
   C. increase
   D. become zero

9. One mole of an ideal gas occupies 22.4 dm³ at
   A. 0°C and 1 atm
   B. 25°C and 1 atm
   C. 0°C and 760 Pa
   D. 273°C and 1 atm

10. The general gas equation is most accurate when gases
    A. are at very high pressure
    B. are at very low temperature
    C. behave ideally
    D. are strongly interacting

11. In the equation ( PV = nRT ), the symbol n represents
    A. number of molecules
    B. number of atoms
    C. amount of gas in moles
    D. mass of gas

12. Which condition is necessary for gases to obey the general gas equation?
    A. High pressure and low temperature
    B. Low pressure and high temperature
    C. Presence of strong intermolecular forces
    D. Gas must be liquid

13. If pressure is doubled and temperature remains constant, the volume will
    A. double
    B. halve
    C. remain the same
    D. become zero

14. The general gas equation is also known as the
    A. combined gas law
    B. Dalton’s law
    C. ideal gas equation
    D. Graham’s law

15. Which of the following is NOT an assumption of the kinetic theory of gases?
    A. Gas particles are in constant motion
    B. Gas molecules occupy negligible volume
    C. There are strong forces between molecules
    D. Collisions are elastic

16. A gas has a volume of 2 dm³ at 300 K. What will be its volume at 600 K if pressure is constant?
    A. 1 dm³
    B. 2 dm³
    C. 3 dm³
    D. 4 dm³

17. Increasing the number of moles of a gas at constant temperature and pressure will
    A. decrease volume
    B. not affect volume
    C. increase volume
    D. decrease pressure

18. Which of the following quantities is directly proportional to pressure according to the gas equation?
    A. Volume
    B. Temperature
    C. Amount of gas
    D. Both B and C

19. The equation ( \frac{PV}{T} = \text{constant} ) is derived from
    A. Boyle’s law
    B. Charles’ law
    C. Combined gas law
    D. Graham’s law

20. The general gas equation is mainly used to
    A. calculate molecular mass only
    B. explain diffusion
    C. relate pressure, volume, temperature and amount of gas
    D. describe electrolysis

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IDEA GAS EQUATION at a glance

 IDEAL GAS EQUATION

The ideal gas equation is:

                                    PV = nRT

Where:

• P = pressure of the gas

• V = volume of the gas

• n = number of moles of the gas

• R = universal gas constant

• T = absolute temperature (in Kelvin)

Value of R 

•  R = 8.31{J mol-1K-1

📌 Important tip:
Temperature must always be converted to Kelvin (K) using

T(K) = t(0C) + 273

This equation states that for an ideal gas PV/T is a constant.

That is, PV = R 
               T

(R = molar gas constant)

             

             PV = RT

That is, for n mole of a gas, the equation becomes

             PV = nRT

 

CALCULATIONS

1.  Calculate the number of moles present in a certain mass of gas occupying 6.5dm³ at     3atm and 15^oC (R = 0.082atmdm³K^-1mol^-1)

    Solution:

    V = 6.5dm³, P = 3atm, T = 15^oC = (15 + 273)K = 288K, n =?

    Using PV = nRT

           n = PV = 3atm x 6.5dm³= 0.8257

                RT    0.082atmdm³K^-1mol^-1 x 288K

    Number of moles = 0.83 mole

Finding the volume of a gas

2. Calculate the volume occupied by 2 moles of an ideal gas at a pressure of 1.0 × 10⁵ Pa and a temperature of 27°C.

R = 8.31J mol^-1,/sup>K-1

Solution:
Convert temperature to Kelvin:

T = 27 + 273 = 300K

Use the formula:

PV = nRT

Make V the subject:

V =          nRT

                   P

Substitute values:

         V =   2 x 8.31 x 300
                    1.0 x 10⁵

     V =        4986
                 100000

V = 0.0499 m³

Answer:
V = 4.99 x10^-2m³

Finding pressure

Question:
3. A gas occupies a volume of 0.02 m³ at 300 K and contains 1 mole of gas. Calculate its pressure.

Solution:

           PV = nRT

Make P the subject:

               P = nRT
                      V

Substitute:

P =          1 x 8.31 x 300
                       0.02

P =       2493
             0.02

P = 124650 Pa

Answer:
            P = 1.25 x 10⁵ Pa

Finding number of moles

Question:
4. Calculate the number of moles of a gas that occupies 0.01 m³ at 27°C and 1.0 × 10⁵ Pa.

Solution:
Convert temperature:

T = 27 + 273 = 300K

Use:

         n =    PV
                  RT

Substitute:

             n = 1.0 x 10⁵ x 0.01
                      8.31 x 300

                n = 1000
                       2493

               n = 0.40 mol

Answer:
                    n = 0.40 mol

Converting cm³ to m³

Question:
5.  A gas occupies 500 cm³ at 27°C and 1.0 × 10⁵ Pa. Find the number of moles.

Solution:
Convert volume:
500 cm³ = 5.0 x 10^-4 x m³

Convert temperature:

T = 300K

n =   PV
        RT

n = 1.0 x 10⁵ x 5.0 x 10^-4
          8.31 x 300

       n =    50
              2493

n = 0.020 mol

Answer:
                n = 0.02 mol

🔑 Important EXAM TIPS

• Always convert °C to K

• Convert cm³ to m³

• Write formula first

• Show substitutions clearly

OBJECTIVE QUESTION

1. The ideal gas equation is written as
   A. PV = RT 
   B. P = VRT 
   C. PV = nRT 
   D.  V = nRP 

2. In the ideal gas equation PV = nRT the symbol n represents
   A. number of molecules
   B. number of particles
   C. number of moles
   D. molar volume

3. Which of the following is the correct unit of the gas constant R?
   A. J K⁻¹
   B. J mol⁻¹
   C. J mol⁻¹ K⁻¹
   D. Pa m³ mol⁻¹

4. In gas calculations, temperature must be expressed in
   A. Celsius
   B. Fahrenheit
   C. Kelvin
   D. Centigrade

5. A gas occupies a volume of 0.02 m³ at 300 K and 1 mole. Calculate the pressure.
   A. ( 4.2 x10^4 ) Pa
   B. ( 8.3 x 10^4 ) Pa
   C. ( 1.25 x10^5 ) Pa
   D. ( 2.49 x 10^5 ) Pa

6. Which of the following is an assumption of an ideal gas?
   A. Gas molecules attract one another
   B. Gas molecules occupy large volumes
   C. Gas molecules are in constant random motion
   D. Gas molecules move in one direction

7. If the temperature of a gas increases while pressure remains constant, the volume will
   A. decrease
   B. remain constant
   C. increase
   D. become zero

8. Which law is combined with Boyle’s and Charles’ laws to give the ideal gas equation?
   A. Dalton’s law
   B. Avogadro’s law
   C. Graham’s law
   D. Faraday’s law

9. The volume of a gas is 500 cm³ at STP. What is this volume in m³?
   A.  5.0 x 10^-2
   B. 5.0 x 10^-3
   C. 5.0  x 10^-4
   D.  5.0 x 10^-5

10. Real gases behave most like ideal gases at
    A. low temperature and high pressure
    B. high temperature and high pressure
    C. low temperature and low pressure
    D. high temperature and low pressure

DALTON’S LAW OF PARTIAL PRESSURE

 DALTON’S LAW OF PARTIAL PRESSURE

This law state that in a mixture of gases which do not react chemically together, the total pressure exerted by the mixture of gases is equal to the sum of the partial pressure of the individual gases that make up the mixture.

Mathematically, the law can be expressed as:

     P_total = P_A + P_B +P_C........P_n

 

Where P_total is the total pressure of the mixture and P_A, P_B, P_C are the partial pressure exerted separately by the individual gases A, B, C that make up the mixture.

The pressure each constituent gas exerts is called partial pressure and is expressed as

Partial pressure of gas A (P_A) = Number of moles of gas A  x P_total

                              Total number of moles of gas in mixture

 

That is, P_A = n_A x P_total

                    n_A + n_B + n_C

If the gas is collected over water, it is likely to be saturated with water vapour and the total pressure becomes

       P_total = P_gas + P_{water vapour}

       P_gas = P_total – P_{water vapour}

 

 

CALCULATION ON THE LAW

A gaseous mixture containing 64g of O₂ and 70g of N₂ exerts a total pressure of 1.8oatm. What is the partial pressure exerted by oxygen in the mixture?

Solution:

Molar mass of O₂ = 16 x 2 = 32gmol^-1

Molar mass of N₂ = 14 x 2 = 28gmol^-1      

Number of mole of O₂ = 64g            = 2.0mole
                                       32gmol^-1

Number of mole of O₂ = 70g     = 2.5mole
                                     28gmol^-1

Total number of moles of gases in mixture = 2.0 + 2.5 = 4.5 mole

Partial pressure of O₂ = 2.0 x 1.80 = 0.80atm

     THEORY QUESTIONS 

1.     State Dalton’s of partial pressure.

2.     Calculate the pressure at 27^oC of 16.0g O₂ gas occupying 2.50dm³

3.     A certain mass of hydrogen gas collected over water at 10^oC and 760mmHg pressure has     a volume of 37cm³. Calculate the volume when it is dry at s.t.p (Saturated vapour pressure of water at 10^oC =1.2mmHg)

 

Charles’ Law at a glance Easy Kemistry. 🧪📘

 CHARLES’ LAW

Charles’ law states that the volume of a fixed mass of gas at constant pressure is directly proportional to its temperature in the Kelvin scale.

This means that:

• When temperature increases, volume increases

• When temperature decreases, volume decreases

Mathematically,

                         V α T

                         V = k/T

                                        V   = k
                                        T

Hence,             V₁ = V₂

                         T₁   T₂

where
(V) = volume of the gas
(T) = absolute temperature (in Kelvin)

The graphical representation of Charles’ law is as shown below:

 

 

 

 

 

 

 

 

 

EXPLANATION OF CHARLES’ LAW USING THE KINETIC THEORY

When a gas is heated, (at constant pressure) its particles gain energy and move faster, pushing the walls of the container outward (to maintain the same number of collisions on the walls of container) so the volume increases. When the gas is cooled, the particles move more slowly and the volume decreases.

Examples

• A balloon expands when heated and shrinks when cooled.

• Hot air causes a hot-air balloon to rise because the air expands.

Conclusion

Charles’ Law shows how the volume of a gas changes with temperature at constant pressure.

Examples of calculation based on Charles law

A certain mass of a gas occupies 300cm³ at 35^oC. At what temperature will it have its volume reduced by half assuming its pressure remains constant?

       Solution:

       V₁ = 300cm³, T₁ = 35^oC = (35 + 273)K = 308K, V₂ = V₁/2 = 300/2 = 150cm³, T₂ = ?

       Using the formula for Charles’ law

                     V₁ = V₂
                      T₁    T₂

       T₂ = V₂T₁ = 150cm³ x 308K = 154K
                  V₁              300cm³

     

EASY KEMISTRY

CHEMISTRY TEST – CHARLES’ LAW
Time: 30 Minutes

Name: __________________________ Class: __________ Date: __________

Choose the correct option from A–D

1. Charles’ law states that for a fixed mass of gas at constant pressure,
A. volume is inversely proportional to temperature
B. volume is directly proportional to temperature
C. pressure is proportional to temperature
D. pressure is inversely proportional to volume

2. Which of the following must remain constant for Charles’ law to apply?
A. Volume
B. Pressure
C. Temperature
D. Mass and volume

3. The temperature used in Charles’ law calculations must be in
A. degrees Celsius
B. degrees Fahrenheit
C. Kelvin
D. Centigrade

4. If the temperature of a gas is increased, its volume will
A. decrease
B. remains constant
C. increase
D. become zero

5. The mathematical expression for Charles’ law is
A. PV = k

B.  V = k 
      T

C.  PV= k 

D.  V = k 
      T

6. A gas has a volume of 200 cm³ at 300 K. What will be its volume at 600 K?
A. 100 cm³
B. 200 cm³
C. 300 cm³
D. 400 cm³

7. Which of the following is an application of Charles’ law?
A. Thermometer
B. Hot-air balloon
C. Barometer
D. Syringe

8. When a gas is cooled, its volume
A. increases
B. decreases
C. remains the same
D. doubles

9. A gas occupies 150 cm³ at 300 K. What will be its volume at 600 K?
A. 75 cm³
B. 100 cm³
C. 300 cm³
D. 450 cm³

10. Charles’ law is valid only when
A. pressure is constant
B. volume is constant
C. temperature is constant
D. mass is constant

11. A gas occupies 50 cm³ at 250 K. What will be its volume at 500 K?
A. 25 cm³
B. 50 cm³
C. 75 cm³
D. 100 cm³

12. According to Charles’ law, volume is directly proportional to
A. pressure
B. temperature
C. mass
D. density

13. What happens to the volume of a gas if its temperature is halved?
A. It doubles
B. It halves
C. It remains constant
D. It becomes zero

14. A gas has a volume of 300 cm³ at 300 K. Find its volume at 150 K.
A. 150 cm³
B. 200 cm³
C. 300 cm³
D. 450 cm³

15. The temperature 0°C is equal to
A. 100 K
B. 273 K
C. 373 K
D. 0 K

16. Which of the following instruments shows the effect of Charles’ law?
A. Syringe
B. Thermometer
C. Hot-air balloon
D. Barometer

17. If a gas expands, its temperature must have
A. decreased
B. increased
C. remained constant
D. become zero

18. A gas has a volume of 120 cm³ at 300 K. What will be its volume at 600 K?
A. 60 cm³
B. 120 cm³
C. 180 cm³
D. 240 cm³

19. Charles’ law applies only to

A. solids
B. liquids
C. gases
D. metals

20. When the temperature of a gas is reduced to zero Kelvin, its volume becomes

A. maximum
B. constant
C. minimum
D. zero

THEORY QUESTIONS  

1.  (a) State Charles’ law.
     (b) Define the terms volume and absolute temperature as used in the law.

................................................................................................................................

2.    Explain why the volume of a gas increases when it is heated at constant pressure.

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3.    A gas has a volume of 200 cm³ at 300 K. Calculate its volume at 600 K, assuming the pressure remains constant.

---

4. (a) Why must temperature be measured in Kelvin when using Charles’ law?
    (b) What would happen if Celsius were used instead?

---

5.    State two everyday applications of Charles’ law and explain one of them.

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Balancing Chemical Reactions at a glance

 

Balancing Chemical Reactions

A chemical equation shows how substances react to form new substances.
A balanced chemical equation has the same number of each type of atom on both sides of the equation.

This is based on the Law of Conservation of Mass, which states that matter is neither created nor destroyed in a chemical reaction.

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Steps for Balancing Chemical Equations

1. Write the correct chemical formula for all reactants and products.

2. Count the number of each atom on both sides.

3. Adjust the coefficients (numbers in front of formulas) to make the atoms equal.

4. Do not change subscripts in the formulas.

5. Check that all atoms are balanced.

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Examples

1. Formation of Water

Unbalanced:
H₂ + O₂ → H₂O

Balanced:
2H₂ + O₂ → 2H₂O

2. Formation of Magnesium Oxide

Unbalanced:
Mg + O₂ → MgO

Balanced:
2Mg + O₂ → 2MgO

3. Reaction of Hydrogen and Chlorine

Unbalanced:
H₂ + Cl₂ → HCl

Balanced:
H₂ + Cl₂ → 2HCl

4. Combustion of Methane

Unbalanced:
CH₄ + O₂ → CO₂ + H₂O

Balanced:
CH₄ + 2O₂ → CO₂ + 2H₂O

5. Reaction of Zinc with Hydrochloric Acid

Unbalanced:
Zn + HCl → ZnCl₂ + H₂

Balanced:
Zn + 2HCl → ZnCl₂ + H₂

Conclusion

Balancing chemical equations ensures that the number of atoms on both sides is equal, showing that mass is conserved in chemical reactions.

EASYKEMISTRY

Chemistry Test – Balancing Chemical Equations
Time: 30 Minutes

Name: __________________________ Class: __________ Date: __________

SECTION A – Balance the following equations

(Each question carries 2 marks)

1.     H₂ + N₂ → NH₃

2.     Fe + H₂O → Fe₃O₄ + H₂

3.     Na + H₂O → NaOH + H₂

4.     Al + O₂ → Al₂O₃

5.     CaCO₃ → CaO + CO₂

6.     KClO₃ → KCl + O₂

7.     Zn + HNO₃ → Zn(NO₃)₂ + H₂

8.     C₂H₆ + O₂ → CO₂ + H₂O

9.     Pb(NO₃)₂ → PbO + NO₂ + O₂

10.     Cu + HNO₃ → Cu(NO₃)₂ + NO₂ + H₂O

SECTION B – Answer all questions

1.      What is meant by a balanced chemical equation?

    

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2.      State the law that is obeyed when chemical equations are balanced.

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3.     Why is it important to balance chemical equations

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4.     Why should subscripts not be changed when balancing equations?

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5. Write a balanced chemical equation for the reaction between magnesium and dilute hydrochloric acid.

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END OF TEST

STATES OF MATTER

Matter can exist in three states,  that is, solid, liquid and gaseous The fundermental difference between these three is the forces of attraction (cohesive forces) between the particles, 

SOLID( strong cohesive forces)	LIQUID(weak cohesive forces)	GASES
Have definite shape and volume	Have no definite shape but definite volume	Have no definite shape and volume
Very dense	Less dense	Least dense
Incompressible	cannot be compressed	Compressible
Fixed mass	Substances have a fixed mass	Fixed mass
Particle vibrate and rotate about a fixed point	Particles can vibrate as well as  move about within a restricted space	Particles move about constantly at great speed and at random

      

CHANGE OF STATE

MELTING

Melting is the physical process where a substance changes from a solid to a liquid. When a solid is heated, the particles acquire greater kinetic energy and move violently. A point is reached when the forces of vibration overcome the cohesive forces holding the solid particles together and the crystalline structure collapses. The particles are no longer held in fixed positions but are free to move about and the liquid state is reached. The temperature at which this occurs is called the melting point of the solid.

 

BOILING

When a liquid is heated, the rate of evaporation increases and the value of the saturated vapour pressure equal the prevailing atmospheric pressure. When this happens, the liquid is said to boil and the temperature at which this happen is known as the boiling point of the liquid.

The boiling point of a liquid change with change in atmospheric pressure. If the pressure is raised, the boiling point will increase and if the pressure is lowered the boiling point will decrease. Also, the presence of impurities increases the boiling point of a liquid.

EVAPORATION

Evaporation is the process of vapourization of liquids at all temperatures. When the surface of a liquid is exposed, the molecules near the surface of the liquid will acquire extra kinetic energy, large enough to enable them break away from the cohesive force binding them to the neighbouring particles. Once free, they escape from the liquid surface to become molecules in the vapour state.

 

Evaporation results in decrease in the volume of liquid and lowering the temperature of the liquid, therefore it causes cooling. Also, it occurs at all temperature but increases with increase in temperature. In addition, it is slower in electrovalent liquids than in covalent liquids.

 

DIFFERENCES BETWWEEN EVAPORATION AND BOILING

EVAPORATION	BOILING
Takes place at the surface of the liquid	Involves the entire volume of the liquid
Takes place at all temperature	Takes place at a fixed temperature

 

CONDENSATION AND FREEZING

Condensation is a process whereby a vapour loses some of its kinetic energy to a colder body and changes into the liquid state.

When a liquid cools, it loses heat energy to its surroundings, causing its temperature to drop. If the cooling continues, the temperature of the liquid keeps dropping until it reaches the freezing point of the liquid. At this temperature, the liquid changes into solid.

 

EVALUATION

1.     Describe the melting process of a solid.

2.     State two differences between evaporation and boiling.

 

KINETIC THEORY OF GASES

The theorypostulates the following for an ideal or perfect gas:

Gas molecules are in constant, rapid, straight motion, colliding with one another and with the walls of the container.

 

The collision of gas molecules is perfectly elastic.

The total volume of the gas molecule is negligible compared to the volume of the container.

The force of attraction between the gas molecules is negligible.

The average kinetic energy of the molecule is a measure of the temperature of the gas molecules.

 

PHENOMENA SUPPORTING THE KINETIC THEORY OF GASES

Brownian motion: This is the constant, irregular movement of particles in a liquid or gas. It shows that gas molecules are in constant motion.

Diffusion: Diffusion is the movement of particles from a region of higher concentration to lower concentration. Diffusion is common in gases and it results from the random movement of particles of a gas.