Matter

Solid State

Solids are a state of matter characterised by closely packed molecules arranged in a regular pattern. This close packing results in solids maintaining a fixed shape, mass, and volume. The strong intermolecular forces between particles prevent them from easily moving past one another, leading to the solid’s rigidity.

Fig 1.13: Closely packed molecules in solids Liquid State Liquids are a state of matter where molecules are relatively close together but not as tightly packed as in solids. This moderate spacing allows liquids to take the shape of their containing vessel, conforming to its contours. While liquids have a fixed volume, they lack a fixed shape, meaning they can flow and change shape depending on their container. Compared to gases, liquids are less compressible due to the intermolecular forces that resist compression.

Fig 1.14: Molecules in a liquid Gaseous State Gases are a state of matter where molecules are spaced far apart and move freely in all directions. These molecules have high kinetic energy, resulting in constant and random motion. Gases have neither a fixed shape nor a fixed volume, as they expand to fill the entirety of their container. The weak intermolecular forces between gas particles allow for easy compression, meaning gases can be compressed into smaller volumes under pressure.

Fig. 1.15: Molecules in a gas Plasma Plasma is a unique form of matter, often forgotten, but it’s actually the fourth state of matter, after solids, liquids, and gases. Just like how water turns into steam when heated, gases can turn into plasma when they get really hot. Plasma is like a mixture made of positively charged particles called ions and negatively charged particles called electrons. It’s full of energy because the electrons are stripped away from their atoms. This makes plasma able to conduct electricity and react to magnetic fields.

Fig 1.16: Plasma state You can find plasma in places like stars, lightning, and fluorescent lights, where there is a lot of heat or energy.

Table 1.8: Various states of matter Solid Liquid Gas Plasma

• Cylinder

• Chalk

• Conical flask

• Cutlass

• Bar magnet

• Wood

• Tree

• Gold

• Stone

• Block

• Book

• Water

• Kerosene

• Petrol

• Diesel

• Turpentine

• Alcohol

• Ethanol

• Beer

• Pito

• Water vapour

• Hydrogen gas

• Natural gas

• Ethane

• Oxygen

• Ozone

• Hydrogen sulphide

• Air

• Helium

• Nitrogen

• Freon

• Carbon dioxide

• Stars,

• Neon signs,

• Lightning

Activity 1.25 Exploration states of matter Materials needed: A list of common objects. Stone, water, gas, lightening, kerosene, helium gas, book, ice cube, water, helium balloon, wooden block, stars, pencil lead, etc What to do:

1. Pair up with a friend and list the three primary states of matter.

2. In pair or individually classify each object into one of the three states of matter.

Liquid Solid Gas

3. Share your classifications with your group of friends and explain your classifications.

4. Compare your classifications with your friends and identify any differences.

Activity 1.26 Exploring the fourth state of matter - Plasma Materials needed: Tablets for research, Poster paper, markers, and other art supplies (optional) What to do:

1. Check on the internet.

2. Search plasma states.

3. Draw your observation on your poster paper.

4. Share your results (10 minutes)

Activity 1.27 Melting ice cubes to demonstrate change of state from solid to liquid Material needed: Ice cubes What to do:

1. Take an ice cube.

2. Notice its coldness and solid state.

3. Place the ice cube in a bowl or cup.

4. Watch closely as the ice cube starts to melt.

5. Observe any changes in its appearance and texture as it turns into water.

6. Optionally, you can use a timer to see how long it takes for the ice cube to melt completely.

7. Reflect on what you observed during the melting process.

8. Think about other examples of solids turning into liquids in everyday life.

9. Share your observations with your friends.

Ice (solid) Water (liquid)

Fig. 1.17

10. Observe and record your observations on the following questions:

a. How does the rate of melting ice cubes vary when placed in different environments (e.g. at room temperature versus in direct sunlight)

b. What factors do you think contribute to these differences?

c. How does the size or shape of ice cubes affect the rate at which they melt, comparing large and small ice cubes or cubes with different shapes under similar conditions?

Activity 1.28 Evaporating Water - Liquid to gas Materials needed:

• water

• pot or kettle

• stove or heat source What to do:

1. Pour some water into the pot or kettle.

2. Place the pot or kettle on the stove or heat source.

3. Turn on the stove and wait for the water to heat up.

4. As the water heats up, observe closely what happens to it.

5. You will start to see steam rising from the surface of the water.

6. This steam is actually water vapour, which is the gaseous form of water.

7. Keep watching as more steam is produced, indicating that the water is evaporating into the air.

8. Share your observations and answers with your friends.

Water Water heated to evaporate Steam (gas) Fig. 1.18

9. Investigate and share your observations and answers with your friend or teacher:

a. How does the rate of water evaporation vary with different temperatures, and what factors influence this rate?

b. How does the surface area affect the rate of water evaporation, comparing the evaporation rates of shallow versus deep containers of water under similar conditions?

Activity 1.29 A simple model to demonstrate how gas changes to plasma.

Materials needed: Balloons, an inflation pump, gas (air, helium), a microwave safe plate or tray, a microwave oven, a timer or stopwatch, pen/pencil and notebook for recording What to do:

1. Inflate a balloon with gas (helium or air).

2. Tie the balloon securely to prevent the gas from escaping.

2. Place the balloon on the microwave-safe plate or tray.

3. Carefully place the plate with the balloon inside the microwave oven.

4. Close the microwave oven door and set the timer for a short duration (e.g., 10-15 seconds).

5. Turn on the microwave oven and observe.

Inside the microwave oven Ionised gas (plasma) Fig. 1.19

6. Share your observations with your friend, teacher or family member.

7. Observation :

The microwave oven generates electromagnetic waves at a specific frequency which interact with the gas molecules inside the balloon to gain energy and move more rapidly as collisions between them become more frequent and intense. This leads to some gas molecules losing their electrons and become ionized. As more and more gas molecules become ionized, the gas transitions into a plasma state.

Characteristics of the State of Matter

a. Solid: Has a definite shape and volume, particles are closely packed and vibrate in fixed positions.

b. Liquid: Has a definite volume but takes the shape of its container, particles are close together but can move and flow.

c. Gas: Has neither a definite shape nor volume, particles are far apart and move freely.

d. Plasma: Similar to a gas but consists of ionized particles, conducts electricity, and is influenced by magnetic fields.

Particle Behaviour in the State of Matter

a. Solid: Particles are tightly packed and vibrate in fixed positions.

b. Liquid: Particles are close together but can move and flow past each other.

c. Gas: Particles are far apart and move randomly at high speeds.

d. Plasma: Particles are highly ionized and move freely, influenced by electric and magnetic fields.

Real-life examples and Applications of the States of Matter

a. Cooking: Understanding the states of matter helps in determining cooking times and techniques for various ingredients.

b. Weather forecasting: Knowledge of states of matter helps in predicting changes in weather patterns, such as precipitation and cloud formation.

c. Pharmaceutical industry: Understanding states of matter is crucial for drug formulation, as it affects drug solubility and stability

Activity 1.30 Molecular arrangement in different states of matter Molecular model kits (or digital simulations) What to do:

1. Organise yourselves into small groups of 3-4 learners.

2. Using a molecular model kit (or access to digital simulations) and a worksheet observe and record the following:

a. How are the molecules arranged in a solid substance?

b. How do the molecule-molecule interactions differ between the solid, liquid, and gas phases?

c. What are the key differences in the energy of the molecules in each state of matter?

d. How does the spacing between molecules change from solid to liquid to gas?

e. How do the shapes and movements of the molecules differ in each state of matter?

Activity 1.31 Modelling the states of matter What to do:

1. Put yourselves into smaller groups of 4-5 students.

2. You will be modelling the behaviour of particles in the three states of matter: solid, liquid, and gas.

3. Stand up and physically represent the particles in a solid state. Stand close together, vibrating slightly in place, with limited movement.

Fig. 1.20: Human model to represent the solid state

4. Repeat the process with the next group, this time modelling the particles in a liquid state. The students should move more freely but still remain somewhat close together, with the ability to flow and change shape.

Fig. 1.21

5. Repeat the process with the next group, this time modelling the particles in a gaseous state. The students should move more freely apart.

6. How has this activity helped deepen your understanding of the behaviour of particles in solids, liquids, and gases?

7. Reflect on the benefits of using the modelling and physical representations in the learning process.

8. Comparison of the molecular force, molecular motion, and molecular attraction in the four states of matter: solids, liquids, gases and plasmas.

Molecular Force: In solids, the intermolecular forces are strong. These forces can include ionic bonds, covalent bonds, metallic bonds, or intermolecular forces such as hydrogen bonding, van der Waals forces, and dipole-dipole interactions.

The strong forces keep the molecules tightly packed and maintain the solid’s rigid structure.

Molecular Motion: In solids, the molecules have limited freedom of movement.

They vibrate around fixed positions, but their overall motion is restricted. The molecules oscillate with small amplitudes around their equilibrium positions.

Molecular Attraction: The strong intermolecular forces result in significant molecular attractions. The molecules are held together in a fixed arrangement, maintaining the solid’s shape and volume. The attractions between molecules give rise to properties such as high density, defined shape, and resistance to compression.

Molecular Force: In liquids, the intermolecular forces are weaker compared to solids. The forces can still include hydrogen bonding, van der Waals forces, and dipole-dipole interactions, but they are not as strong as in solids. The forces are sufficient to keep the molecules in close proximity to each other.

Molecular Motion: In liquids, the molecules have greater freedom of movement compared to solids. They can slide past each other and move more randomly.

The molecules have more kinetic energy, leading to increased molecular motion compared to solids.

Molecular Attraction: The weaker intermolecular forces in liquids result in less rigid molecular attractions. The molecules are attracted to each other but are still able to move and flow past one another. Liquids exhibit properties such as moderate density, ability to flow, and moderate compressibility.

Fig 1.22: The arrangement of molecules in liquid GASES Molecular Force: In gases, the intermolecular forces are very weak compared to solids and liquids. The forces are primarily van der Waals forces and weak dispersion forces. The forces are not strong enough to hold the molecules together closely.

Molecular Motion: In gases, the molecules have high kinetic energy. They move freely and independently, colliding with each other and the container walls. The molecules have a high degree of random motion and travel in straight lines until they collide with other molecules or container boundaries.

Molecular Attraction: The weak intermolecular forces in gases result in negligible molecular attractions. The molecules are not tightly bound to each other, and they can move and expand to fill the entire container. Gases exhibit properties such as high compressibility, low density, and the ability to diffuse and mix rapidly.

Fig 1.23:The arrangement of molecules in gas PLASMAS:

Molecular Force: In plasmas, the molecular forces are significantly weakened or overcome due to high energy and ionization. The typical intermolecular forces found in condensed matter, such as van der Waals forces and hydrogen bonding, are generally negligible. Plasmas are primarily governed by electromagnetic forces. In plasma, molecules are not present but rather a mixture of ions and free electrons. It does not have a fixed molecular arrangement. The charged particles in plasma move independently and do not maintain fixed positions Fig 1.24: The arrangement of molecules in plasma Molecular Motion: In plasmas, the charged particles, such as ions and free electrons, have high kinetic energy and move rapidly. They can travel long distances before colliding with other particles. The motion of particles in a plasma is highly influenced by electric and magnetic fields.

Molecular Attraction: Plasmas are characterized by the interactions between charged particles. These attractions and interactions are dominated by the Coulombic forces between charged particles. Electric fields and magnetic fields play significant roles in shaping the behaviour of plasmas. The interactions in plasmas are primarily electromagnetic rather than molecular.

Heating and cooling can have various effects on different states of matter, including solid, liquid, gas, and plasma. Let us explore the effects individually:

Fig 1.25: Summary of the molecular arrangement of various states of matter

1. Solid:

a. Heating: When you heat a solid, it gets warmer, and its molecules start vibrating faster. This extra energy makes the solid expand, getting a bit bigger. Eventually, if you heat it enough, it reaches its melting point. At this point, the forces holding the solid together weaken, and it turns into a liquid.

b. Cooling: Cooling a solid reduces its temperature, causing the atoms or molecules to vibrate less. As the thermal energy decreases, the solid contracts and its size decreases. Cooling can lead to the solid reaching its freezing point, where it transitions from a liquid to a solid state.

2. Liquid:

a. Heating: When a liquid is heated, its temperature increases, causing the average kinetic energy of its molecules to rise. The increased kinetic energy leads to faster molecular motion, resulting in the liquid expanding.

As the temperature continues to rise, the liquid may eventually reach its boiling point, the liquid changes into a gas

b. Cooling: Cooling a liquid decreases its temperature and reduces the average molecular motion. As the thermal energy decreases, the liquid contracts and its volume decreases. Cooling can cause the liquid to reach its freezing point, leading to the formation of a solid.

3. Gas:

a. Heating: When a gas is heated, its temperature increases, which leads to an increase in the average kinetic energy of its molecules. The increased kinetic energy causes the gas molecules to move more rapidly and with greater force, resulting in an expansion of the gas volume. Further heating can eventually lead to the gas reaching its critical temperature, where it transitions into a supercritical fluid or undergoes a phase change into a plasma.

b. Cooling: Cooling a gas decreases its temperature, causing the average kinetic energy of its molecules to decrease. The reduced kinetic energy leads to slower molecular motion and a decrease in volume. Cooling can cause the gas to reach its condensation point, where it transitions into a liquid state.

4. Plasma:

a. Heating: Plasma is an ionised gas consisting of positively charged ions and negatively charged electrons. Heating a gas to extremely high temperatures causes the atoms to lose electrons, resulting in the formation of a plasma. The added thermal energy provides enough energy for the electrons to break free from their atomic orbits, creating a mixture of free electrons and ions. This process, called ionisation, is what distinguishes plasma from ordinary gases. Plasma conducts electricity and responds to magnetic fields due to its charged particles.

b. Cooling: When plasma cools down, the free electrons begin to recombine with the positively charged ions to form neutral atoms or molecules. As the temperature decreases, the energy available for ionisation reduces, leading to the plasma reverting to a neutral gas state. The recombination of electrons with ions releases energy in the form of light or heat, which can be observed as the plasma transitions back to its non-ionised state Annex 1.1 A: Additional Reading: Density Density is how much substance is packed into an object. It depends on both its mass (how heavy it is) and its volume (how much space it takes up). Density doesn’t change, no matter how big or small something is. Different materials have different densities. When things get hotter, they usually become less dense, and when they’re squished, they become denser. If something is denser than water, it sinks, but if it’s less dense, it floats. In the density column experiment, we layer liquids with different densities, and they separate based on their weight. Ice floats on water because it’s less dense than liquid water, which happens because ice’s molecules spread out more when they freeze.”

Density is defined as the mass per unit volume of a substance at a particular temperature Mathematically:

Density, ρ = Mass, m________ Volume, V

Example

1. A piece of an iron has a volume 15 cm³and a mass of 27 g. calculate the density of the iron in

(a) g/cm³(b)kg/m³Solution v = 15 cm³, m = 27 g ρ = mass______ volume = 27 g/15 cm³= 1.8 gcm

(a) 15 cm³= 15 × 1 m³__________ 1000000 cm³= 0.00015 m³ρ = mass______ volume = 0.0 27 kg/0.00015 m³= 1800 kgm⁻³ Alternatively M = 27 g, v= 15 cm³, 10⁶= 1000000 ρ = mass______ volume = 27____ 1000______ 15/1000000 = 27 x 10³_______ 15 x 10⁶= 1800 kgm⁻³B: Additional Learning on Prefixes for Scientific Notations Prefixes A unit prefix is a specifier or reminder that is attached to units of measurement to indicate multiples or fractions of the units. The use of such prefixes commonly forms units of various sizes. Eg 5μF = 5 × 10⁻⁶F , 5mH = 5 × 10⁻³H. These convections are used in other topics such as alternate current, electrostatics, etc. You need to use it (apply it) without nobody telling you.

The table below shows the multiples and submultiples and their corresponding symbols and magnitudes.

Table 1.9: Prefixes and their multipliers Multiple Prefix Symbol Example 10³Kilo k kilogram, kg(=10³g) 10⁶mega M megavolt, MV(=10⁶V) 10⁹giga G Gigavolt, GV(=10⁹V) 10⁻²centi c Centimetre, cm(=10⁻²m) 10⁻³cilli m Millimetre, mm(=10⁻³m) 10⁻⁶micro μ Micrometre, μm(=10⁻⁶m) 10⁻⁹nano n Nanometre, nm(=10⁻⁹m) 10⁻¹²pico p Picofarad, pF(=10⁻¹²F) 10⁻¹⁵femto f femtometre, fm (=10⁻¹⁵m) 10⁻¹⁸atto a Attometre, am(=10⁻¹⁸m) Submultiples Prefix Symbol 10 ⁻¹Deci D 10 ⁻²Centi C 10 ⁻³Milli M Factor Prefix name Symbol 10²⁴= (10³)⁸Yotta Y 10²¹= (10³)⁷zetta Z 10¹⁸= (10³)⁶exa E 10¹⁵= (10³)⁵peta P 10¹²= (10³)⁴Tera T 10⁹= (10³)³giga G 10⁶= (10³)²mega M 10³= (10³)¹kilo k 10²hector h 10¹deka da

Activity: Exploring Metric Prefixes

Objective: This activity aims to familiarise you with metric prefixes and their corresponding powers of 10. It will also help you understand the relationship between different metric units and their relative size.

Materials needed:

• Pen or pencil

• Paper or notebook What to do:

1. Start by creating a table with three columns. Label the columns: Prefix, Symbol, and Power of 10.

2. Begin with the base unit, which is the meter (m). Write “meter” in the Prefix column, “m” in the Symbol column, and “1” in the Power of 10 column.

3. Research and list the most commonly used metric prefixes. Include prefixes such as kilo-, centi-, milli-, micro-, and nano-. You can also include less commonly used prefixes if you like.

4. In the Symbol column, write down the symbol or abbreviation for each prefix. For example, “k” for kilo-, “c” for centi-, “m” for milli-, “μ” (mu) for micro-, and “n” for nano-.

5. Determine the power of 10 associated with each prefix. For example, kilo- represents 10³, centi- represents 10⁻², milli- represents 10⁻³, micro- represents 10⁻⁶, and nano- represents 10⁻⁹. Write down the corresponding powers of 10 in the Power of 10 column.

6. Complete the table by filling in the remaining rows with the prefixes, symbols, and powers of 10 you have researched.

7. Once you have completed the table, examine the prefixes and their associated powers of 8. Notice how each prefix represents a specific multiplication or division by powers of 10. For example, kilo- represents multiplication by 10³, while milli- represents division by 10³.

9. Use the table to convert between different metric units. For instance, if you have a distance of 500 millimetres, you can use the table to determine that it is equivalent to 0.5 meters.

10. Create conversion problems for yourself or others using the metric prefixes and their corresponding powers of 10. For example, convert 2.5 kilometres to meters or convert 75 milligrams to grams.

11. Share your findings and experiences with a partner or discuss them in a group. Compare different approaches and strategies used during the

activity.