Understanding Dissolving

Dissolving

• Dissolving is the process where a solute mixes evenly with a solvent to form a solution.
• The solute can be a solid, liquid, or gas.

Key Terms

• Solute: The substance that dissolves (e.g., salt, sugar).
• Solvent: The substance in which the solute dissolves (e.g., water).
• Solution: The uniform mixture of solute and solvent.
• Solubility: The maximum amount of solute that can dissolve in a solvent at a given temperature.

The Process of Dissolving

• Solvent particles surround solute particles via a process called solvation.
• Solute particles break apart and mix with solvent molecules, forming a uniform solution where solute particles are evenly distributed.

Factors Affecting Dissolving

• Temperature: Higher temperatures increase the dissolving rate for solids, but gases dissolve less in hot solvents.
• Stirring: Increases the dissolving rate by moving solvent molecules around, bringing them into contact with the solute faster.
• Surface Area: Smaller solute particles dissolve faster than larger ones.
• Nature of Solvent & Solute: Polar solvents dissolve polar solutes, and non-polar solvents dissolve non-polar solutes ("like dissolves like").

Types of Solutions

• Dilute Solution: Contains a small amount of solute dissolved.
• Concentrated Solution: Contains a large amount of solute dissolved.
• Saturated Solution: No more solute can dissolve at a given temperature.
• Unsaturated Solution: More solute can still dissolve.
• Supersaturated Solution: Contains more solute than normally possible and is unstable.

What is a Solution?

• A solution is a homogeneous mixture of two or more substances, with a solute dissolved in a solvent, where particles are evenly spread.

Key Terms Related to Solutions

• Solute: The substance that dissolves (e.g., salt, sugar).
• Solvent: The substance that dissolves the solute (e.g., water).
• Solution: The uniform mixture of solute and solvent.
• Solubility: The maximum amount of solute that can dissolve in a given amount of solvent at a specific temperature.
• Solubility Curve: A graph showing how solubility changes with temperature.

Factors Affecting Solubility

• Temperature: Solubility increases with temperature for solids in liquids, but decreases with temperature for gases in liquids.
• Pressure (for gases only): Higher pressure increases gas solubility in liquids.
• Nature of the Solvent & Solute: Polar solutes dissolve in polar solvents, and non-polar solutes dissolve in non-polar solvents ("like dissolves like").

Types of Solutions Based on Solubility

• Dilute Solution: A small amount of solute dissolved which is like a teaspoon of salt in a glass of water.
• Concentrated Solution: A large amount of solute dissolved which is like sugar syrup for making sweets.
• Saturated Solution: Holds the maximum solute at a given temperature which is like saltwater with no more dissolving salt.
• Unsaturated Solution: More solute can dissolve which is like weak tea where more sugar can be added.
• Supersaturated Solution: More solute dissolved than normally possible (unstable) such as crystallization of sugar from syrup.

Solubility Curve

• A solubility curve shows how the solubility of a substance changes with temperature.
• Upward curve denotes solid solutes dissolve more at higher temperatures.
• Downward curve denotes gaseous solutes dissolve less at higher temperatures.

What is Paper Chromatography?

• Paper chromatography separates mixtures of soluble substances based on differences in solubility and attraction to the paper.

Key Terms for Chromatography

• Chromatography: A method used to separate components of a mixture.
• Mobile Phase: The solvent that moves up the paper (e.g., water or ethanol).
• Stationary Phase: The absorbent material (e.g., chromatography paper).
• Solubility: The ability of a substance to dissolve in a solvent.
• Rf Value: The ratio of the distance traveled by the solute to the distance traveled by the solvent.

How Paper Chromatography Works

• A small spot of the mixture is placed on chromatography paper.
• The paper is dipped in a solvent, and the solvent moves up by capillary action.
• Different components separate into distinct spots as they move at different speeds.
• More soluble substances travel farther, while less soluble substances stay closer to the starting point.

Steps in a Chromatography Experiment

• Materials needed include chromatography paper, a beaker, water or ethanol (solvent), an ink or food dye mixture, a pencil, a ruler, and a dropper.
• A baseline is drawn in pencil near the bottom of the paper, and a drop of ink or dye is placed on the baseline.
• The paper is dipped into the solvent, ensuring the ink spot stays above the liquid.
• Wait for the solvent to rise and separate the components.
• Mark the final solvent level and measure distances traveled.

Rf Value Calculation

• The Rf value (Retention Factor) is calculated as the distance traveled by the solute divided by the distance traveled by the solvent.
• Different substances have unique Rf values, aiding in identifying unknown components.

Uses of Paper Chromatography

• Identifying different dyes in ink or food coloring.
• Checking the purity of a substance.
• Forensic science such as analysing ink in documents.
• Drug testing in medical labs.

What are Exothermic Reactions?

• Exothermic reactions are chemical reactions that release energy to the surroundings, usually as heat, light, or sound.

Key Characteristics of Exothermic Reactions

• Energy is released to the surroundings, causing them to become hotter.
• The products have less energy than the reactants.
• Bond formation releases energy.

Examples of Exothermic Reactions

• Combustion: Burning of fuels (e.g., wood, gasoline, natural gas).
• Respiration: The process where glucose is broken down to release energy.
• Neutralization: Acid-base reactions.
• Condensation & Freezing: Phase changes that release heat to surroundings.

Applications of Exothermic Reactions

• Hand warmers which is oxidation of iron.
• Self-heating cans which is a quicklime reaction.
• Fireworks as combustion reactions.
• Power plants that are burning fuel for electricity.

Endothermic Reactions

• Endothermic reactions absorb energy from their surroundings.
• The products have more energy than the reactants, causing the surroundings to become cooler as heat is absorbed.

Examples of Endothermic Reactions

• Photosynthesis: Plants absorb sunlight to convert carbon dioxide and water into glucose and oxygen.
• Melting Ice: Heat is absorbed to change ice into water.
• Boiling and Evaporation: Water absorbs heat to turn into steam.
• Thermal Decomposition: Certain compounds decompose when heated with example being calcium carbonate..
• Cooking Food: Heat is absorbed to cook raw ingredients.

Reactions of Metals with Oxygen

• Metals react with oxygen to form metal oxides which is a type of oxidation reaction, and part of a combustion or corrosion process.
• General word equation: Metal + Oxygen --> Metal oxide shown in the example of Magnesium + Oxygen --> Magnesium Oxide.

Properties of Metal Oxides

• Most metal oxides are basic and react with acids to form salt and water such as Magnesium Oxide + Hydrochloric Acid --> Magnesium Chloride + water.
• Some metal oxides (e.g., aluminum oxide) are amphoteric and can react with both acids and bases.

Reactivity of Metals with Oxygen

• Highly reactive metals like potassium, sodium, and magnesium burn in oxygen, forming metal oxides quickly.
• Less reactive metals like copper and silver tarnish or form a surface layer of oxide slowly.
• Unreactive metals like gold and platinum do not react with oxygen.

Reaction of Metals with Water

• Some metals react with water to form a metal hydroxide and hydrogen gas.
• General word equation: Metal + Water --> Metal Hydroxide + Hydrogen such as Sodium + Water --> Sodium Hydroxide + Hydrogen.

Reactivity of Metals with Water

• Very reactive metals (e.g., potassium, sodium, calcium) react vigorously with water, producing bubbles of hydrogen gas, and potassium ignites with a lilac flame.
• Less reactive metals (e.g., magnesium, zinc, iron) react slowly with cold water but react with steam to form a metal oxide such as Magnesium + Steam --> Magnesium Oxide + Hydrogen.
• Unreactive metals like copper, silver, gold, and platinum do not react with water.

Observations in Reactions with Water

• Bubbling/fizzing due to hydrogen gas production and the formation of an alkaline solution (metal hydroxide).
• In some cases, heat is released.

Reaction of Metals with Acids

• Metals react with acids to form a salt and hydrogen gas.
• General word equation: Metal + Acid → Salt + Hydrogen demonstrated in the examples Zinc + Hydrochloric Acid → Zinc Chloride + Hydrogen, Magnesium + Hydrochloric Acid → Magnesium Chloride + Hydrogen, and Iron + Sulfuric Acid → Iron(II) Sulfate + Hydrogen

Reactivity of Metals with Acids

• Highly reactive metals like potassium and sodium react explosively with acids and are too dangerous.
• Reactive metals like magnesium, zinc, and iron react steadily, producing hydrogen bubbles.
• Unreactive metals like copper, silver, and gold do not react with acids.

Observations in Reactions with Acids

• Effervescence (bubbling) due to hydrogen gas.
• The metal dissolves, forming a salt solution, and heat may be released indicating an exothermic reaction.