Matter and Energy

Classification of Matter

Matter is anything that has mass and occupies space. It can be classified based on its composition into pure substances and mixtures.

• Pure Substances: Have a fixed or definite composition. They are further divided into:

  • Elements: Composed of only one type of atom (e.g., copper Cu, lead Pb, aluminum Al).

  • Compounds: Composed of two or more elements chemically combined in a definite ratio (e.g., water H2O, table salt NaCl, hydrogen peroxide).

• Mixtures: Contain two or more substances physically mixed but not chemically combined. Mixtures can be separated by physical methods and may have variable composition.

  • Homogeneous Mixtures: Uniform composition throughout; different parts are not visible (e.g., brass, air, solutions).

  • Heterogeneous Mixtures: Composition varies; different parts are visible (e.g., sand and water, peach pie).

Example: Air is a homogeneous mixture, while pasta and tomato sauce is a heterogeneous mixture.

States and Properties of Matter

Matter exists in three primary states: solid, liquid, and gas. Each state has distinct physical properties.

• Solids: Definite shape and volume; particles are closely packed in a fixed arrangement and move slowly.

• Liquids: Indefinite shape but definite volume; particles are close together but mobile and move slowly.

• Gases: Indefinite shape and volume; particles are far apart and move rapidly.

Physical Properties: Characteristics observed or measured without changing the identity of a substance (e.g., shape, state, boiling point, density, color).

Chemical Properties: Describe the ability of a substance to change into a new substance (e.g., flammability, reactivity).

Physical Change: Alters the state or shape but not the identity (e.g., melting, boiling, cutting).

Chemical Change: Produces new substances with different properties (e.g., burning, rusting, caramelizing sugar).

Comparison of States of Matter

Summary of Physical and Chemical Properties and Changes

Temperature

Temperature measures how hot or cold an object is. It is measured in degrees using thermometers and can be expressed in different scales.

• Celsius (°C): Water freezes at 0°C and boils at 100°C.

• Fahrenheit (°F): Water freezes at 32°F and boils at 212°F.

• Kelvin (K): Absolute zero is 0 K; water freezes at 273 K and boils at 373 K. No negative temperatures.

Conversion Formulas:

• From Celsius to Fahrenheit:

• From Fahrenheit to Celsius:

• From Celsius to Kelvin:

Example: Normal body temperature (37°C) in kelvins:

Energy

Energy is the ability to do work and makes objects move. It exists in different forms, including potential and kinetic energy.

• Potential Energy: Stored energy due to position or chemical composition (e.g., water at the top of a dam, gasoline).

• Kinetic Energy: Energy of motion (e.g., rollerblading, mowing the lawn).

Heat: Energy associated with the movement of particles. The faster the particles move, the greater the heat.

Units of Energy

• Joule (J): SI unit of energy.

• Calorie (cal): Amount of energy needed to raise the temperature of 1 g of water by 1°C.

• Kilojoule (kJ): 1 kJ = 1000 J.

• Kilocalorie (kcal): 1 kcal = 1000 cal.

Conversion:

Energy and Nutrition

Food energy is measured in Calories (Cal), kilojoules (kJ), or kilocalories (kcal). Nutritionists use calorimeters to measure the energy content of food.

• Calorimeter: Device used to measure heat transfer during combustion of food.

• Energy Values: Typical energy values per gram for food types:

Example: Whole milk with 13 g carbohydrate, 9.0 g fat, 9.0 g protein:

• Carbohydrate:

• Fat:

• Protein:

• Total: (rounded to 170 kcal)

Specific Heat

Specific heat is the amount of heat required to raise the temperature of 1 g of a substance by 1°C. It varies for different substances.

• Formula:

• Where = heat (J), = mass (g), = specific heat (J/g°C), = temperature change (°C)

Example: If 24.8 g of metal absorbs 275 J and temperature rises by 5°C, specific heat is:

Changes of State

Matter changes state when it is converted from one form to another (solid, liquid, gas). These changes involve energy transfer.

• Melting: Solid to liquid at melting point.

• Freezing: Liquid to solid at freezing point.

• Sublimation: Solid to gas without becoming liquid.

• Deposition: Gas to solid without becoming liquid.

• Evaporation: Liquid to gas at surface.

• Boiling: Liquid to gas throughout the liquid.

• Condensation: Gas to liquid.

Heat of Fusion and Heat of Vaporization

• Heat of Fusion: Amount of heat needed to melt 1 g of solid or released when 1 g of liquid freezes.

• Heat of Vaporization: Amount of heat needed to vaporize 1 g of liquid or released when 1 g of gas condenses.

Example: To melt 32.0 g of ice:

Heating and Cooling Curves

Heating curves show temperature changes and changes of state as heat is added. Plateaus represent changes of state at constant temperature; sloped lines represent temperature changes within a state.

• Plateau: Constant temperature, change of state.

• Slope: Temperature change, no change of state.

Combined Heat Calculations

To calculate total heat for a substance undergoing multiple changes (e.g., heating, melting, vaporizing), sum the heat for each step:

• Heat to warm:

• Heat to melt:

• Heat to vaporize:

• Total heat:

Example: To convert 15.0 g of liquid ethanol at 25°C to gas at its boiling point:

• Calculate heat to warm ethanol to boiling point.

• Calculate heat to vaporize ethanol.

• Add both to get total energy needed.

Concept Map

A concept map summarizes the relationships between matter, energy, states, properties, and changes.

Additional info: Some tables and diagrams referenced in the original notes were not fully available; key content has been logically reconstructed and expanded for completeness.