Chapter 3: Matter and Energy

3.1: Classification of Matter

• Matter: Anything that has mass and occupies space.
• Pure Substances: A matter that has a fixed or definite composition.
• Element: The simplest type of pure substance. It is composed of only one type of material.
• Atoms: These are extremely tiny particles that make up each type of matter.
• Compound: A pure substance that consists of atoms of two or more elements, always chemically combined in the same proportion.
• Bonds: Happens when atoms are held together by attractions.
• Molecules: Small groups of atoms.
• Pure substances that are compounds can be broken down by chemical processes into their elements. They cannot be broken down through physical methods.
• In a mixture, two or more different substances are physically mixed, but not chemically combined.
• Physical processes can be used to separate mixtures because there are no chemical interactions between the components.
• Homogenous Mixture: Also called a solution, the composition is uniform throughout the sample.
• Heterogeneous Mixture: The components do not have a uniform composition throughout the sample.
• Filtration: It helps in separating solids from liquids, which involves pouring a mixture through a filter paper set in a funnel.
• Chromatography: Different components of a liquid mixture separate as they move at different rates up the surface of a piece of chromatography paper.

3.2: States and Properties of Matter

• State of Matter: The physical forms of matter.
• Solid: It has a definite shape and volume.
  • Strong attractive forces hold the particles such as atoms or molecules close together.
  • The particles here are arranged in such a rigid pattern, their only movement is to vibrate slowly in fixed positions.
• Liquid: It has a definite volume, but not a definite shape.
  • The particles here move in random directions but are sufficiently attracted to each other to maintain a definite volume, although not a rigid structure.
• Gas: It does not have a definite shape or volume.
  • The particles here are far apart, have little attraction to each other, and move at high speeds, taking the shape and volume of their container.
• Physical properties
  • These are those characteristics that can be observed or measured without affecting the identity of a substance.
  • When matter undergoes a physical change, its state, size, or appearance will change, but its composition remains the same.
  • Examples:
  • Water boils to form water vapor.
  • Sugar dissolves in water to form a solution.
  • Copper is drawn into thin copper wires.
  • Paper is cut into tiny pieces of confetti.
  • Pepper is ground into flakes.
• Chemical properties
  • These are those that describe the ability of a substance to change into a new substance.
  • When a chemical change takes place, the original substance is converted into one or more new substances, which have different physical and chemical properties.
  • Examples:
  • Shiny, silver metal reacts in air to give a black, grainy coating.
  • A piece of wood burns with a bright flame, and produces heat, ashes, carbon dioxide, and water vapor.
  • Heating white, granular sugar forms a smooth, caramel-colored substance.
  • Iron, which is gray and shiny, combines with oxygen to form orange-red rust.

3.3: Temperature

• Temperatures in science, are measured and reported in Celsius (°C) units.
  • The freezing point of water, is defined as 0 °C, and the boiling point is 100 °C.
• In the United States, everyday temperatures are commonly reported in Fahrenheit (°F) units.
  • Water freezes at exactly 32 °F and boils at 212 °F.
• A typical room temperature of 22 °C would be the same as 72 °F.
• Normal human body temperature is 37.0 °C, which is the same temperature as 98.6 °C.
• Degrees: Smaller units of temperature.
• Scientists have learned that the coldest temperature possible is 273 °C.
• On the Kelvin scale, 273 °C temperature, which is called absolute zero, and has the value of 0 K.
• Kelvins (K): The unit of the Kelvin Scale, no degree symbol is used.

3.4: Energy

• Energy: The ability to do work.
• Kinetic Energy: The energy of motion.
• Potential Energy: Determined by the position of an object or by the chemical composition of a substance.
• Heat: The energy associated with the motion of particles.
• Joule (J): The SI unit of energy and work.
  • It is a small amount of energy and scientists often use the kilojoule (kJ).
• Calorie (cal): Defined as the amount of energy needed to raise the temperature of 1 g of water by 1 °C.

3.5: Energy and Nutrition

• Carbohydrates are the primary fuel for the body.
  • If the carbohydrate reserves are exhausted, fats and then proteins are used for energy.
• In the nutrition laboratory, foods are burned in a calorimeter to determine their energy value.
• The energy values for food are the kilocalories or kilojoules obtained from burning 1 g of carbohydrate, fat, or protein.

3.6: Specific Heat

• Specific Heat

  • The amount of heat needed to raise the temperature of exactly 1 g of a substance by exactly 1 °C.
  • This temperature change is written as ∆T (delta T), where the delta symbol means “change in.”

  

• The high specific heat of the water has a major impact on the temperatures in a coastal city compared to an inland city.

• A large mass of water near a coastal city can absorb or release five times the energy absorbed or released by the same mass of rock near an inland city.

• Heat Equation: Specific heat expression that is arranged to solve for heat.

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3.7: Changes of State

• Change of State: It occurs when the matter is converted from one state to another.

• Melting Point: The particles of a solid gain sufficient energy to overcome the attractive forces that hold them together.

  • The particles in the solid separate and move about in random patterns.
  • The substance is melting, changing from a solid to a liquid.

• Freezing Point: This occurs when a liquid changes to a solid.

  • If the temperature of a liquid is lowered, the reverse process takes place.
  • Kinetic energy is lost, the particles slow down, and attractive forces pull the particles close together; therefore, the substance is freezing.

• Heat Fusion: The energy that must be added to convert exactly 1 g of solid to liquid at the melting point.

  • The heat of fusion is also the quantity of heat that must be removed to freeze exactly 1 g of water at its freezing point.

• Evaporation: It is taking place as water molecules with sufficient energy escape from the liquid surface and enter the gas phase.

• Boiling Point (bp): The molecules within a liquid have enough energy to overcome their attractive forces and become gas.

• Condensation: The water vapor is converted back to liquid as the water molecules lose kinetic energy and slow down.

• Sublimation: The particles on the surface of a solid change directly to a gas with no temperature change and without going through the liquid state.

• Deposition: The reverse of sublimation; the gas particles change directly to a solid.

  • Freezer Burn: This occurs when a solid is left in the freezer for a long time, and so much water sublimes that solids become dry and shrunken.

• Heat of Vaporization: The energy that must be added to convert exactly 1 g of liquid to gas at its boiling point.

• On a heating curve or cooling curve, the temperature is shown on the vertical axis and the loss or gain of heat is shown on the horizontal axis.

  • Cooling Curve: A diagram of the cooling process in which the temperature decreases as heat is removed.

    

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