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This unit covers the fundamental concepts of energy and the states of matter. It explains how inter-molecular forces (IMF) affect solids, liquids, and gases, and outlines the energy transitions during phase changes such as melting, freezing, vaporization, and condensation. The relationship between energy transfer and chemical reactions is emphasized through the Law of Conservation of Energy. Key formulas for calculating heat capacity and specific heat are provided, enabling deeper comprehension of thermal dynamics. Units like calories and joules are defined for practical application.
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The State of Matter of a substance depends on several things • Attraction between particles called IMF or Inter-Molecular Forces • Solids have very high IMF • Gases have no attraction between molecules • Liquids have IMF a bit lower than solids • Energy of the particles • Space between particles
Phase Changes • Energy added • Melting: Solid to Liquid • Vaporizing: Liquid to Gas • Sublimation: Solid to Gas • Energy removed • Freezing: Liquid to Solid • Condensing: Gas to Liquid • Deposition: Gas to Solid
Heating curve Gas Temperature Vaporizing Liquid Melting Solid Time
Gas cooling curve Condensing Temperature Liquid Freezing Solid Time
Energy - capacity for doing work • weightless, odorless, tasteless • All chemical reactions and changes in state involve either: • release of energy, or • absorption of energy
The Law of Conservation of Energy states that in any chemical or physical process, energy is neither created nor destroyed. • All the energy is accounted for as Eth, Eph or Ech.
Endothermic Reactions • Eth flowing into a system from it’s surroundings: • Q has a positive value • system gains heat (gets warmer) as the surroundings cool down
Exothermic reactions • Heat flowing out of a system into it’s surroundings: • Q has a negative value • system loses heat (gets cooler) as the surroundings heat up
Units • A calorie is defined as the quantity of heat needed to raise the temperature of 1 g of pure water 1 oC. • 1 Calorie = 1 kilocalorie = 1000 cal. • The Joule, the SI unit of heat and energy • 4.184 J = 1 cal
Heat Capacity - the amount of heat needed to increase the temperature of an object exactly 1 oC Depends on both the object’s mass and its chemical composition Specific Heat Capacity (abbreviated “C”) - the amount of heat it takes to raise the temperature of 1 gram of the substance by 1 oC
The higher a material’s specific heat the LONGER it take to heat up and the LONGER it takes to cool down. ex. metals have low specific heat they heat up and cool down quickly.
How to solve heat capacity problems: Q = mc∆T Q = mHv Q = mc∆T Q = mHf Q = mc∆T
Meaning of formula symbols Q = heat or energy unit calories or kilocalories or joules or kilojoules (1 cal = 4.18 joules) m = mass in grams c = specific heat – the amount of energy or heat needed to raise 1 gram of something 1°C, it is different for each substance and phase of matter. J/g°C or cal/g°C ∆T = Change in temperature (Tempfinal – Tempinitial) Hf = Heat of fusion or melting. The amount of heat or energy needed to melt or freeze something. Energy needed for a phase change. Hv = Heat of vaporization or condensation. The amount of heat or energy needed to vaporize or condense something. Energy needed for a phase change.
Energy Constants For water c solid = 2.1 J/g°C or 0.5 cal/g°C 1 kcal = 1000 cal c liquid = 4.18 J/g°C or 1 cal/g°C 1 kJ = 1000 J c vapor = 1.8 J/g°C or 0.44 cal/g°C Hf = 334 J/g or 80 cal/g Hv = 2260J/g or 540 cal/g
