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Chemistry: Unit 1SJB

Chemistry: Phases in Detail & Heat/ Heat Equations

Phases in Detail: Gases - molecules spread out and fill spaces they are given - weak intermolecular forces of attraction - large spaces between particles
Evaporation takes place at all temperature on the liquid/ vapor boundary -increases with an increase in temperature
Vapor definition: a gaseous phase of a substance that is a liquid or sold at normal conditions - once liquid particles have absorbed enough energy to overcome attractive forces, they become this
Vapor Pressure gas particles exert pressure on the liquid when they evaporate - In closed system, pressure increases (ex: lid on top of boiling water; nothing in/out) -increases with an increase in temperature
Reference Table H #1 -101.3 kPa= standard pressure (Table A) -1 atm, 760 turr, 760mmHg are values equivalent to Standard Pressure
Reference Table H #2 -Normal boiling point: when the vapor pressure is equal to atmospheric pressure -When a substance boils, evaporation occurs throughout the liquid -Measures the strength of intermolecular forces between molecules -When a substance
Intermolecular forces between molecules 1) If the vapor pressure is high- the attraction between molecules are weak (propanone) 2) If the vapor pressure is low- the attraction between molecules are strong (ethanoic acid)
4 main types of energy 1) Kinetic Energy (KE): Energy in Motion 2) Potential Energy (PE): Stored Energy 3) Chemical Energy (Energy assorted with a chemical change 4) Heat Energy
Heat Energy Amount of energy transferred from one substance to another - can be measures using a calorimeter - calories (cal) or joules (J) measure heat gain or loss - to convert from cal to J to Kcal or KJ: divide by 1,000
Law of Conservation of Energy - energy is neither creates or destroyed 1. energy can be transferred from one substance to another 2. Or energy can be transferred into a new form of energy 3. the TOTAL AMOUNT OF ENERGY WILL REMAIN THE SAME
(Thermometry): Temperature the measure of the average kinetic energy of the particles of a substance
(Thermometry): Heat 1. flows spontaneously from a hot object to cold object ex: body heat to chair, boiling water to hand, burned hand to icepack
(Temperature Scales): Degrees Celcius (°C) -most commonly used scale - temps below zero are negative -2 fixed points - 0°C: melting/ freezing point of water - 100°C: boiling/ condensation of water - values on a thermometer increase by 1°C
(Temperature Scales): Kelvin #1 -contains theoretically the lowest possible temperature/ 0 Kelvin (absolute zero) - has never been exactly reached (1/1000 K) - Absence of all kinetic energy (no motion of particles) - scale is the same as Celcius, just shifted by 273
(Temperature Scales): Kelvin #2 -1 degree change in Celcius is the same as 1 degree change in Kelvin -Table T: K= °C + 273
Thermometers - device used to measure the average kinetic energy of particles and temp - uses liquids like Hg (Mercury) and colored alcohols that expand at high temperature
Measurement of Heat Energy the amount of heat given off or absorbed in a reaction
(Table T): q=mc▲T #1 q= heat (in Joules or calories) m= mass of substance (g) c= specific heat capacity of substance (J/g°C) ▲T= final temp-initial temp (in °C)
(Table T): q=mc▲T #2 *ONLY USED when there is a CHANGE IN TEMP 1. the amount of heat needed to raise 1.0g of a substance by 1.0°C 2. the value for specific heat of water is on Table B
(Table T): Heat of Fusion the amount of heat needed to MELT 1.0g of a substance *q=mHf is used when calculating how much heat is absorbed when a substance MELTS or FREEZES (Table B) **(NO CHANGE IN TEMPERATURE)**
(Table T): Heat of Vaporization the amount of heat needed to VAPORIZE(boil) 1.0g of a substance *q=mHv is used to calculate how much energy absorbed when a substance BOILS or how much energy is released when a substance CONDENSES **(NO CHANGE IN TEMPERATURE)**
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