Chapters 13 & 17 Phases and Heat

Phases of Matter Chapter 13

Phases There are three phases, or states, that we will discuss  Solid  Liquid  Gas

Solids form of matter that has a definite shape and definite volume. Use (s) to denote solids in chemical reactions

Solids In most solids the atoms, ions, or molecules are packed tightly together The particles in solids tend to vibrate around fixed points When you heat a solid, its particles vibrate more rapidly, eventually the solid breaks down and melts.

Types of Solids Crystalline Solids  In a crystal the particles are arranged in an orderly, repeating, three-dimensional pattern called a crystal lattice. There are many different shapes of crystalline solids, pg 397

Types of Solids Non-Crystalline Solids  Amorphous solids lack an orderly internal structure.  Ex – Rubber, plastic, and asphalt.  Glass – transparent fusion product of inorganic substances that have cooled to a rigid state without crystallizing. Sometimes called super-cooled liquids.

Liquids form of matter that has a definite volume, indefinite shape, and flows. Use ( l ) to denote liquids in chemical reactions

Liquids In liquids the atoms or molecules are able to slide past each other. In liquids there are intermolecular attractions between the atoms or molecules, which determine the liquid’s physical properties. When you heat a liquid the particles vibrate more rapidly and start moving past each other faster.

Gases form of matter that takes both the shape and volume of its container Use (g) to denote gases in chemical reactions

Phase Changes Six Changes  Solid  LiquidMelting  Liquid  SolidFreezing  Liquid  GasVaporization  Gas  LiquidCondensation  Solid  GasSublimation  Gas  SolidDeposition

Phase Changes During any given phase change, both phases can exist together in equilibrium Example  At 0°C, water can exist in both the liquid and solid phases in equilibrium

Energy When energy is added to a reaction, or phase change, it is called Endothermic When energy is released during a reaction, or phase change, it is called Exothermic

Phase Changes Which phase changes are endothermic, requiring the addition of energy?  Melting  Vaporization  Sublimation

Phase Changes Which phase changes are exothermic, releasing energy?  Freezing  Condensation  Deposition

Phase Diagram of CO 2

Energy What is energy?  Capacity to do work  Ability to do work Two main types  Kinetic  Potential  Many, many more

Types of Energy Kinetic Energy  Energy of motion  Related to the speed and mass of molecules Potential Energy  Stored energy

Temperature How is energy related to Temperature? What happens to the temperature of a substance when you add energy?  Particles move faster  Temperature increases

Temperature Relationship between energy, particle speed, and temperature Temperature Definition  Average Kinetic Energy

Temperature Scales Kelvin (K) and Celsius (°C) scales Kelvin scale is called the absolute scale  Related to the kinetic energy of a substance Celsius scale is a relative scale  based on the boiling and freezing points of water

Temperature Conversion K = °C + 273

Pressure What is pressure? Physics – Force per unit area Chemistry – related to the number of collisions between particles and container walls

Pressure Conversion 1 atm = kPa

Vapor Pressure Pressure exerted by vapor that has evaporated and remains above a liquid Related to temperature  As temperature increases, vapor pressure increases

Boiling vs. Evaporation Boiling  Vapor pressure equals external, or atmospheric pressure Evaporation  Some molecules gain enough energy to escape the liquid phase  At temp. less than boiling point

Normal Boiling Point Boiling Point at Standard Pressure 1 atm or kPa

Evaporation Why is evaporation considered a cooling process? As the molecules with higher kinetic energy evaporate, the average kinetic energy of the substance decreases

Table H Shows the relationship between temperature and vapor pressure for four specific substances

Thermochemistry Chapter 17

Thermochemistry Heat involved with chemical reactions and phase changes

Heat Energy transferred from one object to another, usually because of a temperature difference Measured in Joules (J) or calories (cal) Heat flows from hot to cold

Heat Transfer Endothermic  Energy being added Exothermic  Energy being released

Specific Heat Capacity Amount of heat needed to raise the temperature of 1 g of a substance by 1°C  Unique for each phase of each substance 4.18 J/(g*°C) for liquid water  Listed in Table B of Reference Tables

Heat What factors affect the amount of heat transferred?  Specific Heat Capacity  Mass  Temperature difference between objects

Heat Equation Heat, q Mass, m Specific Heat Capacity, c Change in Temperature, ΔT q=m*c*ΔT

Example 200g of water is heated from 20°C to 40°C, how much heat is needed? q = m*c*ΔT q= (200g) * (4.18J/g°C) * (20°C) q= J

Example How much energy is required to raise the temperature of 50g of water from 5°C to 50°C? q = m*c*ΔT q= (50g) * (4.18J/g°C) * (45°C) q= 9405 J

Another Example What is the Specific Heat Capacity of Fe, if it takes 180J of energy to raise 10g of Fe from 10°C to 50°C? q = m*c*ΔT 180J = (10g) * c * (40°C) c= 0.45 J/(g*°C)

Phase Change At what temperature does ice melt?  0°C At what temperature does water freeze?  0°C Melting point and freezing point are the same

Phase Change What happens to temperature during phase changes?  Temperature remains constant Temperature remains CONSTANT during a phase change

Phase Change If energy is being added, what kind of energy is it?  Energy being added is potential energy, not kinetic energy  Potential energy is being used to separate or spread the particles apart

Heat of Vaporization, H v Amount of energy needed to vaporize 1g of a substance Water = 2260 J/g q=mH v Use for Liquid  Gas or Gas  Liquid

Heat of Fusion, H f Amount of energy needed to melt 1g of a substance Water = 334 J/g q=mH f Use for Solid  Liquid or Liquid  Solid

Examples How much energy is needed to melt 10g of ice at 0°C?  q = m*H f  q = (10g) * (334J/g)  q = 3340 J

Example How much energy is needed to vaporize 10g of water at 100°C?  q = m*H v  q = (10g) * (2260J/g)  q = J

Phase Change Which requires more energy melting or vaporization?  Vaporization Why?  Molecules are spread farther apart as a gas  It takes more energy to get gas particles spread apart

Heat of Solution Amount of energy released or absorbed when an ionic salt dissolves Δ H is called the heat of reaction  + Δ H means endothermic, energy term on left  - Δ H means exothermic, energy term on right

Heating (Cooling) Curves Shows relationship between temperature and time during constant heating or cooling. Also shows phases, and the phase changes between them.

Heating Curves Diagonal lines are phases Horizontal lines are phase changes Time (s) Temp (˚C) Gas Liquid Solid

Heating Curves Diagonal lines are phases Horizontal lines are phase changes Time (s) Temp (˚C) Vaporization Condensation Melting Freezing

Conservation of Energy Energy can not be created or destroyed, only transferred or converted from one form to another. Energy lost by one object must be gained by another object or the environment  q lost = q gained

Example A chunk of iron at 80°C is dropped into a bucket of water at 20°C. What direction will heat flow?  From the iron to the water  Hot to cold

Example A chunk of iron at 80°C is dropped into a bucket of water at 20°C. What could be the final temperature, when they both come to equilibrium?  Between 20°C and 80°C

Example A 100g block of aluminum, c=0.90J/g*°C at 100°C is placed into 50g of water at 20°C, what will be the final temperature when the aluminum and water reach equilibrium?  q lost = q gained  m*c*ΔT = m*c*ΔT  100g*0.90J/g°C*(100°C-T f ) = 50g*4.18J/g°C*(T f -20°C)  90*(100-T f ) = 209*(T f -20)  T f = 209T f  = 299T f  T f = 44°C