Matter and EnergyPart 2

Concepts to Master• When elevation increases, what happens

to pressure?• Be able to convert between the different

units of heat.• Be able to use q=mC∆t to solve

problems.• Be able to use q=mHf to solve problems• Be able to use q=mHv to solve problems• Understand and be able to label a phase

diagram.• Be able to use Table I to determine the

heat of a reaction.• Be able to use table H to determine the

vapor pressure of a substance when the temperature is given.

• If the temperature of the products is greater than the temperature of the reactants, is the reaction exo or endothermic?

• What is the relationship between the change in temperature and the change in heat flow in an exothermic and endothermic reaction?

• Which type of solid containing a repetitive unit cell?

• What are properties of a liquid?• What are properties of a solid?• What are properties of a gas?• Are the attractions between molecules

within a liquid stronger or weaker than those in a solid?

• Which phase of matter is considered to be disordered?

• How far apart are the particles in a gas?

• What are the allotropes of carbon?• Be able to use table H to determine

the temperature of a substance when the vapor pressure is given.

• What is the relationship between heat, kinetic energy, andattractions between molecules?

• Describe the Kinetic Molecular Theory.• Be able to convert between the

different units for pressure.• Be able to convert between the

different units of temperature.

Vocab

• Absolute zero• Allotropes• Amorphous solids• Aqueous• Atmospheric

pressure• Attraction• Barometric

pressure• Calorimeter• Celsius• Condensation• Crystalline solids• Crystallization

• Deposition• Endothermic• Evaporation• Exothermic• Fahrenheit• Fusion• Gas• Heat• Heat capacity• Heat of fusion• Heat of reaction• Heat of vaporization• Kelvin• Liquid

• Molecular solids• Network solids• Plasma • Solid• Solute• Solution• Solvent• Specific heat• Standard pressure• Sublimation• Surroundings• System• Temperature• Vapor pressure• Vaporization

Labs

• What is a Physical State• Specific heat of Fe and energy in a Cheeto• Exo or Endo• Chocolate Allotropes

States of Matter

And Plasma!

Properties of Gases

• Often transparent• They can be compressed easily.• They expand to fill their containers.• They take the shape of the container.• Flows easily• Particles have complete freedom of

motion (disordered)• Particles are far apart• They exert pressure on the walls of

the container.Pg 23-24 (gases)

Properties of Liquids• Takes the shape of the

container

• Does not expand to fill the container

• Not very compressible

• Flows easily

• No ordered arrangement of particles, disordered

• Particles are free to move relative to each other

• Particles are close together

Pg 44 – 46 (liquids)

Properties of Solids

• Retains it’s own shape and volume

• Can NOT be compressed

• Does not Flow

• Particles are close together

Pg 50-52

Types of Solids• Crystalline solids (or ionic solids)

atoms, molecules or ions arranged in a regular pattern.

Sodium chloride and salts These solids have a fundamental

building block called a unit cells. Unit cells make up a complete crystal

when stacked or repeated in three dimensions.  

• Amorphous solids atoms, molecules or ions have no

regular pattern of arrangement Glass and tar These types of materials have NO

simple repeated unit cell. The structure is more complicated

because they lack the repetitive units.

Types of Solids - Network• Network solids are giant arrangements

of matter in which atoms are bonded together in two or three dimensions. Network solids are like giant molecules. – graphite

• A 2-dimensional network solid is arranged in layers

• weak attractions between the layers• This makes them generally soft and/or

slippery.

– Diamond• A 3-dimensional network is a giant

interlocking design• strong attractions• Giving the substance exceptional hardness

and a high melting and boiling point.

Pg 128-130

Types of Solids -Molecular

• Molecular solids consist of molecules that are held together by weak intermolecular forces.

• This type of solid has a low melting point, none are higher than 400ºC.  Some actually decay before melting. A prime example of this is sugar. 

– Molecules of sugar are held together by weak intermolecular forces

– Sugar caramelizes and/or burns before melting.

Types of Carbon• Allotropes: different forms of the same

element in the same state.

Plasma

• It is what makes up outer space.

• It is formed by heating and ionizing gas. By freeing the ions and electrons in a gas, plasma is created.

• It can conduct electricity and is found naturally on earth – lighting bolts (which are narrow paths of air molecules

and parts of a flame) – free electrons in metal can also be considered plasma– The character, motion and/or shape of plasma

depends on magnetism.

Changes in matter• Are Physical Changes

– Solid to Liquid• Melting

– Liquid to Solid • Crystallization

– Solid to Vapor • Sublimation

– Vapor to Solid• Deposition

– Liquid to Vapor– Boiling

– Evaporation

– Vapor to Liquid • Condensation

Phases of Matter and Freedom to Move

FreedomTo Move

FreedomTo Move

Phases of Matter and Energy

KE

PE

Phases of Matter and Distance

Distance increases

Distance decreases

Phases of Matter and Attraction

Attraction increases

Attraction decreases

Kinetic Molecular Theory• All molecules are moving, some are moving faster than

others.

• All molecules have kinetic and potential energy, some have more of one type than the other.– Kinetic energy is the energy of motion– Potential energy is the energy of position

• Energy in the phases of Matter– Solid molecules are only vibrating therefore they have little

kinetic energy, but lots of potential energy.– Gas molecules are moving all over therefore they have lots of

kinetic energy, and only a little potential energy.– Liquid molecules are in between.

Pg 24-27

Pressure• Pressure is force. A push that

tends to move something in a certain direction.

• Atmospheric Pressure– The pressure exerted by the

atmosphere (air).– Barometric pressure

• Vapor Pressure– When a vapor (gas) is produced

above a liquid while it is in a sealed container, these vapor particles exert pressure on the walls of the container.

– Table H

• Standard Pressure– The pressure exerted by the

atmosphere that will support a column of Hg that is 760 mm tall.

(pg 46-49 +questions)

Pressure Units

• 1 atm = 760 mmHg = 101325 Pascal (Pa) = 101.3 kPa• Conversions

– How many kPa are in 475 mmHg?

– How many atm in 740 mmHg?

– How many kPa in 98765 Pa?

Temperature

• Measure of hotness (or coldness)

• Heat flows from a substances at higher temp to a substances at lower temp.

• Fahrenheit Scale– Water freezes at 32oF– Water boils at 212oF– Not used in science

• Celsius Scale– 0oC is the point at which

water freezes at sea level.– 100oC is the point at which

water boils at sea level.

• Kelvin Scale– Water freezes at 273 K– Water boils at 373 K– Absolute zero

• Lowest temperature possible

• The point at which all molecular vibration stops

• 0 K • -273oC

1 degree K is the same size as 1 degree C.F degrees are not the same size as K or C.1 degree F is 9/5 of a degree C (or K).

Temperature Conversions

• K = oC + 273

• oF = 9/5(oC) + 32

• If the temperature of the day is predicted to reach 31oC. – What is the predicted temperature in K?

– What is the predicted temperature in oF?

Conversion Problems

• Ethylene Glycol freezes at 11.3oF. What is the freezing point in K?

• Convert 147.2 kPa to atm.

STP

• Standard Temperature and Pressure• Standard Temperature = 0oC (273 K)• Standard Pressure = 1 atm (760 mmHg)

– This is the normal atmospheric pressure at sea level.

– We are 937 ft above sea level. There is less atmosphere pressing upon us so the atmospheric pressure is less. It’s about 740 mmHg.

– As elevation above sea level increases, atmosphereic pressure decreases.

Forms of Energy

• Chemical

• Electrical

• Nuclear

• Thermal• Mechanical

• Electromagnetic

Phase ChangesHeating (red) / Cooling (blue) Curve

(pg 13 -18)

Is the phase change endothermic or exothermic?

• Endothermic Process– Heat is absorbed – Particles vibrate faster.– more kinetic energy, less

potential energy – Attractions between

particles become weaker

• Exothermic Process– Heat is released – Particles vibrate slower.– less kinetic energy, more

potential energy – Attractions between

particles become stronger

• Endothermic Process• Melting• Vaporizing, boiling,

evaporating• Sublimimg

• Exothermic Process• Freezing,

solidification, crystalization

• Condensing• Depositing

Table H• When the external pressure is 101.3 kPa, water will boil at what

temperature?

– 12.8°C – 14.5°C – 100°C – 18°C

• According to Reference Table H, when the vapor pressure of water is 101.3 kPa, the temperature is 100°C. When the atmospheric pressure above a liquid equals its vapor pressure, the liquid will boil. If the vapor pressure was 50 kPa, water would boil at about 82°C.

• Find the vapor pressure on Y axis, go across to the appropriate liquid line as indicated by the question. From the line, go down to the X axis and determine the temperature. This temperature is the boiling point of the substance.

Heat (q)• Heat is the energy used to cause the

temperature of an object to increase.• Heat is the energy transferred between a

“system” its surroundings.– If absorbed : endothermic

• T of system?• T of surroundings ?

– If released : exothermic• T of system?• T of surroundings ?

• Units for heat:– Joule (J) and kilojoule (kJ)– calorie (cal) = 4.184J– 1 food calorie = 1000 cal = 1 kcal

(Pg 18-23 and questions)

Calorimetry Measurement of heat flow.

∆H = Hfinal – Hinitial

∆H = Hproducts – Hreactants

• Exothermic when ∆H is negative.– Heat is going out of the

system.– Temp of surrounding

increases– Table I

• Energy within the products< energy within the reactants

• Endothermic when ∆H is positive.– Heat is going into the

system.– Temp of surrounding

decreases.– Table I

• Energy within the products> energy within the reactants

Heat Capacity• The amount of heat needed to raise the temperature of

an object by 1 K (or oC).– If an object has a great heat capacity, it will take more heat to

cause the temp change.– The heat capacity for 1 mol of any substance is its molar heat

capacity.– The heat capacity for 1 gram of any substance is its specific

heat. • The specific heat of water = 4.18 J/goC• The quantity of heat either gained or lost can be

calculated with the following equation:q = mC∆T

– q = heat– m = mass– C = specific heat– ∆T = Tfinal – Tinitial

Heat Problems

• 2000g of water has a temperature change of 3.0oC. How much heat has been transferred?

• A 100g sample of gold is warmed from 18oC to 32oC. The specific heat of gold is 0.129 J/goC. How much heat is required to make this change?

Calorimeter

• Device to measure heat flow

• Used to determine heat capacity of a material or chemical reaction.

Calorimetry Problems

• 200g of water is placed into a calorimeter. An exothermic reaction takes place in the water and cause the temperature to increase by 5oC. How much heat is transferred to the water by the heat of the reaction?

• 56g of water is placed into a calorimeter. An endothermic reaction takes place in the water and cause the temperature to decrease by 6.2oC. How much heat is absorbed from the water by the reaction?

Heat of…

• Fusion (Hf)– The heat absorbed by a substance while melting from a solid to

a liquid. – The temp remains constant during the melting process.– q = m Hf

– Table B: Hf for water = 334J/g

• Vaporization (Hv)– The heat absorbed by a substance in changing from a liquid to a

gas.– The temp remains constant during the vaporization

(evaporation) process.– q = m Hv

– Table B: Hv for water = 2260J/g

Complete Phase Diagram

BP

MP / FP

Hf

Hv

Heat of Sublimation• Sublimation – a phase

change where a solid (with very weak intermolecular forces) becomes a gas without going through the liquid phase.

• Heat of sublimation (Hs)– The heat absorbed by a

substance in changing from a solid to a gas.

– The temp remains constant during the sublimation process.

– q = m Hs

– Not on table B since water does not sublime at STP!

Dry IceSolid carbon dioxide becomes a gas.

Hs ~ 5975 J/g

Deposition

• Opposite of sublimation

• Gas becomes a solid without going through the liquid phase

• Snow

Heat of Fusion and Vaporization

• How many joules are required to melt 100 grams of water?

• How many joules are required to boil 150 grams of water?

• How much ice could be melted by the addition of 576J of heat?

Exo or Endo, More or Less Kinetic Energy, Stronger or Weaker Forces of Attraction?

• Endothermic Process– Heat is absorbed causing molecules to vibrate faster.– more kinetic energy, less potential energy – Resulting in substance with weaker intermolecular forces

• Melting (fusion)• Vaporizing (evaporation)• Subliming

• Exothermic Process– Heat is released causing molecules to vibrate slower.– less kinetic energy, more potential energy – Resulting in substance with stronger intermolecular forces

• Freezing (solidification)• Condensing• Depositing