chapter 3: matter & energy. day one and two separation of mixtures objectives: 1. make a...
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CLASSIFYING MATTER
Chapter 3: Matter & Energy
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Separation of mixtures
Day One
And Two
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Separation of Mixtures
Objectives: 1. Make a mixture out of sand, salt, and iron
2. Separate the sand, salt, and iron mixture
3. Determine percent recovered of iron, salt, and sand
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Separation of Mixtures
You and your partner must submit to me a procedure as to how you are going to separate the salt, sand, and iron mixture before you go into the lab
Materials Used: ANYTHING YOU WANT! ASK!
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Chemical & PhysicalConservation of
Matter
Day Three
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What Is Matter? Matter is defined as
anything that occupies space and has mass.
Even though it appears to be smooth and continuous, matter is actually composed of a lot of tiny little pieces we call atoms and molecules.
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Atoms and Molecules
Atoms are the tiny particles that make up all matter.
In most substances, the atoms are joined together in units called molecules.The atoms are joined in
specific geometric arrangements.
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Classifying Matterby Physical State
Matter can be classified as solid, liquid, or gas based on what properties it exhibits.
State Shape Volume Compress Flow
Solid Fixed Fixed No No
Liquid Indefinite Fixed No Yes
Gas Indefinite Indefinite Yes Yes
• Fixed = Property doesn’t change when placed in a container. • Indefinite = Takes the property of the container.
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Solid Liquid
Gas
1. Arrangement2. Movement
3. Volume
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Physical and Chemical Properties PHYSICAL CHEMICAL
Characteristic that is displayed by the substance WITHOUT changing its composition
Examples: Odor Boiling Point Melting Point Density
Characteristics that is displayed by the substance WITH changing its composition
Examples: Flammability Corrosiveness Acidity Toxicity
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Physical and Chemical ChangePhysical Change Chemical Change
Matter changes its appearance but not its composition
Example: Phase Changes Change in appearance
• Matter DOES change its composition
• Results in a completely NEW substance
• Example: – Burning – Heat exchange– Evolution of a gas – Formation of a precipitate
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Day Four
Classification Activity
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Classifying Matter Separating Mixtures
Day Five
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Types of Matter
Matter
Pure Substance
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Pure Substances vs. MixturesPure Substances
1. All samples have the same physical and chemical properties.
2. Constant composition = All samples have the same pieces in the same percentages.
3. Homogeneous.4. Separate into components
based on chemical properties.
5. Temperature stays constant while melting or boiling.
Mixtures1. Different samples may
show different properties.2. Variable composition =
Samples made with the same pure substances may have different percentages.
3. Homogeneous or heterogeneous.
4. Separate into components based on physical properties.
5. Temperature usually changes while melting or boiling because composition changes.
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Separation of Mixtures Separate mixtures based on different
physical properties of the components.Physical change.
Centrifugation anddecanting
Density
EvaporationVolatility
ChromatographyAdherence to a surface
FiltrationState of matter (solid/liquid/gas)
DistillationBoiling point
TechniqueDifferent Physical Property
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Distillation
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Filtration
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Law of Conservation of Mass Antoine Lavoisier “Matter is neither created nor destroyed in a
chemical reaction.” The total amount of matter present before a
chemical reaction is always the same as the total amount after.
The total mass of all the reactants is equal to the total mass of all the products.
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Conservation of Mass Total amount of matter remains constant in a
chemical reaction. 58 grams of butane burns in 208 grams of
oxygen to form 176 grams of carbon dioxide and 90 grams of water.
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EnergyDay Six
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Energy
There are things that do not have mass and volume.
These things fall into a category we call energy.
Energy is anything that has the capacity to do work.
Although chemistry is the study of matter, matter is effected by energy.It can cause physical and/or chemical changes in
matter.
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Some Forms of Energy Electrical
Kinetic energy associated with the flow of electrical charge. Heat or Thermal Energy
Kinetic energy associated with molecular motion. Light or Radiant Energy
Kinetic energy associated with energy transitions in an atom.
NuclearPotential energy in the nucleus of atoms.
ChemicalPotential energy in the attachment of atoms or because of
their position.
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“Losing” Energy If a process was 100% efficient, we could
theoretically get all the energy transformed into a useful form.
Unfortunately we cannot get a 100% efficient process.
The energy “lost” in the process is energy transformed into a form we cannot use.
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Units of Energy Calorie (cal) is the amount of energy needed
to raise one gram of water by 1 °C.kcal = energy needed to raise 1000 g of water 1 °C.food calories = kcals.
Energy Conversion Factors
1 calorie (cal) = 4.184 joules (J)
1 Calorie (Cal) = 1000 calories (cal)
1 kilowatt-hour (kWh) = 3.60 x 106 joules (J)
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Exothermic Processes When a change results in the release of
energy it is called an exothermic process.
The excess energy is released into the surrounding materials, adding energy to them.Often the surrounding materials get hotter from
the energy released by the reaction.
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Endothermic Processes When a change requires the absorption of
energy it is called an endothermic process.
The required energy is absorbed from the surrounding materials, taking energy from them.Often the surrounding materials get colder due
to the energy being removed by the reaction.
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Temperature Scales Fahrenheit scale, °F.
Used in the U.S. Celsius scale, °C.
Used in all other countries.
A Celsius degree is 1.8 times larger than a Fahrenheit degree.
Kelvin scale, K.Absolute scale.
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Temperature Scales
The Fahrenheit temperature scale used as its two reference points the freezing point of concentrated saltwater (0 °F) and average body temperature (96 °F).More accurate measure now sets average
body temperature at 98.6 °F. Room temperature is about 72 °F.
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Temperature Scales, Continued
The Celsius temperature scale used as its two reference points the freezing point of distilled water (0 °C) and boiling point of distilled water (100 °C).More reproducible standards.Most commonly used in science.
Room temperature is about 22 °C.
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Fahrenheit vs. Celsius
A Celsius degree is 1.8 times larger than a Fahrenheit degree.
The standard used for 0° on the Fahrenheit scale is a lower temperature than the standard used for 0° on the Celsius scale.
F-32C
1.8
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The Kelvin Temperature Scale
Both the Celsius and Fahrenheit scales have negative numbers.Yet, real physical things are always positive amounts!
The Kelvin scale is an absolute scale, meaning it measures the actual temperature of an object.
0 K is called absolute zero. It is too cold for matter to exist because all molecular motion would stop.0 K = -273 °C = -459 °F.Absolute zero is a theoretical value obtained by
following patterns mathematically.
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Kelvin vs. Celsius The size of a “degree” on the Kelvin scale is
the same as on the Celsius scale.Although technically, we don’t call the divisions on
the Kelvin scale degrees; we call them kelvins!That makes 1 K 1.8 times larger than 1 °F.
The 0 standard on the Kelvin scale is a much lower temperature than on the Celsius scale.
When converting between kelvins and °C, remember that the kelvin temperature is always the larger number and always positive!
K C 273
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Heat Capacity Specific Heat
Day Seven
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Change in Heat
ExampleWindow in the winter time
Energy always flows in the same direction
→ When does the energy flow stop?
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Heat Capacity Heat capacity is the amount of heat a substance
must absorb to raise its temperature by 1 °C.cal/°C or J/°C.Metals have low heat capacities; insulators
have high heat capacities. Specific heat = heat capacity of 1 gram of the
substance.cal/g°C or J/g°C.Water’s specific heat = 4.184 J/g°C for liquid.
○ Or 1.000 cal/g°C.○ It is less for ice and steam.
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Heat Capacity Heat capacity is the amount of heat a substance
must absorb to raise its temperature by 1 °C.cal/°C or J/°C.Metals have low heat capacities; insulators
have high heat capacities. Specific heat = heat capacity of 1 gram of the
substance.cal/g°C or J/g°C.Water’s specific heat = 4.184 J/g°C for liquid.
○ Or 1.000 cal/g°C.○ It is less for ice and steam.
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Specific Heat CapacitiesSubstance Specific Heat
J/g°C Aluminum 0.903
Carbon (dia) 0.508
Carbon (gra) 0.708
Copper 0.385
Gold 0.128
Iron 0.449
Lead 0.128
Silver 0.235
Ethanol 2.42
Water (l) 4.184
Water (s) 2.03
Water (g) 2.02
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Heat Gain or Loss by an Object
The amount of heat energy gained or lost by an object depends on 3 factors: how much material there is, what the material is, and how much the temperature changed.
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Hg has a specific heat of 0.139 J/g C. ⁰How much heat is required to raise the temperature of a 22.80 grams sample from 16.2 C to 32.5 C? ⁰ ⁰
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How many joules of heat are required to raise the temperature of 200 grams of water from 20.0 C to 50.0 C? ⁰ ⁰
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CalorimetryDay Eight
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All food has energy, so how can we measure it? Energy remember is a transfer of heat
Some food (Bugles for instance) we can burn and it will continue to burn on its own until it uses up all the energy in the food.
If we can measure the heat it gives off we can calculate the energy.
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Measure energy in food
If its giving off heat then we can measure the temperature change in the surrounding air
However, the energy would dissipate very quickly and it would not be a good way to get the temperature change
We use calorimeters!!!
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What is Heat Capacity?
heat capacity is for objects whose size is predetermined and we can factor out the mass.
Units for Heat Capacity is (J/ C)⁰
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A bomb calorimeter was filled with propane which was then ignited. This reaction released 104,000 J of energy. Initially, the temperature of the calorimeter was 25 C, ⁰after the reaction the temperature was measured at 47.5 C. What is the heat ⁰capacity of this calorimeter?
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Work Day Day Nine
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Bugle Lab Day Ten
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Bugle Lab WriteupDay Eleven
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ReviewDay Twelve
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TestDay Thirteen