CLASSIFYING MATTER

Chapter 3: Matter & Energy

Separation of mixtures

Day One And Two

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

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!

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.

Solid Liquid

Gas

1. Arrangement2. Movement

3. Volume

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

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

Day FourClassification Activity

Classifying Matter Separating Mixtures

Day Five

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.

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 Factors1 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

Heat Capacity Specific Heat

Day Seven

Change in Heat Example

Window 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.

Hg has a specific heat of 0.139 J/g C. ⁰

How much heat is required to raise the temperature of a 22.80 grams sample from 16.2 C to 32.5 C? ⁰ ⁰

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

CalorimetryDay Eight

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.

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!!!

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

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?

Work Day Day Nine

Bugle Lab Day Ten

Bugle Lab WriteupDay Eleven

ReviewDay Twelve

TestDay Thirteen