Chemistry Unit Two

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Chemistry Unit Two

• Matter and Energy

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Matter

• Matter
• anything that has a mass and takes up space.
• Law of Conservation of Mass/Matter
• Matter cannot be created or destroyed in an
  ordinary chemical reaction just rearranged to
  form different substances
• Matter can be described using properties..

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Types of Properties
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Characteristics of Matter

• Physical Properties
• Characteristics of a substance that can be
  observed without the production of a new
  substance.
• Examples
• Color,smell, taste, hardness, density, texture,
  melting/boiling/freezing points, magnetic
  attraction, solubility, electrical conductivity,
  temperature, state or phase

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Two Types of Physical Properties

• Extensive
• Depends on the particular sample
• examples volume, mass, weight, shape, etc
• Intensive
• Depends on the type of matter ? NOT size of
  sample
• examples color, melting point, specific heat,
  density, appearance, etc

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Characteristics of Matter

• Chemical Properties
• describes how a substance reacts or fails to
  react with other substances to produce new
  substances.
• Examples
• Oxidation, Corrosion, Hydrolysis, Combustion,
  Flammability, Reaction to Acid or Base.

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Two Types of Changes

• Physical Change
• an alteration of a substance that only changes
  the physical properties of the substance.
• Does not change the chemical composition of the
  matter!!

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Characteristics of Matter

• Chemical Change
• an alteration of the chemical composition of a
  substance that results in the formation of a new
  substance
• ALWAYS forms a new substance that has different
  physical and chemical properties than the
  original substance.
• Also known as a chemical reaction.

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• Chemical and Physical change
• Practice 1
• Practice 2

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Kinetic Theory

• All matter is made of tiny particles in constant
  motion.
• Potential Energy (PE)
• energy due to the position or condition
• at the atomic level
• the distance between the particles
• closer lower PE farther higher PE
• Kinetic Energy (KE)
• energy due to motion
• Fasterhigher KE slower lower KE

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Phases of Matter
State or Phase Particle level picture Particles description Keep Volume? Keep shape?
Solid
Liquid
Gas

• Arranged in orderly pattern

• Yes

• Yes

• Touching, but not tightly packed

• No

• Yes

• Far apart and rarely touching

• No

• No

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Phases of Matter
State or Phase Particle Movement Amount PE Amount KE Example
Solid
Liquid
Gas
Vibrational only
Low
Very Low
Ice
Vibrational translational
Low
Moderate
Water
High
Vapor
High
Move freely
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Plasma

• extraordinary state of matter
• consists of high energy particles
• electrons are stripped from their nuclei
• examples
• fluorescent light
• Stars
• Lightning
• Most Abundant State of Matter in the Universe!

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Phase Changes Changes of State

• Adding or removing energy (heat) to a substance
  causes phase changes
• The potential energy of the particles is
  increased or decreased
• During a phase change, temperature does NOT change

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Phase Changes

• Melting
• S ? L (adding energy)
• Freezing
• L ? S (removing energy)
• Melting point freezing point of a substance
  occur at the same temperature.

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Phase Changes

• Boiling
• L ? G (adding energy)
• Evaporation
• L ? G (adding energy)
• Condensation
• G ? L (removing energy)
• Difference between boiling evaporation
• Boiling?a specific temp. below the surface
• Evaporation ?any temp. at the surface

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Phase Changes

• Deposition
• G ? S (removing energy)
• Examples Snow, frost
• Sublimation
• S ? G (adding energy)
• Examples solid CO2 (dry ice), solid air
  fresheners

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Phase Change Graphs (T vs t)
Liquid
Melting
Solid
AB -heat ? KE -move faster -temp. ? -solid
BC -heat ? PE -get farther apart -temp. stay
same -melting
CD -heat ? KE -move faster -temp. ? -liquid
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Phase Change Graph (T vs t)
Gas
Boiling
DE -heat ? PE -get farther apart -temp. stay
same -boiling
EF -heat ? KE -move faster -temp. ? -gas
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Phase Change Graph (T vs t)
A
C
B
E
D
F
CD -KE ? -slows down -temp. ? -Liquid
AB -KE ? -slows down -temp. ? -Gas
BC -PE ? -closer together -temp. stays
same -Condensation
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Phase Change Graph (T vs t)
A
B
C
D
E
F
DE -PE ? -closer together -temp. stays
same -Freezing
EF -KE ? -slows down -temp. ? -Solid
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Phase Change Graph (T vs t)
Boiling Point
Boiling
Freezing
Freezing Point
Melting Point
Melting
What is the boiling point? What is the melting
point? What is the freezing point?
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Phase Change Graph (T vs t)
If melting freezing points occur at the
same temperature, how do you know which change is
occurring? -depends on whether adding or
removing energy
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Phase Change Graph (T vs t)
What is this substance? -Water How do you
know? -Boiling melting freezing points of
water (Intensive properties)
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Heat Calculations

• Heat (q)
• Energy transferred from an object at a higher
  temperature to an object at a lower temperature.
  (heat lost -heat gained)
• q mc?T
• qmHfus
• qmHvap

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Heat Calculations

• A 10.0g sample of iron at 50.4oC is cooled to
  25.0oC in 50.0g of water. Calculate the amount
  of heat lost by the iron.
• ciron 0.449 J/goC
• How much heat is gained by the water?
• A 2.1g ice cube at 8.0oC melts completely and
  warms to 12.5oC. How much heat was required?
• Hfus ice 334 J/g
• cice 2.03 J/goC
• cwater 4.18J/goC

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Classification of Matter
Matter
Pure Substances
Mixtures
Elements
Compounds
Homogeneous
Heterogeneous
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Matter

• Pure Substances
• made of only one type of matter
• Mixtures
• a physical combination of two or more substances
• no definite ratio of particles
• Element
• made of only one type of atom
• cannot be broken down into simpler substances
  under normal laboratory conditions

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Matter (contd)

• Compound
• Atoms of two or more elements, chemically
  combined in a definite ratio.
• Homogeneous Mixtures
• Atoms of two or more elements, physically
  combined in no definite ratio.
• The same throughout.
• Must be a SOLUTION
• Heterogeneous Mixture
• Atoms of two or more elements,
• physically combined in no definite ratio.
• Different throughout

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Classifications of Mixtures

• Solutions
• Particles are very tiny, will not separate by
  filtering, will not settle out when allowed to
  stand, particles will not scatter light, (-)
  Tyndall effect.
• Ex. Salt Water, Kool-Aid, Brass
• Colloids
• Particles are tiny, will not separate by
  filtering, will not settle out when allowed to
  stand, particles will scatter light, () Tyndall
  effect.
• Ex. Milk, whipped cream, aerosols
• Suspension
• Particles visible with unaided eye, will separate
  when filtered, will settle out if allowed to
  stand, particles will scatter light, () Tyndall
  effect.
• Ex. Muddy water, snow globe

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Solutions

• SOLUTION
• a solute dissolved in a solvent.
• The solvent is the part in greater quantity.
• For example In a solution of salt water, salt
  is the solute and water is the solvent.
• ELECTROLYTE
• a solution that conducts electricity in water or
  molten form.
• Salt water will conduct electricity.(Electrolyte)
• Sugar water will not.

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Types of Solutions

• Gas-Gas
• Carbon dioxide, Nitrogen,Oxygen (air)
• Liquid-Gas
• Water Vapor in Air (moist air)
• Gas-Liquid
• Carbon dioxide in Water (soda water)
• Liquid-Liquid
• Acetic acid in Water (vinegar)
• Solid-Liquid
• Sodium chloride in Water (brine or salt water)
• Solid-Solid
• Copper in Silver (Sterling Silver)

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Characteristics of Solutions

• Homogeneous Mixture
• Solute / solvent
• Soluble-
• Likes dissolve likes
• Insoluble
• Immiscible
• Miscible

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Solvation

• When a solid solute is placed in a solvent, the
  solvent particles completely surround the surface
  of the solid solute.
• If attractive forces between the solute particles
  and the solvent are greater than the attractive
  forces holding the the solute particles together,
  the solvent particles pull the solute particles
  apart and surround them.

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Process of Solvation
H2O H O
-
-

-

NaCl Na Cl

-
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Water- Universal Solvent

• Polar molecule
• Dipoles allow solvation of ions and polar
  molecules

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Solvation Factors

• Agitation
• Increasing collisions and breaking up solute
  attraction
• Increasing surface area of solute
• Small pieces, more area for solvent to contact
• Increasing temperature of solvent
• Greater kinetic energy creates more collisions

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Heat of Solution

• Endothermic-
• Solute and solvent particles break attractive
  forces holding them together
• Exothermic-
• When solute and solvent particles mix, particles
  now attract each other

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Solubility

• Refers to the maximum amount of solute that will
  dissolve in a given amount of solvent at a
  specified temperature and pressure.
• g solute / 100 g solvent
• Saturated vs Unsaturated vs Supersaturated

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Concentration

• How much solute is dissolved in a specific amount
  of solvent
• Percent by mass
• Percent by volume
• Molarity moles/Liter
• Molality moles solute/kilograms solvent

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Molarity

• moles of solute / Liters of solution
• Calculate the molarity of 1.60 L of a solution
  containing 1.55 g of dissolved KBr?
• How many grams of CaCl2 would be dissolved in 1.0
  L of a 0.10M solution of CaCl2?

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Diluting Solutions

• M1V1 M2V2
• What volume of a 3.00 M KI stock solution would
  you use to make ).300 L of a 1.25 MKI solution?

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Colligative properties of Solns

• Physical properties of solutions that are
  affected by the number of particles but not the
  identity of dissolved solute particles
• Vapor Pressure Lowering
• Boiling point elevation
• Freezing point depression