Chemistry Review Module 2

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Chemistry ReviewModule 2
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Inorganic Compound Classification

• Two main kinds of compounds
• 1. Ionic made up of ions of opposite charge
• A. strong electrostatic force of attraction
• ionic bond
• B. electrons are transferred
• 2. Covalent made up of two or more nonmetals
• A. electrons are shared

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Periodic Table of the Elements
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Chemical Bonds

• Atoms bond with each other to become more
• chemically stable than they were before they
• bonded. To do this their outer electron
• (valence) shell must be complete.
• 1. The Octet Rule states that atoms will either
• gain, lose , or share valence electrons to
• attain 8 electrons in their outer (valence)
  
• shell to become stable.
• 2. The Noble Gases are the only one group of
• elements that are already stable.

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Naming Binary Ionic Compounds

• Ions of opposite charge are bonded together.
• 1. The metal cation () is written first and is
• named by the metals name
• 2. The nonmetal anion (-) is written second and
• is named by the nonmetals name with a
• revised ending of - ide.
• 3. Net charge of ions in the compound 0.
• 4. Subscripts are used to indicate the number
• of ions needed to attain the necessary net
• charge of 0.

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Naming Binary Ionic Compounds
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Examples of Binary Ionic Compounds

• What would the formulas be?
• 1. Sodium chloride
• 2. Lithium nitride
• 3. Calcium Fluoride
• What would the names be?
• 1. Al2S3
• 2. BaO
• 3. MgBr2

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Answers

• 1. NaCl
• 2. Li3N
• 3. CaF2
• Aluminum sulfide
• Barium oxide
• Magnesium bromide

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Periodic Table of the Elements
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Covalent (Molecular) Compounds

• Made up of nonmetals that share
• electrons between atoms.
• 1. This type of bond is called a
• covalent bond.

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Naming Binary Covalent Compounds

• 1. The first nonmetals name is that of the
• element.
• 2. The second nonmetals name has an -ide
• ending, just like with ionic compounds.
• 3. Use prefixes to describe the subscripts
• (1 - mono 2 - di 3 - tri 4 - tetra 5 -
  penta
• 6 - hexa 7- hepta 8 - octa 9 - nona
• 10 - deca)

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Examples of Covalent Compounds

• P2O5 - diphosphorus pentoxide
• How would you write the following formulas?
• 1. Carbon monoxide
• 2. Tetranitrogen decoxide
• 3. Dinitrogen pentoxide

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• 1. CO
• 2. N4O10
• 3. N2O5

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Balancing Chemical Equations

• The Law of Conservation of Matter
• 1. The mass of the products the mass of the
• reactants (matter is not created or
• destroyed).
• 2. Chemical equations are balanced to ensure
• that a chemical reaction follows the law of
• conservation of matter.

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Five Balancing Equation Guidelines

• 1. Count the number of atoms of each element
• on both the reactant and the product side.
• 2. Use coefficients (the numbers in front of the
• chemical symbol or formula)
• 3. Never add or change the subscripts.
• 4. There are 7 Diatomic elements
• (N2, O2, F2, Cl2, Br2, I2, H2)
• 5. Balance the hydrogen atoms and the
• oxygen atoms last.

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Balancing Chemical Equations

• 6. A chemical equation uses symbolic language
• to describe a chemical reaction
• 7. Equation means equal numbers of atoms of
• each element on both sides.
• 8. Quantities of reactants and products are
• expressed in moles by using coefficients

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Balancing Chemical Equations

• Symbols for balancing equations
• (s) solid (g) gas (l) liquid
• (aq) aqueous (dissolved in water)
• Go to www.usaprep.com and practice
• balancing equations.

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Balancing Chemical Equations
Try these equations 1. H2 O2
H2O 2. Ca Br2 CaBr2 3.
Ba(s) O2(g) BaO(s) 4. Mg AuCl3
MgCl2 Au
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Balancing Chemical Equations
The Answers 1. 2H2 O2 2H2O 2.
Ca Br2 CaBr2 3. 2Ba(s) O2(g)
2BaO(s) 4. 3Mg 2AuCl3 3MgCl2 2Au

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Acids and Bases

• Acid Definitions
• 1. Sour taste
• 2. Neutralize the actions of bases
• 3. Blue litmus paper turns red
• 4. Liberates hydrogen gas when reacted with
• certain metals
• 5. Examples Foods and drinks

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Acids and Bases

• Base Definitions
• 1. Bitter taste
• 2. Neutralizes the action of acids
• 3. Slippery to the touch
• 4. Red litmus paper turns blue
• 5. Examples Cleaning solutions

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Acids and Bases

• An acid a base a salt water
• Strong acids
• hydrochloric acid, sulfuric acid, nitric acid
• Strong base
• sodium hydroxide
• Weak acid
• acetic acid
• Weak base
• ammonia

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pH Scale

• The pH scale has
• values from 0 - 14.
• 0 - 6 is acidic
• 7 is neutral
• 8 -14 is basic

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The Mole

• The mole is the SI unit of measure for the
• amount of a substance.
• 1. The mole is a way to measure the mass of
• elements and compounds.
• 2. The molar mass of any element is
• numerically equal to its atomic mass and has
• the units of g/mol.
• 3. Example The mass of one mole of
• potassium (K) is 40 grams.

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Physical and Chemical Changes

• 1. Physical Change - Only the appearance
• changes. The identity is still the same. Some
  
• of the kinds of changes are and
• phase (state).
• A. size and shape (splitting, breaking,
• tearing, or hammering).
• B. phase or state (melting, vaporizing,
• freezing, or condensing).

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Physical and Chemical Changes

• 2. Chemical Change - The appearance and the
• identity changes. A new product is formed.
• The way you know a new product is formed
• is by the following
• A. gas formed B. color change
• C. mass change D. heat change
• E. solid formed F. light released
• Examples are rusting, growing, burning,
• combusting, fermenting, cooking, frying, and
• exploding.

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Nuclear Reactions

• A nuclear reaction results when an unstable
• nucleus breaks down and emits radioactive
• particles.
• 1. There are 3 types of particles released during
  
• radioactive decay.
• A. the alpha particle (a) (released during
  
• alpha decay
• B. the beta particle (ß) (released during
  beta
• decay
• C. the gamma particle (?) (always released
• during radioactive decay of any kind.

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Three Types of Radioactive Decay

• 1. Alpha decay, alpha particle (helium nucleus
• with a 2 charge) released
• A. largest, slowest, least penetrating
  particle
• 2. Beta decay, beta particle (fast moving
• electron with a 1 charge) released
• A. more penetrating than an alpha particle
• 3. Gamma radiation, (no particle, no mass or
• charge - a form of energy) released
• A. the most penetrating and the most
• damaging (shielded by lead)

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Half Life

• The half life of a radioactive isotope is the
  time
• it takes for one half of the isotope to decay.
• 1. Example The half life of mercury - 195 is 31
• hours. If you start with 20 g, how much will
  
• be left after
• (A) 31 hours?
• (B) 62 hours?

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Answers

• 1. 10 grams
• 2. 5 grams

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Particles in an Electric Field
Remember the charges for each particle. Also
remember that like charges repel and
opposite charges attract.
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Nuclear Fission and Nuclear Fusion

• Nuclear Fission is the splitting of nuclei
  resulting
• in a tremendous release of energy.
• Nuclear Fusion is the combining of nuclei
• resulting in an even greater release of energy.
• The sun uses nuclear fusion to produce energy
• 1. 2 hydrogen atoms combine to form 1 helium
• atom.

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Energy/ Heat/ Phase Changes

• 1. Temperature is the measure of the average
• kinetic energy of particles.
• 2. Heat is a form of Energy that may be
• absorbed or released. Heat flows from a
• warm body to a cooler one until equilibrium
  is
• reached.
• 3. A calorie is the amount of heat required to
• raise the temperature of one g of water one
• degree Celsius.

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Types of Phase Changes

• 1. melting - solid to liquid
• 2. freezing - liquid to solid
• 3. evaporation - liquid to gas
• 4. condensation - gas to liquid
• 5. sublimation - solid to gas
• 6. deposition - gas to solid

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Key Terms for Heat

• 1. Energy - the capacity to do work
• 2. Heat - energy transferred from one object to
• another (The SI unit of heat is the Joule
  (4.18
• Joules 1 calorie).
• 3. Thermochemistry the study of heat effects
• in chemical reactions
• 4. Combustion - chemical reactions that
• release heat
• 5. Exothermic reaction - one that releases heat
• 6. Endothermic reaction - one that absorbs heat

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Heat

• Specific heat capacity is the amount of heat
• required to raise the temperature of one gram of
• a substance one degree Celsius. The specific
• heat capacity of water is 1.
• 1. Water has a high specific heat capacity due
• to hydrogen bonds. This means that it
• doesnt change temperature very much
• despite a large amount of addition and
• subtraction of heat energy.
• 2. Metals have low specific heat capacities.

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Energy in Chemical Reactions

• The bond breaking that occurs in reactants
• during a chemical reaction requires energy.
• The bond formation that occurs in products
• during a chemical reaction releases energy.
• An endothermic or exothermic reaction is
• determined by the balance between these two
• processes.

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3 Phases (States) of Matter

• The balance between the attractive forces
• between the particles and the kinetic energy of
• the particles determines the phase of matter.
• 1. High KE of the particles and low attractive
• forces between the particles gas.
• 2. Low KE of the particles and high attractive
• forces between the particles solid.
• 3. Intermediate KE of the particles and
• intermediate attractive forces between the
• particles liquid.

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Solids

• The two main types of solids crystalline solids
• and amorphous solids.
• 1. In a crystalline solid, the atoms, ions, or
• molecules are arranged in an orderly,
• repeating, 3-dimensional pattern (crystal
• lattice)

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Solids

• 2. In an amorphous solid, the internal structure
• lacks order. Atoms, ions, or molecules are
• randomly arranged.
• A. Generally, these substances are cooled
• rapidly. There is not enough time for
  the
• particles to arrange themselves in a
• pattern.
• B. Examples Rubber, glass, plastics,
• polymers

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Allotropes

• Allotropes are substances with the same
• elemental composition, but different geometric
• arrangements.
• 1. Example, carbon has 4 allotropes
• A. diamond - formed under tremendous
• pressure
• B. graphite - more loosely packed
• C. soot - randomly bonded (amorphous
• form)
• D. buckey ball

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Kinetic Molecular Theory of Gases

• 1. Gases are made up of very small particles
• called atoms or molecules.
• 2. Gas particles are separated by large
• distances (low density).
• 3. The particles are in constant, random,
• straight-line motion undergoing thousands of
• collisions per second.
• 4. Collisions are perfectly elastic - total
  kinetic
• energy remains constant.
• 5. Gases exert a pressure due to the collisions
• on each other.

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Collision Theory

• 1. When gas particles collide they exert a
• pressure on their container.
• 2. Temperature is the measure of the average
• kinetic energy of the gas particles.
• 3. There are 4 main properties of gases that
• determine their physical behavior.

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Four Main Properties of Gases

• 1. Pressure force / area
• 2. Temperature the average kinetic energy of
• particles
• 3. Volume space occupied by matter.
• 4. Amount of gas is measured in grams or
• moles.

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What happens to the volume of a gas if the
pressure is increased at constant temperature?
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What happens to the volume of a gas if the
temperature is increased at constant pressure?
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Lewis Dot Structures
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Rules for Lewis Dot Structures

• 1. Count the total number of dots (valence
• electrons) in the structure.
• 2. Spatially arrange the atoms. (More than two
• atoms - locate the central atom)
• 3. Try to obtain 8 dots (valence electrons)
• around each atom (2 dots around hydrogen).
• 4. If single bonds dont work, try double bonds,
• then triple bonds.

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Lewis Dot Structures
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Three Factors Affecting the Rate of Dissolving

• How fast will a solute dissolve in a solvent?
• 1. Stirring or agitation (more solute / solvent
• contact at a faster rate)
• 2. Smaller particles (increases the surface area
• of the solute, therefore there is more solute
  /
• solvent contact at a faster rate)
• 3. Increasing the temperature (increases the
• kinetic energy and faster rate of contact
• between the solute/solvent particles)

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The Solubility of Solids at Different
Temperatures

• 1. What is the solubility
• of sugar at 50oC?
• 2. Which solute is least
• affected by an
• increase in temp?
• 3. Generally, as temp
• increases, Solubility________.

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• 1. 260 grams
• 2. NaBr
• 3. Increases

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GHSGT Chemistry Review

• The End