Ionic Bonding

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Chemistry

• Is the study of matter, its properties and its
  changes or transformations

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Matter

• Anything that has mass and takes up space

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Physical Properties

• a change in the size or form of a substance which
  does not change the chemical properties of the
  substance is called a physical change
• a physical property is a characteristic of a
  substance
• these properties allow us to distinguish or tell
  the difference between substances
• examples of physical properties - state, color,
  odor, luster, texture, hardness, crystal form,
  mass, volume, density, solubility, viscosity,
  malleability, ductility, melting point, boiling
  point

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Chemical Properties

• is a characteristic behavior that occurs when a
  substance changes into a new substance\
• the change itself is called a chemical change
• the starting materials we call reactants - the
  final materials we call products
• examples of chemical properties are reactions
  with water, reactions with acids, combustion
• example  iron oxygen       rust

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Introduction to Naming

• Today, most compounds are known by their IUPAC
  names.  IUPAC stands for International Union of
  Pure and Applied Chemistry. This organization has
  determined a set of rules to be used for naming
  chemicals. Its purpose is to set international
  guidelines so that all scientists follow the same
  rules.

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Classifying Compounds

• There are three main types of compounds that we
  will be dealing with
• Ionic a combination of metals and non-metals
• Molecular or covalent a combinations of two
  nonmetals
• Inter-metallic or metallic a combination of two
  metals

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Compound Properties Examples
Molecular Solid, liquid or gas at STP Relatively low melting and boiling points Do not conduct electricity in aqueous solutions (non electrolyte) May be soluble or insoluble in water Sugar Water Propane
Metallic Ductile, malleable Good conductor of heat and electricity (electrolyte) Shiny when freshly cut or polished Brass Steel
Ionic Crystalline solid at STP High melting and boiling point Usually soluble in water Conducts electricity in aqueous solution Copper sulfate Sodium chloride
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• Molecules are combinations of two or more atoms.
• Molecular element - if the atoms are all the same
  
• For example oxygen gas is a molecule composed of
  two atoms of oxygen. Since there are two atoms
  the molecule is called a diatomic molecule. (just
  remember the gen's)

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Diatomic molecular elements As the heading
suggest these are elements composed of two (di)
nonmetal  atoms.  We seen these in the last
lesson. (just remember the gen's)  
oxygen oxygen O2
hydrogen hydrogen H2
nitrogen nitrogen N2
the halogens (group 17) fluorine F2
the halogens (group 17) chlorine Cl2
the halogens (group 17) bromine Br2
the halogens (group 17) iodine I2

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Mono-atomic elements If we look at the name of
the heading, mono - means one, so these are the
non-metals that that exist in nature as
individual atoms.  Although these are not
compounds we have included them here because we
will reference them many times. Noble Gases
(group 18)
He Helium
Ne Neon
Ar Argon
Kr Krypton
Xe Xenon
Rn Radon
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Polyatomic molecular elements These are
non-metal elements composed of many (poly) atoms.
 
O3 ozone
S8 sulfur
P4 phosphorus (red)
P10 Phosphorus (white)

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• Compound - a molecule that contains two or more
  different types of atoms or ions.
• Water (H2O) is a compound because it contains
  both Hydrogen and Oxygen, two different types of
  atoms.
• The formula for water (H2O) is a combination of
  symbols and subscripts.
• H and O are the symbols for the two types of
  elements (Hydrogen and Oxygen) found in water.
• The number 2 to the lower right hand corner of
  the symbol for Hydrogen is called a subscript.

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• a molecular formula indicates the total number of
  atoms  in one molecule
• an empirical formula is the simplest whole number
  ratio of atoms in the compound.
•  consider hydrogen peroxide (H2O2 ) as an
  example.
• - the molecular formula is H2O2.
• - the empirical formula is HO. (lowest ratio is
  11)
• In some cases the molecular formula and empirical
  formula are the same.
• For example, both the molecular formula and the
  empirical formula for water is H2O. (the lowest
  whole number ratio is 21)
• It is important to recognize however that the
  empirical formula only describes the ratio of one
  atom to another, and not the actual number of
  atoms of each type in the compound.

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Solutions

• Aqueous solutions are those in which the solvent
  is water. They form in at least three ways
  depending on the nature of the solute
• molecular solvation
• dissociation
• ionization

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Electrolytes and Non-electrolytes

• A solution that conducts electrical current is
  said to be electrolytic and the solute is called
  an electrolyte.
• The solute in a solution that does not conduct
  electrical current is a non-electrolyte.
• Generally, dissociated ionic compounds are
  electrolytes whereas dissolved molecular
  compounds are non-electrolytes. The exceptions to
  this rule are the molecular acids. 
• Acids are defined theoretically as species that
  ionize in water to produce hydrogen ions and
  negative ions (anions).

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Bohrs Electron Energy Level Theory

• Bohr was the first to attempt to describe the
  way electrons are distributed in an atom
• The type of bonding that occurs in a substance is
  a function of the way the electrons are
  distributed in an atom

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Electron Energy Level Theory

• Electrons orbit the nucleus in specific pathways
  called energy levels
• There is a fixed number of energy levels
• Each energy level is capable of holding a certain
  number of electrons
• The location of an electron (with respect to the
  nucleus) is an indication of the energy it
  contains (closer more further away less)

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• Electrons can move from one energy level to the
  next by gaining or losing a specific amount of
  energy called a quantum
• Energy levels are not equidistant from each other

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• The part of the energy level or region of space
  in which we find electrons are also called
  orbitals. These are not the fixed pathways we
  normally think of when discussing orbits, but
  rather they are a region or area when the
  electron would be located at some point in time.

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Electron Energy Level Diagrams

• These will demonstrate how electrons are
  distributed in an atom
• Determine the atomic number of the element
• Draw a circle to represent the nucleus and write
  the of protons in the circle
• Determine the period number of the element (ie
  what row it is in). This is equal to the number
  energy levels in the element

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• Each energy level can hold a specific number of
  electrons. We always fill the inner level first.
• For the first three levels it is 2 -8 8. We
  put two hash marks - - for each level and put
  the number of electrons in the level between the
  marks.
• The outermost or last energy level is called the
  valence level.
• 7. The electrons in the valence level are called
  valence electrons.

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Example

• Look up the atomic number of nitrogen and use it
  to determine nitrogen's number of protons and
  electrons. Since nitrogen is atomic number 7, it
  has 7 protons and 7 electrons.
• Draw a circle to represent the nucleus and write
  in the number of protons using the shorthand 7p.
  
• Locate the period number for nitrogen and use it
  to determine the number of electron energy
  levels. The period number (2) equals the number
  of energy levels. Draw two broken lines above the
  circle to represent nitrogen's two electron
  energy levels.

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Bohr theory, the first three energy levels hold a maximum of two, eight, and eight electrons respectively. The energy levels are filled from the inside out, so write "2e-" on the first line to show the number of electrons in the first energy level.            
The outermost energy level of an atom is called the valence level. In a nitrogen atom, the second energy level is the valence level. Since a nitrogen atom has seven electrons, and since two of them are in the first energy level, then the second energy level holds five electrons. Write 5e- in the second energy level. Note that the number of valence electrons in an atom can be determined by looking at the last digit of its group number. Nitrogen is in Group 15.
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Important Point!!!

• Valence electrons are the only ones thought to be
  involved in chemical reactions!

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Your Turn

• Draw energy level diagrams for the following
  elements
• Lithium
• Calcium
• Neon
• Hydrogen

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-- 2 -- 8 -- 8 -- 2 --

• - 1 --
• -- 2 --

-- 8 -- 2 --
-- 1 --
3
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Bohr Diagrams
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Order of Filling Energy Levels

• Lower energy levels are filled before higher
  energy level orbitals
• The electrons will fill the available orbitals
  before pairing up.

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Your Turn

• Construct modified Bohr diagrams for the first
  20 elements

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Why do certain atoms gain or lose electrons?

• The loss or gain of electrons enables atoms to
  achieve an octet of electrons (ie. gain
  electrons) or expose a lower energy octet of
  electrons (ie. lose electrons)

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Nobel Gases and the Octet Rule

• Remember that the outer shell or energy level is
  called the valence level and it contains the
  valence electrons
• It is commonly believed that the valence energy
  level and the valence electrons are responsible
  for chemical bonding
• A group or eight is called an octet
• The octet rule states that chemical stability is
  associated with a group of eight valence
  electrons
• The noble gasses contain a complete outer shell
  of electrons or they have 8 valence electrons
• The noble gasses are believed to be chemically
  stable because their valence level have a
  complement of eight electrons

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Simple Ions

• Simple ions are atoms that no longer have equal
  numbers of electrons and protons.
• As a result of this imbalance they now have a
  positive or negative charge
• Cation
• A cation is  a positve () ion
•  Metal atoms lose electrons to become positive
  thus, metals form cations. 
• Example  Na, Mg 2
• Anion
• An anion is A Negative ION (A N ion)
• Non-metal atoms gain electrons to become more
  negative thus, non-metals form anions 
• Example   F- , O 2-

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Energy level diagrams for IONS Energy level diagrams for IONS
Atoms with 1, 2, or 3 valence electrons lose them to form 1, 2 or 3 ions respectively. naming metallic ions - the full name of the atom is followed by the word ion. Mg2  is the magnesium ion Group 1 (1) (lose 1e) Group 2 (2) (lose 2e) Group 13 (3) (lose 3e)
Atoms with 5, 6, or 7 valence electrons gain electrons to form 3-, 2-, or 1- ions, respectively. naming non-metallic ions - the name of the atom is shortened and the suffix -ide is added. O2-  is oxide Group 15 (3-) (gain 3 e) Group 16 (2-) (gain 2 e) Group 17 (1-) (gain 1 e)
Atoms with valency of 4 generally do not form ions.  These atoms do not gain or lose electrons.  They become stable by sharing electrons - (recall molecular compounds - covalent bonding). Group 14 (do not form ions)
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Your Turn

• Construct modified Bohr diagrams for the ions
  formed by the first 20 elements and name each ion
  formed
• Complete the simple ions sheet

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Your Turn

• Complete sheet u3 s2 l3 Simple Ions

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Chemical Bonding

• A chemical bond will form between two atoms if
  the attractive forces between two atoms is
  greater than the repulsive forces
• Attractive forces are electrical forces that hold
  the atoms, ions, or molecules together
• Bonds are formed through the valence electrons in
  the atom
• Valence electrons are usually transferred from
  the outer shell of one atom to the outer shell of
  another atom or shared among the outer shell of
  combining atoms

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Ionic vs Molecular Compounds

• Ionic Compounds
• Involve the transfer of electrons resulting in
  ionic bonding
• Made up of two oppositely charged ions (metal and
  nonmetal)
• Exist in the form or an ionic crystal lattice
• Binary Molecular Compounds
• Involve the sharing of electrons resulting in
  covalent bonding
• Composed of two nonmetals
• Exist as individual molecules

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Molecular Compounds

• Always form between two or more non-metals
• In order to get full outer energy levels the
  atoms that make up molecular compounds share
  their valence electrons (bonding electrons)
• If they dissolve in water, their solutions do not
  conduct electricity
• Molecular compounds are non - electrolytes

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Molecular Compounds (contd)

• Individual parts of a molecular compound are
  called molecules.
• The molecular formula tell us how many atoms of
  which elements are being shared.
• Example H2O tells us that in every molecule of
  water there are two hydrogen atoms and one oxygen
  atom

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Naming Molecular Compounds (contd)

• When we want to come up with (derive/make) a
  formula for a molecular compound there are two
  routes to follow as well
• 1. It has a trivial name and its formula must
  have been memorized, like water
• 2. We use the prefixes given to tell us how many
  of each element, like silicon dioxide

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Naming Molecular Compounds

• Two basic Ways
• 1. the compound has a trivial (common) name
  that has been used for so long, its too late to
  change it. We must memorize trivial names and
  their formulas.

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Trivial names, formulas, and Special Situations

• The following names and formulas must be
  memorized
• Water H20
• Ammonia NH3
• Glucose C6H12O6
• Sucrose C12H22O11
• Methane CH4
• Propane C3H8
• Octane C8H18
• Methanol CH3OH
• Ethanol C2H5OH
• Hydrogen Peroxide H2O2
• O3 ozone

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Naming Molecular Compounds (contd)

• 2. We use a system of prefixes to signify how
  many of each type of atom is being shared to make
  up the molecules
• 1 mono 6 hexa
• 2 di 7 hepta
• 3 tri 8 octa
• 4 tetra 9 nona
• 5 penta 10 deca

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Naming Molecular Compounds (contd)

• First element in the compound is the regular
  element name
• The last element in the molecule must have its
  name changed to an ide ending
• The only time we are allowed to leave off a
  prefix is if there is only one atom of the first
  element in the molecule

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Naming Molecular Compounds (contd)

• Example
• C02 - carbon dioxide
• CCl4 carbon tetrachloride
• P2O5 diphosphorus pentaoxide

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Your Turn

• Complete sheet u3 s2 l2 Molecular Compounds

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Ionic Bonding

• An ionic bond is a force of attraction between a
  metallic ion and a non-metallic ion
• It is an electron transfer between two elements
  or polyatomic ions
• This type of bond created a crystal with a
  definite, repeated pattern

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• Consider that chemical reactions between metals
  and non-metals occur because metals lose their
  loosely held electrons to highly electronegative
  non-metals.
• This results in ions of the opposite charge
• These oppositely charged ions them mutually
  attract each other to form ionic compounds.
• This attraction between oppositely charged ions
  is called an ionic bond.

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Chlorine atom
Sodium atom
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Writing formulas for ionic compounds

• Write the symbols and charges for the ions that
  make up your ionic compound
• Remember
• The positive ion (cation) is always written
  first, the negative ion (anion) is last
• Almost always your ionic compound will start with
  a metal. (The only time it doesnt is when it
  starts with the word ammonia
• Metal ions has the same name as the metal
  elements
• Non-mental ions always end in ide

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Writing Formulas for Ionic Compounds (Contd)

• Short cut criss-cross the charge and put in
  lowest terms
• Put the two on the charge of calcium as the
  number of chlorines you need, and the one on the
  charge of chlorine as the number of calcium's
  youll need.

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Writing IUPAC names from formulas

• Since every ionic compound is made up of some
  cation in combination with some anion all you do
  is name the cation (positive ion) and then name
  the anion (the one with the negative charge).
• Dont use prefixes
• NaCl sodium chloride
• Li3N lithium nitride

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Writing formulas for ionic compounds (contd)

• Determine the lowest whole number ration of ions
  to give an overall net charge of zero. In other
  words, cancel each others charges.
• Example calcium chloride
• Calcium is the cation, its symbol and charge is
  Ca2
• Chloride is the anion, its symbol and charge is
  Cl-
• It would take two chlorides to cancel the charge
  of one calcium, so the formula is CaCl2

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Your Turn

• Complete sheet u3 s2 l4 Molecular and Binary Ionic

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Writing Formulas for Ionic Compounds (Contd)

• .
• Multi-valent metal ions is the stock system for
  naming. The stock system says that roman number
  in brackets after the metal name is the number of
  the positive charge.
• Example Iron (III) oxide is made up of Fe3
  ions in combination with oxide ions

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Writing IUPAC Names From Formulas (Contd)

• Some ions require special treatment
• Multi valent ions some metals, the ones in B
  groups they can form more that one positive ion
• Example
• Fe can be Fe2 or Fe3
• We have to say which ion was used to make the
  compound by using the Roman numeral for the
  number on the charge in brackets after the name
  of the element.
• Cu2S is copper (II) sulfide
• We know this because it took two copper ions to
  balance the charge on the sulfide ion which is
  always S2-
• PbO2 is lead (IV) oxide. Lead can be a 4 ion or
  a 2 ion. In this case it must be 4 because it
  took two O2- to cancel out the charges

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Example

• Compound Iron(III) oxide
• Ions present Fe3 O2-
• Molecular formula Fe2O3
• Note of oxygen's is same as charge on iron and
  of irons is same as the charge on oxygen

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Writing ionic formulas from names

• First write down the name of the metal
• Then write down the name of the non-metal
• Next determine the charge on the non-metal
• To find the overall negative charge multiple the
  charge by the number of atoms present
• To determine the charge on the metal divide the
  total negative charge by the number of atoms
  present in the molecule

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Example

• RuO2
• Name of metal ? ruthenium
• Name of non metal ? oxygen
• Charge of oxide ion ? 2-
• Number of oxygens present ? 2
• Total negative charge ? 2 x 2 4
• Number of metal atoms present ? 1
• Charge of metal ? 4 divided by the number of
  metal atoms (1) 4
• Name of compound ? ruthenium (IV) oxide

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Example 2

• Nb2O5
• Name of metal ? niobium
• Name of non metal ? oxygen
• Charge of oxide ion ? 2-
• Number of oxygens present ? 5
• Total negative charge ? 2 x 5 10
• Number of metal atoms present ? 2
• Charge of metal ? 10 divided by the number of
  metal atoms (2) 5
• Name of compound ? niobium (V) oxide

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Your Turn

• Complete sheet u3 s2 l4
• Nomenclature Involving Multivalent Ions

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Polyatomic Ions

• Are composed of several elements (more than one
  capital letter) bonded together with an overall
  charge.
• The charge is almost always negative
• You should recognize these special ions

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Writing IUPAC Names From Formulas (Contd)

• We have to recognize the polyatomic ions.
• Once we have identified the cation, and the rest
  of the formula still has more than one capital
  letter in it , there must be a polyatomic ion in
  it. The polyatomic ions all have special names.
• Example
• Na2CO3 is called sodium carbonate. Na is
  positive part (cation) so that means that the CO3
  part has a special name called carbonate

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Writing IUPAC Names From Formulas (Contd)

• Al(OH)3 is aluminum hydroxide hydroxide us the
  name given to the polyatomic ion made up of O and
  H joined together as OH-

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Example 1

• Compound Sodium carbonate
• Ions present Na CO32-
• Molecular formula Na2C03
• Note of sodium's is same as charge on
  carbonate ion and of carbonate ions is same as
  the charge on sodium

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Example 2

• Compound Ammonium sulfide
• Ions present NH4 S2-
• Molecular formula (NH4)2S
• Note of ammoniums is same as charge on sulfate
  and of sulfates is same as the charge on
  ammonium

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Your Turn

• Complete Sheet u3 s2 l5
• Polyatomic Ions (two sheets)

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Special Ionic Compounds

• Some ionic compounds have a number of water
  molecules hanging around with them.
• These special compounds are called hydrates.
• A hydrate or formula has two parts the first
  part is the ionic compound, the second part is
  the water. We separate the two parts with a
  giant asterisk or dot.

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Special Ionic Compounds (Contd)

• CuSO4 5H2O is named in two steps
• 1. First name the ionic compound in the regular
  way (Copper (II) Sulfate)
• 2. Next use your molecular prefixes along with
  the word hydrate to tell how many water molecules
  there are. (pentahydrate)

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Deriving Formulas

• 1. First come up with the formula for the ionic
  compound the usual way by balancing the charges
  on the ion.
• 2. Check the prefix to find out how many water
  molecules are with the ionic compound

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Deriving Formulas

• Example
• magnesium sulfate heptahydrate
• Mg2 with SO42- is MgSO4
• heptahydrate means 7 water
• Full formula MgSO4 7H2O

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Your Turn

• Complete sheet u3 s2 l5 Hydrates

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Nomenclature of Acids

• Acids are hydrogen compounds, that release
  hydrogen ions (H) when they are dissolved in
  water. The formula of an acid is always followed
  by (aq), which means in aqueous solution in
  other words dissolved in water

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Properties of Acids

• There are several key properties that distinguish
  acids from other classes of compounds.
• acids turn blue litmus paper red 
• acids react with metals like zinc, magnesium, and
  iron to produce hydrogen gas. (Hydrogen gas is
  flammable.) 
• acids neutralize bases 
• acids have low pH (pH lt 7) 
• acids that are edible, such as those found in
  foods, taste sour (e.g. vinegar, and citric acid
  found in oranges, lemons, and other citrus
  fruits, sour candy.) 

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Naming Acids

• Naming Acids.
• Look at the formula and name the acid as a
  hydrogen compound. It will be named hydrogen
  ___________. The blank represents some ion.
• The names of hydrogen compounds can only have
  three possible ending
• hydrogen ______ ide.
• hydrogen ______ ate.
• Hydrogen ______ ite.

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Naming Acids (contd)

• The acid name depends on which ion ending you
  have in the formula
• Hydrogen __ ide becomes hydro ___ ic acid
• Hydrogen __ ate becomes ___ ic acid
• Hydrogen __ ite becomes ___ ous acid

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Naming Acids (contd)

• HF(aq) would be called hydrogen flouride so it
  would be called hydrofluoric acid
• HClO4(aq) would be hydrogen perchlorate so it
  would be called perchloric acid
• NNO2(aq) would be called hydrogen nitrite, so it
  would be called nitrous acid

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Writing Formulas for Acids

• 1. Look at the name and work backwards to change
  it to a hydrogen compound.
• Hydro __ ic acid becomes hydrogen ___ ide.
• ______ ic acid becomes hydrogen ____ ate.
• ______ous acid becomes hydrogen __ ite.
• 2. Make your formula by balancing the charges.
  The hydrogen in acids is always 1 and the charge
  on the other ions can be determined from the
  periodic table or the polyatomic chart.

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Writing Formulas for Acids (Contd)

• Example
• Hydrochloric acid must be hydrogen chloride
• Hydrogen is 1 and chlorine is 1- so the formula
  is HCl(aq)
• Nitric acid must be hydrogen nitrate. If
  hydrogen is 1 and nitrate is 1- so nitric acid
  is HNO3 (aq)
• Chlorous acid must be hydrogen chlorite. If
  hydrogen is 1 and chlorite is 1- then the
  formula is HClO2 (aq)

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Special Situations and Exceptions to the Rules

• Any acid names that contain the root word sulf
  or phosph, add an extra syllable to make them
  sound better
• H2SO4 (aq) should be called sulfic acid
  according to our rules but we call it sulfuric
  acid,l with an extra ur syllable
• H3PO4 (aq) should be called phosphic acid but we
  call it phosphoric acid with an extra or in it

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Special Situations (contd)

• The other exception is when you put together a
  formula with certain polyatomic ions that end in
  ____COO. For these acids the hydrogen gets put
  at the end of the formula rather than at the
  beginning. It is still balanced with the
  negative charge, but is put at the end of the
  formula.

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Special Situations (contd)

• Example
• Acetic acid from hydrogen acetate
• Hydrogen is 1
• Acetate ion is 1-
• To balance the charges you need one of each, but
  the hydrogen has to go at the end
• CH3COOH (aq)

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Your Turn

• Complete sheet u3 s2 l7 Acids (2 pages)
• Complets sheet u3 s2 l2 l4 (2 pages)

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Bases

• Bases are substances that behave in opposition to
  acids. In this course, we will restrict our
  discussion of bases to one particular type -
  ionic compounds that contain the hydroxide ion
  (OH-).
• example  Sodium hydroxide - NaOH(aq)

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Properties of Bases

• turn red litmus blue
• neutralize acids
• have high pH (gt 7)
• form slippery solutions
• tend to have a bitter taste
• The pH scale is a measure of the concentration of
  hydrogen ions (H) in aqueous solution, and can
  be used to compare the strength of acid or base.
   The pH scale ranges from pH 1 (strong acid),
  to pH 14 (strong base). The value pH 7 is
  neutral (neither acid nor base)   

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Neutralization Reactions and the Formation of
Salts

• Salts are formed as a result of the reaction
  between an acid and a base.
• The reaction between an acid and a base results
  in the formation of salt and water. The type of
  salt formed would be determined by the type of
  acid and base that react in the neutralization
  reaction.
• ACID BASE ? SALT WATER

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Chemical Reactions

• A chemical reaction is the result of chemical
  species colliding.
• If the collision is successful bonds will be
  broken and new bonds formed.
• This will result in a new chemical species being
  formed.
• Energy is either required (endothermic) or
  released (exothermic) during these reactions.

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Evidence of a Chemical Reaction

• Energy change
• Color change
• Gas evolved
• Precipitate formed (a solid)
• Odor change (never smell directly)
• Every reaction must obey the law of conservation
  of mass matter cannot be created nor destroyed
  only transformed

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5 Basic Reaction Types

• Formation
• Decomposition
• Single replacement
• Double replacement
• Combustion

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Formation

• A formation reaction involves creating a compound
  from the elements it contains
• General formula
• Element element ? compound
• A B ? AB
• Ex.
• Li (s) O2(g) ? LiO2 (g)
• K (s) S8 ? K2S (s)
• F2 (g) Al (s) ? AlF3 (s)

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Formation (contd)

• To correctly predict the product write the metal
  and non-metal ions together and balance the
  charges (ionic compounds)
• NO subscripts from the reactant side are brought
  to the product side
• Ex.
• Na (s) P4 (s) ? Na3P(s)
• Ba (s) O2 (g) ? BaO(s)
• Cl2 (g) K (s) ? KCl(s)

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Formation (contd)

• Sometimes the metal that reacts form a
  multi-valent ion in the ionic compound.
• To determine which ion should be used, use the
  ions that is on top on the periodic table
• Ex.
• V (s) O2 (g) ? V2O5 (s)
• Mn (s) F2 (g) ? MnF2 (s)

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Decomposition

• Involves particles of the same compound colliding
  and forming the elements found in the compound
• Compound ? element 1 element 2 .
• ABC ? A B C
• Example
• C6H12O6 ? C (s) H2 (g) O2 (g)
• VBr5 (s) ? V (s) Br2 (l)

99
Decomposition (contd)

• To correctly predict results
• Write each element that you identify in the
  reactant compound as an element.
• Do not bring any subscripts with the element.
• Write each element as it would normally exist.
• It does not matter what order you write the
  elements.

100
Decomposition (contd)

• U2O 4 (s) ? U (s) O2 (g)
• AlP (s) ? Al (s) P4 (s)

101
Decomposition (contd)

• If the compound contains polyatomic ions, follow
  the same procedure.
• Al2(SO4)3 (s) ? Al (s) S8 (s) O2 (g)
• Examples
• C4H10 (g) ?
• (NH4)2Cr2O7 (s) ?

102
Single Replacement

• Is a reaction that is a result of an element
  colliding with a compound
• The element will switch or replace the
  appropriate element in the compound
• Two possibilities
• a) A BC ? B AC (A is a metal)
• b) A BC ? C BA ( A is a non-metal)

103
Single Replacement (contd)

• Example
• Al (s) BaCl2 (s) ? Ba (s) AlCl3 (s)
• P4 (s) BaCl2 (s) ? Cl2 (g) Ba3P2 (s)

104
Single Replacement (contd)

• The ionic bonds in the compound were broken
• The element then bonds with the oppositely
  charged ion from the ionic compound
• The other ion forms an element
• The new ionic compound must be correctly written
  (criss cross method)
• No subscripts are moved across the equation
• In does not matter the order in which the
  chemical species are written

105
Single Replacement (contd)

• K (s) AlP (s) ? (metal with metal)
• NaF (s) Al (s) ? (metal with metal)
• O2 (g) LiBr (s) ? (nonmetal with nonmetal)
• AgI (s) S8 (s) ? (nonmetal with
  nonmetal)

106
Single Replacement (contd)

• Example
• If the element is a metal you must check to see
  if it is multi-valent and use the most abundant
  charge (the top one)
• U (s) LiF (s) ? Li (s) UF6 (s)
• Ba (s) MnCl4 (s) ? Mn (s) BaCl2 (s)
• If opposite, the metal to be replaced in the
  compound is multi-valent, it does not matter

107
Single Replacement (contd)

• If the compound contains a polyatomic ion, the
  polyatomic ion will not change since it remains
  part of the compound.
• Example
• Ca (s) NaOH (s) ? Na (s) Ca(OH)2 (s)
• Al (s) CuSO4 (s) ? Cu (s) Al2(SO4)3 (s)
• S8 (s) NH4F (s) ? F2 (g) (NH4)2S (s)

108
Double Replacement

• Is the result of two compounds colliding
• Some/all bonds are broken and new bonds are
  formed
• General Form
• compound1 compound 2 ? compound 3 compound
  4
• AB CD ? AD CB
• note A and C were metals
• B and D were nonmetals
• It does not matter which order the reactants
  or products are written in

109
Double Replacement (contd)

• This type of reaction is very similar to a single
  replacement except there is a double switch.
• Simply replace the metal in one compound with the
  metal in the second compound
• Since each product is an ionic compound, balance
  the charges to find the correct products

110
Double Replacement (contd)

• Examples
• BaF2 (aq) AlP ? Ba3P2 (s) AlF3 (aq)
• Na2O (s) Mg(NO3)2 (aq) ? NaNO3 (aq) MgO (s)
• NH4I (aq) Al2(SO4)3 (aq) ? (NH4)2SO4 (aq)
  AlI3 (s)

111
Double Replacement (contd)

• UF6 (s) KMnO4 (aq) ?
• LiCN (aq) BaBr2 (aq) ?
• Ca(NO3)2 (aq) FeSO4 (aq) ?

• U(MnO4)6 (aq) KF (aq)
• LiBr (aq) Ba(CN)2 (aq)
• CaSO4 (s) Fe(NO3)2(aq)

112
Double Replacement (contd)

• A special case of double replacement reactions
  involves the addition of a base with an acid
• This is called a neutralization reaction
• A salt and water are produced
• Example
• HCl(aq) LiOH(aq) ? LiCl(aq) H2O (l)
• Al(OH)3 (aq) H2SO4 (aq) ? Al2(SO4)3 (aq)
  H2O (l)

113
Combustion

• Involves the collision of oxygen with a
  hydrocarbon
• Two types
• A) complete and B) incomplete

114
Complete Combustion

• The reaction has the same products no matter what
  the reactants
• Products always carbon dioxide and water
• Example
• CH4 (g) O2 (g) ? CO2 (g) H20 (g)
• C4H10 (g) O2 (g) ? CO2 (g) H20 (g)
• C6H12O2 (g) O2 (g) ? CO2 (g) H20 (g)

115
Incomplete Combustion

• Same as complete combustion, but not all of the
  reactant is burnt
• This results in two additional products carbon
  and carbon monoxide
• Example
• CH4 (g) O2 (g) ? CO2 (g) H20 (g ) C (s)
  CO (g)
• C4H10 (g) O2 (g) ? CO2 (g) H20 (g) C (s)
  CO(g)
• C6H12O2 (g) O2 (g) ? CO2 (g) H20 (g) C
  (s) CO (g)
• note you will have to be told if the reaction
  is incomplete to know the products

116
Combustion (contd)

• Examples
• CH4 (g) O2 (g) ?
• C20H42 (g) O2 (g) ?)
• C12H22O11 (g) O2 (g) ?

117
Balancing Chemical Equations

• The law of conservation of mass states that
  matter is conserved during a chemical reaction
• Therefore we are required to balance chemical
  equation, we need to have the same amount of
  products and reactants
• To balance chemical equations we use molar
  co-efficients to ensure that we have the same
  number and type of atom on each side of the
  equation
• Most molar co-efficients are simply numbers
  written in front of each reagent and product
  that, when multiplied, result in equal numbers of
  atoms on both sides of the equation

118
Balancing Chemical Equations

1. Start with the largest subscript present (not the
   ones found within the polyatomic ions)
2. Balance the polyatomic ions as a group
3. Balance in a logical progression
4. Leave the elements until last
5. If it seems hard to balance, you may have made a
   mistake in the formulas of your chemical species

119
Balancing Chemical Equations

• Example
• F2 (g) Al (s) ? AlF3 (s)
• The largest subscript is 3
• Compare the number of Fs in the products vs the
  reactants (3 vs 2)
• These numbers are not divisible, therefore, find
  the LCM
• LCM 6
• Determine the number of times that AlF3 goes into
  the LCM
• Place the number 2 in front of AlF3
• Now compare the number of Fs on both sides of
  the equation. F2 would require the number 3 to be
  placed in front of it to make the number of Fs
  equal on both sides (6)
• Next balance the number of Al s. Since there are
  two on the product side, multiple the one on the
  reactant side by 2.
• The balanced equation would be
• 3 F2 (g) 2 Al (s) ? 2 AlF3 (s)

120
Complete the following equations

• Na Cl2 ?
• K O2 ?
• H2 O2 ?
• H2 Cl2 ?
• N2 H2 ?
• N2H4 O2 ?
• CH4 O2 ?

121
Complete the following equations

• Ca HBr ?
• Al O2 ?
• KNO3 HBr ?
• Ba H3PO4 ?
• CaCl2 AL2(SO4)3 ?
• C3H8 O2 ?
• Mg HNO3 ?
• AgNO3 NaCl ?
• K3PO4 CoCl3 ?