AP Chapter 2

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AP Chapter 2

• Chemical Equations and Reactions

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Symbols

• Ag
• C

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Formulas

• H2O
• Ba(C2H3O2)2

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Empirical and Molecular Formulas
Substance Molecular Formula Empirical Formula
Water H2O H2O
Benzene C6H6 CH
Acetylene C2H2 CH
Glucose C6H12O6 CH2O
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Atomic mass unit

• The atomic mass unit (u), is a unit of mass used
  to express atomic and molecular masses.
• It is the approximate mass of a hydrogen atom, a
  proton, or a neutron.
• By definition the atomic mass unit is equal to
  one-twelfth of the mass of a carbon-12 atom.

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The Atom

• The nucleus is very small, dense, and positively
  charged.
• Electrons surround the nucleus.
• Most of the atom is empty space

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Subatomic Particles
PARTICLE SYMBOL CHARGE MASS (u) LOCATION
electron e- -1 ?0 orbit nucleus
proton p 1 ?1 inside nucleus
neutron n0 0 ?1 inside nucleus
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Atomic Number (Z)

• The number of protons in the nucleus of an atom.
• The identifying characteristic of an element.

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Mass Number

• The sum of the protons and neutrons in the
  nucleus of an atom.

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Isotopes
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Isotopes

• This is the symbol for carbon-12.
• Atomic number is 6.
• Mass number is 12.

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Isotopes

• This is the symbol for carbon-12.
• Atomic number is 6.
• Mass number is 12.
• Write the symbols for carbon-13 and carbon-14.

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Isotopes
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What is the average mass of a carbon atom?
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What is the average mass of a carbon atom?

• 12.01

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Atomic Mass

• The atomic mass of carbon is 12.01u.

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Atomic Mass

• The atomic mass of carbon is 12.01u.
• Atomic mass is the average mass of all the
  isotopes of an atom. It takes into account the
  different isotopes of an element and their
  relative abundance.

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• How many electrons, protons and neutrons are in
  an atom of actinium with a mass number of 221?

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• How many electrons, protons and neutrons are in
  an atom of actinium with a mass number of 221?

• 89p
• 89e-
• 132n0

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• How many electrons, protons and neutrons are in
  an atom of rhodium-105?

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• How many electrons, protons and neutrons are in
  an atom of rhodium-105?

• 45p
• 45e-
• 60n0

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Families of the Periodic Table
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The Noble Gases

• Elements in group 18
• All are gases.
• VERY non-reactive.
• Have a full outer energy level.

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The Octet Rule

• The octet rule states that an element's outer
  energy level is full and most stable when it
  contains eight electrons.
• This stability is the reason that the noble gases
  are so non-reactive.

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Exception to the Octet Rule

• The first energy level can only hold two
  electrons and so elements such as Hydrogen and
  Helium that only have one energy level follow a
  duet rule.

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Ion Vocabulary

• An ion is an atom or group of atoms that have a
  charge.
• A monatomic ion is an atom with a charge.
• The charge on the atom is called an oxidation
  number.
• A polyatomic ion is a group of atoms with a
  charge.
• A cation is a positive ion.
• An anion is a negative ion.

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An ion is an atom, or group of atoms, that has a
net positive or negative charge.
cation ion with a positive charge If a neutral
atom loses one or more electrons it becomes a
cation.
anion ion with a negative charge If a neutral
atom gains one or more electrons it becomes an
anion.
2.5
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Ionic

• When an element that easily loses electrons (a
  metal) reacts with an element that easily gains
  electrons (a nonmetal), one or more electrons are
  transferred.
• This creates two ions which are held together by
  an ionic bond.
• A compound that contains ions is called an ionic
  compound.

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a
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Formula Unit
A formula unit is the empirical formula of an
ionic compound. It is the lowest whole number
ratio of ions represented in an ionic compound.
Examples include ionic NaCl and K2O. Ionic
compounds do not exist as individual molecules a
formula unit thus indicates the lowest reduced
ratio of ions in the compound.
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Covalent

• When atoms share electrons the bond created is
  said to be covalent. Covalent bonds often form
  between nonmetal atoms.
• These covalently bonded atoms act as single units
  called molecules.
• A compound made up of molecules is a covalent
  compound.

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Polyatomic Ions
Note that these are charges and not oxidation
numbers.
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Types of monatomic ions and the rules for naming
them

• The periodic table is useful in naming the
  monatomic ions.

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Monatomic cations with one oxidation number

• The cations from the periodic table which have a
  single oxidation number are as follows Group 1
  (1), Group 2 (2), Ag, Cd2, Zn2, and Al3.
• These types of ions are named by using the name
  of the element followed by the word ion.
• Na sodium ion
• Ba2 barium ion
• Zn2 zinc ion
• We can use the roman numeral from the periodic
  table to identify the oxidation number for these
  ions.

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Monatomic cations with multiple oxidation numbers

• All other cations that are not listed in the
  previous category (cations with one oxidation
  number) are considered to have the possibility of
  multiple oxidation numbers.
• These type of ions are named by using the name of
  the element followed by a Roman numeral to
  indicate the oxidation number.
• Cu2 copper (II)
• Pb4 lead (IV)
• Mn7 manganese (VII)

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Monatomic anions

• All anions from the periodic table are named by
  changing the ending of the elements name to
  ide.
• F- fluoride ion
• O2- oxide ion
• N3- nitride ion
• Count back from the noble gases starting at zero
  to determine the oxidation number.

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Name these monatomic ions

• Rb
• P3-
• Fe3
• Br?
• Mn4
• Cd2

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Write the formula for these monatomic ions.

• Sulfide
• Lead (II)
• Barium ion
• Chromium (IV)
• Aluminum ion
• Carbide

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Formulas of Ionic Compounds

• Formulas for ionic compounds can be written by
  the following steps

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Formulas of Ionic Compounds

• Formulas for ionic compounds can be written by
  the following steps
• (1) Write the formula for the cation and anion
  (Dont forget to include the charge of each ion).

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Formulas of Ionic Compounds

• Formulas for ionic compounds can be written by
  the following steps
• (1) Write the formula for the cation and anion
  (Dont forget to include the charge of each ion).
• (2) Decide how many cations and anions are needed
  so that the sum of their charges balances out to
  be zero.

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Formulas of Ionic Compounds

• Formulas for ionic compounds can be written by
  the following steps
• (1) Write the formula for the cation and anion
  (Dont forget to include the charge of each ion).
• (2) Decide how many cations and anions are needed
  so that the sum of their charges balances out to
  be zero.
• (3) Write the formula of the compound by writing
  the number of cations followed by the number of
  anions which you used in step 2. Remember not
  to include the charges of the ions since now they
  balance out to be neutral. (Note when using more
  than one polyatomic ion the polyatomic ion must
  be written in parentheses).

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Write the formula for barium chloride
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Write the formula for iron (II) oxide
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Write the formula for calcium phosphate
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Write the formula for ammonium carbonate
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Write the formulas for the following compounds

• cobalt (II) chloride
• lithium sulfate
• ammonium dichromate
• (d) aluminum oxide
• (e) boron (III) phosphide
• (f) Chromium (V) nitrate

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Naming Ionic Compounds

• When naming ionic compounds the following steps
  are followed

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Naming Ionic Compounds

• When naming ionic compounds the following steps
  are followed
• (1) Separate the compound into its positive and
  negative parts (Note that the positive part of a
  compound will be only the first element with the
  exception of ammonium which is NH4)

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Naming Ionic Compounds

• When naming ionic compounds the following steps
  are followed
• (1) Separate the compound into its positive and
  negative parts (Note that the positive part of a
  compound will be only the first element with the
  exception of ammonium which is NH4)
• (2) Write the name of the cation followed by the
  name of the anion.

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Write the name of ZnO and determine the oxidation
numbers of the elements within this compound.
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Write the name of CuO and determine the oxidation
numbers of the elements within this compound.
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Write the name of MnCO3 and determine the
oxidation numbers of the elements within this
compound.
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Write the name of Fe2(SO4)3 and determine the
oxidation numbers of the elements within this
compound.
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Name the following compounds and determine the
oxidation numbers of each element.

• SrCl2
• (b) Cr(OH)2
• (c) KClO4
• (d) NH4MnO4
• (e) CuP

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Binary Molecular Compounds

• Binary molecular compounds are composed of two
  different nonmetals
• examples CO, SO2, N2H4, P4Cl10
• These compounds are named by using a prefix to
  indicate the number of atoms of each element
  present.

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• The prefix mono- is often omitted especially when
  the first element would have the prefix mono-
• CO
• (example CO is named carbon monoxide, not
  monocarbon monoxide).

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Name the following compounds NF3 N2O4 P4S10

• NF3 is nitrogen trifluoride
• N2O4 is dinitrogen tetraoxide
• P4S10 is tetraphosphorous decasulfide

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• Write formulas for the following compounds
• dichlorine heptaoxide
• carbon hexasulfide
• octaphosphorous pentaoxide

• dichlorine heptaoxide is Cl2O7
• carbon hexasulfide is CS6
• octaphosphorous pentaoxide is P8O5

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Acids

• Acids are compounds that give off hydrogen ions,
  (H) when dissolved in water. When a compound
  has hydrogen as its cation the substance is
  generally an acid
• Examples HCl, H2SO4, H3PO3
• The rules for naming acids are based on the anion
  portion of the acid formula.

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Rules for Naming Acids

• The names of acids are based on the ending of the
  anion name.
• Examples HCl, H2SO4, H3PO3
• Cl? chloride
• SO42? sulfate
• PO33? phophite

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Rules for Naming Acids
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Name the acids

• HNO2
• HCN
• H3PO4

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Write formulas for the following acids

• chromic acid
• hydroiodic acid
• chlorous acid

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Balanced Chemical Equation

• A chemical equation is a written representation
  of a chemical reaction.
• 2Na 2H2O ? H2 2NaOH
• Reactants
• Products
• Coefficients
• You should be able to balance equations using
  coefficients.

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Symbols Used in Equations

• 2Na(s) 2H2O(l) ? H2(g) 2NaOH(aq)
• Solid
• Liquid
• Gas
• Aqueous solution

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Other Symbols Used in Equations

• Solid (cr) or (s)
• Precipitate (?)
• Heated
• Escaping gas (?)
• Catalyst H2SO4
• A word may be written above an arrow to indicate
  something is necessary for the reaction to occur.

?
electricity
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Net Ionic Equations

• Solutions of sodium chloride and silver nitrate
  are mixed.
• Step 1 Change the word equation into a chemical
  equation by writing the formulas for the
  reactants.
• NaCl AgNO3

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NaCl AgNO3

• Step 2 Classify each reactant as a particular
  type of substance.
• Classification of Substances
• Acids compounds with formulas that begin with
  H. (Examples HCl, H2SO4).
• Bases compounds that end with OH. (Examples
  NaOH, Ba(OH)2)
• Metal Oxides binary compounds of a metal and
  oxygen. (Examples CaO, Na2O).
• Nonmetal Oxides binary compounds of a nonmetal
  and oxygen. (Examples SO2, P4O10).
• Salts Ionic compounds other than bases and
  metal oxides. (Examples NaCl, Mg3(PO4)2,
  NH4NO3).
• Other Compounds All compounds not classified as
  one of the five types above. (Examples CH4,
  NH3).
• NaCl and AgNO3 are salts

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Salt
NaCl
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Salts

• A salt is an ionic compound other than a base or
  oxide.

K2Cr2O7
CuSO4
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NaCl AgNO3(salt salt)

• Step 3 Based on your classification of the
  substances determine the type of reaction.

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NaCl AgNO3(salt salt)

• Step 3 Based on your classification of the
  substances determine the type of reaction.
• This is a double replacement reaction.

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NaCl AgNO3 ?

• Step 4 Predict the products of the reaction
  based on the reaction type.
• NaCl AgNO3 ? NaNO3 AgCl

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NaCl AgNO3 ? NaNO3 AgCl

• Step 5 Use solubility rules if necessary.
• Solutions of sodium chloride and silver nitrate
  are mixed.

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Solubility Rules

• Soluble (strong) acids HCl, HBr, HI, HNO3,
  H2SO4, HClO4
• Soluble (strong) bases LiOH, NaOH, KOH, Ca(OH)2,
  Sr(OH)2, Ba(OH)2
• Soluble salts All salts of lithium, sodium,
  potassium, and ammonium cations. All salts of
  nitrate and acetate anions. All chloride,
  bromide, and iodide salts except silver, lead and
  mercury (I). All sulfates except silver, lead,
  mercury (I), calcium, strontium and barium.

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NaCl AgNO3 ? NaNO3 AgCl

• Step 5 Use solubility rules if necessary.
• Na Cl- Ag NO3- ? Na NO3- AgCl

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Na Cl- Ag NO3- ? Na NO3- AgCl

• Step 6 Eliminate all spectator ions.
• A spectator ion appears as both a reactant and a
  product in a chemical equation.
• Na Cl- Ag NO3- ? Na NO3- AgCl

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• Step 7 Write the final net ionic equation
• Ag Cl- ? AgCl

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Ag Cl- ? AgCl
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The AgCl is a precipitate

• A precipitate is a insoluble solid formed when
  solutions are mixed.
• Precipitates are normally formed by reacting two
  salts or by changing the temperature to affect
  the solubility of a compound within a solution.

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Precipitate
Solutions of lead nitrate and potassium iodide
are mixed.
What is the yellow precipitate?
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Solid calcium phosphate is added to excess
hydrochloric acid.

• Ca3(PO4)2 H ? H3PO4 Ca2

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Equal volumes of 0.1M sulfuric acid and 0.1M
sodium hydroxide are mixed.

• H OH- ? H2O

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Solid barium carbonate is added to an excess of
dilute nitric acid.

• BaCO3 H ? Ba2 H2CO3

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Types of Net Ionic Equations

• 1. Double Replacement Reactions These reactions
  start with two compounds and produce two
  different compounds. Such reactions can be
  expected when the two reactants are some
  combination of acid, base, or salt. The products
  can be predicted by exchanging the positive parts
  of the two reactants.
• If carbonic acid, H2CO3 is produced as a product
  it should be written as H2O and CO2. If ammonium
  hydroxide, NH4OH is produced as a product it
  should be written as NH3 and H2O.

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Solid barium carbonate is added to an excess of
dilute nitric acid.

• BaCO3 H ? Ba2 H2CO3
• BaCO3 H ? Ba2 H2O CO2

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Demo this reaction for students

• BaCO3 H ? Ba2 H2O CO2

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Solid ammonium carbonate is added to a saturated
solution of barium hydroxide.

• (NH4)2CO3 Ba2 OH- ? NH4OH BaCO3
• (NH4)2CO3 Ba2 OH- ? NH3 H2O BaCO3

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Single Replacement Reactions

• 2. The reactants are an element and a compound
  and the products are a different element and
  compound. A metallic element will replace the
  positive part of a compound or a nonmetallic
  element will replace the negative part of a
  compound.

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Teacher Example Calcium metal is added to
dilute nitric acid.
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Aluminum metal is added to a solution of copper
(II) chloride.

• Al Cu2 ? Al3 Cu

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Liquid bromine is added to a solution of
potassium iodide.

• Br2 I- ? Br- I2

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Solid calcium is added to warm water.

• Ca HOH ? H2 Ca2 OH-

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Synthesis (Combination) Reactions

• (a) Two elements combine to form a binary
  compound.
• (b) A metal oxide and water combine to form a
  base.
• (c) A nonmetal oxide and water combine to form
  an acid.
• (d) a metal oxide and a nonmetal oxide combine
  to form a salt.
• In these reactions it is necessary to know the
  charges of certain ions in order to predict the
  formulas of your products. You should determine
  these ion charges by using their charges within
  the reacting substances. If this is impossible
  use your prior experience or the periodic table
  to make a prediction.

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Teacher Example Magnesium metal is heated
strongly in nitrogen gas.
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Teacher Example Calcium oxide is added to
water.
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Teacher Example Dinitrogen trioxide gas is
bubbled through water.
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Calcium metal is heated strongly in nitrogen
gas.

• Ca N2 ? Ca3N2

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Teacher ExampleExcess chlorine gas is passed
over hot iron filings.

• Cl2 Fe ? FeCl3

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A piece of lithium metal is dropped into a
container of nitrogen gas.

• Li N2 ? Li3N

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Solid barium oxide is added to distilled water.

• BaO HOH ? Ba2 OH-

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Lithium oxide powder is added to excess water.

• Li2O H2O ? Li OH-

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Solid dinitrogen pentoxide is added to water.

• N2O5 H2O ? H NO3-

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Phosphorus (V) oxide powder is sprinkled over
distilled water.

• P2O5 HOH ? H3PO4

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Metal oxide Nonmetal oxide

• Solid calcium oxide is exposed to a stream of
  carbon dioxide gas.
• CaO CO2 ? CaCO3

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Solid calcium oxide is heated in the presence of
sulfur trioxide gas.

• CaO SO3 ? CaSO4

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Decomposition Reactions

• These reactions begin with a single compound and
  decompose into more than one product. In general
  they are simply the reverse of the synthesis
  reactions listed in 3 above. There are however
  a few other common decomposition reactions that
  you should learn (a) Hydrogen peroxide, H2O2
  will decompose into water, H2O and oxygen, O2.
• (b) Potassium chlorate, KClO3 will decompose
  into potassium chloride, KCl and oxygen O2.

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Decomposition of hydrogen peroxide.

• H2O2 ? H2O O2

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Decomposition of potassium chlorate.

• KClO3 ? KCl O2

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Here are some decompositionreactions from
previous AP Tests

• Solid calcium sulfite is heated in a vacuum.
• A solution of hydrogen peroxide is exposed to an
  iron catalyst.
• Solid potassium chlorate is heated in the
  presence of a manganese dioxide catalyst.

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Complex Ion Formation (Coordination Compounds)

• 5. These reactions involve the reaction of two
  compounds to form a complex ion (coordination
  compound). One of the reacting compounds serves
  as a source of metal ions and the other reacting
  compound serves as a source of ligands. A
  complex ion (coordination compound) is a
  combination of metal ions and ligands. Ligands
  are normally ammonia (NH3), hydroxide (OH-), or
  cyanide (CN-). There needs to be a large supply
  of ligand for a complex ion to form. This is
  normally indicated within a reaction by words
  such as concentrated and/or excess. To form a
  complex ion take the metal ion and add a number
  of ligands which equal twice the metal ions
  oxidation number. For example if the metal ion
  has an oxidation number of 2 you should add four
  of the ligands to it. Then simply add up the
  charges within the complex ion and determine the
  final charge.

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Teacher Example Excess concentrated ammonia is
added to a solution of nickel (II) sulfate.
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An excess of ammonia gas is bubbled through a
solution saturated with silver chloride.

• NH3 AgCl ? Ag(NH3)2 Cl-

What happens to the AgCl in this reaction?
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Excess concentrated potassium hydroxide solution
is added to a precipitate of zinc hydroxide.

• OH- Zn(OH)2 ? Zn(OH)42-

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Excess sodium cyanide solution is added to a
solution of silver nitrate.

• CN- Ag ? Ag(CN)2 -

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Concentrated (15M) ammonia solution is added in
excess to a solution of copper (II) nitrate.

• NH3 Cu2 ? Cu(NH3)42

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Excess concentrated sodium hydroxide solution is
added to solid aluminum hydroxide.

• OH- Al(OH)3 ? Al(OH)63-

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Combustion Reactions
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Combustion

• Hexane is burned in excess oxygen.
• Propanol is burned completely in air.

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Bronsted acid a compound that donates a hydrogen
ion (H) to another compound.

• Bronsted base a compound that accepts a
  hydrogen ion (H) from another compound.

CH3CO2H H2O ? CH3CO2- H3O
H2O NH3 ? OH- NH4
H3O is a hydronium ion.
H2O is amphoteric (it can act like an acid or a
base
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Oxidation-Reduction (Redox) Reaction

• Sodium burns in air.
• Na O2 ? NaO
• In a redox reaction the oxidation numbers of at
  least some of the substances change.

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Reversible Reaction
N2 3H2 ? 2NH3
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Catalyst

• Speeds up a reaction without being permanently
  changed in the process.
• Example enzymes

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Hydrates
CoCl2
CoCl2 6H2O