A Chemist

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A Chemists View of Explosives

• Ionic Bonding and Nomenclature Notes

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I. Chemical bond a mutual electrical attraction
between the nuclei and valence electrons of
different atoms that binds the atoms together.
Another way to describe a chemical bond is to say
the attractive forces between atoms or ions in
compounds. In ionic compounds it is an attractive
force between positive and negative ions.

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• In ionic bonding valence electrons are actually
  TRANSFERRED between a nonmetal and metal. This
  happens because a non-metallic atom is much more
  electronegative and it can pull electrons away
  from the less electronegative metallic atom. In
  an ionic compound the positive and negative ions
  combine so that the overall charge is zero.

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• Sometimes the more electronegative atom is not
  powerful enough to completely take away the
  electrons from another atom so the atoms SHARE
  electrons. This sharing of electrons is called a
  covalent bond.
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Ionic Bonding occurs between metals and
nonmetals.Covalent Bonding occurs between
nonmetals. Bonds (and compounds) form in
order to obtain an electron configuration like
that of noble gases!
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II. Formation of Ionic Bonds and Ionic Compounds

1. Electron Dot Structures show the placement and
   transfer of valence electrons. Rules to remember
   when drawing electron dot structures

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• Only valence electrons are shown. Valence
  electrons are the electrons in the outermost s
  and p sublevels. Transition metals could also
  have d sublevel valence electrons.
• Valence electrons are shown as dots and are not
  drawn randomly! They are arranged around the
  element's symbol to correspond to the elements
  electron configuration. (Only 2 dots or
  electrons per side.)
• 3. Follow the Octet Rule which sates that atoms
  form bonds in order to obtain 0 or 8 valence
  electrons, because of this electron dot
  structures will show no more than 8 electrons for
  each atom or ion. Another way to think of the
  Octet rule Atoms react by changing the number of
  their electrons so as to acquire the stable
  electron configuration of a noble gas.

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B. Electron Dot Structures for Atoms

• Write the element's symbol and place the
  appropriate number of dots to represent the
  valence electrons around the symbol. (The
  electron configuration is given to help you
  understand the idea of valence electrons.)
• a.) Ca Ar4s2
• b.) Li He2s1
• c.) Be He2s2
• d.) O He2s22p4
• e.) Br Ar4s23d104p5

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C. Electron Dot Structures for Ions Ions form
when atoms lose or gain valence electrons.
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• (1.) cations - these form when atoms have LOST
  valence electrons.
• a.) Mg ion
• b.) Li ion
• c.) Al ion
• d.) Ba ion

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• (2.) anions-these form when atoms have GAINED
  valence electrons.
• a.) S ion
  
• b.) Br ion
• c.) N ion
• d.) P ion

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• (3.) Transition and Inner Transition
  Elements-the number of valence electrons for
  these are harder to predict based on their
  position on the periodic table because some of
  these elements have valence electrons in the d
  sublevel. Example
• a.) How many valence electrons does an atom of
• iron have? To answer this question write the
• electron configuration for iron
• Are there any unstable electrons in the d level?
• When iron ionizes what are the possible ions?

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• b.) How many valence electrons does an atom of
  titanium have?
• Electron configuration for titanium
• Are there any unstable electrons in the d level?
• When iron ionizes what are the possible ions?

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• D. Pseudo-noble gas electron configuration-elemen
  ts that cannot acquire a noble gas electron
  configuration, but can become somewhat stable
  with 18 electrons in their outer shell. Examples
  are Hg2, Cd2, Au1, Cu1

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E. Electron Dot Structures for Ionic Compounds

• Write the electron dot structure for each of the
  elements involved.
• Draw arrows from the electrons of the metallic
  atom to the non-metallic atom. This shows the
  transfer of electrons.
• 3. After the Write the dot diagram for the new
  ionic compound, including charges.

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• Lewis Dot Structures for Ionic Compounds
  (compounds held together by ionic bonds usually
  a M
• Example a.) Sodium and Chlorine
• Na Cl ? Na Cl -

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Examples

• b.) Magnesium and Oxygen

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Examples

• c.) Aluminum and Oxygen

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Examples

• d.) Calcium and Fluorine

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Examples

• e.) Sodium and Nitrogen

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• F. Characteristics of ionic compounds (compared
  to molecular compounds)
• -higher melting points
• -higher boiling points
• -generally hard, brittle solids
• -when melted or dissolved in water they can
  conduct electricity
• -shapes are crystalline in nature
• square/cube

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Copper sulfate has a triclinic crystal structure.
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• Lattice Energy the energy released when one mole
  of an ionic crystalline compound is formed from
  gaseous ions.
• Negative values for lattice energy mean that
  energy was released when the ionic crystal is
  formed.

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• Formula unit the smallest number of atoms that
  can make up an ionic compound

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III. Nomenclature Ionic Compounds

• Part 1 Writing Formulas for Ionic Compounds

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• A. Rules for Writing Formulas for Binary Ionic
  Compounds these are compounds containing only 1
  metal and 1 nonmetal.
• 1. Write the cation (metal ion) first and the
  anion (nonmetal ion) second.
• 2. Determine the smallest whole number ratio of
  cations to anions that would make the charge 0.
• To determine the ion formed for main group
  elements look where it is located on the periodic
  table. Many elements in groups 3-12 have either
  a 2 or 3 charge.
• Noble gases to do not form ions except in rare
  cases.
• Hydrogen can either gain, lose, or share an
  electron depending on the other elements with
  which it combines.

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• Example
• Write the formula for sodium chloride.
• Na1 Cl1- NaCl charges equal 0
• Example Write the formula for aluminum oxide.
• Al3 O2- Al2O3 charges equal 0

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• Practice
• Magnesium phosphide____________________
• iron(II) bromide______________________
• Calcium oxide __________________________
• sodium sulfide _____________________
• Copper (II) iodide __________________________
• lead (IV) nitride ____________________
• Aluminum nitride _______________________
• tin (II) chloride _________________
• Potassium fluoride __________________________
  
• Copper (I) phosphide ________________
• Copper (II) oxide ___________________________
• potassium bromide __________________
• Iron (III) fluoride _______________________
• Tin (II) oxide ______________________

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• Rules for Writing Formulas for Ternary Ionic
  Compounds these are compounds containing
  polyatomic ions. (Look at the chart to figure
  out the formula and charge.
• Polyatomic Ions-two or more elements (usually
  nonmetals) bonded together that have collectively
  lost or gained electrons and now have a charge.
  Compounds have a zero charge but a polyatomic ion
  has a charge. You must memorize several
  polyatomic ions (look to Memorize These Ions
  sheet).

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• 1. Write the cation first and the anion second.
  
• 2. Determine the smallest whole number ratio of
  cations to anions that would make the charge 0.
  If a subscript must be added to a polyatomic ion,
  keep the polyatomic ion in parentheses.
• Example
• Write the formula for sodium phosphate.
• Na1 (PO4)3- Na3PO4
• Ammonium sulfide
• (NH4)1 S2- (NH4)2S

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• Practice
• aluminum sulfate _______________
• potassium chlorate ______________
• Copper (II) acetate ______________________
• plumbous nitrate _______________________
• Iron (III) oxalate _______________________
• magnesium chlorate _____________________
• Magnesium dichromate __________________
• tin (II) hypochlorite ____________________
• Lead (II) perchlorate__________________
• tin (II) nitrite _________________________
• Ammonium carbonate___________________
• iron (II) sulfite _________________________
• Sodium cyanide ___________________
• Lithium phosphite _____________________

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Part 2 Writing Names for Ionic Compounds

• There are two naming systems currently
  acceptable
• IUPAC International Union of Practical and
  Applied Chemistry is the newest system-this
  system uses Roman numerals to give the charges or
  oxidation number of positive ions ONLY if the
  positive ion has variable charges. This is the
  system we will use.

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• 2. ous and ic system oldest system and
  still very commonly used. May be used ONLY if
  the positive ion has a variable charge and
  exhibits only 2 oxidation numbers.
• Review the 4 ions with which we may use this
  system.
• Cu ____ _________ _____
  ________
• Sn ____ _________ _____
  ________
• Fe ____ _________ _____
  ________
• Pb ____ _________ _____
  ________

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• A. Rules for Writing Names for Binary Ionic
  Compounds these are compounds containing only 1
  metal and 1 nonmetal.
• 1. The correct full name of the cation (metal
  ion) is written first. (Do not forget about the
  roman numerals if it is a Cu, Fe, Sn, or Pb ion!)
  
• The last syllable in the anion (nonmetal ion) is
  dropped and ide is added.
• Example NaCl Sodium Chloride
• Example CuS Copper (II) Sulfide or Cupric
  Sulfide

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• Practice
• MgCl2 ___________________________
• AlI3 _____________________________
• Na3P _____________________________
• Ca3N2 _______________________________
• FeN _____________________________
• PbCl2 ________________________________
• CuF ____________________________
• CuCl2 _________________________________
• ZnS ______________________________
• Pb02 _________________________________
• Fe203 _____________________________
• KI ___________________________________
• Cu0 ______________________________
• Cs3N ________________________________

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• B. Rules for Writing Names for Ternary Ionic
  Compounds - these are the compound containing
  polyatomic ions. (It is imperative that you know
  the correct names of these ions!!!!)
• 1. The correct full name of the cation (metal
  ion or polyatomic ion) is written first. (Do not
  forget about the roman numerals if it is a Cu,
  Fe, Sn, or Pb ion!)
• 2. The correct full name of the anion
  (polyatomic ion or nonmetal ion) is written
  second. If the anion is a polyatomic ion do not
  change the ending. If the anion is a nonmetal
  ion then the ending is dropped and ide is added.
  
• Example KNO3 potassium nitrate
• Example Cu2CrO4 copper (I) chromate or
  cuprous chromate

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• Practice
• Na3PO4_______________________
• Al2(SO4)3 _________________________________
• CuNO3 _______________________
• PbCO3 __________________________________
• Li2SO3 ________________________
• CaCr207 _______________________________
• NH4Cl_________________________
• CsClO4 __________________________________
• K2CN ________________________
• Fe(HSO3)3 _______________________________

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Mixed Review

• Write the name for the following
• CuO BaO CaCl2
• NaBr K2O Mg3N2
• AgOH Pb(ClO)4 CaSO3
• Sr(NO3)2
• Write the formula for the following
• potassium iodide iron (II) chloride
• sodium sulfide aluminum sulfide
• copper (II) nitride potassium oxide
• lead (IV) oxalate magnesium phosphite
• sodium bicarbonate

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Part 3 Writing Names and Formulas for Acids

• A. Rules for Writing Names for Acids
• Acids compounds that produce hydrogen ions when
  dissolved in water an anion bonded to a hydrogen
  cation is an acid. Any compound starting with
  hydrogen is an acid.
• Even though acids usually contain only nonmetals,
  they are treated as ionic compounds because the
  hydrogen present is the H1 ion.

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Rules for writing the name of an acid

• Does the compound contain an oxygen atom?
• If No always start with hydro, then the root
  of the second element, and end with -ic.
• Example HCl hydrochloric acid
• If Yes oxygen is present start with the root
  of the anion (usually a polyatomic ion).
• If it ends in ate, change it to ic.
• If the anion ends in ite, change it to ous.
• Example HNO3 Nitric acid

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Common Roots for the Polyatomic Ions

• Acetate acet- Bromate brom-
• Chlorate chlor- Nitrate nitr-
• Nitrite nitr- Chromate chrom
• Oxalate oxal- Sulfate sulfur-
• Sulfite sulfur-
• Phosphate phosphor-
• Perchlorate perchlor-
• Hyporchlorite hypochlor-
• Carboante carbon-

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Practice
Practice writing the name for the following acids

• HNO2 HCl H2SO4
• HBr H3N HF
• HI H2S

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B. Rules for Writing the Formula for an Acid

• Hydrogen will always be the cation (H1). Write
  the cation first and the anion second.
• Determine the smallest whole number ratio of
  cations to anions that would make the charge 0.
• Example phosphoric acid
• H1 (PO4)3- H3PO4
• Example phosphoric acid H3PO4

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• nitric acid_________________
• acetic acid______________________
• carbonic acid _______________
• chloric acid ___________________
• chlorous acid__________________
• phosphorous acid _____________

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• Lets Review
• Binary Ionic Compounds (BIC) 2 elements, one is
  a metal and one a nonmetal, will end with ide,
  use periodic table to look up ions formed to
  determine the formula
• Ternary Ionic Compounds (TIC) 3 or more
  elements, at least one is a metal and at least
  one is a nonmetal, these will contain a
  polyatomic ion (memorize the polyatomic ions),
  use the ions charge to determine the formula
• Acids (A) will begin with H, hydrogens charge
  when forming an acid is H1,
• If the acid does NOT contain oxygen then start
  the name with hydro-, followed by the root of the
  second element, and end with ic.
• If the acid does contain oxygen start with the
  root of the anion (usually a polyatomic ion) and
  then change the ending to ic if it was ate
• and ous if it was ite.

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• CaCO3 Mg3P2
• Cu(NO2)2
• CuCl HCl
• Magnesium permanganate
• Barium fluoride
• Iron (III) nitrate
• Sulfuric acid
• Lead (IV) fluoride

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Part 4 Percent Composition, Empirical, and
Molecular Formula Problems

• A. Molar Mass Practice
• Find the molar mass of ammonium sulfate (also
  called the formula mass)
• Find the molar mass of copper (II) chloride

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• B. Percent Composition by Mass the percent, by
  mass, of each element in a compound.
• If you have a box containing 100 golf balls and
  100 ping pong balls, which type of ball
  contributes the most to the mass of the box?
• The same principle applies to finding the
  composition of a compound. Different elements
  have different masses and this must be taken into
  consideration.

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• How to find the percent composition of a
  compound
• Write a correct formula for the compound
• Find the molar mass of the compound
• 3. Divide the total atomic mass of EACH ELEMENT
  by the molar mass
• 4. Multiply by 100 to convert your results to a
  percent
• 5. Since you have no significant figures to go
  by, express your answer to TWO decimal places
  with the sign.

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• Practice
• Find the percentage composition by mass of zinc
  carbonate.
• Find the percentage of nitrogen by mass in
  ammonium nitrate.
• Find the percentage composition by mass of
  aluminum oxalate.
• Silver is worth 7.40 per Troy ounce ( 1 lb
  12 troy oz). How many grams of silver nitrate
  must be decomposed in order to sell the PURE
  SILVER for 200.00?

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Part 3 Empirical and Molecular Formulas

• C. Empirical Formulas - are formulas which have
  been reduced to their lowest terms (we write all
  formulas for ionic compounds as empirical
  formulas).
• Empirical formulas can be calculated from 2 types
  of data
• a. you may be given the percentage of each
  element present in the compound
• b. you may be given the actual number of grams
  of each element present in the compound

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• Here are the steps to follow to work this type of
  problems
• Take either the GRAMS of each element given or
  the PERCENTAGE of each element given and divide
  it by the atomic mass of that element. Leave the
  result of each division in AT LEAST 4 SIG FIGS.
• Compare all the results from Step 1 by selecting
  the SMALLEST value and dividing ALL values by
  this smallest one.
• The results of Step 2 will either be VERY close
  to whole numbers or will be recognizable mixed
  number fractions (decimals such as 1.5, 2.333,
  6.67, etc).
• If any result from Step 3 is a mixed number, you
  must multiply ALL values by some number to make
  it a whole number. Ex 1.33 x 3, 2.25 x 4,
  2.50 x 2, etc.
• 5. Use these whole number results as SUBSCRIPTS
  and write the empirical formula, listing the
  elements in the order they are given in the
  problem. (HINT dont be surprised if the
  subscripts in some formulas are VERY large-many
  organic molecules are huge)

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• Here is a way to remember the steps percent to
  mass, mass to mole, divide by small, multiply
  till whole
• Example 100. grams of a compound contains
  22.430 grams of carbon, 6.542 grams of hydrogen,
  44.8598 grams of oxygen and 26.1682 grams of
  nitrogen. Find the empirical formula.
• Example Find the empirical formula for a
  compound which contains 26.8 Sn, 16.0 Cl and
  57.2 I.

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• D. Molecular Formulas are either the same as
  its experimentally determined empirical formula
  or its some whole number multiple of it.
• To determine the molecular formula, you must know
  the compounds empirical formula AND the molar
  mass of the molecular compound.
• How to find the molecular formula
• Calculate the mass of the empirical formula
  (which you have already found or it will be given
  to you )
• Divide the known molar mass by the mass of the
  empirical formula.
• Multiply that number by the subscripts of the
  empirical formula to get the subscripts for the
  molecular formula.

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• Example The molar mass of a compound is 181.50
  g/mol and the empirical formula is C2HCl. What
  is the molecular formula?
• Example Find the empirical formula for a
  compound containing only carbon and hydrogen if
  it is known to contain 84.21 carbon.
• b. If the molar mass is 114 g/mol, what is the
  molecular formula of this compound?