Molecules and Compounds

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Molecules and Compounds

• Chapter Three

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Molecules

• Molecules are groups of atoms chemically bonded
  together.
• Molecules may be elements (the diatomic, sulfur
  and phosphorus) or compounds.

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Allotropes/Molecular Elements

• Some Elements exist as molecules
• Hydrogen, fluorine, chlorine, bromine, iodine and
  oxygen exist as diatomic molecules HOFBrINCls
• Phosphorus exists as a tetratomic molecule
• Some elements exist in a variety of forms
• Carbon graphite diamond buckminsterfullerine
• Phosphorus - red and white
• Sulfur - S6 and S4

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Molecules and compounds

• Molecular compounds - molecules containing atoms
  from two or more different elements
• Covalent bonds - the force holding the atoms
  together in a molecular compound

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Formulas

• A compound is represented by using the symbols
  for the elements of which it is composed
• Subscripts are used to indicate how many atoms of
  a particular element exist in the compound
• If there is only one atom of a particular
  element, the one is assumed

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Formulas, cont

• Changing the subscripts changes the compound
• consider H2O and H2O2
• Two different compounds can, however, share the
  same chemical formula
• dimethyl ether and ethyl alcohol both have the
  formula C2H6O

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Ethyl alcohol and dimethyl ether

• Ethyl alcohol on the left dimethyl ether on the
  right
• These species are termed geometric isomers
• Formulas that show the order and arrangement of
  specific atoms are known as structural formulas

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Electrical nature of matter

• Electrostatic forces
• attraction between opposite charges
• repulsion between same charges
• Charged atoms or molecules are known as ions
• cations - positively charged
• anions - negatively charged

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Ionic Compounds

• Compounds consisting of ions are known as ionic
  compounds.
• The forces holding them together are called
  ionic bonds

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How charged species arise

• Neutral atoms and molecules have the same number
  of protons and electrons
• Cations have more protons than electrons
  resulting from the loss of an electron
• Anions have more electrons than protons resulting
  from the gain of an electron

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Formulas of Ionic compound

• Formula unit - simplest whole-number ratio of
  ions in an ionic compound
• For example Ca2 Br-
• you need to have the resulting formula be
  electrically neutral
• so two Br- are needed for each Ca2
• the resulting formula is CaBr2

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Types of Ionic Compounds

• Ionic compounds will always consist of one of the
  following combinations a metal and a nonmetal,
  a polyatomic ion and a nonmetal, a metal and a
  polyatomic ion or two polyatomic ions. Ionic
  compounds can be distinguished from molecular
  compounds by the kinds of elements they contain.

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Polyatomic ions

• cations or anions consisting of groups of atoms
  that are covalently bonded to each other
• examples are NO3-, SO42-
• when more than one appears in a formula unit, the
  polyatomic ion is put in between parentheses, and
  a subscript is used to indication the number of
  the ions that appear in the formula unit
• example Ba(ClO3)2

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Guidelines for determining if a compound is Ionic

• Metals almost always form positive ions and form
  ionic compounds
• Nonmetals form monatomic ions in ionic compounds
  only when combined with a metal.
• It is difficult to predict when the metalloids
  form ions
• The farther apart two elements are in the
  periodic table, the more likely they are to from
  an ionic compound on reaction

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Ionic Crystal Lattice

• Ionic compounds are generally solids and have
  their ions arranged in extended three-dimensional
  networks. This regular array of positive and
  negative ions is called a crystal lattice.

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Properties of Ionic Compounds

• High melting points that correlate with charges
  on ions
• Most ionic solids do not conduct electricity but
  molten ionic compounds do.
• Most ionic compounds dissolve in water

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Properties cont.

• Solutions of ionic compounds in water conduct
  electricity (electrolytes)
• In ionic substances, each ion has its own
  characteristics, and these are different from the
  characteristics of the atom from which the ion
  was derived (NaCl)

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Binary compound nomenclature

• Five types of binary compounds
• Metals exhibiting only one oxidation state
  forming a compound with a nonmetal or polyatomic
  ion
• Metals exhibiting two or more oxidation states
  forming a compound with a nonmetal or polyatomic
  ion
• Ammonium ion with nonmetal
• Two polyatomic ions
• Compounds of nonmetals and nonmetals

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Metals with only one oxidation state

• Groups of metals with only one common oxidation
  state
• alkali metals - 1
• alkaline earths - 2
• Zn - 2
• Al - 3
• All other metals can exhibit more that one
  oxidation state

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Anions in negative oxidation states

• Nonmetallic anions usually exhibit one negative
  oxidation state
• halogens -1
• chalcogens -2
• N, P -3

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Forming compounds

• Metal and nonmetal combine to neutralize charge
• Consider - Al3, O2-
• cross multiply charges
• 2 Al3 3 O2- Al2O3

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Naming binary compounds

• Use name of metal with no changes
• Change the name of the anion by taking the stem
  and add the suffix -ide
• Examples
• NaCl - sodium chloride
• MgCl2 - magnesium chloride
• AlCl3 - aluminum chloride

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Metals with multiple oxidation states

• Two systems Stock and IUPAC
• IUPAC system
• metal name and the oxidation state in Roman
  numbers in parenthesis
• Fe2 iron(II)
• Form compound by balance charge of metal with
  correct number of nonmetals
• CoCl3 cobalt(III) chloride

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Stock System Nomenclature

• Metals in multiple oxidation states usually have
  one or two common oxidation states
• First row transition metals are 2 and 3 (except
  Cu2 and Cu)
• use -ous suffix for lower common oxidation state
• use -ic suffix for higher common oxidation state

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Examples

• CoCl3 - cobaltic chloride
• NiCl2 - nickelous chloride
• For metals with Latin names, use them
• CuCl - cuprous chloride
• FeBr3 - ferric bromide

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Oxy anions

• anions composed of oxygen and another elements
• other elements can be a metal or a nonmetals
• Examples
• SO42-, NO2-, MnO4-

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Naming

• Need common oxidation states
• most common oxidation state for nonmetals is the
  group number (except for the halogens)
• next most common oxidation state is the group
  number minus one
• use -ate suffix for higher oxidation state and
  -ite suffix for next higher oxidation state

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Examples

• SO42- - sulfate
• SO32- - sulfite
• NO3- - nitrate
• NO2- - nitrite
• Salts with these oxyanions
• Na2SO4 - sodium sulfate
• KNO3 - potassium nitrate

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Binary Molecular Nomenclature
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Nonmetals nonmetals

• Name nonmetal further to the left of the periodic
  table first with no changes
• Name nonmetal further to the right of the
  periodic table second with the -ide suffix
• Use Greek prefixes to indicate the number of each
  one

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Greek prefixes
Number Prefixes
1 mono
2 di
3 tri
4 tetra
5 penta
6 hexa
7 hepta
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Examples

• N2O3 - dinitrogen trioxide
• CO2 - carbon dioxide
• P2O5 - diphosphorus pentoxide

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Acids

• Binary acids
• name begins with hydro
• then add stem of nonmetal plus -ic
• end with acid
• Examples
• HCl - hydrochloric acid
• H2S - hydrosulfuric acid

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Oxyacids

• Take oxyanion suffix and convert
• change -ate to -ic
• change -ite to -ous
• Do not use hydro- in the beginning
• Examples
• H2SO4 - sulfuric acid
• H2SO3 - sulfurous acid

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Hydrates

• Some ionic compounds can have water molecules
  attached within the structure
• These compounds are termed hydrates and have
  properties distinct from the unhydrated form

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Naming Hydrates

• Hydrates are named by naming the ionic compound
  and then using a Greek prefix to indicate the
  number of water molecules followed by the word
  hydrate

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Chemical Formulas

• Molecular compounds
• chemical formula represents a discrete molecular
  unit (e. g. CO2)
• Ionic compounds
• chemical formula represents a formula unit (the
  whole number ratio of cations to anions e. g.
  K2SO4)

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Empirical Formula

• Simplest whole number ratio of atoms in the
  compound
• All ionic formulas are empirical
• Molecular formulas are either equal to the
  empirical or a whole number multiple

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The composition of compounds

• Mole composition is the number of moles of each
  of the elements that make up the compound
• CO2 - one mole of C and two moles of O
• Mass composition is the mass of each element in
  the compound
• CO2 - 12.0 g of C and 32.0 g of O

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Percent composition

• mass of each element per 100 mass units of
  compound
• in 100 g of NH3, there is 82.0 g of N
• therefore, the mass percentage of N is 82.0 N

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CO2

• Calculation of composition of carbon dioxide
  requires determining the number of grams of each
  element (C and O) in one mole

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Determination of Empirical formula

• Convert percent composition to an actual mass
• Convert mass to moles of each element
• Find the whole number ratio of the moles of
  different elements

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Laughing gas

• Contains 63.6 N and 36.4 O
• Assume 100 g of substance, so you have 63.6 g of
  N and 36.4 g of O
• Calculation gives an empirical formula of N2O

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Empirical Formula from Experimental Analysis

• Combustion Method
• Combust compound in oxygen. Carbon is converted
  into carbon dioxide, hydrogen is converted into
  water, remaining element is found by difference.

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Example

• An acetic acid sample with a mass of 1.000 g
  combusts to give 1.466 g CO2 and 0.6001 g H2O.
  The compound is known to contain C, H, and O.
• 0.4001 g C from CO2
• 0.0673 g H from H2O
• 0.533 g O by difference

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Empirical Formula From Experiment

• Actual measurement of masses is determined from
  experiment.
• Mass is converted to moles
• Simplest whole number ratio is determined

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Molecular formula

• The actual number of each atom in a formula unit
• Consider acetylene and benzene
• both have the empirical formula CH
• acetylene is actually C2H2
• benzene is actually C6H6

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Molecular Formula from Empirical

• Molecular formula must be integral multiple of
  empirical formula therefore the mass of the
  molecular formula must be the same integral
  multiple of the mass of the empirical formula.

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Example

• Acetic acid
• mass of molecular 60 g/mol
• mass of empirical formula 30 g/mol
• ratio 2
• empirical formula CH2.O
• molecular formula C2H4O2