Ionic Compounds

1
Ionic Compounds

• Ionic compounds usually consist of metal and
  non-metal elements
• Formation of ionic compounds can be considered an
  electron exchange process
• The metal gives up electrons and the non-metal
  gains electrons
• Na Na e-
• Cl e- Cl-
• Na Cl- NaCl
• The result is the compound produced has no net
  charge
• The metal gives up as many electrons as are
  gained by the nonmetal
• Writing Formulas for Ionic Compounds
• Empirical Formulas are generally written
• From the ionic charges the formulas can be
  readily written
• The charge on the positive ion (cation) becomes
  the subscript after the negative ion (anion)
• The charge on the negative ion (anion) becomes
  the subscript after the positive ion (cation)
• Na2S

2
Ionic Compounds

• Problem solving tips
• Figure out the charge on a monatomic ion from its
  position in the periodic table
• You need to know the charges on polyatomic ions
  Learn them!
• If you dont know the charge on a polyatomic ion,
  perhaps you can figure it out from the formula
  of a compound of that ion
• Youre asked to write the formula for magnesium
  bismuthate
• Looking up the formula for sodium bismuthate
  gives NaBiO3
• Na is the formula for sodium ion, so bismuthate
  is BiO3- and magnesium bismuthate must be
  Mg(BiO3)2.
• Most metal-containing compounds are ionic. Learn
  where metallic elements occur in the periodic
  table.
• Compounds without metals are usually not ionic
  but molecular
• Exceptions are compounds of polyatomic ions
  NH4Cl, (NH4)3PO4
• Learn the formulas, charges and names of the
  common polyatomic ions.

3
Ionic Compounds

• Ionic compounds and Coulombs Law
• Electrostatic forces of attraction and repulsion
  between objects results from the electric charge
  they possess.
• This is the kind of interaction responsible for
  the bonding between the ions making up ionic
  compounds
• The quantitative relationship describing the
  force of attraction between ions is given by
  Coulombs Law
• where n is the charge on the cation, n- is the
  charge on the anion, and d is the distance
  between them
• If the charge on one of the ions doubles, the
  force of attraction doubles
• If the charge on both ions doubles, the force of
  attraction quadruples
• If the distance is cut in half, the force of
  attraction quadruples

• Crystal Lattices are the arrangements of ions in
  ionic compounds
• In the solid, the ions are constrained to fixed
  locations
• The lattice structure must be consistent with the
  compounds formula
• or formula unit

4
Lecture 5
Slide 4
5
Naming Inorganic Compounds

• Inorganic compounds generally do not contain
  carbon
• Naming Cations
• Cations formed from metal atoms are given the
  name of the element
• Na sodium ion Ca2 calcium ion Al3 aluminum
  ion
• Some metals - particularly transition metals -
  can have more than one charge
• Fe2 iron(II) ion Fe3 iron(III) ion
• or
• Fe2 ferrous ion Fe3 ferric ion
• Cu copper(I) ion Cu2 copper(II) ion
• or
• Cu cuprous ion Cu2 cupric ion
• Cations from non-metal elements have names that
  end in -ium
• NH4 ammonium ion H3O hydronium ion
• Learn the Formulas - including charge - and names
  of the ions in Figure 3.6, p. 107

6
Naming Inorganic Compounds

• Naming Anions
• Monatomic anions have names formed by dropping
  the ending of the elements name and adding
  -ide
• H- hydride S2- sulfide P3- phosphide
• Some polyatomic anions have names that end in
  -ide
• OH- hydroxide CN- cyanide O22- peroxide
• Oxyanions - polyatomic anions having oxygen -
  have names ending in -ate or -ite
• NO3- nitrate NO2- nitrite
• SO42- sulfate SO32- sulfite
• Oxyanions of the halogens have prefixes in their
  names
• ClO4- perchlorate ClO3- chlorate ClO2-
  chlorite ClO- hypochlorite
• Anions formed by adding H to the oxyanion add a
  prefix word hydrogen or dihydrogen or bi-
• CO32- carbonate HCO3- hydrogen carbonate or
  bicarbonate
• PO43- phosphate H2PO4- dihydrogen phosphate
• Learn the Formulas - including charge - and names
  of the ions in Table 3.1, page 110

7
Names and Formulas of Binary Molecular Compounds

• Molecular compounds are formed by the non-metal
  representative elements
• Rules for naming binary compounds
• The name of the element farthest to the left in
  the periodic table is usually written first.
• If both elements are in the same group of the
  periodic table, the lower one is usually
  written first.
• The name of the second element is given an -ide
  suffix.
• Greek prefixes indicate the number of each
  element in the molecule.
• mono is not used with the first element
• when the prefix ends in a or o and the name of
  the anion begins with a vowel, the a or o is
  often dropped
• N2O dinitrogen monoxide NF3 nitrogen
  trifluoride
• NO nitrogen monoxide H2O water - a common
  name
• NO2 nitrogen dioxide NH3 ammonia - a common
  name
• N2O5 dinitrogen pentoxide HCl hydrogen
  chloride
• H2S(aq) hydrosulfuric acid HCl(aq) hydrochloric
  acid

8
Names and Formulas of Binary Molecular Compounds

• Some oddities in formulas
• Formulas for binary compounds containing H are
  usually written with H second, after the other
  element
• CH4, C2H6, NH3, N2H4
• Compounds of H and O or the halogens have H
  written first
• H2O, H2O2, HF, HCl, HBr, HI
• Common Names
• CH4 methane C2H6 ethane C3H8
  propane C4H10 butane
• NH3 ammonia N2H4 hydrazine
• PH3 phosphine
• NO nitric oxide N2O nitrous oxide
• H2O water

9
Atoms, Molecules and the Mole

• The Mole
• The mole is the SI unit of amount of matter
• A mole contains the same number of fundamental
  particles as are contained in exactly 12 g of
  .
• These fundamental particles could be atoms,
  molecules, protons, electrons, people, houses,
  or any other object.
• This number of particles is 6.022 x 1023
  particles - Avogadros number
• This number must be committed to memory!
• For elements, the Molar mass is the mass in grams
  of one mole of atoms of the element
• The molar mass of an element is numerically equal
  to the atomic mass in amu.
• One atom of has mass 12 amu one mol
  has mass 12 g
• The molar mass of H is 1.0079 amu one mol of H
  has mass 1.0079 g
• The molar mass of Cu is 63.546 amu one mol of Cu
  has mass 63.546 g

10
Atoms, Molecules and the Mole
Conversions between moles to mass and mass to
moles
x Molar mass
Molar mass
Moles Mass Mass
Moles

• Example Calculate the number of g of C in 0.432
  mol of C
• Example Calculate the number of moles of C in
  32.4 g C
• Example Calculate the number of atoms in a given
  mass of an element.
• How many atoms of C are in 32.4 g C?

11
Atoms, Molecules and the Mole

• Molecular and empirical formulas, moles and mass
• Empirical formulas are used to describe ionic
  compounds
• Al2O3 says that the atom ratio of the elements in
  this compound is
• 2 atoms of Al to 3 atoms of O
• It also gives the mole ratio of the atoms in this
  compound
• 2 mol Al to 3 mol of O
• One mol Al2O3 contains 2.00 mol Al and 3.00 mol O
• The molar mass of Al2O3 formula weight Al2O3
  102.0 g/mol

• Be careful to understand the chemical state of an
  element
• N vs. N2 H vs. H2 O vs. O2 Cl vs. Cl2
• one mol N is 14.0 g
• one mol N2 is 28.0 g

12
Atoms, Molecules and the Mole

• Molecular and empirical formulas, moles and mass
• Molecular formulas are used to describe molecular
  substances
• Example
• Calculate the molar mass of ethanol, C2H6O

• Mass Percentage Composition
• Gives the by mass of each element in a formula
• The relative mass of an element in a formula is
  its atomic mass x number of atoms in the formula
• The by mass of an element is 100 x the ratio of
  the its mass to the total mass of all the
  elements in the formula, which is the formula wt.
• For C8H18,

13
Atoms, Molecules and the Mole

• Empirical Formulas from Mass of Elements
• The empirical formula gives the simplest whole
  number ratio of the number of atoms of each
  element in a compound.
• The empirical formula also gives the simplest
  whole number ratio of the number of moles of
  each element in one FW of compound.
• Given the mass of each element
• The mass s will sum to 100
• Assume the respective mass s are grams of each
  element
• Calculate the number of moles of each element
  from these masses
• Find the ratio of moles by dividing each number
  of moles by the smallest number of moles
• Experimental uncertainty and round-off errors
  will often produce uncertainties in the mole
  ratio

14
Atoms, Molecules and the Mole
Example - isopropyl alcohol is the major
component of rubbing alcohol. It has 59.9
C, 13.4 H, and 26.7 O. What is the
empirical formula of isopropyl alcohol?
15
Atoms, Molecules and the Mole

• Molecular formula from the Empirical Formula
• Determine the molecular formula from the molar
  mass of the compound
• The subscripts in the molecular formula must be
  an integral number times the subscripts in an
  empirical formula.
• This multiplier is equal to the ratio of the
  molar mass to the empirical formula mass.
• Example
• Experimental determination of the molar mass
  of isopropyl alcohol gives 60.0 g/mol
• The empirical formula mass is 3(12.0)
  8(1.01) 1(16.0)60.1 g/mol
• The ratio of the molar mass to empirical
  formula mass is 0.998
• Therefore, there is one empirical formula unit
  per molecular formula unit and the molecular
  formula is C3H8O.

16
Atoms, Molecules and the Mole

• Some useful results from formulas and molar
  masses
• From the molecular formula of a compound and some
  knowledge of how atoms can be arranged in
  molecules, one can come up with how the atoms
  are connected and arranged in three space.
• From empirical formulas, one can obtain important
  quantitative information
• Example How much iron metal can be produced from
  100.0 g Fe3O4?