Ionic

1
Ionic Covalent Compounds
2
Bonding

• Chemical bonds are forces that cause a group of
  atoms to behave as a unit.
• Bonds result from the tendency of a system to
  seek its lowest possible energy.
• Bond breaking always requires energy, and bond
  formation always releases energy.

3
Types of Bonds

• The type of bonding depends upon the nature of
  the atoms that are combined.
• A metal and a non-metal will form ionic bonds
  when electrons are transferred from the metal to
  the non-metal. The resulting attraction between
  oppositely charged ions forms a stable crystal.

4
Types of Bonds

• When metals bond with each other, the valence
  electrons are shared by the atoms in the entire
  crystal. The electrons are no longer associated
  with a specific nucleus, and are free to move
  throughout the sample.

5
Lewis Structures

• Lewis Structures, also known as Lewis dot
  diagrams, show how the valence electrons are
  arranged among the atoms in the molecule.
• For ionic compounds, it shows the end result
  when the metal loses its electrons to the
  non-metal.

6
Lewis Structures Ionic Compounds

• A dot (. ) is used to represent each valence
  electron. Consider sodium chloride. The dot
  diagram for the sodium atom is
• Na
• The dot diagram for the chlorine atom is
• Cl

.

.


7
Lewis Structures Ionic Compounds

• Note that electrons (dots) representing valence
  electrons are written singly or in pairs on
  either side of the atomic symbol, and above and
  below the symbol.
• Na Cl
• It doesnt matter which side has the unpaired
  electron.

.
.


8
Lewis Structures Ionic Compounds

• Sodium donates its single valence electron to
  the chorine atom, and chlorine accepts the extra
  electron.
• Na Cl
• After transfer, both ions have a noble gas
  electron configuration.
• Na Cl

.
.


-




9
Lattice Energy

• When metals react with non-metals, a large
  amount of energy is usually released. These
  reactions are usually exothermic.
• The electron transfer from the metal to the
  non-metal usually requires energy, and is
  endothermic.
• The large release of energy comes from forming
  an ionic crystal.

10
Lattice Energy

• The crystal will have oppositely charged ions
  in contact with each other and like charged ions
  separated from each other. The alternating
  charged ions form a crystal lattice. As the
  lattice forms, large amounts of energy are
  released.

11
Lattice Energy
12
Lattice Energy

• Lattice energy is defined as the energy
  required to convert a mole of an ionic solid into
  its constituent gaseous ions.
• The greater the lattice energy, the stronger
  the bonding, and the more stable the compound.

13
Lattice Energy

• Based on Coulombs Law, the lattice energy
  increases with smaller ionic size.

14
Lattice Energy

• Based on Coulombs Law, the lattice energy
  increases significantly with greater ionic
  charge.
• E constant (q1)(q2)/r

15
Melting Points

• As ionic crystals melt, the lattice is
  disrupted, and the ions become free to move, and
  can conduct electricity.

16
Ionic Bonding

• The strength of the attraction between ions
  increases significantly with increasing ionic
  charge, and results in high melting points.

17
Naming Inorganic Compounds

• 1. Binary Compounds
• Binary compounds contain only two elements. The
  elements are either a metal with a non-metal
  (ionic bonding), or two non-metals (covalent
  bonding).

18
Naming Binary Compounds

• a) Metal Non-metal
• When metals react with non-metals, the metal
  loses electrons and the non-metal gains
  electrons. The resulting attraction between
  oppositely charged ions creates ionic bonds.

19
Common Ionic Charges

• The charges of ions of elements in groups 1A-7A
  (the main groups) are usually predictable.
• Group 1A metals form 1 ions, group 2A metals
  form 2 ions, etc.
• The non-metals of group 5A have a -3 charge,
  those of group 6A have a -2 charge, and the
  halogens form ions with a -1 charge.

20
Typical Ionic Charges
21
Naming Binary Compounds

• For example, NaCl is called
• sodium chloride
• Where chlor is the root for the element
  chlorine.

22
Naming Binary Compounds

• Three common transition metals also have only
  one ionic charge, and are also named the same
  way.
• They are zinc ion (always 2), silver ion
  (1) and cadmium ion (2)
• ZnS is zinc sulfide, as sulf is the root for
  sulfur.

23
Writing Formulas of Binary Compounds

• Compounds have no net charges, so the formulas
  of ionic compounds must contain equal numbers of
  positive and negative charges.
• Magnesium bromide, made from magnesium ion
  (Mg2) and bromide ion (Br1-) has the formula
• MgBr2

24
Binary Compounds with Variable Charge Metals

• Most transition metals and the metals on the
  lower right side of the periodic table can have
  several ionic charges.
• The properties of the ion vary greatly with
  charge, so the charge must be specified in naming
  the ion or its compounds.

25
Typical Ionic Charges
26
Binary Compounds with Variable Charge Metals
27
Binary Compounds with Variable Charge Metals

• If an ion has variable charges, you must
  specify the charge in naming the metal.
• If an ion has only one charge, it is incorrect
  to specify its charge.

28
Naming Fe2O3

• Fe2O3 is an iron oxide, but we must specify the
  charge of the iron ion.
• We know each oxide has a -2 charge, so three
  oxide ions have a total charge of -6.
• The two iron ions therefore have a charge of
  6, with each iron having a charge of 3.
• The name of the compound is iron(III) oxide.

29
Covalent Bonding
Covalent compounds exist as discrete molecules,
whereas ionic compounds consist of an aggregate
of cations and anions.
30
Covalent Bonds

• When two (or more) non-metals form bonds,
  electrons are shared. The result is a covalent
  bond.
• Covalent bonds form because the attraction of
  electrons for the nuclei in the atoms is greater
  than the electron-electron repulsion or the
  nucleus-nucleus repulsion.

31
Types of Bonds

• There is usually an optimum bond length or
  internuclear distance where attractions between
  electrons and the nuclei are optimized and
  repulsions are minimized.

32
Covalent Bonding
33
Bond Energy
Bond formation releases energy, and bond
breaking requires energy.
34
Molecules

• Molecules are neutral combinations of two or
  more atoms that are bonded together and act as a
  unit.
• Molecules may be elements (H2, O2, O3, F2), or
  compounds containing atoms of different elements
  bonded together. Molecules typically contain
  elements that are non-metals.

35
Molecules

• Scientists studying the nature of matter
  focused much of their research in 1800s on the
  composition of compounds.
• Since molecules are much too small to observe,
  they typically observed the reactions of larger
  amounts of matter and used mass measurements to
  develop their theories of matter.

36
The Law of Definite Proportion

• Joseph Proust (1754-1826) determined the chemical
  composition of many compounds. He found that a
  given compound always contains the exact same
  proportion of elements by mass. This is known as
  the law of definite proportion.
• For example, all samples of water contain 88.8
  oxygen by mass, and 11.2 hydrogen by mass.

37
The Law of Multiple Proportions

• This chemical law applies when two (or more)
  elements can combine to form different compounds.
• Common examples are carbon monoxide and carbon
  dioxide, or water and hydrogen peroxide.
• John Dalton (1766-1844) conducted experiments on
  these types of compounds, and determined that
  there is a simple relationship between the masses
  of one element relative to the others.

38
The Law of Multiple Proportions

• When two elements form a series of compounds, the
  ratios of the masses of one element that combine
  with a fixed mass of the other element are always
  in a ratio of small whole numbers.
• The meaning of this law is difficult to
  understand unless it is illustrated using a
  specific series of compounds.

39
The Law of Multiple Proportions

• Consider the compounds of water and hydrogen
  peroxide. At this point in history, chemists
  knew the compounds were different, and that they
  both contain (or can be broken down into) the
  elements hydrogen and oxygen. They did not yet
  know the formulas for either compound, nor was
  the concept of atoms fully developed.

40
The Law of Multiple Proportions

• Analysis of 100 grams of the compounds produced
  the following data

Compound Mass of oxygen/100g of compound Mass of hydrogen/100g of compound Grams of oxygen/gram of hydrogen
water 88.8 grams O 11.2 grams H 7.93 gO/gH
Hydrogen peroxide 94.06 grams O 5.94 grams H 15.8 gO/gH
41
The Law of Multiple Proportions

• The Law of Multiple Proportions is illustrated
  when the numbers in the last column are compared.

Compound Grams of oxygen/gram of hydrogen
water 7.93 gO/gH
Hydrogen peroxide 15.8 gO/gH
15.8/7.93 2/1 The small whole number ratio
suggests that there is twice as much oxygen in
hydrogen peroxide as there is in water.
42
The Law of Multiple Proportions

• The key feature is that small whole numbers are
  generated. The results support the hypothesis
  that molecules consist of various combinations of
  atoms, and that atoms are the smallest unit of
  matter. The ratio doesnt produce fractions,
  since there is no such thing as a fraction of an
  atom.
• For the example cited, we would propose that
  hydrogen peroxide contains twice as many oxygen
  atoms/hydrogen atoms than does water. We cannot,
  however, determine the actual formula of either
  compound.

43
Empirical Formulas

• The studies of water and hydrogen peroxide lead
  to empirical formulas. These are based on
  experiment, and represent to simplest way of
  expressing the ratio of atoms in a compound.
• The early scientists analyzed new chemical
  compounds to determine their composition and
  chemical formulas. Modern analytical
  laboratories still provide this service.

44
Molecular Formulas

• Molecular formulas show the exact number of
  each type of atom in a molecule. For example,
  hydrogen peroxide has a molecular formula of
  H2O2. Its empirical formula shows that there is
  one hydrogen for every oxygen, so it is OH or HO.
  Neither of these formulas provides the structure
  or arrangement of the bonds in the molecule.

45
Structural Formulas

• Structural formulas provide the arrangement of
  atoms in the molecule. The structural formula
  for hydrogen peroxide is
• H-O-O-H
• This formula shows the arrangement of the
  atoms, but doesnt show bond angles or the shape
  of the molecule.

46
Naming Covalent Binary Compounds

• When two non-metals form a compound, they share
  electrons, rather than transfer them. The
  resulting bond is called a covalent bond.
• The naming of these compounds is fairly simple.
  The first element is named first, and the second
  element is named as the root ide.
• Prefixes are used to indicate the number of
  each atom present.

47
Naming Covalent Binary Compounds
These prefixes are used only for compounds
containing two non-metals. The prefix mono is
never used for the first element in the compound.
48
Naming Covalent Binary Compounds

• The prefix mono is never used for the first
  element. CO2 is carbon dioxide.
• If the prefix ends in an a or o, and the element
  that follows begins with a vowel, the last letter
  of the prefix is usually dropped. N2O5 is called
  dintrogen pentoxide (and not pentaoxide).

49
Naming Covalent Binary Compounds

• Note that these prefixes are only used for
  binary covalent compounds. It is incorrect to
  use them for compounds containing a metal and a
  non-metal.

50
Naming Covalent Binary Compounds

• There are some compounds of metalloids or
  metals in very high oxidation states that are
  sometimes named using this system.

51
Naming Covalent Binary Compounds

• For example, compounds such as SnCl4 or PbCl4
  are covalent in nature, and not ionic solids.
  They may be called tin(IV) chloride or tin
  etrachloride or lead(IV) chloride or lead
  tetrachloride.

52
Binary Compounds with Hydrogen

• With metals, hydrogen can form ionic
  compounds in which the hydrogen has a -1 ionic
  charge. These compounds are named like any
  binary ionic compound.
• NaH is sodium hydride
• CaH2 is calcium hydride

53
Binary Compounds with Hydrogen

• With non-metals, the bonding is covalent.
  Hydrogen never forms a positive ion in nature.
• Many of the compounds containing hydrogen have
  common names that do not follow the usual
  nomenclature rules.

54
Binary Compounds with Hydrogen

• Examples include
• water H2O
• ammonia NH3
• phosphine PH3
• hydrogen sulfide H2S
• Note that the order in which the elements are
  written is also irregular.

55
Binary Compounds with Hydrogen- Acids

• Hydrogen also forms binary compounds that act
  as acids in water. These compounds dissociate in
  water to donate a proton to water.
• HCl(g) H2O(l) ? H3O(aq) Cl(aq)
• hydrogen chloride hydronium

56
Naming Binary Acids

• The naming of the pure compound and its aqueous
  acid solution differ.
• HCl is a gas called hydrogen chloride. HCl(aq)
  is an acid called hydrochoric acid.

57
Naming Binary Acids

• Name the following acids
• H2S(aq) , HBr(aq)

58
Organic Nomenclature

• Compounds containing carbon, hydrogen and
  sometimes oxygen, nitrogen, sulfur and the
  halogens, have a separate system of nomenclature.
  

59
Unusual Ions

• Mercury forms two ions, mercury(I) and
  mercury(II). The mercury(I) ion is polyatomic,
  and exists as two mercury(I) ions bonded
  together. Its formula is Hg22.
• Oxygen in compounds usually exists as the oxide
  ion, O2-. Oxygen also exists as the peroxide
  ion, O22-, with each oxygen having a -1 charge.

60
Naming Polyatomic Ions

• There are many ions, such as sulfate or
  nitrate, that contain more than one element.
• Many of these ions contain oxygen and a
  non-metal.
• These ions can be found in a group of acids
  called the oxy acids (such as sulfuric acid,
  nitric acid, etc.).

61
Polyatomic Ions

• The bonding within these polyatomic ions (such
  as nitrate, sulfate and phosphate) is covalent.
  The ionic charge results from the loss of one or
  more H ions to water, resulting in a negative
  charge on the anion formed.
• In water, the covalently bonded hydrogen is
  donated to water, forming hydronium ions and the
  corresponding anion.

62
Naming the Oxy Acids

• The easiest way to learn the names of the ions
  is to memorize a short list of oxy acid names and
  their formulas.
• The names of the ions are derived from the
  names of the acids.
• Keep in mind that the acids must be aqueous
  solutions.

63
Common Oxy Acids

• Acid Name
• HNO3 Nitric acid
• H2SO4 Sulfuric acid
• HClO3 Chloric acid (or iodic or bromic acid)
• H3PO4 Phosphoric acid
• H2CO3 Carbonic acid

64
Naming Complex Ions

• Once the list of acids is learned, the names of
  other acids and ions can be derived.
• Removal of the hydrogens in the acid as H ions
  results in ions that end in ate.
• HNO3 minus one H ion gives NO31-, the
  nitrate ion.
• The oxy acids that end in ic, produce ions that
  end ate.

65
Naming Complex Ions

• Sulfuric acid is H2SO4. Removing two H ions
  produces SO42-, the sulfate ion.
• Keep in mind that the formula of the ions must
  include the charge.
• If only one of the H ions is removed from
  sulfuric acid, HSO41- is produced. This is
  called the hydrogen sulfate ion, also commonly
  known as the bisulfate ion.

66
Naming Complex Ions

• Carbonic acid, H2CO3, produces two ions
• HCO31-, the hydrogen carbonate or
  bicarbonate ion
• and
• CO32-, the carbonate ion

67
Naming Complex Ions

• Some of the oxy acids previously listed also
  exist with one more oxygen in the formula.
• HClO3, HBrO3 and HIO3 , in aqueous solution are
  chloric, bromic and iodic acid respectively.
• Adding an oxygen to the formulas provides the
  formulas for the per root ic acid.
• HClO4 is perchloric acid. The ion, ClO41- is
  the perchlorate ion.

68
Naming Complex Ions

• Several of the oxy acids listed previously can
  have one less oxygen atom in the formula. These
  acids have names that end in ous, and ions that
  end in ite.
• HNO3 is nitric acid. HNO2(aq) is nitrous acid.
  The ion NO21- is the nitrite ion.

69
Naming Complex Ions

• Sulfuric acid, phosphoric acid, chloric, bromic
  and iodic acids all can have one less oxygen
  atom. The acids are sulfurous acid, phosphorous
  acid, chlorous acid, bromous acid and iodous
  acid.
• The ions are called sulfite, phosphite,
  chlorite, bromite and iodite ion.

70
Naming Complex Ions

• The halogen oxy acids HClO3, HBrO3, and HIO3
  also exist with two less oxygen atoms in the
  formula. The name of the resulting acid has the
  name
• hypo root ous acid.
• HClO(aq) is hypochlorous acid, and ClO1- is the
  hypochlorite ion.

71
Naming Complex Ions

• If you memorize the list of acids ending in ic,
  you can derive the names and formulas for many
  other acids and ions.
• Acid Name
• HNO3 Nitric acid
• H2SO4 Sulfuric acid
• HClO3 Chloric acid (or iodic or bromic acid)
• H3PO4 Phosphoric acid
• H2CO3 Carbonic acid

72
Naming Complex Ions

• In naming the ions from the acids on the list,
  remember that ic ? ate.
• If there is one additional oxygen atom, the acid
  has the name per root ic, and the ion has the
  name per root ate.
• If there is one less oxygen atom, the acid has a
  name ending in ous. The ions will have names
  ending in ite. (ous? ite)

73
Naming Complex Ions

• If an acid has two less oxygen atoms than the
  ic list, its name has the form hypo root ous.
  The ion will have the name hypo root ite.

74
Other Common Formulas

• CH3COOH Acetic acid
• CH3COO1- Acetate ion
• NH3 Ammonia
• NH4 Ammonium ion
• OH1- Hydroxide ion
• H3O Hydronium ion
• MnO41- Permanganate ion
• CrO42- Chromate ion
• Cr2O72- Dichromate ion

75
Percent Composition

• A chemical formula can be used to calculate the
  percent composition of a compound. Likewise, the
  percent composition can be used to determine the
  empirical formula of a compound. This is
  extremely useful when trying to determine the
  formula of a new, or unknown compound.

76
Chemical Composition

• Usually, the compound is combusted in the
  presence of oxygen. Any carbon in the compound
  is collected as carbon dioxide (CO2), and any
  hydrogen is collected as water (H2O).

77
Chemical Composition

• Similar techniques exist to analyze for other
  elements.
• The formula obtained for the compound is the
  simplest whole number ratio of the elements in
  the compound, or the empirical formula. It may
  differ from the actual formula. For example,
  hydrogen peroxide is H2O2, but chemical analysis
  will provide an empirical formula of HO.

78
Percent Composition

• To calculate the composition of a known
  compound, you determine the total mass of the
  molecule, and the mass due to each of the
  elements in the compound.
• by mass of element A
• total mass of A (100)
• molecular mass

79
Percent Composition

• To calculate the composition of a known
  compound, you determine the total mass of the
  molecule, and the mass due to each of the
  elements in the compound.
• by mass of element A
• total mass of A (100)
• molecular mass

80
Composition Problem

• Problem Calculate the percent composition of
  ammonia.

81
Determining Formulas

• It is generally more useful to obtain percent
  composition data (usually from a laboratory), and
  determine the empirical formula of a compound.
  This will be the simplest whole number ratio of
  the elements, and provides no information about
  the structure of the compound.

82
Determining Empirical Formulas

• If given composition
• 1. Assume a quantity of 100 grams of the
  compound.
• 2. Determine the number of moles of each element
  in the compound by dividing the grams of each
  element by the appropriate atomic mass.
• 3. To simplify the formula into small whole
  numbers, divide the moles of each element by the
  smallest number of moles.

83
Determining Empirical Formulas

• 4. If necessary, multiply each number of moles
  by a factor that produces whole number
  subscripts.
• 5. If you know the approximate molar mass of the
  compound, determine the molecular formula.

84
Composition Problem

• Caffeine contains 49.48 carbon, 5.15 hydrogen,
  28.87 nitrogen, and 16.49 oxygen. The compound
  has a molar mass of 194.2. Determine the
  empirical and molecular formula of caffeine.