Nomenclature Binary Ionic Compounds

1
Nomenclature- Binary Ionic Compounds

• Binary compounds are compounds with two different
  kinds of atoms.
• Naming Binary Ionic Compounds
• name metal name stem of nonmetal name -ide
• The stem names and ionic symbols for some common
  nonmetals are given in the following table

2
Examples of Binary Ionic Compounds

• Name K2O
• name metal name nonmetal stem -ide
• name potassium ox -ide potassium oxide
• Name Mg3N2
• name metal name nonmetal stem -ide
• name magnesium nitr -ide magnesium
  nitride
• Name BeS
• name metal name nonmetal stem -ide
• name beryllium sulf -ide beryllium
  sulfide
• Name AlBr3
• name metal name nonmetal stem -ide
• name aluminum brom -ide
• name aluminum bromide

3
Metals forming ions more than one charge

• Some metal atoms (transition and inner-transition
  elements) more than one type of charged
  ionCobalt - Co2 and Co3
• Binary compounds same pattern and
• the number of positive charges on the metal ion
  indicated by a Roman numeral in parentheses
  following the metal name.
• CoO and Co2O3 - cobalt ions with 2 and 3
  charges respectively.
• cobalt (II) oxide and cobalt (III) oxide.
• An older system only 2 different ions-ous and
  ic attached to the stem of the metal
  name.(Non-English names for elements with
  symbols derived from non-English names)

4
Nomenclature- Covalent Compounds

• Covalent compounds most often between
  representative elements classified as non-metals.
• name name of least electronegative element
  stem of more electronegative element -ide.

• Number of each type of atom in the molecule
  indicated by Greek prefixes.
• The prefix mono is not used when it appears at
  the beginning of the name.

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Examples of binary covalent compounds

• SO2 name sulfur di- ox -ide sulfur
  dioxide
• XeF6 name xenon hexa- fluor -ide xenon
  hexafluoride
• H2O name di- hydrogen mono- ox -ide
• dihydrogen monoxide (also known as water).
  Note, the final o of mono- was dropped for ease
  of pronunciation.

6
Ionic Compounds with polyatomic ions

• Formula rules same as those for binary ionic
  compounds.
• Symbol for the metal first, followed by the
  formula for the negative polyatomic ion. Equal
  numbers of positive and negative charges must be
  represented by the formula.

When more than one polyatomic ion is required in
the formula, parentheses are placed around the
polyatomic ion before the subscript is inserted.
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Ionic Compounds with polyatomic ions

• Examples
• Compound containing K and ClO3-
• Compound containing Ca2 and ClO3-
• Compound containing Ca2 and PO43-
• Names of ionic compounds with a polyatomic ion
  name name of metal name of polyatomic ion
• KClO3 - potassium chlorate
• Ca(ClO3)2 - calcium chlorate
• Ca3(PO4)2 - calcium phosphate
• CaHPO4 - calcium hydrogen phosphate

KClO3
Ca(ClO3)2
Ca3(PO4)2
8
Noble gas configuration

• Noble gases have completely filled s and p
  subshells of their valence (outermost) shell
  (except He)
• Octet rule - rule for predicting electron
  behavior in reacting atoms. Atoms will gain or
  lose electrons to achieve an outer electron
  arrangement identical to that of a noble gas.
• Learning check 4.1

9
Lewis structure

• Lewis structure (also called electron-dot
  formula)-
• Nucleus and all electrons around the nucleus
  except those in the valence shell- Elemental
  symbol
• Valence shell electrons- Dots arranged around the
  symbol.
• Number of valence electrons
• Valence electrons as those having the largest n
  value in the configuration.
• For representative elements-
• Number of valence electrons same as the Roman
  numeral group number.
• Example 4.2, Learning check 4.2

10
Ionic bonding

• Simple Ions- Atoms acquire a net positive or
  negative charge by losing or gaining one or more
  electrons
• The attractive force that holds together ions of
  opposite charge is called as ionic bond.
• Metals lose electrons
• Non-metals gain electrons
• Example 4.4
• Magnesium, Mg, has two valence electrons which it
  loses to form a simple ion with a 2 electrical
  charge. The ion is written as Mg2.
• Oxygen, O, has six valence electrons. It tends
  to gain two electrons to form a simple ion with a
  -2 electrical charge. The ion is written as O2-.
• Bromine, Br, has seven valence electrons. It
  tends to gain one electron to form a simple ion
  with a -1 electrical charge. The ion is written
  as Br -.

11
Ionic bonding

• Representative metals form vely charged
  ionsCharge number (Roman numeral) of the
  group.
• Representative nonmetals form vely charged
  ionsCharge 8 - the number (Roman numeral) of
  the group.
• Example 4.5, Learning check 4.5
• Atoms are changed into ions with noble gas
  configurations
• Ionic bonds- attractive force between oppositely
  charged ions
• Ionic Compounds- The substances formed when ionic
  bonds form between positive and negative ions are
  called ionic compounds.
• Formulas- ratio of the ve and ve ions ? total
  charge zero
• Example 4.6

12
Covalent Bonding

• Covalent bonding- sharing of electrons to satisfy
  octet rule of each atom
• Example- F2
• Representation shared pair of dots or single
  line
• Bonding due to overlap of atomic orbitals

molecular orbital formed
13
Covalent bonding

• Covalent bond- Attractive force between two atoms
  both attracted to a shared pair of electrons
• Sharing between identical atoms (homoatomic)
  Cl2, O2 and N2
• Sharing between different atoms (heteroatomic)
  H2O, and CH4
• Example 4.10, Learning check 4.10, O O ? O2
  (more than two electrons shared)

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Examples of covalent bonding
15
Lewis structures for covalent molecules

• Step 1 Molecular formula and number of atoms
• CO2 (1 atom C, 2 atoms O)
• Step 2 Initial structure using the connecting
  pattern
• O C O
• Step 3 Total number of valence shell electrons
  (Use periodic table group number)
• O6, C4, O6. Total 16
• Step 4 Place a pair of electrons between each
  bonded pair of atoms. Subtract from total valence
  electrons. Use remaining to complete octet of
  atoms
• OCO ? 16-2-2 12 electrons left

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Lewis structures for covalent molecules

• Step 5 Check octet rule for all atoms. If yes,
  Lewis structure is complete. If not, move
  unbonded pairs to positions between bonded atoms
  to complete octets.
• Octet of O atoms not complete
• Moving one unbonded pair of electrons from C atom
  between the O and C atoms
• or
• Double bonds- Sharing of two pairs of electrons
• Triple bonds- Sharing of three pairs of electrons
• Example 4.11, Learning check 4.11

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Shapes of molecules and polyatomic ions

• not flat 2-D
• distinctive 3-D shapes
• Valence-shell electron-pair repulsion theory or
  VSEPR theory
• Mutual repulsion of electron pairs in valence
  shells
• Applied to central atom (bonded to two or more
  other atoms) of a molecule or an ion to predict
  the shape
• Step 1 Draw the Lewis structure
• Rules for VSEPR
• Rule 1 All valence-shell electron-pairs
  considered (bonding and non-bonding)
• Rule 2 Double or triple bonds with central atom
  treated as a single pair for prediction

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VSEPR theory

• Two pairs of electrons One pair on each opposite
  side of central atom
• Example CO2
• Central atom carbon
• Rule 1 Consider the two electron-pairs each side
• Rule 2 Treat each double bond as single pair
• Shape Two pairs located on opposite sides of C
  atom. A linear molecule

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VSEPR theory

• Three pairs of electrons Triangle
• Example ozone (O3)
• Central atom oxygen
• Rule 1 Consider the four electron-pairs (3
  bonding, 1 non-bonding)
• Rule 2 Treat the double bond as a single pair
• Shape The double-bonded pairs, 1 bonding pair
  and 1 non-bonding pair arrange as a flat triangle

20
VSEPR theory

• Four pairs of electrons Tetrahedral
• (four corners of a tetrahedron)
• Example water (H2O)
• Central atom oxygen
• Rule 1 Consider the four electron-pairs (2
  bonding, 2 non-bonding)
• Rule 2 No double bond
• Shape The two bonding pairs and 2 non-bonding
  pairs arrange as four corners of tetrahedron

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VSEPR theory

• Five electron-pairs
• Six electron-pairs
• Learning check 4.13
• Example 4.14

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Polarity of molecules

• Homoatomic molecules- shared electron pair
  attracted equally.
• Nonpolar covalent bonds
• Electronegativity is measure of tendency of an
  atom to attract shared electrons of a covalent
  bond
• Results in shifting of shared
• electrons towards the more
• electronegative atom called
• Bond polarization
• Polar covalent bonds
• More electronegative- partial negative charge
  (d-)
• Less electronegative- partial positive charge (d)

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Polarity of molecules

• Example 4.15, learning check 4.15
• Extent of bond polarization- electronegativity
  differences (?EN)
• (?EN)0.0 ? Nonpolar covalent
• (?EN) 2.1 ? Ionic
• 0.0lt(?EN) lt2.1 ? Polar covalent
• Example 4.16
• Learning check 4.16