Chemical Evidence

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Chemical Evidence
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Principles of Chemical Bonding

• Bonds form when electrons are exchanged or shared
• Only valence (outer) electrons are used
• Each atom tries to achieve a noble gas electron
  configuration

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Chemical Evidence

• Ions and Ionic Compounds

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Ionic Compounds

• Ionic solids are a vast 3D array of positive and
  negative ions that has no discrete molecular
  units
• The electrostatic attraction of positive and
  negative charges holds it together

• Example KBr

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Ions from Metals

• Metals lose electrons to match the number of
  valence electrons of their nearest noble gas
• Mg --gt Mg2 Ne3s2 --gt Ne
• Names of ions are the same as their elements
• Positive ions form when the number of electrons
  are less than the number of protons
• Group 1 metals --gt ion 1
• Group 2 metals --gt ion 2
• Group 13 metals --gt ion 3
• Group 14 metals --gt ion 4

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Some Typical Ions with Positive Charges (Cations)

• Group 1 Group 2 Group 13 Group 14
• H Mg2 Al3 Sn4 and Sn2
• Li Ca2 Pb4 and Pb2
• Na Sr2
• K Ba2
• Cations have the same names as the element
• Examples name of Mg2 magnesium, name of Sn4
  tin(IV) and Sn2 tin(II)

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Transition Metal Ions Names

• 1 2 1 or 2
  2 or 3
• Ag Zn2, Cd2 Cu, Cu2 Fe2, Fe3
• Hg2, Ni2 Co2, Co3
• Cr2, Cr3
• Names of ions with more than one possible charge
  include the charge as a roman numeral in
  parenthesis
• For example tin(II) Sn2 and tin(IV) Sn4
• or lead(II) Pb2 and lead(IV) Pb4

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Ions from Nonmetals

• Nonmetals gain electrons to match the number of
  valence electrons of their nearest noble gas
• Cl --gt Cl-1 Ne3s23p5 --gt Ar
• Names of ions end in -ide Cl-1 chloride
• Negative ions form when the number of electrons
  is greater than the number of protons
• Group 15 nonmetals --gt ion -3
• Group 16 nonmetals --gt ion -2
• Group 17 nonmetals --gt ion -1

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Some Typical Anions Names

• Group 15 Group 16 Group 17
• N-3 (nitride) O-2 (oxide) F-1 (fluoride)
• P-3 (phosphide) S-2 (sulfide) Cl-1
  (chloride)
• Se-2 (selenide) Br -1 (bromide)
• I-1 (iodide)

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Formulas of Ionic Compounds

• The empirical (simplest) formula of an ionic
  compound is one in which the ions of opposite
  charge combine in a ratio that results in an
  electrically neutral substance.
• Example of the net charge approach Mg2 2
  Cl-1 --gt MgCl2
• 1 x Mg2 2 and 2 x Cl-1 -2, so the net
  charge 0 2 (-2)
• Example of the cross over approach
• Mg2 O-2 --gt Mg2O2 ---gt MgO

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Formulas of Ionic Compounds

• Polyatomic ions are placed in parentheses
• Example Ca2 PO4-3 --gt Ca3(PO4)2
• 3 x Ca2 6 and 2 x PO4-3 -6, net charge 0
• Example Ca2 PO4-3 --gt Ca3(PO4)2

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Naming Ionic Compounds

• The cations are named first, then the anions
• No indication of the number of each ion is given
• Example Ca3(PO4)2 calcium phosphate
• Memorize the names and formulas of all ions

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Learning Check

• Write the missing name or formula for each
• A. ZnCl2 _______________
• B. ______ aluminum oxide
• C. (NH4)2SO4 _______________
• D. ______ copper(II) nitrate

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Chemical Evidence

• Covalent Compounds

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Covalent Compounds

• Formed from non-metallic elements
• The sharing of electrons to create covalent bonds
  is called valence bond theory. Atoms share
  electrons to achieve a noble gas electron
  configuration
• Example H H ---gt H2 or HH or H-H 1s1 1s1
  ---gt 1s2 (He electron configuration for each
  nucleus)

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Naming Covalent Compounds

• In covalent compounds, the elements are named in
  the order that they appear in the formula
• Greek prefixes are used to indicate the number of
  each type of atom (the mono prefix is not used
  for the first element)
• The last element in the formula ends in -ide
• Examples CO carbon monoxide and N2O4
  dinitrogen tetraoxide

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Learning Check

• Write the missing name or formula for each
• A. CO2 _______________
• B. ______ phosphorus trichloride
• C. N2O4 _______________
• D. ______ dihydrogen monoxide

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Covalent Bonding Structure
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Covalent Bonding

• Three basic principles of bonding
• bonds form when electrons are exchanged or shared
• only valence (outer) electrons are used
• in this process each atom tries to achieve a
  noble gas electron configuration
• Covalent Bonds
• formed from non-metallic elements
• the sharing of electrons to create covalent bonds
  is called valence bond theory
• Example H H ---gt H2 or HH or H-H 1s1 1s1
  ---gt 1s2

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Lewis Symbols

• The number of valence electrons for an atom can
  be obtained by using the second or only digit of
  elements Group number.
• Electron-dot symbols (Lewis symbols) are used to
  represent the valence electrons around an atom.
• One dot for each valence electron is placed
  around the element symbol at the four compass
  points.
• All elements in a Group have the same
  electron-dot structures. They are isoelectronic.

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Lewis Structures of Compounds

• Example Cl2 gt Ne3s23p5 Ne3s23p5
• Only the unpaired electrons in the unfilled 3p
  orbitals creates a bond.
• The electrons not used in bonding are called
  lone pairs.

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Lewis Structures of Compounds

• Example O2 gt He2s22p4 He2s22p4
• Shared pairs between the unfilled 2p orbitals
  creates two bonds.
• More than one bond between two atoms is possible
  if this completes an octet.Multiple bonds are
  common between period two elements, particularly
  C, N and O.

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Drawing Lewis Structures

• 1. Count the total number of valence electrons in
  the molecule. To write the electron-dot structure
  of ions you must add or subtract electrons from
  the total, corresponding to the charge on the
  ion. Electrons are added to anion and subtracted
  from cations.
• 2. Identify the center atom using one of the
  following rules
• a. The element of which there is the fewest in
  the formula may be the center atom.
• b. The element closest to the center of its
  period may be the center atom since it can form
  the largest number of bonds.
• c. Elements of the same group but with a
  higher atomic number tend to be center atoms.
• 3. Write the element symbol of the center atom
  and place the outer atom element symbols around
  it. If the molecule has a chain or ring, write
  the element symbols for all the chain or ring
  atoms in a linear or circular pattern. Chains
  and rings are the most common for compounds
  containing C and H in the same formula or for
  compounds with H and more than one B, Si, N, O, P
  or S.
• 4. Draw single bonds from the center atom(s) to
  each of the outer atoms. Subtract the number of
  electrons used to create the bonds (2 per bond)
  from the total. Note that because its valence
  (number of bonds) is 1, hydrogen is always an
  outer atom.
• 5. If there are any electrons left, give the
  outer atoms enough electron pairs to fill their
  octets. Subtract the electrons used in this step
  from the total. Note that hydrogen requires only
  two electrons to complete its noble gas shell and
  never gets an octet.
• 6. If there are still electrons left, give the
  center atom the remaining electrons to help
  complete its octet.
• 7. If the center atom still does not have an
  octet at this point, create multiple bonds
  between it and the outer atom(s) by sharing one
  of the lone pairs on the outer atoms with the
  atom at the center.

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Learning Check

• Draw the Lewis Structures of the following
  molecules
• A. CO3-2
• B. PCl3
• C. SiH4

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3-D Shapes of Molecules

• Determined by the positions of the valence
  electrons around the central atom
• The 3-D shape is a result of repulsions of bonded
  pairs and lone pairs of electrons
• Example H2O or H-O-H is bent and NOT linear
• You must use VSEPR theory (valence-shell-electron-
  pair repulsion) to determine 3-D shape of
  molecules

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VSEPR Electron Pair Geometry

• Valence Shell Electrons wish to minimize the
  Electron Pair Repulsions between them.
• The repulsions between independent regions of
  electron density around an atom produce
  electron pair geometries.
• Electron pair geometries include linear(2),
  trigonal planar(3), tetrahedral(4).
• The electron pair geometries determine the
  approximate bond angles for a structure.

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VSEPR Molecular Geometry

• A single bond, a multiple bond and a lone pair
  each count as one independent region of electron
  density.
• Electron pair geometries are further broken down
  into molecular geometry subclasses.
• The names for these subclasses is determined by
  the 3-D arrangement of atoms.

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Using the VSEPR Model

• 1. Draw the Lewis structure.
• 2. Count the number of each type of electron
  region around the center atom. Note that
• a. Each single, double or triple bond counts
  as only one region (a bonding region).
• b. Each lone pair counts as one region, an
  unshared region
• 3. Use the table provided to determine the
  molecular geometry.

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Learning Check

• Determine the electron pair and molecular
  geometry of each of the following molecules
• A. SiH4
• 1) tetrahedral 2) pyramidal 3) angular
• B. PCl3
• 1) trigonal planar 2) pyramidal 3) angular
• C. CO3-2
• 1) trigonal planar 2) pyramidal 3)
  angular