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Chemical Evidence

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... electron configuration by having filled s and p subshells (the octet/duet rule) ... Ionic solids are a vast 3D array of positive and negative ions that has no ... – PowerPoint PPT presentation

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Title: Chemical Evidence


1
Chemical Evidence
2
Principles of Chemical Bonding
  • Bonds form when electrons are exchanged (ionic)
    or shared (covalent)
  • Only valence (outer) electrons are used
  • Each atom tries to achieve a noble gas electron
    configuration by having filled s and p subshells
    (the octet/duet rule).

3
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

4
Formation of Monoatomic Ions
5
Formation of Monoatomic Ions
  • Metals lose electrons to match the number of
    valence electrons of their nearest noble gas
  • Mg --gt Mg2 Ne3s2 --gt Ne
  • Names of monoatomic cations are the same as their
    elements Mg2 magnesium
  • If more than one possible charge Sn2 tin(II)
  • Nonmetals gain electrons to match the number of
    valence electrons of their nearest noble gas
  • Cl --gt Cl-1 Ne3s23p5 --gt Ar
  • Names of anions end in -ide Cl-1 chloride

6
Formation of Monoatomic Ions
7
Other Ions Their Names
Transition Metal Ions Cu2 copper(II) Cu1
copper(I) Ag1 silver Au1 gold(I) Au3
gold(III) Fe3 iron(III) Fe2 iron(II) Hg22
mercury(I) Hg2 mercury(II)
8
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

9
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

10
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

11
Learning Check
  • Write the missing name or formula for each
  • A. ZnCl2 _______________
  • B. ______ aluminum oxide
  • C. (NH4)2SO4 _______________
  • D. ______ copper(II) nitrate

12
Covalent Bonding Structure
13
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
  • the shared electrons are concentrated between the
    two atoms

14
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

15
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16
Learning Check
  • Write the missing name or formula for each
  • A. CO2 _______________
  • B. ______ phosphorus trichloride
  • C. N2O4 _______________
  • D. ______ dihydrogen monoxide

17
Lewis Symbols
  • The number of valence electrons 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 Lewis
    symbol.

18
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.

19
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.

20
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.

21
Learning Check
  • Draw the Lewis Structures of the following
    molecules
  • A. SO2
  • SiH4

22
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 is NOT linear
  • You must use VSEPR theory (valence-shell-electron-
    pair repulsion) to determine 3-D shape of
    molecules

23
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.

24
VSEPR Electron Pair Geometry
25
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.

26
VSEPR Molecular Geometry
27
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.

28
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29
Learning Check
  • Determine the electron pair and molecular
    geometry of each of the following molecules
  • SO2
  • B. SiH4
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