Transcript Ch. 3 HW- 3.18, 3.21, 3.32, 3.33, 3.38, 3.39, 3.43, 3.52, 3.53, 3.56, 3.59, 3.61

Ch. 3
HW- 3.18, 3.21, 3.32, 3.33, 3.38,
3.39, 3.43, 3.52, 3.53, 3.56, 3.59,
3.61
• Compounds are held together with
chemical bonds
• There are two types of bonds, ionic bonds
and covalent bonds.
• Before we talk about them, lets talk about
why they form.
• Octet Rule- Main group elements tend to
undergo reactions to achieve a complete
outer shell of electrons. Normally this is 8.
Ions
• If an atom gains electrons, it becomes
negatively charged and called an Anion
• If an atom loses electrons, it becomes
positively charged and called an Cation
• Remember, the number of protons and
neutrons stay the same during most
chemical reactions
Exceptions to the Octet Rule
1) Elements of periods 1 and 2 cannot have
a charge greater than +2
2) Octet rule does not apply to transition
elements
Atoms and Ions
• Atoms and their ions are completely
different with different physical and
chemical properties
• Examples:
Naming Ions
• Most names come from a system
developed by the International Union of
Pure and Applied Chemistry (IUPAC)
• These names are called systematic names
• Some ions are referred to by their
common names as well
Naming Cations
• Elements of Group 1, 2, and 3 form only
one type of ion, so we can name the
element and add the word “ion”
• Examples:
Naming Cations
• Transition Metals form more than one ion
so we name the element and add roman
numerals to designate the charge
• Examples:
• Exception:
Older system (common names)
• When a metal can form 2 different ions:
• The suffix “-ous” is used to designate the
one with the smaller charge
• The suffix “-ic” is used to designate the
one with the large charge
• Example:
• See table 4.2, page 97!!!! (know these)
Naming Anions
• Monoatomic Anions are named by adding
“-ide” to the stem part of the name
• Examples:
• Polyatomic anions, anions that contain
more than one element, are more
complicated.
• See/Know Table 4.4. page 97
Back to Bonding
• Atoms bond together in such a way that
each atom achieves a valence electron
configuration of the nearest noble gas.
• There are two ways this can happen:
Ionic Bonds
1) Atoms may gain or lose electrons to
achieve a filled valence shell, thus
becoming ions.
An Ionic Bond results from the force of
attraction between a cation and an anion
Covalent Bond
2) Atoms may share electrons to fill the
valence shell
A Covalent Bond results from the force of
attraction between 2 atoms that share
one or more pairs of electrons
• Ionic bonds usually form between metals
and nonmetals.
• Example:
• Covalent bonds usually form between
nonmetals or between a nonmetal and a
metalloid.
• Example:
Electronegativity
• Electronegativity-a measure of an atoms
attraction for the electrons it shares in a
chemical bond.
• Fluorine is the most electronegative element.
• All other elements are assigned values in
relation to fluorine
• Trend in periodic table
• NOTE: Ionization energy and electronegativity
are different!!!!
Ionic Compounds
• These are compounds formed by ionic bonds
• Ionic bonds form by the transfer of one or more
valence electrons from an atom of lower
electronegativity to the valence of an atom with
higher electronegativity
• The atom that gains the electron becomes an
anion
• The atom that loses the electron becomes a
cation
• The compound formed by the combination
of positive and negative ions is called an
ionic compound
• Ionic bonds usually occur when the
difference in electronegativity is 1.9 or
more.
• Example:
• Single head arrows are used to show the
transfer of 1 electron
• Ionic compounds do not consist of
molecules, but their formula still gives the
definite ratio of atoms
• Ionic compounds are overall neutral so we
can predict the formula by balancing the
charges
• Examples
– Write the formula for:
• Lithium Ion and Bromide Ion
• Barium (Ba) and Iodide Ion
Naming Ionic Compounds
• Simply give the name of the cation first,
then the anion.
• For Binary Ionic Compounds (contain only
two elements)
– Example:
• Binary Compounds that contain metal with
multiple ions, use the Roman Numerals
• Ionic Compounds containing Polyatomic
ions, name positive one first, followed by
negative one, as separate words
• Examples:
Covalent Bonds
• Covalent bonds- formed when electrons
pairs are shared between two atoms
• These usually occur between two
nonmetals or between a nonmetal and
metalloid
• The electron pair is shared by two atoms
and at the same time fills the valence of
each atom
• A bond formed by sharing a pair electrons
is called a single bond
• Example:
• Lines are used to represent a shared pair
of electrons
Two Types of Covalent Bonds
• Nonpolar Covalent Bonds- electrons in
the bond are shared equally
• Polar Covalent Bonds- electrons in the
bond are not shared equally
• As a result of polar covalent bonds one
atom ends up with a partial negative
charge and the other ends up with a partial
positive charge.
• Example:
• These molecules are said to have a dipole
Drawing Lewis Structures
1) Determine the number of valence
electrons in the molecule
2) Decide on the arrangement of atoms in
molecule
3) Connect the atoms with single bonds
then arrange the remaining electrons in
lewis dot pairs so that each atom has a
complete outer shell
4) A pair of electrons involved in a covalent
bond, called bonding electrons, are shown
with a single line. Nonbonding electrons
are shown as pairs of lewis dots
5) Single bonds share one pair of electrons
and are shown with one line. Double
bonds share two pairs of electrons and are
shown with two lines. Triple bonds share
three pairs of electrons and are shown
with three lines.
Examples
Naming Binary Molecular
Compounds
• 1) Name the less electronegative element
first.
• 2) Use prefixes mono-, di-, tri-, tetra-, etc, to
show the number of atoms of each element.
(Note: Mono is omitted for the 1st word!)
• Examples:
Bonds Angles and Shapes of
Molecules
• We can predict bond angles using the
Valence-Shell Electron-Pair Repulsion
Theory, VSEPR.
• Valence electron repel each other.
• We treat each single bond, double bond,
triple bond, and lone pair as a single
electron density region.
VSEPR
Regions
4
Geometry
tetrahedral
Angles
109.5
3
Trigonal Planar
120
2
Linear
180
Geometry vs Shape
• Geometry and Shape of molecules are
different!!
• Geometry considers all electron density
regions.
• Shape only considers bonded atoms
Examples
Polar and Nonpolar Compounds
• Polar Compounds- posses an overall
dipole over the entire molecule
• A molecule is polar if:
– It has polar bonds
– It has a center of partial positive charge and a
center of partial negative charge in different
locations