Bonding
Ionic and covalent
Key Terms1
 Chemical formula– the combination of
chemical symbols and subscripts to
indicate what the elements are in the
compound and how many atoms of each
element are in the compound
 Example: H2O= two hydrogen atoms and 1
oxygen atom.
Key Terms2
 The octet rule– each atom wants to have
8 electron in its outer most energy level.
 Atoms can share, take, or give away
electrons to accomplish this.
 Valence electrons– electrons in the outer
most energy level that are responsible for
the reactivity of that atom.
Key Terms3
 Lewis structure (electron-dot notation)–
valence electrons are shown as dots around
the element’s symbol. Label the # of Ve- on
your periodic table. Only used for main block
elements
 Because each orbital can hold two electrons,
electrons are grouped in pairs forming the shape of a
box around the element’s symbol.
 Paired electrons can also be represented by a dash
instead of dots if they are being shared in a
compound.
Key terms4
Ions
 Ion– any atom that has given up or taken
electrons to create a positive or negative
charge.
 This is done to fill the highest energy level.
 Cations (cat-ion) – any element that has given
away its electrons to become a positively
charged ion. Cations are metals.
 Anions– any element that has taken electrons
to become a negatively charged ion. Anions
are nonmetals.
Atoms and charges
 What happens when an atom gains or looses
an electron?
 The atom becomes charged!
 Example: Copper has an atomic number of 29.
This means copper has 29 electrons(-) and 29
protons(+). If copper were to loose two
electrons, what would copper’s charge be?
29 Protons (+)
+ 27 Electrons (-)
2 Protons left over, each proton has a positive
charge so the charge of copper would be +2!
Practice Problem
 Oxygen has an atomic number of 8 and
an atomic mass of 16. How many
neutrons does oxygen have?
 Answer: 16= N+8, N=8
 What would Oxygen’s charge be if it
gained two electrons?
 Answer: 8(+ protons) + 10 (- electrons)=2
Ionic Bonding
Ionic Bonding
 Ionic bonding– any bond between metals and
nonmetals (cations and anions)
 Charges are based on how many Ve- are
needed to fill the outer shell or drop to the
previous full shell. Label this on your table.
 The charges must cancel each other out.
 Example: Na (+1) and Cl (-1)=NaCl (0)
 Example: Ca (+2) and F (-1)=CaF2 (0)
 Ionic compounds are usually solids and in a
crystal structure (crystal lattice).
Ionic Compounds
 Both ions should have complete outer shells
after bonding.
 Both elements should have noble gas
electron configurations
 When naming, the first element always
stays the same, but the last element should
end with –ide
 Ex. MgO= Magnesium Oxide instead of
Magnesium Oxygen
 Ex. CaCl2= Calcium Chloride vs Chlorine
Covalent Bond Key Terms1
 Molecule: a group of atoms held together by
covalent bonds
 Covalent bond: when atoms share electrons
 Nonpolar covalent bond: electrons are
equally shared by all atoms and the
electrical charge is balanced
 Polar covalent bond: electrons are not
shared equally and there is an imbalance in
the electrical charge surrounding the
molecule.
Polar vs. nonpolar
 Polar
nonpolar
Key Terms2
 Polar bonds: when atoms in a molecule
have an uneven electron distribution.
 Bond length: average distance between two
bonded atoms
 Bond energy: energy required to break a
chemical bond and form neutral isolated
atoms.
 When the bond length gets shorter, the
bond energy gets higher
Key Terms3
 Lewis structures must be used to create
covalent compounds (molecules). Sorry no short
cuts this time.
 Single bond: when only two electrons are
being shared between two atoms
 Double bond: when 4 electrons are being
shared between two atoms
 Triple bond: when 6 electrons are being
shared between two atoms
Covalent Bonds
Covalent Bonding2
 There are no ions involved with
covalent bonds which means no
charges.
 This is a bond between two
nonmetals.
 Electrons are shared.
 The magic number is still 8.
Covalent properties
 covalent bonds can produce solids, liquids,
or gaseous molecules
 they are poor conductors of electricity
 They have low melting points and boiling
points
 They are usually very dull in appearance
Covalent Nomenclature
 When naming covalent compounds, you
MUST use prefixes for the first and second
words
 The only exception is if you only have one
atom for the first element.
 the less electronegative (furthest to the left
on the p/t) element is given first and its full
element name is written
 -ide is still needed at the end of the second
element as well as a prefix
Covalent molecule
prefixes
1.
2.
3.
4.
5.
6.
7.
8.
9.
10.
MonoDiTriTetraPentaHexaHeptaOctaNonaDeca-
Covalent nomenclature
 Examples
 Dihydrogen monoxide (H2O)
 Sulfur trioxide (SO3)*there is only one sulfur
so no prefix is needed.
 Trisilicon tetranitride (Si3N4)
 Dinitrogen Pentaoxide (N2O5)