This ppt explains the difference between ionic and covalent compounds

2. What is a Chemical Formula?
• When elements combine to form compounds a
chemical reaction occurs creating a new compound
with properties that are different from the original
elements.
• When a new compound is formed it has a new
chemical formula. A chemical formula tells what
elements a compound contains and the exact
number of atoms of each element in a unit of that
compound.
• Example: 2Na + Cl2  2NaCl

3. Why do atoms form compounds?
• The electric forces between electrons and
protons (which are oppositely charged) hold
atoms and molecules together which cause
compounds to form.
• But, not all atoms form compounds readily.
Noble gases already have all the electrons they
need! In other words, their outer energy level is
completely filled.
• When an atoms outer energy level is filled, they
are considered chemically stable.

4. How do atoms become chemically
stable?
• When an atom is not chemically stable, it
must bond with another atom to become
stable. To do this, an atom must gain or lose
electrons becoming an ion. An ion is a
charged particle because it now has either
more or fewer electrons than protons.
• A cation is an atom that has given away an
electron in its outer energy level (valence
electron) and is now a positively charged.

5. Howdo atoms become chemically
stable?
• An anion is an atom that has gained an
electron in its outer energy level and is
now a negatively charged particle.

6. Types of Chemical Bonds
• A chemical bond is the force that holds
atoms together in a compound.
• There are 2 main types of chemical
bonds: Ionic and Covalent bonds

7. Ionic Bonds
• An ionic bond is the bond that occurs when a
metal gives away a valence electron (becomes a
cation) and a nonmetal gains the valence
electron (becomes an anion). In other words,
it’s a bond between a metal and a nonmetal!
• Form crystalline solids
• High melting and boiling points
• Conduct electricity when melted
• Many are soluble in water
• Strong attraction between positive and negative
ions
• Brittle and hard

8. Bozeman video
• http://www.bozemanscience.com/chemi
cal-bonds-covalent-vs-ionic

9. Ionic Bonds
• There are 2 ways to show ionic bonding:
• Electron – dot method:
• The criss cross method:
Make sure to reduce when necessary!

10. Ionic Bonds
• Sometimes compounds only involve 2
elements. These are called Binary
compounds.
• To name binary ionic compounds:
1. Write the name of the metal first
(capitalized).
2. Write the name of the nonmetal (lower
case).
3. Change the ending of the nonmetal to –ide.
• Example: NaCl  Sodium chloride
• MgCl2  Magnesium chloride

11. Ionic Bonds
• Oxidation number is the charge on the
ion. This is what we use for the criss
cross method. For metals, it is the
number of valence electrons available for
bonding. For nonmetals, it is the number
of valence electrons needed to fill the
outer energy level.
• Example: Na+1 and Cl-1
• Mg+2 and I-1

12. Ionic Bonds
• Some compounds involve more than 2 elements. These
compounds contain polyatomic ions. Polyatomic ions are
a covalently bonded group of atoms that ionically bond
with an ion of opposite charge.
• To name compounds with polyatomic ions:
1. Write the name of the positively charged ion first. If it is
a single metal, write it’s name capitalized. If it is a
polyatomic ion, use it’s name also capitalized.
2. Write the name of the negatively charged ion second. If
it is a single non-metal, remember to change the ending
to –ide. If it is a polyatomic ion, use it’s name and DO
NOT change the ending.
• Example: NaSO4 Sodium sulfate
• NH4Cl Ammonium chloride

13. Covalent Bonds
• Covalent bonds form between
nonmetals. It is the attraction they form
when sharing electrons.
• Once a covalent bond is formed, the
resulting compound is referred to as a
molecule.
• Single, Double or Triple bonds can form
between 2 nonmetals.

14. Covalent Bonds
• Electrons are not always shared equally
between the atoms in a covalent bond.
• Polar molecule is a molecule that has a
slightly positive end and a slightly negative
end.
• Nonpolar molecule is a molecule in which
electrons are shared equally.
•

15. Covalent Bonds
• Properties of Covalent bonds include:
• Poor conductors of heat and electricity
• Lower melting and boiling points
• Solid, liquids or gases
• Strong bonds between atoms but weak
attraction between molecules
• Soft
• Do not dissolve in water

16. Covalent bonds
• To draw covalent bonds:
• Place the element that is LEAST represented
in the middle and radiate all other elements
around the central atom.
• Only place valence electrons that will be
shared between the 2 atoms. Any electrons
not being shared must go on the “outside”
of the atom.
• Connect shared electrons with dashes.
• Covelant bonds DO NOT REDUCE!

17. Covalent Bonds
Covalent compounds use prefixes to
indicate the number of atoms per element
is in the compound.

18. Covalent Bonds
1. The first element NEVER uses the prefix mono but
does use all other prefixes. Make sure to capitalize
it!
2. Write the appropriate prefix for the second
element.
3. Write the root of the 2nd element’s name. If the
prefix ends is o/a and the element name also
begins with o/a, drop one of the o/a. Ex) mono &
oxygen becomes monoxide
4. Change the ending of the 2nd element to –ide.
Example) CO  Carbon monoxide
CO2  Carbon dioxide
P2O5  Diphosphorus pentoxide

19. Bozeman videos
• http://www.bozemanscience.com/namin
g-compounds-part-1
• http://www.bozemanscience.com/namin
g-compounds-part-2

20. Diatomic Molecules
• There are 7 nonmetals that are highly
unstable. Because of this, they bond
readily with another like atom. These are
referred to as Diatomic Molecules
because they are found naturally in pairs.
• Atoms include: H, N, O, F, Cl, Br, and I. If
you look closely (with the exception of
H), they make a “7” on the periodic table!

21. What is a hydrate?
• A hydrate is a compound that has water
chemically attached to its ions and
written into its chemical formula.
• To name, name the ionic compound to
which it is attached. Then use the prefix
for the number of hydrates and then
simply end with the word “hydrate”.
• Example: CoCl2*6H2O  Cobalt chloride
hexahydrate