7.1 – Putting Atoms Together
Def. Molecule –
Ex. Air = roughly 80% _______________ molecules, _____ oxygen molecules, and trace amounts of H2O and CO2
Def. Chemical Formula –
– small numbers written ____________________, in chemistry usually indicates a charge 
– small number ________________________, tells us how many atoms there are 
Diatomic Elements –
Molecular Compounds
Def.
 Involves two _______________ (includes hydrogen) elements joining together and sharing their ________
electrons to fill their outer __________________ (shell).
2H+1O

H2O
Ex. Illustration to the right shows two
_______________ molecules and a single
_______________ molecule. Hydrogen wants
to have a full _______________ shell (first
shell can hold a max. of __) meanwhile
the oxygen atom also wants to have a
full valence shell (2nd shell can hold __).
Therefore a _______________________________
__________________ occurs.
 A key aspect of molecular compounds is the fact they are ______________________ ____ ____________________.
Only certain combinations in ________________ _______________________ are found in nature.
Ex.
 The above example highlights the fact that a change in ______________________________________
dramatically changes the ______________________________________________________________________________.
 These molecules can be drawn as __________________________________________ diagrams (ex. Above) or as
____________________________________ (right) with
the shared pair of electrons represented by
a _____, 2 pairs by a ______, or 3 sets of shared electrons
by _____. The most that will ever be shared is 3.
Ionic Compounds
Def.
 Involves a ___________ + a ________________________ joining together, with the metal ______________________
its _______________ _____________________ and the __________________________ taking them.
Def. Ion –__________________________________________  2 forms 1.) _________ = Positive 2.) _________ = Negative
Sodium Atom, Na
Sodium Ion, Na+
Chloride Ion, Cl-
Chloride Atom, Cl+
+ Charge (Protons)
- Charge (Electron)
Ionic charge
Why does the metal always give away its valence electrons and the non-metal always take them?
In order to illustrate how this exchange of electrons occurs, take a look at the bohr-rutherford diagrams
on page 260. Copy the diagram into the note. Remember the third shell can fit a maximum of 8 electrons.
Bohr-Rutherford Diagrams Above… Lewis structure below
HW  Q 1-8 on page 261
Periodic Table Worksheet
REVIEW
1. ________________ are located on the left side of the periodic table, with _____________________ located on
the right side. They are separated by a “_________________” that touches the ________________________.
2. Horizontal rows are known as __________________ and tell us the number of _____________________ or
______________ an element has.
3. _______________ or groups are vertical __________________ on the periodic table and tell us how many
valence __________________ an element has.
4. ATOMS have the same number of Protons as they do __________________, so their charge is neutral,
this information is detailed in the __________________.
5. According to your periodic table, the most common isotope of carbon has ___ neutrons.
6. IONS have a positive or negative __________________ because they have given up their __________________
__________________ or taken valence electrons from an atom of another element in order to fill/empty
their outer energy shell.
7. Positive ions are known as __________________ (think the t looks like a +) while negative ions are
called __________________ (has two n’s for negative).
PRACTICE
Magnesium Atom
a) Period # =
b) # of energy levels =
c) Group # =
d) # of Valence Electrons =
e) Member of the _____________ Family
f) Draw a Bohr-Rutherford Diagram
Lithium Atom
a) Period # =
b) # of energy levels =
c) Group # =
d) # of Valence Electrons =
e) Member of the _____________ Family
f) Draw a Lewis-Dot Diagram
Argon Atom
a) Period # =
b) # of energy levels =
c) Group # =
d) # of Valence Electrons =
e) Member of the _____________ Family
f) Number of Neutrons =
g) Draw a Bohr-Rutherford Diagram
Fluorine Atom
a) Period # =
b) # of energy levels =
c) Group # =
d) # of Valence Electrons =
e) Member of the _____________ Family
f) Number of Neutrons =
g) Draw a Lewis Dot diagram
ION REVIEW
a) An ION of fluorine is going to mimic (look like) an atom of __________________.
b) An ION of beryllium is going to mimic an atom of __________________.
c) An ION of __________________ has to gain 3 electrons to mimic argon. Its ionic charge will be ____.
d) An ION of __________________ has to give up 3 electrons in order to mimic an atom of sodium. Its
charge will be ___.
e) An Ion of Oxygen has to _________ ___ electrons to mimic _____________. Its ionic charge will be _____.
Calcium Atom
a) Atomic # =
b) # of Protons =
c) # of Electrons =
d) Ionic Charge =
e) Atomic Mass =
f) Number of Neutrons =
g) Draw a Bohr-Rutherford Diagram
Oxygen ION
h) Atomic # =
i) # of Protons =
j) # of Electrons =
k) Ionic Charge =
l) Atomic Mass =
m) Number of Neutrons =
n) Draw a Bohr-Rutherford Diagram
Theories of the Atom Review – Questions 1, 2, 3, 5, & 7 on page 233
Naming Ionic Compounds
Def.
 Involves a ___________ + a ________________________ joining together, with the metal ______________________
its _______________ _____________________ and the __________________________ taking them.
Def. Ion –__________________________________________  2 forms 1.) _________ = Positive 2.) _________ = Negative
NAMING
Ionic compounds are easily identified by the presence of a ___________ (first term). Any time
you see a metal as the first term, you should automatically be thinking IONIC! Additionally,
there is never a __________________ reference in the compound name (so no _________________).
The nice thing about naming Ionic molecules is we __________ need to worry about numbers.
Example  K2O
1) STEP 1 – determine how many elements (each capital letter represents a new element) are
present in the compound (_______) and locate them on the periodic table. Is there a metal..? If so it
must be a Ionic Compound.
2) STEP 2 - Identify which of the elements is the metal (it must go first). __________________
3) STEP 3 – Identify the non-metal element (it goes 2nd). ____________________
4) STEP 4 –Write the name of the metal, than the name of the non-metal (but change the ending of
the non-metal to “IDE”). ____________________ _____________
Try these…
Li2O
MgF2
NaCl
K3N
CaS
Be3P2
LiBr
Na2O
WRITING chemical FORMULAS for Ionic Compounds
Since the electrons are ______________________________ and electrical ___________________ are
present we need to reference the Periodic Table in order to determine how many of each
element there is going to be. _____________________________________________________________!
Step 1: Locate the elements in the Periodic Table- ensure one is a metal, and one is a non-metal (notice
“ide” ending as a clue). Write down their chemical symbol.
Step 2: Determine the “charge” each element carries when its forms an ionic bond (remember: STABLE)
Step 3: Backcross the “charges“ to SUBSCRIPTS for each element
Ex. Rubidium fluoride
Compound
CHARGE
FORMULA
Sodium chloride
 “charges it up”
Magnesium nitride
_____ _____  “backcross it down”
Calcium fluoride
Francium phosphide
Naming Molecular Compounds
Def.
 This means there is no taking or giving of electrons, therefore there is no _________
These compounds are…
a)
b) Have a low melting point
c)
d) Do not conduct electricity
e)
NAMING
Molecular compounds are _________________________ by the use of PREFIXES. Anytime you notice a numerical
prefix in a compound, you should automatically be thinking MOLECULAR and NON-METAL atoms.
Mono
Di
Tetra
Penta
Hexa
Hepta
Example  H2O2
1) STEP 1 – determine how many elements (each capital letter represents a new element) are
present in the compound (_______) and locate them on the periodic table.
2) STEP 2 – how many HYDROGEN (H) atoms are there? ________
3) STEP 3 – how many OXYGEN (O) atoms are there? ________
4) STEP 4 – write the corresponding prefixes first (on the smaller lines), then the name of the
element on the larger lines, remember to always change the ending of the second element to “IDE”
just like we did yesterday with IONIC compounds.
TRY THESE
________ _______________ ____ ______________
N2O
ICl2
BrI
XeF4
ClO2
S2Cl2
UF6
P2O5
WRITING chemical FORMULAS for Molecular Compounds
The ______________ in-front of the element tells you how many ______________ (units) of that element there are
in the formula. These numbers are to be written as SUBSCRIPTS after the elements symbol.
1) STEP 1 – write the symbol for each element on the top lines
2) STEP 2 – write the corresponding prefix number below and after each elements symbol
TRY THESE
Nitrogen dioxide
Ex. Dinitrogen tetroxide 
____
____
Dinitrogen monoxide
Sulphur trioxide
Diphosphorus pentasulphide