Download 2 - Ms. Schuld`s Scientists

Bonding
Unit 4, Day 3
12/09/14
Door
SMARTBOARD
Door
SMARTBOARD
Go to your ASSIGNED seat!!
Go to your ASSIGNED seat!!
Door
SMARTBOARD
Bonding
WEEK 15  12/09/14
CATALYST
Can two anions form an
ionic bond? Why or why
not?
Have out on
your desk:
• Binder
• Sheets from front
*Answer in complete sentences!
Homework:
Ionic bonding practice
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End
4
…
Bonding
Ionic bonding
12/09
Bonding
12/09/14
Learning Intention:
We will demonstrate
knowledge of ionic
bonding.
Success Criteria:
We will be successful when we
model the bonding
relationships that would occur
between various elements.
Today’s Agenda
 Lake
sturgeon
 Ionic bonding NOTES
 Homework time
Lots of work today!  Let’s ROCK!
Lake sturgeon bowl
participants!!!
 Meet
TODAY after school in this room (326)
for a brief informational meeting!
 Anyone
who wants to
participate but can not stop by?


GET
PUMPED
Must know this!
Lewis Structures
Atoms
and ions may be
represented by Lewis structures
Models that show ONLY the
valence electrons
Electrons are represented by dots
surrounding the element symbol.
Electrons
are typically paired.
Example: Oxygen
Practice: Draw the following
Lewis Structures
Aluminum
Lithium
Bonding
Chemical
Bond:
two or more atoms sharing or
transferring electrons in order to
link the elements together to
create a compound or molecule
Octet Rule:
atoms
want to have an electron
configuration the same as a
noble gas (8 valence electrons).
This means atoms will give up or
share electrons in order to
achieve this configuration.
Types of bonds…
1) Ionic:
electron is transferred
from one atom to another
2) Covalent:
electron(s) are
shared between two atoms
Ionic Bonding & Nomenclature
When
ionic bonds are formed,
valence electron(s) are
transferred from the atom with
lesser electronegativity to the one
with greater electronegativity
Electronegativity is the tendency of
an atom to attract an electron.
Must
follow the octet rule during
bonding
Remember!
Ions are formed by
gaining or losing electrons!
Unless
the overall compound has
a charge, make sure all of the
charges of cations balance out
the charges of your anions.
Example: Formation of sodium
sulfide (Na + S  _______)
1.
Identify the charge of
each component based
on its location in the
periodic table.
2.
Example: Formation of sodium
sulfide (Na + S  _______)
Balance the charges so that the
total charge of the compound is 0.
Criss-cross method: the charge of the cation will be
equal to the number of anions you have and the
charge of the anion will be equal to the number of
cations you have. If they have the same number,
you do not need to do this!
Example: Formation of sodium
sulfide (Na + S  _______)
3.
Use the balanced
numbers as
subscripts to identify
how many of each
atom are in the
compound (you do
not need to identify
with subscripts if
there is only one of
the atom in the
compound)
Practice: Write the formula for
each compound
 Potassium
+ Chlorine  Potassium Chloride
 Magnesium
 Aluminum
+ Chlorine  Magnesium Chloride
+ Sulfur  Aluminum Sulfide
Ionic Lewis Structures
Use
Lewis structures to show the
movement of electrons and the
bonds they form.
Note:
The central atom is the one
with the lowest electronegativity
Ex: Draw the Lewis Structure for NaCl
1.
Draw the Lewis structure for each of the atoms
2.
Identify how the electrons move in order to form
the bond
Draw each atom with its new electrons & charge
3.
Practice: Draw the Lewis structure
for the following ionic compounds
MgCl2
You try!
NaF
Al2S3
FYI!
Polyatomic Ions
Poly
= Many
Atomic = Atoms
Multiple atoms that make up a
specified ion that forms ionic
compounds
See list of polyatomic ions!
Example: Write the formula for
calcium hydroxide
1.
Identify the ions present and
their charges
Ex: Write the formula for
calcium hydroxide
2.
Balance the charges
3.
Use the balanced charges as the
subscripts for the compund
**Be sure if you have multiples of the
polyatomic ion that the whole ion is in
parentheses with the subscript outside of the
parentheses!
Practice  Write the formula!
 Sodium
Carbonate
 Potassium
Permanganate
Homework time!
 Quietly
complete as much as you can.
 Work with your neighbor.
 Before asking Ms. Schuld
1) Check with a classmate
2) Consult your notes
 You
ROCK :)
Tomorrow = Nomenclature of
ionic compounds!
 Nomenclature:
scientific term for naming
Bonding!

https://www.youtube.com/watch?v=_M9khs87xQ8

https://www.youtube.com/watch?v=NgD9yHSJ29I

http://www.pbslearningmedia.org/resource/lsps07.
sci.phys.matter.chembonds/chemical-bonds/
Lab review!
12/09/14
On a Tuesdayyyy
Success Criteria
Question
In order for an
ionic compound to
be formed, what
must occur?
Nomenclature for Ionic
Compounds
Nomenclature:
scientific
term for naming
Type I Ionic Compounds
 Compounds
that contain a metal
cation and a non-metal anion
1.
2.
3.
Cation is named first. Cation takes
its name directly from its element.
Anion is named second.
Replace the suffix of the anion with
“-ide”
Write the name all together
Example: NaCl
1.
Cation is named first. Cation takes
its name directly from its element.
Anion is named second.
2.
Replace the suffix of the anion with
“-ide”
3.
Write the name all together
Compounds
that contain a
Typecation
II Compounds
metal
with no fixed
charge, and a non-metal anion
The cation is named the same
way as type I, but now a
Roman numeral is used in
order to designate the charge
of the metal
Ex.
Iron (III) for Fe3+ vs. Iron (II) for
Fe2+
The
anion is written the
same way as type I – take
the root of the element and
replace the suffix with “-ide”
Example: Name the
compound HgO
1.
Cation is named first. Cation takes
its name directly from its element.
Anion is named second.
2.
Determine the charge of the
anion to help you identify the
charge of the cation
Example: Name the
compound HgO
3.
Write the name of the cation with
the roman numeral
4.
Write the name of the anion
following the cation and roman
numeral, making sure the suffix is
replaced with “-ide”
Name the following
compounds.
1.
2.
3.
CuI2
AuCl3
Ag2S
Properties of Ionic Compounds
Form
crystals
Very hard – each ion is
bonded to several
oppositely charged ions
High melting points
Brittle
Conduct electricity when
Covalent Bonding and
Nomenclature
Covalent
Bonding: electrons are
shared between two atoms,
creating a chemical connection
between the two atoms
 Rather
than having charges that
attract two atoms together, it is
the mutual need of additional
electrons to fulfill their octet
Atom
H
Electrons
Electrons Needed Electron
in Outer
To Fill
Pairs
Shell
Outer
Shared
Shell
1
1
1
Cl
7
1
1
O
6
2
2
N
5
3
3
C
4
4
4
S
6
2
2
 Lewis
structures are used to
represent these compounds
 Lines connecting the atoms show
that the pairs of electrons are being
shared between the two atoms
 Must always follow octet rule
 When
counting, lines count as 2
electrons for each atom
Three main types of covalent
bonds
Single: when two
atoms share a single
pair of electrons
Example:
F2
Each fluorine is
missing one
electron, so they
share a pair of
electrons, making
both outer shells full
Three main types of covalent
Double:
bondswhen two
atoms share two pairs
of electrons
Example:
O2
Each atom has 6
valence electrons, so
if each atom shares
two electrons with
the other, then both
atoms’ valence shells
will be full
Three main types of covalent
Triple:
when two
bonds
atoms share three
pairs of electrons
Example:
N2
Each atom has 5
valence electrons, if
each atom shares
three electrons with
the other its need for
additional electrons
will be satisfied.
Bonding between different
elements
1.
Determine how many
electrons each element
needs to fill the octet rule
Different
elements have
different needs! This is solved
the same way as when atoms
had different charges.
The crossover rule still works!
Example: make a compound
where oxygen and hydrogen are
covalently bonded
Oxygen
needs 2e
Hydrogen needs 1e
We
need 2
hydrogen for every 1
oxygen OR you
could use 2 oxygen
and 2 hydrogen
Helpful Hints
1.
2.
3.
The atoms that need the most
electrons should be placed in the
center
The atoms that need the fewest
electrons should be placed
around the ones that need most
Attempt to satisfy octets with
single bonds before making
double or triple bonds
Draw Lewis Structures for the
following compounds
CCl4
CO2
SO2
Nomenclature: Type III NonMetal Compounds
This
is only for compounds
containing only non-metals!
Covalently bonded
1.
2.
3.
The less electronegative atom is
named first, then the more
electronegative
The suffix of the more
electronegative atom is replaced
with “-ide”
The number of atoms of each
element is given a Greek prefix
** If there is only one of the less
electronegative atom, then
you
do not have to add the
monoprefix
1.
Example: Name the
The
less electronegative
compound
CCl4 atom is named
first, then the more electronegative
2.
The suffix of the more electronegative
atom is replaced with “-ide”
3.
The number of atoms of each element
is given a Greek prefix
Polar Covalent Compounds
Polar
covalent bonds occur
when the difference in
electronegativity is 0.5-1.5
Electrons are shared
unequally between atoms
The
atom with the greater
electronegativity will pull
the electrons closer to it,
therefore causing the
compound to be polar
Example:
Water
Properties of Covalent
Compounds
 May
be solid, liquid or gas at room
temperature
 Tend
 The
to be soft solids
more atoms that each
compound has, the higher the
melting point will be
 Brittle
 Poor conductors of heat and
electricity
VSEPR: Valence Shell Electron
Theory
 Covalent
bonds and lone pair electrons
like to stay as far apart from each other
under all conditions
 The shape of the molecule (molecular
geometry) is determined by how many
bonding and non-bonding electron
pairs there are. This identifies the overall
shape, and the angles between the
bonds
Steps to Predict Molecular
Geometry
1.
2.
Draw the Lewis Structure
Count the number of lone pairs
there are in the compound
 Double
3.
4.
and triple bonds count as 1!
Count the number of bonds
there are in the compound
Use the reference table to
predict geometry
Example: Predict the shape of
H2 O
1.
Draw the Lewis structure of the
compound
2.
Count the number of lone pairs there
are in the compound
3.
Count the number of bonds there are in
the compound
4.
Use the table for reference to determine
the geometry.
Review: IONS
When
atoms will
lose electrons they
become cations
When
atoms will
gain electrons
they become
anions
2
Na Mg
2
S
Cl
Steps for ions:
1.
2.
3.
4.
Find number of valence electrons
Determine how many electrons the
element will gain or lose to get to 0 or 8
Write that number as a superscript on
the right side of the ion
Add a + after if positive or a – after if
negative
Potassium
(K)
Carbon
Iodine
(I)
Review: Ionic Charge
+
Na
The Ionic Charge for an
atom is the charge on the
atom (how many
electrons it gains or loses)
What is the Ionic Charge for the
following ions?
2O
2+
Ca
3B
+
K
Ionic Bonds
Ionic
bonds are bonds between
cations & anions
If
atoms transfer e- to create
opposite ions, the opposite charges
attract and form an ionic bond
Example!
Na+
+ Cl-  NaCl
(sodium chloride = salt!)
Key Question…
Why
do atoms transfer
electrons in the first place?
(Why would they want to
gain or lose electrons?)
Octet Rule!
 Atoms
want ________ electrons in their
outermost shell!
 They
will be happy when they have 8
valence electrons!
(Or if they are closer to the atomic
structure of He they will be satisfied with 2)
So which
elements bond with which?!
 Cations
and Anions come together to
make compounds!
Example from yesterday’s lab!
Chlorine and calcium ions came together to
form calcium chloride.
2+
Ca
+
2Cl
 CaCl2
But where does that 2 come from?!!?!
Criss-Cross Method
1.
2.
3.
4.
5.
List the Metal first, non-metal second.
Write the ionic charge above the symbol.
Criss-cross the numbers. (a 1 does not
need to be written, drop the + or -)
Final answer lists only the subscripts
(numbers below the line)
If they are the same number, they cancel
out
IE: Combine hydrogen and
chlorine
IE: Now magnesium and fluorine
Charges must be balanced!
Na+ + Cl-  NaCl
Mg2+ + F- + F-  MgF2
Predict the compound!
Na and O
Ca and S
Be and F
Missing safety contracts!!!
 Alfredo
 LJ
 Famo
 Thandi
 Christian
 Constance
 Demond
 Alejandro
 Emileye
 Tara
 Josh
 Jazmin
 Daisy
 Juana.
 Jubi
Hehehehe
 http://www.huffingtonpost.c
om/2014/02/19/alanandersen-bacontfm-iceslip_n_4816325.html
The families of the periodic table
Alkalai Metals: Group 1
Have a low density, low melting point, are good
conductors, and are extremely reactive with
oxygen, water and other halogens
The families of the periodic table
Alkalai Earth Metals: Group 2
Have very similar properties to group 1, however
slightly less reactive
The families of the periodic table
Halogens: Group 17
All metals, and diatomic
Diatomic: always found bonded to another
atom
The families of the periodic table
Noble Gases: Group 18
Inert (unreactive) because their electron shells are
full, making it almost impossible for these elements
to react with others
Why to we care about any of this??
 To

understand this!
http://www.nuffieldfoundation.org/practicalchemistry/ammonium-dichromate-volcano
 Ammonium
dichromate