Transcript Document 7542385

As you come in,
Materials:
– Grab a packet, paper and pencil
Plan:
– Return lab practical and cumulative quiz
– Learn about ionic bonding
Homework:
– 7-1 & 7-3 Practice Problems
Assessment:
– Polyatomic Ion Quiz - FRIDAY
Unit Four:
Bonding/Nomenclature
GPS SC1: Students will analyze the
nature of matter and its classifications.
Classification of Matter
Matter
Mixtures
Pure Substances
Elements
Compounds
Ionic
Metal
Cation
Nonmetal
Anion
Covalent
Nonmetal
Nonmetal
Ionic Bond
Ionic bond: the electrostatic force that
holds oppositely charged particles
together in an ionic compound
Compounds that contain ionic bonds
are called ionic compounds.
Ionic compounds form between metals
and nonmetals.
Criss-Cross Method for
Writing Chemical Formulas
Let’s predict the formula for the stable
ionic compound that contains:
–calcium ion and phosphide
–aluminum ion and sulfate
–tin(II) and carbonate
Valence Electrons &
Stability Review
Let’s talk about the formulas that you
just completed.
– Questions to ALWAYS ask yourself after
criss-crossing:
•
•
•
•
How many of the cations are in the compound?
How many of the anions are in the compound?
Is the ratio (AKA formula unit) reduced?
Are transition metals involved?
What is the chemical formula for
sodium hypochlorite?
A.) NaCl
B.) NaClO
C.) NaClO2
D.) NaClO3
E.) Na(ClO)
What is the chemical formula for
sodium acetate?
A.) Na2C2H3O
B.) (Na)(C2H3O2)
C.) NaC2H3O2
D.) Na3C2H3O2
E.) NaCHO
What is the chemical formula for
tin(IV) oxide?
A.) Sn4O
B.) Sn4O2
C.) Sn2O
D.) Sn2O4
E.) SnO2
What is the chemical formula for
magnesium oxide?
A.) MnO
B.) MnO2
C.) MgO
D.) Mg2O2
E.) MgO2
What is the chemical formula for
nickel(II) nitrate?
A.) Ni2NO3
B.) Ni1NO32
C.) NiNO3
D.) (Ni)(NO3)2
E.) Ni(NO3)2
What is the chemical formula for
magnesium phosphate?
A.) MgP
B.) Mg3P2
C.) MgPO4
D.) Mg2(PO4)3
E.) Mg3(PO4)2
Writing Ionic Compound
Names from the Formulas
When given a formula, analyze the ions.
Identify the two ORIGINAL ions that
were criss-crossed to give the formula.
“Backwards criss-crossing” or “un-crisscrossing” will NOT work in many cases!
You’ll have to work harder than that!
Practice Writing Names from
Chemical Formulas
Write the name of the compound that is
represented by the following formulas:
– MgCl2
– Na2SO4
– PbS
– Cu2SO3
A few more…
FeSO4
NiNO3
CrClO3
FeO
Sn2S
Name the following compound:
PbS
A.) lead sulfur
B.) lead sulfide
C.) lead (I) sulfide
D.) lead (II) sulfide
E.) lead sulfate
Name the following compound:
Cu2CO3
A.) copper carbonate
B.) copper (I) carbonate
C.) copper (II) carbonate
D.) copper (II) carbonate (III)
E.) copper carbontrioxide
Name the following compound:
CrPO4
A.) chromium phosphate
B.) chromium (I) phosphate
C.) chromium (II) phosphate
D.) chromium (III) phoshate
E.) chromium (IV) phosphate
Name the following compound:
NiSO4
A.) nickel sulfate
B.) nickel (I) sulfate
C.) nickel (II) sulfate
D.) nickel (III) sulfate
E.) nickel (IV) sulfate
Name the following compound:
Al2(SO4)3
A.) aluminum (III) sulfate
B.) aluminum sulfate
C.) aluminum (II) sulfate
D.) aluminum sulfite
E.) aluminum sulfide
Name the following compound:
KBr
A.) potassium bromide
B.) potassium (I) bromide
C.) potassium bromite
D.) potassium bromate
Name the following compound:
TiO2
A.) titanium oxide
B.) titanium (II) oxide
C.) titanium (IV) oxide
D.) titanium oxide (II)
Name the following compound:
Cr2SO3
A.) chromium (II) sulfate
B.) chromium (II) sulfite
C.) chromium (I) sulfite
D.) chromium sulfate
Write the chemical formula for zinc
selenide.
A.) ZnSe
B.) Zn2Se
C.) ZnSe2
D.) Zn2Se2
Write the chemical formula for
calcium nitrate.
A.) CaNO3
B.) CaNO2
C.) Ca(NO3)2
D.) Ca(NO2)2
E.) Ca3N2
Write the chemical formula for
cobalt(III) phosphide.
A.) Co3P
B.) Co3P3
C.) CoP
D.) CoPO3
E.) CoPO4
Formula unit ratios form
beautiful crystals…Ionic
compounds are crystals!
What You’ll See in Our Lab
Ionic compound solids that
form when two solutions are
mixed
“Snow Globe” effect
Separation
line
Paint
effect
Forming Ionic Compounds
Pre-Lab
Lab Purpose: To write formulas and names of 20 ionic
compounds; to see ionic compounds formed in the lab
Lab Safety: Wear goggles and apron. Avoid contact with
chemicals. Wash your hands after the lab. Do not mix up
the pipettes.
Lab Procedure: Note the layout of the test tubes. Place
4-5 drops of each chemical into the well-plate.
Lab Report: Turn in your data table on Monday.
Post-Lab Clean-up: Place your well-plate on the table
beneath the goggle cabinet. Wipe down your station.
Physical Properties of Ionic
Compounds
Applying heat to melt
sodium chloride
Applying a current to
test conductivity of
sodium chloride
Covalent Bonding
Why do atoms bond?
Remember that ionic
compounds form by
gaining and losing
electrons.
Covalent compounds
form when 2 atoms both
need electrons.
They can both achieve
the full octet by sharing
electrons.
What is a covalent bond?
Covalent bond: chemical bond that
results from sharing electrons
Covalent bonding generally occurs
when elements are relatively close
together on the periodic table.
The majority of covalent bonds form
between nonmetallic elements…often
groups 4A to 7A.
Naming Covalent Compounds
First, you must be sure that the
compound is COVALENT by analyzing
the formula.
Second, you use prefixes to
communicate the number of atoms of
each element that form the compound.
Example: H2O = dihydrogen monoxide
Prefixes
One = mono
Two = di
Three = tri
Four = tetra
Five = penta
Six = hexa
Seven = hepta
Eight = octa
Nine = nona
Ten = deca
NEVER USE MONO ON THE FIRST ELEMENT NAME.
ALWAYS USE THE -IDE ENDING ON THE SECOND
ELEMENT NAME.
Name the following compound:
P4 S5
A.) phosphorus sulfide
B.) tetraphosphorus pentasulfur
C.) tetraphosphorus pentasulfide
D.) tetraphosphide pentasulfide
E.) phosphorus (IV) sulfide
Name the following compound:
SeF6
A.) selenide fluoride
B.) selenium fluorine
C.) monoselenium hexafluoride
D.) selenium hexafluoride
E.) monoselenide hexafluoride
What is the chemical formula for
dinitrogen trioxide?
A.) NO
B.) N2O
C.) N2O2
D.) N2O3
E.) N3O2
As you come in,
Materials:
– Keep your phone. If you don’t have a smart phone, grab a
laptop and login.
Plan:
– Socrative app and Edmodo app
– Take Ionic Compound Quiz
– Get test results back
– Learn how to name acids (Take Socrative quiz)
– Learn how to draw Lewis structures
Assessments:
– Covalent & Acids Naming/Formula-writing Quiz - Tomorrow
Add
App
It’s FREE. Don’t worry!
Search app store for “Socrative Student”
Install app
If you don’t have a smart phone, don’t worry! You’ll
use a laptop in class later to do the same activities.
Add
App
It’s FREE, too.
Enable you to do practice quizzes or watch
preview videos on your phone.
Search app store for “Edmodo”
Install app
Again, you can do the same thing on a computer.
Don’t worry if you don’t have a smart phone!
DEVICE DOWN
Cell phone goes FACE DOWN on the
top corner of your desk.
It stays there until Mrs. Tarvin gives the
DEVICE UP signal.
Any deviation from the rule: phone to
the pocket at front of room & use a
laptop instead
DEVICE DOWN
Ionic Compound Quiz
– Need: Pencil/pen & periodic table
After the quiz, look over your Periodic Table test
results. We’ll go over the test on Friday since
your cumulative quiz is Monday.
If you have time, practice naming and writing
formulas for covalent compounds in your practice
packet.
Naming Acids
Acids are compounds that contain hydrogen
ion bonded to an anion, except water.
Two types of acids:
– Binary Acids - “Hydro____ic acid”
• Examples: HCl, HBr, H2S,
– Oxyacids - “_____ic acid” OR “____ous acid”
• Examples: H2SO4, HNO3, HC2H3O2
• ic = ate anion; ous = ite anion
DEVICE UP
Open the Socrative App.
– On a laptop: m.socrative.com
Enter the room #: 189001
Wait for me to start the activity.
You’ll practice naming acids with 5
multiple choice questions.
Lewis Structures
Lewis Structures
Lewis structures: electron-dot diagrams show how
electrons are arranged in molecules (AKA covalent
compounds)
–
–
–
–
Diagrams show ONLY valence electrons.
Dashes represent “bonding pairs”
Dots represent “lone pairs” or “nonbonding pairs”.
Brackets and charge signify polyatomic ions.
How to Draw Lewis Structures
1. Find the total # of valence electrons in the molecule.
2. Divide this number by 2. This is the number of bonding
pairs.
3. Arrange the elements to show a central atom and terminal
atoms. The central atom is usually first in the formula, and
hydrogen is always terminal.
4. Place one bonding pairs between the central atom and
each terminal atom.
5. Place the remaining pairs around the terminal atoms…if
they need more electrons. Put the rest around the central
atom.
6. Check to be sure that multiple bonds don’t exist.
Analyze the Structure
Lewis structures are drawn to predict the type of
covalent bond.
– Polar covalent: One atom has a significantly
higher electronegativity than the other atom. The
bonding electrons are shared UNEVENLY as a
result.
– Nonpolar covalent: The atoms have similar
electronegativity values, so the bonding electrons
are shared EVENLY.
Determine Polarity
Draw the Lewis Structure.
Look at the central atom.
– Does it have a surrounding
atom that differs from the
others?
– Does it have a lone pair of
electrons?
If yes to either, then the molecule
is polar.
Polar molecules WILL dissolve in water
which is significant in MANY situations.
Example #1: CH4
– Name the compound
• Carbon tetrahydride
– Total valence electrons
• 4+1+1+1+1=8
– Bonding pairs
• 8/2 = 4 bonding pairs
– Arrange the atoms
• Remember the H is always terminal.
– Place bonding pairs between the C & each H
– How many pairs remain?
• O remaining pairs
– Check for multiple bonds. (Does each atom have 8 valence
electrons?)
• Yes, so there are no multiple bonds needed.
Example #2: NH3
– Name the compound
• Nitrogen trihydride (AKA ammonia)
– Total valence electrons
• 5+1+1+1=8
– Bonding pairs
• 8/2 = 4 bonding pairs
– Arrange the atoms
• Remember the H is always terminal.
– Place bonding pairs between the N & each H When the central
atom has a lone pair
– How many pairs remain?
of electrons, the
• 4 pairs - 3 pairs used = 1 remaining pair
molecule
bends.
• Add the remaining pair to the terminal first...then
the central.
pairs of
– Check for multiple bonds. (Does each atom haveLone
8 valence
electrons take up
electrons?)
• Yes, so there are no multiple bonds needed. lots of space.
Example #3: Br2
– Name the compound
• Bromine
– Total valence electrons
• 7 + 7 = 14
– Bonding pairs
• 14/2 = 7 bonding pairs
– Arrange the atoms
• Two atoms are arranged next to each other.
– Place bonding pairs between the two Br atoms
– How many pairs remain?
• 7 pairs - 1 pairs used = 6 remaining pairs
• Add the remaining pair to the terminal first...then the central.
– Check for multiple bonds. (Does each atom have 8 valence
electrons?)
• Yes, so there are no multiple bonds needed.
Example #4: O2
– Name the compound
• Oxygen gas
– Total valence electrons
• 6 + 6 = 12
– Bonding pairs
• 12/2 = 6 bonding pairs
– Arrange the atoms
• Two atoms are arranged next to each other.
– Place bonding pairs between the two O atoms
– How many pairs remain?
• 6 pairs - 1 pair used = 5 remaining pairs
• Add the remaining pair to the terminal first...then the central.
– Check for multiple bonds. (Does each atom have 8 valence
electrons?)
• No, so there ARE multiple bonds needed.
Example #5: CO2
– Name the compound
• Carbon dioxide
– Total valence electrons
• 4 + 6 + 6 = 16
– Bonding pairs
• 16/2 = 8 bonding pairs
– Arrange the atoms
• C should be central with an O on each side. (symmetrical)
– Place bonding pairs between the C and each O atom
– How many pairs remain?
• 8 pairs - 2 pairs used = 6 remaining pairs
• Add the remaining pairs to the terminal first...then the central.
– Check for multiple bonds. (Does each atom have 8 valence
electrons?)
• No, so there ARE multiple bonds needed.
Example #6: CO
– Name the compound
• Carbon monoxide
– Total valence electrons
• 4 + 6 = 10
– Bonding pairs
• 10/2 = 5 bonding pairs
– Arrange the atoms
• Two atoms are arranged side by side.
– Place a bonding pair between the C and O atom
– How many pairs remain?
• 5 pairs - 1 pair used = 4 remaining pairs
• Add the remaining pairs to the terminal first...then the central.
– Check for multiple bonds. (Does each atom have 8 valence
electrons?)
• No, so there ARE multiple bonds needed.
As you come in,
• Use your cell phone or a laptop to access the
socrative app.
(m.socrative.com)
• Take the practice covalent and acid quiz. You may
use a periodic table as the resource.
• Do you have any questions? We’ll be taking the
REAL quiz in a few minutes.
• When you have finished the practice quiz, please put
your cell phone in the organizer pocket to prepare for
the real quiz.
After the quiz
Draw the following Lewis structures into
your notes. Predict the polarity of each.
– I2
– SO3
– CF4
– PCl3
– SiO2
Advanced Lewis Structures
Draw the Lewis structure for CH3Cl.
– Check the polarity of the overall molecule.
– Will this molecule mix with water?
Advanced Lewis Structures
Draw the Lewis structure for silicate.
– Check the polarity of the overall molecule.
– Will this molecule mix with water?
Advanced Lewis Structures
Draw the Lewis structure for CH3OH.
– Check the polarity of the overall molecule.
– Will this molecule mix with water?
VSEPR Theory
Three-dimensional representation of
molecules
Valence Shell Electron Pair Repulsion
Theory
Counting areas of electron density around
the central atom.
Electron density repels!
This bends the molecule into interesting
shapes.
Arranging Areas of High E- Density
# regions of high
electron density
molecular shape
bond angle
2
linear
180°
3
trigonal planar
120°
4
bent
<109.5°
4
tetrahedral
109.5°
4
trigonal pyramidal
<109.5°
Predict Shapes & Angles
Go back to the Lewis structures that you drew
earlier. Let’s check them and make some
predictions.
– I2
– SO3
– CF4
– PCl3
– SiO2
Examples to Build
at Your Station
1.
2.
3.
4.
5.
6.
7.
8.
H2O
O2
HCl
CO2
NH3
CH4
NO3CH3CH3
Draw a data table of 7 columns.
Column 1 - Formula
Column 2 - Name
Column 3 - Lewis structure
Column 4 - Shape
Column 5 - Bond angle
Column 6 - Polarity
Column 7 - Type of Intermolecular Force
Intermolecular Forces
Abbreviated - IMF
Forces of attraction and repulsion that
exist BETWEEN molecules
Three types of IMF:
– London dispersion forces
– Dipole-dipole forces
– Hydrogen bonding
Strength
Increasing
Three Types of IMF
London dispersion forces (LDF) - weakest of all
intermolecular forces; temporarily exists between any
two molecules; ONLY one that occurs in nonpolar
molecules (low melting & boiling points)
Dipole-dipole forces - forces of attraction/repulsion
that exist between POLAR molecules as a result of
the partial charges (high melting & boiling points)
Hydrogen bonding - STRONGEST intermolecular
force that occurs when H is bonded to FON creating
REALLY strong partial charges (highest melting &
boiling points)