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Adventures of
Oxygen Clip
Chapters 7,8
1. Compare and contrast types of chemical bonds (i.e. ionic, covalent). 2. Predict formulas for stable ionic compounds (binary and tertiary) based on balance of charges.
GOALS
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3. Use IUPAC nomenclature for both chemical names and formulas:
•Ionic compounds (Binary and tertiary) •Covalent compounds (Binary and tertiary)
4. Apply concepts of the mole and Avogadro’s number to conceptualize and calculate empirical/molecular formulas, mass, moles and molecules relationships.
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5. Identify substances based on chemical and physical properties
Why do Atoms Form Compounds?
•Stability.•What makes an atom stable?
•Full outer energy level.•Eight.
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•A Chemical Bond holds atoms together in a compound.
•Two basic types:1-Ionic2-Covalent
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Ionic Bonding
Transfer of electrons from one atom
to another atom.
Occurs between metals & nonmetals.
Electrically neutral
Called compounds.
Compound composed of cations and anions.
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OPPOSITS ATTRACT!
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Ionic Bondin
g
CLIP
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Properties of Ionic Compounds• Crystalline solids at room
temperature.• Arranged in repeating three-
dimensional patterns• Have high melting points• Can conduct electricity
when melted or dissolved in water
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Covalent BondingThe sharing of
electrons between
atoms.Each atom
attempts to fill
their valence shell.
Occurs between nonmetals
and nonmetals.
Called Molecules: Neutral group of atoms joined by a
covalent bond
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Hydrogen and Fluorine
Hydrogen and Chlorine
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Single, Double, Triple
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Single Covalent Bonds (2e-)
Structural Formula: dashes
Unshared pair
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Double and Triple Covalent Bonds• Double bond- 2 pairs (for a total of 4)
• Triple bond- 3 pair (for a total 6)
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Clip
Clip
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• The element that has a greater electronegativity attract the electrons more
• So, the electronegativity difference between two atoms tells you what kinds of bond is likely to form
Unequal Sharing of Electrons
Polar molecules
happen when one
atom has a greater positive charge
Called Polar Molecules
Unequal Sharing of Electrons
δ+Called Polar Moleculesδ_
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Practice Quiz
• The shape may affect the polarity of an entire molecule
• Ex CO2 (2 polar bonds cancel each other)
• The presence of a polar bond in a molecule often makes the entire molecules polar. (Water molecule)
• A molecule that has 2 poles is called a dipolar molecules, or dipole.
Properties of Covalent Molecules• Many are
gases or liquids at room temperature
• Composed of two nonmetals.
• Have low melting and boiling points
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Properties of Ionic and Covalent Compounds/Molecules
1.CO2
2.H2O3.NaCl4.MgCl2
5.NO2
6.Li2S7.NaF
9.BeO 10.HCl11.NaF12.KCl13.H2O2
14.N2
15.Cl2
Covalent or Ionic?20
Metallic Bonds• Valence electrons (1-
3) can be thought of as a sea of electrons. They are “mobile” and can easily drift freely from one part of the metal to another.
• Metallic bonds consist of the attraction of the free-floating valence electrons for positively charges metal ions.
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Other Atomic Attractions• Intermolecular attractions are weaker
than either ionic or covalent bonds.
• Van der Waals Forces– Weak attraction consisting of dipole
interactions and dispersion forces– Dipole interactions: when polar
molecules are attracted to another.– Dispersion Forces: weakest of all
interactions. Caused by motion of electrons. Occurs between nonpolar molecules. Temporary polarity.
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Hydrogen bonding• Found in many
biological molecules• Important in the
properties of water.• Attraction between
hydrogen (when bonded to a very electronegative element) and another molecule.
• About 5% the strength of an average covalent bond. 23
Goals revisited
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Ionic Bonding- Formula Units
A formula unit is the lowest whole-number ratio of the ions in an ionic compound.
A chemical Formula shows the kinds and numbers of atoms in the smallest representative unit of a substance.
How do you figure out the “Chemical Formula?”
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• Writing chemical formulas is a shorthand way of indicating what a substance is made of.
• These formulas also let you know how many atoms of each type are found in a molecule.
The chemical formula for water is H2O. Carbon Dioxide is CO2. Why does oxygen combine in different ratios, in different compounds? The chemical formula for table salt is NaCl. Calcium Chloride is CaCl2. Why does chlorine combine in different ratios, in different compounds?
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The simplest compounds are ones
with only two elementsThese are called binary
KI, CO, H2O, NaCl
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+1
+2
-1
-2-
3+3
+4 -4
0Oxidation numbers
Tell you how many electrons an atom must gain, lose or share to become
stable.28
We can predict the ratio of atoms in ionic compounds based on
their oxidation numbers
Oxidation numbers
K Cl+1
-1
KClTells you how many electrons an atom must gain, lose or share to become
stable.
1 valence electron
7 valence electronAll
compounds are neutralThat means the
overall charge is ZERO!
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Subscripts show the number of atoms of that kind in the compound
Na
Br
+1
-1
NaBr
Ca Br
+2
-1
CaBr2
To make it ZERO, you
need 1 Ca & 2 Br.
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Some elements have more than one oxidation number
Fe O+3
-2
Fe2O3
Fe O+2
-2
FeOWe call these elements- Multivalent
Elements
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Now You Try writing Binary Ionic formulas
1.K + Br2.Mg + Cl3.Ca + I4.K + O5.K + I6.Sr + Br7.Na + O
8.Ga + Br9.Fe+2 + O10.Fe+3 + O11.Cu+2 + F12.Cr+3 + O13.Mg + O14.Al + P
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Cations: ammonium, NH4+
Anions: nitrate, NO3-
sulfate, SO42-
hydroxide, OH-
phosphate, PO43-
carbonate, CO32-
chlorate, ClO3-
permanganate, MnO4-
chromate, CrO42-
Polyatomic Ions:
-a tightly bound group of covalently bonded atoms that has a positive or negative charge and behaves AS A UNIT.
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Polyatomic Ions-Compounds containing polyatomic ions include both ionic and covalent bonding
Writing Formulas Examples:Sodium and NitrateMagnesium and ChlorateAmmonium and Sulfate
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Try these1.Na + SO4
2.Mg + PO4
3.Ca + CO3
4.Na + OH5.Mg + OH6.NH4 + OH
7.K + PO4
8.NH4 + NO3
9.H + SO4
10.Ca + SO4
11.K + NO3
12. Na + PO4
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Naming Binary Compounds and Molecules
• Steps:– If it is Binary-1. Decide if it is an ionic or covalent
bond.– Metal- nonmetal…..
» Ionic– Nonmetal- nonmetal….
» Covalent
Example:• NaCl
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If ionic …….2. Only 2 elements3. Check to see if
any elements are multivalent.
4. If all single valent, write the name of the positive ion first.
5. Write the root of the negative ion and add –ide.
Examples:1.NaCl2.K2O3.AlCl34.BaF25.KI6.Li2O
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If ionic …….5. Check to see if any
elements are multivalent.
6. If multivalent ions, determine the oxidation number of the element.
7. Use Roman numerals in parentheses after the name of the element.
8. Write the root of the negative ion and add –ide.
Examples:
1.FeO2.Fe2O3
3.CuO4.Cu2O5.PbCl4
6.PbI2
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If Covalent... (Molecular Formula)2. Use Greek prefix to indicate how many atoms of each element are in the molecule
3. Add -ide to the more electronegative element
Greek Prefixes1- mono-2- di-3- tri-4- tetra-5- penta-6- hexa-7- hepta-8- octa-
Example:• NO• Nitrogen Monoxide
• PCl3• Phosphorous trichloride
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If it contains a polyatomic ion...2. Write the name
of the positive ion.
3. Write the name of the polyatomic ion.
Examples:1. NaCO3
2. KNO3
3. NaC2H3O2
Example:• KOH• Potassium Hydroxide
• CaCO3• Calcium Carbonate
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Name the following:1.KBr2.HCl3.MgO4.CaCl25.H2O6.NO2
7.CuSO4
8.CaSO4
9.NH4OH10.CaCO3
11.Cu(ClO3) 2
12.Cr2O 3
3342
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Drawing Lewis Structures• Step #1: Add up the number of
valence electrons that should be included in the Lewis Structure. (TVE)
• Step #2: Calculate # of bonds.– Determine TOE: Theoretical Octet
Electrons– TOE- TVE from step1– Divide by 2 ( 2 electrons for each
bond) • Step #3: Draw the “skeleton
structure” with the central atoms and the other atoms, each connected with a single bond.
• Step #4: Any “leftover” electrons so that all elements meet octet rule (or full outer energy).
NH31. 5 + 3(1) = 8 (nitrogen
has five; each hydrogen has one)
2. . N-8, H (2 each x 3=) 6…– so TOE=14– 14-8= 6– 6/2= 3 bonds
3. .
4. .
Double, triple bonds.• Same as last except…• Step #4: If there are no
electrons left, move electrons from a different atom to form another bond…double
• Side note: When more than one Lewis structure can be drawn, the molecule or ion is said to have resonance.
CO32-
Drawing Lewis Structures
Try these…1. CCl42. NF3
3. SH2
4. H2O5. CH4
6. CO2
7. BF3
8. F2O9. SO2
10. SO3
11.NF3
12.N2
13.NH4+ (notice the + charge)
14.NO3- (notice the - charge)
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Molecules Have Shapes.
• VSEPR theory proposes that the geometric arrangement of terminal atoms, or groups of atoms about a central atom in a covalent compound, or charged ion, is determined solely by the repulsions between electron pairs present in the valence shell of the central atom.
• The number of electron pairs around the central atom can be determined by writing the Lewis structure for the molecule. The geometry of the molecule depends on the number of bonding groups (pairs of electrons) and the number of nonbonding electrons on the central atom.
Molecular Shapes• VSEPR Theory:
(Valence electron-pair repulsion theory)
• The repulsion between electron pairs causes molecular shapes to adjust so that the valence-electron pairs stay as far apart as possible
• Lone pairs have more repulsive force than do shared electron pairs, and thus they force the shared pairs to squeeze more closely together.
Linear
Tetrahedral
PyrimidalTrigonal Planar
Practice: Go back to Lewis Structure Practice, and
predict shapes.
Bent
Shapes and Polarity• Molecules can be polar, and
when they are polar, they are called dipoles.
• Dipoles are molecules that have a slightly positive charge on one end and a slightly negative charge on the other
• Shape can help determine polarity
• Molecules that are symmetrical tend to be nonpolar. Molecules that are asymmetrical tend to be polar
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Practice: Go back to Lewis Structure
Practice, and predict polarity.
H2O- Bent-Polar SO3 -trigonal planar-nonpolar.
BF3- trigonal planar-nonpolar.
SO2 -bent-polar
Molecules in Motion Website
Starter 7