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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 Bonding

The 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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10

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

• Animation

• 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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• Ionic and Covalent Bonding Review Clip

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Properties of Ionic and Covalent Compounds/Molecules

1.CO2

2.H2O

3.NaCl4.MgCl2

5.NO2

6.Li2S

7.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 elements

These 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 electron

All 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

Fe2O

3

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.BaF2

5.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.Cu2O

5.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.H2O

6.NO2

7.CuSO4

8.CaSO4

9.NH4OH

10.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. H2O

5. CH4

6. CO2

7. BF3

8. F2O

9. 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

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