chemicalbondingpresentation20120813.notebook …yooschem1314.pbworks.com/w/file/fetch/69567690...chemical bonding presentation 2012 08 13.notebook 2 August 17, 2012 The Octet Rule

chemicalbondingpresentation20120813.notebook

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August 17, 2012

Ionic Compoundsand

Ionic Bonding

Chemical Bonds

There are three basic types of bonds:

Ionic The electrostatic attraction between ions

Covalent The sharing of electrons between atoms

Metallic Each metal atom bonds to other metals atoms within a mobile "sea" of electrons (not covered until AP Chemistry)

Electronegativity is how strongly an atom attracts electrons.

Atoms with a high electronegativity will be able to attract electrons away from atoms with a much lower electronegativity.

This removal of electrons can occur when the difference in electronegativity between the two atoms is approximately 1.7 or higher.

Once a positive and negative ion are formed, they will be attracted to each other via the electrostatic force:

F = k q1 q2

r2

Ionic Bonding

Note: The heavier nonmetals from 4,6,5 th groups( In, Tl, Sn, Pb, Sb Bi )may act like metals

An electronegativity difference of approximately 1.7 can only occur between a metal and a nonmetal.

Ionic Bonding

1 Which pair of atoms will form an ionic bond?

A Li and Ne

B K and Br

C K and Cs

D S and Cl

2 Which pair of atoms will form an ionic bond?

A Li and Be

B Na and Mg

C K and Ca

D Na and Cl


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The Octet RuleIn forming compounds, atoms tend towards the electron configuration of a noble gas (s2p6 configuration with 8 electons in the outer shell)

Atoms in the first three groups (1 3) will lose electrons and become cations (think: metals).

Atoms in groups 5A 7A (or 1517) will gain electrons and become anions (think: nonmetals).

Atoms in group 8 (or 18) are the noble gases, which do not generally gain or lose electrons.

Note: The heavier nonmetals from 4,6,5 th groups( In, Tl, Sn, Pb, Sb Bi )may act like metals

• The charges of the ions of representative elements (sblock and pblock elements) can be determined by the position on the periodic table.

• Recall that elements shown here gain or lose electrons to establish a noble gas electron configuration.

Ions formed by Representative Elements

Valence Electrons

Valence electrons are the electrons in the highest occupied energy level of an element’s atoms.

The number of valence electrons largely determines the chemical properties of an element.

To find the number of valence electrons in an atom of a representative element, simply look at its group number.

Atoms in group 3 have 3 valence electrons, atoms in group 7 have 7 valence electrons, etc.

3 How many valence electrons does Aluminum have?

A 5

B 7

C 3

D 27

4 How many valence electrons does Barium have?

A 1B 2

C 52

D 3

The Formation of Cations

• Metals usually give up valence electrons • This results in a noble gas (8 electron) outer shell.

The configuration of the Sodium ion is the same as Neon

Na : 1s2 2s2 2p6 3s1 Na+1 : 1s2 2s2 2p6

Loss of valence electrons

Ne atom


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Na atom Na+ ion

loses e

11p11e

11p10e

• Cations of Group 1A elements always have a charge of 1+.

• Cations of Group 2A elements always have a charge of 2+.


The gain of negatively charged electrons by a neutral atom produces an anion.

• An anion is an atom or a group of atoms with a negative charge.

• The name of a monoatomic anion typically ends in ide.

• Nonmetals from Group 5 and above gain electrons to form anion.

• Group 4 elements do not form cations or anions

The Formation of AnionsA gain of one electron gives chlorine an octet and converts a chlorine atom into a chloride ion.

The Formation of Anions

A chloride ion has the same electron configuration as argon.

Cl: 1s2 2s2 2p6 3s2 3p5 Cl 1s2, 2s2, 2p6, 3s2, 3p6 Ar atom


Cl atom Cl ionGains an e

17P17e

17p18e

• Anions of Group 5A elements have a charge of 3

• Anions of Group 6A elements always have a charge of 2

• Anions of Group 7A elements have a charge of 1



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5 The ion formed from a calcium atom

A Ca+

B Ca2+

C Ca

D Ca2

6 Nonmetals tend to lose electrons forming ions

True

False

7 Metals lose electrons to form cations

True

False

8 The ion formed from nitrogen

A N

B N2

C N3+

D N3

9 Anions are formed from nonmetals

True

False

Formation of Ionic Compounds

Compounds composed of cations and anions are called ionic compounds.

Although they are composed of ions, ionic compounds are electrically neutral.


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The electrostatic forces that hold ions together in ionic compounds are called ionic bonds.

Ionic bonds are formed by the "harpoon method".

Ionic BondsFor instance, when sodium and chlorine are close together, sodium's valance electron flies off and "harpoons" the chlorine atom.

The result is a sodium cation (+) next to a chloride anion ()

These oppositely charged two ions attract: they reel one another together to form an ionic bond.

Ionic Bonds

1s2 2s2 2s2 2p6 3s1 1s2 2s2 2p6 3s2 3p5 1s2 2s2 2p6 1s2 2s2 2p6 3s2 3p6Na Cl Na+ Cl

1s2 2s2 2p6 1s2 2s2 2p6 3s2 3p6

Ne Ar

Na Cl

Ionic Bonds

The electron transfer process in creating an ionic bond:

The dots represents the valence electrons in an atom.

http://www.tutorvista.com/content/chemistry/chemistryii/chemicalbonding/ionicbondinganimation.php

http://www.visionlearning.com/library/flash_viewer.php?oid=1349&mid=55

Ionic Bonds

• A chemical formula shows the kinds and numbers of atoms in the smallest representative unit of a substance.

• A formula unit is the lowest wholenumber ratio of ions in an ionic compound.

• Every ionic compound has a 3D array of positive and negative ions.

Formula Units Properties of Ionic Compounds

• They are crystalline solids at room temperature• They have high melting points• They conduct electricity when melted (molten) or dissolved in water (aqueous)

∗∗

http://www.tutorvista.com/content/chemistry/chemistry-ii/chemical-bonding/ionic-bonding-animation.php

http://www.visionlearning.com/library/flash_viewer.php?oid=1349&mid=55


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Predicting an Ionic Compound FormulaPotassium (K) with an electronegativity of 0.8 and oxygen (O) with an electrogegativity of 3.5 will form an ionic compound.

What is the formula for an ionic compound of potassium and oxygen?

How many additional valence electrons does oxygen want?2

How many valence electrons does potassium have?1

How many potassium atoms will it take to give oxygen the electrons it needs? 2

The formula unit is K2OK

OK Always Metal First( low electonegativity)

Predicting an Ionic Compound Formula

What is the formula for an ionic compound of Mg and N?

How many additional valence electrons does N want?3

How many valence electrons does Mg have to offer?2

How many Mg atoms will it take to give how many N the electrons it needs? (Find the lowest common multiplier first.)

3 Mg : 2 N

The formula unit is Mg3N2

MgN

N

Mg

Mg

If you don't like finding least common multipliers, you can use this alternative method:

1. Write down the ions side by side along with their charge. Always write the metal first.

2. "Crisscross" the numerical values of the charges.

3. Reduce subscripts to lowest ratio.


MgN

N

Mg

Mg

Example: Write the formula for calcium sulfide.

Step 1: Identify the cation & write its common ion

Calcium is in group 2 Ca2+

Step 2: Identify the anion & write its formula

Sulfur is in group 6 S2

Step 3: Crisscross; reduce subscripts if necessary

Ca2+ S2

Ca2S2

CaS


What is the compound formed between Mg and S?

Mg+2 S2 Mg2S2

Always use the lowest ratio of the ions! = MgS

Predicting an Ionic Compound Formula 10 The formula for the ionic between Cs and O

A CsO2

B OCs2

C Cs2O

D OCs2


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11 The ionic compound between Ca and N

A CaNB Ca2N2

C Ca3N2

D Ca2N3

12 The ionic compound between Al and O

A Al3O2

B Al2O3

C AlOD Al2O2

13 What is the ionic compound formed between Ca and Al?

A CaAl

B Ca3Al2

C Al2Ca3

D No compound

14 What is the ionic compound formed between P and Br?

A P3BrB BrP

C no ionic compoundD (BrP)2

15 What is the formula for sodium phosphide?

A SP3B NaPC Na3PD NaP3

16 What is the formula forstrontium bromide?

A SrBrB SrBr2C Sr2BrD BrSr2


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17 The formula for barium sulfide is Ba2S2.

True False

Many cations have the same name as the original, neutral atom.

Naming Binary Ionic Compounds Cations

Charge formula name_________________________________

1+ H+ Hydrogen ion Li+ Lithium ion K+ Potassium ion Cs+ Cesium ion Ag+ Silver ion

2+ Mg2+ Magnesium ion Ca2+ Calcium ion Ba2+ Barium ion Zn2+ Zinc ion Cd2+ Cadmium ion

3+ Al3+ Aluminum ion

• All monoatomic anions all end in "ide".

• The ions that are produced from Group 7A (or 17) elements are called halide ions.

Group 15 Group 16 Group 17

Nitride N3 Oxide O2 Fluoride F

Phosphide P3 Sulfide S2 Chloride Cl

Bromide Br

Iodide I

Naming Binary Ionic Compounds Anions

• Binary (twoelement) compounds are named by writing the name of the cation followed by the name of the anion.

• The name of the cation is the same as the metal name.

• The name of the anion is the name of the nonmetal with the suffix changed to ide.

• Binary compounds end in "ide."

Examples:NaCl = sodium chloride KI = potassium iodideLi2S = lithium sulfide

Naming Binary Ionic Compounds

18 Na2S is

A Sodium sulfateB Sodium sulfideC Disodium sulfide

D Sulfur nitride

19 The correct name for SrO is __________.

A strontium oxide B strontium hydroxide C strontium peroxide D strontium monoxide E strontium dioxide


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20 The correct name for Al2O3 is __________.

A aluminum trioxide B dialuminum oxide C dialuminum trioxide D aluminum oxide E aluminum hydroxide

• Recall that sblock metals have only one possible ionic charge, based on the Octet Rule.

• However, most transition metals can have more than one ionic charge. For this reason, there is a system for designating each ion.

• Sn, Pb from the p block will form more than one type of ions and behave like transition metals.

Cations formed by Transition Elements

Cations Formed by Transition Elements

• Only common transition metals are shown.

• Silver, cadmium and zinc only form one cation, Ag+, Cd2+ and Zn2+

• Note the mercury cations.

• Tin and Lead act like transition metals

Cations formed by Transition ElementsWe will use the Stock naming system (Roman numerals) to name transition metals.

Formula Name_____________________________Cu+1 Copper (I) ionCo+1 Cobalt (II) ion

Fe+2 Iron (II) ionMn+2 Manganese (II) ionPb+2 lead (II) ion

Cr+3 Chromium (III) ionFe+3 Iron (III) ion

Writing Formulas with Transition Metals The charge on the cation is indicated by the Roman numeral, as shown in these examples.

Iron (III) oxide: Fe3+ O2 Write ion formulas.

Fe3+ O2 Crisscross charges.

Fe2O3 Reduce if necessary.

Tin (IV) oxide Sn4+ O2 Write ion formulas.

Sn4+ O2 Crisscross charges. Sn2O4 Reduce if necessary.

SnO2

21 The name of FeCl3 is

A iron chlorideB iron (II) chlorideC iron (III) chloride


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Formulas with Transition Metals In order to correctly name a formula containing a transition metal, it is necessary to first determine the charge on the cation.

Since all compounds are neutral, then the total positive cation charge must equal the total negative anion charge. In other words,

Total cation charge + Total anion charge = 0

(charge of cation) (# of cations) + (charge of anion) (# of anions) = 0

In the case of FeCl3, we make the following substitutions:

(charge of cation) (# of cations) +(charge of anion) (# of anions) = 0(x) (1) + (1) (3) = 0

Thus x = 3 and the cation is Fe3+ or iron(III).

Formulas with Transition Metals A short cut method is to "uncrisscross" the ions, but you must always double check your ions (or you'll get in trouble!).

FeCl3: Fe1 Cl3 Uncrisscross.

Fe3+ Cl1 Check the ions Cl does form a 1 ion and Fe3+ is Iron (III)

Iron (III) Cloride

CrO: Cr1 O1 Uncrisscross.

Cr+ O Check the ions O forms a 2 ion and Cr+ does not exist! (this formula

had to be reduced from Cr2O2)Cromium (II) Oxide

22 Which metal is capable of forming more than one cation?

A KB CsC BaD AlE Sn

23 What is the charge on zirconium ion in ZrO2

A 2+B 4+ C 1+D 2

∗

24 The formula for tin (IV) oxide is

A SnO B Sn2O C SnO2 D SnO

∗ 25 The formula for copper (II) sulfide is

A CuS2

B CuS C Cu2 S2

D (CuS)2


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26 Which one of the following compounds is copper(I) chloride?

A CuCl

B CuCl2C Cu2ClD Cu2Cl2E Cu3Cl2

27 The charge on the cation in the salt Fe2O3 is __________. A 1+

B 2+

C 3+

D 5

E 6

A polyatomic ion is a group of atoms bonded together that have a charge and acts like a single unit or ion

• They are not free compounds and are not found free in nature. • They are found combined with other ions.

Eg: Sulfate = (SO4 )2

Nitrate = (NO3)

Carbonate = (CO3) 2

• Use " ( ) " to keep the atoms together. • Do not change the subscripts inside the "( )"

Polyatomic Ions

• Most of the polyatomic ions contain oxygen atoms

• Many anions names end with “ite” or “ate”

In “ite/ate” pairs, the ion with fewer oxygen atoms will have the “ite” ending

• Examples:sulfite /sulfate nitrite /nitrate

• Note that the suffix does not indicate the actual number of O atoms.

Polyatomic Ions (con't)

Polyatomic Ions (con't)• Familiarize yourself with the polyatomic ions on your reference sheet• Be careful of ide, ite, and ate!

H+ = proton or hydrogen ion

or bicarbonate

]

• Ternary ionic compounds contain three or more different elements due to the presence of polyatomic ion(s).

• Just as in binary ionic compounds, the name of the cation is given first, followed by the name of the anion.

• Names of ternary compounds often end in ite or ate.Examples

CaCO3 = calcium carbonate (in eggshells)Zn(C2H3O2)2 = zinc acetateAgNO3 = silver nitrateNa2SO3 = sodium sulfite

Naming Ternary Ionic Compounds


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Writing Formulas for Ternary Ionic Compounds

Ternary ionic compounds are neutral, just like binary ionic compounds. Therefore, the goal is to find the lowest ratio of cations to anions that will yield a neutral compound. This ratio is represented in a formula unit.

Examples of formula units

Ca(CO3 ) Zn(C2H3O2)2

Ag(NO3) Na2(SO3)

Writing Formulas for Ternary Ionic Compounds (con't)

To write a formula, the crisscross method can again be used.

Example: Write the formula for lithium phosphate.

Step 1: Identify the cation & write its formula

Lithium is in group 1 > Li +

Step 2: Identify the anion & write its formula

Phosphate is a polyatomic ion > PO43

Step 3: Crisscross; reduce subscripts if necessary

Li1+ PO43 Li3(PO4)1 or simply Li3(PO4)

When writing formulas with polyatomic ions, there are two important things to remember:

1) It is helpful to use " ( ) " to keep the atoms together, keeping the charge OUTSIDE the ( )

For examplenitrate = (NO3)1

carbonate = (CO3) 2

2) NEVER alter any symbols or subscripts INSIDE inside the "( )"

Writing Formulas for Ternary Ionic Compounds

Ca2+ (NO3) Ca(NO3)2

Example: Write the formula for calcium nitrite.

28 The formula for sodium hydroxide is

A Na (OH)2

B Na(OH)

C Na(OH2)

D Na(HO)

29 The formula for aluminum phosphate is

A Al(PO4 )

B Al3(PO4)

C Al2(PO4)3

D Al3(PO4)3

30 The formula for magnesiumcarbonate is

A Mg2(CO3)B Mg(CO3) C Mg2(CO3)2

D Mg(CO3)2


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31 The formula for calcium sulfate is

A Ca(SO4)

B Ca2(SO4)2C Ca(SO3) D Ca2(SO3)2

32 NaClO is

A sodium chlorate

B sodium chloride C sodium chlorite

D sodium hypochlorite

33 Mg(HCO3)2 is

A Magnesium carbonate

B Magnesium hydrogen carbonate

C Magnesium hydroxide

D Magnesium carboxide

34 Ammonium carbonate is

A (NH4)(CO3)

B (NH4)2(CO3)

C (NH4)(CO3)2

D (NH4)2(CO2)

PRACTICERecognizing and Writing Symbols

Complete the table by filling in the gaps.

PRACTICEWriting Formulas for Ionic Compounds

Complete the table by filling in the formula for the ionic compound formed by each pair of cations and anions, as shown for the first pair.


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Write the formula for the following compounds:

1. Magnesium iodide

2. Calcium sulfite

3. Barium hydrogen carbonate

4. Iron (III) phosphate

PRACTICEWriting Formulas for Ionic Compounds

Notebook Covalent Bonding

and Molecular Compounds

Chemical Bonds

Recall that there are three basic types of bonds:

Ionic The electrostatic attraction between ions

Covalent The sharing of electrons between atoms

Metallic Each metal atom bonds to other metals atoms within a mobile "sea" of electrons (not covered until

AP Chemistry)

Chemical Bonds

Ionic BondingIonic bonds occur when the difference in electronegativity between two atoms is more than 1.7.

Covalent BondingIf the difference of electronegativity is less than 1.7, neither atom takes electrons from the other; they share electrons.

This type of bonding typically takes place between two nonmetals.

In the case of ionic bonding, a 3D lattice of ions is the result...not individual molecules. The chemical formula for an ionic compound is just the ratio of each type of ion in the lattice, not a particular number of ions in a molecule.

In contrast, covalent bonding results in individual molecules; each with its own unique shape. These shapes help determine the physical and chemical properties of everything around us!

Covalent Bonding


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1 Which pair of atoms will form a covalent bond?

A Li and Ne

B K and Br

C C and O

D Na and Cl

2 Which pair of atoms will form a covalent bond?

A Li and I

B Na and Cl

C K and Fl

D H and O

Molecular Compounds

Covalent compounds are formed between two nonmetals.

When atoms are bonded covalently, the atoms are held together by sharing electrons. Such a compound is called a molecular compound (molecule).

In covalent bonds, electron sharing usually occurs so that atoms attain the electron configurations of noble gases. Both atoms used the shared electrons to reach that goal.

Naming Binary Molecular Compounds

• Use prefixes to indicate the number the atoms

• All end in "ide"

Examples

NO2 = nitrogen dioxide

P2O5 = diphosphorous pentoxide ( pentaoxide>pentoxide)

Naming Binary Molecular Compounds

Look on your reference sheets for the prefixes.

The atom with the lower electronegativity is usually written first.

If there is only one of the first atom, the mono is left off.

Examples

CO = carbon monoxide

CO2 = carbon dioxide

3 Chlorine monoxide is

A ClO2

B ClO

C OCl

D O2Cl


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4 Dinitrogen tetroxide is

A NO2

B N2O4

C NO3

D N4O2

5 H2O is

A Hydrogen monoxide

B Dihydrogen monoxide

C Hydrogen oxide

D Hydrogen dioxide

6 SO3 is

A sulfate

B sulfur oxide

C sulfur trioxide

D sulfite

7 MgO is

A monomagnesium monoxide

B magnesium monoxide

C monomagnesium oxide

D magnesium oxide

8 P4O10 is

A Diphosphorous pentoxide

B Tetraphosphorous decoxide

C Phosphorous oxide

D Phosphate

Lewis structures are diagrams that show valence electrons as dots. Lewis structures are also known as Lewis dot or electron dot diagrams.

Note that no electrons are paired until after the fourth one.

Lewis Structures


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9 How many valence electrons does nitrogen have?

A 2

B 3C 4D 5E 7

10 The Lewis structure for

nitrogen is N

True

False

Recall that atoms tend towards having the electron configuration of a noble gas.

For most atoms, that means having 8 valence electrons.

The Octet Rule also applies to molecular compounds.

So, in covalent bonding, an atom will share electrons in an effort to obtain eight electrons around it (except hydrogen which will attempt to obtain 2 valence electrons).

A pair of valence electrons that is not shared between atoms is called an unshared pair, also known as a lonepair or a nonbonding pair.

The Octet Rule How do electron dot structures represent shared electrons?

Two atoms held together by sharing a pair of electrons are joined by a single covalent bond.

H + H H H

Hydrogen atom

Hydrogen atom

Hydrogen molecule

Shared pair of electrons

H

H

1s

1sHydrogen molecule

How do electron dot structures represent shared electrons?

An electron dot structure such as H:H represents the shared pair of electrons of the covalent bond by two dots.

H + H H H

Hydrogen atom

Hydrogen atom

Hydrogen molecule


H

H

1s

1sHydrogen molecule

Structural Formulas

A structural formula represents the covalent bonds by dashes and shows the arrangement of covalently bonded atoms. As in the example below, one shared pair of electrons is represented by one dash.

HH

Hydrogen molecule


H Hmov

e me


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11 How many electrons are shared by two atoms to create a single covalent bond?

A 2

B 1

The halogens form single covalent bonds in their diatomic molecules. Fluorine is one example.

Single Covalent Bonds

F F F F F F+ −−> OR

Fluorineatom

Fluorine molecule

Fluorineatom

1s

2s

2p

1s

2s

2p

Fluorine molecule

In a water molecule, each hydrogen and oxygen atom attains a noblegas configuration by sharing electrons.

Lewis Structures

The water molecule has two unshared, or lone, pairs of electrons.

2 H + O > O H or O HH

HHydrogen atoms

Oxygen atom

Water molecule

1s 2p2s

1s 1s

O

H H

Water molecule

In the ammonia molecule, NH3, each atom attains a noblegas configuration by sharing electrons. This molecule has one unshared pair of electrons.

Lewis Structures

3 H + N > N H or N H

H

H H

HHydrogen atom

Nitrogen atom

Ammonia molecule

1s 2p2s

1s 1sH

N

H1s

Ammonia molecule

Drawing Lewis Structures

First, find the total number of valence electrons in the polyatomic ion or molecule.

If it is an anion, add an electron for each negative charge.

If it is a cation, subtract an electron for each positive charge.

The P atom has 5 valence electrons.

A Cl atom has 7, and there are three of them.

The total number of valence electrons is:


The central atom is the least electronegative element (excluding hydrogen).

Connect the other atoms to it by single bonds.

P has an electronegativity of 2.1 and Cl has an electronegativity of 3.0,

P will be the central atom.

The Cl atoms will surround the P atom.

The single bonds are shown as single lines.


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1. Count each single bond as a pair (two) of electrons.

2. Add electons to the outer atoms to give each one 8 (a full shell), or just 2 electrons for hydrogen.

3. Do the same for the central atom.

4. Check: Does each atom have a full outer shell (8 except, 2 for hydrogen)?

Have you used up all the valence electrons? Have you used too many electrons?

Drawing Lewis Structures Drawing Lewis Structures

First, find the total number of valence electrons in the polyatomic ion or molecule.

If it is an anion, add an electron for each negative charge.

If it is a cation, subtract an electron for each positive charge.

The N atom has 5 valence electrons

and each of the three H atoms has 1 so the total number of valence electrons is,

NH3

5 + 3(1) = 8


The central atom is the least electronegative element (excluding hydrogen because it can only have one bond).

Connect the other atoms to it by single bonds.

H can never be the central atom so N must be

The H atoms will surround the N atom.

The single bonds are shown as single lines.

HN HH

NH3


HN HH

Each H already has two electrons, so that's done. But we have to add electrons to N to make 8.

HN HH

Count each single bond as a pair (two) electrons. Now add electons to the outer atoms to give each one a full shell (2 in the case of H).

Next, do the same for the central atom.

Check:Does each atom have a full outer shell ?

Have you used up all the valence electrons you started with? Have you used too many electrons?

VSEPR Theory Valence Shell Electron Pair Repulsion

According to VSEPR theory, the molecules will adopt a shape/geometry so as to reduce the repulsion between the bonded electrons.

The VSEPR number of a molecule is a three digit number that can be used to determine a molecule's shape. Here's how you find it.

1. Draw the Lewis structure for the molecule. Locate the central atom, if applicable.

2. The first digit of the VSEPR number is the total number of electrondomains around the central atom.

Electron domains are either shared pairs of electronsor lone pairs of electrons

Multiple bonds (i.e. double or triple bonds) count as only ONE electron domain.

VSEPR Numbers


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3. The second digit of the VSEPR number is the total number of bondingdomains around the central atom.

Bonding domains are single, double or triple bonds.

4. The third digit of the VSEPR number is the total number of lone pairs around the central atom.

Each pair of electrons that are not involved in bondscounts as one lone pair.

5. Check your work the first digit is equal to the sum of the second and third.

VSEPR Numbers

C

NCl

F

OSB

P

I

H

Si

SeXe

CH4 Draw a Lewis Structure and use that to determine the VSEPR number

H

H

HC

H

Check to make sure that each atom has a full outer shell. Now calculate the VSEPR #.

Electron domains = 4Bonding domains = 4Lone pairs of electrons = 0

Its VSEPR number is 4 4 0

Slide for Answer

C

NCl

F

OSB

P

I

H

Si

SeXe

NF3Draw a Lewis Structure and use that to determine the VSEPR number

N F

F

F




Slide for Answer

C

NCl

F

OSB

P

I

H

Si

SeXe

SiF4Draw a Lewis Structure and use that to determine the VSEPR number

F

Si

F

F F




Slide for Answer

Lewis Structures

If you are drawing the Lewis Structure for an ION...

A negative ion has extra electrons, add the charge of the ion to your valence electron count.

ClO2 has 1(7) + 2(6) + 1 = 20 electrons

A positive ion is missing electrons, subtract the charge of the ion to your valence electron count.

NH4+ has 1(5) + 4(1) 1 = 8 electrons

C

NCl

F

OSB

P

I

H

Si

SeXe

PO43Draw a Lewis Structure and use that to determine the VSEPR number

O

P

O

O O




Slide for Answer


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August 17, 2012

Lewis Structures

Draw the Lewis dot structure for the sulfate ion, SO4 2

Lewis StructuresDraw the Lewis dot structure for the hydronium ion, H3O+

C

NCl

F

OSB

P

I

H

CO OSi

SeXe

CO2

Draw a Lewis Structure and use that to determine the VSEPR number

We ran out of electrons, but carbon does not have an octet yet!

Now What?

Slide for Answer

Double and Triple Covalent Bonds

Atoms form double or triple covalent bonds if they can attain a noble gas structure by sharing two pairs or three pairs of electrons.

A bond that involves two shared pairs of electrons is a double covalent bond.

A bond formed by sharing three pairs of electrons is a triple covalent bond.

Carbon Dioxide, CO2

1. Determine the # of valence electrons.

1 (4) + 2 (6) = 16 e

2. Form single bonds.This leaves 12 electrons, 6 pairs

3. Place lone pairs on oxygen atoms to give each 8.


O C O

O C OO C O

Carbon Dioxide, CO2

4. Check: We had 16 electrons to work with; how many have we used?

5. There are too many electrons in our drawing. We must form DOUBLE BONDS between C and O.

Instead of sharing only 1 pair, a double bond shares 2 pairs. So one pair is taken away from each atom and replaced with another bond.

O C O

O C O


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Double and Triple Covalent BondsDouble and Triple Covalent Bonds

It requires more energy to break double and triple bonds compared to single bonds. Triple bonds are the strongest of the three.

Bond Type Bond Energy

C C

C C

C C

348 kJ

614 kJ

839 kJ

Bond Type Bond Energy

N N 163 kJ

418 kJ

941 kJ

N N

N N

Double and Triple Covalent BondsDouble and Triple Covalent Bonds

This chart compares bond strength and bond length for single, double, and triple bonds.

Type of Bond

Electrons Shared

Bond Strength

Bond Length

Single 2 weakest longest

Double 4 intermediate

intermediate

Triple 6 strongest

shortest

12 As the number of bonds between a pair of atoms increases, the distance between the atoms:

A increases

B decreases

C remains unchanged

D varies, depending on the atoms

13 As the number of bonds between a pair of atoms increases, the strength of the bond between the atoms:

A increases

B decreases

C remains unchanged



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14 As the number of bonds between a pair of atoms increases, the energy of the bond between the atoms:

A increases

B decreases

C remains unchanged


15 How many electrons are shared by two atoms to create a single bond?

16 How many electrons are shared by two atoms to create a double bond?

17 How many electrons are shared by two atoms to create a triple bond?

• If you run out of electrons before the central atom has an octet…

…form multiple bonds until it does.

Writing Lewis Structures

Oxygen molecule


1s

2s

2p

1s

2s

2p

O + O > O O or O O

O

O

Oxygenatom

Oxygenatom

Oxygenmolecule

Oxygenmolecule


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August 17, 2012

C

NCl

F

OSB

P

I

H

CSi

SeXe

CO


Carbon has the lower electronegativity, so we will consider it the "central" atom...

Electron domains = 2Bonding domains = 1Unpaired electrons = 1


O

Slide for Answer

Coordinate Covalent Bonds


In carbon monoxide, oxygen has a stable configuration but the carbon does not.

1s 2p2s

2s1s 2p

C + O −−> C O

Carbonatom

Oxygenatom

Carbonmonoxide

C

OCarbon monoxide molecule

A coordinate covalent bond is a covalent bond in which one atom contributes both bonding electrons.

In a structural formula, you can show coordinate covalent bonds as arrows that point from the atom donating the pair of electrons to the atom receiving them.

In a coordinate covalent bond, the shared electron pair comes from one of the bonding atoms.

Carbon has 4 valence electrons, oxygen has 6.


C

NCl

F

OSB

P

I

H

Si

SeXe

F2


Electron domains = 4Bonding domains = 1Unpaired electrons = 3


F F

Slide for Answer

A molecule is a neutral group of atoms joined together by covalent bonds. Air contains oxygen molecules.

A diatomic molecule is a molecule consisting of two atoms. Certain elements do not exist as single atoms; they always appear as pairs.

When atoms turn into ions, this NO LONGER HAPPENS!

• Hydrogen• Nitrogen• Oxygen• Fluorine• Chlorine• Bromine• Iodine

Remember:HONClBrIF

Diatomic Molecules


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August 17, 2012

C

NCl

F

OSB

P

I

H

Si

SeXe

O3


For the central oxygen:Electron domains = 3Bonding domains = 2Unpaired electrons = 1


O

O OSlide for Answer

Consider the Lewis structure we would draw for ozone, O3:

We would expect the double bond to have a shorter bond length than the single bond.

However, the true, observed structure of ozone shows that both OO bonds are the same length. How can this be?

Resonance

O

OO

O

O

O

∗

One Lewis structure cannot accurately depict a molecule like ozone. Therefore, we use multiple structures, called resonance structures, to describe the molecule.

Ozone has two resonance structures.

Resonance

O

O

OO

O

O

∗ ResonanceThe actual ozone molecule is a synthesis of these two resonance structures.

The bond length for both outer oxygen atoms falls somewhere between the single and double bond length.

O

O

OO

O

O

Resonancestructure

Resonancestructure

Ozone molecule

∗

Resonance

The nitrate ion, NO31 also requires resonance structures to explain its covalent bonding.

There are three resonance structures for the nitrate ion:

∗

Draw the Lewis dot structure for SO3:

Lewis Structures


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18 How many resonance structures can be drawn for the carbonate ion, CO32 ?

A 1

B 2

C 3

D 4

E 5

∗The benzene molecule is a regular hexagon of carbon atoms with a hydrogen atom bonded to each one. There are two resonance structures for benzene.

Resonance

Benzene, C6H6, is obtained from the distillation of fossil fuels. More than 4 billion pounds of benzene is produced annually in the United States. Because benzene is a carcinogen, its use is closely regulated.

∗

ResonanceIn truth, the shared pairs of electrons do not always remain between adjacent C atoms. They are not localized.

Instead, the electrons are said to be delocalized, meaning that they they can move around the 6carbon ring.

Benzene is commonly depicted as a hexagon with a circle inside to signify the delocalized electrons in the ring... we will talk more about this at the end of the year when we study organic chemistry.

<−−> or

∗ Exceptions to the Octet Rule

There are three types of ions or molecules that do not follow the octet rule:

• Ions or molecules with an odd number of electrons

• Ions or molecules with less than an octet

• Ions or molecules with more than eight valence electrons (an expanded octet)

Though relatively rare and usually quite unstable and reactive, there are ions and molecules with an odd number of electrons.

NO is an example:

Odd Number of ElectronsFewer Than Eight Electrons

Beryllium (Be) this metal is shown to form molecular compounds, rather than ionic compounds as expected; only needs 4 electrons to be stable

Boron (B) only needs 6 electrons to be stable

Memorize these exceptions


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Exceptions to the Octet Rule

Draw the Lewis dot structure for boron trifluoride, BF3:

Slide for Answer

Exceptions to the Octet Rule Expanded Octet

Draw the Lewis dot structure for phosphorous pentachloride, PCl5:

• The only way PCl5 exists is if phosphorus has 10 electrons around it.

• This is called an expanded octet.

• Atoms on the third energy level or higher are allowed to expand their octet to 10 or 12 electrons.

• The d orbitals in these atoms participate in bonding, allowing the expanded octet.

Exceptions to the Octet Rule Expanded Octet Exceptions to the Octet Rule

How many electrons do these central atoms have around them?

Draw the Lewis dot structure for the xenon tetrafluoride, XeF4.

Lewis Structures

Draw the Lewis dot structure for the iodine tricholoride, ICl3.

Lewis Structures


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VSEPR stands for ValenceShell ElectronPair Repulsion

The VSEPR model uses electrondot diagrams to help predict molecular shapes based on this fact:

In a molecule, each electron pair will position itself as far away as possible from other electron pairs in order to

minimize repulsion.

VSEPR Model How does VSEPR theory help predict the shapes of molecules?

According to VSEPR theory, the repulsion between electron pairs causes molecular shapes to adjust so that the valenceelectron pairs stay as far apart as possible.

The shape of a molecule plays an important role in determining its chemical and physical properties.

To determine a molecule's shape, i.e. its molecular geometry, we must first determine its electrondomain geometry.

Recall:

• Electron domains are either shared pairs of electrons or lone pairs of electrons

• Bonding domains are single, double or triple bonds.

• Each pair of electrons that are not involved in bonds counts as one lone pair.

To determine the electrondomain geometry, look at the first number and use the following chart...

How does VSEPR theory help predict the shapes of molecules? How does VSEPR theory help predict

the shapes of molecules?

ElectronDomain Geometry (EDG)

The EDG (2,3,4,5,or 6) gives us the general shape of the molecule, as shown here.

However, these domains do not have to be bonds.

The molecular geometry tells us if there is a bond or lone pair of electrons present, thereby specializing the general shape.

Let's take a closer look...

Linear ElectronDomain Geometry

Linear

Two atoms around a central one will form a linear shape with bond angles of 180o


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Linear Molecular Geometry

There is only one molecular geometry for linear electrondomain: linear molecular geometry (220).

Trigonal Planar ElectronDomain Geometry

trigonal planar

Three atoms around a central one will form a trigonal planar shape with bond angles of 120o

Trigonal Planar Molecular Geometry

There are two molecular geometries:• Trigonal planar, if all the electron domains are bonding (330)• Bent, if one of the domains is a nonbonding pair (321)

trigonal planar(330)

bent(321)

Trigonal Planar Molecular Geometry

Tetrahedral ElectronDomain Geometry

Four atoms around a central one will form a tetrahedral shape with bond angles of 109.5o

tetrahedral

Tetrahedral Molecular Geometry

There are three molecular geometries:• Tetrahedral, if all are bonding pairs (440)• Trigonal pyramidal, if one is a nonbonding pair (431)• Bent, if there are two nonbonding pairs (422)


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Tetrahedral Molecular Geometry

tetrahedral (440)

trigonal pyramidal(431)

bent (422)

Five atoms around a central one will form a trigonal bipyramidal shape with bond angles of 120o and 90o

trigonal bipyramidal

Trigonal Bipyramidal ElectronDomain Geometry

Trigonal Bipyramidal Molecular Geometry

Trigonal bipyramidal

Seesaw

Tshaped

Linear

Trigonal Bipyramidal

(550)

SeeSaw (541)

TShape (532)

Linear (523)

Trigonal Bipyramidal Molecular GeometryThere are four molecular geometries for the trigonal bipyramidal electron domain geometry:

Six atoms around a central one will form an octahedral shape with bond angles of 90ο

octahedral

Octahedral ElectronDomain GeometryOctahedral Molecular Geometry

Square Planar

Octahedral

Square Pyramidal


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Octahedral (660)

Square Pyramidal

(651)

Square Planar (642)

Octahedral Molecular Geometry

There are only three molecular geometries for the octahedral electron domain geometry:

How does VSEPR theory help predict the shapes of molecules?

Using VSEPR numbers, you can determine molecular geometry.VSEPR numbers are a set of 3 numbers.

1) the total number of electron domains2) the number of bonding domains*3) the number of unshared pairs of electrons

(* Remember that multiple bonds count as ONE domain)

Electrondomain geometry has the same name as the first shape.

Draw the Lewis structure for ammonia, NH3.

What are the VSEPR numbers for NH3?

4,3,1

What is the electrondomain geometry of NH3?tetrahedral

What is the molecular shape of NH3?trigonal

pyramidal

VSEPR Numbers and Molecular Geometries

N

H

HHSlide for Answer

Draw the Lewis structure for ClF3. Note that molecules containing ONLY halogens will usually violate the octet rule.

What are the VSEPR numbers for ClF3?

5,3,2

What is the electrondomain geometry of ClF3?trigonal

bipyramidal

What is the molecular shape of ClF3?Tshape

VSEPR Numbers and Molecular Geometries

Slide for Answer

19 The methane molecule (CH4) has which geometry?

A linear

B trigonal bipyramidal

C trigonal planar

D tetrahedral

20 Give the VSEPR number for this molecule.

∗


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F Xe F

23 Which compound below contains an atom that is surrounded by more than an octet of electrons?

A PF5

B CH4

C NBr3

D OF2

∗

Polarity of BondsThough atoms often form compounds by sharing electrons, the electrons are not always shared equally. In a covalent bond, one atom has a greater ability to pull the shared pair toward it.

Polarity of Bonds

Identical atoms will have an electronegativity difference of ZERO. As a result, the bond is NONPOLAR.

Polarity of Bonds

BOND TYPE ELECTRONEGATIVITY DIFFERENCE

NonPolar Covalent zero, or very small

Polar Covalent about 0.2 to 1.6

Ionic

above 1.7(between metal & nonmetal)


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Therefore, the fluorine end of the molecule has more electron density than the hydrogen end.

Polarity of Bonds

H F

We use the symbol to designate a dipole (2 poles). The "+" end is on the more positive end of the molecule and the arrow points towards the more negative end.

• When two atoms share electrons unequally, a bond dipole results.

Electronegativities

Polarity of Bonds

Compound Bond Electronegativity Dipole length (A0) Differenece Moment (D)

HF 0.92 1.9 1.82HCl 1.27 0.9 1.08HBr 1.41 0.7 0.82HI 1.61 0.4 0.44

Bond lengths, Electronegativity, Differences and Dipole Moments of the Hydrogen Halides

But just because a molecule possesses polar bonds does not mean the molecule as a whole will be polar.

Polarity of Molecules

For instance, in the case of CO2:

The polar bond is shown as a dipole, the arrow points to the more negative atom. Dipoles add as vectors.

Polarity of MoleculesBy adding the individual bond dipoles, one can determine the overall dipole moment for the molecule.

For a molecule to be polar, it must a) contain one or more dipoles AND b) have these polar bonds arranged asymmetrically

In other words, if all the dipoles are symmetrical, they will cancel each other out and the molecule will be NONPOLAR.

Many molecules with lone pairs of electrons will be POLAR.

These are some examples of polar & nonpolar molecules. What are their VSEPR numbers?

Polarity of Molecules

330, nonpolar

440, nonpolar

440, polar

431, polar110(?), polar

Slide for Answer

Slide for Answer Slide for Answer

Slide for AnswerSlide for Answer


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24 Which of these are polar molecules?

A a & b

B a, b & c

C a & c

D a, c & d

E c & e

• Hybridization is the process of mixing up of different orbitals to produce a hybrid orbital.

Hybridization of Atomic Orbitals

• The valence shell electronpair repulsion model (VESPR) was devised to account for the molecular shapes. In this model, atoms and pairs of electrons will be arranged to minimize the repulsion of these atoms and pairs of electrons.

• The number of these new hybrid orbitals must be equal to the numbers of atoms and nonbonded electron pairs surrounding the central atom.

• The process of hybridization will tell you where the electron pairs are located, which orbitals are participating and how the molecule gets the shape.

Hybridization

• Mixing the s and p orbitals yields two degenerate orbitals that are hybrids of the two orbitals.• These sp hybrid orbitals have two lobes like a p orbital.• One of the lobes is larger and more rounded as is the s orbital.

s orbital p orbital

sp hybrid orbitalsthe two sp hybrid orbitals together

+

• These two degenerate orbitals would align themselves 180° from each other.• This is consistent with the observed geometry of beryllium compounds: linear.

• The sp orbitals are higher in energy than the 1s orbital but lower than the 2p.

Hybridization

F 2p orbital F 2p orbital

overlap region

sp hybrid orbital

Sigma (σ) Bonds

• Sigma bonds are characterized by• Headtohead overlap.• Cylindrical symmetry of electron density about the internuclear axis.

1s 1s 1s 3p

3p 3p

• Pi bonds are characterized by• Sidetoside overlap.• Electron density above and below the internuclear axis.

Pi bonds (π)

p p

Internuclear axis

π bond


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CH4• we envision that the s electron in Carbon is promoted to the p level in order to make room for the 4 Hydrogen electons to share.

• 4 sp3 hybrid orbitals of same energy and shape are created.

Hybridization

25 what is the hybridization in BeF2

A sp3

B sp2

C sp

D spd

26 what type of hybridization is there in SiCl4

A sp3

B sp2

C sp2dD sp

Hybridization

Hybridization27 Hybridization in water

molecule is

A sp2B sp3C spD sp3d


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August 17, 2012

28 what is the hybridization in Dichloro methane

A sp3B sp2C spD sp3d

Hybridization Involving Multiple BondsThe sigmabond is defined as the linear overlap of ( axial or along the axis) atomic orbitals (hybrids except for hydrogen) in which two electrons are directly between the two bonded nuclei.

Pibonds are defined as the parallel overlap of porbitals. A double bond has one sigmabond and one pibond. A triple bond thus consists of a sigmabond and two pibonds with the pibonds in different planes.

Ethylene molecule, CH2=CH2

Hybridization Involving Multiple Bonds

Acetylene, C2H2


C2H2


Benzene

29 each carbon atom in the above molecule is sp3 hybridized.

True False

Count the total number of electron domains around the atom. That many hybrid orbitals are needed via combination of s, p or d remember double bonds are counted only once.


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30 Name the hybridization in each carbon

A sp3

B sp2

31 CH3CH=CH2 what is the hybridization of the middle carbon atom?

A sp3

B sp

C sp2

D sp4

32 What is the hybridization of the Xe atom in XeF4?

A SP

B Sp3d2

C sp3d

D sp3

33 What is the geometry of a molecule with sp3d hybridization?

A tetrahedron

B octahedron

C trigonal pyramid

D trigonal bipyramid

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