COMPOUNDS AND BONDING. Objectives 1 Identify characteristics in atoms involved in chemical bonding Compare the physical and chemical properties of ionic

COMPOUNDS AND

BONDING

Objectives 1

• Identify characteristics in atoms involved in chemical bonding

• Compare the physical and chemical properties of ionic and covalent compounds

• Compare the arrangement of atoms in molecules, ionic crystals, polymers, and metallic substances

Vocabulary: define in your journal

• Valence electrons

• Electron dot structures

• Octet rule

• Ionic bond

• Anion

• Cation

• Metallic bonds

• Covalent bond

• Structural formula

• Unshared pairs

• Double covalent bonds

• Triple covalent bonds

• Coordinate covalent bonds

• Bonds

• Bond dissociation energy

• Resonance structures

• Nonpolar covalent bond

• Polar covalent bond

• Polar bond

• Polar molecule

• Dipole

• Van der Waals Forces

• Dispersion forces

• Dipole interactions

• Hydrogen bonds

Van der Waals forces

Consist of two kinds1. Dispersion forces: weak force caused by

the motion of electrons, increases with increasing electrons

Seen in diatomic molecules, weak in fluorine and chlorine making them gases

Stronger in bromine causing it to be a liquid

2. Dipole interactions: caused by attraction of polar molecules for each other; occurs between weakly positive and negative molecules

Water is held together by the dipole interactions of adjacent oxygen (-) and hydrogen (+) atoms when water molecules come close together

Hydrogen bonds (+) are strongest of dipole interactions

I. COMPOUND: A. The Chemistry meaning:

Notes: Handout

Materials made from atoms of two or more combined

elements.

Notes

B. Compounds have properties that are unlike the

elements that form them.

Notes

1. Na (Sodium) is a silvery metal that reacts violently

with water.

2. Cl (Chlorine) is a green gas that can kill any animal.

Notes

3. Na + Cl NaCl (table salt) which we eat with little harm.

Notes

II. BONDING:A. The Chemistry meaning:

Notes

Interaction between atoms that results in the formation of

a compound.

Notes

B. CHEMICAL BOND: Strong attractive force between atoms

or ions in a compound.

Notes

C. BOND ENERGY: Energy involved in the making and

breaking of chemical bonds.

Notes

1. Energy is released to make a bond.

exothermic2. Energy has to be applied to

break the bond. endothermic

Notes

3. The bond energy is the lowest potential energy for

the compound.

Notes

D. IONIC BONDS: Atoms gain or lose electrons

Notes

1. Usually happens between a metal and a nonmetal.

Notes

2. Ion - atom which has:a. Gained one or more

electrons - negative charge (anion)

b. Lost one or more electrons - positive charge (cation)

Notes

Lewis Dot Diagram

• Complete the handout Ionic Bonds and Covalent bonds

Quiz 1: Ionic and Covalent Bonds

Draw a Lewis dot diagram showing the following bonds

1.An ionic bond forming calcium phosphide

2.A covalent bond forming carbon monoxide (make sure Carbon and Oxygen both have 8 valence electrons available to them)

Hybrid orbitalsHybrid orbitals• Two overlapping orbitals form what is known

as a hybrid or molecular orbital• Just as in a s,p,d, or f orbital the electrons

can be anywhere in the orbital (even though the electron has started out in one atom, at times, it may be closer to the other nucleus)

• Each hybrid orbital has a specific shape • You do not need to know shapes• You need to know that hybrid orbitals exist

and that they are formed from overlapping orbitals

Overlapping orbitalsOverlapping orbitals• Draw orbital diagrams for F + F, H + O, Li + F

1s 2s 2p

1s2s2p

1s 2s 2p

1s

1s

F2

H2O

1s 2s 1s2s2p

LiF is ionic (metal + non-metal)

Lewis diagramsLewis diagramsDraw Lewis dot diagrams for Ne, Sb, Rb, F.

How many variations of the Lewis diagram for P can be drawn?

Ne Sb Rb F

P P P P

• Lewis diagrams follow the octet rule: atoms when forming ions, or bonding to other atoms in compounds have 8 outer electrons

• Q - How can the octet rule be explained?• A - s (2 e–) and p (6 e–) orbitals are filled

Cl– Na+ Cl–

Ionic bondingIonic bonding• Recall: Ionic bonding involves 3 steps:

1) loss of e-, 2) gain of e-, 3) +ve, -ve attract

Na Cl

e–1) 2)

3)

Na+

This can be represented via Lewis diagrams…• Diagram the reaction between Li + Cl and Mg + O (PE 3)

The octet rule (ionic compounds)The octet rule (ionic compounds)• Draw Li + Cl and Mg + O (PE 3, pg. 230)

Li Cl [ Cl ]–[Li]+

[ O ]2–[Mg]2+OMg

• Note also that the charge on an ion can be determined by the number of places removed from a noble gas (Ca, N, Al?)

• Ca2+, N3-, Al3+

Covalent bondingCovalent bonding• Covalent bonds can also be shown via Lewis

diagrams - E.g draw Lewis diagrams showing the combination of 1) H+Cl, 2) C+Cl, 3) H+O, 4) Mg+F, 5) N+H, 6) Do PE 4 (pg. 234)

HCl

H Cl

CCl4

C

Cl

Cl

Cl

Cl H2OH O H

MgF2 - Ionic

[ F ]2– [Mg]2+

NH3

H N H

H

• Note bonds can also be drawn with a dash to represent two electrons

E. COVALENT BONDS: Atoms share electrons rather

than losing them

Notes

1. Usually between nonmetals

2. Sharing can be unequal causing the compound to seem to have a charge.

Notes

a. One nucleus is larger(has more protons)

and acts like a bigger magnet.(And because it has

a higher electronegativity)

Notes

b. The electrons spend more time around the larger

nucleus than around the smaller nucleus.

Notes

3. Water for example seems to have a charge, but it is a

covalent compound.

Notes

a. The oxygen nucleus is larger(has more protons)

and acts like a bigger magnet. (actually it has a higher

electronegativity)

Notes

b. The electrons spend more time around the oxygen nucleus than around the

hydrogen nuclei.

Notes

Structure Determination Structure Determination by VSEPRby VSEPR

Structure Determination Structure Determination by VSEPRby VSEPR

Water, HWater, H22OO The electron pair The electron pair geometry is geometry is TETRAHEDRALTETRAHEDRAL

The electron pair The electron pair geometry is geometry is TETRAHEDRALTETRAHEDRAL

The molecular The molecular geometry is geometry is BENTBENT..

The molecular The molecular geometry is geometry is BENTBENT..

H O H••

••

H O H••

••

2 bond 2 bond pairspairs

2 lone 2 lone pairspairs

Answer the following

1. Is the electronegativity values of nonmetallic elements greater or less than the electronegativity values of metallic elements?

2. In a polar covalent bond does the more electronegative atom have a slight positive or negative charge?

3. How many valence electrons do Oxygen and Sulfur have?

4. How many valence electrons will Nitrogen gain or lose to be isoelectronic with a noble gas?

5. What is the name of an ion that has gained electrons?

Types of covalent bonds(not on paper)

• Single covalent bond: a bond formed when a pair of electrons is shared between two atoms

• Double covalent bond: a bond formed when two pairs of electrons are shared between two atoms

• Triple covalent bonds: a bond formed when three pairs of electrons are shared between two atoms

• Coordinate covalent bond: a covalent bond formed when one atom contributes both bonding electrons

Answer the following

Double covalent bond electronegativity isoelectronic covalent bond triple covalent bond

• Attraction for a shared pair of electrons• Formed by sharing electrons• Having the same number of electrons• Sharing four electrons• Sharing six electrons

Pi Bond: caused by overlap of p orbitals

Hydrogen bonds

#11. Try this one

• Draw a structure for phosphate PO4-3

#12 Practice

• Mg + F2 --> Mg F2

- Draw a Bohr model showing the reaction

- Draw a Lewis dot diagram showing the reaction

- Draw a model of the reaction using the electron configuration diagrams including the boxes

#13. Try this

VSEPR: Lone PairsVSEPR: Lone Pairs

Lone pairsLone pairs• Thus far we have considered (built) only

structures where there are no free electrons around the central atom

• These electrons that are not involved in bonds are called “lone pairs”

• Essentially, they have the same influence on molecular structure as electron pairs in bonds

• The result is some weird shapes and names…

orVs.

orH C

H

H

H

H C H

H

H

H N H

H

H N H

H

Variations on Tetrahedral MoleculeVariations on Tetrahedral Molecule• The tetrahedral molecule is AX4

• Lone pairs can be indicated with AXYEZ, where Z is the number of lone pairs

• By replacing 1 bond with a lone pair the tetrahedral shape becomes “trigonal pyramidal”

• AX3E

• By replacing two bonds with lone pairs we get a “bent” (non-linear) shape (AX2E2 )

Variations on Trigonal BipyramidalVariations on Trigonal Bipyramidal• AX5 is trigonal

bipyramidal

• AX4E is unsymmetrical tetrahedron

• AX3E2 is T-shaped

• AX2E3 is linear

Variations on Octahedral ShapeVariations on Octahedral Shape• AX6 is octahedral

• AX5E is square pyramidal

• AX4E2 is square planar

For more lessons, visit www.chalkbored.com

http://www.chalkbored.com/

Answer these

Given elements with electronegativities as follows: L: 1.9, M:3.0, Q: 0.9, R: 2.5

14. Which combination would be the least polar?

15. Which combination would be the most polar?

16. How many valence electrons in SO4-2?

Compounds

• Compounds are pure substances made of more than one kind of atom

• Obey the law of definite proportions by always combining in the same proportions by mass

• A molecule is a neutral group of atoms that act as a unit

• Compounds composed of molecules are molecular compounds– Tend to have relatively low melting and boiling

points– Many exist as gases or liquids at room

temperatures– Most are composed of two or more

nonmetallic elements– Examples: water and carbon dioxide

• Compounds composed of positive and negative ions are called ionic compounds– Arranged in orderly 3-dimensional pattern

– Each positive ion between two or more positive ions and at the same time each negative ion is between two or more positive ions.

– Are electrically neutral

– Most are crystalline solids at room temperature

– Usually formed from a metallic and nonmetallic element

– Example: sodium chloride (table salt)

Define the following as we go through the lesson

17. Chemical formula18. Molecular formula

19. Subscript

20. Ionic compound

21. Simplest formula

22. Formula unit

23 to 25. Describe ionic and molecular compound characteristics

Chemical Formulas

• Chemical formulas show the kinds and numbers of atoms in the smallest representative unit of the substance

• The chemical formula for a molecular compound is called the molecular formula– The number of atoms of each kind is

indicated by a subscript written after the formula

– Examples: water = H20, carbon dioxide = CO2

• Requires a diagram to show the arrangement of the atoms

• Ionic compounds are composed of equal amounts of each type of atom arranged in an orderly pattern

• Chemists use the simplest formula to represent an ionic compound

• Example: sodium chloride could be NaCl, Na2Cl2, Na3Cl3, etc., so the simplest formula, NaCl is used

• Formula unit is the lowest whole-number ratio of ions in an ionic compound

Characteristics of Ionic and Molecular Compounds

Characteristic Ionic Compound Molecular Compound

Representative unit

Formula unit Molecule

Type of elements Metallic combined with nonmetallic

Nonmetallic

Physical state Solid Solid, liquid or gas

Melting point High: above 300C Low: below 300C

IV. THE OCTET RULE: atoms form bonds to achieve a noble

gas configuration

Notes

VII. ELECTRON DOT SYMBOLS or Lewis electron

dot symbols

Notes

A. Each atom would then have 8 electrons in the outermost

energy level.

B. Hydrogen is an exception, it would have 2 electrons.

Notes

V. Bond formation is predicted by two

observations:

Notes

A. Noble gases are unreactive and form very few

compounds.

Notes

B. A filled outermost orbital is very stable

Notes

1. Helium has 2 valence electrons.

2. All the other noble gases have 8 valence electrons.

Notes

IV. Most stable compounds containing representative elements contain ions that have acquired an electron

configuration that is the same as the electron configuration of

a noble gas.

Notes

A. Developed by G. N. Lewis in 1916

Notes

B. The element symbolrepresents the nucleus and the core electrons

C. Dots represent the valence electrons - each dot is an electron

Notes

The number of valence electrons is the group number for all group A elements. (An element in group 2A has 2

valence electrons)

This statement is not on your paper.

Notes

D. Imagine a box around the symbol

FFluorine is in family 7A. It has 7 valence electrons

Notes

1. Place one dot on each side of the symbol until there is

one dot on each side.

F.

..

.

Notes

2. Then pair the dots on each side

.... .

.

.F

Notes

E. Exact placement of dots depends upon how the symbol is being used.

Notes

F. Try to group the dots so that the electron pairs are on opposite sides of the symbol to illustrate the way electrons

spread out around the nucleus.

Se.. . ...

Se = family 6A = 6 valence electrons

Notes

G. Do electron dot symbols for oxygen, chlorine, carbon and

neon:

O Cl C Ne

Notes

CO.

.

.. ...

. ..Notes

ClNe.. .

...

..

... ..

..Notes

VIII. Writing electron dot structures of compounds:

answer these questions:

Notes

A. How many kinds of atom are in the compound?

Notes

1. Determine this from the compound formula.

2. MgCl2 = 2 types of atoms –

Magnesium and Chlorine

Notes

B. How many valence electrons are available?

Notes

1. Use the position of each atom in the periodic table to

determine the number of valence electrons.

Notes

2. Group A families: use the family/group number.

3. Group B families: we are not doing these.

Notes

4. Mg is in group 2A = 2 valence electrons

Notes

5. Cl is group 7A = 7 valence electrons but there are two Cl, so multiply the valence number by the number of atoms in the formula 2 X 7 = 14 valence

electrons

Notes

6. Add up the total electrons for the compound

2 (from Mg) + 14 (from Cl) = 16valence electrons

Notes

7. Use the total number of valence electrons in the electron dot structure.

Notes

C. What is the skeleton structure?

Notes

1. The skeleton structure shows which atoms are bonded to

each other.

Notes

2. The single atom in the formula is usually the central atom.

3. Carbon always goes in the middle. (Even if there are more than one

carbon in the formula.)

Notes

Cl Mg Cl

D. Where do the dots go in the structure?

1. Place the dots around the element symbols so that each

symbol has 8 electrons.

SKELETON = Cl Mg Cladd in the 14 electrons

Cl ClMg .. ......

....

. ...

2. Hydrogen is an exception it will have 2 electrons.

Practice: Do each Lewis dot diagram

MOLECULE: H2ONUMBER OF EACH TYPE

OF ATOM

H = 2

O = 1

NUMBER OF VALENCE ELECTRONS FOR EACH

ATOM

H = 1

O = 6

TOTAL VALENCE ELECTRONS

H: 21=2+ O: 16=6

8

SKELETON STRUCTURE

H O H

LEWIS DOT STRUCTURE

H O H.. .. ....

MOLECULE: CaCL2NUMBER OF EACH TYPE

OF ATOM

Ca = 1

Cl = 2


ATOM

Ca = 2

Cl = 7


Ca: 12= 2+ Cl: 27=14

16

SKELETON STRUCTURE

Cl Ca Cl

LEWIS DOT STRUCTURE

Cl Ca Cl.. ....

...... ....

MOLECULE: H2SNUMBER OF EACH TYPE

OF ATOM

H = 2

S = 1


ATOM

H = 1

S = 6


H: 21=2+ S: 16=6

8

SKELETON STRUCTURE

H S H

LEWIS DOT STRUCTURE

H S H.. .. ....

MOLECULE: SiF4NUMBER OF EACH TYPE

OF ATOM

Si = 1

F = 4


ATOM

Si = 4

F = 7


Si: 14= 4+ F: 47=28

32

SKELETON STRUCTURE

FF Si F

F

LEWIS DOT STRUCTURE

FF Si F

F

. ......

. . .....

... ... ..

......

.. . .

MOLECULE: CCl4NUMBER OF EACH TYPE

OF ATOM

C = 1

Cl = 4


ATOM

C = 4

Cl = 7


C: 14= 4+ Cl: 47=28

32

SKELETON STRUCTURE

ClCl C Cl

Cl

LEWIS DOT STRUCTURE

ClCl C Cl

Cl

. ......

. . . ....

... ... ..

......

.. . .

XI. LIMITATIONS OF THE OCTET RULE:

A. The octet rule is a rule of thumb and compounds do

not obey rules.

B. Some compounds cannot be explained by the octet rule.

C. Nitric oxide (NO) has an odd number of valence

electrons (5 + 6 = 11) and does not satisfy the octet rule

and is not an ion.

D. Atoms without octets

1. Some stable compounds have a central atom without

an octet.

2. Boron trifluoride BF3 is stable and has been shown

experimentally to have only 6 electrons around the Boron

atom.

E. Some stable compounds have a central atom with more

than an octet of electrons.

1. Compounds synthesized from noble gases are an

example.

2. More than an octet is possible because of the larger

size of some central atoms.

3. Sometimes the bond includes electrons from the core electrons in addition to

the valence electrons.

F. Equivalent electron dot structures

1. Some compounds and polyatomic ions can have more

than one structure.

2. These are resonance structures.

3. Writing each of the possible structures and putting double headed arrows () between

the structures shows the resonance.

Molecular orbitals

• Bonding orbitals: molecular orbital with an energy lower than that of the atomic orbitals from which it is formed (all electrons seek lowest energy level)

• Sigma bond (): a bond formed when two atomic orbitals combine to form a molecular orbital that is symmetrical along the axis connecting the two atomic nuclei

• Pi bond (): a bond in which the bonding electrons are most likely to be found in the sausage-shaped regions above and below the nuclei of the bonded atoms

VSEPR theory

• Valence-shell electron-pair repulsion theory: because electron pairs repel, molecules adjust their shapes so that valence-electron pairs are as far apart as possible.

• Hybridization: several orbitals mix to form the same total number of equivalent hybrid orbitals

• ex: methane CH4 : One 2s orbital and three 2p orbitals of a carbon atom mix to form four sp3 hybrid orbitals. These four sp3 hybrid orbitals then overlap the 1s orbital of each hydrogen atom.

Face centered

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COMPOUNDS AND BONDING. Objectives 1 Identify characteristics in atoms involved in chemical bonding Compare the physical and chemical properties of ionic