unit 7: bonding what previous knowledge will help us understand bonding? how can we describe energy...

UNIT 7: BONDING What previous knowledge will help us understand bonding?

How can we describe energy involved in a chemical bond?

How can we explain and draw ionic bonds?

How can we explain and draw covalent bonds?

What are metallic bonds and why are they good conductors?

What is the difference between bond polarity and molecule polarity?

How are molecules geometrically arranged?

How does the VESPR theory influence the geometry of molecules?

What are the different forces that hold molecules together?

What are sigma and pi bonds?

AIM: Prior Knowledge

ENDOTHERMIC: chemical reaction that absorbs heat, producing products with more PE than the reactants

EXOTHERMIC: chemical reaction that produces heat with less PE than the reactants

POTENTIAL ENERGY: stored energy based on position or composition

Why do atoms become ions? To become stable

How do atoms become ions? Gain or loss of electrons

How do metals form ions? MELPS

How do nonmetals form ions? Opposite of MELPS

BONDING

Chemical bonds provide the glue that holds all compounds together

The electron structure of atoms helps explain many aspects of chemical bonding

ENERGY AND CHEMICAL BONDS

Chemical bonds are the forces that holds atoms together.

Energy is required to overcome these attractive forces and separate the atoms in a compound

Breaking a chemical bond is an endothermic process

ENDOTHERMIC: chemical reaction that absorbs energy producing products with more potential energy than the reactants

Ex) N2 + ENERGY N + N

ENERGY AND CHEMICAL BONDS

Formation of a bond is an exothermic process.

EXOTHERMIC: chemical reaction that releases energy producing products with less potential energy than the reactants

Ex) N + N N2 + ENERGY

Bonding and stability When bonds are formed their products are more

stable

The compounds have smaller amount of potential energy

POTENTIAL ENERGY: stored energy based on position or composition

The bonded elements of a compound are more stable than the individual atom or ions because the atoms have filled their valence electron shell

Types of bonds

There are three types of bonds:1. Ionic2. Covalent3. Metallic

They differ in the types of elements involved.

Also how the valence electrons are handled

IONS

Atoms become ions so that they can become stable. Atoms become ions by either gaining or losing electron

IONS

Metals form ions by losing electrons and becoming a positive ion with a smaller atomic radius. Positive ions are called cations. Nonmetals form ions by gaining electrons to become a negative ion with a larger atomic radius. Negative ions are also called anions

IONIC BONDING

The bond that involves the transfer of one or more electrons from a metal atom to a nonmetal atom to form ions. The positive ion and negative ion they attract each other and create a bond.

IONIC BONDINGThere is a large electronegativity difference (E.D.) between a nonmetal and a metal. The nonmetal rips away the valence electron from the metal atom. Nonmetal becomes a negative ion or anion and the metal becomes a positive ion or cation

IONIC BONDINGExamples:

1 - EX) when Na and Cl atoms come together

Na loses electrons and becomes Na+1

Cl gains electrons and becomes Cl-1

They attract to form ionic compound NaCl(aka: salts or electrolytes)

Ionic bonds have the highest polarity (most unequal type of bonding) and the most ionic character

CLUE FOR RECOGNIZING IONIC

BONDS1. Metal and nonmetal 2. Electronegativity difference in greater than 1.7 (approx)

Properties of Ionic Bonds

1. High melting and boiling point.

2. Good electrical conductor as a liquid or when dissolved in water

3. Not good electrical conductor as solid

4. Hard substances

DRAWING LEWIS DOT STRUCTURES FOR IONIC

BONDS Write the metal symbol with no dots in

brackets

Place the charge at the top right of the bracket

Write the nonmetal symbol with 8 dots around it (except H!)

Draw brackets around the symbol and place the charge of the ion at the top right of the bracket

DRAWING LEWIS DOT STRUCTURES FOR IONIC

BONDS Example: Draw the Lew dot structure of

the following elements Na and F

QUESTIONS:

What is an ionic bond?

What are the clues for recognizing an ionic bond?

What happened to electrons in an ionic bond?

We will be using dot structures in order to draw ionic bonds. Know the information from class about the roles of metals and nonmetals in forming ionic bonds; predict what the dot structure of a metal will look like.

Predict what the dot structure of a nonmetal will look like

PRACTICE

BondingCOVALENT/METALLIC

ionic Covalentmolecular metallic

Bonding

m,nm all nm all m

+,-

transfer share M-SOME

Polar covalent

Non polar covalent

AIM - What is it about the structure of noble gases that

leads to their stability? Noble gases (group 18) are stable and undergo few

chemical reactions – lack reactivity

(Argon and Xenon combine with fluorine – rare)

Why? – All have 8 valence electrons except He with 2

Octet – configuration of 8 valence electrons represents max # of valence electrons an atom can have (except H and He – max 2)

Octet Rule – states atoms generally react by gaining, losing, or sharing electrons in order to achieve a complete octet of 8 valence electrons – noble gas configuration

Aim- How can we understand covalent (molecular) bonds?

Covalent bond – formed when two nuclei share electrons to achieve a stable arrangement of electrons – between all nonmetals

Diatomic – covalent bond formed between two nonmetals of the same element:

Br2, I2, N2, Cl2 H2, O2, F2,

PROPERTIES

Exist in gas, liquid, or solid state

Good insulators

Poor conductors

Low melting points

Many are soft substances

CONSTRUCTING LEWIS DOT STRUCTURE FOR SINGLE BINARY COVALENT MOLECULAR COMPOUNDS

1. Determine valence electrons in total (add them up for each element in the compound

2. Divide by 2 to determine the number of pairs of electrons in total for the compound

3. Place first pair between the two elements (use a dash – to represent the shared pair)

4. Place remaining pairs around each elements making sure not to violate the octet rule (Remember H can have a max of 2 electrons)

H2 Cl2

HCl HBr

CONSTRUCTING LEWIS DOT STRUCTURES FOR SINGLE TERNANRY COVALENT

MOLECULAR COMPOUNDS (more than two elements involved)



3. Determine the most electronegative element and place it in the middle

4. Place the other elements around it

5. Start placing pairs (as dash lines) between the central atom and the terminal atoms

6. Place remaining around each elements making sure not to violate the octet rule (Remember H can have a max of 2 electrons)

NH3 CH4 CCl4

CONSTRUCTING LEWIS DOT STRUCTURES FOR MULTIPLE COVALENT MOLECULAR COMPOUNDS (more than two elements involved)



3. Determine the most electronegative element and place it in the middle

4. Place the other elements around it

5. Start placing pairs (as dash lines) between the central atom and the terminal atoms

6. Place remaining around each elements making sure not to violate the octet rule (Remember H can have a max of 2 electrons)

7. If octet rule is not yet reached you can make additional pairs of electrons into double or triple bonds until octet rule is obeyed by all elements

*can only be done with CNOPS

CO2 O2 N2

Coordinate Covalent Bond

Sigma () Bonds

Sigma bonds are characterized by Head-to-head overlap. Cylindrical symmetry of electron density

about the internuclear axis.

Pi () Bonds Pi bonds are

characterized by Side-to-side

overlap. Electron density

above and below the internuclear axis.

© 2009, Prentice-Hall, Inc.

Sigma and Pi Bonds

Single bonds are always bonds, because overlap is greater, resulting in a stronger bond and more energy

Double bonds – contain one sigma and one pi bond

Triple bonds – contain one sigma and two pi bonds

Triple bonds are the shortest and strongest type of bond then double and single bonds are the longest and weakest type of bond

How can we use the 1.7 rule to predict bond polarity?

1.7 rule is applied primarily to binary compounds.

Determine the electronegativities of all atoms in the bond.

Take the difference between the bonded atoms

If the difference is:

>1.7 then ionic

< 1.7 but not “0”, then polar covalent

=0 non polar covalent

Nonpolar covalent bond

Polar covalent bond

Ionic bond

Covalent bonds

MOLECULE POLARITY: Nonpolar/polar shapes

SNAP PAD SNAP : Symmetrical Nonpolar Asymmetrical Polar

PAD: Polar Asymmetrical Dipole

OPEN: Odd Polar Even Nonpolar

DIPOLE MOMENT

dipole moment--separation of the charge in a molecule; product of the size of the charge and the distance of separation• align themselves with an electric field (figure b at right)• align with each other as well in the absence of an electric field

• water—2 lone pairs establish a strong negative pole• ammonia—has a lone pair which establishes a neg. pole• note that the direction of the “arrow” indicating the dipole

moment always points to the negative pole with the cross hatch on the arrow (looks sort of like we’re trying to make a + sign) is at the positive pole.

Compound

Polar or Nonpolar

NH3

H20

How can a molecule be both polar and non polar?

CX4 tetrahedral shape. They are non polar.

The individual ligands or bonds are polar

Conclusion – nonpolar/polar.

What Determines the Shape of a Molecule?

Simply put, electron pairs, whether they be bonding or nonbonding, repel each other.

By assuming the electron pairs are placed as far as possible from each other, we can predict the shape of the molecule.


Valence Shell Electron Pair Repulsion Theory (VSEPR)

“The best arrangement of a given number of electron domains is the one that minimizes the repulsions among them.”


THE VSPER MODEL AND MOLECULAR

SHAPE molecular geometry--the arrangement in space of the

atoms bonded to a central atom not necessarily the same as the structural pair geometry lone pairs have a different repulsion since they are experiencing an attraction or “pull” from only one nucleus as opposed to two nuclei. They also take up more space around an atom as you can see on the left.

Each lone pair or bond pair repels all other lone pairs and bond pairs--they try to avoid each other making as wide an angle as possible.

- works well for elements of the s and p-blocks

- VSEPR does not apply to transition element compounds (exceptions)

MOLECULAR SHAPES- Linear

- Bent (angular)

- Pyramidal

- Tetrahedral

Intermolecular Forces

The attractions between molecules are not nearly as strong as the intramolecular attractions that hold compounds together.



They are, however, strong enough to control physical properties such as boiling and melting points, vapor pressures, and viscosities.



These intermolecular forces as a group are referred to as van der Waals forces.


Ion-Dipole Interactions Ion-dipole interactions (a fourth type of

force), are important in solutions of ions.

The strength of these forces are what make it possible for ionic substances to dissolve in polar solvents.


van der Waals Forces

Dipole-dipole interactions

Hydrogen bonding

London dispersion forces


Dipole-Dipole Interactions

Molecules that have permanent dipoles are attracted to each other. The positive end of one

is attracted to the negative end of the other and vice-versa.

These forces are only important when the molecules are close to each other.


Dipole-Dipole Interactions

The more polar the molecule, the higher is its boiling point.


London Dispersion Forces

While the electrons in the 1s orbital of helium would repel each other (and, therefore, tend to stay far away from each other), it does happen that they occasionally wind up on the same side of the atom.



At that instant, then, the helium atom is polar, with an excess of electrons on the left side and a shortage on the right side.



Another helium nearby, then, would have a dipole induced in it, as the electrons on the left side of helium atom 2 repel the electrons in the cloud on helium atom 1.



London dispersion forces, or dispersion forces, are attractions between an instantaneous dipole and an induced dipole.



These forces are present in all molecules, whether they are polar or nonpolar.

The tendency of an electron cloud to distort in this way is called polarizability.


Factors Affecting London Forces

The shape of the molecule affects the strength of dispersion forces: long, skinny molecules (like n-pentane tend to have stronger dispersion forces than short, fat ones (like neopentane).

This is due to the increased surface area in n-pentane.


Factors Affecting London Forces

The strength of dispersion forces tends to increase with increased molecular weight.

Larger atoms have larger electron clouds which are easier to polarize.


Which Will Have a Greater Effect?Dipole-Dipole Interactions or Dispersion Forces

If two molecules are of comparable size and shape, dipole-dipole interactions will likely the dominating force.

If one molecule is much larger than another, dispersion forces will likely determine its physical properties.


How Do We Explain This?

The nonpolar series (SnH4 to CH4) follow the expected trend.

The polar series follows the trend from H2Te through H2S, but water is quite an anomaly.


Hydrogen Bonding

The dipole-dipole interactions experienced when H is bonded to N, O, or F are unusually strong.

We call these interactions hydrogen bonds.


Hydrogen Bonding Hydrogen bonding

arises in part from the high electronegativity of nitrogen, oxygen, and fluorine.


Also, when hydrogen is bonded to one of those very electronegative elements, the hydrogen nucleus is exposed.

Summarizing Intermolecular Forces


Intermolecular Forces Affect Many Physical Properties

The strength of the attractions between particles can greatly affect the properties of a substance or solution.


Single Bonds

Single bonds are always bonds, because overlap is greater, resulting in a stronger bond and more energy


How can we understand and recognize metallic bonding?

Contains positively charged metals

Metals are arranged in a crystalline lattice structure immersed in a:sea of mobile electrons

Electrons are delocalized. This means no one atom owns any electrons they belong to the whole crystal.

M-SOME

Metals are good conductors because of mobile ions

Sea of Mobile Electrons

ionic Covalentmolecular

metallic

Bonding

m,nm all nm all m

+,- Polar Covalent

Non Polar Covalent

unit 7: bonding what previous knowledge will help us understand bonding? how can we describe energy...

Documents

stored energy

positive ions

negative ions

reactants potential

chemical reaction

nonmetals form ions

ionic bonds

chemical bondschemical