Unit 2: Chemical Bonding Chemistry2202. Outline Bohr diagrams Lewis Diagrams Types of Bonding ...

transcript

Unit 2: Chemical Bonding

Chemistry2202

Outline

Bohr diagrams Lewis Diagrams Types of Bonding

Ionic bonding Covalent bonding (Molecular) Metallic bonding Network covalent bonding

Types of Bonding (cont’d) London Dispersion forces Dipole-Dipole forces Hydrogen Bonding

VSEPR Theory (Shapes) Physical Properties

Bohr Diagrams (Review)

How do we draw a Bohr Diagram for - The F atom? - The F ion?

Draw Bohr diagrams for the atom and the ion for the following:

Al S C l Be

Lewis Diagrams

LD provide a method for keeping track of electrons in atoms, ions, or molecules

Also called Electron Dot diagrams the nucleus (P& N) and filled energy

levels are represented by the element symbol

Lewis Diagrams

dots are placed around the element symbol to represent valence electrons

Lewis Diagrams

eg. Lewis Diagram for F

• ••••

lone pair

bonding electron

Lewis Diagrams

lone pair – a pair of electrons not available for bonding

bonding electron – a single electron that may be shared with another atom

Lewis Diagrams

eg. Lewis Diagram for C

• ••

Lewis Diagrams

eg. Lewis Diagram for P

• •••

Lewis Diagrams

eg. Lewis Diagram for Na

Lewis DiagramsFor each atom draw the Lewis diagram

and state the number of lone pairs and number of bonding electrons

Li Be Al Si

Mg N B O

Lewis Diagrams for Compounds draw the LD for each atom in the

compound The atom with the most bonding

electrons is the central atom Connect the other atoms using single

bonds (1 pair of shared electrons) In some cases there may be double

bonds or triple bonds

Lewis Diagrams for Compounds eg. Draw the LD for:

NH3 SiCl4 N2H4 HCN

SI2 CO2 N2H2 CH2O

POI CH3OH

N2 H2 O2

Lewis Diagrams for CompoundsA structural formula shows how the atoms are connected in a molecule.

To draw a structural formula: replace the bonded pairs of electrons

with short lines omit the lone pairs of electrons

Why is propane (C3H8) a gas at STP while kerosene (C10H22) a liquid?

Why is graphite soft enough to write with while diamond is the hardest substance known even though both substances are made of pure carbon?

Why can you tell if it is ‘real gold’ or just ‘fool’s gold’ (pyrite) by hitting it with a rock?

‘As Slow As Cold Molasses’

‘All Because of Bonding’

‘liquids’ @ -30 ºC

Bonding

Bonding between atoms, ions and molecules determines the physical and chemical properties of substances.

Bonding can be divided into two categories:

- Intramolecular forces

- Intermolecular forces

BondingIntramolecular forces are forces of attraction between atoms or ions.

Intramolecular forces include:

1. ionic bonding

2. covalent bonding

3. metallic bonding

4. network covalent bonding

BondingIntermolecular forces are forces of attraction between molecules.

Intermolecular forces include:

5. London Dispersion Forces

6. Dipole-Dipole forces

7. Hydrogen Bonding

Ionic and Covalent Bonding

ThoughtLab p. 161

Identify #’s 1 - 6

Ionic Bonding Occurs between cations and anions –

usually metals and non-metals. An ionic bond is the force of attraction

between positive and negative ions. Properties:

conduct electricity as liquids and in solution hard crystalline solids high melting points and boiling points brittle

In an ionic crystal the ions pack tightly together.

The repeating 3-D distribution of cations and anions is called an ionic crystal lattice.

Ionic Bonding

Each anion can be attracted to six or more cations at once.

The same is true for the individual cations.

Ionic Bonding

Covalent Bonding

Occurs between non-metals in molecular compounds.

Atoms share bonding electrons to become more stable (noble gas structure).

A covalent bond is a simultaneous attraction by two atoms for a common pair of valence electrons.

Covalent Bonding

Molecular compounds have low melting and boiling points.

Exist as distinct molecules.

Covalent Bonding

Molecular compounds do not conduct electric current in any form

Property Ionic Molecular

Type of elements

Metals and nonmetals

Non-Metals

Force of Attraction

Positive ions attract negative ions

Atoms attract a shared electron

Electron movement

Electrons move from the metal to

the nonmetal

Electrons are shared

between atoms

State at room temperature

Always solids Solids, liquids, or gas

Property Ionic Molecular

Solubility Soluble or low solubility

Soluble or insoluble

Conductivity in solid state

None None

Conductivity in liquid state

Conducts None

Conductivity in solution

Conducts None

Metallic Bonding (p. 171) metals tend to lose valence electrons. valence electrons are loosely held and

frequently lost from metal atoms. This results in metal ions surrounded by

freely moving valence electrons. metallic bonding is the force of attraction

between the positive metal ions and the mobile or delocalised valence electrons

Metallic Bonding

Metallic Bonding This theory of metallic bonding is called

the ‘Sea of Electrons’ Model or ‘Free Electron’ Model

Metallic Bonding This theory accounts for properties of metals

1. electrical conductivity

- electric current is the flow of electrons

- metals are the only solids in which electrons are free to move

2. solids- Attractive forces between positive cations

and negative electrons are very strong

Metallic Bonding3. malleability and ductility- metals can be hammered into thin

sheets(malleable) or drawn into thin wires(ductile).

- metallic bonding is non-directional such that layers of metal atoms slide past each other under pressure.

Network Covalent Bonding (p. 199)

occurs in 3 compounds (memorize these) diamond – Cn

carborundum – SiC quartz – SiO2

large molecules with covalent bonding in 3-d

each atom is held in place in 3-d by a network of other atoms

Network Covalent bonding Properties:

the highest melting and boiling points the hardest substances brittle do not conduct electric current in any

Strongest

1. Network Covalent (Cn ,SiO2 , SiC)

2. Ionic bonding(metal & nonmetal)

3. Metallic bonding (metals)

4. Molecular (nonmetals)Weakest

P decreases

Valence Shell Electron Pair Repulsion theory (VSEPR)

The shape of molecules is determined by the arrangement of valence electron pairs around the atoms in a compound.

There are 5 shapes that can be determined by the # of bonds and # of lone pairs on the central atom.

1. Tetrahedral (4 bonds; 0 lone pairs)

2. Pyramidal (3 bonds; 1 lone pair)

3. V-shaped (2 bonds; 2 lone pairs)

4. Trigonal Planar (3 bonds; 0 lone pairs)

5. Linear (2 bonds; 0 lone pairs)

For each molecule below draw the Lewis diagram and the shape diagram.

HOCl H2Se H2O2

NBr3 C2F4 C2H6

CHCl3 CH3OH

I2 SiH4

HSiHO C2H2

HCN PBr3

Electronegativity (EN - p. 174) EN is a measure of the attraction that an atom has for shared electrons.

A higher EN means a stronger attraction or electrostatic pull on valence electrons

EN values increase as you move:- from left to right in a period- up in a group or family

Increases

Electronegativity & Covalent Bonds1. polar covalent bond

- a bond between atoms with different EN- the shared electron pair is attracted

more strongly to the atom with the higher EN

δ−δ+

Electronegativity (p. 174) polar covalent bond

a covalent bond between atoms with different EN

the shared electron pair is attracted more strongly to the atom with the higher EN

nonpolar covalent bondbond dipole

Complete: #’s 7 – 9 on p.178

Weakest

Covalent (nonmetals)

→ London Dispersion

(all molecules)

→ Dipole-Dipole

(polar molecules)

→ H bonding

(H-N, H-O, H-F)

p. 226 #13

Omit parts g), j) – o), q), u), & v)

- Answers on p. 815 for #13

- Incorrect answers

c), d), & s)

Electronegativity and Ionic Bonds Because the EN of metals is so

low, metals lose electrons to form

cations Nonmetals gain electrons to form

anions because the EN of nonmetals is relatively high

Electronegativity and Ionic Bonds When ions form, the resulting

electrostatic force is an ionic bond

Electronegativity and Covalent Bonds Atoms in covalent compounds can

either have: the same EN

eg. Cl2 , PH3, NCl3 different EN

eg. HCl

Electronegativity and Covalent Bonds

Atoms that have the same EN attract the shared valence electrons to the same extent.

Covalent bonds resulting from equal sharing of the bonding electron pairs are called Nonpolar Covalent Bonds

Atoms that have different EN attract the shared pair of valence electrons at different strengths

The atom with the higher EN exerts a stronger attraction on the shared electron pair

eg. H2O

Since the oxygen atom has a higher EN the bonding electrons will be pulled closer to the oxygen atom

This results in slight positive and negative charges within the bond.

These charges are referred to as “partial charges” and are denoted

with the Greek letter delta (δ).

The region around the oxygen atom will be slightly negative, and around the hydrogens will be slightly positive

The symbol, δ+ represents a partial positive charge (less than +1) and δ− represents a partial negative charge (less than −1).

Since the bond is polarized into a positive area and a negative area the bond has a “bond dipole”.

Electronegativity and Covalent Bonds The arrow points to

the atom with the higher EN.

Electronegativity and Covalent Bonds Covalent bonds resulting from

unequal (electronegativities) sharing of bonding electron pairs are called Polar Covalent Bonds

Electronegativity Homework

#’s 7, 8, & 9 - p. 178 #’s 1, 2, & 3 - p. 180

Bond Energy (pp. 179-180)

1. Describe the forces of attraction and repulsion present in all bonds.

2. What is bond length?3. Define bond energy.4. Which type of bond has the most energy?5. How can bond energy be used to predict

whether a reaction is endothermic or exothermic?

Test Outline Bohr Diagrams (atoms & ions) Lewis Diagrams (Electron Dot) Ion Formation Ionic Bonding, Structures & Properties Covalent Bonding, Structures & Properties

Test Outline Metallic Bonding Theory& Properties Network Covalent Bonding & Properties Electronegativity Bond Dipoles & Polar Molecules VSEPR Theory LD, DD, & H-bonding Predicting properties (bp, mp, etc.)

Molecular Dipoles

The vector sum of all the bond dipoles in a molecule is a Molecular Dipole

A Polar Molecule has a molecular dipole that points toward the more electronegative end of the molecule.

eg. H2O

Molecular Dipoles

NonPolar Molecules DO NOT have molecular dipoles. This occurs when:

- the bond dipoles cancel

- there are no bond dipoles

eg. CO2 PH3

Molecular Dipoles

To determine whether a molecule is polar:

- draw the Lewis diagram and the shape diagram

- draw the bond dipoles and determine whether they cancel

Intermolecular Forces

Strongest bonds; Highest mp and bp

1. Network Covalent (Cn SiO2 SiC)

2. Ionic bonding(metal & nonmetal)

3. Metallic bonding (metals)

4. Molecular (nonmetals)Weakest bonds; Lowest mp and bp- Intermolecular forces present

To compare mp and bp in covalent compounds you must use:

- London Dispersion forces (p. 204)

(all molecules)

- Dipole-Dipole forces (pp. 202, 203)

(polar molecules)

- Hydrogen Bonding (pp. 205, 206)

(H bonded to N, O, or F)

Intermolecular Forces (p. 202)

Intermolecular Forces Covalent compounds have low mp and

bp because forces between molecules in covalent compounds are very weak.

Intermolecular forces were studied extensively by the Dutch physicist Johannes van der Waals

In his honor, two types of intermolecular force are called Van der Waals forces.

Intermolecular forces can be used to account for the physical properties of covalent compounds.

Van der

• LD forces exist in ALL molecular elements & compounds.

•The positive charges in one molecule attract the negative charges in a second molecule.

• The temporary dipoles caused by electron movement in one molecule attract the temporary dipoles of another molecule.

The strength of these forces depends on:a)the number of electrons

more electrons produce stronger LD forces that result in higher mp and bpeg. CH4 is a gas at room temperature.

C8H18 is a liquid at room temperature.C25H52 is a solid at room temperature.Account for the difference.

Two molecules that have the same number of electrons are isoelectronic

eg. C2H6 and CH3F

b) shape of the molecule molecules that “fit together” better will experience stronger LD forces

eg. Cl2 vaporizes at -35 ºC while C4H10 vaporizes at -1 ºC. Use bonding to account for the difference.

2. Dipole-dipole Forces

- occur between polar molecules

- the δ+ end of one polar molecule is attracted to the δ- end of another polar molecule (& vice-versa)

eg. Which has the higher boiling point CH3F or C2H6 ?

p. 202

3. Hydrogen Bonds

- a special type of dipole-dipole force (about 10 times stronger) - only occurs between molecules that contain a H atom which is directly bonded to F, O, or N ie. the molecule contains at least one H-F, H-O, or H-N covalent bond.

3. Hydrogen Bonds

-the hydrogen bond occurs between the H atom of one molecule and the N, O, or F of a second molecule.

eg. Arrange these from highest to lowest boiling point

C3H8 C2H5OH C2H5F

p. 206

NOTE: To compare covalent compounds you must use:

- London Dispersion forces

(all molecules)

- Dipole-Dipole forces

(polar molecules)

- Hydrogen Bonding

Alchem worksheet pp. G32, 33

p. 210

Intermolecular Forces1. Use intermolecular forces to explain the following:

a) Ar boils at -186 °C and F2 boils at -188 °C .

b) Kr boils at -152 °C and HBr boils at -67 °C.

c) Cl2 boils at -35 °C and C2H5Cl boils at 13 °C .

2. Examine the graph on p. 210:

a) Account for the increase in boiling point for the hydrogen compounds of the Group IV elements.

b) Why is the trend different for the hydrogen compounds of the Group V, VI, and VII elements?

c) Why are the boiling points of the Group IVA compounds consistently lower than the others.

3.Which substance in each pair has the higher boiling point. Justify your answers.

(a) SiC or KCl

(b) RbBr or C6H12O6

(c) C3H8 or C2H5OH

(d) C4H10 or C2H5Cl

2. Examine the graph on p. 210:

a) Account for the increase in boiling point for the hydrogen compounds of the Group IV elements.

b) Why is the trend different for the hydrogen compounds of the Group V, VI, and VII elements?

c) Why are the boiling points of the Group IVA compounds consistently lower than the other compounds.

Dipole-Dipole Forces In the liquid

state, polar molecules (dipoles) orient themselves so that oppositely charged ends of the molecules are near to one another.

Summary

The types of bonding/forces ranked from strongest to weakest are:

Strongest - Network Covalent

- Ionic

- Metallic

Weakest - Covalent

NOTE: To compare covalent compounds you must use:

- London Dispersion forces

(all molecules)

- Dipole-Dipole forces

(polar molecules)

- Hydrogen Bonding

p. 226 #’s 13 & 14

Dipole-Dipole Forces

The electrostatic attractions between these oppositely charged ends of the polar molecules are called dipole-dipole forces.

Results of dipole-dipole attractions: polar molecules will tend to

attract one another more at room temperature than similarly sized non-polar molecules

energy needed to separate polar molecules is therefore higher than for non-polar molecules of similar molar mass

Results of dipole-dipole attractions: The melting points and boiling

points of substances made of polar molecules are higher than for substances made of non-polar molecules.

Ion-Dipole Forces

An ion-dipole force is the force of attraction between an ion and a polar molecule (a dipole).

Ion-Dipole Forces NaCl dissolves in water because the

attractions between the Na+ and Cl- ions and the partial charges on the H2O molecules are strong enough to overcome the forces that bind the ions together.

Induced Intermolecular Forces

Induced Intermolecular Forces Induction of electric charge occurs

when a charge on one object causes a change in the distribution of charge on a nearby object. (for example, the balloon)

Induced Intermolecular ForcesThere are two types of charge -

induced dipole forces:

1. An ion-induced dipole force results when an ion in close proximity to a non-polar molecule distorts the electron density of the non-polar molecule

Induced Intermolecular Forces The molecule then becomes

momentarily polarized, and the two species are attracted to each other. (ie. hemoglobin)

2. In a dipole-induced dipole force the charge on a polar molecule is responsible for inducing the charge on the non-polar molecule.

Dispersion (London) Forces Bond vibrations, which are part of

the normal condition of a non-polar molecule, cause momentary, uneven distribution of charge;

a non-polar becomes slightly polar for an instant, and continues to do so in a random but constant basis.

Dispersion (London) Forces At the instant that one non-polar

molecule is in a slightly polar condition, it is capable of inducing a dipole in a nearby molecule

This force of attraction is called a dispersion force.

Dispersion (London) Forces Two factors affecting the

magnitude of dispersion forces are:

1. The number of electrons in the molecule:

Vibrations within larger molecules that have more electrons than smaller molecules can easily cause an uneven distribution of charge.

Dispersion (London) ForcesThe dispersion forces between these

larger molecules are thus stronger, which has the effect of raising the boiling point for larger molecules.

2. The shape of the molecule:

A molecule with a spherical shape has a smaller surface area than a straight chain molecule that has the same number of electrons

Dispersion (London) ForcesTherefore, the substance with

molecules that have a more spherical shape will have weaker dispersion forces and a lower boiling point.

London dispersion forces are responsible for the formation and stabilization of the biological membranes surrounding every living cell.

Hydrogen BondingIn order to form a hydrogen bond, a

hydrogen atom must be bonded to a highly electronegative atom such as oxygen, nitrogen, or fluorine.

Hydrogen BondingThese bonds are very polar, and

since hydrogen has no other electrons, the positive proton, H+, is exposed and can become strongly attracted to the negative end of another dipole nearby

Hydrogen BondingA hydrogen bond is an electrostatic

attraction between the nucleus of a hydrogen atom, bonded to fluorine, oxygen, or nitrogen and the negative end of a dipole nearby.

Hδ+δ+δ+δ+

δ+δ+δ+δ+ δ−δ−δ−δ−

Hydrogen Bonding In biological systems, these polar

bonds are often parts of much larger molecules (ie. N H bonds and C O bonds found in biological molecules)

Hydrogen Bonding in WaterHydrogen bonds between the

hydrogen atoms in one water molecule and the oxygen atom in another account for many unique properties of water.

Hδ+δ+δ+δ+

δ+δ+δ+δ+ δ−δ−δ−δ−

Hydrogen Bonding in Water In liquid water, each

water molecule is hydrogen bonded to at least four other water molecules.

The large number of bonds between water molecules makes the net attractive force quite strong

Hydrogen Bonding in Water the strong attractive forces are

responsible for the relatively high boiling point of water.

The water molecules are farther apart in ice then they are in liquid water making ice less dense than liquid water.

Hydrogen Bonding in Water Hydrogen bonds

force water molecules into the special hexagonal, crystalline structure of ice when the temperature is below 4 degrees celcius.

Assignment # 4

Electronegativity

Electronegativity is a result of the space between the nucleus and the electrons

As the number of protons in the nucleus increases, the attractive force on the electrons increases, pulling them closer to the nucleus

Unit 2: Chemical Bonding Chemistry2202. Outline Bohr diagrams Lewis Diagrams Types of Bonding ...

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