Chemistry Module2.ppt

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UNIT 1: Structure and Properties of Matter

Atomic Models and Properties of

Atoms

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Chapter 4: Chemical Bonding

and Properties of Matter

UNIT 2

The chemical bonding in a substance influences the shape of its molecules, and molecular shape influences the properties of that

substance. One of the properties of iron is its strength, which makes it ideal for use in support structures.

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The strength of iron makes it useful in items such as horseshoes.

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UNIT 2 Section 4.1

4.1 Models of Chemical Bonding

Chapter 4: Chemical Bonding and Properties of Matter

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UNIT 2 Section 4.1

Electronegativity

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Electronegativity is the relative ability of an atom to attract shared electrons in a chemical bond.

What general trends in electronegativity are shown in the periodic table?

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UNIT 2 Section 4.1

Electron Sharing and Electronegativity

Electronegativity difference, ΔEN, between two atoms bonded together can be low, intermediate, or high. The electron density diagrams below show the differences in the bonds.

• when ΔEN is 0: electrons are equally shared

• when ΔEN is 1: electrons are more closely associated with the more electronegative atom

• when ΔEN is high, there is little sharing of electrons

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UNIT 2 Section 4.1

Scientists have categorized types of bonds according to ΔEN.

• ΔEN between 1.7 and 3.3:

mostly ionic

polar covalent

mostly covalent (non-polar)

Three categories of bonds have been set based on ΔEN .

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Chemists use the electron-sea model to describe metallic bonding. The model proposes that the valence electrons of metal atoms move freely among the ions, forming a “sea” of

delocalized electrons that hold the metal ions rigidly in place.

UNIT 2 Section 4.1

Metallic Bonding

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Microscopic analysis shows that the

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UNIT 2 Section 4.1

Properties of Metals

Melting and Boiling Points

• the stronger the bonding forces, the higher the melting and boiling points of pure metals

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Periodic table trends include:

1. For Group 1, melting points decrease as the atomic number increases.

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UNIT 2 Section 4.1

Properties of Metals

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Malleability and Ductility

• Based on the electron-sea model, metals can be shaped because, when struck, the metal ions can slide by one another while the electrons still surround them.

Hardness

•The variation between metals is due to differences in crystal size (smaller ones make harder metals).

Electrical and Thermal Conductivity

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UNIT 2 Section 4.1

Alloys

Alloys are solid mixtures of two or more metals.

• the addition of the second metal, even in a very small

amount, can significantly affect the properties of a substance • in some cases, non-metal atoms, such as carbon, are added

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If atoms of the second metal are similar in size to the first metal, they take the place of those atoms.

If atoms of the second metal are much smaller than atoms of the first metal, they will fit in spaces

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• occurs when ΔEN is between 1.7 and 3.3

• essentially, involves one atom losing one or more electrons and another atom gaining those electron(s)

UNIT 2 Section 4.1

Ionic Bonding

There are different ways to show the transfer of electrons in the formation of ionic compounds.

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Ionic compounds exist as crystal lattice structures with

particular patterns of alternating positive and negative ions. The unit cell is the smallest group of ions that is repeated.

UNIT 2 Section 4.1

Ionic Crystals

NaCl forms a cubic crystal lattice structure.

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Different types of crystal structures can form.

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UNIT 2 Section 4.1

Properties of Ionic Compounds

Melting and Boiling Points

• high due to very strong attractions between ions

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Solubility

• ionic compounds are soluble in water when the attractive forces between the ions and water molecules are stronger than the attractive forces among the ions themselves

When sodium chloride (NaCl) dissolves in water, attractive forces between water

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UNIT 2 Section 4.1

Properties of Ionic Compounds

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Conductivity

• solids do not conduct because ions cannot move

• compounds conduct when dissolved in water and ions can move

Mechanical Properties

• hard and brittle, so will break apart when struck

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Ionic Bond

Between atoms of metals and nonmetals with very different electronegativity

Bond formed by transfer of electrons

Produce charged ions all states. Conductors and have high melting point.

An electronegativity difference of 2 is essential for a compound to be ionic.

Ionic compounds are solids at room temperature and are hard and brittle.

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Size of Na

+

ion is smaller

than Cl

-

ion.

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Why Ionic Compounds exist as Crystals?

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Why Ionic Compounds are hard and

brittle?

Crystals are made of alternate positive and negative ions such that opposite ions lie over one another. When external

force is applied even a slight shift brings like ions close to one another. This make the ionic compounds hard and

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Why Ionic Compounds have high melting

points?

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Solubility of Ionic compounds in Water

When a crystal of an ionic substance is placed in water, the polar water molecules detach the positive and negative ions from the crystal lattice by their electrostatic pull. These ions then get

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UNIT 2 Section 4.1

Covalent Bonding

The length of a covalent bond is determined by

different electrostatic forces.

Forces in covalent bonds:

• both attractive and repulsive forces play a role

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• occurs when ΔEN is less than 1.7

• covalent bonds are classified into two types:

• polar covalent: atoms do not share electrons equally

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Describe the chemical bonding and

structure of NaCl. How do bonding and structure influence the general properties of the substance?

UNIT 2 Section 4.1

Answer on the next slide

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NaCl is composed of a metal atom bonded to a non-metal atom with ΔEN > 1.7. As such, the bond is classified as ionic. It exists as a cubic crystal lattice structure, with an

alternating pattern of chloride ions and sodium ions.

Properties of NaCl include high melting and boiling points; solubility in water; hard and

brittle; a poor conductor as a solid, but it does conduct electricity when dissolved in water.

Section 4.1

UNIT 2

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UNIT 2 Section 4.1

Quantum Mechanics and Bonding

Valence Bond (VB) Theory explains bond formation and molecular shapes based on orbital overlap.

•The region of overlap has a maximum capacity of two electrons, which have opposite spins.

•There should be maximum overlap of orbitals, since the greater the overlap, the stronger and more stable the bond. •Atomic orbital hybridization is used to help explain the shapes of some molecules.

Quantum mechanics is used to explain and describe chemical bonding. It is also used to account for shapes of molecules.

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UNIT 2 Section 4.1

Quantum Mechanics and Bonding

According to MO theory:

• Covalent bond formation involves atomic orbital overlap that results in formation of new orbitals called molecular orbitals.

• Molecular orbitals have shapes and energy levels that are different from those of atomic orbitals.

• The electrons in molecular orbitals are delocalized throughout the orbital.

Molecular Orbital (MO) Theory explains bond

formation and molecular shapes based on the formation of new molecular orbitals.

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UNIT 2 Section 4.1

Explaining Single Bonds

For molecules like hydrogen fluoride:

• the 1s orbital of H overlaps with the half-filled 2p orbital of F

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UNIT 2 Section 4.1

Explaining Single Bonds

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For molecules like methane:

• the VB theory of hybrid orbitals is used to explain molecular shape

• carbon forms four hybrid orbitals (sp3) by combining

three 2p orbitals and a 2s orbital so that four identical bonds can be created

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UNIT 2 Section 4.1

Explaining Double Bonds

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Hybrid orbitals are used to

explain the structure of ethene or molecules like ethene.

• it is planar with ~120º bond angles

• the structure is explained by formation of 3 sp2_{hybrid orbitals}

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UNIT 2 Section 4.1

Explaining Double Bonds

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For bond formation in ethene:

• one sp2 orbital of each carbon overlaps to form a σ bond

between the carbons

• two sp2 orbitals of each carbon overlap with the 1s

orbitals of the hydrogens to form σ bonds

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UNIT 2 Section 4.1

Explaining Triple Bonds

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For molecules like ethyne:

• the linear structure is explained by formation of 2 sp hybrid orbitals for each carbon (a 2s orbital + a 2p orbital)

• sigma bonds form from overlap between sp of each carbon and between sp of carbons and 1s of hydrogens • two pi bonds form from overlap of the two 2p orbitals of

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Allotropes are compounds that consist of the same element but have different physical properties.

UNIT 2 Section 4.1

Allotropes

Allotropes of carbon: A graphite, B diamond,

C buckyballs, D nanotubes

An example is allotropes of carbon, which differ in the pattern of covalent bonds between carbon atoms.

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Network solids are substances that consist of atoms bonded covalently in a continuous two- or three-dimensional array. There is no natural beginning or end to the chains of atoms.

UNIT 2 Section 4.1

Covalent Network Solids

Silicon dioxide, SiO₂, exists as a network solid that is represented as (SiO₂)_n.

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Section 4.1 Review

UNIT 2 Section 4.1

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UNIT 2 Section 4.2

4.2 Shapes, Intermolecular Forces, and

Properties of Molecules

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Molecular compounds form a much greater variety of structures than ionic compounds form.

Understanding the properties of molecules requires an understanding of their three-dimensional shapes.

Different theories and models are used to predict molecular shapes.

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UNIT 2 Section 4.2

Depicting Two-Dimensional Structures of

Molecules with Lewis Structures

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UNIT 2 Section 4.2

Some Exceptions When Drawing

Lewis Structures

The ammonium ion has a co-ordinate covalent bond.

Co-ordinate Covalent Bonds:

• one atom contributes both electrons • bonds behave the same way as other

covalent bonds and therefore are not indicated in Lewis structures

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Expanded Octet (Expanded Valence): •central atom has more than an octet of electrons

•a feature of some Period 3 and higher

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UNIT 2 Section 4.2

Some Exceptions When Drawing

Lewis Structures

In BF3(g), boron has an incomplete octet.

An Incomplete Octet:

• central atom has fewer than an octet of electrons

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Resonance Structures:

• measured bond lengths may not support Lewis structures

• one of two or more Lewis structures that show same relative position of atoms but different positions of

electron pairs

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UNIT 2 Section 4.2

Predicting the Shapes of Molecules

Using VSEPR Theory

The valence-shell electron pair repulsion (VSEPR) theory

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For VSEPR, there are five electron-group arrangements. (Electron groups are represented by bars).

• is a model used to predict molecular shape

• is based on electron groups around a central atom being positioned as far apart as possible (repulsion)

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UNIT 2 Section 4.2

Electron Groups and Molecular Shapes

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UNIT 2 Section 4.2

Electron Groups and Molecular Shapes

If one or more electron groups around a central atom is a lone pair, different strengths of repulsive forces will alter bond angles to differing degrees.

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UNIT 2 Section 4.2

Summarizing Molecular Shapes

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UNIT 2 Section 4.2

Summarizing Molecular Shapes

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UNIT 2 Section 4.2

Guidelines for Using VSEPR Theory to

Predict Molecular Shape

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What is the electron-group arrangement and molecular shape of HCN?

UNIT 2 Section 3.2

Answer on the next slide

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Chapter 3: Atomic Models and Properties of Atoms

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HCN has two bonding groups and no lone pairs.

The electron-group arrangement is linear, and the shape of the

molecule is also linear.

Section 3.2

UNIT 2

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Chapter 3: Atomic Models and Properties of Atoms

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UNIT 2 Section 4.2

The Influence of Molecular Shape

on Polarity

The shape of a molecule affects that molecule’s polarity. • polar bonds have a bond dipole

• bond dipoles are indicated using vectors that point in the direction of higher electron density

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UNIT 2 Section 4.2

Determining Whether a

Molecule is Polar

A molecule with one or more polar bonds is not necessarily a polar molecule. The molecule’s shape must be considered. The polarity as a whole can be determined by adding the vectors.

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UNIT 2 Section 4.2

Molecular Shapes and Polarities

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UNIT 2 Section 4.2

How Intermolecular Forces Affect the

Properties of Solids and Liquids

Intermolecular forces exist between ions and molecules and influence the physical properties of substances.

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Categories of forces: • dipole-dipole

• ion-dipole

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UNIT 2 Section 4.2

Dipole-Dipole

Dipole-dipole forces:

• are forces of attraction between polar molecules, which have a region of partial positive charge and a region of partial negative charge

• are a main reason for melting and boiling point differences between polar and non-polar molecules • include hydrogen bonding, as an example of one type

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UNIT 2 Section 4.2

Ion-Dipole

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Ion-dipole intermolecular forces.

Ion-dipole forces:

• are forces of attraction between partial charges on polar molecules and ions

• depend on the size and charge of the ion and the

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UNIT 2 Section 4.2

Induced Dipoles

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A dipole can be induced in a non-polar molecule.

Dipole-induced dipole forces:

• are forces of attraction between a polar molecule and a non-polar molecule that has an induced (temporary) dipole due to the nearby polar molecule

Ion-induced dipole forces:

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UNIT 2 Section 4.2

Dispersion Forces

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The more linear molecule has a higher boiling point because the dispersion forces are greater.

Dispersion forces:

• are forces of attraction between all molecules, including non-polar molecules

• are due to spontaneous temporary dipoles that form due to the constant motion of electrons in covalent bonds

• depend on the size and shape of the molecules

• _{the larger and more linear the molecule, the greater the}

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Section 4.2 Review

UNIT 2 Section 4.2

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Hydrogen Bonding

A

hydrogen bond

is the attractive interaction of

a hydrogen atom and an electronegative atom, such

as nitrogen, oxygen or fluorine, that comes from

another molecule or chemical group. It is not a true

chemical bond. The hydrogen has a polar bonding to

another electronegative atom to create the bond.

These bonds can occur between molecules

(

intermolecularly

), or within different parts of a single

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Examples of intermolecular H- bonding: Water, HF

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HF, HCl. HBr, HI – boiling points

HF > HI > HBr > HCl

HF has highest boiling point because of intermolecular hydrogen bonding.

HI has the higher boiling point compared to HCl because of dipole-dipole interactions due to large size of Iodide molecule. Due to larger size of