UNIT 1 structure and properties of matter

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 th

UNIT 1: Structure and Properties of Matter

Atomic Models and Properties of

Atoms

Chemical Bonding and Properties of

Matter

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

UNIT 2 Section 4.1

4.1 Models of Chemical Bonding

Chapter 4: Chemical Bonding and Properties of Matter

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Three types of chemical bonding are ionic, covalent, and metallic

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?

Chapter 4: Chemical Bonding and Properties of Matter

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

the more electronegative atom

UNIT 2 Section 4.1

Scientists have categorized types of bonds according to ΔEN

mostly ionic

polar covalent

mostly covalent (non-polar)

Three categories of bonds have

been set based on ΔEN

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.

Microscopic analysis shows that the

structure of metals consists of aggregates

of crystals.

UNIT 2 Section 4.1

Properties of Metals

Melting and Boiling Points

boiling points of pure metals

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Chapter 4: Chemical Bonding and Properties of Matter

Periodic table trends include:

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

2 For Groups 1 to 6, across a period, melting points increase as atomic number increases.

Malleability and Ductility

because, when struck, the metal ions can slide by one

another while the electrons still surround them

Hardness

crystal size (smaller ones make harder metals)

Electrical and Thermal Conductivity

free to move from one atom to the next

UNIT 2 Section 4.1

Alloys

Alloys are solid mixtures of two or more metals

amount, can significantly affect the properties of a substance

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

between the larger atoms.

• occurs when ΔEN is between 1.7 and 3.3

electrons and another atom gaining those electron(s)

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

Different types of crystal structures can form

of crystal structure that an ionic compound will form

UNIT 2 Section 4.1

Properties of Ionic Compounds

Melting and Boiling Points

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

molecules and NaCl ions act

to break apart the ionic bonds.

Mechanical Properties

Ionic crystal will break on smooth planes, where like charges become aligned.

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

Size of Na+ ion is smaller

than Cl- ion.

WHY??

Why Ionic Compounds exist as Crystals?

Positive and negative ions attract each other with electrostatic force that extends in all directions This means that ions will

be bonded to a number of oppositely charged ions around them This leads to formation of an alternating cation-anion pattern of crystal lattice

Why Ionic Compounds are hard and

Why Ionic Compounds have high melting points?

Since in crystal lattice the positive and negative ions are tightly held in their positions in the lattice, only at high temperature

do the ions acquire sufficient kinetic energy to overcome their attractive forces and attain the freedom of movement as in a liquid

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

surrounded by water molecules and can lead an independent

existence and are thus dissolved in water

UNIT 2 Section 4.1

Covalent Bonding

The length of a covalent bond is determined by different electrostatic forces.

Forces in covalent bonds:

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Chapter 4: Chemical Bonding and Properties of Matter

• polar covalent: atoms do not share electrons equally

• non-polar covalent: atoms share electrons almost equally

Describe the chemical bonding and structure of NaCl How do bonding and structure influence the general properties

of the substance?

Answer on the next slide

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

UNIT 2 Section 4.1

Quantum Mechanics and Bonding

Valence Bond (VB) Theory explains bond formation and

molecular shapes based on orbital overlap

electrons, which have opposite spins

greater the overlap, the stronger and more stable the bond

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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Chapter 4: Chemical Bonding and Properties of Matter

UNIT 2 Section 4.1

Quantum Mechanics and Bonding

According to MO theory:

that results in formation of new orbitals called molecular orbitals

different from those of atomic orbitals

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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Chapter 4: Chemical Bonding and Properties of Matter

UNIT 2 Section 4.1

Explaining Single Bonds

For molecules like hydrogen fluoride:

According to MO theory, the bond is a sigma (σ) bond, ) bond,

which is symmetrical and freely rotates

For molecules like methane:

molecular shape

three 2p orbitals and a 2s orbital so that four identical

bonds can be created

The four sp3 orbitals of C overlap with the s orbitals of H to form methane.

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

for each carbon (a 2s orbital mixes with two 2p orbitals)

For bond formation in ethene:

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

between the carbons

orbitals of the hydrogens to form σ) bond, bonds

and below the plane to form a pi (π) bond) bond

For molecules like ethyne:

hybrid orbitals for each carbon (a 2s orbital + a 2p

orbital)

carbon and between sp of carbons and 1s of hydrogens

each carbon

Allotropes are compounds that consist of the same element

but have different physical properties

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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Chapter 4: Chemical Bonding and Properties of Matter

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

Covalent Network Solids

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

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Chapter 4: Chemical Bonding and Properties of Matter

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

The shape of a molecule is the result

of the presence of atoms, bonding electrons, and non-bonding electrons,

as well as forces of attraction and repulsion.

UNIT 2 Section 4.2

Depicting Two-Dimensional Structures of

Molecules with Lewis Structures

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Chapter 4: Chemical Bonding and Properties of Matter

Co-ordinate Covalent Bonds:

covalent bonds and therefore are not

indicated in Lewis structures

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Chapter 4: Chemical Bonding and Properties of Matter

Expanded Octet (Expanded Valence):

electrons

around the central atom.

UNIT 2 Section 4.2

Some Exceptions When Drawing

Lewis Structures

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

support Lewis structures

that show same relative position of

atoms but different positions of

electron pairs

Actual bond lengths in ozone are between those of single and double bonds.

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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Chapter 4: Chemical Bonding and Properties of Matter

For VSEPR, there are five electron-group arrangements

(Electron groups are represented by bars).

positioned as far apart as possible (repulsion)

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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Chapter 4: Chemical Bonding and Properties of Matter

UNIT 2 Section 4.2

Guidelines for Using VSEPR Theory to

Predict Molecular Shape

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Chapter 4: Chemical Bonding and Properties of Matter

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

Answer on the next slide

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Chapter 3: Atomic Models and Properties of Atoms

LEARNING CHECK

HCN has two bonding groups and

UNIT 2 Section 4.2

The Influence of Molecular Shape

on Polarity

direction of higher electron density

UNIT 2 Section 4.2

Determining Whether a

Molecule is Polar

A molecule with one or more polar bonds is not necessarily a

polarity as a whole can be determined by adding the vectors

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Chapter 4: Chemical Bonding and Properties of Matter

Both water and carbon dioxide have two polar bonds But water’s bent shape results in

a polar molecule, while carbon dioxide’s linear shape results in a non-polar molecule

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

UNIT 2 Section 4.2

Dipole-Dipole

Dipole-dipole forces:

have a region of partial positive charge and a region of partial negative charge

differences between polar and non-polar molecules

molecules and ions

magnitude of the partial charge and size of the 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:

• are forces of attraction between an ion and a non-polar

molecule that has an induced dipole due to the nearby ion

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 force of attraction

Hydrogen Bonding

A hydrogen bond is the attractive interaction of a

hydrogen atom and an electronegative atom, such

another electronegative atom to create the bond These bonds can occur between molecules

(intermolecularly), or within different parts of a single molecule (intramolecularly).

Examples of intermolecular H- bonding: Water, HFExamples of intramolecular H – Bonding:

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

iodide, induced polarity is greater and hence

stronger the bond formed, and thus higher the

boiling point