Chemical Bonding

Bonding

Chemical Bond: mutual electrical attraction

between nuclei and valence electrons of different

atoms

Type of bond depends on electron configuration

and electronegativity

Why do atoms bond?

TO BECOME STABLE!!Key: 8 valence electrons

Types of bonding

Ionic: an electrostatic force of attraction between

positive and negative ions; ions for when atoms lose/gain

electrons

Covalent: sharing of electron pairs between atoms

polar unequal sharing of electrons

nonpolar equal sharing of electrons

Metallic: an attraction between metal atoms and outer

mobile electrons

Ionic Bonding

compound formed by the electrostatic force of attraction

between positive and negative ions. It involves a transfer of

electrons

# (+) = # (-) (metals – nonmetals)

Formula unit: simplest combining ratio of ions in a compound

does not exist independently

http://www.visionlearning.com/library/module_viewer.php?mid=55

Dot Diagram for NaCl

Dot Diagram for CaBr2

Warm up Write the compound that forms between the following ions

A) Mg and F

B) Zinc (positive 2) and nitrogen

C) phosphorous and cesium

Which of the following are properties of ionic compounds?

High melting points shares electrons

non electrolyte transfers electron

Electrolyte Low melting points

Crystal lattice: 3-D crystal structure (arrangement of ions)

Lattice energy: energy released when one mole of ionic crystalline compound is formed from gaseous ions

http://www.visionlearning.com/library/module_viewer.php?mid=55

Properties

high melting point

most soluble in water

crystalline solid

good conductors in liquid or aqueous state

hard solids but will fracture

Metallic Bonding

Metallic bonding: an attraction of metallic atoms for

“delocalized electrons;” electrons roam traveling freely

from one atom to another

positive ions in a sea of electrons

http://www.revisioncentre.co.uk/gcse/chemistry/bonding.html

http://www.ndt-ed.org/EducationResources/CommunityCollege/Materials/Structure/metallic.htm

Fun Facts

The strength of a metallic bond is determined by

number of outer electrons softer metals can be

combined with harder metals. (forming alloys)

The strongest metals are the transition metals.

some “d” electrons are delocalized, involved in

bonding

Properties

good conductors of heat and electricity

shiny (luster)

hard, metallic crystals

malleable

ductile

high melting point

Covalent bonding

Covalent Bonding

Occurs between NONMETALS

Nonmetals have high electronegativities and want

to gain more electrons. They cannot lose valence

electrons. In order to bond, nonmetals must then

SHARE valence electrons.

The atoms share enough electrons to obtain 8

valence electrons (including the shared electrons)

Terms to know!

Molecular compound: neutral compound consisting of

nonmetals covalently bonded in which the electrons are

shared

Molecule: smallest representative unit of a molecular

compound, can exist independently

Diatomic Molecules (Memorize these!)

Diatomic molecules: molecules consisting of two atoms of the

same element. They are ALWAYS found in pairs.

Ex: H2 O2 F2 Br2 I2 N2 Cl2 “HOFBrINCl”

Form a “7” starting at atomic number (Z) 7 and include H

Properties

low melting & boiling points

brittle, dull solids or gases

poor conductors of heat and electricity

polyatomic ions

Polyatomic Ions are covalently bonded group of atoms that

have a charge.

Examples: Nitrate (NO3-), Hydroxide (OH-), Ammonium

(NH4+)

Nitrate Structure

Let’s try it!

Calcium and nitrate

Sodium and phosphate

Definitions

Bond length: average distance between nuclei of

two bonded atoms (sum of atomic radii)

Bond angle: angle between two bonds in a

molecule

Bond(strength) energy: energy needed to break a

bond and form neutral atoms

Facts As bond length increases, bond energy decreases. More bonding regions increases the amount of bond

energy

As electronegativity differences increase, bond energy increases.

**Think about it!!!**What relationship do each of the above have?

Apply That Information!

These diagrams are not

drawn to scale

Answer: What can you say about the bond

strength and bond length of these 2

compounds?

Octet Rule

Definition: atoms tend to gain, lose or share electrons so

that they have eight valence electrons

exceptions Hydrogen (2e-),

Beryllium (4e-), Boron (6e-)

Some elements can have expanded octets (more than 8)

example: Sulfur (can have 16 e-)

http://chemed.chem.wisc.edu/chempaths/GenChem-Textbook/Exceptions-to-the-Octet-Rule-573.html

Other covalent bonds

Covalent network bonding: 3-D network of

covalently bonded atoms macromolecules

Very high melting points, very hard solids

Ex: diamond, graphite, quartz

Network Covalent Bonding

Coordinate Covalent bond

When both of the shared electrons in a covalent bond come

from the same atom

http://www.chemguide.co.uk/atoms/bonding/dative.html

Equal sharing of electron between atoms in a compound exist

in NON POLAR MOLECULES

Unequal sharing of electrons is a POLAR MOLECULE

Determine Bond from Electronegativity

Differences

0 0.3 1.7 3.3

Covalent Ionic

Non-polar Polar

Determine the type of bond

Use the following electronegativity values to determine the

bond character of each:

H – S (2.2, 2.6)

S – Cl (2.6,3.2)

Rb – Se (0.8,2.6)

C – I (2.6,2.7)

Cs – S (0.8, 2.6)

O – O (3.4,3.4)

F – I (4.0,2.7)

Ca – N (1.0,3.0)

Molecular geometry

a.k.a. Molecular Shape

Polar vs. Nonpolar

Nonpolar molecule: equal sharing of electrons between atoms in a compound, no positive or negative poles exist

ex: H2 H : H

Polar molecule: unequal sharing of electrons within a molecule positive and negative ends exist

ex: H2O

Molecular polarity the distribution of molecular

charge (even or uneven)

Molecular polarity depends on:

symmetry

molecular shape

Lone pairs

Molecular polarity influences intermolecular

forces.

Polar vs Non Polar Structures

Water is polar

Methane is

non polar

Hydrogen

Fluoride is

polar

VSEPR theory

Valence Shell Electron Pair Repulsion

electron pairs (clouds) spread as far apart as

possible to minimize repulsive forces

Important VSEPR terms

Lone Pair: A pair of electrons that do NOT create a chemical

bond

Shared Pair: A pair of electrons that make a chemical bond

between elementsLone Pair of

electrons in an

electron cloud

Shared Pair of

electrons

VSEPRShape # Bonds to

Central Atom

# Lone Pairs

to Central

Atom

Examples Bond Angles

Linear 2 atoms

together or 2

bonds to central

atom

0 H-Cl

O = C = O

180°

Bent 2

2

2

1

104.5°

Trigonal Planar 3 0 120°

Trigonal

Pyramidal

3 1 107.5°

Tetrahedral 4 0 109.5°

Drawing Structural Formulas

1. Place the least electronegative substance in the middle (Hydrogen will

be at the end because it only wants 2 e-’s)

2. Calculate the number of valence electrons available (this should always

be an even number)

3. Divide that number by 2 to get your bonding pairs

4. Connect all of your atoms in your molecule

5. Use remaining bonding pairs for lone pairs, until all of the bonding

pairs are used

6. Check if all atoms are stable(following the octet rule, unless it is an

exception), if all happy your structure is done

7. Make double and triple bonds to make remaining atoms stable

Let’s Try these examples!

NH3

CH4

H2O

BF3

CO2

Molecular Polarity

Nonpolar molecules are usually symmetrical. (about their

bonds)

Polar molecules are usually nonsymmetrical. (about their

bonds)

Polar molecules are called dipoles. (positive and negative

ends)

**As soon as you have a lone pair on your central

atom, your molecule is polar.**

Two shapes are always polar: bent and trigonal pyramidal.

Showing molecular polarity

To denote a polar bond, an arrow is drawn pointing to the

more electronegative substance, with a plus sign at the end

that is less electronegative

To denote a polar molecule, a lowercase sigma (δ-) is placed

at the more negative end, and a lowercase sigma (δ+) is

placed at the more positive end

Intermolecular forces

Intermolecular Forces

Intermolecular Forces

Definitions: weak forces of attraction between

molecules

Types

Dipole-dipole

Dipole-induced dipole

Hydrogen

London Dispersion

Dipole-Dipole

Exist between polar molecules, higher melting point and

boiling point than expected substances exist mostly as

solids or liquids

http://www.chem.purdue.edu/gchelp/liquids/dipdip.html

Hydrogen bonding

Special dipole-dipole force occurring when

Hydrogen on one molecule is attracted to N, O, F

of another molecule

Ex: gives water its unusual properties, ice

floats in liquid water, higher melting/boiling

point, surface tension

Hydrogen Bonding

http://www.chemguide.co.uk/atoms/bonding/hbond.html

London Dispersion

Nonpolar-Nonpolar instantaneous dipole due to a

shift in electron strength increase with an increase

in number of electrons; low melting

point/boiling point, mostly gases

Ex: diatomic molecules, noble gases

Determine the IMF(s) Present

1) H2O

2) SCl23) PF3

Helpful Hint: Draw the molecular

structure and determine if the bond is polar

or non polar.

Physical properties and bonding

Melting point: temperature at which

solidliquid

Boiling point: temperature at which liquidgas

Density: mass/volume, units: g/mL or g/cm3

Color: some transition metals produce colored

ions

Ex: copper blue-green

Solubility: ability of a solute to dissolve in a given

amount of solvent makes a solution

How do these properties relate to bond

type or intermolecular forces?

Stronger bonds/imf’s

higher melting point/boiling point

Weaker bonds/imf’s

lower melting point/boiling point

Solubility “likes dissolves like” NP-NP no

attractive forces exist between solute and

solvent, random mixing