Ionic- Stealing Covalent- Sharing is caring

Ionic- Stealing

Covalent- Sharing is caring

Metallic- Sea of Electrons

Ionic Bond – chemical bond resulting from the

electrostatic attraction between positive and

negative ions

How to determine:

*Look at electronegativity to determine the type of

bond.

*Difference in electronegativity of 3.3 to 1.7 is

considered an ionic bond.

Pure ionic bond – one atom has completely given

up one or more electrons; another atom has

gained them. (electrons are transferred)

**The bond is the opposite charge attraction that

occurs after each element has reached noble

gas stability **

Covalent bonds – chemical bonds resulting from

the sharing of electrons between two atoms

How to determine:

Difference in electronegativity

1.7 – 0.3 – 0.0

Polar covalent bond – united atoms have an

unequal attraction for shared electrons

Nonpolar covalent bond – bonding electrons are

shared equally by the bonded atoms, resulting in

a balanced distribution of electrical charge

Molecule – a group of two or more atoms held

together by covalent bonds and able to exist

independently

Diatomic molecule - molecule containing two

atoms (H2 , HCl)

Chemical formula – shorthand representation of

the composition of a substance using atomic

symbols and numerical subscripts.

A.K.A- Molecular formula

Ionic compound – composed of positive and

negative ions combined so that the positive and

negative charges are equal.

i.e. The positive and negative must be the same for

the compound to be stable

Other Ionic Compound Need-To-Know’s

Formula Unit – simplest unit indicated by the

formula of any ionic compound.

Formation of crystal structure involves many

ions. [orderly array of ions called a crystal

lattice].

Lattice energy – the energy released when one

mole of an ionic crystalline compound is formed

from gaseous ions.

Properties of Ionic compounds

+- attraction causes the macroscopic properties

-Relatively high melting points

- Brittle solids

-Conduct electricity in the molten state

(free charged ions)

Metallic Bonding

-Bonding that occurs within the atoms of a metal is different from ionic and covalent bonding.

-*It is a way that the atoms can bond with themselves*

-Metals can eject their two valence electrons becoming a positively charged ion but stay bonded through the localization of electrons

Metallic bond – a chemical bond resulting from the attraction between positive ions and mobile electrons.

Metallic bonding accounts for many properties of metals.

Examples:

Good conductors of electricity

Good conductors of heat

[free movement of electrons]

Luster-absorbs and re-emits light due to the many energies of electrons

Malleability (hammered into shape)

Ductility (drawn into a wire)

[metallic bonding is not directional]

What is an ionic bond?

What types of elements are involved in an ionic

bond?

Are Ionic bonds strong or weak?

(Hint: Melting Point)

What are the names of the following ionic

compounds

NaCl, MgCl2, Al2O3,Na3N

Covalent bonding- formation of a molecule that

requires the sharing of electrons between two or

more atoms.

Single bond – covalent bonds produced by the

sharing of one pair of electrons between two

atoms.

Double bond – covalent bond produced by the

sharing of two electron pairs

Triple bond – covalent bond produced by the

sharing of three electron pair

Unshared pair – a pair of electrons that is not involved in the bonding. These electrons belong exclusively to one atom.

Lewis structures:

Shows valence electrons

Structural formula:

Shows the bonds between the atoms

**Same as Ionic but the numbers matter now!

Step 1: Leave name of first element alone

Step 2: Change nane of second element to end

in –ide

Step 3: Add prefixes to signify the subscripts

Mono=1

Di=2

Tri=3

Step 1: Add up all of the valence electrons of EVERY atom in the chemical formula E.g. CH4 (C= 4, H=1, H=1, H=1, H=1;)

4+1+1+1+1= 8

Step 2: Determine a central atom (generally the first atom written in the compound)

**Never Hydrogen**

Step 3: Use 2 electrons or a – line to create single bonds between atoms.

Step 4: Add extra electrons as lone pairs to unstable atoms if needed**Remember the octet rule and that Hydrogen and Helium are weird**

Step 5: If atom is still unstable after lone pairs

are added. Look to move pairs of electrons to

form either double or triple bonds.

Step 6: Consider whether or not the compound

is a polyatomic ion

i.e. it cannot be stable and will have a charge

Covalent compounds

-Sharing electrons means that forces between the

molecules are weak

-Relatively low melting points

-Do not conduct electricity [No free ions]

Polyatomic ion – a charged group of covalently

bonded atoms

VSEPR Theory – electrostatic

repulsion between valence

level electron pairs

surrounding an atom causes

these pairs to be oriented as

far apart as possible

**i.e. the 3-d shape of the

molecule**

*Based on the fact that bonded

atoms and electrons want to

be as far away from each

other as possible*

Intermolecular Forces

A.K.A Van der Waals forces –attractive forces between molecules.

-Forces of attraction between molecules

i.e. the forces between all of the water molecules in a glass of water

** 2 types**

-Dipole-Dipole force- which occurs between the molecules of polar covalently bonded molecules

-London Dispersion Forces- which occur between the molecules of non-polar covalently bonded molecules

Dipole – equal but opposite charges separated by

a short distance.

Dipole – dipole force – forces of attraction between

polar molecules

Dipole-Dipole Forces are very common with

compounds containing Hydrogen

Hydrogen bonding – a dipole-dipole force;

London dispersion forces – weak intermolecular

forces responsible for the attraction between

nonpolar molecules.

London dispersion forces result from the constant

motion of electrons [creates instantaneous and

induced dipoles]