Bonding and Structure

• Ionic bonds occur between metals and non-metals
• Electrons are transferred from metal atoms to non-metal atoms
• Positive and negative ions are formed
• Ionic compounds always exist in a lattice

PROPERTIES:

Solid at room temperature

giant structures- high melting temperatures ( in order to melt an ionic compound, energy must be supplied to break up the lattice of ions)

*conduct electricity when molten or dissolved in water - ions are free to move in liquid state but not solid

*Britter and shatter easily - oppositely charged ions which a blow may move the ions and produce contact between ions with like charges

What is an Ionic Bond?

A chemical bond in which an electron/electrons are transferred from one atom to another resulting in the formation of oppositely charged ions with electrostatic forces of attraction between them

Why do Ionic compounds have high melting temperatures?

Because they have strong electrostatic attraction between the ions that extend through the whole structure which are difficult to break

Name properties of Ionic compounds

• Solid at room temperature
• Giant structures- high melting temperatures
• Conduct electricity when molten or dissolved in water - ions are free to move in liquid state but not solid
• Britter and shatter easily

A covalent bond forms between a pair of non-metal atoms. It is when a pair of electrons are shared between two atoms.

- Atoms with covalent bonds are held together by the elctroostatic attraction between the nuclei and the shared electrons.

DATIVE/ CO-ORDINATE

- an electron deficient atom accepts an electron pair (lone pair) from another atom.

An arrow is used to represent Dative Covalent bonding. The arrow points towards the atom that is accepting the lone pair

Co-ordinate bonds have exactly the same strength and length as ordinary covalent bonds

Electronegative- better at attracting electrons then the other. The power of an atom to attract the electron density in a covalent bond towards itself.

- measured using the Pauling scale ( runs from 0-4)

Electronegativity depends on:

• the nuclear charge
• the distance between the nucleus and the outer shell electrons
• the shielding of the nuclear charge by electrons in inner shells; the smaller the atom the closer the nucleus is to the shared outer main level elctrons and the greater its electronegativity; the greater the nuclear charge the greater the electronegativity

As you go up a group in the periodic table the electronegativity increases (the atoms get smaller) and there is less shielding by electrons in inner shells.

As you go across a period the electronegativity increases- the nuclear charge increases.

The greater the difference in electronegativity, the more polar the covalent bond is

In a metal element the outer main levels of the arom merge. Metals consist of a lattice of positive ions existing in a sea of ouer electrons which are delocalised.

-the number of delocalised electrons depends on the number of electrons lost by each metal atom

-metals have giant structures

PROPERTIES:

• Good conductors of electricity and heat; energy is spread by vigorous vibrations of the closely packed ions.
• Malleable and ductile; after a small distortion, each metal ion is till in the same environment as before so the new shape is retained
• High melting points and boiling points; giant structures- strong attraction between metal ions and the sea of delocalised electrons, this makes the atoms difficult to separate

.

The strength depends on-

-the charge of the ion; the greater the charge the greater the number of delocalised electrons and the stronger the attraction between positive ions and the electrons

-the size of the ion; the smaller the ion, the closer the electrons are to the positive nucleus and the stronger the bond

Metals tend to be strong as delocalised electrons extend throughout the solid so there are no individual bonds to break.

[1s2  2s2  2p6  3s2  3p6  4s2  3d10  4p6]

Molecules ans seperate atoms are held together by other, wekaer forces called intermolecular forces

There are three types of intermolecular forces:

+ van der Waals forces  ; act between all atoms and molecules

+ Dipole-dipole forces   ; act only between certain types of molecules

+ Hydrogen bonding   ; acts only between certain types of molecules

Molecules with polar bonds may have a dipole moment

in molecules with more then one polar bond the effects of each bond may cancel out to leave the molecule with no dipole moment.

EXAMPLE:

carbon dioxide is linear so the dipoles cancel out

Dipole-dipole forces act between molecules that have permanent dipoles.

E.G. hydrogen chloride - chlorine is more electronegative than hydrogen so the electrons are pulled towards the chlorine atom rather than the hydrogen atom. The molecule therefore has a dipole (H&+---Cl&-)

All atoms and molecules are made up of positive and negative charges even though they are neutral overall. These charges produce very weak electrostatic attractions between all atoms and molecules called van der Waals forces.

The distribution of the charge on an atom is changing at every instant.  An instant later the dipole may be in a different direction, but it is almost certain the atom will have a dipole at any point in time, even though it will be for an instant - a temporary dipole. This dipole affects the electron distribution in nearby atoms so that they are attracted to the original atom- the original atom has induced dipoles in nearby atoms

As the electron distribution of the original atom changes it will induce new dipoles in the atoms around it, which will be attracted to the original one. These forces are sometimes called instantaneous dipole-induced dipole forces; van der Waals

The more electrons there are the larger the instantaneous dipole will be therefore the size of the van der Waals forces increases with the number of electrons present.The boiling points of the noble gases increases as the atomic numbers of the noble gases increases and the boiling points of hydrocarbons increas with increased chain length as a result.

Hydrogen bonding has some characteristics of dipole-dipole attraction and some of a covalent bond. For a hydrogen bond to occur we need a very electronegative atom with a lone pair of electrons covalently bonded to a hydrogen atom. The atom must have a patial charge

OXYGEN         NITROGEN          FLUORINE

hydrogen- high boiling point because hydrogen bonding is present between the molecules and these strong intermolecular forces make the molecules more difficult to seperate

noble gases - gradual increase in boiling point because the only force is van der Waals which increase as electrons increase.

ICE: in liquid state hydrogen bonds easily break and reform because molecules can move about however when water freezes, the water molecules are no longer free to move which means the hydrogen bonds hold the molecules in fixed positions. The molecules are less closey packed as ice than in liquid watre (there are large gaps). This means that ice is less dense than water.

When we heat a solid it supplies energy to the particles making them vibrate more about a fixed position. This slightly increases the average distance between the particles and so the solid expands.

When changing a solid to a liquid more energy is needed to weaken the forces between the particles. The energy needed is called enthalpy change of fusion. while the solid is melting the temperature does not change because the heat energy provided is absorbed.

Enthalpy is the heat measured under constant pressure.

When changing a liquid to a gas we need to supply enough energy to break all the intermolecular forces between the particles. The energy needed is called the enthalpy change of vaporistaion.

Crystals are solid and the particles have a regular arrangment. The strength of the forces of attraction between the particles in the crystal affect the physical properties of the crystal.

They are called polymorphs or allotropes of carbon. They are macromolecular structures.

DIAMOND

• Giant covalent molecule. Each carbon atom is joined to 4 other carbon atoms in a tetrahedral arrangement- this makes a very strong and rigid structure.
• solid at room temperature, high melting + boiling points, does not conduct electricity
• crystalline- hard material
• insoluble in water abd non-aqueous solvents

GRAPHITE

• Giant covalent molecule arranged in layers held by van der Waals forces. The carbon atoms in the layers are arranged in hexagins ( connected to 3 others). These delocalised electrons flow along layers but not between them
• Solid at room temp, High MP + BP, conducts electricity in 1 direction, feels soft when rubbed- pressure makes the layers move across each other as the weak bonding is broken, insoluble in water and non-aqueous solvents.

Lone pairs repel as much as possible and sqeeze the shape of a molecule by approximately 2 degrees.