chemistry - unit 1 - section 3: bonding

Definition of ionic compounds

Oppositely charged ions held together by electrostatic potential

Examples of ionic bonding

To gain a full shell of electrons sodium gives an electron to chlorine. They both have full electrons shells. They form oppositely charged ions and attract to each other

What are the molecular ions

OH- hydroxide

NO3 nitrate

NH4+ ammonium

SO4 2- sulfate

CO3 2- carbonate

How to work out the formula for an ionic compound

Use the swap and drop method: same the powers and drop it to become the number of that specific element

Structures and feature of an ionic compound

• giant ionic structure

  1. regular structure

  2. Cubic shape

  3. Giant repeating pattern

  4. Positive ion with negative ion

• Dissolvable in water because water is polar and can attract the positive and negative ions

• Conducts electricity when molten or dissolved in solution as ions are free to move around

• High melting point as it has strong electrostatic force between the oppositely charged ions. Lots energy is required to overcome these forces

Definition of a covalent bond

the sharing of the outer electrons in order for atoms to obtain a full shell. Electrostatic attraction between the shared electrons and the positive nucleus

Definition of dative covalent bonds

When one atom donates two electrons to an atom or ion to form a bond

• an arrow is its representation when in displayed formula

Structure and features of diamond

• 4 carbon giant covalent structure

• Tetrahedral structure

• Bcs of it tightly packed and rigid it conduct heat well

• It can be cut into gem stones

• High melting point bcs of the strong covalent bonds

• Insoluble bcs it is difficult to break the covalent bond

• It has no delocalised electrons so it cannot conduct electricity

Structure and features of Graphite

• 3 carbon ionic compound

• 4th carbon electrons are delocalised

• It can conduct electricity bcs of the delocalised electrons

• It has high melting point as it has strong covalent bonds

• It is insoluble as it has strong covalent bonds that are require a lot of energy to break the bonds

• The layers slid over each other due to the weak forces

• It has low density as the distance between the layers is greater than the length of the covalent bonds

Shapes of molecules

The number of bonding pairs and lone pair determines the shape and angle of the compound

• bond pair bond pair < lone pair bond pair < lone pair lone pair

The electrons in the bonding pair repel equally away from each other as far apart as possible

• lone pairs change the shape and angles of the molecule; pushing the bond pairs closer together

• For every lone pair subtract 2.5 degrees

names of shapes of molecules with no lone pairs

Linear - 2 bonding pairs - 180 degrees

Trigonal planar - 3 bonding pairs - 120 degrees

Tetrahedral - 4 bonding pairs - 109.5 degrees

Trigonal bipyramidal - 5 bonding pairs - 120 and 90 degrees

Octahedral - 6 bonding pairs - 90 degrees

Shapes of molecules with lone pair

Pyramidal - 3 bonding pairs 1 lone pairs - 107 degrees

Bent - 2 bonding pairs 2 lone pairs - 104.5 degrees

Trigonal planar - 3 bonding pairs 2 lone pairs - 120 degrees

Square planar - 4 bonding pairs 2 lone pairs - 90 degrees

Definition of electronegativity

The ability of an atom to attract an electron towards itself in a covalent bond

• the further up and right you go in the periods the more electronegativity the element is

• The Pauling scale helps us to quantify the electronegative an element is

Definition of Polar bonds

Covalent bonds can become polar if the atoms attached to it have a difference in electronegativity

• The bigger the difference in electronegativity, the more polar the bond will be

• To show the polarity we use δ+ next to the the least electronegative element and the δ- next to the most electronegative

• Atoms that are bonded to the same electronegative value or similar electronegative values are not polar they share electrons in the middle

• Polar bonds that are symmetrical are not polar

• Uneven distribution of charge leads to polar molecules e.g H2O

Intermolecular force

Hydrogen bond

Permeates dipole - dipole

Van der waals

van der waals

Any molecules or atom with electrons can form a dipole when they move near the to another atom or molecule

• as electrons in a molecule can move from one end to another. Hence creating a temporary dipole

• When they move away the dipole interaction is destroyed

• Van der waals forces can hold some molecules in crystal structure

• The bigger the molecule the bigger the van der waals force as you have larger electron cloud

• Long straight chains hydrocarbons have more van der waals forces and so more energy is needed to over come these forces. This means the boiling point increases

• Branched hydrocarbon = can’t pack close together; weakens the van der waals forces between the chain and lowers the boiling point

Permanent dipole dipole

• interactions exist in molecules with a polarity

• Weak electrostatic forces that exist between molecules with a polarity

• Molecules have vander waals and permeate dipole dipole

• Polar molecules can be tested by placing a charged rod near a steady stream of polar liquid. The liquid bends towards the rod as the molecules align to face the oppositely charged rod

Hydrogen bonding

• the strongest intermolecular force and occurs when you have very electronegative elements

• Occurs when hydrogen on one molecule forms a bond with the lone pair on nitrogen, oxygen and fluorine

• Mole that have hydrogen bond and dipole - dipole and van der waals

Definition of metallic bonding

• giant metallic lattice structure

• Positive metallic ions are formed as metals donate electrons to form a sea of delocalised electrons

• There is an electrostatic attraction between positive metal ion and the negative delocalised electron

• Metals are good electrical conductor as the delocalised electrons are mobile and carry a current

• Good thermal conductor as the delocalised electrons can transfer kinetic energy

• High melting points due to strong electrostatic attractions

• Insoluble as the metallic bond too strong to break

Bonding summary