Chemical Bonding For Leaving Certificate Chemistry This

Chemical Bonding For Leaving Certificate Chemistry

This presentation is divided in 5 major areas. 1. 2. 3. 4. 5. Chemical Compounds. Ionic Bonding. Covalent Bonding. Electronegativity. Shapes of Molecules and Intermolecular forces.

1 Chemical compounds. • A compound is formed when two or more elements combine in a chemical reaction. E. G. hydrogen gas burning in oxygen will result in formation of a compound of water. 2 H 2(G) + O 2 (G) 2 H 20(L) • Compounds can be broken down into their elements. If an electric current is passed through water the compound splits into its elements of hydrogen and oxygen.

Chemical formulas • Na. Cl, H 2 O, CO 2 are all compounds • When written as Na. Cl, H 2 O, and CO 2, they are chemical formulas representing these compounds (like shorthand). • Let’s examine these formulas: • Na. Cl This means that there is: • 1 atom of Na ratio • 1 atom of Cl 1: 1

• • H 2 O This means that there is: 2 atoms of H ratio 1 atom of O 2: 1 • CO 2 • This means that there • • is: 1 atom of C 2 atoms of O ratio 1: 2 • Fe 2 O 3 (rust) • This means that there • • is: 2 atoms of Fe 3 atoms of O ratio 2: 3

• Noble gases are • chemically stable because their outer energy level is filled. The first two elements in group 0 are inert and do not form bonds, however, the higher members form some bonds. Atom Electron arrangement He 2 Ne 2, 8 Ar 2, 8, 8 Kr 2, 8, 18, 8

The octet rule • Elements will try to lose gain or share electrons to • • achieve 8 electrons in their outer shell. The octet rule states that atoms on reaction tend to reach an electron arrangement with 8 electrons in the outermost energy level. This outermost energy level is also known as the valence shell.

Valency Group number No. of Electrons in outer shell Electrons needed/to be lost for octet rule Valency 1 1 e- (lost) +1 2 2 e- (lost) +2 3 3 e- (lost) +3 4 4 e- (shared) +/-4 5 5 e- 3 e- (gained) -3 6 6 e- 2 e- (gained) -2 7 7 e- 1 e- (gained) -1 8 8 e- 0

Valency of Transition Elements. • Transition elements have a variable valency. • The valency depends on the other elements the • • transition element s bonding to. E. g. Copper(1) oxide, copper(2) oxide, iron (3) oxide. Similarly chromium and manganese have variable valency, chromium (3), chromium (6), manganese (2) and manganese (7).

• Helium is a much safer alternative to hydrogen as it is stable and is therefore used in weather balloons and blimps. • Argon is used in electric light bulbs to prevent the tungsten filament from evaporating or reacting.

2 Ionic Bonding • Definition: • • • ionic bonding is the electrostatic force of attraction between ions of opposite charge. What are ions? These are elements which have a positive or negative charge. E. G. Na has 11 e- (E. C. = 2, 8, 1. ) When Na gives away this one e- it now has more protons than electrons so it has an overall positive charge. Ionic bonds generally form between metals and non-metals.

• In the case of sodium • chloride sodium donates 1 electron to achieve a full stable 8 electron outer valence shell. Chlorine accepts this 1 electron as it only requires 1 e- to have a complete stable outer shell. 1 electron transfers sodium chlorine Na+ Cl-

Crystal lattice structures • Ionic bonds result in a crystal lattice structure.

Properties of Ionic Compounds Contain positive and negative ions (Na+Cl-) Solids such as table salt (Na. Cl(s)) High melting and boiling points Strong force of attraction between particles Separate into charged particles in water to give a solution that conducts electricity.

Uses of ionic materials • Salt tablets are taken to replace lost salt in perspiration. • Brine is used to cure bacon in a preservation process. • Fluoridation of water supplies to prevent tooth decay.

Drawing Lewis Structures (Dot and Cross) • Sum the valence electrons from all atoms. Add one for each negative charge and subtract one for each positive charge. • Draw a skeleton structure with atoms attached by single bonds. • • • Complete the octets of atoms bound to the central atom. Place extra electrons on the central atom. If the central atom doesn’t have an octet, try forming multiple bonds.

Dot and cross diagrams for ionic bonding • E. G. 2 – Mg. CL • E. G. 1 – Na. Cl. 2

3 Covalent Bonding • Definition: • • • A covalent bond is a shared pair of electrons. A single bond has 1 shared pair of electrons. A double bond has 2 shared pairs of electrons. A triple bond has 3 shared pairs of electrons. E. G. H-H O=O N N Covalent bonds are typical of non-metal elements. (Metals mix to form alloys. )

Dot and Cross Diagrams for Covalent Bonding O 2: ·· O =O ·· ·· ·· N N ·· N 2: ·· The number of electron pairs is the bond order.

Sigma bonds vs Pi bonds Sigma bonds. • A sigma bond is formed when electrons are shared in line with the nuclei. (a head-on overlap of orbitals. ) Pi bonds. • A Pi bond is formed when the shared orbitals are side on i. e. (not in line with the nuclei. ) N. B. Sigma bonds are stronger. In a covalent single bond it is a sigma bond, however, in a double or triple bond there is 1 sigma bond and the others are pi bonds.

Polar and Non-Polar Covalent Bonding • A polar bond can be pictured using partial charges: + H 2. 1 = Cl 3. 0 0. 9

Polar and non-polar materials in everyday life. • Polar – Water and Glucose. • Non-polar – Cooking oil and Petrol.

Covalent Compounds Exist as neutral molecules (C 6 H 12 O 2) Solids, liquids, or gases (C 6 H 12 O 6(s), H 2 O(l), CO 2(g)) Lower melting and boiling points (i. e. , often exist as a liquid or gas at room temperature) Relatively weak force of attraction between molecules Remain as same molecule in water and will not conduct electricity

4 Electronegativity • Polarity refers to a separation of positive and negative charge. In a nonpolar bond, the bonding electrons are shared equally: H 2, Cl 2: • In a polar bond, electrons are shared unequally because of the difference in electron density.

Electronegativity • Electronegativity refers to the ability of an atom in a molecule to attract shared electrons. • The Pauling scale of electronegativity:

Trends in Electronegativity • Down a group: electronegativity decreases 1. An extra level of electrons shields the outer e 2. • 1. 2. from the nucleus. An increase in atomic radius. Across a period: Electronegativity increases Increase in the size of nucleus (nuclear charge increases). The atomic radius decreases.

0. 0 – 0. 4 – 1. 7

5 Shapes of Molecules The Electron Pair Repulsion Theory EPR • The shape of a molecule depends on the number of pairs • of electrons that lie around the central atom of the molecule. The following shapes will arise: Linear – one or two pairs of e- around the central atom. Bond angle is 180º Trigonal planar – Three pairs of e-, bond angle is 120º Tetrahedral – Four pairs of e-, bond angle is 109. 5º

Linear shape Trigonal planar Click on the image to view the videos. Tetrahedral

Loan Pairs and EPR • If un-bonded pairs of electrons remain • 1. 2. on the central atom they will distort the shape of the molecule. Two examples include: H 2 O NH 3

Water is V-Shaped Ammonia is pyramidal Click on the image to view the videos.

Polarity in Molecules • Polar molecules have a centre of positive charge • separated from the centre of the negative charge. Even though the individual bonds may be non- polar the overall molecule is polar for this reason. E. G. NH 3 and H 2 O Non-polar molecules have the centre of the positive and negative charges in one location. There is an equal pull on the e- even though the individual bonds may be polar. E. G. CCl 4 and CO 2

Inter and Intra Molecular Forces • The force of attraction between ions is stronger than between molecules. • Inter: are forces between molecules. • Intra: are forces within a molecule. • There are 3 kinds of forces that can attract molecules together

1. Van der Waal’s: • These are the weakest forces caused by • • the movement of e- within a molecule. The electrons move randomly within the bond so at 1 point in time they are nearer to 1 atom than the other. This induces a temporary dipole force. Temporary dipoles will result in increased boiling points. The greater number of e- in a molecule the greater the number of temporary dipoles.

2. Dipole-dipole: • The positively charged end of a polar • • molecule is attracted to the negative end of another molecule. The dipoles in this case are permanent. As a result they are stronger than Van der Waal’s forces.

3. Hydrogen Bonding • When H is bonded to F, O or N these elements are sufficiently electronegative to make the bond polar. • H has only 1 e- in its atom, a strong • partial positive charge will result. This means it is very strongly attracted to the negative atom and as a result H 2 O is a liquid at room temperature with a fairly high boiling point.