Bonding and Structure

Bond angle

Bond angle

Angle between two covalent bonds on a molecule or giant covalent structure

Bond energy

Energy required to break one mole of a particular covalent bond in the gaseous state

Bond length

Internuclear distance of two covalently bonded atoms

Bond length and strength in covalent bonds

Bond strength is inversely proportional to bond length. Bond length depends on; Size of atom -> larger atom = longer bond (more electrons = more shielding = less attraction) Number of pairs of electrons shared -> more pairs = shorter bond. More electrons in shared electron cloud overcomes the repulsion of the nuclei meaning nuclei can remain closer together

Bonding in Ethene (C2H6)

Each carbon atom uses three of its four electrons to form sigma bonds Two sigma bonds are formed w/ the Hydrogen atoms and one sigma bonds formed with the other carbon atom Fourth electron from each carbon atom occupies a p orbital which overlaps sideways with another p orbital on the other carbon atom to form a pi bond Therefore C-C is a double bond: one sigma and one pi bond

Bonding in Hydrogen cyanide (HCN)

Contains a triple bond One sigma bond is formed between the H and C atom (overlap of an sp C hybridized orbital with the 1s H orbital) Second sigma bond formed between the c and N atom (overlap of an sp C hybridized orbital with an sp orbital of N) The remaining 2 sets of p orbitals of nitrogen and carbon will overlap to form two pi bonds at right angles to each other

Bonding in Nitrogen

Contains a triple bond Triple bond formed from overlap of the sp orbitals on each N to form a sigma bond and the overlap of two sets of p orbitals on the nitrogen atoms to form two pi bonds These pi bonds are at right angles to each other

Bonding spectrum

Pure covalent (electrons shared halfway between two atoms) Polar covalent (attracted more strongly to electronegative atom - has a permanent dipole) Ionic - electron is completely removed and transferred to the more electronegative atom.

charge density formula and relationship to EFofA

Charge density. = charge of ion / ionic radius Greater charge density = greater EFofA

Covalent bond

Strong electrostatic attraction between two nuclei and the shared pair of electrons between them.

Covalent bond properties - simple molecular

Usually gases, liquids or soft solids at room temperature - weak intermolecular forces Low melting and boiling points Cannot conduct electricity - no free electrons to carry charge Usually more soluble in non-polar solvents

Dative covalent bonds

A bond in which two atoms share a pair of electrons, both of which are donated by one atom. Behaves the same way as a covalent bond once formed. E.g. Al2Cl6 (lone pair on chlorine is donated to empty orbital on aluminium) and NH4+ (ammonia molecule shares its lone pair with a hydrogen ion H+)

Electron pair repulsion theory

Electron pairs in the outer shell of atoms and ions repel each other and get as far apart as possible.

Electronegativity

The ability of an atom to attract a shared pair of electrons in a covalent bond

Evidence of ions

Migration of ions e.g electrolysis of green copper chromate. ~ Cathode - blue (Cu2+ cations) ~ Anode - yellow (CrO4 2- anions) Electron density maps (from x-ray diffraction) Physical properties of ionic compounds

Giant lattice

Giant ionic lattices Giant covalent lattices e.g. Diamond and graphite Giant metallic lattices

How are pi bonds formed?

Sideways overlap of adjacent p-orbitals above and below the sigma bond

How are sigma bonds formed?

Direct overlap of orbitals (s AND p) between the bonding atoms

How is a single pi bond represented?

Single pi bond drawn as TWO electron clouds, one arising from each lobe of the p-orbitals

How to use Pauling electronegativity values to predict the formation of ionic and covalent bonds

IONIC = large E.N difference COVALENT = small/no E.N difference

Hybridization

The mixing of s orbitals with p orbitals to form molecular bonds

Hydrogen bonds

A strong intermolecular force between a delta positive hydrogen covalently bonded to fluorine, oxygen or nitrogen and a lone pair of electrons on the delta - atom (O,F,N) of a nearby molecule. Accounts for;

• high Bp of ammonia, water and HF

• open structure and low density of water

• solubility of alcohol in water Hydrogen is always delta + and has a 180 degree bond angle

IM force spectrum (strongest to weakest)

1. H-bonding (H-F, O-H, N-H bonds) + permanent dipole-dipole + LF

2. Permanent dipole-dipole + LF

3. LF

Compare elements w/ other elements w/ similar no. Of electrons to compare strength of IMF

Instantaneous dipole-induced dipole interactions (London forces)

Intermolecular forces that exist between all molecules. They arise from attractions between temporary instantaneous dipoles and the dipoles they induce in neighbouring molecules. More electrons in atom/molecule = stronger London force

Ionic bonding

The strong electrostatic forces of attraction between oppositely charged ions in a lattice

Lone pairs

A pair of of electrons in the outer shell of one of the atoms in a molecule or ion which is not involved in bonding Reduces bond angle by 2.5 degrees Counts as an area of electron density. E.g H20 = 109.5 - (2x2.5) = 104.5

Metallic bonding

The strong electrostatic attraction between metal cations and a sea of delocalised electrons

Multiple bonds

Double or triple bonds. Area of high electron density Counts as one area of electron density when calculating bond angles. Reduces bond length - high area of electron density = reduces repulsion of nuclei

Permanent dipole- permanent dipole interaction

Attractive forces between the positive pole of one molecule and the negative of another. Exist between polar molecules E.g. ICl

Physical properties of ionic compounds

Hard, brittle, crystalline - repulsion between ions if structure is distorted High melting and boiling points - strong attraction between oppositely charged ions which require lots of heat energy to overcome Soluble in polar substances (eg. H2O) - contain positive and negative ions which are broken apart by polar substances Conduct electricity in molten or aqueous - ions free to move

Polar molecules

Bonds between atoms of different elements with significant difference in electronegativity Are asymmetric - contain a net dipole moment. Not all molecules with polar bonds are polar

Properties of metals

• high mp - strong bonds in giant lattice require lots of heat energy to overcome/break

• high density - close packing with little space

• heat and electrical conductors - delocalised electrons free to move + transfer charge

• malleable - regular layers of metal cations can easily slide over each other. If a force causes a slip, ions can resettle into close packing.

Shapes of molecules

Electron d shape angle example 2. Linear. 180. BCl2 3. Trigonal planar 120 BeCl3 4. Tetrahedral 109.5. CH4 5. Trigonal bipyramid 180, 90, 120 PCl5 6. Octahedral 180, 90 SF6

Strength of ionic bonding

Ionic radius: The greater the size of the ion, the larger the distance between the two charges, so the smaller the force Ionic charge: Larger the charge, the stronger the bond

Trend in BP group 7 hydrides

HF has a much higher boiling point than HCl due to HF's ability to form hydrogen bonds. Hydrogen bonds are the strongest type of inter molecular forces and require the most energy to break. HCl, HBr, and HI all have permanent dipoles. The polarity of the H-X bond decreases down the group as E.N decreases so dipole-dipole forces decrease down the group. This would suggest the boiling temperature should fall from HCl to HI, however the number of electrons per molecule increases down the group, so HI is more polarisable than HBr and HBr more than HCL and therefore London forces increase down the group. The increase in London forces has more effect on the boiling temperature than dipole-dipole so the general trend is of increasing boiling point. London Forces win over dipole-dipole in this case.

Trend in melting points of metals (across the period and down the group)

Across the period, melting points of metal generally increase because the charge of the metal cation increases, thus the electrostatic force of attraction to the delocalised electrons increases, requiring more heat energy to overcome Down the group, melting points generally decrease because the shielding of the metal cation increases whilst charge stays the same, thus there is a lower force of attraction to delocalised electrons meaning less heat energy is required to be overcome

Trends and factors in electronegativity (across and period and down a group)

Across period: Increase in E.N due to an increase in nuclear charge but shielding being generally constant Down group: decrease in E.N as shielding and radius increases, weakling ability of nucleus to attract an e- pair

Trends in bond length and strength (down the group and as no. Of bonds increase)

Down the group: atomic radius increases so bond length increases and attraction between nucleus and shared electron pair decreases with distance so bond strength decreases As no. Of bonds increase: bond length decreases as EFofA increases so bond strength increases In pi bonds, orbital overlap is less efficient, therefore there is weaker bond enthalpy

Trends in ionic radii down a group - e.g. N3- to Al3+

The ionic radii decreases. The ions are isoelectronic as they have the same number and arrangement of electrons. The proton number increases down the group, so nuclear charge increases. Adding electrons to anions causes repulsion, so the ionic radius swells. Removing electrons decreases repulsion, causing the ion to contract.

When can a substance dissolve?

When the force of attraction between the solute and the solvent are strong enough to overcome the solute-solute and solvent-solvent forces of attraction For non-aqueous solvents the rule is 'like dissolves like'

When is a compound likely to dissolve in water?

Must be polar If it is ionic If it forms hydrogen bonds If it reacts with water (For soluble organic substances) if the chain length is not too long which would make the molecule more non-polar

Why are hydrogen bonds different to other permanent dipole- permanent dipole interactions

Hydrogen atom has no shielding electron and so has a stronger interaction with the lone pair on another small, highly electronegative atom

Why does water have higher melting and boiling points that would otherwise be expected?

Strong H- bonds in the -OH group require more thermal energy to overcome

Why does water have relatively high surface tension?

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Why is ice less dense than water?

H- bonds in ice are permanent meaning the water molecules are on average further apart because they are not making or breaking