Electronic Structure

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  • Created by: SophieHW
  • Created on: 04-03-15 11:09

Electronic Structure

Energy Levels

  • s - holds up to 2 electrons - spherical
  • p - holds up to 6 electrons - dumbbell shape
  • d - holds up to 10 electrons
  • f - holds up to 14 electrons

3d is higher in energy than 4s and so gets filled after the 4s

Shapes of orbitals

  • Orbitals represent the mathematical probabilities of finding an electron at any point within certain spatial distributions around the nucleus.
  • Each orbital has its own approximate, three dimensional shape.
  • It is not possible to draw the shape of orbitals precisely.

When a positive ion is formed electrons are lost

When a negative ion is formed electrons are gained

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Ionic and Metallic Bonding

Ionic Bonding

  • Metal atoms lose electrons to form +ve ions
  • Non-metal atoms gain electrons to form -ve ions.
  • Ionic bonding is stronger and the melting points higher when the ions are smaller and/or have higher charges. 

Metallic Bonding

The three main factors that affect the strength of the metallic bond are:

1. Number of protons/strength of nuclear attraction

The more protons, the stronger the bond

2. Number of delocalised electrons per atom (the outer shell electrons are delocalised)

The more delocalised electrons, the stronger the bond

3. Size of ions

The smaller the ion, the stronger the bond

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Dative Covalent Bonding

A dative covalent bond forms when the shared pair of electrons in the covalent bond come from only one of the bonding atoms. A dative covalent bond is also called co-ordinate bonding. 

Common examples you should be able to draw that contain a dative covalent bond: NH4H3ONH3BF3

The dative covalent bond acts like an ordinary covalent bond when thinking about the shape so in NH4the shape is tetrahedral

The direction of the arrow goes from the atom that is providing the lone pair to the atom that is deficient

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Bonding and Structure

Bonding and Structure

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Organic: Drawing Displayed Formulae

  • When drawing organic compounds add the hydrogens atoms so that each carbon has 4 bonds
  • Remember that the shape around the carbon atom in saturated hydrocarbons is tetrahedral and the bond angle is 109.5
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Organic: Drawing Displayed Formulae

General rules for naming carbon chains

  • Count the longest carbon chain and name appropriately 
  • Find any branched chains and count how many carbons they contain
  • Add the appropriate prefix for each branch chain 
  • E.g. -CHmethyl or -C3Hpropyl
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Basic Rules For Naming Functional Groups

The functional group is named by a prefix or a suffix, e.g. bromoethane, ethanol, propene

The position of the functional group on the carbon chain is given by a number - counting from the end of the molecule that gives the functional group the lowest number.

We only include numbers, however, if they are needed to avoid ambiguity.

Where there are two or more of the same groups, di-, tri- or tetra are put before the the suffix/prefix.

Words are separated by numbers with dashes.

Numbers are separated by commas.

If there is more than one functional group or side chain, the groups are listed in alphabetical order (ignoring any di, tri etc) 

The suffix for alkenes can go in front of other suffixes. 

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Isomers

Structural Isomers: same molecular formula but different structures (or structural formulae)

Structiral isomers can arise from:

  • Chain isomerism
  • Position isomerism
  • Functional group isomerism

Chain isomers: compounds with the same molecular formula but different structures of the carbon skeleton

Position isomers: compounds with the same molecular formula but different strucutres due to different positions of the same functional group on the same carbon skeleton

Functional group isomers: compounds with the same molecular formula but with atoms arranged to give different functional groups

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Electronegativity and Intermediate Bonding

Factors affecting electronegativity:

  • Electronegativity increases across a period as the number of protons increases and the atomic radius decreaes because the electrons in the same shell are pulled in more. 
  • It decreases down a group because the distance between the nucleus and the outer electrons increases and the shielding of inner shell electrons increases.

A compound cotaining elements of similar electronegativity and hence a small electronegativity difference will be purely covalent. 

Formation of a permanent dipole - (polar covalent) bond

  • A polar covalent bond forms when the elements in the bond have different electronegativities. 
  • When a bond is a polar covalent bond it has an unequal distribution of electrons in the bond and produces a charge separation, dipole δ+ δ- ends 

If all the bonds in a compound are the same polar bond and there are no lone pairs then the dipoles cancel out and the substance will be non polar, e.g. CClwill be non polar where as CH3Cl will be polar

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Electronegativity and Intermediate Bonding (cont.)

  • F, O, N and Cl are the most electronegative atoms
  • The element with the larger electronegativity in a polar compound will be the δ- end, e.g.

δ+  δ-

H - Cl

A compound containing elements of very different electronegativity and hence a very large electronegativity difference will be ionic.

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Intermolecular Bonding - Van der Waals Forces

Van der Waals forces occur between all molecular substances and noble gases. They do not occur in ionic substances.

  • These are also called transient, induced dipole-dipole interactions. They occur between all simple covalent molecules and the separate atoms in noble gases. 
  • In any moecule the electrons are moving constantly and randomly. As this happens the electron density can fluctuate and parts of the molecule become more or less negative i.e. small temporary or transient dipoles form.
  • These instantaneous dipoles can cause dipoles to form in neighbouring molecules. These are called induced dipoles. The induced dipole is always the opposite sign to the original one.

Main factor affecting the size of Van der Waals

  • The more electrons there are in the molecule the higher the chance that temporary dipoles will form. This makes the Van der Waals stronger between the molecules and so boiling points will be greater. 

The increasing boiling points of the halogens down the group 7 series can be explained by the increasing number of electrons in the bigger molecules causing an increase in the size of the Van der Waals between the molecules. This is why Iis a solid whereas Clis a gas.

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