Bonding

Metal and non-metal

Metal atoms transfer electrons to non-metals

Lattice structure

• The size of the ions

• Charge of the ions

The smaller the ion they stronger the attraction to the electrons

• Solid

• They have a giant structure therefore a high melting temperature as the forces require a lot of energy to overcome

• Molten and aqueous

• Ions that carry the current are free to move in the liquid state

• They form a lattice of alternating positive and negative ions, a blow may move the ions and produce a contact between like charges

When two or more atoms bond together

• Atoms are held together by covalent bonds

• Molecules are held together by weak intermolecular forces

• Carbon atoms are arranged in sheets of flat hexagons covalently bonded with 3 other atoms

• The fourth electron is delocalised

• The sheets are held together by a weak van der Waals force

• Weak bonds between the layers mean that it is slippery and can slide over each other

• Delocalised electrons in graphite are free to move along the sheets

• Layers are quite far apart compared to its length of the bond so it has a low density

• High melting point

• Insoluble in any solvent

• Each carbon atom has 4 bonds

• They arrange themselves in a tetrahedral shape

• High melting point

• Hard

• Good thermal conductor as particles allow vibrations to pass through

• Can’t conduct electricity

• Won’t dissolve in any solvent

• This is where one of the atoms provides both of the shared electrons

• This occurs when an atom has a lone pair of electrons and the other has none to share

• Electron pairs repel each other

• This results in each pair of electrons being at positions at the greatest possible distance from each other

• Full line represents bond in no direction

• Wedge represents bond pointing towards you

• Many lines represents bond pointing away from you

1. Find the central atom

2. Work out how many electrons are in the outer shell

3. Add 1 electron for every atom that the central atom is bonded to (take ions into account)

4. Add up the electrons, Divide by 2

5. Compare the number of electron pairs to the number of bonds to find lone pairs

• Linear

• 2 electron pairs

• 2 bond pairs

• 0 lone pairs

• 180 bond angle

• Trigonal planar

• 3 electron pairs

• 3 bond pairs

• 0 lone pairs

• 120 bond angle

• Tetrahedral

• 4 electron pairs

• 4 bond pairs

• 0 lone pairs

• 109.5 bond angle

• Trigonal pyramidal

• 4 electron pairs

• 3 bond pairs

• 1 lone pair

• 107 bond angle

• Bent (V-shape)

• 4 electron pairs

• 2 bond pairs

• 2 lone pairs

• 104.5 bond angle

• Trigonal Bipyramidal

• 5 electron pairs

• 5 bond pairs

• 0 lone pairs

• 120/90 bond angles

• Seesaw

• 5 electron pairs

• 4 bond pairs

• 1 lone pair

• 86/102 bond angles

• T-shape

• 5 electron pairs

• 3 bond pairs

• 2 lone pairs

• 90 bond angles

• Octahedral

• 6 electron pairs

• 6 bond pairs

• 0 lone pairs

• 90 bond angles

• Square pyramidal

• 6 electron pairs

• 5 bond pairs

• 1 lone pair

• 90 bond angles

• Square planar

• 6 electron pairs

• 4 bond pairs

• 2 lone pairs

• 90 bond angles

The power of an atom to attract the electrons in a covalent bond

F, O, N, Cl

A higher number means an element is better able to attract bonding electrons

• When atoms have equal number of electronegativity

• Or when both aren’t very electronegative

• This means that the electrons are shared equally

• An example is F2

• When one atom is more electronegative than another

• Electrons shared will be closer to the more electronegative one

• An example is H2O

It is a difference in charge between two atoms caused by a shift in electrons density in the bond

If charge is distributed unevenly over a whole molecule then the molecule will have a permanent dipole

• Causes all atoms and molecules to be attracted to each other

• Electrons in charge clouds are always moving quickly

• These move and and cause temporary dipoles

• These are attracted to each other

• Dipoles destroy and create themselves all the time, but the overall effect is the same

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  • larger molecules have larger electrons clouds which mean they have stronger van de Waals

  • Branched molecules are weaker

• In molecules that have permanent dipoles, there will be weak electrostatic forces of attraction between the δ+ and δ- charges on neighbouring molecules

• Strongest intermolecular forces

• It only happens when hydrogen is covalently bonded to F, N or O

• The bond is polarised, and hydrogen has such a high chance density that the hydrogen atoms form weak bonds with lone pairs of electrons

• They don’t conduct electricity

• Low melting points

• Can dissolve in water depending on low polarised they are

• As you go down group 7 hydrides, two factors complete

  • Polarity of the molecule

  • Number of electrons

• General trend is that it increases

• The outermost electrons of a metal atom becomes delocalised

• This leaves a positive metal ion surrounded by a sea of delocalised electrons

• These exist as giant metallic lattice structures

• Metals have high melting points

• They are malleable

• They are ductile

• They can conduct electricity

• They are mostly insoluble