An attraction between oppositely charged ions which are formed by the transfer of electrons from one atom to another: Generally involves attaction between oppositely charged ions in a lattice. Can form giant ionic structures, usually metal to non-metal.
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- High boiling and melting points (needed to break strong electrostatic forces)
- Can conduct electricity when molten or in solution (free ions)
- Usually hard and brittle
- Soluble in water
- Solid at room temperature
- Low thermal conductivity
A pair of electrons shared between two atoms: normally non-emtal atoms
- Usually liquid or gas and room temperature Pic space :)
- Low melting and boiling points
- Usually insoluble in water
- Poor conductors
- Soft in comparison to ionic
- Can form simple molecules or giant structures
In a standard covalent bond, each atom provides 1 electron to the bond represented as a single short straight line between the two atoms
Covalent bonding happens as the electrons are more stable when attracted to two nuclei rather than one, not fixed, in a state of constant motion.
Examples of Giant Covalent Strcutures
- Carbons bonded to four other carbons
- Hard and shiny
- Tetrahedral structure
- Carbon atoms bonded to three other carbons
- Fourth electron is delocalised so graphite can conduct electricity
- Layers can slide over each other
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Co-ordinate Bonding/Dative Covalent Bond
This forms when a pair of electrons is shared between two atoms, but only one atom provides both of these electrons
- Atoms that can form co-ordinate bonds must have at least one spare pair of electrons, known as a lone pair
- The electron providing the pair is known as the electron pair donor
- The atom accepting pair must be electron deficient, this atom is known as the electron pair accepter
- The 1s energy level is vaccant and hence can accept a pair of electrons
Can be shown as a short arrow from the donor to the acceptor
H(3)N: --------> BF(3)
An attraction between a lattice cations and a sea of electrons
- Metallic bonds are formed when atoms lose electrons and so the resulting electrons are all attracted to the resulting cations
- Happens as electrons are attracted to more than one nucleus so are stable and so the electrons are delocalized, not attached to any particular ion but are free to move among them
- Good conductors of heat and electricity Pic space :)
- Very malleable
- Tensile strength
- Hard and shiny
The ability of an atom to attract a bonding pair of electrons in a covalent bond
- Electronegativity increases across the periods as the nuclear charge increases and as the shielding stays the same this results in the electrons being strongly attracted to the atom
In turn, it will decrease down a group as the shielding increases and so the electrons are less strongly attracted to the atom
An electron with a high electronegativity is said to be electronegative, whereas the opposite is said to be electropositive
- Electronegativity is measured on a scale from 0.7 to 4.0. In the Periodic Table, the pattern is a diagonal line from left to right, with F being the most electronegative. Noble gases don't apply as they are stable, full outer shell
Electronegativity in Bonding
Covalent: If both atoms have a similar electronegativity, they both attract the bonding pair with the same strength and so the bond is covalent and it remains midway between the two, shared equally.
Polar Covalent: If one more more electronegative than the other, the bonding pair will be more attracted to this atom, so the electronegative atom gains a slight negative charge and the other atom a slight positive charge. This is called a polar covalent bond.
Ionic: If the electronegativity difference is large, the 'sharing' is so uneven, the electronegative atom now owns the pair, they have been transferred. They both gain full charges and this is an ionic bond.
Metallic: If both atoms are electropositive, neither can attract the electrons and so the electrons are free to move about, both atoms gain a positive charge and the bond is metallic.
Deciding the bond type according to electronegativ
- Less/between 1.6 - 1.9 = Metallic
- If either atom has a value greater than 1.9 with a difference between the atoms less than 0.5 = Covalent
- If either atom has a value greater than 1.9 with a difference between the atoms more than 0.5 but less than 2.1 = Polar covalent
- If the difference is greater than 2.1 = Ionic
There are exceptions, these are guide values and you shouldn't be actually asked a question on this, not on the AQA spec, but my guide website included it, so i did.
Molecular Shapes Basics
- Electron pairs in covalent bonds will repel each other, this is known as mutual repulsion. They will repel as far apart as possible
- The shape of the molecule depends on the number of pairs around a central atom, these pairs can be lone or shared (bond) pairs
Electron Pair Repulsion
In order of mutual repulsion: Lone Pair: Lone Pair. Lone Pair: Shared Pair. Shared Pair: Shared Pair.
Molecular Shapes: Linear and Trigonal Planar
Two shared pairs
Bond Angle: 180 degrees Space here for a pic :)
Three shared pairs
Bond Angle: 120 degrees Space here for a pic :)
Molecular Shapes: Tetrahedral and Trigonal Bipyram
Four shared pairs
Bond Angle: 109.5 degrees space here for a pic :)
Five shared pairs
Bond Angles: 90 degrees and 120 degrees Space for a pic :)
Molecular Shapes: Octahedral
Six shared pairs
Bond Angles: 90 degrees and 180 degrees
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Molecular Shapes with lone pairs
Non bonded pairs have a higher force of repulsion than bonding pairs, this is because bonding pairs feel the pull of two nuclei, not one, so less repulsion
Two shared pairs, two lone pairs
Bond Angle: 105 degrees Pic space :)
Molecular shapes with lone pairs
Three shared pairs, one lone pair
Bond Angle: 107 degrees
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Working out molecular shape
Best shown by example, here i'm using CH(4)
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1. How many electrons around central atom (how many in outer shell) = C = 4
2. Add one electron for each bond formed/ How many electrons shared by H? = 4
3. Allow for ion charge, if any. + 1 charge, deduct electron, opposite, add one = N/a
4. Divide by two to find total number of electron pairs around central atom = 4 + 4/2 = 4
5. Work out how many lone pairs, if there are 4 electron pairs but three bonds there must be a lone pair = 0 here So total number of electron pairs = 4 = Tetrahedral
Intermolecular Forces: Hydrogen Bonding
Intermolecular forces= The forces of attraction between molecules. Hydrogen Bonding:
Some covalent molecules are polar because of the difference in electronegativity of atoms
Hydrogen bonding occurs between H atoms that a covalently bonded to either N, O or F atoms, (three most electronegative elements)
This results in a very polar covalent bond
The H atom of one molecule is strongly attracted to the electronegative element of another molecule
This is the strongest type of intermolecular force, but weaker than an ionic/covalent bond
Examples: H(2)O, NH(3) and HF
Hydrogen Bonding in Water: How it affects its prop
High Boiling Point: Due to high number of hydrogen bonds each molecule can have relative to its low molecular mass = strong bonds. Also due to H(2)O's high bond number (4), so it is harder to break the connection.
Structure: Compact and dense, due to hydrogen bonding pulling water molecules together, results in floating of ice.
Floating of Ice: Molecules frozen in place, arranged in rigid lattice structure, large holes, so less molecules in ice than in same volume of water, so ice wil float on water.
Surface Tension: Molecules at surface of the liquid have fewer neighbours and so a greater attraction to those nearby, making surface more difficult to break, so if small object placed on surface, it can remain suspended due to this.
Water as a solvent: The partial charge makes it an excellant solvent, water dissolves many substances with charged particles by surronding them and 'pulling' them into solution, moving towards the atoms of opposite charge. To dissolve in H2O substances must have a net electrical charge like ionic compunds and polar covalent.
Dipole: Dipole Attractions
Formed when there is a large difference in electronegativity between two atoms bonded together in a covalent bond, this causes the electrons to be shared unequally, they and are pulled towardsmost electronegative atom
Generally covalent bonds that have N, O, F, Cl, Br or I atoms. C-H atoms are considered non-polar.
Dipoles always exist in molecules (Van der Waals) as electron movement is random, so charge is bound to be unequal at times. Permanent dipoles form as in normal dipoles, there are mutiple of them and they cancel out, but they don't cancel out in a permanent dipole, this is because of the sig. difference in electronegativity.
Results in a higher boiling point than those of temporay dipoles as it slightly increases the intermolecular bonding strength
Van Der Waals Forces
Molecules that do not have a permanent dipole has only this intermolecular force
They come about as the electrons are in constant motion and at any one point the electron charge cloud may not be symmetrical and so a temporary dipole is produced, other dipoles are then induced. Weakest kind of intermolecular force.
Strength of Van Der Waals depends on:
- Size of molecule
- Number of electrons
- Shape of molecule
None of the intermolecular forces are stronger than a covalent bond, when a liquid is boiled it is the intermolecular forces that need to be broken not the covalent bonds.
States of Matter Basics
Particles packed closely together, ordered patern with strong bonds Atoms only vibrate gently. When heated, heat makes solid particles vibrate till this breaks the bonds and the solid begins to melt. Temp constant once solid begins to melt.
Particles in constant motion, free to pass over each other. Forces of attraction exist, but are weaker than in solid When heated, forces between molecules overcome, so they are more widely seperated, this is called vapourisation, particles now in gaseous state. Temp constant once liquid starts to boil.
Particles free to move randomly, forces are weak, steady temp during reverse process, when bonds are made not broken.
States of Matter: Types of Solid: Metallic
Five solid types: metallic, giant ionic, giant covalent (macromolecular), simple molecular, hydrogen bonded.
High melting and boiling points: due to strong electrostatic attraction, the more free electrons, the higher the melting/boiling point as the attraction is stronger.
Good conductors of heat and electricity: Delocilaised electrons
Pic Space (you may need to draw this in exam)
States Of Matter: Types of Solid: Ionic Lattices
Closely packed ions, cations and anions, force holding ions together is electrostatic attraction.
Hard, brittle solids
High melting and boiling points: due to strong electrostatic forces, the the higher the attraction, the harder to break.
Soluble in polar solvents: water on ions attracts polar water molecules
Good conductors of electricity when molten or in liquid form: for electricity to flow there must be a moving charge, free electrons/freely moving ions. So not possible with solids.
Giant Covalent: Basics and Diamond
Graphite, ( metal form of carbon), silicon and the non-metal form of carbon, (diamond) have giant covalent structures. This is because they have medium electronegativity and the ablity to form 4 covalent bonds, each atom essentially in the centre of a tetrahedron strongly bond by four other atoms. As this is continued throughout the entire lattice, this is why they have such high melting and boiling points.
Diamond (example): Allotrope of carbon, different properties and structure to graphite, no free electrons, due to covalent bond strength and quantity. Hence diamond is hard with a high melting and boiling point and so a weak electrical conductor
Giant Covalent: Graphite
Graphite (example): Unique, carbons arranged in flat, parallel layers, each layer containing million of hexagonally arranged carbon atoms. Each carbon bonded to three in its layer, strong covalent bonds account for high melting point.
Between layers are weak Van Der Wall forces which allow slippage between layers, accounts for use as a lubricant. Delocalised electron, graphite is good electrical conductor.
Governed by weak Van Der Waals holding molecules in the lattice
Example = Iodine
Low melting and boiling points: Van Der Waals easily borken
Non electrical conductors: when as solid or in solution, unless they react with solvent to form ions.
Insoluble in polar solvents, water: Unless molecules themselves are polar. Water molecules attract each other strongly and so don't mix with Iodine ones s attraction is weak. However, of attractions within solvent are about the same as potential solute, they do mix.
Example = Ice
High melting and boiling point: due to hydrogen bonding in comparison to similar compounds.
Less dense as solid than liquid: Solid has free space . As melting happens, H bonds are broken and structure becomes less ordered, molecules pack together = increased density.
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