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Atomic Structure

Electrons

  • Electrons has a -1 charge
  • They whizz around the nucleus in orbitals
  • The mass of an electtron is 1/2000

Protons

  • Have a +1 charge
  • Make up the nucleus along with electrons
  • Have a mass of 1

Nuetrons

  • Have no overall charge, they are said to be Neutral
  • Make up the nucleus along with Protons
  • Have a mass of 1

 

 

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Atomic structure

The mass of an atom is the number of Protons and Neutrons added togther.

 

An ISOTOPE is an atom of the same elements with the same number or Protons and different number of Neutrons

 

Relative atomic mass = Average mass of an atom 

                             1/12th mass of an atom of carbon 12

 

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Mass spectrometer

The mass spectrometer can be used for loads of stuff including:

  • Relative atomic mass
  • Relative molecular mass
  • Realtive isotopic mass
  • Molecular structure.  

(http://www.chemguide.co.uk/analysis/masspec/masspec.GIF)

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Mass spectrometer

1. Vapourisation - The sample is turned into a gas using an electrical heater.

2. Ionisation - The sample is bombarded with high speed electrons that knock off an electron from each atom created ions with a +1 charge. The sample is ionised so that later on the sample can be Accelerated, Detected and Deflected.

3. Acceleration - The positive ions are attracted towards a negatively charged plate and are accelerated to high speeds. The lighter the ions are the faster they travel.

4. Deflection - The ions pathways are altered by a magnetic field, Lighter ions have less momentum so are delfected more. The stronger the magnetic field the more the deflection. Any ions with a +2 charger are deflected twice as much.

5. Detection - The magnetic field strength is increased gradually so that ions of increasing mass enter the detector one after another. As they strike the detector they accept electrons and loser their charge and create a current that is proportional to their abundance.

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Electronic structure

(http://chemwiki.ucdavis.edu/@api/deki/files/4242/=subshells.jpg?size=webview)

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Electronic structure

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Ionisation energy

The first ionisation energy is the amount of energy required to remove 1 mole of electrons from 1 mole of a gaseous element.

We can write equations to show the process:

 

  • The more protons ther are in the nucleus, the more positively charged the nucleus is so the electrons are attracted more strongly to it, so ionisation energy increases.
  •  The further away electrons are from the nucleus the less strongly they are attracted and so ionisation energy decreases.  
  • As shielding increases, the ionisation energy decreases as electrons are further away from the nucleus and are attracted less strongly.
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Ionisation energy decreases down Group 2

  •  Each element in group 2 has an extra electron shell to the previous. This therefore means there is more shielding so ionisation energy decreases.
  • The outer electrons are also further away from the nucleus (atomic radii increases) which mean electrons arent as strongly attracted to the nucleus, alos decreasing ionisation energy.

(http://www.creative-chemistry.org.uk/alevel/module1/images/trends2chart.gif)

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Ionisation energy increases across a period

  • Ionisation energy increases across a period because there are an increases number of protons which increases the attraction of electrons to the nucleus, increasing ionisation energy.
  • There is the same/similar shielding but more electrons, decreaseing atomic radii, and therefore increasing ionisation energy.

(http://www.creative-chemistry.org.uk/alevel/module1/images/trends6chart.gif)

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The drop between groups 2&3 and groups 5&6

The drop between groups 2&3:

Aluminuim has an extra electron in its 3p orbital (magnesium has electrons in the 3s orbital). The 3p orbital is at a lower energy level than the 3s orbital and is therefore removed easier, decreasing ionisation energy.

The drop between groups 5&6:

The shielding is identical in Phosphorus and Sulfur, however in Sulfur there is a lone pair of electrons (electrons fill up singally before pairing up). This lone pair of electrons repel each other meaning one becomes further away from the nucleus and is therefore easier to remove and redusces the ionisation energy.

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Ionic Bonding

An ionic bond occurs when electrons are transferred from one atom to another resulting in ions with electrostatic forced of attraction between them due to opposite charges.

We show ionic bonds using a dot and cross diagram:

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The behaviour of ionic compounds

1. Ionic compounds conduct electricity only when they are molten or dissolved. This is because in a liquid the ions are free to move and are able to carry the charge whereas in a solid they are fixed.

 

2. Ionic compounds have high melting points.Giant ionic lattices are held togther by strong forces of electrostatic attraction. (NaCl is an example of a giant ionic lattice).

 

3. Ionic compounds tend to dissolve in water. This is because water is polar, so the +ve parts of the water attract the -ve ions, and the -ve parts of the water attract the +ve ions. This break the ionic comound down.

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Covalent Bonding

Covalent bonds are the sharing of a pair of electrons.

We can again show covalent bonding using dot and cross diagrams:

 

 

 

 

T melt or boil a simple covalent compound, you only need to overcome the Van der Waals or hydrogen bonds that hold the molecules together. You don't need to break the covalent bonds between atoms.Covalent coumpunds have relatively low melting and boiling points.

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Graphite and Diamond

 

Graphite:

  • The carbon atoms in Graphite are arranged in sheets of flat hexagons. The weak van Der Waals forces between the layers means that the sheets can slide over each other, making it perfect to use in pencils 
  •  Each carbon forms 3 covalent bonds, and the spare electron in its outer shell acts as a delocalised electron. Beacuse of this delocalised electron, Graphite can conduct electricity.
  • The layers are quite far apart compared to the length pf the covalent bonds which makes Graphote low density. It is used to make stong, lightweight sports equipment.
  • Graphite has a very high melting point due to strong covalent bonds in the hexagon sheets (it sublimes at over 3900K)
  • Graphite is insoluble in any solvent becasue the covalent bonds are too strong to break.
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Graphite and Diamond

Diamond:

  • Each carbon atom is covalently bonded to 4 other carbon atoms, its the hardest known substance on the planet.
  • It has a very high melting point due to strong covalent bonds (it sublimes at over 3800K)
  • It is a good thermal conductor becuase vibrations travel well through its stiff lattice.
  • It cant conduct electricity because all of its electrons are held in bonds, they arent delocalised.
  • Like Graphite, diamond wont dissolve in any solvent.
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Dative Covalent Bonding

Dative Covalent bonding is where one atom provides a pair of electrons to share.

We can show Dative Covalent bonding by drawing dot and cross diagrams.

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Shapes of molecules

  • The shape of a molecule depends on the number of pairs of electrons in the central atom.
  • Electrons are all negatively charged so the repel each other.
  • Lone pairs of electrons repel more than bonding pairs.
  • The greatest angles are between lone pairs of electrons.
  • Angles between bonding pairs are pften reduced by lone-pair repulsion.

Lone pair/lone-pair bond angles are the biggets

Lone-pair/bonding angles are the second biggest

bonding pair/bonding pair angles are the smallest

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Polarisation and Intermolecular forces

Electronegativity is the ability of an atom to attractbelectrons from a covalent bond.

The greater the difference in electronegativities, the more polar the bond will be.

  • The covalent bonds in Diatomic gases are non-polar because the electronegativities are the same. 
  • Some elements (Like carbon & hydrogen) have pretty similiar electronegativities and so their bonds are essentially non-polar
  • In a polar bond, the difference in electronegativities causes a dipole.

A Dipole is a difference in charge between 2 atoms caused by a shift in electron density in the bond. 

δ+ and δ- charges on Polar molecules cause weak electrostatic forces of attraction between molecues.

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Polarisation and Intermolecular forces

Intermolecular forces are forces between molecules.

There are 3 types of intermolecular forces:

  • Van der Waals - the weakest type and are found in everything
  • Dipole-Dipole - caused by polar molecules
  • Hydrogen Bonding - The strongest type. Between Nitrogen, Oxygen, Fluorine and Hydrogen.

Van der Waals are responsible for holding Iodine molecules together in a lattice. Iodine atoms are held together by strong covalent bonds to form I2 molecules. The molecules are then held together by weak Van der Waals.

  • Larger molecules have larger electron clouds meaning stronger Van der Waals.
  • Molecules with a greater surface area have larger Van der Waals.
  • Liquids with stronger Van der Waals have higher boiling points because to boil the liquid you need to overcome these forces.
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Polarisation and Intermolecular forces

Hydrogen Bonding:

  • Flourine, Nitrogen and Oxygen are very electronegative. They draw the bonding electrons away from the Hydrogen atom so the bond is polarised.Hydrogen has such a high charge density becuase it is so small that it forms weak bonds with the lone pair of electrons on the Fluorine, Nitrogen or Oxygen atoms of the other molecule.
  • Substances with hydrogen bonding have higher melting and boiling points than other similar molecules because of the extra energy needed to break the Hydrogen bonds.
  • Ice has more Hydrogen bonds than liquid water, and hydrogen bonds are relatively long. So the water molecules in ice are further apart on average, making ice less dense than water.
  • 
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Metallic Bonding

Metal elements exist as giant metallic lattice structures.

  • The outermost shell of electrons of a metal atom is delocalised and are free to move about the metal. This leaves a positive metal ion.
  • The positive metal ions are attracted to the delocalised elctrons forming forces of electrostatic attraction. This is metallic bonding.
  • The number of delocalised electrons per atom affects the melting point. The more there are so the stronger the bonding will be. The size of the metal ion and the lattice structure also affect the melting point.
  • As there are no bonds holding specific ions together, the metal ions can slide over each other when the strcutre is pulled, so metals are said to be malleable and ductile.
  • Delocalised electrons can pass kinetic energy to each other making metals good thermal conductors.
  • Metals are good electrical conductors becuause the delocalised electrons are able to move and carry a charge.
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Periodicity

Period 3: 

  • Atomic radii decreases across a period because the number of protons increases across the period but there is the same amount of shielding. This means the nucleus has a greater +ve charge and electrons are pulled in more strongly.
  • Ionisation energy increases across a period due to the decreasing atomic radii. Because the atomic radii is decreasing, the electrons are being attracted to the nucleus more strongly making them more difficult to remove.

Melting and boiling points:

  • Na, Mg and Al are metals. Their boiling points increase across the period becuase their metallic bonds gets stronger. This is because there are more delocalised electrons and a decreasing atomic radius. This attracts the ions togther more strongly.
  • Si is has a macromolecular structure - many stong covalent bonds hold all its atoms in place . A lot of energy is required to break these bonds so Si has high melting a boiling points.
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Periodicity

  • P(P4), S(S8) and Cl(Cl2) all have molecular structures. Their melting and boiling points depend on the strength of the Van der Waals forces between their molecules. These foces are weak so these substances have relativelt low melting and boiling points.
  • S is the biggest molecule out of the 3 so has the most Van der Waals forces and therefore the greatest melting and boiling points.
  • Ar has very low melting and boiling points as it exists as single atoms, resulting in very weak Van der Waals forces.

(http://www.docbrown.info/page07/periodgraphs/period3mptsbpts.gif)

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Isomers

Structual Isomers have different structural formulars but the same Molecular formular. 

 

Chain Isomers- Chain isomers have different arrangements of the carbon skeleton. Some are straight chains and other are branched.

Positional Isomers Have the same skeleton and the same atoms attached. The difference is that the atoms are attached to a different carbon atom.

Functional group isomer - Isomers have the same atoms arranged into different functional groups.

 

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Alkanes and Petroleum

Alkanes are saturated Hydrocarbons (they are made of Carbon and Hyrdrogen only). They have a general formular CnH2n+2. They are saturated because they have no double bonds.

  • Petroleum (crude oil) is made of mostly alkanes. They range from smallish alkanes to massive alkanes. Crude oil is seperated into fractions using Fractional Distillation.

1. Crude oil is vapourised at about 350 degrees. The largest hydrocarbons dont vapourise at all because their boiling points are too high. Instead they just run to the bottom and form a qooey residue.

2. The vapourised crude oil goes into the fractionating column. As they go up the column, it gets cooler. Because of the different chain lengths, each fraction condenses at different temperatures.

3. The smallest chain condense at the lower temperatures, at the top of the column. The larger ones condense at the bottom.

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Cracking

Light fractions like Petrol and Naptha are much more useful than heavier fractions like Butimen. To meet the demand for lighter fractions, heavier fractions can be cracked. Cracking is breaking long-chain hydrocarbons into smaller hydrocarbons (which could include alkenes).

Thermal cracking:

  • High temperatures and high pressures (1000 degrees and 70 atm) 
  • Produces a lot of alkenes (used to make polymers such as polyethene)

Catalytic cracking:

  • Makes mostly motor fuels and aromatic hydrocarbons
  • Uses a Zeolite catalyst (hydrated aluminosilicate), at a slight pressure and high temperature (about 450 degrees)
  • Catalyst cuts the cost of the reaction, the reaction can be done at lower temperature and it also speess up the reaction.
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Alkanes as fuels

Complete combustion of alkanes produces Carbon Dioxide and Water:

Incomplete combustion of alkanes prodcues Carbon monoxide and Water:

Nitrogen Oxides are produced when there is enough energy in the air for Nitrogen and Oxygen to combine:

This happens in a petrol engine at high temperatures when a spark ignites the fuel. They are contributors of acid rain and smog

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Alkanes as fuels

Sulphur doixide also contributes to acid rain. It is produced from the Sulfur impurities in crude oil. It combines with Water vapour and Oxygen in the air to form Sulfuric acid:

Carbon particles make asthma worse and cause cancer. They also contribute to global dimming.

Unburned Hydrocarbons cause smog.

Water vapour, Methane and Carbon Dioxide are all greenhouse gases.

To remove Sulfur, some chimneys now contain CaO  which absorbs the SO2. This produces Gypsum (CaSO4) which is used for plaster. This is called 'Flue gas desulfurisation'.

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Catalytic converters

Catalytic converters reduce the amount of CO, NO and unburned Hydrocarbons released from the exhaust of cars into the atmosphere.

It is a honeycomb structure that is coated with Platinum and Rhodium, these act as catalysts. The honeycomb structure provides a larger surface area.

The polluting gases pass over the catalysts and react with each other to form less harmful substances:

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Comments

:) PurpleJaguar (: - Team GR

These cards are really helpful, thanks a lot 

:)

Asawer

First flash card, Protons, bullet point #2 is neutrons, not electrons in the nucleus :)

Asawer

These flash cards are excellent, good job

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