CHEM 1

Atomic Structure

Nucleons are particles inside the nucleus: protons and neutrons

  • Mass of nucleon: 1.67 x 10-27 kg
  • Mass of electron: 9.11 x 10-31 kg
  • Protons and electrons have charge of magnitude 1.6 x10-19 C (+ and - respectively)
  • The maximum number of electrons in the nth shell = 2n
  • Atomic no. = no. of protons

(http://images.tutorvista.com/content/atoms-molecules/relative-atomic-mass-formula.jpeg)

  • Mass number = total relative mass of protons and neutrons
  • Avogradro's constant: 6.0223 , is the number of atoms in 12g of C-12
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Mass Spectrometer

Vaporisation

  • The sample is vaporised so that it can be investigated in its gaseous state

Ionisation

  • An electron is knocked off each of the atoms in a sample by collision with high-energy electrons, fired from an electron gun
  • Definition of first ionisation energy: The enthalpy change that occurs when an electron is removed from 1 mole of gaseous atoms to leave 1 mole of gaseous +1 ions in standard conditions: 100kPa, 298K

Acceleration

  • Ions are accelerated by an electric field

Deflection

  • The ions are then deflected by a magnetic field according to their masses. The lighter they are, the more they are deflected.The amount of deflection also depends on the number of positive charges on the ion


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Mass spectrometer diagram

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

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Electron Shells

Althougth not learnt a GCSE, shells in fact contain orbitals. And an outer shell can hold more than 8 electrons.

  • The formula 2n2 indicates the maximum no. of electrons the nth shell can hold
  • The shells are divided into subshells
  • The subshells contain orbitals s,p,d or f
  • Orbitals can only contain a maximum of 2 electrons
  • orbitals fill singly before pairing
  • p-orbitals can hold up to 2 electrons but always come in groups of 3
  • The 1st shell a 1s orbital
  • The 2nd shell contains 1s orbital and up to 3 p orbitals
  • The 3rd shell can contain up to 1s orbital, 3 p orbitals and 5 d orbitals
  • The 4th shell can contain up to 1s orbital, 3 p orbitals, 5d orbitals and 7f orbitals
  • 4s orbital is slightly lower in energy than 3d orbital
  • Calcium – 1s2, 2s2, 2p6, 3s1, 4s2 (note 4s filled before 3d)

  • Iridium - 1s2, 2s2, 2p6, 3s2, 3p6, 4s2, 3d10, 4p6, 5s2, 4d10, 5p6, 6s2, 4d14, 5d7 (over lapping orbitals)

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Ionisation Energy

Ionisation energy is the energy required to remove 1 mole of electrons from 1 mole of gaseous atoms to produce one mole of gaseous +1 ions

(http://scienceaid.co.uk/chemistry/fundamental/images/ionisation.jpg)

  • Electrons are removed from a gaseous sample by fired high energy electrons from an electron gun at it
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IE Trends

There are many trends and deviations which you could be tested on so take note and make sure you know why.

Trends across a period

  • General increase due to increasing nuclear charge making it harder to remove an electron
  • Mg to Al there is a decrease in IE because the outer electron is in a 3p orbital which is higher in energy
  • P to S there is a decrease in IE because of a shared pair of electrons which causes repulsion

Trend down a group

  • General decrease down the group due to increased shielding
  • Electron is in a higher energy level which is further aways from the nucleus
  • Less attraction between protons and electrons
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IE Table

Ionsisation Energy Table (http://faculty.sdmiramar.edu/fgarces/zCourse/All_Year/Ch100_OL/aMy_FileLec/04OL_LecNotes_Ch100/03_AtomsElements/306_PeriodicTable/306_pic/ionization2.gif)

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Ideal gas equation

(http://mrdchemgwiki.wikispaces.com/file/view/pv-nrt.jpg/226501032/pv-nrt.jpg)

Pressure, Pa

Volume, m3

Temperature, K

R= 8.31 (gas constant)

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Empirical and molecular formulae

Empirical Formula

  • Use n = mass/Mr to calculate number of no. of moles of each element
  • Divide each one by the smallest no. of moles present(http://content.tutorvista.com/chemistry_11/content/us/class11chemistry/chapter16/images/img163.gif)

Relative Molecular Formula

  • To find RMF divide mass of relative molecular formula by mass of empirical formula
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Moles in Solution

  • The concentration is how much solute is in a known volume of solution.
  • Concentration is measured in moles per cubic decimetre, 1 mol dm-3
  • 1mol dm-3 means that there is 1 mole of solute in a litre/decimetre of solution

  • To make a 1mol soution you add solvent to 1 mole of solute until you have 1 litre of solution

  • n = CV

  • n = m/Mr

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n = CV triangle

(http://scienceaid.co.uk/chemistry/applied/images/moletriangle.png)

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Balanced Equations

There must be the same no. of moles of each element before and after the reation has taken place/

Atom Economy

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Ionic bonding

Bonds between atoms always involve their valence (outer) electrons

  • Noble gases have a full outer shell and are very unreactive
  • When atoms bond they share or transfer electrons to achieve a more stable arrangement
  • There are 3 types of strong chemical bonds: covalent, ionic and metallic

Ionic Bonding

  • Occurs between metals and non-metals
  • Metals have 1,2, or 3 electrons in their outer shell so the easiest way for them to gain a full outer shell is to lose these electrons
  • Electrons are transferred from the metal to the non-metal
  • The metal loses an electron and is now a positive ion
  • The non-metal has gained an electron and is now a negative ion
  • The two oppositely charged ions are attracted to each other by strong electrostatic forces of attraction
  • Charges of ions balance to form a neutral molecule
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Ionic bonding 2

Properties of ionic compounds

  • Always solid at room temperature
  • Giant ionic structures result in high melting points
  • Can conduct electricity when molten or aqueous as the ions are free to move and act has charge carriers
  • They are brittle due to contact with like charge when given a blow
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Covalent Bonding

Non-metals need to gain electrons to a achieve a full outer shell. A covalent bond is a shared pair of electrons and usually occurs between non-metals

  • Atoms are held together by the electrostatic attraction between both nucluei and shared pair of electrons
  • In a double bond 4 electrons are shared (like O2)
  • The atoms within a covalent bond are strong bonded, however the molecules are not strongly attracted

Properties of covalently- bonded substances

  • Low melting point - due to weak intermolecular forces
  • Poor conductors of electricity - no charged particles to carry current

Dative covalent bonding

  • One atom provides both the electrons - a lone pair
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Metallic Bonding

In a metal element the outer electrons are no longer associated with the metal ion. They are delocalised.

  • A metallic bonding consists of a lattice of metal ions and a sea of delocalised electrons
  • There are electrostatic forces of attraction between the positive metal ions and the delocalised electrons

Strength of metals

  • The charge on the ion: greater charge on the metal ion and more delocalised electrons so stronger electrostatic forces of attraction
  • The size of the ion: The smaller the ion, the closer the electrons are to ther nucleus, the stronger the attractive forces between them

Properties of metals

  • Strong: metallic bonding extends throughout the solid, there are no individual bonds to break
  • Malleable: when a metal is beaten it is still in the same environment, attraction remains
  • High melting point: strong electrostatic forces
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Electronegativity

Electronegativety is the power of an atom to pull the electron density towards itself in a covalent bond

Electronegativity depends on:

  • The nuclear charge
  • The distance between the nucleus and the outer shell electrons
  • The shielding of the nucleus by inner electron shells

Trends in electronegativity

  • As we go up a group electronegativity increases because there is less shielding
  • As we go across a group electronegativity also increases because nuclear charge increases, shielding stays the same

(http://www.chemguide.co.uk/atoms/bonding/pteneg.GIF)

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Polarity

Polarity is the unequal sharing of electrons in a covalent bond and it is a property of the bond.

  • In diatomic molecules like halogens, the electron pair is shared equally because each atom pulls the electron density equally. These bonds are non-polar
  • In a covalent bond with atoms of different electronegativity, the electron density is unequally shared, resulting in partial charges on each of the atoms. These bonds are polar. E.g. Hydrogen chloride

(http://www.chemguide.co.uk/atoms/bonding/ab2.GIF)polar

      (http://www.chemguide.co.uk/atoms/bonding/ab1.GIF)non-polar

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Intermolecular forces

Moles and atoms in a substance are attracted to one an other by intermolecular forces

van der Waals

  • These are the weakest type of intermolecular force and occurs between all molecules
  • The electrons distribution is changing every second creating temporary dipoles
  •  The instaneous dipole induces dipoles in nearby molecule
  • They act in addition to any other intermolecular force
  • More electrons induce a larger instantaneous dipole

Dipole-dipole forces

  • Resultant of all the polar bonds in a molecule may result in a dipole moment
  • The polar bonds of a molecule may cancel to produce no dipole moment or reinforce. This depends on the shape of the molecule
  • Dipole-dipole forces act between molecules that have a permanent dipole.
  • Molecules with permanent dipoles will orientate themselves to attract
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Intermolecular forces - Hydrogen bonding

A hydrogen atom sandwiched between 2 very electronegative atoms.

Conditions

  • Either O, N or F covalently bonded to a hydrogen atom

Process

  • The electronegative atom will cause a permanent dipole between itself and the H atom
  • The H is delta + and the electronegative atom is delta -
  • The delta - hydrogen is electron-deficient
  • The hydrogen is attracted to the lone pair of electrons of the electronegative atom of another nearby molecule

That why ammonia and water have high boiling points

  • When water freezes the hydrogen bonds hold the molecules in fixed positions, resulting in a 3d structure
  • In this structure the molecules are less packed 
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Crystals

Crystals are solids held together by forces of attraction

  • They have a regular arrangement
  • Held together by ionic, metallic or covalent bonds or intermolecular forces
  • The strength of attraction between molecules affects the physical properties of the crystal
  • Crystals can be ionic, metallic, molecular or macromolecular

Ionic Crystals

  • Ionic compounds have strong forces of attraction between oppositely-charged ions
  • They extend throughout the structure

Metallic crystals

  • Metals exist as a lattice of positive ions embedded in a sea of delocalised electrons

Molecular crystals

  • Have low boiling/melting points because of weak vdW forces. IMF are weaker than covalent bonds
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Crystals (cont.)

Macromolecular crystals

  • Covalent bonds extend throughout the structure
  • They have high melting/boiling points

Diamond

  • Pure carbon
  • Covalent bonding between every carbon atom
  • Each carbon forms 4 covalent bonds with other carbon atoms
  • Bond angle 109.5, tetrahedral
  • +3700 K melting point

Graphite

  • Pure carbon
  • Two sorts of bonding: strong covalent and vdW
  • Each carbon forms 3 covalent bonds with another carbon atom, trigonal planar
  • Spare delocalised electron in p orbital allow graphite to conduct electricity
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Shapes of molecules

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Shapes of molecules (cont.)

Method for identifiying shape

1.       Identify group number of central atom and add to this the number of bonds it makes. This equals n

2.       For negative charges add bonds and for positive charges remove bonds

3.       Now divide the total by 2

4.       This number will indicate the parental shape e.g. 4 indicates tetrahedral

5.       Now compare this number with the number of bonds the central atom makes. The difference is the number of lone pairs

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Periodic Table Orbitals

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Periodicity of Period 3

Period 3:  Na, Mg, Al, Si, P, S, Cl

  • Elements in Groups 1,2 and 3 have giant metallic structures
  • In Group 4, Si forms 4 covalent bonds, classed as a semi-metal
  • Elements in Groups 5,6 and 7 either form ionic or covalent bonds
  • Ar is unreactive

Melting/boiling points

  • The giant structures in groups 1,2,3 have high melting boiling points
  • Melting point of the metal increases across the period because of a greater charge on the metal ion and more delocalised electrons
  • The molecular structures to the right tend to have low melting and boiling points
  • Melting points of the molecular structures depend on the size of the van der Waals forces     S8>P4>Cl2 (Silicon has much higher bp than these due to giant structure)
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Period 3 Trends

Atomic radii

  • Is a periodicity property because on a radii graph we can see it decreases across the period and there is a jump to the start of the next
  • Decreases because of increasing nuclear charge with the same shielding effect 

1st Ionisation Energy

  • General increase due to increasing nuclear charge (making it harder to remove an electron) with the same shielding effect
  • Down the group -  IE decreases due to greater shielding from the electrons and valence electron is further from the nucleus in a higher energy level
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Period 3 - Ionisation energy detail

Drop Mg ---> Al is because of electron in 3p orbital in Mg which is easier to remove than 3s. Drop S ----> P due to paired electrons in a 3p orbital

(http://www.docbrown.info/page07/periodgraphs/period3_1stIE.gif)

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Organic Chemistry

Organic chemistry is the chemistry  of carbon compounds

  • Carbon can form rings and long chains
  • Carbon has 4 electrons in its outer shell
  • C-C bonds are relatively strong and polar
  • In all stable carbon compounds carbon forms 4 bonds and has 8 electrons in its outer shell

Display Formula

(http://www.bbc.co.uk/schools/gcsebitesize/science/images/butane_chem_struc.gif)

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Different Formula

(http://www.swotrevision.com/pages/alevel/chemistry/images/img_31.GIF)

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Nomenclature

The root indicates how many carbons the longest carbon chain consists of

(http://www.wissensdrang.com/media/oc03.gif)

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Nomenclature (cont.)

  • functional group is an atom or group of atoms which gives the whole molecule its properties i.e. the reactive bit.
  • A family of molecules all with the same functional group, but increasing numbers of caron atoms is called a homologous series.
  • They can be represented by a general formula which uses algebra to tell you how to work out the molecular formula of any molecule in the family.

Prefixes

  • methyl - CH3ethyl - C2H5propyl - C3H8butyl - C4H10...
  • With multiple functional groups they are written in alphabetical order
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Isomerism

Isomerism is when molecules have the same molecular formula, but a different structure

  • Positional Isomerism - functional group is in different places around the carbon skeleton. The carbon skeleton itself remains unchanged.

  • Chain Isomerism - the arrangement pf the carbon chain is different e.g. branching
  • Function Group Isomerism - the functional group is different

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Alkanes

Alkanes are saturated hydrocarbons - they only contain single c-c bonds

General formula: CnH2n+2

  • Alkanes are almost non-polar due to hydrogen and carbon having similar electronegativity
  • Only forces are weak van der Waals
  • Longer-chained hydrocarbons have higher boiling points due to greater van der Waals
  • Branched hydrocarbons have lower mp/bp as they can't pack together as closely so and van der Waals aren't so effective
  • Alkanes are insouble in water because h-bonds are stronger than the van der Waals
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Fractional Distillation of crude oil

Crude oil is a mixture of mostly alkanes

Fractions (ascending temp): Gases, petrol, naptha,kerosene, gas oil, fuel oil

  • Crude oil is heated and the different fractions are collected
  • Each fraction is a mixture of hydrocarbons of similar chain length

Process

  • Crude oil is heated in a furnace
  • A mixture of liquid and vapours pass into the tower that are cooler at the top than bottom
  • Vapours pass up through a series of trays until one is cool enough to condense it
  • Short-chained hydrocarbons are at the top where it is cooler because they have a low bp
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Industrial Cracking

Naptha fraction is in high demand by the chemical industry and for petrol

  • Naptha is a short-chain hydrocarbon and there is usually a lot of long-chained hydrocarbons and less short-chained ones in crude oil.
  • To produce more short-chain hydrocarbons which are more valuable, longer chained materials are cracked.
  • When longer fractions are cracked they can produce alkenes as well as alkanes
  • Alkenes are used as chemical feedstock

Thermal Cracking

  • Alkanes are heated at high pressure
  • C-C bonds break and electron goes to another carbon, producing a free radicals
  • High proportion of alkenes are produced

Catalytic Cracking

  • Low temp and pressure, zeolite (Al2O3/ SiO) used which has honeycomb structure with large SA, acidic
  • Mostly gas products with short chains
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Combustion of alkanes

  • Alkanes are quite unreactive

Complete combustion equations

(http://www.chemguide.co.uk/organicprops/alkanes/c3h8-1.gif)

(http://www.chemguide.co.uk/organicprops/alkanes/c4h10-1.gif)

Pollution

  • CO is produced as a result of incomplete combustion. It is poisonous
  • NOx (NO, NO2, N2O4) is produced if N2 has enough energy to combine with oxygen. This reacts with water vapour and therfore contributes to acid rain
  • Carbon particles exacerbate asthma

Removing sulfur

  • Sulfur dioxide contributes to acid rain
  • CaO or limstone absorbs sulfur dioxide in a process called flue gas desulfurisation
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Catalytic converters

Internal combustion engines produce a lot of pollutants so vehicles are increasingly being issued with catalytic converters

  • They reduce output of carbon monoxide and nitrogen oxides
  • It is a ceramic honeycomb coated in Rhodium or Platinum which act as catalysts
  • Honeycomb shape provides a large surface area
  • Reaction takes place on the surface of the catalyst

Reaction with Pollutants

2CO + 2NO à N2 +2CO2

 Hydrocarbons + NO à N2 + CO2 + H2O

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Greenhouse effect

Greenhouse effect

  • Radiation from the sun is absorbed and re-radiated as infra-red
  • The greenhouse gases trap the infra-red radiation so it cannot escape
  • The temperature on earth increases
  • The temperature increase leads to more water vapour, which is an even more effective green house gas
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Greenhouse effect diagram

(http://science-at-home.org/wp-content/uploads/2009/09/greenhouse-effect.jpg)

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Comments

:) PurpleJaguar (: - Team GR

This is helpful

Thanks a lot

kafia12

thank you so much :D helped a lot

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