CHEM 1

?
  • Created by: Franklin
  • Created on: 02-05-14 12:12

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
1 of 40

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


2 of 40

Mass spectrometer diagram

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

3 of 40

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)

4 of 40

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
5 of 40

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
6 of 40

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)

7 of 40

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)

8 of 40

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
9 of 40

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

10 of 40

n = CV triangle

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

11 of 40

Balanced Equations

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

Atom Economy

12 of 40

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
13 of 40

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
14 of 40

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
15 of 40

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
16 of 40

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)

17 of 40

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

18 of 40

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
19 of 40

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 
20 of 40

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
21 of 40

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
22 of 40

Shapes of molecules

23 of 40

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

24 of 40

Periodic Table Orbitals

25 of 40

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)
26 of 40

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
27 of 40

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)

28 of 40

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)

29 of 40

Different Formula

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

30 of 40

Nomenclature

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

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

31 of 40

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
32 of 40

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

33 of 40

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
34 of 40

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
35 of 40

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
36 of 40

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
37 of 40

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

38 of 40

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
39 of 40

Greenhouse effect diagram

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

40 of 40

Comments

:) PurpleJaguar (: - Team GR

Report

This is helpful

Thanks a lot

kafia12

Report

thank you so much :D helped a lot

Similar Chemistry resources:

See all Chemistry resources »See all Bonding & shapes resources »