1.5 Concentrations of solutions
- Concentration is how much of a substance is dissolved up in a known volume
- When we make a solution, the concentration depends on the amount (or volume) of solute and the final volume of the solution
- Concentration is always for moldm-3, sometimes abbreviated to M
- 1000 cm cubed= 1 dm cubed
- 1. Write the balanced symbol equation for the reaction (use state symbols)
- 2. Find the ratio of reactants
- 3. Calculate the number of moles of the acid used to neutralise the alkali
- 4. Use the ratio from step 2 to work out the number of moles of alkali used
- 5. Calculate the concentration of the alkali using the answer from step 4 and the concentration triangle to help you
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2.4 Electronic structure: Sub-shells and
- Electron shells are split up into sub-shells. They are labelled s,p,d and f. n= 1 shell has a s sub-shell, n=2 has s and p, n=3 has s, p, and d and n=4 has s,p,d and f
- The sub-shells are divided into atomic orbitals. An electron in a given orbital can be found in a particular region of space around the nucleus.
- Each orbital can hold a max of two electrons. Electrons have a spin, each electron spins at the same rate clockwise (↑) or anticlockwise (↓). Electrons can only occupy the same spin if they have opposite, or paired, spins
- The arrangement of electrons in shells and orbitals is called the electronic configuration of an atom.
- The orbitals are filled in order of increasing energy, so the lowest possible energy arrangement is produced. The orbitals are first occupied singly by electrons, then they start to pair up.
- The chemical properties of an element are decided by the electrons in the incomplete outer shells
- When sub-shells are fully occupied by electrons, the arrangement is closed shell
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2.5 Atoms and Ions
- Ionisation is the process of an atom gaining energy in order to lose an electron, and become a positive ion
- When one electron is pulled out of an atom, the energy required is called the first ionisation enthalpy. Defined as 'the energy needed to remove one electron from every atom in one mole of a isolated gaseous compounds of an element'. General equation- X (g) -> X+(g) + e-
- As you go across the period or down a group, first ionisation enthalpy increases. Nuclear charge becomes more positive (as protons are being added), so electrons become more tightly held and it gets harder to pull one away
- Going down group 7, ionisation enthalpies decrease, as the increasing electron shells shield the outermost electron from the nucleus
- Successive enthalpies increase- remaining electrons will be attracted more strongly
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5.1 Ions in Solids and Solutions
- In solids, ions are held by their opposite electrical charges. The +ve cations attract several -ve anions, forming a giant ionic lattice.
- The crystals of some solids have water molecules fitted into the lattice. The crystals are called hydrated crystals, and the water inside is called the water of crystallisation.
- When hydrated crystals are heated, the water removed is driven off as steam, leaving an anhydrous solid.
- Many ionic substances dissolve in water. The ions become surrounded by water molecules and spread out
- Water is a bent, polar molecule, the oxygen is delta - and the hydrogens are delta +. The hydrogen (+ve) end is attracted to negative ions, and the oxygen end (-ve) is attracted to positive ions. Each ion is surrounded by its own sphere of water molecules- hydration
- Ions in solution behave independently. Reactions often involve two types of ions.
- Some of the ions do not take part in the reaction- spectator ions.
- An ionic equation only shows the ions that take part in a reaction- do not show spectator ions
- State symbols are important to identify if a reaction involves ionic precipitation.
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5.3 Bonds Between Molecules: Temporary and Permane
- The forces between molecules, causing them to be attracted to each other, are called intermolecular bonds. When a solid melts/boils, the intermolecular bonds break, but the covalent bonds within the molecule remain intact
- A dipole is a molecule with a negative end and a positive end. When a molecule has a dipole, it is polarised. Ways a molecule can be polarised-
- Permanent dipole- a molecule has two atoms bonded which have different electronegativities, one atom attracts e- more than the other eg. HCl. Polar molecules
- Instantaneous dipole- No permanent dipole, but a instantaneous (or temporary) dipole can arise. At some instant, the electrons within the molecule shift to one end of the molecule, so one end has a more -ve charge. Normally only lasts for a short time, but if any other molecules nearby, the instantaneous dipole may affect them and produce induced dipoles.
- Induced dipole- when an unpolarised molecule gets a dipole induced when it meets a dipole.
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5.3 Bonds Between Molecules: Temporary and Permane
- Three types of intermolecular bond arise from attractive forces between dipoles-
- Instantaneous dipole- induced dipole bond (ID-ID)- Weakest type, act between all molecules. Occurs between Cl2 molecules, alkanes and nobel gases most clearly- have no other intermolecular bonds. The polarisation is not permanent, and in gases and liquids, the dipoles change position as the molecules move around.
- Bigger atoms with more electrons are more polarisable- bigger induced dipoles -> higher boiling point. Longer chain molecules- more electrons- instantaneous dipoles more likely
- Permanent dipole- permanent dipole bond (PD-PD)- A molecular dipole depends on electronegatively differences and shape of molecule. There needs to be two definite ends of the molecule with electronegativity differences. Stronger than ID-ID. In HCl. Attraction between opposite charges and repulsion between like charges occur in liquids, but attractions are stronger. In solids, the attractions hold the molecules
- Permanent dipole- induced dipole (ID-PD)- Permanent dipole induces a dipole in another molecule. Between HCl and Cl2
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9.1 Oxidation and Reduction
- The above reaction is made up of two half-reactions, can be described by two half-equations.
- The copper is losing electrons, and the oxygen is gaining in the other. Oxidation is about loss of electrons. Reduction is gain of electrons
- A reduction/oxidation reaction is a redox reaction
- An oxidising agent removes electrons from something else, and a reducing agent gives electrons to something else
- Each atom in a molecule or ion is assigned an oxidation state to show how much it has been oxidised or reduced. Atoms in elements have oxidation state 0. In simple ions the oxidation state is same as the charge on the ion
- Since compounds have no overall charge, the oxidation state must add up to 0
- Oxidation states are used in systematic naming of compounds and ions existing in more than one oxidation state eg. iron(II) oxide (FeO) and iron(III) oxide (Fe2O3). They are also used to clarify the names of oxyanions (negative oxygen cont. ions) eg. chlorate(I) (ClO-) and chlorate(V) (ClO3-)
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11.4 The p block
- The halogens y have 7 electrons in the outer shell, and are the most reactive group of non-metals, and do not occur naturally in the element form.
- They are all found in compounds, often as halide ions (a single negatively charged ion, eg Br-). They all occur as diatomic molecules- atoms linked with a covalent bond
- In compounds, a halogen atom gains stability by gaining an electron from a metal to form a halide ion (ionically bonded compound) or by sharing an electron from another atom (covalently bonded compound). In both cases, the halogen has a oxidation state of -1
- The physical states moves from gas to liquid to solid as you go down group 7. This change is caused by an increase in the strength of intermolecular bonds of the diatomic molecule (instantaneous dipole- induced dipole attraction)
- Fluorine is the most volatile halogen- this is because it has the smallest molecules, with the least number of electrons, so has the weakest intermolecular forces
- The molecules get bigger as the group is descended -> more electrons per molecule -> stronger intermolecular forces -> change in physical state
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11.4 The p block 2
- The halogens remove electrons from other elements- they are oxidising agents. The elements at the top of the group are the most reactive and are the strongest oxidising agents
- Fluorine (at the top) has the smallest atoms, so the attraction between the nucleus and the extra electron is needed is strong, so the atom gains the e- more readily to become a negative ion
- With many metals, the halogens react to form halide ions eg, KBr and CaCl2 (oxidation state -1)
- With non-metals (and some p-block and transition metals), halogens form molecular compounds containing covalent bonds (oxidation state -1)
- If two halide ions are mixed in solution, a displacement reaction takes place. For example-
- Cl2 (aq) + 2K+I+ (aq) -> 2K+Cl- + I2 (aq)
- These are also redox reactions- chlorine is reduced (0 to -1) and iodine is oxidised (-1 to 0). They are non-reversible, iodine will not liberate chlorine as it is less reactive
- Oxidation of halides to halogen- 2X- -> X2 + 2e-
- Silver halides are precipitated when a solution of silver ions is added to a solution containing Cl-, Br- or I-. These are precipitation reactions- Ag+ (aq) + X- (aq) -> AgX (s)
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- Halogenoalkanes are alkanes chains with a halogen atom attached. They are named after the parent alkane, with the halogen atom added as a prefix eg. 1-cloropropane.The prefixes are listed in alphabetic order
- The carbon-halogen bond is slightly polar, but not enough to affect properties. All halogenoalkanes are immiscible with water. The bigger the halogen and the more halogen atoms, the higher the boiling point
- Homolytic fission forms radicals. Happens when halogenoalkanes reach the atmosphere, where they are exposed to strong uv radiation. This is how Cl radicals deplete ozone
- Heterolytic fission is more common, halogenoalkanes are reacted in a polar solvent, forming a negative halide ion and a positive carbocation.
- As you go down the group, the C-Hal bond gets weaker, so compounds get more reactive
- In a substitution reaction a functional group in a compound is replaced by another group. Hydrolysis a reaction involving breaking a bond with water
- Halogenoalkanes perform subsitution reactions with nucleophiles, molecules or -ve ions with a lone pair of electrons that can be donated to a +ve atom, forming a covalent bond
- Curly arrows show movement of electrons
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15.1- 15.2 Greener Industry
- Batch- starting materials put in a vessel to react. When reaction complete- terminated. Process repeated until required amount of product made. Cost effective for small quantities and range of products can be made in same vessel, but is more time consuming, bigger workforce and contamination more likely
- Continuous- Staring materials fed in at one end, and withdrawn at other in continuous flow. Suited to high tonnage work, low contamination risk, smaller workforce and easily controlled/automated, but higher capital cost and not cost effective when run below capacity
- Raw materials-starting materials. Feedstocks- reactants. By-product- secondary product derived from a process
- Fixed costs- not dependant upon yield eg. labour, land, telephone bills. Capital costs- relate to establishing a plant/building. Variable costs- relate to unit of production eg. raw materials, treatment, disposal + distribution
- Percentage yield- used to see if a reaction is economically viable
- Atom economy- used to see how effectively all the reactant atoms are used in a chemical reaction
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CS Halogens in industry
- Fluorine- making PTFE, HCFC's and in toothpaste. So reactive that it's almost impossible to store- reacts with container. Need to generate fluorine and make product required immediately
- Chlorine- making PVC, bleach. Most plants making chlorine have a plant to make bleach on site
- Bromine- medicines, flame retardants. Great care needed when transporting- carried in lead-lined steel tanks, with a metal cage around the tanker for protection in the case of an accident
- Iodine- medicines, human nutrient
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