Chemistry Revision Cards

The core principles of chemistry

Elements: An element= a substance that cannot be broken down into two or more substances.                                                  An atom= smallest uncharged particle in an element                   A compound= made up of two or more elements chemically bonded together

The periodic table: metals= in left and middle, conduct electricity when molten & solid, malleable, form +ve cations in compounds, more reactive down group                                                 non-metals= form -ve anions, covalent bonds with other non-metals, more reactive up group.

Amount of substance (the mole):

AVAGADRO CONSTANT= 6.02 X 1023 mol-1, one mole is theamount of substance containing the avagadro no. of atoms

The atomic number= number of protons in the nucleus of an atom of that element

The relative atomic mass= weighed average mass of an atom divided by 1/12 the mass of a carbon-12 atom

The relative molecular mass= weighed average mass of a molecule of an element/compound divided by 1/12 mass of a carbon-12 atom.

Mass number= sum of protons & neutrons in the nucleus of an atom of that isotope

Isotopes= are atoms of the same element with the same number of protons, but different number of neutrons.

1) All air pumped out= vacuum                                       2) Sample of element in gaseous form injected into spec. It is bombarded with high-energy electrons from electron gun. Energy from electron gun strips off electrons from atoms forming positive ions.                                                               3) Positive ions accelerated by high electric potential. All ions have same energy, so all ions of the same mass will have the same speed.                                                            4)Ions then deflected by magnetic field. Greater masses deflected less than smaller masses. Smaller masses detected first- depends on m/e ratio.                                                         5) Detector linked to computer which calculates m/e ratio and relative abundance of each positive ion.

So... BOMBARDMENT, ACCELERATION, DEFLECTION & DETECTION

Electronegativity= The extent to which it attracts a pair of electrons in a covalent bond towards itself.

Shielding & Effective Nuclear Charge (ENC):                         Across periods- greater ENC due to same amount of shielding but more protons in nucleus- greater force of attraction on outer electrons. Electrons repel each other as much as possible within an atom, stronger positive forces bring them closer together.                                                       Down groups- weaker force of attraction on outermost electrons due to increased shielding from full shells.                     Effective nuclear charge= net charge on the nucleus, allowing for electrons in orbit around nucleus shielding its full charge.

THERFORE:Atomic radius: across period=decrease, down group=increase

Ionic radius: +ve= smaller than neutral atom- less -ve charge, so greater ENC. -ve= larger than neutral atom

First Ionisation energy: the energy required to remove one electron from the ground state of each atom in a mole of gaseous atoms of that element.  A(g)--> A+(g)+ e-

Second Ionisation energy: the energy required to remove one electron from each ion in a mole of gaseous singly positively charged ions of that element. A+(g)--> A2+(g)+ e-                            Successive ionisation energies: e.g. Al, 13 electrons. Easy to remove first 3 as in outer shell. Big jump to 4 IE as changing shells so decreased shielding and stronger ENC. Closer shells are to nucleus, more energy required to remove the electrons.                                                 Variation across & down periods and groups:General increase in ENC across groups so noble gases have highest ionisation energies- don't need to lose electrons. Group 1 elements have lowest ionisation energies- want to lose electron to get a full shell

First electron affinity= the energy change when one electron is added to each atom in a mole of neutral gaseous atoms.

A(g)+e- --> A-(g)

Second electron affinity= the energy change when one electron is added to each ion in a mole of singly negatively charged gaseous ions.

A-(g)+e- --> A2-(g)

Values of first electron affinities are usually exothermic (negative)

From page 39

From page 39

From page 39/40

Molecular formula= the number of atoms of  each element in one molecule of the substance

Molar Mass= the mass of 1 mol of the substance g mol-1

Empirical formula=the simplest whole number ratio of the atoms of each element in the substance       

Calculation= step 1: divide each % by RAM of element Step 2: divide results of step 1 by smallest value obtained. Step 3: If step 2 does not produce intergers, multiply until they do.

Ionic Compounds: made up of cations & anions = charged

Solubility of Ionic solids: all grp 1 compounds soluble in water, all nitrates soluble, all carbonates insoluble apart from grp 1 & ammonium carbonate

Rule 1: Write full equation & balance. For next steps write different equations

Rule 2:For dissolved ionic substances, write ions separately

Rule 3: For all solids, liquids & gases write full ionic formula

Rule 4:cross out all spectator ions not involved in the reaction.

Example: Magnesium with dilute hydrochloric acid

Balance equation: Mg(s)+2HCl(aq)--> MgCl2(aq)+H2(g)

Split(aq)into ions:Mg(s)+2H+(aq)+2Cl-(aq)--> Mg2+(aq)+2Cl-(aq)+H2(g) 

Remove Spectator Ions: in this case= 2Cl- Mg(s)+2H+-->Mg2+(aq)+H2(g) 

A mole= the amount of a substance containg the avagadro constant of atoms, molecules or groups of ions.

1 mol of any substance = relative formula mass in (g)

Molar mass= mass of 1 mol of substance

MOLES= MASS / MOLAR MASS

Examples: pg 57

Examples:

Calculate number of molecules:

NUMBER OF MOLECULES=AMOUNT OF SUBSTANCE (moles) X AVAGADRO CONSTANT

Pg 60

Pg 60

Mass & Volume calculations

Mass to mass:step 1- calculate no. of moles of reactant step 2) balance using stochiometry to calculate no. moles of product in relation. Step 3) convert the moles of product to mass

MASS OF A-->MOLES OF A-->MOLES OF B-->MASS OF B

Volume of gas:

Molar volume of a gas is the volume occupied by 1 mol of gas under specified conditions of temp and pressure. = 24000 cm3 mol-1

moles of gas = volume of gas/24000

MASS OF A-->MOLES OF A-->MOLES OF GAS B-->VOLUME OF GAS B

pg 66/67

% yield calculation:

% YIELD= ACTUAL YIELD / THEORETICAL YIELD X 100

Calculation of reaction stochiometry: If masses of both reactants, or product and reactant known stochiometry can be worked out. Basically the ratio of moles.

Examples:

When reaction takes place, ratio not always correct, one reactant used completely (limiting reagent as stops reaction happening anymore) the other reagent is said to be in excess.

A limiting reagent= the substance that determines the theorectical yield of a product in a reactionn                                  Step 1: calculate amount of moles of one reagent and use stochiometry to calculate how much product in moles could be formed from this reagent.                                                 Step 2:calculate moles of other reagent, work out how much product in moles, could be formed from this reagent.                    Step 3:reagent that produces LEAST moles = limiting reagent

Step 4:calculate theoretical yield from least amount

Step 1 & 2= convert to moles, use mole ratio to calculate amount of product. Step 4=conversion of moles to mass/volume.

Units of concentration= mol dm-3

MOLES = CONCENTRATION X VOLUME (solution)

Examples:pg 73

Pg 74

Pg 74

Specific heat capacity=the heat required to increase the temperature of 1g of a substance by 1'C

HEAT REQUIRED= MASS X SPECIFIC HEAT CAPACITY X RISE IN TEMPERATURE

Example:

Exothermic reaction= Chemical energy converted to heat energy and temperature of system (surroundings) rises

Endothermic reaction=Heat energy converted into chemical energy and temperature of system decreases

Enthalpy= 'H' is the chemical energy in the system at a constant pressure that can be converted to heat.

Enthalpy level diagrams: change in enthalpy= ^H=H products - H reactants

IF EXOTHERMIC, PRODUCTS HAVE LESS ENTHALPY THAN REACTANTS (GIVEN IT OFF) SO ^H IS NEGATIVE

IF ENDOTHERMIC, PRODUCTS HAVE MORE ENTHALPY THAN REACTANTS (TAKEN IT IN, SO ^H IS POSITIVE.

Standard Conditions: pressure of 1 atm, temp. 25'C, solutions conc of 1 mol dm-3

The enthalpy change for any reaction is independent of the route taken from reactants to products.

Examples:

^H1=^H2+^H3     OR       ^H3=^H1-^H2

Energy level diagram:

Definitions:                                     measure= kJ mol-1

Standard Enthalpy of... Formation (^Hf): Enthalpy change when 1 mol of substance formed from its elements in their standard states under standard conditions of 1 atm pressure and 298 K (25'C).        Reaction (^Hr):Enthalpy change when number of moles written in reaction react under standard conditions of 1 atm pressure & 298 K Combustion (^Hc):Enthalpy change when 1 mol of substance burned in excess oxygen under standard conditions of 1 atm pressure & 298K Neutralisation (^Hneut):enthalpy change when 1 mol of water is produced by neutralisation of acid by excess base under standard conditions. All solutions at 1 mol dm-3, temp 298K, 1 atm pressure.

Atomisation (^Ha): Enthalpy change when 1 mol of gaseous atoms formed from elements in their standard states under standard conditions 298K and 1 atm pressure.

Reaction: Hess's cycle arrows point up from constituent elements to products & reactants.

Experiment= iron reacting with copper (II) sulfate in polystyrene cup (insulator minimises heat loss). Source of error, some heat lost to surroundings.

Examples:

Combustion: Experiment= calorimetry, spirit burner, heat produced used to warm known vol of water.

All combustion reactions= exothermic ^H= -ve. Arrows from products & reactants go DOWN towards combustion products (CO2, H20)

Examples:

Formation: Heat change in a reaction

Examples: Ethanol- 2C(s)+3H2(g)--> C2H50H(l)

Neutralisation: ionic equation for neut of any acid by base= H+(aq)+OH-(aq)-->H2O(l)

Examples:

Pipette 25cm3 of 1.00moldm-3 acid into polystyrene cup. Record temp. Measure temp of alkali calculate mean of two temps. Measure 25cm3 of alkali using pipette, then add to acid, stir and record highest temp reached. 

Temp rises so reaction= exothermic

Bond enthalpy= enthalpy change when a bond in a gaseous molecule is broken (kJ mol-1)

If a bond enthalpy is worked out for a particular compound, the value obtained is slightly different from the average value.

MEXOBENDO- making bonds is exothermic (-ve) breaking bonds is endothermic (+ve)

Calculation of enthalpy from average bond enthalpies:

Step 1:list all bonds broken. Write down energy to break each bond (+ve). Add to find total energy required.                          Step 2:list all bonds made. Write down energy to make each bond (-ve). Add to find total energy released.                            Step 3:Add two totals to give ^Hr. kJ mol-1

Examples:

Pg 97

pg 97

Pg 97/98

Pg 98

Metallic bonding & properties

Structure:Cations arranged in regular lattice, in layers, surrounded by a sea of electrons that can move through the lattice. They are delocalised throughout the entire structure. Bonds between he positive ions whicih are held in a fixed position by the electrons. Strength depends on metal ion, metallic radius and structure of lattice.                                              Electrical conductivity: flow of charge. Metals conduct when molten and solid. Delocalised electrons mobile & moves through lattice of metal ions. Electric current= flow of electrons.             Thermal conductivity: Good heat conductors- free moving electrons, pass kinetic energy along piece of metal.              Malleability: Can be hammered into shapes. Layer of metal ions can slide over another layer- because always electrons between layers preventing repulsion between +ve ions. (d block metals hard- lots of electrons binding layers together.)

Ionisation energy: ENC smaller than for elements that occur later in period as metals have 1,2 or 3 electrons in outer orbits= held less firmly & 1st ionisation is smaller

Electronegativity: lowest electronegativity values on periodic table, decreases down group & increases across period= caesium lowest.

Formation of +ve ions: low values of ionisation energies & electronegativities mean it's easy for metals to lode electrons in bonding- forming +ve ions.                                     Reactions:

Acids: Zn(s)+2H+--> Zn2+(aq)+H2(g)

Water: Ca(s)+2H2O(aq)-->Ca2+(aq)+20H-+H2(g)

Less reactive metal: Fe(S)+Cu2+-->Fe2+(aq)+Cu(s)

Ionic bond=form between atoms with very different electronegativites. Neutral overall, but consist of positive & negative ions.                                               Cation/Anion=Cation= +ve, formed when metal loses electrons. Anion=-ve, formed when non metal gains electrons                  Strength of ionic bonds: depends on charge and radii of ions. Larger difference in charge=stronger ionic bond. Smaller radii=stronger bonds. Strong bonds=higher melting temp.                                                                  Ionic radius:Down group=Increases as more orbitals being added. From grp 5 in one period to grp 3 in the next period: N3-,O2-,F-,Mg2+ & Al3+ all have same electron config. =ISOELECTRONIC. Radii decrease from N3-(largest) to Al3+ (smallest)- due to same number of electrons, greater ENC experienced, so held more tightly.                     

Evidence for existence of ions: Only conduct electricity when molten (lattice breaks down, freeing electrons) - ions normally held in strong lattice, no delocalised electrons. ELECTROLYSIS= cations move to anode (+ve->-ve), anions moce to cathode (-ve->+ve)

 e.g. Copper (II) chromate. PLATINUM ELECTRODE, IN HCl

 CuCrO4- copper is oxidised forming Cu2+ ions chromate is reduced to Cr042-. This is seen through colour change around the cathode & anode.  At the anode, solution turns yellow as -ve chromate ions attracted to +ve terminal. At the cathode, solution turns blue green as +ve copper ions attracted to -ve terminal.

Melting temp & solubility: very high melting temps due to strong forces of attraction between ions. Many ionic compounds soluble in water- when hydrated, energy made breaks bonds between +ve & -ve ions so they dissolve.  

Ion pair formation: eg sodium chloride--> enthalpy change= sum of energies for endothermic removal of an electron from sodium & exothermic addition of the electron to chlorine.

Na(g)-->Na+(g)+e-: ^H=+494 kJmol-1

Cl(g)+e- --> Cl-(g): ^H=-364 kJmol-1

SO... Na(g)+ Cl(g)-->Na+(g)+Cl-(g): ^H=+130 kJmol-1

The lattice energy (^Hlatt)= the energy change when 1 mol of the solid is formed from its constituent gasoues ions that start infintely far apart.

Enthalpy of atomisation (^Ha)=is the enthalpy change when one mol of gaseous atoms is made from an element in its standard state.

Born Haber cycle:of NaCl

Step 1:enthalpy of atomisation of Na(s). Turn solid Na to gas.

Step 2:enthalpy of atomisation of Cl(g).Spilt Cl molecules into atoms

Step 3:First ionisation energy of Na. Remove an electron from each gaseous sodium atom

Step 4:First electron affinity of Cl. Add an electron to each gaseous chlorine atom.

Step 5: The lattice energy= bringing the ions together in an ionic lattice.

Energy level diagrams:

Factors that affect lattice energy: magnitude of charges, sum of radii of anion & cation,arrangement of ions in lattice, relative sizes of ions, extent of covalency (explained on next cards)

Examples:

Ions take up an arrangement that maximises the lattice energy. The ions are arranged in the position of minimal potential energy, so as little repulsion occurs between the ions as possible.

Polarisation of the anion: +ve cation exerts force on electrons in -ve anion. If significant, the anion is polarised.

Cations with high charge, small radius have high polarising power- Smaller Mg 2+ ion more polarising than Ca2+ ion because its smaller. Mg2+ ion more polarising than Na+ ion because it has a greater charge & smaller radius.

Polarising power measured by charge density of cation. Anion with high charge, large radius easily polarised. I- ion more easily polarised than smaller Cl- ion. S2- ion more easliy polarised than Cl- ion because of greater charge & larger radius.

If either type of ion is present, the ionic bond will have a degree of covlency about it as the anion will be significantly polarised. (This causes Born Haber/experimental lattice energy to be greater than theoretcial value as it's not 100% ionic.)  

Electrons involved in covalent bonding= VALENCE ELECTRONS

Covalent bond= overlap between two orbitals resulting in sharing of pair of electrons by two atoms. Happens between atoms with similar electronegativities. Sigma bond formed between orbitals.

Expansion of the octet: explains why phosphorous can form 5 covalent bonds. has three outer electrons, but has 5 empty 3d- orbitals with similar energy to a 3s-orbital. So one 3s electron promoted into empty 3d-orbital= 3d1, now has 5 unpaired electrons, so can form 5 covalent bonds.

Covalent bond strength: depends on sum of atomic radii of two bonded atoms. Small atoms=stronger bonds & no. of electron pairs being shared- double bond (2 pairs shared) stronger than single bond (1 pair shared).  

Single bond form:

sigma bond between two atoms, sharing of one pair of electrons.

Double bond form:sharing of two electron pairs. Orbital in one atom overlaps antoher forming sigma bond, but p orbital (at right angles to sigma bond) overlaps sideways with another p orbital forming a pi bond.

The presence of a pi bond makes rotation around the double bond impossible.

Lone Pairs= a pair of valence electrons that is not used in bonding.

A dative covalent bond= a covalent bond in which both electrons in the shared pair are provided by one atom

E.g. Dative bonds with nitrogen:

Ammonium ion (NH4+) -lone pair on nitrogen atom used to form dative covalent bond with a H+ ion. Once dative covalent bond formed, all four N-H bonds are identical.

Dative bonds with chlorine: PCl6- ion- solid Phosphorus pentachloride is not PCl5 but ionic compound of PCl4+ & PCl6-. One PCl5 molecule uses one of its lone pairs to form a dative covalent bond with an empty orbital of a phosphorus atom in another PCl5 molecule. On heating dative bond breaks = two PCl5 molecules.

Pg 122

A homologous series= a series of compounds containing the same functional group & have same general formula. They only differ by CH2.

A functional group=a small group of atoms or a single halogen atom that gives the compounds in the series particular characteristics.

Hydrocarbons= compounds containing hydrogen & carbon only. 

Alkanes= a hydrocarbon in which all C-C bonds are single. Example of a homologous series.

General formula: CnH2n+2

Naming alkanes:

1)Identify longest unbranched Carbon chain to establish stem name (meth, eth,prop,but,pent,hex,hept,oct,non,dec)

2)Name any substituent groups bonded to carbon atom (if more than 1, must be alphabetical)- methyl, ethyl, propyl, butyl.

3)Identify position of each substituent by number. If two identical groups on same carbon= 'di' (more= tri, tetra)

Cycloalkane= a saturated hydrocarbon with the carbon atoms in a ring.

Alkenes= a hydrocarbon with one double bond between carbon atoms. (An unsaturated compound). Identified like alkanes only with 'ene' at end. General formula: CnH2n

Structural formula= must include any double bonds in molecule.

Example:

Displayed/Full structural formula= must show all atoms separately and all bonds between them.

Example:

Skeletal formula= carbon atoms not drawn, straight line= single bond, double line=double bond.

Example:

Isomers= compounds with the same molecular formula but different structural formulae.

Structural isomerism:SAME MOLECULAR FORMULA, DIFFERENT MOLECULAR STRUCTURE. (carbon chain, positional & functional group)

Carbon chain: only differ in length of longest unbranched carbon chain e.g. methylpropane & butane both have 4 carbon atoms, but butane arranged in straight line, whereas (mepro) has one alkyl group.                                                        Positional Isomerism: same functional group, different location on carbon framework. e.g. Propan-1-ol and Propan-2-ol.

Geometric Isomers= are isomers that only differ in the spatial arrangement of atoms in the planar part of the molecule or above and below a ring.

Why do they exist? Because there is restricted rotation around the double bond. One isomer cannot spontaneously convert into the other unless sufficient energy is supplied to break the double bond.

E-Z notation:priority assigned to atoms attached to C=C carbon atoms(higher atomic no.=higher priority).

If the 2 higher priority atoms are on opposite sides of the double bond the isomer is labelled 'E', if they're on the same side labelled 'Z'--> on ze zame zide.

Alkanes= CnH2n+2

Properties:Boiling temp= increases as number of electrons increases- stronger intermolecular forces (instantaneous induced dipole-induced dipole) require more energy to separate molecules. BRANCHED ALKANES: lower boiling temps- they dont pack so well together=less energy to break apart.

Solubility: Insoluble in water, but dissolve in each other- crude oil=lots of alkanes dissolved in each other.

Density: less dense than water, increase as molar mass of alkane increases.

Chemical reactions:

With halogens: in presence of UV light react with Cl and Br. Hydrgoen atoms replaced with halogen atoms= Free radical substitution

free radical= an uncharged species with an unpaired electron used to form a covalent bond.

Substitution reaction= one atom or group is replaced by another atom or group. There are always 2 reactants & 2 products.

Reaction mechanism: half headed arrow used to show movement of a single electron.

Mechanisms:

Initation:UV light energy splits chlorine molecule into radicals-causes sigma bond to break homolytically- Cl atoms both receive one electron.

Propagation: Cl radicals VERY REACTIVE. When Cl radical collides with methane molecule with enough energy, hydrogen atom removed to form methyl radical (CH3*) and molecule of Hydrogen chloride. This then causes a chain reaction to occur as different radicals are produced and collide.

Termination:chain reaction broken when 2 radicals collide.   CH3*+CH3*-->C2H6(ethane) or Cl*+Cl*-->Cl2 or CH3*+Cl2*-->CH3Cl

Alkenes= unsaturated compounds CnH2n

Properties:lower melting points than alkanes- double bond does not allow molecules to pack together as effectively. Insoluble in water-form hydrogen bonds with water molecules.

Chemical reactions: much more reactive, most reactions are addition. Pi bond breaks leaving sigma bond between 2 carbon atoms, two new atoms/groups add on to each carbon atom (1 to each).

Addition reaction=occurs when two substances react together to form a single substance.                                              Reaction with hydrogen: reagent=hydrogen, conditions=nickel catalyst, 150'C, product=ethane, reaction type=addition.     Reaction with halogens: reagent=bromine, conditions=mix at room temp., product= 1,2-dibromoethane, reaction type= electrophillic addition.

An electrophile=a species that bonds to an electron rich site in a molecule. It gets a pair of electrons from that site & forms a new covalent bond.                                                    Test for alkenes: shake with bromine water- brown colour of bromine disappears- addition reaction.                                 Reaction with hydrogen halides: reagent= hydrogen bromide, conditions=mix gases at room temp., product=bromomethane, reaction type= electrophillic addition.                                                                                 Addition of hydrogen bromide: uneven breaking of H-Br bond= heterolytic

Step 1:

Step 2:

Addition to asymmetric alkenes- the hydrogen atom goes to the carbon which already has the most hydrogen atoms directly attached. The remaining carbon gains a positive charge = carbocation and attracts a negative molecule.

Step 1:

Step 2:

A polymer= a chain of covalently bonded molecules

Polymer production:usual procedure is to cause free radicals to be formed that initiate polymerisation.

Example:poly(ethene) heated under pressre of 1000atm in presence of trace of oxygen. (Oxygen produces radicals)

Monomer:CH2=CH2

Polymer name:POLYETHENE

Uses:Low density plastic bags, water pipes, bottles, buckets.

Uses: polypropene- ropes, containers for boiling temps, polychloroethene(window frames,guttering,pipes)polytetrafluoroethene(non stick coatings, low friction bearings)

Disposal:addition polymers resistant to chemical & biological attack. Last a long time, but can litter & fill up disposal sites, also if combusted can produce toxic fumes.

Summary of alkene reactions:

Hazards and risks:toxicity, irritation, corrosive, flammable, carcinogenic.

Pg 150