Periodic table

 (a) describe the Periodic Table in terms of the arrangement of elements: (i) by increasing atomic (proton) number, (ii) in periods showing repeating trends in physical and chemical properties, (iii) in groups having similar physical and chemical properties;            

Elements are arranged in order of increaing atomic number.                                                               A period is a horizonatal row of elements in the periodic table showing trends in properties across a period                                                                                                                                                     A group is the vertical column in the periodic table having similar chemical properties and the same number of outer shell electrons                                                                                                           b) describe periodicity in terms of a repeating  pattern across different periods;  Periodicity is a regular variation of properties of elements with atomic number and position in the periodic table(c) explain that atoms of elements in a group have similar outer shell electron configurations, resulting in similar properties; elements in the same group have atoms with the same number of electrons on their outer shells - similar chemical behaviour this is because elements within the same group have similar electron configurations and the same orbitals 

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Periodicity of physical properties of elements

d) describe and explain the variation of the first ionisation energies of elements shown by: (i) a general increase across a period, in terms of increasing nuclear charge, 

Across a period : general increase                                                                     This is because the number off protons increases  so more attraction is acting on the electrons                                                                                                   -Electrons are added to  the same shell so electron shielding does not change The shells are drawn inwards slightly because of the increased nuclear attraction so the atomic radius decreases                                                                             So across a period more energy is required to remove an electron                   

 There is a noticeable decrease in ionisation energies at the start of each new period because a new shell has been added so there is increase atomic radius and increased electron shielding.                                              

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Periodicity of physical properties of elements

(ii) a decrease down a group in terms of increasing atomic radius and increasing electron shielding outweighing increasing nuclear charge; 

Down a group: ionisation energies decrease                                                           The number of shells increase so atomic radius increases so there is decrease nuclear attraction on the outer electrons                                                             There are more inner shells so electron shielding increases                                  The number of protons incresease but the increased attraction is aoutweighed by the distance and shielding .So less energy is required to remove an electron

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Periodicity of physical properties of elements

(e) for the elements of Periods 2 and 3: (i) describe the variation in electron  configurations, atomic radii, melting points and boiling points,  Group 1 - 4 general increase in boiling points  ,4-5 sharp decrease ,5-8 low boiling points ,same goes for melting point                                                              

 ii) explain variations in melting and boiling points in terms of structure and bonding;The change between group 4 and five is because there is a shift from giant structures to simple molecular structures , strong to weak forces                                                                                                                                              Group 1 - 4 Giant metallic structure, metallic bonding, strong forces between positive ions and negative delocalised electrons                                                                                                 Group 5 (4 and 5 in period 2) - Giant covalent structure, Covalent bonding ,strong forces between atoms                                                                                                                             Group 4 -8 - Simple molecular structures, van der Waals' forces between molecules  Metals : melting and boiling points increase across a period because ionic charge increases so ionic size decreases as the number of outer shell electorns increases so the attraction between ions and electrons increases

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Redox reactions of Group 2 metals

(a) describe the redox reactions of the Group 2 elements Mg  Ba: (i) with oxygen, (ii) with water;

The alkaline earth metals                                                                                        - High melting and boiling poits, light metals with low densities that form colourless compounds                                                                                             -  They have an outer shell that has two electrons so electron configuration ends in s2  - Strong reducing agents so they are oxidised in reactions to form a 2+ ion, reactivity increses down the group because of the decrease in ionisation energies down a group

Redox reaction with oxygen makes an ionic oxide                                                                        2Ca(s) + o2(g) -- 2CaO(s)

React with water to form hydroxides and hydrogen

Ca(s) +2H2O(l) -- Ca(OH)2(aq) + H2 (g)

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Reactions of Group 2 compounds

(c) describe the action of water on oxides of elements in Group 2 and state the approximate pH of any resulting solution;

Group 2 oxides and hydroxides are bases thats are neutralised by acids to form a salt and water   MgO(s)+HCl(aq) -- MgCl2(aq)+H2o(l) for example ,the solid will dissolve , this is the observation

Oxides: react with water to form a metal hydroxide of a pH of 10-12                                              MgO(s)+H2O(l)--Mg(OH)2(aq)

Hydroxides : dissolve in water to form alkaline solution ,solubility increases down the group so the solutions get more alkaline  Ca(OH)2(s)+aq--Ca 2+ (aq)+2OH-

(d) describe the thermal decomposition of the carbonates of elements in Group 2 and the trend in their ease of decomposition

Carbonates decomposed by heat to form a metal oxide and carbon dioxide, they become more difficult to decompose with heat as you move down the group.Thermal decomposition is the breaking up of a chemical substance with heat into at least two chemical substances.

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Reactions of Group 2 compounds

(f) explain the use of Ca(OH)2 in agriculture to neutralise acid soils; the use of Mg(OH)2 in some indigestion tablets as an antacid. 

Because they are alkaline they are used against acidity

Calcium hydroxide used by farmers to neutralise acid soils

Magnesium hydroxide used in milk of magnesia to neutralise stomach acid 

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Group 7 Elements

Candidates should be able to: (a) explain, in terms of van der Waals’ forces, the trend in the boiling points of Cl2, Br2 and I2;

The halogens  have low melting and boiling points and exist as diatomic molecules.                  

As you move down the group there is an increase in van der Waals' because there is an increase in electrons so the boiling points increases down the group and the physical states at room temperature an pressure show a gas-liquid-solid trend.

Have 7 electrons in their outer shell

Their p-subshells with 5

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Redox reactions and trends in reactivity of Group

(b) describe the redox reactions, including ionic equations, of the Group 7 elements Cl2, Br2 and I2 with other halide ions, in the presence of an organic solvent, to illustrate the relative reactivity of Group 7 elements; 

They are strong oxidising agents they gain one electron to form a halide ion                                    1/2Cl2 +  e- ---> Cl-

As they gain an electron the reactivity decreases down the group because the atomic radius increase,Electron shielding increases and the ability to gain an electron decreases.A more reactive halogen will displace a less reactive one in a displacement reaction. A change of colour will show a redox reaction and the oragmic solvent cyclohexane will show what type of ions they are. Chlorine oxidises both Br and I                                                                                          Cl2(aq)+2Br-(aq)--->2Cl-(aq)+Br2(aq) Orange in water and in cyclohexane                                                                                                        Cl2(aq)+2I-(aq)--->2Cl-(aq)+I2(aq) Brown in water and purple in cyclohexane  Br2(aq)+2I-(aq)--->2Br-(aq)+I2(aq) brown in water and purple in cyclohexane

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Redox reactions and trends in reactivity of Group

(c) explain the trend in reactivity of Group 7 elements down the group from the decreasing ease of forming negative ions, in terms of atomic size, shielding and nuclear attraction;

Reactivity decreases down the group because the atomic radius increases and so does the electron shielding and also the ability to gain an electron into the p sub shell

(d) describe the term disproportionation as a reaction in which an element is simultaneously oxidised and reduced,illustrated by: (i) the reaction of chlorine with water as used in water purification,                  Disproportionation is the reation in which an element is both oxidised and reduced. Chlorine reacts with water to form Chloric(I) acid and HCl                    

  Cl2(aq)+H2O(l)--->HCIO(aq)+HCl(aq)                                                                 Chlorine is oxided and reduced

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Redox reactions and trends in reactivity of Group

(ii) the reaction of chlorine with cold, dilute aqueous sodium hydroxide, as used to form bleach, 

Chlorine reacts with sodium hydroxide to form bleach                                             Cl2(aq)+2NaOH(aq)--->NaCl(aq)+NaClO(aq)+H2O(l) THis is another disproportionation reaction of chlorine

(e) interpret and make predictions from the chemical and physical properties of the Group 7 elements and their compounds;

Covalent diatomic molecule X2 ,Simple molecular structure with weak van der Waals' between molecules,Oxidising agent that forms a halide ion when they gain electron and reactivity decreases down the group.

Fluorine is very reactive, too reactive

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Redox reactions and trends in reactivity of Group

(f) contrast the benefits of chlorine use in water treatment (killing bacteria) with associated risks (hazards of toxic chlorine gas and possible risks from formation of chlorinated hydrocarbons);

Chlorine is used to kill bacteria in water but it forms chlorinated hydrocarbons that are suspected of causing cancer

(g) describe the precipitation reactions, including ionic equations, of the aqueous anions Cl–, Br– and I–with aqueous silver ions, followed by aqueous ammonia;

Add silver nitarate to an aqueous solution of the hallide ions, a siver halide precipate willbe formed. If unsure of the jallide add diluted ammonia and then concentrated                                                                                                       AgCl(s) White precipitate and soluble in dilute ammonia                                     AgBr(s) Cream precipitate and soluble in concentrated ammonia                       AgI(s) Yellow Precipitate and insolube in both         

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Electrons. Bonding and Structure

(a) Define the terms first ionisation energy and successive ionisation energy

First ionisation energy : the energy required to remove one electron from each atom in one mole of gaseous atoms to form one mole of gaseous 1+ ions.                                                                         Succesive ionisation energies : the measure of energy required to remove each electron in turn.

(b) Explain that ionisation energies are influenced by nuclear charge, electron shielding and the distance of the outermost electron from the nucleus; 

Factors affecting Ionisation energies :                                                                                                  Atomic radius : the greater the radius the smaller the nuclear attaraction experienced by the outer electrons Nuclear Charge : The greater the nuclear charge the greater the attractive force on the outer electrons. Electron shielding : The repulsion between electrons in different inner shells.Inner shells repel outer shell electrons ,the greater the number of inner shells present,the larger the shielding , the smaller the nuclear atrraction experienced by outer electrons

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Electrons. Bonding and Structure

(c) predict from successive ionisation energies of an element: (i) the number of electrons in each shell of an atom, (ii) the group of the element; 

Li(g) ---> Li+ (g) + e-

Each succesive ionisation energy is larger than before because as each electron is removed there is less repulsion so the shell is drawn closer to nucleaus to the atomic radius decreases and the nuclear attraction increases soo more energy is required.

You can tell the number of electrons in a shelll by the massive increase in ionisation energy when there is a change from one shell to another. This shows that the electron was closer to the nucleus and had less shielding. You can tell the grup of the element by countion the ionisation number before moving to the next shell

(d) state the number of electrons that can fill the first four shells

A shell is a group of atomic orbitals with the same principal quantum number , n. each shell holds up to 2n squared electrons  when n =1 its the 1st shell and has 2 electrons when n =2 its the 2nd shell and has 8 electrons when n = 3 its the 3rd shell and has 18 electrons when n = 4 its the 4th shell and and 32 electrons

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Electrons. Bonding and Structure

(e) describe an orbital as a region that can hold up to two electrons, with opposite spins;

An atomic orbital is a region within an atom that can  hold up to two electrons with opposite spins

(f) describe the shapes of s and p orbitals; 

S - orbitals have a spherical shape P - orbitals have a 3D dumb-bell shape (g) state the number of :electrons that occupy the s , p and d subshells and the orbitals S orbital - each shell contains one s0 1x 2 = 2 electrons in each shell P orbital - each shell from n=2 upwards contains 3 so 3x2 = 6 electrons D orbital - each shel from n=3 upwards contains 5 so 5 x 2 = 10 electrons F orbital - each shell from n=4 upwards  contains 7 s0 7 x 2 = 14 electrons A sub shell is a group  of the same type of atomic orbitals within a shell The sub shells have different energy levels that increase in order. The electrons are added one at a time withe the lowest available energy filled first and each energy level must be full before starting the next Remember singly before pairing.

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Electrons. Bonding and Structure

i) deduce the electron configurations of: (i) atoms, given the atomic number, up to Z = 36, (ii) ions, given the atomic number and ionic charge, limited to s and p blocks up to Z= 36

The 4s subshell strarts to fillbefore the 3d is filled

(j) classify the elements into s, p and d blocks

Example oxygen is the 4th element in the 2p block so it ends in 2p4


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Electrons. Bonding and Structure

(a) describe the term ionic boding

Ionic bonding is the the electrostatic attaraction between opppositely charged ions (metal and non metal). Electrons are transferred from the metal to the non metal forming a positive metal ion and a negative non metal ion.

This forms giant ionic lattices with each ion being surrounded by oppositely charged ions and the attraction being from all directions.

(b) construct ‘dot-and-cross’ diagrams, to describe ionic bonding; 

(c) predict ionic charge from the position of an element in the Periodic Table; Atoms of metals groups 1-3 lose electrons to form positive ions with the electron configuration of the previous noble gas Atoms of non metals groups 5-7 gain electrons to form negative ions with the electron configuration of the next noble gas Atoms of Be,B,C,Si do not form ions as a lot energy is needed to transfer the outer shell electron to form ions

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Electrons. Bonding and Structure

(d) state the formulae for the following ions: NO3, CO32–, SO42– and NH4+

ammonium NH4+                                                                                                                                  Hydroxide OH-   Nitrate NO3- Nitrite NO2- Hydrogencarbonate HCO3- Carbonate CO32- Sulfate SO42- Sulfite SO32- Dichromate Cr2O72- Phosphate PO43-

(e) describe the term covalent bond as a shared pair of electrons;

Covalent bond is a bond formed by a shared pair of electrons                                                            Between two non metals

f) construct ‘dot-and-cross’ diagrams to describe: (i) single covalent bonding, eg as in H2, 

Single covalent bond is when each atom contributes one electron to a covalent bond

(ii) multiple covalent bonding, eg as in O2, N2 and CO2

Multiple covalent bonds is when there is more than one pair of electrons in the covalent bond.

(iii) dative covalent (coordinate) bonding, eg as in NH4+

dative covalent bond is a shared pair of electrons which has been provided by one of the bonding atoms only such as in NH4+ or in H3O+


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Electrons. Bonding and Structure

(g) explain that the shape of a simple molecule is determined by repulsion between electron pairs surrounding a central atom;

The shape of a molecule or ion is determined by the number of electron pairs in thr outer shell surrounding the central atom because each electron pair repels outher electron pairs and they push each other as far as possible

(h) state that lone pairs of electrons repel more than bonded pairs;

Lone pairs are more electron dense than bonded pairs so they repel more ,they reduce the bond angle by 2.5 degrees

(i) explain the shapes of, and bond angles in, molecules and ions with up to six electron pairs (including lone pairs) surrounding a central atom, eg as in: (i) BF3 (trigonal planar), (ii) CH4 and NH4+ (tetrahedral), (iii) SF6 (octahedral), (iv) NH3 (pyramidal),  (v) H2O (non-linear), (vi) CO2 (linear); 

3 electron pairs = trigonal planar 120 degres BF3 4 electron pairs = tetrahedral 109.5 Degrees CH4 6 electron pairs = octahedral 90 degrees SF6 3 electron pairs + 1 lone pair = Pyramidal 107 degrees NH3 2 electron pairs + 2 lone pairs = Non linear 104.5 degrees H2o 2 double bonds = linear 180 degrees CO2

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Electrons. Bonding and Structure

(k) describe the term electronegativity as the ability of an atom to attract the bonding electrons in a covalent bond;

Electronegativity is a meausre of the attraction of a bonded atom for the pair of electrons in a covalent bond                                                                                                                                       One of the atoms is more attracted to the electrons than the other.

(l) explain that a permanent dipole may arise when covalently-bonded atoms have different electronegativities, resulting in a polar bond;

This small difference in charge across a bond that results in the electronegativities of the bonded atoms is a permanent dipole. A polar covalent bond has a permanent dipole

molecules that are symmetrical may have dipole which cancel out.

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Electrons. Bonding and Structure

(m) describe intermolecular forces based on permanent dipoles, as in hydrogen chloride, and induced dipoles (van der Waals’ forces), as in the noble gases; 

An intermolecular force is an atrractive force between neighbouring molecules 1) permanent dipole- dipole interactions is when two polar molecules with permanent dipoles attracts each other to form a weak dipole-dipoel force 2)van der Waals' between induced dipoles in neighbouring molecules When electrons are moving in a shell the distribution of charge in electron shells is unbalanced ,at any moment there will be an instantaneous dipole across the molecule and this induces a dipole in neighbouring molecules which induces further dipoles .The induced dipoles attarct causing van der Waals' V.d.Waals' increases with an increase in electrons  because the greater the number of electrons the greater the induced dipoles and the greater the van der waals' formed and the boiling point increases as well

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Electrons. Bonding and Structure

n) describe hydrogen bonding, including the role of a lone pair, between molecules containing –OH and –NH groups, ie as in H2O, NH3 and analogous molecules 3) Hydrogen Bonds are very strong dipole-dipole attraction between an electron defficient hydrogen on one molecule and a lone pair on the other. Such as in water. (o) describe and explain the anomalous properties of H2O resulting from hydrogen bonding, eg: (i) the density of ice compared with water(ii) its relatively high freezing point and boiling point; Ice is less dense than water - ice has an open lattice with hydrogen bonds holding the water molecules apart so when it melts the hydrogen bonds collapse allowing the water molecules to move closer together Water has high melting and boiling points : Hydrogen bonds have to be overcome in order to melt or boil water and a lot of energy is required to do this as they are relatively strong (p) describe metallic bonding as the attraction of positive ions to delocalised electrons;  Metallic bonding is the electrostatic attraction between positive metal ions aand delocalised electrons

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Electrons. Bonding and Structure

(q) describe structures as: (i) giant ionic lattices, with strong ionic bonding, ie as in NaCl,

 Because each ion is surrounded by oppositely charged ions and they all attract each other forming a giant ionic lattice                                                                                                               Ionic compounds : Have high melting and boiling points because they are solid at room temperature and a lot of energy is required to break the electrostatic attraction between the oppositely charged ions.                                                                                                                   Electrical Conductivity : In a solid the ions are in a fixed position and cannot moved so it is a non conductor but when it is liquid or in solution the lattice breaks down the ions are free to move now it conducts.                                                                                                                                Solubility : Ionic lattice dissolves in polar solvents such as water . The postive ions are attacted to the oxygen and the negative ions are attarcted to the hygrogen.

(ii) giant covalent lattices, ie as in diamond 

Simple molecular structures : made of simple molecules , molecules are held together by weak forces and each molecule has atoms bonded by covalent bonds.                                                    Melting and boiling points : low melting and boiling points because not a lot of energy i required to break the weak van der waals                                                                                                      Electrical conductivity : No charged particles so do not conduct electricity                                      Solubility : soluble in non polar subsatnces such as hexan as van der waals' are formed between the structure and the solvent and this weakens the structure.

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Electrons. Bonding and Structure

Giant covalent structures: 3D structure of atoms bonded together by covalent bonds                      High melting and boiling points because a lot of energy is requiredto break the covalent bonds in a lattice                                                                                                                                         Electrical coonductivity: no free charged particles so they conduct electricity                                  Solubility : insoluble in polar and non polar solvents because the covalent bonds are too strong to be broken by both.

Diamond :Tetrahedral structure held together by a covalent bond that has poor conductivity because ethere are no delocalised electrons as all outer-shell electrons are used for covalent bonds. It is hard because the tetrahdral shape allows external forces to spread through the  shape.

Graphite : strong hexagonal layer structure but with weak van der Waals' forces between the layers. that has good conductivity because there are delocalised electrons between the layers and the electrons can move parallel to the layers when voltage is applied. It is soft becuse although the bonding within each layer is strong the weak forces between the layers allow them to slide easily

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Electrons. Bonding and Structure

(iii) giant metallic lattices, 

Delocalised electrons throught the metallic structure thats can move so you can not tell which electron came from which ionbut the charges must balance.

High melting and boiling points :: a lot of energy is required to break the attraction between the positive ions and the negative electrons to help dislodge the ions from their rigid positions               Electrical conductivity: Good conductors because the electrons can move in the lattice                      Malleability and ductility: Ductile means that they can be drawn out or stretched and malleable means they can be hammered into shape because the electrons are delocalised electrons

Alloys : mixtures of metals, positive ions of one can replace the others'

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