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electrons, the attraction between the other shell electrons and the positive nucleus is weaker. Ionic radius across the period of metals and non metals: Across a period of metals the ionic configuration is the same however there is one more proton each time which creates a stronger force which attracts the outer electron closer making a smaller ion. Ionic radius vs strength of the electrostatic attractions: lonic charge - the higher the charge the stronger the electrostatic attractions. lonic radius if the ions are smaller, then the opposite charge is closer so the electrostatic attraction will be stronger. Attraction increases 2+ Bottom, left has the lowest melting point due to lowest charge. Top right has the highest melting point due to having the most charge. Covalent bonding Covalent bonding includes the sharing pair of electrons between two atoms. HYDROGEN CHLORINE Dative bonds form when the shared electrons in the dative covalent bonds have both come from I atom. Covalent Bond Unpaired HYDROGEN CHLORIDE valence electrons Atom 1 Atom 2 (nonmetal) (nonmetal or metalloid) Sharing of available valence electrons Covalent molecule Chemistry Learner.com Tetrahedral basis of organic chemistry: methane Dr Phil Brown H H с XC H Chemistry H Electron pairs in bonds repel each other in order to be as far apart as possible around the central atom. HA 109.5° C..!! CH H 108.70 pm DIAMOND Tetrahedral arrangement a tetrahedral is a regular triangle based pyramid. In diagrams like above, an ordinary line represents a bond in the plane of the screen or paper. A dotted line shows a bond going away you into the screen or paper, A wedge shows a bond coming out towards you. from GRAPHITE giant covalent structures Graphite Carbon atoms, 4 atoms per bond, High melting point, Diamond: 6个个个个个个 Graphite: Diamond Carbon atoms 3 bonds (per atom) Conduct electricity Made of layers, High melting point, Delocalised electrons, soft, slippery (wif between layers) Diamond has a high melting point due to having strong covalent bonds between 4 carbon atoms which require a lot of to break the bonds, Diamonds can not conduct electricity because there are no free moving electrons. energy Graphite has 3 bonded carbon atoms which form a layer (hexagonal) which have no covalent bonds between them, making the intermolecular forces between them weak, which is the reason graphite is soft and slippery meaning it can conduct electricity. Silicon dioxide: The giant covalent structure is made of silicon and oxygen, each silicon atom is bonded to 4 oxygen atoms making it have a high melting point, + metallic bonding ++++ + + SEA OF DELOCALISED ELECTRONS POSITIVE METAL IONS STRONG ATTRACTIONS BETWEEN POSITIVE IONS + DELOCALISED ELECTRONS (IN) LATTICE OUTER SHELL OF THE METAL ATOMS ELECTRONS IN THE OUTER SHELL Chemistry THE DELOCALISED (MOBILE) ELECTRONS ARE FREE TO MOVE THROUGH THE METAL) ONCE THE OUTER ELECTRON ARE DELOCALISED, THE POSITIVE METAL CENTRES ARE SUSPENDED IN THE SEA OF ELECTRONS and electrons. When you go down group one and group melting points of the elements decrease. +LAYERS WHICH CAN SUIDE OVER EACH OTHER + ELECTRONS WHICH CAN CONDUCT ELECTRICITY SABPIDEL) WHEN METALS ARE CLOSELY PACKED IN A LATTICE, THER OUTER ELECTRONS DISSOCIATE FROM THE POSITIVE CENTRE METALLIC BONDING INVOLVES STRONG ELECTROSTATIC FORCES OF ATTRACTION BETWEEN THE METAL CENTRES AND DELOCALISED ELECTRONS two the Intermolecular forces Electronegativity - the tendency of an atom to attack a bonding pair of electrons. This is measured on the pauling scale. 0.9 0.8 Rb 0.8 Cs 0.7 07 They can conduct electricity because it is malleable and they have delocalised electrons, High melting and boiling points due to strong Why are the noble gases not electrostatic attractions between positive ions Ba 11 1.5 16 16 12 14 Ra 0.9 1.1 4 A21 2.5 Cu Zn Ga Ge 20 19 16 16 18 18 19 22 22 22 15 Cd In Sn Sb Te 25 W 1.9 Ir Pt electronegative: H-Cl 8 9 10 11 12 13 14 15 16 17 group go to the right side of the period. 20 What is the trend of the electronegativity in the periodic table: 23.0 The electronegativity is weaker as you go down the and gets stronger as you The least electronegativity is caesium The most electronegative is fluorine. се 18 1.9 1.9 2.0 O O се P The noble gases are not electronegative. because they have a full outer shell, this means that there is no tendency to attack a bonding pair of electrons. How does it have an effect on bonding in covalent molecules: Cl - Cl се a 3.0 Be 2.8 At Both have the same tendency to attack a bonding pair Cl has slightly more tendency to attack a banding pair, el is more dense therefore more negative Polar and non polar molecule: No electronegativity difference between two atoms leads to a pure. A non polar molecule is a molecule where the electrons are distributed evenly through the structure. Chemistry A small electronegativity difference leads to a polar covalent bond. A polar molecule is a molecule with particularly positive charge in one part of the molecule and a negative charge in the other due to uneven electrons. A large electronegativity difference leads to an ionic bond 00 Cl Cl Nonpolar covalent bond Bonding electrons shared equally between two atoms 8+ HCI 8- Trends in electronegativ Polar covalent bond Bonding electrons shared unequally between two atoms The number of protons in the nucleus - (nuclear charge) more protons will give a stronger attraction to the bonding pair of electrons. The distance from the nucleus closer to the nucleus means there will be a stronger attraction. The amount of screening by inner electrons (shielding) - less shielding gives a stronger attraction. Dipole - dipole forces: Molecules that are polar contain dipoles within the molecule. A dipole is a separation of charges within a covalent molecule. Permanent dipole interactions occur when partial charges form within the molecule due to uneven distribution of electrons. Polar molecules align so that the positive end of the molecule interacts with the negative end of the other molecule. Lots of molecules join in a manner these are weak forces. 8 8+ 8+ 8 CI-H -----H- CI H3C Hydrogen bonding: This is the strongest type of intermolecular force, however it is 10 times weaker than the covalent bond. These occur when the molecule has A very electronegative atom (F, 0, N). that has a lone pair of electrons, A hydrogen atom covalently bonded to this electronegative atom. There is an attraction between the lone pair of electrons on oxygen, nitrogen or fluorine attached to one of these atoms in a different molecule. Lone pair H Hydrogen bond 0: CH 3 H CH3 Hydrogen bonding in methanol Van der waals forces: Ⓒ Chemistry These occur between all molecules, even if there are other intermolecular forces acting on that molecule, these also occur around atoms. Helium has two electrons in its neutral atom, but these electrons could be anywhere at any time. Van der Waals Forces Atom 1 O Step 1 O Atom 2 O Step 2 Atoms are polarized and attract one another If they happen on one side, a temporary dipole is created. This can induce a dipole in a neighbouring atom, causing the two atoms to be temporarily attached to each other, until the electrons are able to move again. Example: Examples of Van der Waals Forces 1. London dispersion forces: Chlorine (Cl₂) 8+ 2. Dipole-dipole interactions : Hydrogen chloride (HCI) H CI 5+ CI NON-POLAR MOLECULE 6+ ChemistryLearner.com 5- INSTANTANEOUS (TEMPORARY) DIPOLE CI INSTANTANEOUS (TEMPORARY) DIPOLE H CI ELECTRON CHARGE CLOUD MOVES 8- ELECTRON CHARGE CLOUD MOVES S+ INTER MOLECULAR ID-ID FORCES INSTANTANEOUS (TEMPORARY) DIPOLE NON-POLAR MOLECULE AS THE INSTANTANEOUS DIPOLE APPROACHES THE NON-POLAR MOLECULE IT INDUCES A DIPOLE IN THE NEIGHBOURING NON-POLAR MOLECULE 6- 8+ INDUCED DIPOLE 8- Identifying what is present: DIPOLE-DIPOLE-IS THERE A LARGE DIFFERENCE IN ELECTRONEGATIVITY? HYDROGEN BONDING-IS H ATTACHED TO O.N.F? VAN DER WAALS ARE THERE ANY ELECTRONS? Equation: 3. 30 moles mass 48= 2mol 24 Example: How many moles are present in 3g of CQH3? Mr 2(12) = 24+6 Mr = 30 Mr X moles Chemistry olmoles Avogadro's number = 6.02x10² Reacting r Mg+ 2HCl₂ 1:2 4x(7 +35.5) 146g 4 mol 23 mass Side reactions Lost in separation, Mgcl2 + H₂ Percentage yield Percentage yield = actual yield X100 Theoretical yield Why wouldn't you get 100% yield? Reversible product, A mole ratio is the ratio between the amount in moles of any two compounds involved in a chemical reaction. mole ratio The amount of moles given x Example: 2H₂ + 0₂ - What is the h2/h20 mole ratio? 2:2 = 1:1 Suppose you had 20 moles of h2 on hand and plenty of 02 how many moles of h20 would u have? H2 X H20 = 20 x 2 = 20 moles H20 IM - Equations: Concentration Cone x vol dm 3 Converting into dm³ X1000 Cm - dm Mass = 2 = /1000 worked out moles Moles given in the equation 27,0₂ moles Cone x vol Dm Cm X1000 dm - m = /1000 Empirical formula Example: Empirical formula - the smallest ratio of atoms in the molecular formula. Molecular formula the actual number of each atom in a molecule, Equation: Molecular mass 1316 1.316 Chemistry Mass of empirical formula Example: Determine the empirical formula for a compound with 15.8% carbon and 84.2% sulphur. с 15.8 15.8 12 =1.316 = S 84.2 84.2 32 =263125 2,63125 1.316 =1.99 (2 sf) CS2 - Limiting reagents A limiting reagent is completely reacted during a reaction, A reagent in excess doesn't fully react has some left at the end. 19 of H₂ reacts with 5g of Cl₂ in the following reaction: H₂ + Cl₂ -2HCl state structu re Period 2: Period 2 elements can form oxides, often burnt in pure oxygen, acid / alkali 1 Reactivity of group 2 and 3 metals with oxygen state S S 11/12 group formula Li₂0 Beo B₂O3 CO₂ NO₂ F₂0 2 GIL Alk = 0.5mol 2 S 2 n 3 GIL 5. 71 = 0.0704mol S GCL 3 S 4 5 7 G S SM Period 3: Period 3 can also form oxides by being burnt with pure oxygen. group 1 form Na₂0 Mgo Al2O3 SiO₂ P₂O5 SO CL₂0 1g H₂ 0.5Mol = excess G in SM S 4 5 6 7 S G struc gil gil gil gel sm sm ac/al SM 9 sm 3 5g Cl₂ 0.0704mol = limited Chemistry Reactivity of metals Reactivity with oxygen: Reactive metals will react vigorously and less reactive metals will burn quickly when powdered. Many metals will react slowly, forming a layer of oxide on the metals surface (tarnishing) even metals like silver this will occur for, Reactions with water: Group I and I group 2 metals will react directly with water to form hydroxides, Group one reacts immediately on contact and fizzing will occur as hydrogen is being produced to form a alkali solution, Group 2 reacts in a similar way but some of the elements like magnesium need to be reacted with steam to form a hydroxide. Reaction with dilute acids: Metals which are more reactive than hydrogen will react in dilute acid to form a salt and hydrogen gas, The salt depends on the dilute acid used for example hydrochloric acid will form a chloride salt, Cedor reactions Oxidation is a loss of electrons Reduction is gain of electrons Oilrig Displacement Displacement reactions are redox reactions, Example: Mgis) + CuSO4 (aq) In displacement reactions the more reactive metal loses electrons and the least reactive atom gains an electron. Mg atoms have lost electrons to form Mg²+ions in 4 MgSO this change can be shown using an half equations, Mg(s) Cu (aq) + 2e 2. 3. 4. 5. 2. MgSO4(aq) + Cuis) 2+ 4 Cu²+ ions in Cuso have gained electrons to form the orange solid, Oxidation numbers rules: I. + e (aq) Cu (s) The oxidation number of an atom in a element. or molecule is always=0 The oxidation state of atoms in compounds are usually the ions charge, The oxidation state of fl in a compound is - The oxidation state of oxygen is -2 expect in peroxides and fo where it is - and + Chlorine oxidation state is unless bonded with oxygen, 6. Hydrogen is + unless bonded with a metal - 7. Group I metals are +1, group 2 are +2 Aluminium is +3 8. q The sum oxidation states in a compound are always o More reactive Reactivity series Potassium Sodium Calcium Magnesium Aluminium Carbon Zine Iron Tin Lead Copper Sliver Gold Platinium Chemistry Less reactive Please stop calling me a careless zebra instead try learning how copper saves gold Transition metals Transition metals often act as a catalyst to form coloured compounds, Less reactive than group I and 2 however they do have high melting points, They have 2 outer shell electrons unless they are specialised. Hard, strong, shiny metals Good conductors of heat/ electricity. High melting/ boiling point, They have a high density. They have more than one ion, Malleable and ductile, They have variable oxidation states, Transition metals are a subset of d block elements, an element must have an incomplete d-subshell either in the atom or common ion for it to be classed a transition element. Electron configuration Energy level 4 3 2 46 4d 2n² Energy level 1 = 2(1)² =2 Energy level 2 = 2(2)²=8 Energy level 3 = 3(3)=18 4p -3d 4s 3p 3s 2p 14 Each energy level is given a number known as the quantum number, Electrons found in each energy level are found in of orbital, regions Energy orbitals can hold a maximum of 2 electrons, Aufubawe principle - diagonal rule: Electrons must fill the lowest energy level first, 18 28 2P 38 3P 3D AS AP AD AF Electronic rules: F al Br I At Chemistry Electrons are represented by 1 If possible electrons want to remain unpaired, electrons repel A full orbital must contain an opposite spin (1) When transition metal forms ions they lose electrons from 4s before 3d. Electron affinity Decreasing electron affinity $3-83 Shielding increases Distance between nucleus and outer electron increases Attraction of nucleus and outer electron is weaker, Harder to gain an electron, more energy is released Fluorine does not fit the trend due to It is too small, -2 F ions is formed but electrons are all too close together, These electrons repel each other, electron affinity (KJ mal-1) Electron affinity of the Group 7 elements 400 300 - 200- 100 0+ II fluorine chlorine 'bromine iodine Electron affinity - the energy released when one mole of gaseous atoms gain on mole of electrons to form a ion епенду Jonisation Decreasing the first ionisation energy Increase first ionisation energy Increases energy needed to remove electrons Nuclear charge increases. Shieling has ne overall change. The distance of electrons from nucleus decreases as you to the right The shielding and distance from the electrons to the positive nucleus and shells increases, Decreases attraction, Energy needed to lose an electron decreases. Endothermic energy in from the surroundings, Exothermic energy existing from the surroundings lonisation energy - the energy needed for one mole of electrons to be removed from one mole of gaseous atoms to form a gaseous ion (endothermic reaction) Element Li Be B C N Ne Na Log ie Total Electrons 3 4 5 6 7 10 Chemistry Orbital Diagram 1s 2s 10 1 16 16 12 12 1 2p 11111 14 16 14 12 1 12 3s 12 12 16 11 Be is higher than li due to increased amount of nuclear charge with no extra shells, Group 3 has an extra electron in 2p. increasing shielding making it easier to lose and electron, Group 6 has an extra electron pairs therefore the repulsion between the electrons make it easier to lose, Group 7 the volume increases due to increased amount of nuclear charge so 2p orbital is almost full, Ionisation energy graph, na Number of electrons removed