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When you gain an electron, you gain a positive charge

Electrons in outermost shell are called valence electrons and the valence shell is the outermost occupied shell Electron Configuration of Ions CATIONS are formed when a neutral atom loses one or more electrons: np electrons are lost first ns electrons are lost second

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When you gain an electron, you gain a positive charge

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  1. Electrons in outermost shell are called valence electrons and the valence shell is the outermost occupied shell Electron Configuration of Ions CATIONS are formed when a neutral atom loses one or more electrons: np electrons are lost first ns electrons are lost second nd electrons of previous shell lost third (if there’s any) Fe [Ar] 3d6 4s2 Fe3+ [Ar] 3d5

  2. Anions: to form monoatomic anions, add electrons until the next noble gas configuration has been reached N [He] 2s2 2p3 (room for 3 more electrons) N3- [He] 2s2 2p6 O [He] 2s2 2p4 (room for 2 more electrons) O2- [He] 2s2 2p6 When you gain an electron, you gain a positive charge When you lose an electron, you lose a negative charge

  3. Periodic Table and Electronic Configuration Periodic table is divided into s,p,d & f blocks which are named for the last subshell of the element that’s occupied Groups 1 & 2 – s block [no. of group] Groups 13-18 – p block [no. of group minus 10] Transition metals – d block Lanthanides & actinides – f block Periods (horizontal rows) are numbered according to the principle quantum number of the valence shell Exceptions: Helium 1s2 put in with the noble gases in group 18 Hydrogen 1s1 ( can act like group 1 & 17 )

  4. Periodicity of Atomic Properties • Atomic Radius of an element is defined as half the distance between the centres of neighbouring atoms. • Values generally increase down a group and decrease from left to right across a period • Ionic Radius of an element is its share of the distance between neighbouring ions in an ionic solution • Values generally increase down a group and decrease from left to right across a period • Cations are smaller and anions larger than their parent atoms

  5. Ionisation energy of an element is the energy needed to remove an electron from an atom of an element in the gas-phase • 1st & 2nd ionisation energies • generally higher for elements close to He and lower for elements close to Cs • Very high value if electron is expelled from a closed shell • Electron Affinity, Eea,of an element is the energy released when an electron is added to a gas-phase atom • generally positive values - generally energetically favourable • elements with highest values are those close to O, F, Cl.

  6. Chemical Bonds Chemical bond is a link between atoms • A bond forms if the resulting arrangement of atoms has a lower energy than the sum of the energies of the separate atoms. • Chemical bonds result from changes in location of electrons in the outermost shells of atoms (valence shells) Ionic Bonds An ionic bond is the electrical attraction between the opposite charges of cations and anions. Lewis Symbols for atoms and ions - keeps track of valence electrons

  7. N1s2 2s2 2px1 2py1 2pz1 O1s2 2s2 2px2 2py1 2pz1 Lewis Symbol : Symbol of Element + dot for each valence electron H 1s1 He1s2 A single dot: represents an electron alone in an orbital A pair of dots: represents two paired electrons in the same orbital

  8. Cl [Ne] 3s2 3p5Cl Cl - - Ca Cl Cl Cl To work out the formula of an ionic compound: 1. remove dots from Lewis symbol for metal atom. 2. transfer the dots to the Lewis symbol for the nonmetal atom and complete its valence shell 3. adjust the no.s of atoms to ensure all dots removed from metals are accommodated by nonmetals. 4. add charges of ions. Ca [Ar] 4s2 Ca + + Ca 2+ Chemical formula for calcium chloride is CaCl2

  9. Octet of electrons formed in each case • OCTET RULE: atoms of the reactive representative elements tend to undergo those chemical reactions that most directly give them the electronic configuration of the nearest noble gas Energy is needed to produce ions and when they pack together there’s a net overall lowering of energy - lattice enthalpy is a measure of this attraction Properties of Ionic Solids • ions stack together in regular crystalline structures (crystalline solids) • High m.p.s and b.p.s • Brittle • Form electrolyte solutions when dissolved in water

  10. Covalent Bonds • Compounds of nonmetals are not ionic - ionic compounds need ions of both positive and negative charge. • Atoms of nonmetals don’t become cations - too many electrons have to be lost to achieve noble gas configurations • Nonmetals form covalent bonds to one another by sharing pairs of electrons • A covalent bond is a pair of electrons between two atoms. • In covalent bonds, atoms share electrons to reach a noble gas configuration.

  11. The valence of an element refers to the number of covalent bonds an atom of the element forms H H F F H has one valence electron 1s1 H forms one bond, therefore its valence is 1 Line represents shared pair of electrons + H H H H F has 7 valence electrons [He] 2s2 2p5 F F F forms one bond, therefore its valence is 1 F2 has lone pairs of electrons - pairs of valence electrons not involved in bonding. H2 has no lone pairs

  12. H H H Structure of Polyatomic Species Methane CH4 Carbon is tetravalent H C Single bond: single shared pair of electrons (2) Double bond: double shared pair of electrons (4) Triple bond: triple shared pair of electrons (6) Bond Order: the number of electron pair bonds that link 2 atoms Single bond has a bond order of 1

  13. H N H H • choose atom with lowest I.E. for central atom • H is never central • arrange atoms symmetrically around central atom • central atom often written first in chemical formula (CH4) 1. Count total no. of valence electrons on each atom & divide by 2 to get no. of electron pairs Ammonia NH3 N 1 x 5 = 5 H 3 x 1 = 3 Total = 8 8  2 = 4 4 valence e- pairs 2. Write chemical symbols of atoms to show layout in molecule

  14. H H N N H H H H 3. Place one electron pair between each pair of bonded atoms ( one pair remains) 4. Complete octet of each atom (or duplet for H) by placing electron pairs as lone pairs around atoms

  15. O O O H H H H H H • If don’t have enough electron pairs to form octets, form multiple bonds • Check each atom has an octet or duplet Water H20H 2 x 1 = 2 O 1 x 6 = 6 Total = 8 8  2 = 4 4 valence e- pairs

  16. NH3 + H+ NH4+ (ammonium ion) H H N N H H H H Polyatomic Ions A shared electron pair can originate from one atom + H + H+ • If both shared electrons come from one atom- called a coordinate covalent bond

  17. Resonance 2 structures have exactly same energy - blend together as a resonance hybrid - electron density is spread evenly around the ring Benzene C6H6 Kekulé structure Resonance stabilises a molecule by lowering its total energy Resonance occurs between structures with the same arrangement of atoms, but different arrangement of electron pairs

  18. Acid + Base Complex O H O H H H H Lewis Acids & Bases Lewis Acid: electron pair acceptor Lewis Base : electron pair donor + H+ + Complex Lewis acid Lewis base

  19. Electronegativity • most bonds lie somewhere between pure ionic and pure covalent • Electronegativity is the electron withdrawing power of an atom A measure of the electron pulling power of an atom on an electron pair in a molecule Highest at upper right hand corner of periodic table and lowest at bottom left hand corner The greater the difference in electronegativities of 2 elements, the greater the extent of ionic character

  20. Molecular Structure Lewis structures : 2D diagrams & generally don’t show how molecules are arranged in space VSEPR Model: Valence Shell Electron-Pair Repulsion Model • molecules consist of central atom & attached atoms • attached atoms lie to corners of different shapes - describe the shapes of the molecules • Bond Angles: angles between the bonds • regions of high electron concentration (found in bonds and lone pairs) repel one another and take up positions as far away from each other as possible

  21. Ca Cl Cl F F F B Predicting Shapes of Molecules • write down Lewis structure & decide how electron pairs can be arranged around each “central” atom to minimise repulsions CASE A: Central atom with no lone pairs 1. CaCl2 Linear with bond angle of 180o 2. BF3 Trigonal Planar 120o

  22. Tetrahedral 109.5o 3. CH4 Trigonal bipyramidal 120o & 90o 4. PCl5 5. SF6 Octahedral 90o

  23. CASE B: Molecules with multiple bonds • VESPR theory doesn’t distinguish between single & double bonds • Electron pairs in a double bond act as a single unit of high electron concentration Linear CO2 CASE C: Molecules with lone pairs on central atom VESPR formula; A = central atom X = atom bonded to central atom E = lone pair of electrons on central atom NB: lone pairs on attached atoms not included

  24. e.g. BF3 = AX3 species (no lone pairs on central atom) NH3 = AX3E species (1 lone pair on nitrogen) Electron arrangement: 3D arrangement of all regions of high electron concentration (bonds and lone pairs) around central atom. AX3E species has 4 regions of high electron density Strengths of repulsions are in the order: lone pair-lone pair > lone-pair-bonding pair > bonding pair - bonding pair

  25. H O H AX2E2 O H H Water H2O • 4 electron pairs adopt a tetrahedral arrangement • only 2 positions are occupied by atoms - classified as angular or “bent” • lone pairs push away from each other • bonding atoms forced closer together Bond angle is less than that of tetrahedron ( 104.5o compared to 109o)

  26. H N H H Ammonia NH3 AX3E • 4 pairs adopt a tetrahedral arrangement • 3 positions occupied by atoms • H atoms have moved slightly towards each other Trigonal pyramidal ( 107o compared to 109.5o ) N H H H

  27. No. of e - pairs in bonds No. of lone pairs VSEPR type Structure Example BeCl2 CaCl2 AX2 2 0 linear nonlinear bent H2OH2S AX2E2 2 2 BCl3 BF3 planar triangular AX3 3 0 trigonal pyramidal AX3E 1 NH3 3 AX4 tetrahedral CH4 4 0 trigonal bipyramidal AX5 5 0 PCl5 SF6 AX6 6 0 octahedral

  28. Polar Molecules • Polar Covalent Bond: electron pair shared unequally between 2 atoms resulting in some partial ionic character • arises from differences in electronegativities of atoms O more electronegative than H & gains greater share in the bonding e- pair O H Polar bond - = O has a partial negative charge O- H + + = H has a partial positive charge Two atoms in a polar bond with partial charges give rise to an electric dipole moment

  29. When bonded atoms have different electronegativities, the bond is polar Nonpolar molecule has zero dipole moment e.g. H2 & Cl2 O C O O H H + - - Polar bonds, but overall a non-polar molecule - Polar bonds and polar molecule + + AX2, AX3, AX4, AX5 & AX6 - non-polar (where X = same element)

  30. Bond Strengths • strength of bonds is measured by bond enthalpy, HB • bond breaking requires energy, so all HB values are positive • HB typically increase as bond order increases and decrease as atomic radius and no. of lone pairs increases Bond Lengths • Bond length: distance between the centres of 2 atoms joined by a chemical bond • for bonds between same atoms, the shorter the bond, the stronger it is • e.g. C C is shorter & stronger than C C • covalent radii- added to estimate bond lengths in molecules

  31. 1s 1s Molecular Orbitals - another view of the covalent bond H2  bond (sigma bond) Looking along internuclear axis, e- distribution resembles that of an s orbital

  32. Molecular Orbitals - another view of the covalent bond • a shared e- pair resides in a molecular orbital formed by the partial overlapping of two atomic orbitals • space created by the overlapping of 2 atomic orbitals is called a molecular orbital • M.O. Can hold a maximum of 2 e- with opposite spins (like an A.O.) • A.O. That overlap are generally those of valence shell electrons • p-orbitals can also form sigma bonds

  33. N 1s2 2s2 2px1 2py1 2pz1 N2  Two 2pz orbitals form a “head-on” sigma bond  Looking along internuclear axis, e- distribution resembles that of an p orbital Two 2px orbitals form a pi bond

  34. remaining two 2py orbitals form another  bond • N2 has one sigma bond ( from both 2pz orbitals) and two pi bonds (from both 2px and 2py orbitals) A single bond is a  bond A double bond is a  & a  bond A triple bond is a  & two  bonds

  35. 2p 2s 1s 1s Promotion & Hybridisation • Bonding in polyatomic orbitals more complex to explain e.g. CH4 • C 1s2 2s2 2px1 2py1 (looks like C can only form 2 bonds) • Know carbon can nearly always form 4 bonds. HOW? • NB carbon has one empty p orbital (2pz) • If we promote an e- from the 2s orbital into this empty 2pz orbital, get 4 unpaired electrons. 2p 2s After promotion (can form 4 bonds) Before promotion (can only form 2 bonds)

  36. sp3 orbital BUT now looks like there are 2 different bond types in CH4 Type 1H 1s C 2s H 1s C 2px H 1s C 2py H 1s C 2pz Type 2 • In fact, four sp3 hybrid (mixture) orbitals formed from combination of s orbital and 3 p orbitals • These orbitals have equal energies lying between the energies of s and p orbitals

  37. Bonding in Methane (CH4) • One unpaired electron occupies each of carbon’s sp3 hybrid orbitals • Each of these 4 electrons can pair with an electron in a H 1s orbital • 4 sigma () bonds are formed • Structure of methane is tetrahedral

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