Nomenclature of inorganic compounds Ionic equations. MUDr. Jan Pláteník, PhD. Structure. of matter: Atom Molecule Ion Compound

Transcription

1 Nomenclature of inorganic compounds Ionic equations MUDr. Jan Pláteník, PhD Structure of matter: Atom Molecule Ion Compound Mixture (dispersionion system) 1

2 Atom Smallest particle of a pure element having its chemical properties Positively charged nucleus (neutrons, protons) Electron shell: Electron is wave/particle Behavior of electron described by quantum mechanics (... wave function, quantum numbers) Orbital: space area within the atom shell where occurrence of an electron or pair of electrons is more probable Molecule smallest particle of a substance having its chemical properties Atoms connected via covalent bonds Examples: noble gases: monoatomic other gases: diatomic 2 O, N 3 etc. e molecular crystals: diamond...many thousands of atoms in proteins and nucleic acids O N 2

3 Molecul ular crystal of diamond Ion atom or molecule with non zero charge (number of electrons does not match number of protons) tendency to form ions depends on electronegativity of element cations (+) or anions (-) monoatomic: Na +, Cl -, +, Fe 2+ molecular: NO 3-, SO 4 complex: [Fe(CN) 6 ] 4-3

4 Molecular ions of oxo-acids: e.g. sulfate, SO 4 : resonance stabilization of sulfate ion..similar is nitrate NO 3-, phosphate PO 4 3-, carbonate CO 3, etc. Compound Chemically pure substance consisting of the same molecules that containing two or more different atoms Atoms are held together in a defined spatial arrangement by chemical bonds Independent molecules (e.g. CO 2 ) or crystalline structures 4

5 Chemical formulas Stoichiometric (empirical) e.g.: sodium chloride NaCl e.g.: glucose C 2 O Molecular e.g.: sodium chloride NaCl e.g.: glucose C 6 12 O 6 Structural Chemical bond Cohesive force holding together atoms in molecules and crystals Ionic bond: electrostatic forces among ions having opposite charges Covalent bond: sharing a couple of electrons by the two bonding atoms 5

6 Polarity of chemical bond Determined by difference in electronegativity of the two connected atoms: < 0.4 nonpolar covalent bond e.g. -, carbon-hydrogen polar covalent bond e.g. -O-, N 3, carbon-oxygen, carbon-nitrogen >1.7 ionic bond e.g NaCl... Gradual transition! Molecular shape affects electric dipole moments: δ+ δ+ δ δ+ δ O C O O CO 2 : linear, nonpolar δ 2 O: angled, polar... Water as polar solvent: 6

7 Coordination covalent bond (also dative, donor-acceptor bond) Both bonding electrons provided by one of the atoms (donor), whereas the other atoms provides an empty orbital (akceptor) Coordination (complex) ions central atom of transition metal providing empty orbitals ligands providing free electron pairs Number of ligands (coordination number) is usually 4 or 6 e.g. [Fe(CN) 6 ] 4-7

8 Valence Number of covalent bonds formed by an atom octet rule: tendency to achieve electron configuration of the nearest noble gas e.g. -F: achieves configuration of e, F configuration of Ne Therefore O usually bivalent, N trivalent, C tetravalent etc. Oxidation number (formal valency) Oxidation number of element in compound equals its charge after giving all bonding electron pairs to the more electronegative atom Can be zero, positive or negative integer Basis for nomenclature of inorganic compounds Redox reactions: oxidation number increases in oxidation, decreases in reduction 8

9 Rules for determination of oxidation numbers Free electroneutral atom, or atom in molecule of pure element: oxidation number = 0 Oxidation number of a monoatomic ion equals its charge In heteroatomic compounds the bonding electrons are given to the more electronegative atom, practically: has nearly always oxidation number I (only in metallic hydrides -I) O almost always -II (only in peroxides -I) F always -I Alkaline metals (Na, K..) always I Alkaline earth elements (Ca, Mg..) always II Examples: Rules for determination of oxidation numbers CO 2 : C IV, O -II 2 SO 4 : I, S VI, O -II Sum of oxidation numbers of all atoms in electroneutral molecule is 0, in polyatomic ion equals the ion charge e.g: CO 3 : C IV, O -II 1 IV + 3 (-II) = -2 9

10 Czech nomenclature of oxides: Oxidation number I II III IV V VI VII VIII Suffix -ný -natý -itý -ičitý -ečný/-ičný -ový -istý -ičelý General formula X 2 O XO X 2 O 3 XO 2 X 2 O 5 XO 3 X 2 O 7 XO 4 English (international) nomenclature Example 1: Two oxidation numbers II and III possible for Fe: FeCl 2 Iron(II) chloride, ferrous chloride FeCl 3 Iron(III) chloride, ferric chloride Example 2: oxoacids of chlorine, four oxidation numbers I, III, V and VII possible for Cl: ClO hypochlorous acid ClO 2 chlorous acid ClO 3 chloric acid ClO 4 perchloric acid 10

11 IONIC EQUATIONS Ionic salts: no true molecule Crystal lattice of NaCl: Dissolution of NaCl in water: electrolytic dissociation producing hydrated independent ions Na +, Cl - 11

12 Reaction I Stoichiometric equation: AgNO 3 + KCl AgCl + KNO 3 Ionic equation: Ag + + NO K + + Cl - AgCl + K + + NO 3 - Net ionic equation: Ag + + NO K + + Cl - AgCl + K + + NO 3 - Ag + + Cl - AgCl white ppt Also possible: AgNO 3 (aq) + KCl(aq) AgCl(s) + KNO 3 (aq) Ag + (aq) + Cl - (aq) AgCl(s) (aq) (s) (l) (g)... aqueous... solid... liquid... gaseous 12

13 Reaction II ionic: CuSO 4 + 2NaO Cu(O) 2 + Na 2 SO 4 Cu 2+ + SO 4 + 2Na O - Cu(O) 2 + SO Na + net ionic: Cu 2+ + SO 4 + 2Na O - Cu(O) 2 + SO Na + Cu O - Cu(O) 2 pale blue ppt Reaction III 2 NaO + (N 4 ) 2 SO 4 2 N 3 + Na 2 SO O ionic: 2Na + + 2O - + 2N 4+ + SO 4 2N 3 + 2Na + + SO O net ionic: 2Na + + 2O - + 2N 4+ + SO 4 2N 3 + 2Na + + SO O O - + N 4 + N O 13

14 Ammonia gas: N 3, N 3 (g) Aqueous ammonia: N 3 (aq), N 3. 2 O, N 4 O N 3. 2 O N 4 O N 4 O N 4+ + O - Reaction IV Cu(O) 2 + 2(N 4 ) 2 SO 4 + 2NaO [Cu(N 3 ) 4 ]SO Na 2 SO O Ionic: Cu(O) 2 + 4N SO 4 + 2Na + + 2O - [Cu(N 3 ) 4 ] SO 4 + 2Na O Net ionic: Cu(O) 2 + 4N O - [Cu(N 3 ) 4 ] O Dark blue complex 14

15 Writing Ionic equations: Summary 1. write correct and balanced stoichiometric equation first 2. rewrite to ionic: write separately any species that exist separately and indicate its charge if present, but write together what exists joined (usually a precipitate of insoluble salt, or a soluble coordination complex) 3. Cancel out all species not involved in the reaction 4. Check that the equation is still balanced What combinations of cations and anions are insoluble? All nitrates (NO 3- ) and acetates (C3COO - ) are soluble All salts of Na, K, Li, and N 4+ are soluble All chlorides, bromides and iodides are soluble except salts of Pb 2+, Ag +, and g 2 2+ Most sulfate salts are soluble except BaSO 4, PbSO 4, gso 4, and CaSO 4. Most hydroxides are insoluble. Soluble are only NaO and KO. Ba(O) 2, and Ca(O) 2 are marginally soluble. Most sulfides (S ), carbonates (CO 3 ) and phosphates (PO 4 3- ) are insoluble. 15

16 Names of coordination compounds Names of neutral ligands: 2 O aqua N 3 ammin NO nitrosyl CO carbonyl Names of anionic ligands always end in o: F fluoro Cl chloro Br bromo I iodo O hydroxo CN cyano etc.. Names of coordination compounds 1. Complex particle is cation: e.g. [Cu(N 3 ) 4 ]SO 4 [Cu(N 3 ) 4 ] 2+ + SO 2 4 Tetraamminecopper(II) sulfate 2. Complex particle is anion: e.g. K 3 [CoF 6 ] 3 K + + [CoF 6 ] 3 Potassium hexafluorocobaltate(iii) 16

17 Names of coordination compounds 3. Both cation and anion are complexes: e.g. [Pt(N 3 ) 4 ][PtCl 4 ] [Pt(N 3 ) 4 ] 2+ + [PtCl 4 ] 2 Tetraammineplatinum(II) tetrachloroplatinate(ii) 4. Neutral complexes: e.g. [CrCl 3 ( 2 O) 3 ] Triaquatrichlorochromium(III) complex 17