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IONIC BONDING A. CHEMICAL BONDS: Attractive forces between atoms in a compound or molecule 1. Formed by VALENCE ELECTRONS B. IONIC BONDS 1. Formed by TRANSFER OF ELECTRONS FROM ONE ATOM TO ANOTHER 2. Formed between METALS with low ionization energies and NONMETALS of high ionization energies 3. Best examples of ionic bonding: between Groups 1A & 2A with Groups 6A & 7A 4. Electronegativity difference: 1.7 or greater Ex. KF 4.0 - .8 = 3.2 Na2O 3.5 - .9 = 2.6 Which compound has greater ionic character? KF C. Energy Changes in the formation of an ionic compound 1. Elements converted to the gaseous state Na(s) + ½ Cl2(g) Na(g) + Cl(g) ΔH = +55 kcal 2. Electron transfer Ionization of Na Na Na +1 + 1eΔH = 119 kcal Formation of Chloride Ion Cl + 1e Cl -1 ΔH = - 83 kcal 3. Ions combine to form solid (Bond Formation) Na+1 + Cl -1 Na+1 Cl -1 ΔH = - 184 kcal LATTICE ENERGY TOTAL ENERGY: 55+119-83-184 = -93 Kcal LATTICE ENERGY Energy released when ions arrange themselves in a crystal lattice FORMATION OF IONIC BONDS (Key steps) 1. 2. Electron transfer Attraction of positive and negative ions (Electrostatic attraction) WHY DO ATOMS REACT? Elements tend to react to acquire a stable electron configuration, usually a noble gas configuration. (OCTET RULE) Charges on the ions formed depend upon the # of electrons lost or gained to acquire the noble gas configuration IONS t 1. Positive ions (metallic ions): CA IONS smaller than their corresponding atoms are _______ Why? a. Outer energy level is removed Na Na+ 1s22s22p63s1 1s22s22p6 b. Greater positive than negative charge draws remaining energy levels closer to nucleus Ions 2. Negative (nonmetallic) ions: ANIONS are larger ______ than their corresponding atoms Why? Extra repulsion in the electron cloud weakens pull of nucleus on e- cloud F F-1 1s22s22p5 1s22s22p6 F F- 3. While metal atoms are usually larger than nonmetal atoms in the same period, POSITIVE ions are usually SMALLER than negative ions (of the same period). 4. Periodic trends in ion size Compare the sizes of the ions of Mg , Al , Na Na+ Mg2+ Al3+ Of the ions of P , S , and Cl P3- S2- Cl 1increases b. Within a group, ion size ____________ moving down. CHARGES OF IONS WITH NOBLE GAS CONFIGURATIONS Br: Charge: -1 S: Charge: -2 Sr: Charge: +2 Cs Y Al N Attains e- configuration of Kr Attains e- configuration of Ar Attains e- configuration of Kr NAMING IONIC COMPOUNDS 1. Monatomic positive ions have the name of their metal. 2. In a binary compound, the negative ion is named by dropping the element’s ending and adding IDE. NaCl = Sodium chloride BaF2 = Barium fluoride K3N = Potassium nitride FORMULAS OF IONIC COMPOUNDS COMPOUNDS ARE ELECTRICALLY NEUTRAL! Magnesium & oxygen: Mg2+O2- = MgO Scandium & hydroxide: Sc3+(OH)1- = Sc(OH)3 Cobalt(II) & bromine: Co2+Br1- = CoBr2 Cobalt(III) & bromine: Co3+Br1- = CoBr3 Many transition elements form ions with more than one positive charge (oxidation state) Except: Ag and Zn Ex. Iron(II): Fe2+ Iron(III): Fe3+ 3. Certain metals form positive ions with more than one positive charge. Determine the charge on the positive ion and write it in Roman numerals, in parentheses, after the name of the positive ion. Iron(III) chloride FeCl3 Iron(II) chloride FeCl2 Cr(NO3)2 Chromium(II) nitrate Cr(NO3)3 Chromium(III) nitrate Polyatomic Ions Charged species containing more than one atom (Ex. NH4+, SO42-) 4. Compounds containing three or more elements contain polyatomic ions. Name the positive ion and then the polyatomic ion. Al(NO3)3 Aluminum nitrate NH4ClO3 Ammonium chlorate Na2SO4 Sodium sulfate Calcium & sulfate Ca 2+(SO4)2- = CaSO4 Aluminum & sulfite Al3+(SO3)2- = Al2(SO3)3 Gallium & hydrogen carbonate Ga3+(HCO3)-1= Ga(HCO3)3 PRACTICE! BaI2 KNO2 NiS KClO4 CuBr Au(ClO2)3 CoI3 Al2(CO3)2 AgNO3 Sn(SO4)2 GaCl3 (NH4)2O Sr3(PO4)2 Ca(HCO3)2 HYDRATES 1. Ionic compounds which retain water in their crystal lattice Ex. CuSO4* 5 H2O CoCl2 * 6 H2O BaCl2 * 2 H2O 2. When heated, they decompose, releasing water. CuSO4* 5 H2O(s) CuSO4(s) + 5 H2O(l) CuSO4* 5 H2O(s) CuSO4(s) + 5 H2O(l) Hydrate: CuSO4 * 5 H2O Water of hydration: H2O Anhydrous salt: CuSO4 EFFLORESCENCE Loss of water by a hydrate a. Endothermic process Calculate the percent of water in BaCl2*2H2O. Write the equation for the efflorescence of BaCl2*2H2O. Label: Water of hydration, hydrate, anhydrous compound WRITING FORMULAS FROM NAMES Pay attention to ENDINGS! Potassium bromide K+1Br-1 = KBr Aluminum nitrite Al+3(NO2)-1 = Al(NO2)3 Mercury(II) sulfite Hg2+(SO3)2- = HgSO3 Iron(III) hydroxide Fe3+(OH)1- = Fe(OH)3 WRITING EQUATIONS FOR THE FORMATION OF IONIC COMPOUNDS Write the balanced equation for the formation of copper(II) oxide from its elements. 1. Write the correct formula for the compound. Cu2+O2- CuO 2. Write the EQUATION for the formation of the compound. Then balance. Cu(s) + O2(g) → CuO(s) 2 Cu(s) + O2(g) → 2 CuO(s) Write the balanced equation for the formation of iron(III)chloride from its elements. 1. Fe3+Cl1- = FeCl3 2. Fe(s) + Cl2(g) → FeCl3(s) 3. 2 Fe(s) + 3 Cl2(g) → 2 FeCl3(s) Oxidation Reduction Chemistry: Redox Chemistry Oxidation and Reduction reactions always take place simultaneously. Loss of electrons – oxidation (Increase in Oxidation Number) Ex: Na ------> Na+1 + e-1 Gain of electrons - reduction ( Decrease in Oxidation Number) EX: Cl2 + 2 e-1 ------> 2 Cl-1 Oxidation occurs when a molecule does any of the following: Loses electrons Loses hydrogen Gains oxygen If a molecule undergoes oxidation, it has been oxidized and it is the reducing agent Reduction occurs when a molecule does any of the following: Gains electrons Gains hydrogen Loses oxygen If a molecule undergoes reduction, it has been reduced and it is the oxidizing agent. Leo the Lion! LEO the lion says GER Loss of Electrons is Oxidation, Gain of Electrons is Reduction Example Redox reactions involve electron transfer: Lose e - =Oxidation Cu (s) + 2 Ag (aq) Cu2+ (aq) + 2 Ag(s) Gain e - =Reduction Oxidizing and Reducing Agents Now the confusing part… CuO + H2 Cu + H2O Cu goes from +2 to 0 Cu is reduced, therefore it is called an oxidizing agent because it causes some other substance to be oxidized H goes from 0 to +1 H is oxidized, therefore it is called a reducing agent because it causes some other substance to be reduced. Identifying Agents in an Equation Reduction: CuO is the oxidizing agent CuO + H2 Cu + H2O Oxidation: H2 is the reducing agent Oxidation Numbers A count of the electrons transferred or shared in the formation or breaking of chemical bonds You must assign each element in the reaction an oxidation number Follow a set of rules… Oxidation Number Rules 1. The total of the oxidation numbers of all the atoms in a neutral molecule, an isolated atom, or a formula unit is 0 2. In their compounds, the Group 1A metals all have an oxidation number of +1, and the Group 2A metals have an oxidation number of 2+. Rules Continued 3. In its compounds, hydrogen has an oxidation number of +1 (except in metal hydrides such as NaH, where it is -1) 4. In its compounds, oxygen has an oxidation number of -2 (except in peroxides such as H2O2, where it is -1) 5. In their binary compounds with metals, Group7A elements have an oxidation number of -1. Group 6A elements have an oxidation number of -2, and Groups 5A elements have an oxidation number of -3. Examples Is the reactant oxidized or reduced? Pb PbO3 SnO2 SnO KClO3 KCl Oxidized Reduced Reduced Identifying Redox Reactions 0 +3 -2 0 +3 -2 2 Al + Fe2O3 2 Fe + Al2O3 Al increases from 0 to +3, it is ______ Oxidized! Fe decreases from +3 to 0, it is _______ Reduced! Problems What is the oxidation number of each element? I 0 Why? Its alone! I2 Cr2O3 Cr +3 O -2 Al +3 Cl -1 AlCl3 Na2SO4 Na +1 S+6 CaH2 Ca +2 H -1 O4-2 Oxidation Numbers with Charges (PO PO43)31P 4O @ 5+ 3- = 5+ 3@ 2- = 8113- ? Fluorine and oxygen are highly electronegative and will attract electrons very strongly meaning they will have a negative oxidation number. Generally, phosphorus will be 3- oxidation state: however, when combining with oxygen, phosphorus will lose five electrons and take on a 5+ oxidation charge. SYNTHESIS (COMBINATION) One or more elements or compounds combine to from a single product (compound). A B AB One Product Examples: 2 Cu + CO2 + O2 H2O 2 CuO H2CO3 DECOMPOSITION A single compound is broken down into 2 or more simpler substances AB A B One Reactant Examples: 2 H2O2(aq) 2 KClO3(s) 2 H2O(l) + O2(g) 2 KCl(s) +3 O2(g) SINGLE REPLACEMENT A more reactive element replaces another from a compound A BC Examples: Zn(s) + 2 HCl (aq) Cl2 (g)+ 2 KI(aq) AC B ZnCl2 (aq) + H2(g) 2 KCl(aq) + I2(s) DOUBLE REPLACEMENT Ions in a solution combine to form a product that leaves the scene of the reaction AB CD AD CB Example: AgNO3 (aq) + NaCl (aq) AgCl (s) + NaNO3 (aq) ONLY occurs if one product is removed from the solution (precipitate (s), gas (g) or water forms.) AgNO3 (aq) + NaCl (aq) AgCl (s) + NaNO3 (aq) Actual reaction: Ag+1 + Cl-1 Spectator Ions: Na+1 & NO3-1 AgCl COMBUSTION Element or compound reacts with oxygen, producing light and heat CxHy+ O2 1 CO2 + H2O C2H2+ 2.5 O2 * 2 C2H2 + 5 O2 2 CO2 + 1 H2O 4 CO2 + 2 H 2O Nuclear Reactions The composition of the nucleus is changed. Stable Nuclei Stable nuclei are NOT radioactive Stable nuclei are elements #1-83 (#84are radioactive) Strong nuclear forces = attraction between particles in nucleus that hold it together VERY STRONG! P P • #1-20 equal number of protons & neutrons for stable nuclei • #21-83 nuclei need more & more neutrons to be stable • #84 radioactive (all isotopes Types of Radioactive Decay 1. Alpha particles 2. Beta particles 3. Gamma particles Alpha Particles = α 2 4 4 2 Consists of 2 protons & 2 neutrons Has a +2 charge Identical to a He-4 nucleus Stopped by paper 226 88 Ra α + Rn 4 2 222 86 He Alpha decay problems Write the nuclear equation for the alpha decay of 231 Pa 91 231 Pa 91 4 2 α 227 89 + Ac Write the nuclear equation for the alpha decay of 244 Pu 244 94 Pu 94 α 4 2 + 240 92 U Beta particles 0 -1 0 -1 β= e High speed electron is emitted out from atom -1 charge Stopped by heavy clothing Neutron changes into a proton & an electron I e + Xe 131 53 0 -1 131 54 Beta Decay problems Write a nuclear equation for the beta decay of 223 Fr 87 223 87 Fr 0 -1 223 e + Ra 88 Write a nuclear equation for the beta decay of 50 Ti 22 50 22 Ti 0 -1 50 23 e+ V Gamma Radiation 0 0 High energy Radiant energy 0 charge, 0 mass Most penetrating Stopped by lead or concrete Recap Other Nuclear Reactions Fission is splitting of the nucleus (atomic bomb) Fusion is joining of nuclei (H-bomb) Fission Chain Reaction Fusion