Chapter 4 The Effects of Chemical Reactions. • Introduction to Chemical Reactions. - Chemical reaction: a process in which one or more substances change into one or more new substances. - Clues that a chemical reaction has occurred : 1. Color change Example: two colorless aqueous solutions mix together to produce a bright yellow precipitate. 2. A precipitate (solid) is formed when mixing two solutions together. 3. Gas formation. Bubbles of gas (effervescence) are produced when mixing substances together (solid – liquid or aqueous – aqueous ….) 4. Heat is produced. - Chemical reactions are described by using word equations or chemical equations. - Chemical equations need to be balanced when written because it shows the correct proportions (amounts) of chemicals in a reaction. - A balanced chemical equation has equal number of atoms of each element in the reactants (left hand side) and the products (right hand side). - Exercise: Balance the following equations. a) KClO3→KCl + O2 b) Na2O + H2O NaOH c) Cu + AgNO3 Cu(NO3)2 + Ag d) C3H7OH + O2 CO2 + H2O • Synthesis and Decomposition Reactions. Synthesis: Two or more substances (elements and / or compounds) combine to form one larger compound. General pattern: A + B → C Examples: N2 + 3 H2 → 2 NH3 CaO + CO2 → CaCO3 2 P + 3 Cl2 → 2 PCl3 Decomposition: This is opposite to synthesis; that is, one large compound breaks down (decomposes) into 2 or more simpler substances. Example: 2 KClO3 → 2 KCl + 3 O2 General pattern: R → S + T Remark: Usually decomposition happens due to heat or electricity. - Predicting the product of decomposition or synthesis reactions. 2 AlCl3 (s) → 2 Al (s) + 3 Cl2 (g) Zn (s) + S (s) → ZnS (s) 2 Zn (s) + O2 (g) → 2 ZnO(s) - Single Displacement (Replacement) Reactions. Definition: A reaction in which an element displaces (replaces) another element in a compound, producing a new compound and a new element. General pattern: A + BC → AC + B Example: Mg (s) + CuSO4 (aq) → MgSO4 (aq) + Cu (s) Zn (s) + 2 AgNO3 (aq) → Zn(NO3)2 (aq) + 2 Ag (s) Fe (s) + MgCl2 (aq) → no reaction. Remark: The element that displaces the other element in a compound must be more reactive (active) than that element, otherwise no reaction takes place. In the general pattern above, A should be more reactive than B for the reaction to proceed. The following reactivity (activity) series lists the chemical strength (reactivity) of the metals in order from the more reactive to the less reactive. KPlease stop calling my amazing zebra in the long Nahungry class. sorry !! Ca Mg Al Zn Fe Sn Pb H Cu Ag Examples of single displacement reactions : 2 Al (s) + 3 CuSO4 (aq) → Al2(SO4)3 (aq) + 3 Cu (s) Sn (s) + Zn(NO3)2 (aq) → no reaction Exercise: Complete and balance the following equations. If there is no reaction occurring write no reaction. a) 2 Al (s) + 6 HCl (aq) → 2 AlCl3 (aq) + 3 H2 (g) b) Cu (s) + H2SO4 (aq) → no reaction c) 2 AlCl3 (aq) + 3 Ca (s) → 3 CaCl2 (aq) + 2 Al (s) d) Mg (s) + 2 HNO3 (aq) → Mg(NO3)2 (aq) + H2(g) - Reactivity of halogens decreases down the group. F2> Cl2> Br2> I2 The reactions taking place for the halogens or their compounds are in solution (aqueous) Examples: Cl2 (aq) + 2 KBr (aq) → 2 KCl (aq) + Br2 (l) Cl2 (aq) + NaF (aq) → no reaction. Exercise: F2 (aq) + 2 LiCl (aq) → 2 LiF (aq) + Cl2 (g) I2 (aq) + NaCl (aq) → no reaction • Double displacement reactions. - Definition: A reaction in which two compounds mix together and an exchange of ions (elements) occurs which results in the formation of 2 new compounds. - General pattern: AB + CD → AD + CB - Solubility: the amount of solute that dissolves in a given amount of solvent at a given temperature. - When we say a substance is soluble, it means it dissolves in water; whereas if it is insoluble it means it doesn’t dissolve in water. - The compound in a reaction that is soluble is in aqueous (aq) phase, whereas the compound which is insoluble is in the solid state (s). - The solid which is formed in a double displacement reaction is called the precipitate and it is insoluble. - Solubility rules (used in double displacement reactions). 1. All alkali metal ions and ammonium ion (NH4+) are soluble. 2. All nitrates (NO3-) are soluble. 3. All sulfates (SO4-2) are solubleexceptwith Ba+2 , Pb+2 , Ca+2 , Sr+2 , Ag+ . 4. All chlorides, bromides and iodides(Cl-, Br-, I-) aresolubleexcept with Ag+ , Pb+2 , Hg+, Cu+ 5. All OH- are insolubleexceptwith rule 1, and Ba+2 and Sr+2 . 6. All oxides (O2-), sulfides (S2-), sulfites (SO32-), carbonates (CO32-), phosphates (PO43-) are insoluble except with rule 1 Remark: If all compounds formed in a double displacement reaction are soluble (aqueous) then no reaction takes place. Exercise: State whether each of the following compounds is soluble or insoluble ? Na2SO4 : Fe(NO3)2: LiOH: ZnSO4: PbBr2: BaSO4: Mg(OH)2: PbO: NH4Cl: Na2S: Cu(OH)2: KF: Exercise: Complete and balance the following chemical equations: - KNO3 (aq) + NaCl (aq) → - LiCl (aq) + AgNO3 (aq) → - Zn (s) + FeSO4 (aq) → - NaOH (aq) + CuCl2 (aq) → - ZnCl2 (aq) + Na3PO4 (aq) → - Pb(NO3)2 (aq) + K2S (aq) → • Net ionic equation: a chemical equation which shows ONLY the ions that are involved in the formation of the precipitate (solid). Examples: Pb+2 (aq) + S-2 (aq) → PbS (s) Ag+ (aq) + Cl- (aq) → AgCl (s) Cu+2 (aq) + 2 OH- (aq) → Cu(OH)2 (s) • Full ionic equation: an equation which shows All the ions in the soluble (aqueous compounds) in both reactants and products. Example: - 2 NaOH (aq) + CuCl2 (aq) → 2 NaCl (aq) + Cu(OH)2 (s) 2 Na+ (aq) + 2 OH- (aq) + Cu+2 (aq) + 2 Cl- (aq) → 2 Na+ (aq) + 2 Cl- (aq) + Cu(OH)2 (s) - 3 ZnCl2 (aq) + 2 Na3PO4 (aq) → Zn3(PO4)2 (s) + 6 NaCl (aq) Full ionicequation: 3 Zn+2(aq) + 6 Cl-(aq) + 6 Na+ (aq) + 2 PO4-3 (aq) → Zn3(PO4)2 (s) + 6 Na+ (aq) + 6 Cl- (aq) Net ionic equation: 3 Zn+2 (aq) + 2 PO4-3 (aq) → Zn3(PO4)2 (s) Exercise: Complete and balance the following equation, then write full ionic and net ionic equations for the reaction. Pb(NO3)2 (aq) + 2 NaI (aq) → Full ionic equation: Net ionic equation: Spectator ions: the ions that are not involved in the formation of the precipitate (solid). Note that the spectator ions appear on both sides of the full ionic equation. For example, in the above reaction, Na+ (sodium ions) and NO3- (nitrate ions) are the spectator ions. Exercise: Complete and balance the following equation, then write the net ionic equation and identify the spectator ions. BaCl2 (aq) + K2SO4 (aq) → Net ionic equation: Ba+2 (aq) + SO4-2 (aq) → Spectator ions: - Combustion reaction is a special type of (synthesis) reaction in which the substance reacts with (burns in) oxygen. Examples: C(s) + O2(g) → CO2(g) • Production of gases (lab scale): 1. CO2 2. SO2 3. H2 4. H2S (hydrogen sulfide) 5. NH3 (ammonia) General pattern of the chemical reactions to produce the above gases: 1. Metal carbonate + acid → CO2 Example: Na2CO3 (aq) + 2 HCl (aq) → 2 NaCl(aq) + CO2(g) + H2O(l) 2. Metal sulfite + acid → SO2 K2SO3 (aq) + 2 HCl (aq) → 2 KCl(aq) + SO2(g) + H2O(l) 3. Metal + acid → H2 Remark: This is a single displacement reaction therefore the metal used in the reaction should be higher in the reactivity series than hydrogen. Zn (s) + 2 HCl (aq) → ZnCl2 (aq) + H2(g) 4. Metal sulfide + acid → H2S Na2S (aq) + 2 HCl (aq) → 2 NaCl (aq) + H2S (g) 5. Ammonium compound + base (alkaline solution) → NH3 NH4Cl (aq) + NaOH (aq) → NaCl (aq) + NH3 (g) + H2O (l) Exercise: Write the net ionic equations for each of the above 5 reactions. Answers 1. 2 H+ (aq) + CO3-2(aq) → CO2(g) + H2O (l) 2. 2 H+ (aq) + SO3-2(aq) → SO2(g) + H2O (l) 3. Zn(s) + 2 H+(aq) → Zn+2(aq) + H2(g) 4. 2H+ (aq) + S-2 (aq) → H2S (g)

Ch 2 Chemical Compounds and Bondings Chemical bond: force of attraction between 2 atoms or 2 ions. There are 2 main types of chemical bonds: 1. Ionic bond: a bond between a metallic atom (metal) and a non-metallic atom (non-metal) in which there is a complete transfer of electrons from the metal to the non-metal. The compound which is formed is called an ionic compound. Ex.: NaCl , MgCl2, Al2O3 To write the formula of an ionic compound we use the criss-cross method (we down cross multiply the charges without the sign, only the numbers of the charges) Remark: if the charges are the same then they cancel each other in the formula so there will be one atom of the metal and one atom of the non-metal in the compound. Exercise: Write the formula of the compound which is formed between the following elements, and name each compound. a) Rb and S: b) Ca and Se: c) Al and Br: d) Na and N: - Draw a Bohr diagram to show the transfer of electrons (loss / gain) in an ionic compound. Example: Na2O (sodium oxide) Remark: The ionic bond is also described as an electrostatic force of attraction between a positive ion and a negative ion (eg: Na+ Cl- ). 2. Covalent bond: a bond between a non-metal and another non-metal in which there is a sharing of electrons between the non-metallic atoms. The compound that is formed is called a covalent or molecular compound. Molecule: 2 or more atoms (non-metallic) bonded together; the atoms can be of the same element such as O2 or of different elements such as HCl, CH4, H2O,….. We show the sharing of electrons between non-metallic atoms by using the Lewis diagram. In addition we can make intersecting circles for the atoms to show the sharing. Reminder: Lewis diagram of an atom shows only the valence electrons of that atom. Most of the atoms follow the octet rule (there are very few exceptions), that is each atom will have eight electrons in the valence shell (same as noble gases) except hydrogen will have 2 electrons after sharing (same as helium). Remark: 2 electrons that are not bonded to any other atom is called a lone pair (non-bonding pair) of electrons. Exercise: Draw lewis diagrams to show the sharing of electrons in each of the following compounds: 1. NF3 2. CH4 3. CO2 4. CCl4 5. CH2O Remark:In drawing Lewis structure, we show the bonds between the atoms and we also show all lone pairs (if present) on any atom. - Naming molecular compounds: 1: mono 2: di 3: tri 4: tetra 5: penta 6: hexa 7: hepta 8: octa 9: nona 10:deca Example: PCl5 : phosphorus pentachloride Remark: If the first element contains only one atom we don’t write mono before it; If the second element contains only one atom we have to write mono before it. Ex.: NO : nitrogen monoxide CO: carbon monoxide Exercise: Fill in the table below Name Chemical Formula diphosphorus pentoxide SO3 CO Aluminum sulfide Al2S3 SF6 Calcium oxide Lithium nitride Li3N Remark: If the compound contains a transition metal, then we have to mention the type of charge of the transition metal by inserting a roman numeral in brackets (I, II, III, IV ….) after the symbol of the transition metal. Example: Name the following compounds: - FeCl2 : Iron (II) chloride - Cu(NO3)2 : Copper (II) nitrate Remark: There are few transition metals that have only one type of charge such as zinc, nickel, and silver; in this case no roman numeral is required. Zinc : Zn+2 Silver: Ag+ Nickel: Ni+2 - ZnSO4 : zinc sulfate - AgNO3 : silver nitrate - CuSO4 : copper (II) sulfate - Co(NO3)3 : cobalt (III) nitrate Exercise: Write the formula of the following compounds. Calcium phosphate: Iron(III) hydroxide: Sodium hydroxide: Manganese(II) hydroxide: Barium sulfate: Zinc carbonate: Ammonium nitrate: Remark: We must enclose the polyatomic ion in brackets if the number after it is more than 1. Note: If the polyatomic ion that ends with the prefix –ate decreases by one oxygen atom then the prefix changes to -ite. If the prefix ending with – ite decreases by one oxygen atom then we precede the prefix by hypo, whereas if the prefix ending with – ate increases by one oxygen atom then we precede the prefix by per. Example: ClO3- is called chlorate; if we reduce one oxygen atom then the ion becomes ClO2- and is called chlorite, however if we increase by one oxygen atom then the ion becomes ClO4- and is called perchlorate; and if the chlorite is reduced by one oxygen atom then the ion becomes ClO- and is called hypochlorite. Exercise: Name the following compounds: K2SO3 : NaNO2 : Mg(ClO4)2 : LiBrO2 : - Comparison Table between ionic and covalent (molecular) compounds: property Ionic Compounds Covalent Compounds State (at room temperature) solids Solids, liquids or gases Melting point and boiling point Very high Usually low Involvement of electrons Loss and gain (transfer) of electrons Sharing of electrons Electric conductivity When dissolved in water (in solution) , electric conductivity is high