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Bonding and Reactions A compound is two or more elements chemically combined. Compounds are formed from elements in chemical reactions. The forces that hold atoms together are called bonds. Atoms form bonds in two main ways: (1) An ionic bond is a bond formed by the transfer of electrons (between oppositely charged ions). Example: Na+1 + Cl-1 -- NaCl (1) Covalent bonding is bonding involving a sharing of electrons. A bond is a shared pair of electrons. Example: C + 4H --- H C H A molecule is the smallest part of a covalent compound that retains all of its characteristics. The term formula unit is used to describe the smallest part of an ionic compound that retains all of the characteristics of the compound. A formula is used to represent a single molecule or formula unit of a compound. Example: H2O – 2 atoms of hydrogen with one atom of oxygen A subscript is a small number indicating the number of atoms or ions of each element that are present in a compound. The subscript one is understood. An atom that gains or loses electrons is in a charged condition and is called an ion. Atoms that form negative ions gain electrons (anions). Example: O-2 Atoms that form positive ions lose electrons (cations). Example: Na+1 An ion made of more than one atom is a polyatomic ion. Examples: OH-1, SO4-2, NO3-1 (Refer to the complete list on your worksheet or in your text. You must learn the names, formulas, and charges.) When an atom or group of atoms has an overall charge, that charge is called an oxidation number. Oxidation numbers must be used to write a formula correctly. Rules For Determining Oxidation Numbers (1) Metals (positive) in groups IA to IIIA have an oxidation number equal to the group number. Example: Ca is in group IIA so it is +2. (2) Many B group metals have more than one charge and must be memorized. Example: Sn can be +2 or +4. (3) Nonmetals in groups IVA to VIIA form ions with negative charges equal to the group number – 8. Example: O is in groupVIA , so 6-8 = -2 (4) Polyatomic groups – refer to chart on worksheet or in your text Atoms and ions combine in definite whole number ratios. If the charges on the ions forming a compound add to zero, the formula is written correctly. In a binary compound (made of 2 elements) the positive ion is written first and the negative ion is written second. Example: Ca+2 + Br –1 ----- CaBr2 Br –1 Ca+2 + Br –1 --- CaBr2 (calcium bromide) Example: Al +3 + S -2 Al2S3 (aluminum sulfide) Non-binary compounds contain at least 3 elements. Polyatomic ions must be written in parenthesis, when expressing more than one. Example: Al +3 + SO4-2 Na+1 + NO3-1 Al2(SO4)3 NaNO3 Naming Compounds (1) For binary compounds: (A) Write the name of the element with the positive charge. (B) Add the name of the negative element modified to end in -ide. Examples: NaCl – sodium chloride Al2S3 – aluminum sulfide Roman numerals are used in naming compounds that contain elements with more than one oxidation number. Examples: Cu2S – copper (I) sulfide CuS – copper (II) sulfide FeF2 – iron (II) fluoride FeF3 – iron (III) fluoride PbCl2 – lead (II) chloride PbCl4 – lead (IV) chloride An older system for naming cations having more than one possible charge uses an –ic ending for the ion with the higher charge and an –ous ending for the ion with the lower charge. In addition, the Latin names for some of the elements are used. Examples: stannous chloride – SnCl2 cupric oxide – CuO cobaltous fluoride – CoF2 auric chloride – AuCl3 Occasionally, covalent compounds are named using a system in which Greek prefixes are used to indicate the number of atoms present. Mono – 1 Di – 2 Tri – 3 Tetra – 4 Penta – 5 Hexa – 6 Hepta – 7 Octa – 8 Nona – 9 Deca – 10 Hints: (1) The prefix mono- may be omitted for the first element (PCl3 – phosphorus trichloride). (2) For oxides, the ending “a” or “o” in the prefix is sometime omitted (N2O4 – dinitrogen tetroxide). Examples: CO2 –carbon dioxide CO - carbon monoxide CCl4 – carbon tetrachloride SO3 – sulfur trioxide Some compounds have names ending in –ide and are not binary. These compounds contain polyatomic groups. Examples: OH-1 – hydroxide CN-1 – cyanide O2-2 – peroxide In naming a compound containing more than 2 elements, write the name of the first element or polyatomic group (NH4+1). Then write the name of the second element (-ide ending) or the polyatomic group (do not change the ending). Polyatomic ions that contain oxygen end in -ite or –ate. Examples: CuSO4 –copper(II) sulfate ZnCO3 – zinc carbonate NH4Cl – ammonium chloride BaOH – barium hydroxide Chemical equations – are used to represent chemical reactions. The starting substances are called reactants. The substances formed are called products. Example: hydrochloric acid and sodium hydroxide react to form sodium chloride and water HCl(aq) + NaOH(aq) - NaCl(aq) + H2O(l) phase labels (physical state symbols) (g) – gas (l) – liquid (cr) – crystalline solid (aq) – aqueous (water solution) (s) – solid There are 3 steps to writing an equation correctly: (1) Determine what the reactants and products are. (2) Assemble the parts of the equation: (a) Write the formulas correctly. (b) Connect reactants and products using an arrow to show the direction of the reaction. ------- means yields or produces (c) Reactants are written to the left, products on the right. (3) Balance the equation by changing the coefficients (stoichiometric coefficients). Never change subscripts. Example: Propane reacts with oxygen to give carbon dioxide and water. (1) reactants – propane and oxygen (2) products – carbon dioxide and water (3) formulas/equation – C3H8(g) + O2(g) ---- CO2(g) + H2O(l) (4) balance – to get the same number of atoms of each element on both sides of the equation C3H8(g) + 5 O2(g) ---- 3 CO2(g) + 4 H2O(l) Practice: Balance the following (1) __ Al2O3 -- __ Al + __ O2 (2) __ Na2S + __ O2 -- __ Na2O + __ SO2 (3) __ C + __ H2O -- __ H2 + __ CO (1) 2 Al2O3 -- 4 Al + 3 O2 (2) 2 Na2S + 3 O2 -- 2 Na2O + 2 SO2 (4) C + H2O -- H2 + CO (balanced) Four Types of Reactions (1) Single displacement – one element displaces another in a compound M + M1Nm -- MNm + M1 Nm + MNm1 --MNm + Nm1 Examples: Cl2 + 2 KBr -- 2 KCl + Br2 Na + AgCl -- NaCl + Ag (2) Double displacement – the positive and negative portions of 2 compounds are interchanged MNm + M1Nm1 --- MNm1 + M1Nm Examples: PbCl2 + Li2SO4 -- 2 LiCl + PbSO4 ZnBr2 + 2 AgNO3 -- Zn(NO3)2 + 2 AgBr (3) Decomposition – substances break into simpler substances when energy (heat, light, electricity) is supplied A -- B + C + D Examples: H2CO3 -- H2O + CO2 2 KClO3 -- 2 KCl + O2 (4) Synthesis – 2 or more substances combine to form one new substance B + C + D -- A Examples: 2 H2 + O2 -- 2 H2O NH3 + HCl -- NH4Cl Law of Conservation of Matter Matter can neither be created nor destroyed, but can be converted from one form to another. Therefore, the mass of the reactants in an ordinary chemical reaction must equal the mass of the products. All chemical reactions involve changes in energy. The law of conservation of energy states that the total energy in a system is constant. If the energy of the products is less than the energy of the reactants, energy is released during the reaction. This is an exothermic reaction. reactants -- products + energy example: the combustion of propane If the energy of the products is greater than the energy of the reactants, energy must be absorbed during the reaction. This is an endothermic reaction. Example: instant ice pack Molecular mass is the sum of the atomic masses for all elements (atoms) in a compound. Examples: NaCl 1 Na – 23 1 Cl – 35.5 58.8 AMU Al(OH)3 1 Al – 27 3 O – 48 (3 x 16) 3 H – 3 (3 x 1) 78 AMU