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UNIT 1 – PHYSICAL SCIENCE AND CHEMICAL REACTIONS CHAPTER 5: CHEMICALS IN ACTION chemistry: the study of matter matter: anything that has mass and occupies space MATTER mixtures SUBSTANCES PURE COMPOUNDS ELEMENTS molecular ionic metal non-metal pure substance: is one in which all the particles that make up the substance are the same; has constant properties. Atoms are in a fixed ratio. eg. (H2O(ℓ)) mixture: contains two or more pure substances not in a fixed ratio eg. sugar + water sweet water (solute) + (solvent) (solution) element: a pure substance composed of only one type of atom (eg. Iron (Fe)) Cannot be broken down by ordinary chemical means (eg. heating, electricity, filtration, distillation, magnets) Metals Non- metals Found on left hand side of staircase line of P.T. Shiny Dull Ductile, malleable Brittle R.H.S. Hard soft Ductile – the ability to be formed into a wire Malleable – the ability to be stretched by hammering or rolling compound: pure substance composed of at least two types of elements (or atoms) eg. water (H2O(ℓ)); salt (NaCl(s)) Ionic Molecular Composed of 2 or more ions ( metal + non-metal) 2 or more non-metallic atoms Held together by an ionic bond covalent bond (e.g.) NaCl- sodium chloride CO2 – carbon dioxide The Modern Periodic Table Facts about hydrogen Group IA 1 valence electron Loses 1 electron to form a hydrogen ion Group VIIA 1 less electron than a Noble Gas A gas Can gain 1 electron to form a hydride ion Periodic Table Demitri Mendeleev created the first useable periodic table It became the most predictive device in all chemistry Elements are arranged in increasing order of Atomic Number (AN) The periodic table is divided into vertical columns (families/groups) and horizontal rows (periods) Atomic Number: (AN) the number of protons (p+) in the nucleus of an atom Group A Elements IA ~ Alkali Metals Soft, reactive IIA ~ Alkaline Earth Metals When compared to IA; harder or not as reactive IIIA ~ Aluminum Group IVA-VIA ~ Name according to first element in group VIIA ~ Halogens Very reactive VIIIA ~ Noble (inert) Gases All gases Non-reactive Group B Elements (Transition Elements) When compared to Group A elements; harder and more colourful, with higher boiling and melting points metalloids: located on both sides of the staircase line and have characteristics of metals and nonmetals Matter has a well defined underlying structure composed of the following: 1. Atoms The smallest whole part of an element that is still representative of the element A neutral particle composed of a nucleus containing protons (p+) and neutrons, and electrons (e-) outside the nucleus (the number of electrons equals the number of protons) 2. Ions A charged atom (The charge is caused by the transfer of an electron (or electrons) from one atom to another) The metallic atom (which loses an electron) has a positive charge and is called a cation - eg. Na+, Al3+ The non-metallic atom (which gains an electron) has a negative charge and is called an anion - eg. Clˉ, N3ˉ 3. Molecules Occurs when non-metallic atoms share electrons (which then bond together into a group) - eg. Methane (CH4) H . . This is a Lewis Dot Diagram. represent molecules. C : H : C It is used to . . H Development of Atomic Theory 1. Dalton “Billiard Ball” Model (1803) States that the atom is composed of uniform matter, not divisible (cannot be broken down into smaller parts) 2. Thompson “Raisin Bun” Model (1897) Evidence from experiments (electricity through gases in a vacuum tube) led Thompson to believe that negative (-ve) charged particles were scattered throughout the positive atom like raisins in a bun He called these negatively charged particles electrons. (He was the first) 3. Rutherford’s “Nuclear” Model (1914) By releasing (+ve) alpha particles at a thin gold foil Rutherford was astonished that a small percentage were deflected at large angles. He suggested that the atom must have a small, dense, positively charged inner core called a nucleus. He called these positively charged particles inside the nucleus protons. (He was the first) Outside the nucleus was a large area of mostly empty space containing electrons This work led one of Rutherford’s students (H. G. Mosely) to discover that the positive charge inside the nucleus increased by one from element to element in Mendeleev’s periodic table. From this he coined the phrase atomic number (the number of protons inside the nucleus of an atom). This discovery led to modifications of Mendeleev’s periodic table – instead of listing the elements in order of increasing atomic mass it would now be in order of increasing atomic number The mass of the protons was too small to account for the total atomic mass of the atom so Rutherford predicted that there must be a neutral particle in the nucleus similar in mass to the proton In 1932 James Chadwick demonstrated this particle, which he called a neutron. (He was the first to name this particle) 4. Bohr “ Orbit” Model (1921) Electrons surround the nucleus of an atom in circular orbits, with each electron having a fixed quantity of energy Electrons cannot exist between orbits but can move to unfilled orbits The higher the energy level, the further from the nucleus the electron is The maximum numbers of electrons in the first three orbits are 2, 8, and 8 respectively (Octet Rule) An atom with the maximum number of electrons in the outermost energy level is stable (unreactive) ep+ e- 5. Quantum Mechanical Model (1920’s) Similar to Bohr’s model, except that the electrons are said to be found in orbitals (not orbits) (an orbital is an area of space where an electron is likely to be found) p+ Other main features include: - Atomic Number = Number of protons - Number of protons = Number of electrons - An energy level represents a specific value of energy of an electron and corresponds to a general location - Period number = Number of energy levels occupied by electrons - The first three energy levels will have 2, 8, and 8 structures of electrons NOTE: 2, 8, and 8 should be written as follows started -8e-- -8e-- -8e-- -2e-- -2e-- -2e-- Each lower energy level must be filled to its maximum before the next level is - The electrons in the highest energy level are called valence electrons - Group number = Number of valence electrons (one exception is helium) (Group A only) - Max. number of electrons in each energy level = Max. number of atoms in each period Energy Level Diagrams For Atoms valence level: the outer-most energy level of an energy level diagram NOTE: It is important to include atom when writing the name ex. Lithium atom (Students will be expected to format all energy level diagrams for atoms as follows) Valence Electrons (Valence Level): - 1e-- 2e-Atomic Number: 3p+ Name: Lithium atom Symbol: Li Energy Level Diagrams For Ions ion: a charged atom Formed when a metal loses electrons to form a positive ion (cation) or when non-metals gain electrons to form a negative ion (anion) All atoms gain/lose electrons to become like the nearest noble gas (become stable) When naming ions: Metal (cation) names: stay the same Non-metal (anion) names: change (last three letters become “ide”) ex. Lithium ion (Students will be expected to format all energy level diagrams for ions as follows) Valence Electrons (Valence Level): - 2e-3p+ Atomic Number: Name: Symbol: Li+ Lithium ion ex. Fluoride ion (The atom Fluorine becomes the ion Fluoride) - 8e-- Valence Electrons (Valence Level): - 2e-Atomic Number: 9p+ Name: Fluoride ion Fˉ Symbol: Isotopes Atomic number = Number of protons in the nucleus of an atom Atomic mass (mass number) = Sum of the protons and neutrons in the nucleus of an atom isotopes: forms of an element that have the same number of protons but different numbers of neutrons Two Ways to Write Isotopes ex. Argon Atomic mass (AM) ex. Argon Atomic number (AN) argon-40 40 18 Ar Atomic mass Retrieving Information About Isotopes Atomic mass is given in both naming forms Atomic number is either given or must be looked up on the Periodic Table Number of protons = Atomic number Number of electrons = Number of protons Number of neutrons = Atomic mass minus atomic number (n = AM – AN) Molecular Formulas chemical formulas: a group of symbols representing the number or type of atoms/ions in a chemical substance ex. methane CH4 (molecular) ex. salt NaCl (ionic) Molecular Formula The chemical formula for a binary (two element) molecular compound Composed of two non-metals held together by the sharing of electrons forming a covalent bond NOTE: Non-metals are to the right of the staircase line on the Periodic Table When Naming Molecular Compounds: Prefixes must be used for naming both non-metals (memorize) mono: 1 tri: 3 penta: 5 hepta: 7 di: 2 tetra: 4 hexa: 6 octa: 8 nona: 9 deca: 10 The first name stays the same but the last three letters of the second name change to “ide” - ex. N2O6 dinitrogen hexaoxide NOT dinitrogen hexaoxygen Subscripts indicate prefixes. It also tells the number of atoms in a molecule. If the subscript for the first name of a compound is (1) then the prefix is omitted. - ex. CO2 carbon dioxide NOT monocarbon dioxide * if the name of a compound is given the prefixes become subscripts (e.g) trisulfur tetraoxide becomes S3O4 The following groups do not follow these rules: ( Memorize ) Trivial or Common Names methane CH4 (g) ethanol C2H5OH(ℓ) ethane C2H6 (g) water H2O(ℓ) propane C3H8 (g) hydrogen peroxide H2O2(ℓ) butane C4H10 (g) hydrogen sulfide H2S(g) glucose C6H1206 (s) ozone O3(g) sucrose C12H22011 (s) silicon carbide SiC(s) methanol CH3OH(ℓ) ammonia NH3(g) Diatomic Molecules All Group VII molecules (H2, F2, Cl2, Br2, I2, At2), oxygen (O2), and nitrogen (N2) Polyatomic Molecules sulfur S8 (s) phosphorous P4 (s) Ionic Compounds Formed when a metal is joined to a non-metal (formed when a cation [+ charged atom] is joined to to an anion [- charged atom]) Held together by an ionic bond caused by the transfer of electrons Ionic Compounds: ( Properties ) Are all solids at room temperature All dissolve in water to some degree - ie. Are aqueous (aq) Will conduct electricity, but only when dissolved in water - ie. Are good electrolytes Have relatively high melting point and boiling point-higher than molecular compounds How to Write Ionic Compounds When metals (cations) and non-metals (anions) join, the net electrical charge of an ionic compound must be zero. - ie. The sum of all positive charges must equal the sum of all negative charges Never change the charge on an ion from the Periodic Table - To find the net charge, multiply the charge by using a subscript Never Never Never Never use the prefix system - The first name of the ionic compound stays the same - The second name of the ionic compound changes the last three letters to “ide” All subscripts in an ionic compound must be in lowest terms NOTE: When the subscript is (1) it is not written in - When the charges of both the cation and the anion are the same, simply leave out the subscripts - When the charges of the cation and the anion are different, write the charge of the cation as the subscript of the anion and vice versa ( criss-cross ) NOTE: non-metal + non-metal = molecular compound NOTE: metal(write first) + non-metal(write second) = ionic compound Multivalent Metals or (the Stock System) Multivalent metals are metals that have more than one charge Of the charges, the top one listed is the most common. When using multivalent metals, if no charge is mentioned (if you are not told which charge to use) use the most common one. To write the name of a compound which includes a multivalent metal, you must write the ion charge (on the metal only) in Roman numerals in brackets. Always start with the anion (non-metal) when given the chem. formula eg. Fe2O3: iron (lll) oxide eg. SnCl4: tin (lV) chloride eg. nickel (ll) oxide: NiO eg. platinum (lV) sulfide: PtS2 NOTE: Know Roman numerals I – X Polyatomic Ions (Complex Ions) A group of atoms that become stable by losing/gaining electrons When naming, they are treated identically to monatomic ions Put polyatomic ions in brackets when using subscripts most end in ( ate ) or ( ite ) all are negative except for ammonium ( NH4+ ) eg. lithium sulfate: LiSO4ˉ eg. calcium nitrate: Ca(NO3)ˉ2 eg. Fe2CO3: iron (l) carbonate Complex ions cannot exist alone Compounds which contain complex ions are ionic compounds are solids in pure form but break into ions in water. NOTES: Break down of compounds ~ molecular compound molecules atoms ionic compound formula units ions Hydrates Hydrates are ionic compounds that decompose at relatively low temperatures to produce water and an associated compound - eg. CuSO4 • 5H2O(s) If heated 5 molecules of water are released and copper (ll) sulfate remains The CuSO4 is now said to be anhydrous (without water) Naming Hydrates: Two Methods prefix system always written as one word CuSO4 • 5H2O(s) the ‘hydrate’ is the H2O means ‘in the presence of’ naming . . . is named as copper (ll) sulfate pentahydrate This is the preferred method of CuSO4 • 5H2O(s) . . . can also be named as copper (ll) sulfate-5-water Hydrogen Compounds and Acids hydrogen compounds: Different from most molecular compounds because they form conducting solutions. Most acids contain hydrogen atoms bonded to some non-metal (eg. HCl) or complex ion (eg. H2SO4). Nomenclature of H-Compounds Most are named as acids except for the pure compounds (before dissolving) The pure compounds are named as though they were ionic - eg. HCl(g): hydrogen chloride - eg. HCN(g): hydrogen cyanide - eg. H2S(g): hydrogen sulfide Nomenclature of Acids H-compounds must be dissolved in water before an acid can be formed NOTE: Any solution made when an ionic/molecular compound dissolves in water is said to be aqueous and the subscript (aq) is added - eg. HCl(g) + water HCl(aq) hydrogen chloride + water hydrochloric acid In naming acids, you first name the H-compound as if it is ionic, and then convert to the acid name by following these three rules: IONIC NAME ACID NAME hydrogen (root)ide hydro(root)ic acid hydrogen (root)ate (root)ic acid hydrogen (root)ite (root)ous acid NOTE: Two exceptions to these rules are acids which have sulf or phosph as roots. Syllables are added to these acid names to make them sound better. Thus, H2SO4(aq) is sulfuric acid not sulfic acid and H3PO4(aq) is phosphoric acid not phosphic acid. Classification of Acids binary: if the acid contains a H-compound with hydrogen and one other kind of atom eg. HCl(aq): hydrochloric acid oxo: if the acid contains a H-compound with hydrogen, oxygen, and one other kind of atom eg. HClO3(aq): chloric acid NOTE: If an acid contains a COO- group then the (H) atom is placed at the end of the group, not in the front eg. CH3COOH(aq) (not HCH3COO(aq)): acetic acid (vinegar) Properties of Acids 1. are solids, liquids, and gases as pure substances at room temperature (like molecular substances) 2. are soluble in water (like all ionic and some molecular substances) 3. form coloured and colourless solutions (like ionic compounds) 4. form conducting solutions (like ionic compounds) 5. form solutions which turn blue litmus red 6. taste sour and are usually corrosive (burn) 7. pH lower than 7 (0-6.9-acid; 7-neutral;7.1-14-base) 8. will neutralize basic solutions (acid+base = salt +water) 9. react with alkali metals to produce salt and H2 gas ex. 2K(s) + 2HCl(aq) 2KCl(s) + 2H2(g) salt: any ionic compound that is not an acid or a base Chemical Reactions Occur when old bonds within reactants are broken and new bonds within products are produced - ie. reactants products Why does matter chemically react? Atoms want to become stable. They share electrons to form a covalent bond (molecular compound) and transfer electrons to form an ionic bond (ionic compound). How can you tell when a chemical reaction occurs? You cannot observe it directly, but there are five general observation that indicate a chemical reaction: An unexpected colour change A gas is produced A solid (precipitate) is produced or dissolved A new odour is detected A temperature or energy change - If energy (heat, sound, light, etc.) is absorbed then it is an endothermic reaction. - If energy reaction. (heat, sound, light, etc.) is released then it is an exothermic Can a chemical reaction create or use up energy? No. The law of conservation of energy states that in any chemical process energy is neither created nor destroyed, just transformed from one form to another. - eg. Electrical energy from batteries produces heat and light Balancing Chemical Equations Law of Conservation of Atoms: states that in any chemical reaction atoms are neither created nor destroyed but conserved in the reaction The total number of atoms on the reactant side must equal the total number of atoms on the products side. (Must also have the same type of atoms.) 3H2O(ℓ) unit subscript: gives the number of atoms in a molecule or ions in a formula coefficient: gives the number of molecules/formula units in a compound Always change the coefficient; never change the subscript once it is balanced for charge All coefficients must be in lowest terms and must be whole numbers Na(s) + HOH(ℓ) NaOH(aq) + H2(g) reactant reacts with reactant to yield/produce product along with product Na(s) + HOH(ℓ) NaOH(aq) + H2(g) (unbalanced) Na – 1 Na - 1 H–2 H-3 O–1 O–1 2Na(s) + 2HOH(ℓ) 2NaOH(aq) + H2(g) (balanced) Na – 2 Na - 2 H–4 H-4 O–2 O–2 Types of Chemical Reactions 1. Simple Decomposition (SD) Occurs when a compound is broken down into it’s elements General Form: compound element1 + element2 + element3 . . . eg. hydrogen chloride gas is decomposed into it’s elements HCl(g) H2(g) + Cl2(g) eg. sodium hydroxide is decomposed 2NaOH(s) 2Na(s) + O2(g) + H2(g) eg. sodium phosphate 4Na3PO4(s) 12Na(s) + P4(s) + 8O2(g) 2. Simple Composition (SC) Occurs when elements react with each other to form a compound General Form: element1 + element2 + element3 . . . compound eg. glucose is formed from it’s elements 6C(s) + 6H(g) + 3O(g) C6H12O6(s) NOTE: Molten means liquid 3. Single Replacement Occurs when an element and a compound react to form a different element and compound These reactions involve ions (Remember that a cation must bond with an anion) Write water as HOH(ℓ) Use the most common charge when more than one is possible General Form: element1 + compound1 element2 + compound2 eg. 2NaCl(aq) + Br2(ℓ) 2NaBr(aq) + Cl2(g) eg. 2Na(g) + 2HOH(ℓ) 2NaOH(aq) + H2(g) 4. Double Replacement (DR) Occurs when two compounds react and form two new compounds The cation of compound1 reacts with the anion of compound2 and the cation of compound2 reacts with the anion of compound1 One of the reactants must be (l) or (aq) General Form: compound1 + compound2 compound3 + compound4 eg. 2AgBr(aq) + ZnCl2(l) 2AgCl(g) + ZnBr2(aq) eg. Pb(NO3)2(aq) + 2KI(s) PbI2(aq) + 2KNO3(aq) NOTE: If the state of matter for an ionic reactant is unknown, assume it’s aqueous NOTE: Coefficients must always be whole numbers 5. Hydrocarbon Combustion Occurs when a hydrocarbon (a compound composed of carbon and hydrogen) reacts with oxygen gas to produce carbon dioxide gas and water vapour General Form: hydrocarbon + O2(g) CO2(g) + H2O(g) In general, balance carbon first, hydrogen second, and oxygen third eg. methane combusts in the air CH4(g) + 2O2(g) CO2(g) + 2H2O(g) eg. 2C2H6 + 7O2(g) 4CO2(g) + 6H2O(g) 6. Other Any reaction that is not one of the first five Always be given both reactants and products eg. CO2(g) + Ca(OH)2(aq) CaCO3(s) + H2O(ℓ)