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CHEM 110 QUANTITATIVE CHEMISTRY Chapter One Introduction to Matter and Measurement Dr V Paideya 2014 Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia • Intranet http://cheminnerweb.ukzn.ac.za/Firstyear/ oneten.aspx • Mastering Chemistry https://secure.ecollege.com/ukznmlp/inde x.learn?action=welcome Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia School of Chemistry, University of KwaZulu-Natal, Westville Campus, Durban CHEM110: General Principles of Chemistry Worksheet 1 Matter, Measurements and Molecules AIM: • To provide a background to understanding the properties of matter in terms of atoms, molecules and ions including scientific measurements CONTENT: Units, significant figures and scientific notation, basic nomenclature, atoms and molecules, elements and compounds, atomic structure and isotopes LEARNING OBJECTIVES - You should be able to: • Distinguish between elements, compounds and mixtures • Recognise symbols of common elements and common prefixes for units • Use significant figures, scientific notation and SI units. Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia CHEMISTRY “study of matter & changes it undergoes” “matter is anything that has mass and takes up space” - study of physical & chemical properties of matter - what changes occur in these properties, in the course of/as the result of a chemical reaction, & how these changes may be observed - why the reaction involved does (or doesn’t…) occur be able to understand & explain such macroscopic changes from an atomic/molecular (submicroscopic) perspective States (Phases) of Matter - solid, H2O(s); liquid, H2O(l); gas, H2O(g) - phase transitions occur @ specific P/T values, governed by properties of atoms/molecules Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia Matter - atoms are building blocks of matter - each element is made of same kind of atom/molecules* - compounds made of two or more different kinds of elements bonded together Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia Pure Substances, Elements & Compounds pure substance -has distinct properties & unvarying/constant composition eg. NaCl(s), H2O(l), HCN(g) element -substance that cannot be decomposed into simpler substances eg. Cl2(g), Br2(l), I2(s); Ne(g), Hg(l), Au(s) compound -substance composed of 2 or more different elements 2 or more different kinds of atoms eg. UF6(g), H2O(l), CaCO3(s) Law of Constant Composition/Definite Proportions (Joseph Proust ca 1800) “...elemental composition of pure substance is always the same…” - different samples of pure compound have the same elemental composition - elements present in such samples have same proportion by mass Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia Classification of Matter Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia Mixtures - combination of 2 or more substances, in which each substance retains own chemical identity & can thus be separated from each other - 2 types: heterogeneous: - mixture of visibly different composition, properties or appearance eg. sand in H2O(l) (s, l), sand & NaCl (s, s), petrol & H2O(l) (l, l) homogeneous: - mixture of visibly uniform composition, properties & appearance throughout eg. NaCl(aq) (s,l), air (g,g), stainless steel (s,s), soda water (g,l) Properties of Matter: - physical: measurement without changing identity/composition eg. Change in state, temperature, volume - chemical: must involve change in chemical identity eg. Combustion, oxidation - extensive: dependent on quantity of sample involved eg. mass, volume - intensive: independent of quantity eg. , colour, m.p.; useful for identification of substances Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia Physical and Chemical Changes Separation of Mixtures Figure 1.6 In the course of a chemical reaction, the reacting substances are converted to new substances. Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia Separation of Mixtures 1. Distillation Separates a homogeneous mixture on the basis of differences in boiling point. Figure 1.8 Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia Separation of Mixtures 2. Filtration Separates solid substances from liquids and solutions. 3. Chromatography Separates substances on the basis of differences in solubility in a solvent. Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia The Scientific Method A systematic approach to solving problems Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia SI Units Système International d’Unités Uses a different base unit for each quantity Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia Metric System Prefixes convert the base units into units that are appropriate for the item being measured. Tera T 1012 Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia SI Units - Temperature • The Kelvin is the SI unit of temperature. • It is based on the properties of gases. • There are no negative Kelvin temperatures. • K = C + 273.15 Figure 1.10 Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia Derived SI Units Volume • The most commonly used metric units for volume are the litre (L) and the millilitre (mL). – A litre is a cube 1 dm long on each side. – A millilitre is a cube 1 cm long on each side. 1dm3 = (1 dm) x (1 dm) x (1 dm) = 10 cm x 10 cm x 10 cm = 1000 cm3 = 1 L Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia Derived SI Units Density Density is a physical property of a substance and is determined through the following formula: mass density = volume or symbolically m = V Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia Common Volumetric glassware Different measuring devices have different uses and different degrees of accuracy. Figure 1.12 Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia Uncertainty in Measurement Precision and Accuracy • Accuracy refers to the proximity of a measurement to the true value of a quantity. • Precision refers to the proximity of several measurements to each other. Figure 1.15 Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia Significant Figures • All digits of a measured quantity, including the uncertain, are called significant figures. • The greater the number of significant figures, the greater the certainty of the measurement. Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia Significant Figures To determine s.f in a measurement read no. from left to right, counting the digits starting with first digit that is not zero All nonzero digits are significant, e.g. 123.45 Zeros between two significant figures are themselves significant, e.g. 103.405 Zeros at the beginning of a number are never significant, e.g. 00123.45 = 123.45 Zeros at the end of a number are significant if a decimal point is written in the number, e.g. 123.450 has six significant figures but If no. ends in zero but contains no decimal can be problem – exponential notation used to indicate if zeros at the end are significant, e.g. 1.03 x 104 g Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia Handling Significant Figures in Calculations Addition and Subtraction The answer has the least no. of digits to the right of the decimal pt .in comparison to the original nos. e.g.(a) 89.332 + 1.1 = 90.432 = 90.4 e.g. (b) 2.097 – 0.12 = 1.977 = 1.98 Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia Rounding off Procedure 1. Drop off the digit that follows if it is less than 5 e.g. 8.724 -- 8.72 2. Add 1 to the preceding digit if it is equal or greater than 5 e.g. 8.727 -- 8.73 Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia Multiplication and Division The number of significant figures in the final product or quotient is determined by the original number that has the smallest no. of significant figures e.g. 1. 2.8 x 4.5039 = 12.61092 = 13 e.g. 2. 6.85 ÷ 112.04 = 0.0611388789 = 0.0611 Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia Infinite number of significant figures Exact numbers by definition or counting nos. of objects can be considered to have an infinite no. of significant figures. If an object has a mass of 0.2786, then the mass of 8 such objects = 0.2786 x 8 = 2.229 g NB. We don’t round off this product to one significant fig. because 8 is 8. 0000 ..... Similarly the average of two measured lengths 6.64 cm and 6.68cm = (6.64 + 6.68) ÷ 2 = 6.66 cm Bec. 2 is 2.0000 ....... Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia Significant Figures 1. any figure that is not zero is significant: 845 mL _____ s.f.1243.29 mg _____ s.f. 2. zeroes between non-zero figures are significant: 1906 mL _____ s.f.40501.09 J _____ s.f. 3. exact (“counting”) numbers by definition have an number of s.f., so physical constants defined to be exact numbers do so also...:1 atm 101.325 kPa 760 mmHg; 0 OC 32 OF 273.15 K all _____ s.f. 4. leading zeroes (to the left of the first non-zero figure) are not significant: 0.008 kg _____ s.f. 0.003798 L _____ s.f. 5. trailing zeroes (to the right of the last non-zero figure) are significant only if the number has a d.p.: 300.0 mm _____ s.f.0.0300 mm _____ s.f. 6. in measurements without a d.p., the number of s.f. is ambiguous: 1200mm ?? OR either: i) use scientific notation ii) 1200.0 Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia Using Significant Figures in Calculations all calculations governed by two fundamental rules multiplication/division number of s.f. in final answer is the LEAST of numbers of s.f. in each of original measurements addition/subtraction number of d.p. in final answer is the LEAST of numbers of d.p. in each of original measurements Eg. 1 Calculate i) volume, in mm3, of a box of length 6.741 cm, breadth 2.441 x 10-1 m, & height 4.2 mm i) 6.9 x 104 mm3; ii) density () of a pure liquid, in g cm-3, if 103.67 g of it is needed to fill the box completely ii) 1.5 g cm-3 Eg. 2 An empty container of mass 23.29 g has a mass of 86.1 g when filled with 0.5000 dm3 of a pure liquid. Determine the of this liquid in g cm-3. 0.126 g cm-3 Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia CHAPTER TWO ATOMS, MOLECULES AND IONS Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia Early Atomic Theory John Dalton 1803 - 1807 -each element is composed of very small, indestructible, particles called atoms* -all atoms* of given element are physically & chemically identical to each other, but atoms of a particular element are different from atoms of all other elements Law of Conservation of Mass -atoms are neither created or destroyed in chemical reactions - mass reactants present @ start = mass products formed @ completion* Law of Constant Composition -different samples of a pure compound have the same elemental composition -elements present in such samples have same proportion by mass at the end of a chemical process as before the process took place. Law of Multiple Proportions -if 2 elements (C & O) can combine to form 2 or more different compounds (CO & CO2), the different masses of one element (O) combining with a fixed mass of the other (C) can be expressed as a simple integral ratio Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia Atomic Theory Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia The Law of Multiple Proportions • Was deduced by Dalton from the preceding four postulates and states that: • If two elements A and B combine to form more than one compound, the masses of B that can combine with a given mass of A are in the ratio of small whole numbers. Examples • H2O consists of 2 hydrogens and 1 oxygen • H2O2 consists of 1 hydrogen and 1 oxygen Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia The Discovery of Atomic Structure The Electron (JJ Thomson, 1897) - electrical discharges from cathode originally thought to be new form of radiation - showed that radiation emitted was - independent of cathode material used - deflected by magnetic/electric fields - findings consistent with model in which “beam” /”rays” composed of negatively charged “particles”(-) with charge/mass ratio of 1.78 x 108 C g-1 Electron Charge & Mass (Robert Millikan, 1909) - oil drop experiment - (-) charge on oil drops found always to be a multiple of minimum value of 1.6(02) x 10-19 C ie. 1.602 x 10-19 C must be charge of single electron (e-) -mass of single e- determined to be 9.109 x 10-28 g - only 1/1836 of mass of an H atom - first subatomic particle Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia Radioactivity The spontaneous emission of radiation by an atom was first observed by Henri Becquerel. It was also studied by Marie and Pierre Curie. Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia Radioactivity Three types of radiation were discovered by Ernest Rutherford – particles – particles – rays Figure 1.21 Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia Discovery of the Nucleus Ernest Rutherford shot particles at a thin sheet of gold foil and observed the pattern of scatter of the particles . The Nuclear Atom Some particles were deflected at large angles. This led Rutherford to postulate that the atom had a nucleus with positively charged Particles called protons.Protons Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia Modern Atomic Structure - more than 99.99 % of atom mass & entire Q+ centred in atomic nucleus, where nucleons (protons, p+ (Rutherford, 1919) & neutrons, nO (Chadwick, 1932) are collectively bound together by strong nuclear force - atomic nucleus surrounded by much larger atomic volume, containing as many e- as p+, so atom is electrically neutral & held together by force of Coulombic/ electrostatic attraction - Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia Atomic (Z) & Mass (A) Numbers - atoms of different elements have different numbers of p+ in their nuclei mass number (A) number of p+ & nO Z AE element symbol atomic number (Z) number of p + (number of e - in neutral atom) Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia Isotopes Atoms with identical atomic numbers (Z) but different mass numbers (A), or atoms with the same number of protons which differ only in the number of neutrons are called isotopes. Examples: 11 6C 12 6C carbon-12 isotope 13 6C 14 6C carbon-14 isotope Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia Isotopes - atoms of same element having d different numbers of nO in their nuclei ie. same Z, different A, or same Z, different N - chemical properties largely similar, but physical properties, & particularly the ones involving radioactive “nuclei”, can be very different - each Mg atom is one of three naturally occurring isotopes 25Mg; 26Mg - 24Mg; Mg-24 Mg-25; Mg-26 Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia Complete the following table: Element Eg. 3 CompleteSymbol the table below: name p+ (Z) 79 No e- A 118 138 Ba 143 Pb 92 126 35 Krypton 46 36 Experimentally.. High Resolution Mass Spectrometry (p. 39) used for very precise (4-6 d.p.; 7-10 s.f. in total) measurements of the masses of an element’s isotopes & their naturally occurring abundances Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia Atomic Mass Atomic and molecular masses can be measured with great accuracy with a mass spectrometer. Figure 1.23 Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia Average Atomic Mass (commonly called Atomic Mass) • We use average masses in calculations, because we use large amounts of atoms and molecules in the real world. • Average atomic mass is calculated from the fractional abundance of each isotope and mass of that isotope. For example, the average atomic mass of C made up mostly of 12C (98.93%) and 13C (1.07%) - is 12.01 u. • extremely small SI masses of individual atoms (~4 x 10-22 g) too awkward for everyday usage, so masses expressed in unified atomic mass units (amu, u): 1 amu (u) = 1.66054 x 10-24 g 1 g = 6.02214 x 1023 amu (u) Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia Average Atomic Masses of Naturally Occurring Elements -use average masses in “real world” calculations, as even smallest weighable sample (~1 g 10-6 g) involves large (~1015, or 10 quadrillion) numbers of atoms -no AAMs calculated as “weighted average” of an element’s isotopic masses (IMs) & naturally occurring abundances AAM = (IM x % ab/100) or (IM x fr ab) Eg. 4 a)Naturally occurring Mg has three isotopes: Calculate its AAM. b) Naturally occurring Pb has four isotopes: Calculate its AAM. 24Mg 78.99 %, 23.9850 u 25Mg 10.00 %, 24.9858 u 26Mg 11.10 %, 25.9826 u 204Pb 1.40 %, 203.973037 u 206Pb 24.10 %,205.974455 u 207Pb 22.10 %, 206.975885 u 208Pb 52.40 %, 207.976641 u Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia Eg. 4 a) Naturally occurring Mg has three isotopes: Calculate its AM. 24Mg 25Mg 26Mg 78.90 %, 23.9850 u 10.00 %, 24.9858 u 11.10 %, 25.9826 u AM = (IM x % ab/100) = {[23.9850 x (78.90/100)] + [24.9858 x (10.00/100)] + [25.9826 x (11.10/100)] } = 18.924 + 2.4986 + 2.8841 = 24.306 u b) Naturally occurring Pb has four: 204Pb 1.40 %, 203.973037 u 206Pb 24.10 %, 205.974455 u 207Pb 22.10 %, 206.975885 u 208Pb 52.40 %, 207.976641 u Calculate its AM. AM = (IM x % ab/100) = 2.856 + 49.640 + 45.742 + 108.98 = 207.22 u Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia Eg. 5 Chlorine has two naturally occurring stable isotopes: 35Cl 37Cl 34.968853 u 36.965803 u If the (average) atomic mass of naturally occurring elemental Cl is 35.453 u, what are the % abundances of the two isotopes? Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia The Periodic Table - rapid post-Dalton growth in experiment-based chemical knowledge showed very quickly that many elements could be grouped together on basis of similarities in their physical & chemical properties - arrangement of elements in order of Z showed that these similarities occurred in repetitive/periodic patterns, & agreed so closely with experimentally acquired data, that phys/chem properties of 2 “missing” elements were accurately predicted before their being reported as formally discovered, &/or phys/chem properties characterized Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia The Periodic Table When one looks at the chemical properties of elements, one notices a repeating pattern of reactivities. Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia Periodic Table • The rows are called periods. • The columns are called groups. • Elements in the same group have similar chemical properties. • Nonmetals are on the right side of the periodic table (with the exception of H). • Metalloids border the stair-step line (with the exception of Al and Po). • Metals are on the left side of the chart. Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia Groups Table 1.7 The above five groups are known by their names. Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia The Periodic Table Metals, Non-Metals, & Metalloids Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia Molecules and Chemical Formulae The subscript to the right of the symbol of an element tells the number of atoms of that element in one molecule of the compound. Notice how the composition of each compound is given by its chemical formula. Figure 1.29 Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia Diatomic Molecules Figure 1.28 These seven elements occur naturally as molecules containing two atoms. Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia Molecular Compounds Molecular compounds are composed of molecules and almost always contain only nonmetals. Types of Formulae Empirical formulae give the lowest whole-number ratio of atoms of each element in a compound, e.g. HO. Molecular formulae give the exact number of atoms of each element in a compound, e.g. H2O2. Structural formulae show which atoms are attached to which within the molecule, e.g. H-O-O-H. Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia Picturing Molecules Different representations of the methane (CH4) molecule. Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia Ions and Ionic Compounds • When atoms lose or gain electrons, they become ions. – Cations are positive and are formed by elements on the left side of the periodic chart. – Anions are negative and are formed by elements on the right side of the periodic chart. Anion formation + ee.g. Cl atom ----------------------> Cl ion (anion) 17 protons 17 protons 17 electrons 18 electrons Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia Ionic Compounds Ionic compounds (such as NaCl) are generally formed between metals and nonmetals. Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia Using Ionic charge to write empirical Formulae • Because compounds are electrically neutral, one can determine the formula of a compound by: – writing the value of the charge on the cation as the subscript on the anion. – writing the value of the charge on the anion as the subscript on the cation. Note: if the subscripts are not in the lowest whole number ratio, simplify it, e.g. Ca2O2 would become CaO. Ex. What is the empirical formula of the compound formed by (a) Al3+ and S2- and (b) Zn2+ and PO42Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia Chemical Nomenclature Positive Ions (Cations) a) Cations formed from metal atoms have the same name as the metal, e.g. Na+ is the sodium ion. b) If a metal can form different cations, the positive charge is indicated by a Roman numeral in parentheses following the name of the metal, e.g. Au+ is the gold(I) ion and Au3+ is the gold(III) ion. c) Cations formed from nonmetal atoms have names that end in -ium, e.g. NH4+ is the ammonium ion. Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia Chemical Nomenclature Common Cations Table 2.4 Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia Chemical Nomenclature Negative Ions (Anions) a) The names of the monatomic anions are formed by replacing the ending of the name of the element with ide, e.g. O2- is the oxide ion. b) Polyatomic anions containing oxygen (called oxyanions) have names ending in -ate or -ite, e.g. SO42- is the sulfate ion and SO32- is the sulfite ion. c) Anions derived by adding H+ to an oxyanion are named by adding the prefix hydrogen or dihydrogen, e.g. HCO3is the hydrogen carbonate ion. Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia Chemical Nomenclature Common Anions Table 2.5 Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia Chemical Nomenclature More on naming oxyanions Examples: ClO4perchlorate ion (one more O atom than chlorate) ClO3 chlorate ClO2 chlorite ion (one less O atom than chlorate) ClO hypochlorite ion (one O atom less than chlorite) Names of ionic compounds consist of the cation followed by the anion name, e.g. Cu(ClO4)2 is copper(II) perchlorate, and CaCO3 is calcium carbonate. Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia Chemical Nomenclature Name and Formulae of Acids 1. Acids containing anions whose names end in -ide are named by changing the -ide ending to -ic, adding the prefix hydro- to this anion name, and then following with the word acid. 2. Acids containing anions whose names end in -ate or -ite are named by changing the -ate ending to -ic and the -ite ending to -ous and then adding the word acid. Figure 2.22 Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia Chemical Nomenclature Binary Molecular Compounds 1. The name of the element farther to the left in the periodic table is written first. 2. If both elements are in the same group in the periodic table, the one having the higher atomic number is written first. 3. The name of the second element is given an -ide ending. 4. Greek prefixes are used to indicate the number of atoms of each element. Examples N2O4 is dinitrogen tetroxide P4S10 is tetraphosphorus decasulfide. Table 2.6 Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia Derivatives of Alkanes Alcohol An alcohol is an example of an alkane that has some hydrocarbon groups replaced with an –OH group. Methanol (CH3OH) Ethanol (C2H5OH) Propanol (C3H7OH) Octanol (C8H17OH) Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia Naming of simple Organic Compounds Alkanes Compounds contain only carbon and hydrogen and are called hydrocarbons. Simplest class of hydrocarbons are alkanes- names of compounds end in -ane Methane (CH4) Ethane (C2H6) Propane (C3H8) Octane (C8H18) Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia EXERCISE • Name the following ionic compounds: (a) MgO (b) AlCl3 (c) (NH4)2SO4 • Write Chemical Formulae for the following compounds: (a) Copper(I) oxide (b) Sodium hydoxide (c) Zinc sulphate Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia