CHEM 110 QUANTITATIVE CHEMISTRY. Chapter One. Introduction to Matter and Measurement. Dr V Paideya 2014

Transcription

1 CHEM 110 QUANTITATIVE CHEMISTRY Chapter One Introduction to Matter and Measurement Dr V Paideya 2014

2 Intranet oneten.aspx Mastering Chemistry x.learn?action=welcome

3 AIM: School of Chemistry, University of KwaZulu-Natal, Westville Campus, Durban CHEM110: General Principles of Chemistry Worksheet 1 Matter, Measurements and Molecules 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.

4 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, H 2 O(s); liquid, H 2 O(l); gas, H 2 O(g) - phase transitions specific P/T values, governed by properties of atoms/molecules

5 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

6 Pure Substances, Elements & Compounds pure substance -has distinct properties & unvarying/constant composition eg. NaCl(s), H 2 O(l), HCN(g) element -substance that cannot be decomposed into simpler substances eg. Cl 2 (g), Br 2 (l), I 2 (s); Ne(g), Hg(l), Au(s) compound -substance composed of 2 or more different elements 2 or more different kinds of atoms eg. UF 6 (g), H 2 O(l), CaCO 3 (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

7 Classification of Matter

8 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 H 2 O(l) (s, l), sand & NaCl (s, s), petrol & H 2 O(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

9 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.

10 Separation of Mixtures 1. Distillation Separates a homogeneous mixture on the basis of differences in boiling point. Figure 1.8

11 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.

12 The Scientific Method A systematic approach to solving problems

13 SI Units Système International d Unités Uses a different base unit for each quantity

14 Metric System Prefixes convert the base units into units that are appropriate for the item being measured. Tera T 10 12

15 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 Figure 1.10

16 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. 1dm 3 = (1 dm) x (1 dm) x (1 dm) = 10 cm x 10 cm x 10 cm = 1000 cm 3 = 1 L

17 Derived SI Units Density Density is a physical property of a substance and is determined through the following formula: density = mass volume or symbolically = m V

18 Common Volumetric glassware Different measuring devices have different uses and different degrees of accuracy. Figure 1.12

19 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

20 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.

21 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 Zeros between two significant figures are themselves significant, e.g Zeros at the beginning of a number are never significant, e.g = Zeros at the end of a number are significant if a decimal point is written in the number, e.g 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 x 10 4 g

22 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) = = 90.4 e.g. (b) = = 1.98

23 Rounding off Procedure 1. Drop off the digit that follows if it is less than 5 e.g Add 1 to the preceding digit if it is equal or greater than 5 e.g

24 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 x = = 13 e.g = =

25 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 , then the mass of 8 such objects = x 8 = g NB. We don t round off this product to one significant fig. because 8 is Similarly the average of two measured lengths 6.64 cm and 6.68cm = ( ) 2 = 6.66 cm Bec. 2 is

26 Significant Figures 1. any figure that is not zero is significant: 845 ml s.f mg s.f. 2. zeroes between non-zero figures are significant: 1906 ml s.f 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 kpa 760 mmhg; 0 O C 32 O F K all s.f. 4. leading zeroes (to the left of the first non-zero figure) are not significant: kg s.f 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.: mm s.f mm s.f. 6. in measurements without a d.p., the number of s.f. is ambiguous: 1200mm?? either: i) use scientific notation OR ii)

27 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 mm 3, of a box of length cm, breadth x 10-1 m, & height 4.2 mm i) 6.9 x 10 4 mm 3; ii) density ( ) of a pure liquid, in g cm -3, if g of it is needed to fill the box completely ii) 1.5 g cm -3 Eg. 2 An empty container of mass g has a mass of 86.1 g when filled with dm 3 of a pure liquid. Determine the of this liquid in g cm g cm -3

28 CHAPTER TWO ATOMS, MOLECULES AND IONS

29 Early Atomic Theory John Dalton 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 start = mass products 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 & CO 2 ), the different masses of one element (O) combining with a fixed mass of the other (C) can be expressed as a simple integral ratio

30 Atomic Theory

31 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 H 2 O consists of 2 hydrogens and 1 oxygen H 2 O 2 consists of 1 hydrogen and 1 oxygen

32 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 10 8 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 C ie x C must be charge of single electron (e - ) -mass of single e - determined to be x g - only 1/1836 of mass of an H atom - first subatomic particle

33 Radioactivity The spontaneous emission of radiation by an atom was first observed by Henri Becquerel. It was also studied by Marie and Pierre Curie.

34 Radioactivity Three types of radiation were discovered by Ernest Rutherford particles particles rays Figure 1.21

35 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

36 Modern Atomic Structure - more than % of atom mass & entire Q + centred in atomic nucleus, where nucleons (protons, p + (Rutherford, 1919) & neutrons, n O (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 -

37 Atomic (Z) & Mass (A) Numbers - atoms of different elements have different numbers of p + in their nuclei mass number (A) number of p + & n O atomic number (Z) number of p + (number of e - in neutral atom) Z A E element symbol

38 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 6 C 12 6 C 13 6 C 14 6 C carbon-12 isotope carbon-14 isotope

39 Isotopes - atoms of same element having d different numbers of n O 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 - 24 Mg; 25 Mg; 26 Mg Mg-24 Mg-25; Mg-26

40 Complete the following table: Element Eg. 3 name Complete Symbol the table below: p+ (Z) No e- A Ba Pb Krypton 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

41 Atomic Mass Atomic and molecular masses can be measured with great accuracy with a mass spectrometer. Figure 1.23

42 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 u. extremely small SI masses of individual atoms (~4 x g) too awkward for everyday usage, so masses expressed in unified atomic mass units (amu, u): 1 amu (u) = x g 1 g = x amu (u)

43 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 (~10 15, 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. 24 Mg %, u 25 Mg %, u 26 Mg %, u b) Naturally occurring Pb has four isotopes: 204 Pb 1.40 %, u 206 Pb %, u Calculate its AAM. 207 Pb %, u 208 Pb %, u

44 Eg. 4 a) Naturally occurring Mg has three isotopes: 24 Mg %, u 25 Mg %, u Calculate its AM. 26 Mg %, u AM = (IM x % ab/100) = {[ x (78.90/100)] + [ x (10.00/100)] + [ x (11.10/100)] } = = u b) Naturally occurring Pb has four: 204 Pb 1.40 %, u 206 Pb %, u 207 Pb %, u Calculate its AM. AM 208 Pb %, u = (IM x % ab/100) = = u

45 Eg. 5 Chlorine has two naturally occurring stable isotopes: 35 Cl u 37 Cl u If the (average) atomic mass of naturally occurring elemental Cl is u, what are the % abundances of the two isotopes?

46 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

47 The Periodic Table When one looks at the chemical properties of elements, one notices a repeating pattern of reactivities.

48 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.

49 Groups Table 1.7 The above five groups are known by their names.

50 The Periodic Table Metals, Non-Metals, & Metalloids

51 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

52 Diatomic Molecules Figure 1.28 These seven elements occur naturally as molecules containing two atoms.

53 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. H 2 O 2. Structural formulae show which atoms are attached to which within the molecule, e.g. H-O-O-H.

54 Picturing Molecules Different representations of the methane (CH 4 ) molecule.

55 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 + e- e.g. Cl atom > Cl ion (anion) 17 protons 17 protons 17 electrons 18 electrons

56 Ionic Compounds Ionic compounds (such as NaCl) are generally formed between metals and nonmetals.

57 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. Ca 2 O 2 would become CaO. Ex. What is the empirical formula of the compound formed by (a) Al 3+ and S 2- and (b) Zn 2+ and PO 4 2-

58 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 Au 3+ is the gold(iii) ion. c) Cations formed from nonmetal atoms have names that end in -ium, e.g. NH 4 + is the ammonium ion.

59 Chemical Nomenclature Common Cations Table 2.4

60 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. O 2- is the oxide ion. b) Polyatomic anions containing oxygen (called oxyanions) have names ending in -ate or -ite, e.g. SO 4 2- is the sulfate ion and SO 3 2- is the sulfite ion. c) Anions derived by adding H + to an oxyanion are named by adding the prefix hydrogen or dihydrogen, e.g. HCO 3 - is the hydrogen carbonate ion.

61 Chemical Nomenclature Common Anions Table 2.5

62 Chemical Nomenclature More on naming oxyanions Examples: - ClO 4 perchlorate ion (one more O atom than chlorate) - ClO 3 chlorate - ClO 2 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(ClO 4 ) 2 is copper(ii) perchlorate, and CaCO 3 is calcium carbonate.

63 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

64 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 N 2 O 4 is dinitrogen tetroxide P 4 S 10 is tetraphosphorus decasulfide. Table 2.6

65 Derivatives of Alkanes Alcohol An alcohol is an example of an alkane that has some hydrocarbon groups replaced with an OH group. Methanol (CH 3 OH) Ethanol Propanol (C 2 H 5 OH) (C 3 H 7 OH) Octanol (C 8 H 17 OH)

66 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 (CH 4 ) Ethane (C 2 H 6 ) Propane (C 3 H 8 ) Octane (C 8 H 18 )

67 EXERCISE Name the following ionic compounds: (a) MgO (b) AlCl 3 (c) (NH 4 ) 2 SO 4 Write Chemical Formulae for the following compounds: (a) Copper(I) oxide (b) Sodium hydoxide (c) Zinc sulphate