MATTER and MEASUREMENTS [MH5; Ch. 1] What is Chemistry? The Study of Matter and its Properties at the Molecular Level. What is Matter? Matter is anything that has mass and occupies space. It can exist in three phases... Solid: - fixed shape and volume Liquid: - fixed volume, but takes the shape of its container Gas: - takes both the shape and volume of its container Classification of Matter Matter Pure Substances Mixtures Homogeneous Elements Compounds Mixtures Heterogeneous Mixtures 1
Pure Substances Pure substances have a fixed composition and a unique set of properties; these properties distinguish it from other substances. Elements Cannot be broken down into two or more pure substances. Identified by a symbol; some of which make sense: C is Carbon, Ba is Barium, Ar is Argon Some don t make sense (in English, anyway): Na is Sodium, Fe is Iron, Hg is Mercury Compounds A pure substance that contains two or more elements. In any given compound, the proportions of each element are always the same. EXAMPLE: Hydrocarbons are compounds containing Carbon and Hydrogen. Methane is CH 4 Propane is C3H8 Octane is C H 8 18 The properties of compounds are different from those elements which make up the compound... EXAMPLE: Ethanol, C2H5OH, is the alcohol in alcoholic beverages. Ethanol is a liquid and it contains: 52.2 % Carbon (a solid) (These % s are given 13.0 % Hydrogen ( a gas) by masses.) 34.8 % Oxygen (a gas) 2
Mixtures Two or more substances combined together; each substance retains its own chemical identity. The substances in a mixture do not react with each other. Two types of mixtures... Homogeneous Composition is the same throughout the mixture. May also be called a solution, which consists of one or more solutes dissolved in a solvent: * The solute is usually a solid, and the smaller amount. * The solvent is usually a liquid, and the larger amount. Alloys are solid solutions which are usually metallic in nature: * An example is brass, which contains copper (Cu); 67-90 % and zinc (Zn); 10-33 % Heterogeneous Composition varies throughout the mixture. Rocks are a common example...often many different minerals are readily observable. * Granite rock is often reddish in colour with regions of white and black interspersed throughout...these different colours are due to the different minerals. 3
What does the term properties mean? Chemical properties tell us how a substance will participate in chemical reactions...will it change into a different substance or substances? For example, Hydrogen gas (H 2) is explosive; Helium gas (He) is unreactive. Chemical properties also tell us how many bonds a certain element will form when it makes compounds in chemical reactions. Physical properties can be useful for identification; these may be observed without chemical reaction. Things like colour, melting point, boiling point and density of a substance are all physical properties. They do not depend on the amount of substance present. Symbols of State We describe the physical state of a material by attaching the appropriate one of the following: (s) solid e.g. H2O(s) [ice] ( ) liquid e.g. H2O ( ) [liquid water] (aq) aqueous e.g. NaC (aq) [NaC in solution] (g) gaseous e.g. O 2 (g) [oxygen gas] (g,at) a gaseous atom e.g. O (g,at) [gaseous O atom} 4
Chemistry is both a qualitative and quantitative science... Qualitative aspects How do we prepare X? What happens when X reacts with Y? How can we convert X into Z? How can we separate X and Y? Quantitative aspects involve numbers. How much X reacts with 1 g of Y? How much Z is produced when X reacts with Y? How fast does X react with Y? How much energy is evolved? Quantitative measurements need units. We use mainly SI units (le Système International; meaning metric!) [MH5; 1.2] Units for each type of measurement and calculation will be discussed when applicable... Examples of some one dimensional units include: Mass (grams, g;milligrams, mg; kilograms, kg; etc) Length (metres, m; centimetres, cm; millimetres, mm; kilometres,km) 3 Volume (litres, L; millilitres, ml, cubic centimetres, cm or cc) Other units are multi dimensional...for example; Density (mass per unit of volume) Concentration (grams or moles of solute per volume of solution) 5
Significant Figures [MH5; 1.2] An indication of the reliability of an experimental measurement as most experimental results contain some error or uncertainty. We have two different ways of expressing this uncertainty... Accuracy: How close is the result agree with the true value? Precision: How well do repeated measurements agree with each other? The number of significant figures used in reporting a result gives us an estimate of the limits of error/uncertainty. The error is assumed to be ± 1 in the last (least significant) digit; the extreme right-hand digit. EXAMPLES: 2.35 g really means: 2.35 ± 0.01 g 2.3568 g really means: 2.3568 ± 0.0001 g Are zeros significant? Sometimes! 2.350 has 4 sig figs: this means 2.350 ± 0.001 - so a trailing zero has significance. But zeroes are not significant when used to fix the position of the decimal point, i.e. when leading: 2 0.0235 has only 3 sig figs (2.35 x 10 ) What about 235,000? When a number is this large, it is unlikely that we mean exactly 235,000; the trailing zeros are really only indicating the power of 10. It would be much better would be to write it in exponential or scientific notation: 235,000 = 2.35 x 10 5 6
Note that we often use defined or exact numbers for which the concept of significant figures does not arise; there is no uncertainty: EXAMPLES: 1 m = 100 cm (exactly) 1 L = 1000 ml (exactly) 1 kg = 1000 g (exactly) Coefficients in equations are exact numbers: P 4 + 6 Cl 2 4 PCl 3 (exact) Significant figures may change in a calculation, and how they change depends on what mathematical operations are performed. A discussion of errors and carrying significant figures through calculations appears in the Lab Manual, but this is a more sophisticated treatment than we need to use at the moment. The following method is described in MH5; pages 11-13. Addition and Subtraction of Measured Quantities Find the quantity with the fewest figures to the right of the decimal point. The number of figures to the right of the decimal point in your answer should be the same as this. There will likely be some rounding off needed at some point during the calculation! Addition: 2.3568 Subtraction: 4.15 + 0.026 3.0815 2.383 1.07 In the addition example, 2.3568 could be rounded to 2.357 before the calculation, or 2.3828 could be rounded to 2.383 afterwards. In the subtraction example, 3.0815 could be rounded to 3.08 before the calculation, or 1.0685 could be rounded to 1.07 afterwards. 7
Multiplication and Division of Measured Quantities Find the quantity with the fewest significant figures; this is the number of significant figures you should have in your answer. Multiplication: 17.15 0.0977 = 1.675555 ±? 17.15 has 4 sig figs and 0.0977 has 3 sig figs Round off after the calculation to get the answer 1.68, with 3 significant figures. If you rounded off before the calculation, you could introduce round off errors. Division: 4.383 g 2.72 L = 1.611397 ±? 4.383 has 4 sig figs and 2.72 has 3 sig figs Although your calculator may give 10 (!!) sig figs, only 3 are justified: 1.61 g L 1 Never copy numbers off your calculator without thinking about them!! We will look at significant figures again when we begin examples involving calculations and in our discussions of laboratory pre-lab exercises. 8