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Unit 1 Introduction: Matter and Measurement

CHM 1045 : General Chemistry and Qualitative Analysis. Unit 1 Introduction: Matter and Measurement. Dr. Jorge L. Alonso Miami-Dade College – Kendall Campus Miami, FL. Textbook References : Module #1. Scientific Method:. A systematic approach to solving problems.

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Unit 1 Introduction: Matter and Measurement

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  1. CHM 1045: General Chemistry and Qualitative Analysis Unit 1Introduction:Matter and Measurement Dr. Jorge L. Alonso Miami-Dade College – Kendall Campus Miami, FL • Textbook References: • Module #1

  2. Scientific Method: A systematic approach to solving problems. Observation: the detection of a phenom enon by our sensory organs or their extensions (instruments). Scientist study CAUSE  EFFECT Relationships Which factors affect the behavior of gasses? η,P, V,& T {GasVariables} Hypothesis: initial or tentative explanation of the causes of a phenomenon. Experiment: carefully designed hypothesis testing; done by controlling all the variables except your suspected CAUSE (independent variable, x) which is manipulated in order to observe its EFFECT (dependent variable, y).

  3. Also, P ↑V ↓ = k 1 P k P V V= Experiment: carefully designed hypothesis testing; done by controlling all the variables except your suspected CAUSE (independent variable, x) which is manipulated in order to observe its EFFECT (dependent variable, y). OR Dependent Variable Boyle’s Law Independent Variable Law: concise verbal or mathematical summary for a variety of observations and experiences. Theory: a comprehensive explanation for natural phenomena that has withstood repeated analysis and experimentation. Kinetic Molecular Theory

  4. What is the universe composed of? Matter: Anything that has mass and takes up space. • Chemicals • Substances • Things Chemistry Energy: the ability to perform an activity (work). Physics • Kinetic (motion & heat) • Electromagnetism (light, electricity & chemical bonds) • Nuclear • Gravity {Matter with Energy}

  5. Describing Matter: Physically and Chemically Chemistry: The study of matter and the changes it undergoes. (1) Physical Properties (3) Chemical Properties (2) Physical Change (4) Chemical Change (reactions & equations)

  6. Let’s get Physical! Physical Properties: The physical characteristics (appearance) of matter. State or phase (s,l,g), color, mass, volume, density, melting & boiling points, solubility, etc H2O (g) H2O (s) H2O (l) • Physical Changes: • Changes in the physical properties of matter. • Changes of: state (s l g), density, temperature, shape, volume, etc. • Physical changes occur without changes inthe composition of matter. Physical change: H2O (s)H2O (l)H2O (g) Chemical change: 2 H2O (s) 2 H2(g) + O2(g)

  7. Heat Energy& Phase (State) Changes (fusion) Phase Changes = Changes of State

  8. Kinetic Energy& States of Matter Heat = Kinetic Energy  Temperature c. pt. (condensation point) f. pt. (freezing point) m. pt (melting point) b. pt. (boiling point) Density (mass per unit volume) For H2O: m.pt. = f.pt = 0OC b.pt. = c.pt = 100OC For Methanol: m.pt. = f.pt = -98OC b.pt. = c.pt = 65OC {KineticMolecularTheory: PhaseChange}

  9. Heating Curve: Energy & Phase Changes Heating water vapor Heat of Vaporization Heating liquid water Heating solid ice Heat of Fusion For H2O: m.pt. = f.pt = 0OC b.pt. = c.pt = 100OC For Methanol: m.pt. = f.pt = -98OC b.pt. = c.pt = 65OC (∆Hv in kJ/mol)

  10. Let’s get Chemical! What happens when you add Na to water? K? • Chemical Properties:Can only be observed when a substance reacts and is changed into another substance. • Does it react? With which substance does it react? Flammability? Corrosiveness?, etc. {Na & K in H2O} • Chemical Changes: The changes that occur in the process of producing new substances.Combustion, oxidation, decomposition, etc. Is this chemical or physical change? Chemical change is always accompanied by physical change!

  11. Describing Chemical Change Chemical Reaction: the actual phenomenon that occurs when chemicals change in composition. Chemical Equation: a symbolic representation of a chemical reaction. Based on the Atomic Theory. Quiz Question: Write the balanced chemical equations for the reactions of (1) Sodium (Na) + Water (HOH) (2) Potassium (K) + Water

  12. Chemical Reactions and Equations + - + 2 Na + 2 HOH  2 NaOH + H2 2 K + 2 HOH  2 KOH + H2 (Hint: single displacement or replacement reaction)

  13. Matter and the Atomic Theory • Atomsare the building blocks of all matter. • Elements are made of the same kind of atom. • Compounds are made of two or more different kinds of atoms. • Mixtures are composed of different elements/compounds together.

  14. Heterogeneous * Classification of Matter Cu(NO3)2 Mixtures (Heterogeneous) Physical Separation Mixtures Homogeneous Solutions (Homogeneous) Physical Separation Cu(NO3)2 (aq) Cu(NO3)2 (s) Compounds Chemical Decomposition PureSubstances Elements

  15. Separatory Techniques • based on differences in physical properties of the substances present in the mixture/solution. methods of physically separating (purifying) substances from mixtures and solutions into pure substances. • Filtration – by solubility vs. insolubility • Metal Smelting & Refining- by differences in melting point (ability to form a liquid) • Distillation –by differences in boiling points (ability to form a gas) • Chromatography – by differences in degree of solubility

  16. (1) Filtration: Separates insoluble solid substances from liquids and solutions. Mixture: K2Cr2O7(s) +NaNO3(s) + H2O K2Cr2O7(s) NaNO3(aq)

  17. (2) Metal Smelting & Refining These techniques are used to differentially melt mixtures of metals (alloys) by means of their different melting points (ability to form a liquid when heated). The sweat furnace operates at a temp at which one metal is selectively melted from a component, leaving the metal with the higher melting point, usually a ferrous metal as a recoverable solid. Example: mixture of Cu + Zn, heated to 500°C

  18. * (3) Distillation: Separates homogeneous mixture on the basis of differences in boiling points (ability to form a gas). Substanceb.pt. Ethyl Alcohol 77oC Water 100oC Sodium Chloride 1413oC Solution: Alcohol +

  19. Distillation of Hydrocarbons: Petroleum Refinery Towers compounds composed of molecules arranged in a long chain of carbon atoms with hydrogen atoms attached to the carbon chain. • Name (b.pt. C) # C Structural Formula • Methane (-162) 1CH4 • Ethane (-89) 2CH3CH3 • Propane (-42) 3CH3CH2CH3 • Butane (-0.5) 4CH3CH2CH2CH3 • Pentane (36) 5 CH3CH2CH2CH2CH3 • Hexane (69) 6 CH3CH2CH2CH2CH2CH3 • Heptane (98) 7 CH3CH2CH2CH2CH2CH2CH3 • Octane (126) 8 CH3CH2CH2CH2CH2CH2CH2CH3 • Nonane (151) 9 CH3 CH2 CH2CH2CH2CH2CH2CH2CH3 • Decane (174) 10 CH3CH2CH2CH2CH2CH2CH2CH2CH2CH3

  20. Distillation of Hydrocarbons: Petroleum Refinery Towers 0 C 120 C 200 C 250 C 300 C

  21. Separation by differences in ability to form a gas (boiling points) Mixture / Solution or Pure Substance? H2O vapor Physical Separation Cu(NO3)2(s) • Mixturescan be separated by differences in the physical properties of the substances they are composed of. • Pure substancescannot be separated by physical methods.

  22. Separation by differential solubility, filtration and evaporation H2O vapor Mixture: CdS (yellow, insoluble substance), Cu(NO3)2 (blue soluble substance), H2O(clear liquid).

  23. (4) Chromatography: Separates substances on the basis of their differences in their solubility in a specific solvent. Filter paper Substance to be separated (black ink) Solvent: 50: 50 Water: Alcohol {Paper Chromatography}

  24. Heterogeneous * Classification of Matter Cu(NO3)2 Mixtures (Heterogeneous) Physical Separation Mixtures Homogeneous Solutions (Homogeneous) Physical Separation Cu(NO3)2 (aq) Cu(NO3)2 (s) Compounds Chemical Decomposition PureSubstances Elements

  25. Chemical Decomposition of Pure Substances • Cannot be separated by physical means. • Composed of one substance only, which can be either an element or a compound. • Compounds can be broken down by chemical means, elements cannot. Examples of pure substances: Gold (Au), Oxygen (O2), Water (H20), Methanol (CH3OH), Table salt (NaCl) Each has its specific physical properties (m. pt., density, etc.)

  26. Compounds Electrolysis of Water: can be broken down into more elemental particles (elements) by chemical decomposition reactions. 2 H2O (l)→ 2 H2(g) + O2(g) elect. {Electrolysis}

  27. How do we get pure Sodium? NaCl is electrolyzed in a Downs cell. • Gaseous Cl2 allowed to disperse • Molten Na siphoned off 2 NaCl (l)→2 Na(l) + Cl2(g) elect

  28. Elements cannot be broken down into more elemental particles by ordinary chemical means.

  29. Classification of Matter Heterogeneous Physical Separation Homogeneous Mixture Physical Separation Solution Chemical Decomposition {mixture vs. compound} Element Compound

  30. Units of Measurementlength (m)mass (g, kg)volume (mL, L)temperature (oC, oK)time (s)

  31. Metric System When using dimensional analysis for metric problems:always consider the larger unit as having a valueof1, then the smaller unit would contain a large multiple of that unit. X 1000 X 10 X 10 X 1000 Example: 1 m compared to cm.

  32. Atomic Dimensions Atoms Tenth of a nanometer (10 -9 m) Nuclei of atoms Hundredth of a picometer (10 -12 m) Protons & Neutrons Fentometer (10-15 m) Quarks & electrons Attometer (10-18 m)

  33. Metric Conversions Always convert PREFIXES to UNITS (not PREFIXES to other PREFIXES) Example: Mm compared to pm. meter, liter, gram Factors, ratios, equivalences. Example: cm compared to m.

  34. Metric Conversion Problems • How many pm are there in 0.0023 cm? • Change 60. mph to km/s. {Hint: 1 mi. = 1.6 km} • How many m3 of water are there in 25 ft3 ? 3 3 3

  35. 10 cm 10 cm Volume: Liter (L) and the milliliter (mL) 10 cm • A liter is a cube 1 dm3 = 10 cm long on each side. 1 L = dm3 = (10 cm)3 = (10 X 10 X 10) cm3 = 1000 cm3 = 1000 mL or 1mL = 1/1000 L Cubic centimeter • A milliliter (mL) is a cube 1 cm long on each side. = milliliter

  36. * Temperature: measure of the average kinetic energy (motion caused by heat) of the particles in a sample. {K.E ∝ Temp} T = change in temp As KE increases molecules vibrate more and their volume expands (Temp). • 373 • 273 • 100 100 - 0 100 • 212 • 32 • 180 C = (F − 32) 1.8 F = 1.8(C) + 32 K =C + 273.15

  37. Measured vs Exact Numbers • Measured Numbers: (1) Accuracy & Precision (2) Uncertainty (3) Significant figures & • rounding-off • Exact Numbers: from formulas, definitions & counting For sphere 1 mile = 5,280 ft 1 km = 1,000 m

  38. Measured Numbers: Accuracy versus Precision • Accuracyrefers to the proximity of a measurement to the true value of a quantity. • Precisionrefers to the proximity of several measurements to each other.

  39. Measured vs. Exact Numbers Measured numbers are obtained when a measuring instrument (ruler, balance, thermometer) is used to determine a physical property of a substance. 13.7 +0.1 uncertainty The number of significant figures these measurements contain depend on the accuracy of the instrument being used. 7.63 +0.01 uncertainty

  40. Uncertainty in Measurements +0.01 +0.1 uncertainty Different instruments have different degrees of accuracy, uncertainty is + 1 of estimated digit. 89.5 mL 2.65 mL

  41. Measuredvs. ExactNumbers METRIC METRIC-ENGLISH ENGLISH CONVERSIONS Exact Numbers Exact Numbers Measured Numbers (1 in is exact,the 2.54 cm is measured) How many km are there in a Marathon (26 miles)?

  42. Significant Figures • Significant figures refers to digits that were accuratelymeasured by an instrument. Example: 220g, 220. g, 220.5g, 220.50g, 220.507g. (all numbers above are measures of the same object, what is the difference?) accuracy

  43. Rules for determining the number of Significant Figures • All nonzero digits (NZD) are always significant. • Zeroesbetween NZD are always significant. Ex: 103 • Zeroes to the left of NZD are never significant. Ex: 0.0103 • Zeroes to the right of NZD are significant ifa decimal point is written anywhere in the number. Ex. 0.01030

  44. Rounding-off • Round-off your calculated numbers, to the correct number of significant figures, so we do not overstate the accuracy of our answers. • Example: 23g + 23.632g = 46.632 = 47g You cannot add an inaccurate measurement to a accurate measurement and get and accurate answer.

  45. * Significant Figures in Addition & Subtraction • When addition or subtraction is performed, answers are rounded to the least significant decimal place. Example: add the following numbers 34 231.678 0.00354 265.68154 266

  46. * Significant Figures in Multiplication & Division • Answers are rounded to the number of digits that corresponds to the least number of significant figures in any of the numbers used in the calculation. Example: (29.2 – 20.0) (1.79 x 105) 1.39 (29.2 – 20.0) = 9.2 Calculator answer = 1.1847482 x 106 Correct answer = 1.2 x 106

  47. Uncertainty in Measurements • Piece of Black Paper – with rulers beside the edges: Determine the Area of Black Paper! Let’s look more accurately ! Area = Length x Width

  48. Uncertainty in Measurements • Piece of Paper – Side A enlarged • How long is the paper to the best of your ability to measure it? 13.6 cm + 0.1 cm When using an instrument your last digit recorded should be a significant digit estimated between the two smallest measurement lines of your instrument. Your precision would be + 1 of that digit.

  49. Uncertainty in Measurements • Piece of Paper Side B – enlarged • How wide is the paper to the best of your ability to measure it? 7.63 cm + 0.01 cm When using an instrument your last digit recorded should be a significant digit estimated between the two smallest measurement lines of your instrument. Your precision would be + 1 of that digit.

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