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Lab Safety

Lab Safety. Lab Safety Begins Before You Go to the Lab!. Always read through the lab instructions the day before you go to the lab. Ask any questions you may have concerning the lab BEFORE you begin. Proper Lab Behavior.

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Lab Safety

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  1. Lab Safety

  2. Lab Safety Begins Before You Go to the Lab! • Always read through the lab instructions the day before you go to the lab. • Ask any questions you may have concerning the lab BEFORE you begin.

  3. Proper Lab Behavior • Never indulge in horseplay or behavior that could lead to injury of others.

  4. Dressing for Lab • gloves, and lab aprons when instructed to do so. (** Note: Contact lenses should not be worn in the lab!) • Due to the dangers of broken glass and corrosive liquid spills in the lab, open sandals or bare feet are not permitted in the lab.

  5. Safe Lab Practices • Learn the location and proper usage of the eyewash fountain, fire extinguisher, safety shower, fire alarm box, office intercom button, evacuation routes, clean-up brush and dust pan, glass/chemical disposal can.

  6. Safe Lab Practices (cont.) • Never look directly into a test tube. View the contents from the side. • Point test tubes that are being heated away from you and others. • Never taste any material in the lab. • Food, drink and gum are prohibited in lab.

  7. Safe Lab Practices (cont.) • Never smell a material in a test tube or flask directly.

  8. In the Event of a Lab Accident…… • Report all accidents regardless of how minor to your teacher. • For minor skin burns, immediately plunge the burned area into cold water and notify the instructor.

  9. In the Event of a Lab Accident…… (cont.) • If you get any chemical in your eye, immediately wash the eye with the eye-wash fountain and notify the instructor. • Immediately notify the teacher of any chemical spill and clean up the spill as directed.

  10. At the End of Your Lab Time… • Return all lab materials and equipment to their proper places after use. • Dispose of all chemicals AS DIRECTED BY YOUR instructor! • Wash and dry all equipment, your lab bench and your clean-up area.

  11. Remeber the concepts

  12. Weight/Volume Percent • Amount of solute = mass of solute in grams • Amount of solution = volume in milliliters • Express concentration as a percentage by multiplying ratio by 100% = weight/volume percent or % (W/V) 6.2 Concentration Based on Mass

  13. Calculating Weight/Volume Percent Calculate the percent composition or % (W/V) of 2.00 x 102mL containing 20.0 g sodium chloride 20.0 g NaCl, mass of solute 2.00 x 102mL, total volume of solution % (W/V) = 20.0g NaCl / 2.00 x 102 mL x 100% = 10.0% (W/V) sodium chloride 6.2 Concentration Based on Mass

  14. Calculate Weight of Solute from Weight/Volume Percent Calculate the number of grams of glucose in 7.50 x 102 mL of a 15.0% solution 15.0% (W/V) = Xg glucose/7.50 x 102mL x 100% Xg glucose x 100% = (15.0% W/V)(7.50 x 102mL) Xg glucose = 113 g glucose 6.2 Concentration Based on Mass

  15. Weight/Weight Percent • Weight/weight percent is most useful for solutions of 2 solids whose masses are easily obtained • Calculate % (W/W) of platinum in gold ring with 14.00 g Au and 4.500 g Pt [4.500 g Pt / (4.500 g Pt + 14.00 g Au)] x 100% = 4.500 g / 18.50 g x 100% = 24.32% Pt 6.2 Concentration Based on Mass

  16. 6.3 Concentration of Solutions: Moles and Equivalents • Chemical equations represent the relative number of moles of reactants producing products • Many chemical reactions occur in solution where it is most useful to represent concentrations on a molar basis

  17. Molarity • The most common mole-based concentration unit is molarity • Molarity • Symbolized M • Defined as the number of moles of solute per liter of solution 6.3 Moles and Equivalents

  18. Calculating Molarity from Moles • Calculate the molarity of 2.0 L of solution containing 5.0 mol NaOH • Use the equation • Substitute into the equation: MNaOH = 5.0 mol solute 2.0 L solution = 2.5 M 6.3 Moles and Equivalents

  19. Calculating Molarity From Mass • If 5.00 g glucose are dissolved in 1.00 x 102mL of solution, calculate molarity, M, of the glucose solution • Convert from g glucose to moles glucose • Molar mass of glucose = 1.80 x 102 g/mol 5.00 g x 1 mol / 1.80 x 102g = 2.78 x 10-2 mol glucose • Convert volume from mL to L 1.00 x 102mL x 1 L / 103mL = 1.00 x 10-1 L • Substitute into the equation: Mglucose = 2.78 x 10-2 mol glucose 1.00 x 10-1 L solution = 2.78 x 10-1M 6.3 Moles and Equivalents

  20. Dilution Dilution is required to prepare a less concentrated solution from a more concentrated one • M1 = molarity of solution before dilution • M2 = molarity of solution after dilution • V1 = volume of solution before dilution • V2 = volume of solution after dilution 6.3 Moles and Equivalents moles solute = (M)(L solution)

  21. Dilution • In a dilution will the number of moles of solute change? • No, only fewer per unit volume • So, • Knowing any three terms permits calculation of the fourth 6.3 Moles and Equivalents M1V1 = M2V2

  22. Calculating Molarity After Dilution • Calculate the molarity of a solution made by diluting 0.050 L of 0.10 M HCl solution to a volume of 1.0 L • M1 = 0.10 Mmolarity of solution before dilution • M2 = XMmolarity of solution after dilution • V1 = 0.050 L volume of solution before dilution • V2 = 1.0 L volume of solution after dilution • Use dilution expression • XM = (0.10 M) (0.050 L) / (1.0 L) 0.0050 M HCl OR 5.0 x 10-3 M HCl 6.3 Moles and Equivalents M1V1 = M2V2

  23. Representation of Concentration of Ions in Solution Two common ways of expressing concentration of ions in solution: • Moles per liter (molarity) • Molarity emphasizes the number of individual ions • Equivalents per liter (eq/L) • Emphasis on charge 6.3 Moles and Equivalents

  24. Comparison of Molarity and Equivalents 1 M Na3PO4 • What would the concentration of PO43- ions be? • 1 M • Equivalent is defined by the charge • One Equivalent of an ion is the number of grams of the ion corresponding to Avogadro’s number of electrical charges 6.3 Moles and Equivalents

  25. Experiment 1 Determination of pka for a weekacid

  26. Acids and Bases • Acid – increases the [H+] of a solution • Base – decreases the [H+] of a solution • May form OH- ion or absorb H+ ions • Strong acid/base • Weak acid/base

  27. The Buffer • What do we mean by BUFFER

  28. Carbonic Acid – Bicarbonate Buffer System Respiratory component Renal component CO2 +H2O H2CO3H + + HCO3– (H2CO3 is a ‘volatile’ acid as  CO2 exhaled ) Principal buffer system in the body. Provides 95% of buffering capacity in plasma.

  29. The importance of pH control • The pH of the ECF remains between 7.35 and 7.45 • If plasma levels fall below 7.35 (acidemia), acidosis results • If plasma levels rise above 7.45 (alkalemia), alkalosis results • Alteration outside these boundaries affects all body systems • Can result in coma, cardiac failure, and circulatory collapse

  30. Weak acids - HA H + A + Weak acids Conjugate base

  31. Henderson–Hasselbalch equation - + HA H + A Ka , acid dissociation constant (also known as acidity constant, or acid-ionization constant)

  32. Henderson–Hasselbalch equation log log log log

  33. Henderson–Hasselbalch equation X log log + log log log

  34. Henderson–Hasselbalch equation - - + log log pKa = -log Ka pH = -log

  35. pH

  36. Titration a method of estimating the amount of solute in a solution. The solution is added in small, measured quantities to a known volume of a standard solution until a reaction occurs, as indicated by a change in color or pH or the liberation of a chemical product

  37. Titration of acetic acid OH- H2O CH3COO− + H+ CH3COOH OH- added pKa = 4.8 pH 3 4 5 6 7

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