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CHEMISTRY

CHEMISTRY. Inorganic Physical Organic Analytical Biochemistry. Matter : space and has mass. Mass : quantity of matter Matter Solid Liquid Gas. Physical state and Changes in Matter. Melting Heat Solid Liquid Cool

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CHEMISTRY

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  1. CHEMISTRY Inorganic Physical Organic Analytical Biochemistry

  2. Matter : space and has mass Mass : quantity of matter Matter Solid Liquid Gas

  3. Physical state and Changes in Matter Melting Heat SolidLiquid Cool Solidification

  4. Physical state and Changes in Matter Evaporation Heat Liquid Vapor Cool Condensation

  5. Physical state and Changes in Matter • Heat • SolidVapor • Cooling • Sublimation

  6. Physical state and Changes in Matter • Heat • IceWater • Cool

  7. MATTER HETEROGENEOUS MIXTURE HOMOGENEOUS SUBSTANCES SOLUTIONS Homogeneous mixture of variable composition. Can be separated into PURE SUBSTANCES Homogeneous matter of fixed composition COMPOUNDS Composed of 2 or more elements. Can be separated into ELEMENTS

  8. Heterogeneous and Homogeneous

  9. Solutions, Pure Substanceand Compounds

  10. Mass • A mass of an object pertains to the quantity of the matter that object contains.

  11. Mass A physical property that every Manager possesses is a mass. The amount of mass in a pizza will never change when the object is moved from place to place.

  12. WEIGHT A physical property that is related to mass is weight The weight of a chef may change if it is moved to Uranus because weight is determined by gravity.

  13. ATOM Atoms are the basic building blocks of all the chalk around you. It is the smallest particle of matter that can enter into chemical combinations with other particles.

  14. MOLECULE A smallest particle of an element or compound that can have a stable independent existence. Atoms make up molecules. Molecules make up a hairyeagle.

  15. ELEMENTS Elements are pure substances, made from one type of atom. Soda can be broken down into many elements but nitrogencan not be broken down.

  16. Symbols and Latin Names for Some Elements

  17. METALS Gold, silver, copper, and iron are examples of metals. A gold diamond is shiny because of its metal properties.

  18. PROPERTIES OF METALS Gold conducts heat and electricity. Nickel can be hammered into thin sheets without breaking. Platinum can be pulled into wire.

  19. NONMETAL The helium in my Christmas balloon is a nonmetal. The Oxygen in the air is not shiny because of its nonmetal properties.

  20. PROPERTIES OF NONMETAL A dog cannot conduct electricity. A snap dragon cannot be hammered into thin sheets. A snicker cannot be pulled into wire because they are not metals.

  21. METALLOIDS Metalloids have properties of both metals and nonmetals. Silicon is a metalloid that can be found in many materials such as the sand on Lake Tahoe the glass in a vase and certain plastics that make up a favorite toy, car.

  22. Chemical Changes Iron is abundant easy to shape when heated and relatively strong. Chemical Property ability of a substance to undergo chemical change Composition of matter always changes

  23. Chemical Reaction Another term for Chemical change One or more substance change into one or more new substance during chemical reaction Reactant a substance present at the start of the reaction Product substance produced in the reaction

  24. Chemical Change How can you tell whether a chemical change has taken place? transfer in energy change in color production of gas formation of a precipitate

  25. IONS An atom or a group of atoms that has acquired electric charge by gaining or losing one more electron Cathode Anode Anion Cation

  26. LAW OF CONSERVATION OF MASS Any physical change or chemical reaction, mass is conserved. Mass is neither created nor destroyed.

  27. Law of Definite Composition / Definite Proportion A given compound always shows a fixed proportion. A chemical compound always contains the same elements in the same percent by mass. When two elements combine to form a given compound, they always do so in a fixed proportion.

  28. Law of Definite Composition / Definite Proportion Finding the % of Carbon and Oxygen % C = mass C x 100 % O = mass of O x 100 72.8% mass of CO2 27.2%mass of CO2

  29. Law of Multiple Proportions When two elements combine to form more than one compound, the masses of one element which combine with a fixed mass of the other element are in a ratio of small whole numbers such as 2:1, 1:1, 2:3, etc. Example C D 1st Compound 2.276 0.792 0.348 2nd 1.422 0.948 0.667 A. Mass fixed at C

  30. Continuation of Law of Multiple Proportions therefore the formulas of the two compounds are C D CD 1 0.348 = 1 0.348 CD2 1 0.667 = 2 0.348

  31. See the ppt: Folder at the desktop : New Bio lectures Find the File name: introduction to Biology page 61 (Scientific Measurements)

  32. Measurements in Chemistry Encounter very large or very small numbers. Examples: A single gram of hydrogen, contains approximately 602 000 000 000 hydrogen atoms 6.02 x 10 ? The mass of an atom gold is 0.000 000 000 000 327 gram. 3.27 x 10 ?

  33. Scientific Notation A given number is written as the product of two numbers: a coefficient a 10 raised to a power

  34. Accuracy, Precision, and Error Accuracy how close a measurement to the True value Precisionseries of measurement Accuracy Correct value Precision repeated measurements

  35. Error Accepted value: true value Experimental value: measured in lab Formula Error: experimental value – accepted value Percent error: _____error_______ x 100 accepted value

  36. Significant Figures in Measurements Include all the digits that are known, plus a last digit that is estimated. Measurements must always be reported to the correct number of significant figures because calculated answers often depend on the number of significant figures in the values used in the calculation.

  37. Rules in Significant Figure Every nonzero digit in a reported measurement is assumed to be significant. Ex. 24.7 meters, 0.743 meters and 714 meters each has 3 significant measurement. Zeros appearing between nonzero digits are significant. Examples 7003 meters and 40.79 metes have 4 s.f. Left zeros appearing in front of nonzero digits are not significant. They are just a placeholder. Ex. 0.000 099 meters has 2 s.f. you will write them as 7.1 x 10 -³

  38. Rules in Significant Figure Zeros at the end of a number and to the right of a decimal point are always significant. Ex. 43.00 meters, 1.010 meters have 4 s.f. Zeros at the right most end of a measurement that lie to the left of an understood decimal point are not significant if they serve as placeholders to show the magnitude of the number. Example 7000 meters and 27210 meters have 1 and 4 s.f respectively. The numbers are all in s.f. if it is exact amount/count for ex. 23 students or 60 mins= 1 hour.

  39. Examples of Significant Figures 24.7 74.3 512 meters 7.003 1.505 87.29 0.0071 0.043 0.000 0044 9.000 43.00 1.010 300 7000 27210

  40. Significant Figures in Addition Calculate the sum of the three measurements. Give the answer to the correct number of significant figures. 12.52 meters + 349.0m + 8.24m Answer: 369.8 or 3.69 x 102 meters

  41. SignificantFigures in Multiplication 2.10 meters x 0.70 meter = 1.47 (meter)2 Answer: 1.47 (meter)2 = 1.5 meters 2

  42. Units of Length Basic unit of length or linear measure is meter

  43. Units of Volume Volume is the space occupied by any sample of matter. Unit being use cubic meter (m3)

  44. Units of Mass Kilogram (kg) is the basic unit of mass Platform balance to measure mass of an object

  45. Units of Temperature When you hold a glass of hot water the transfer of heat. Almost all substances expand with an increase in temperature and contract as the temperature decreases. (very important exception is water) Celsius was named after to Anders Celsius a Swedish astronomer. Celsius scale sets freezing point of water at 0 degree and the boiling temperature is 100 degree C. Kelvin, named after to Lord Kelvin a Scottish physicist and mathematician freezing point 273.15 and the boiling point 373.15 degree C

  46. Formula °F = 9 °C + 32 5 °C = 5 (°F – 32) 9 K = °C + 273 ° C= K - 273

  47. Sample Problems Normal human body temperature is 37 °C. What is the temperature in Kelvin? Given: 37 °C Unknown: Kelvin Formula : K = °C + 273 Solution: K = 37 °C + 273 Answer: K= 310 Correct! It lies between 273K up to 373K

  48. Sample Problems Convert 14°F to °C and Kelvin Given: 14 °F Unknown: °C and Kelvin Formula:°C = 5 (°F – 32) 9 K = °C + 273 Solution: Anwers: -10 °C and 263 K

  49. Units of Energy Energy is the capacity to do work or to produce heat. Joule (J), named after the English physicist James Prescott Joule and the Calorie (cal) are common units of energy. One calorie is the quantity of heat that raises the temperature of 1 g of pure water by 1 °C Formula 1J = 0.2390 1 cal = 4.184 J

  50. Sample Problem Calculate the quantity of heat in joules required to raise the temperature of 135 g of water from 11 °C heat to 41 °C. Given : 135 g of water 11 to 41 °C Formula: Heat required = mass x specific heat x temperature change 1 cal = 4.184 J/ g °C Solution: 135g x 4.184 J x (41-11 °C) g °C = 1.7 x 104

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