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Chemistry 10 general chemistry

Chemistry 10 general chemistry. Deomila A. Basnig. Grading system. Definitions. Chemistry is the study of matter, its properties, composition, structure, and the changes it undergoes Matter is anything that occupies space and has mass.

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Chemistry 10 general chemistry

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  1. Chemistry 10general chemistry Deomila A. Basnig

  2. Grading system

  3. Definitions • Chemistry is the study of matter, • its properties, • composition, • structure, • and the changes it undergoes • Matteris anything that occupies space and has mass. • A substance is matter that has a definite or constant composition and distinct properties Examples are water, silver, sugar, table salt, etc.

  4. Pure Substances Pure substances consist of matter that has a definite composition and which always occur in fixed ratios. For instance, table salt always has the same components combined in the same ratio (or it would no longer be table salt!)

  5. Elements A substance that cannot be separated into simpler substances by chemical means. Currently 109 are known, of which eighty-three occur naturally on the Earth. It is these which are found on the periodic table. Examples: Gold (Au) Carbon (C) Zinc (Zn) Cobalt (Co) Note: Only the first letter of an element’s symbol is capitalized. Co is an element whereas CO is a compound To Compounds

  6. Compounds A compound is a substance composed of atoms of two or more elements chemically united in fixed proportions. Compounds have distinct properties from their components. Compound may only be separated by chemical means (reactions). Examples: H2O (water) CO2(Carbon dioxide) H2SO4(Sulfuric Acid)

  7. Mixtures • A combination of two or more substancesin which the substances retain their distinct identities. • - Has a variable composition • Example: juice, softdrinks, air,seawater, • Mixtures may be separated by physicalmeans (evaporation, magnet, distillation, etc.) and do not have constant composition.

  8. Types of Mixtures Homogeneous mixtures (solutions) - When the composition of a mixture, after sufficient stirring, is the same throughout the solution. • Examples: • Air • Soft drinks • Sugar or salt in water

  9. Types of Mixtures Heterogeneous mixtures - A type of mixture in which the composition is not uniform and the components remain visible and separate. • Examples: • Iron filings and sand • Milk (colloidal suspension) • Concrete

  10. Classification of Matter Matter Mixtures Pure Substances Separation by Physical Methods Homogeneous Mixtures Heterogeneous Mixtures Compounds Elements Separation by Chemical Methods

  11. Physical methods of separation for mixtures: • 9. extraction/ solvent extraction • 10. Centrifugation • 11. Chromatographic separation • e.g. paper chrom. • 12. Sublimation & deposition • 13. adsorption • Evaporation • Distillation • Filtration • Decantation • Scooping • Use of magnet • Crystallization • Seiving

  12. Three States of Matter • Solid - molecules are held close together in an orderly fashion with little freedom of motion • Liquid – molecules are close together but are not held so rigidly in position and can move past one another • Gas – molecules are separated by distance that are large compared with the size of the molecules

  13. Three states of Matter

  14. Properties of Matter Physical Property Chemical Property Extensive Property Intensive Property

  15. Physical Property A physical property can be measured and observed without changing the composition or identity of a substance. Examples: Boiling Point Density Conductivity He gas is lighter than air.

  16. Chemical Property A chemical property refers to the ability of a substance to react with other substances. In order to observe this property a chemical change must take place. Examples: Sugar ferments to form alcohol Hydrogen burns in oxygen to create water.

  17. Extensive Property Measurable properties which depend on the amount of substance present are called extensive properties. Examples: Mass Length Volume

  18. Intensive Property Measurable or observable properties which are independent of the amount of substance present are called intensive properties. Examples: Color Density Boiling Point Melting point

  19. measurement

  20. Temperature Comparisons and Conversions oF = 9/5 oC + 32 oC = (oF - 32)5/9 K = oC + 273 373 100 212 Water Boils Body Temperature 310 37 98.6 298 25 77 Room Temperature 273 0 32 Water Freezes Kelvin Celsius Fahrenheit K oC oF

  21. oF = 9/5 oC + 32 oC = (oF - 32)5/9 K = oC + 273 0C = K -273 • What is the temperature in Farenheit for a 100 K? Solution: 0C= (100 K – 273 K) 10C/K = -173 0C 0C = 100 – 273 = -1730C 0F= 9/5 (-173) +32 = -279.4 0F What is the equivalent temp scale in Kelvin for a 00F? Given: 0 F= _ K 0C= (0-32) 5/9 = -17.78 C K = -17.78 c + 273 = 255 K

  22. Metric system & english system

  23. Scientific Notation Allows representation of large or small numbers accurately. Removes possible ambiguity about significant figures. Numbers are expressed follows; N x 10n N is a number between 1 and 10 and n is an integer exponent

  24. Examples The number 5,876.73 is expressed in scientific notation as; (move to the left) 5.87673 x 103 2. The number 0.000034785 is expressed as; (move to the right) 3.4785 x 10-5

  25. Fish : true value is 5.0000 g

  26. Guidelines for using Significant Figures • Any digit that is not zero is significant • 845 cm has 3 SF • 1.234 kg has 4 SF • Zeros between nonzero digits are significant • 606 m has 3 SF • 40,501 s 5SF • Zeros to the left of the first nonzero digit are not significant • 0.0000349 g has 3 SF • 0.13 0.004008 4 SF

  27. Guidelines for using Significant Figures • If a number is greater than 1, then all the zeros written to the right of the decimal point count as significant figures • 2.0 mg has 2 SF • 0.3005 L has 4 SF • 40.00 4 • 0.0035600 5 • 0.050603 5

  28. Guidelines for using Significant Figures • For numbers that do not contain decimal points, the trailing zeros (zeros after the last nonzero digit) may or may not be significant. This is one reason why it is important to use scientific notation. • 400 cm • May have 1 or 2 or 3 SF • We can express as: • 1 SF 4 x 102 • 2 SF 4.0 x 102 • 3 SF 4.00 x 102

  29. Addition and Subtraction 1. Write each number so that n has the same exponent 2. Add or subtract the N parts of the numbers 3. The exponent n remains the same Example: 2.3x104m+ 1.5x103m would be rewritten as 2.3 x104 m .15x104 m and the final answer would be 2.45 x 104 2.5 x 104 m

  30. Example: Addition 2.3x104 m+ 1.5x103 m • 2.3x104 + 1.5x103 23 x 103 + 1.5 x 103 24.5 x 103 or 25 x 103 2.5 x 104 • 2.3x104 + 1.5x103 2.3 x104 m 0.15x104 m 2.45 x 104 2.5 x 104 m

  31. Dimensional Analysis in Solving Problems • Conversion of Units • Mass, volume, Density Problems • Rates • Percentages • Specific Gravity

  32. 1 g = 1000 mg OR 1000 mg or 1 g 1 g 1000 mg 1 mg = 10-3 g = 1 x 10-3 g = 0.001 g ? g = 325 mg x 1 g = 0.325 g 1000mg

  33. Conversion of units The mass of the Earth is 5.98 × 1024 kg. What is the mass expressed in megagrams, Mg? Note: in using scientific calc: X10 6- EXP6 Factors: 1 kg1000 g 1000 g 1 kg Factors: 1 kg = 1000 g 1 Mg = 1 x 106 g

  34. Conversion of units The mass of the Earth is 5.98 × 1024 kg. What is the mass expressed in megagrams, Mg? Factors: 1 L = 1000 ml 1 dL = 0.1 L or 1 x 10-1 L

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