1. Chemistry Introduction

2. Menu Definitions Classification of Matter Properties of Matter Measurement and SI Units Working with Numbers Quit

3. Definitions MatterMatter is anything that occupies space and has mass. ChemistryChemistry is the study of matter and the changes it undergoes substanceA substance is matter that has a definite or constant composition and distinct properties Examples are water, silver, sugar, table salt, etc.

4. Matter Mixtures Pure Substances Homogeneous Mixtures Heterogeneous Mixtures CompoundsElements Separation by Chemical Methods Separation by Physical Methods

5. A combination of two or more substances in which the substances retain their distinct identities.substances Mixtures may be separated by physical means (evaporation, magnet, distillation, etc.) and do not have constant composition. Mixtures

6. Types of Mixtures Homogeneous mixtures (solutions) 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 To Heterogeneous

7. Types of Mixtures Heterogeneous 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 To Homogeneous

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

9. 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

10. Compounds To Elements 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: Examples: H 2 O (water), CO 2 (Carbon dioxide), H 2 SO 4 (Sulfuric Acid)

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

12. Physical Property A physical property can be measured and observed without changing the composition of a substance. Examples: Boiling Point Density Conductivity

13. 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.

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

15. Intensive Property Measurable or observable properties which are independent of the amount of substance present are called intensive properties. Examples: Color Density Temperature

16. Measurement and SI Units SI units are an international standard of units developed in 1960 based on the decimal (base 10) system. Base QuantityName of UnitSymbol Lengthmeterm MassKilogramkg Timeseconds TemperaturekelvinK Amount of Substance molemol

17. Length Length measures the extent of an object. Length can be used to determine derived units such as area and volume. Area = m x m = m 2 Volume = m x m x m = m 3 1 Liter (L) = 1dm 3 (One cubic decimeter) 1 milliliter (mL) = 1cm 3 (One cubic centimeter) 1 L = 1000 mL Density d = m/V (mass per unit volume)

18. Mass Mass is a measure of the quantity of matter inside of a substance or object. It should not be confused with the term weight, which is a measure of the force that gravity exerts on an object. They are related by the following equation; F = mg where g is the acceleration due to gravity, m is the mass and F is the force in Newtons In chemistry, the smaller unit of mass grams (g) is preferable to kilograms (kg). 1kg = 1000g

19. Temperature Temperature measures the average kinetic energy of the particles contained within a system or object. Although Kelvin are the accepted SI unit, the Celsius scale is often used. Both are based on the decimal system. The Fahrenheit scale is seldom used for scientific measurement. Refer to the next frame for a comparison of temperature scales and conversion factors.

20. Temperature Comparisons and Conversions KelvinCelsiusFahrenheit 100212373Water Boils 98.6 2577 37 0 298 310 32273 K oCoC oFoF Body Temperature Water Freezes Room Temperature o F = 9/5 o C + 32 o C = ( o F - 32)5/9 K = o C + 273

21. Working With Numbers Scientific Notation Significant Figures Accuracy and Precision Factor-Label Method of Solving Problems

22. Scientific Notation Allows representation of large or small numbers accurately. Removes possible ambiguity about significant figures. Numbers are expressed follows; N x 10 n where N is a number between 1 and 10 and n is an integer exponent that is positive if the decimal point is moved to the left to make N between 1 and 10, and negative if it must be moved to the right.

23. Examples 1. The number 5,876.73 is expressed in scientific notation as; 5.87673 x 10 3 2. The number.000034785 is expressed as; 3.4785 x 10 -5

24. 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: 2.3x10 4 + 1.5x10 3 would be rewritten as 2.3x10 4 +.15x10 4 and the final answer would be 2.45 x 10 4.

25. Multiplication and Division 1. Multiply or divide the N parts of the numbers together 2. Add the exponents, n, if multiplying 3. Subtract exponents if dividing Example: 3.0x10 3 x 4.0x10 4 = 12x10 7 = 1.2x10 8

26. Significant Figures Significant figures refer to the meaningful digits in a measured or calculated quantity The last digit is understood to be uncertain when significant figures are counted

27. Guidelines Any digit that is not zero is significant Zeros between nonzero digits are significant Zeros to the left of the first nonzero digit are not significant If a number is greater than 1, then all the zeros written to the right of the decimal point count as 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

28. Calculations Involving Sig Figs Addition and Subtraction: The number of digits to the right of the decimal point in the final answer is determined by the lowest number of significant figures to the right of the decimal in any of the original numbers Multiplication and Division: The number of significant figures in the final answer is determined by the original number that has the smallest number of significant figures Exact numbers (from definitions or by counting) are considered to have an infinite number of significant figures For chain (multiple) calculations, carry the intermediate answers to one extra decimal place and round the final answer to the correct digits

29. Accuracy and Precision Accuracy tells how close a measurement is to the true value of the quantity that was measured Precision refers to how closely two or more measurements of the same quantity agree with one another Precise and accurate Precise but not accurate Neither precise nor accurate

30. Dimensional Analysis (Factor Label Method) Allows accurate conversion between units of similar types It utilizes the fact that equivalent quantities using different units may be set up as a ratio to convert from one type of unit to another Algebraically, labels are treated exactly the same way as the numbers they refer to The unit you are converting to should always be placed in the ratio such that the old units cancel out and the new unit is in the desired position whether numerator or denominator

31. Examples 1in = 2.54cm therefore the ratio 1in/2.54cm or 2.54cm/1in may by used to convert centimeters to inches or inches to centimeters, respectively 100in x (2.54cm/1in) = 254cm 1km = 0.6215mi 10km x (0.6215mi/1km) = 6.215mi