Nature of Chemistry-Classifying Matter

1. Nature of Chemistry-Classifying Matter

2. Matter matter: has mass and takes up space 3 states of matter: • solid: fixed shape & volume • liquid: fixed volume, but indefinite shape • gas: indefinite shape and volume The state depends on the forces of attraction (intermolecular force) between the particles in the substance.

3. Matter Matter is classified as a pure substance or a mixture. -pure substance: fixed composition; cannot be separated by physical means -mixture: made up of two or more substances that can be easily separated by physical means such as distillation, filtering, chromatography, etc.

4. Pure Substances elements: pure substance that cannot be separated into smaller substances by physical or chemical means -consist of only one type of atom -found on the periodic table, which organizes the elements based on chemical properties Periodic Table Dmitri Mendeleev (Russian) first organized the known elements based on properties, especially atomic mass -was able to accurately predict the properties of unknown elements

5. There was one problem with Mendeleev’s periodic table: some elements were swapped (Te & I as one ex.) Henry Moseley improved Mendeleev’s periodic table by arranging the elements based on atomic number (# protons) and not on atomic mass (# protons & neutrons) -technology had improved Elements are classified on the periodic table as: 1. metals: hard, shiny solid that is malleable, ductile and a good conductor of electricity -usually a solid -located on the left side of the periodic table

6. 2. nonmetals: usually gases or brittle solids -do not conduct and are not malleable or ductile -located on the right of the periodic table (except H) 3. metalloids: located between the metals and nonmetals that have properties of both metals and nonmetals compounds, another pure substance, are two or more different elements chemically combined in specific proportions. -can be broken down into simpler substances -have characteristics that are different from the individual components.

7. Element Groups Elements are grouped according to their characteristics (to be discussed later). Group 1: alkali metals Group 2: alkaline earth metals Groups 3-12: transition metals Group 17: halogens Group 18: noble gases The metalloids are along the stairstep in groups 13-17: -B, Si, Ge, As, Sb, Te, Po, At (also a halogen) -Al is not included

8. Mixtures Mixtures can be classified as heterogeneous or homogeneous. 1. heterogeneous: mixture where the different components are easily distinguished -may have to use a microscope to distinguish -salad, pizza -there are two types of heterogeneous mixtures: a. colloid: particles are larger than those in solution, but not heavy enough to settle out -can distinguish colloids by the Tyndall effect, scattering of light by a colloid (solutions will not scatter light) to allow a beam to be seen through it -ex: gelatin, milk

9. Mixtures b. suspension: mixture containing a liquid in which visible particles settle -muddy water, tea 2. homogeneous: contains 2 or more gaseous, liquid, or solid substances evenly blended throughout -also called a solution, where the particles are so small they cannot be seen with a microscope and don’t settle out -table 3.3, p 67

10. Separating Mixtures Four ways to separate mixtures: 1. filtration: uses a porous barrier to separate a solid from a liquid 2. distillation: separation technique based on differences of boiling points of liquids -the liquid with the lowest boiling point is vaporized first and condensed and collected 3. crystallization: separation technique that results in the formation of pure solid particles from a solution containing the dissolved substance -produces highly pure solids

11. Separating Mixtures Four ways to separate mixtures: 4. chromatography: technique that separates the components of a mixture (mobile phase) on the basis of the tendency of each to travel across the surface of another substance (called the stationary phase) at different rates

12. Physical and Chemical Properties Physical Properties of Matter physical property: characteristic that can be observed or measured without changing the substance’s composition -density, color, taste, odor, hardness, melting point, boiling point Can be classified as intensive and extensive: intensive property: independent of the amount of substance present -density, color, melting point…

13. Physical and Chemical Properties extensive property: dependent on amount of substance present -mass, length, volume Chemical Properties of Matter chemical property: ability of a substance to combine with or change into new substances -ability to rust or change colors

14. Physical and Chemical Changes Physical Changes physical change: altering a substance without changing its composition -depends on temperature and pressure -shape change, color change, change of state Chemical Changes chemical change: process that involves changing substances into something new -also called a chemical reaction -rusting, corrosion, flammability

15. Signs a chemical change occurs: 1. heat/light produced 2. bubbling (production of a gas) 3. new substance formed 4. color change (usually accompanies another sign) 5. change in odor Law of Conservation of Mass: mass is neither created or destroyed during a chemical reaction. mass of reactants = mass of products

16. Law of Conservation of Mass Is mass conserved in the following HgO  Hg + O2

17. Law of Conservation of Mass Now is mass conserved in the following 2HgO  2Hg + O2

18. Energy energy: capacity to do work temperature: the average kinetic energy of an objects’ atoms or molecules -higher temp = higher kinetic energy -SI Unit, the Kelvin (K) heat: flow of energy due to a temperature difference -measured in Joules (J) -flows from high temperature to low temperature

19. Energy Flow Endergonic: process that absorbs any form of energy -endothermic: absorbs heat Exergonic: process that releases any form of energy -exothermic: releases heat Other forms of energy: light, sound, electricity

20. Energy Changes As you raise the temperature of a material, the kinetic energy increases. calorie: amount of energy required to raise the temperature of one gram of water by one degree Celcius 1 calorie = 4.184 J Practice: 1. Convert 60.1 calories to Joules 2. Calculate 28.4 Joules to calories

21. Specific Heat specific heat: amount of energy required to raise the temperature of one gram of that substance by one degree Celcius (J/goC) -calculated using the formula q = mcDT where q = heat absorbed (+) or released (-) m = mass of sample in grams c = specific heat of the substance DT = change in temperature (Tf - Ti)

22. Specific Heat-Example The temperature of a sample of iron with a mass of 10.0g changed from 50.4oC to 25.0oC with a release of 114J of heat. What is the specific heat of iron? q = -114 J (heat is released-exothermic) m = 10.0 g c = ? DT = 25.0oC – 50.4oC = -25.4oC c = q = -114 J = 0.449J/goC mDT (10.0g)(-25.4oC)

23. Specific Heat-You Try If the temperature of 34.4g of ethanol increases from 25.0oC to 78.8oC, how much heat has been absorbed by the ethanol? The specific heat of ethanol is 2.44J/goC. q = ? will it be + or -? why? m = 34.4 g c = 2.44J/goC DT = 78.8oC – 25.0oC = 23.8oC q = (34.4g)(2.44J/goC)(23.8oC) = 1997.7 = 2.00x103 J

24. Specific Heat-Practice p 495 # 5-6, 11 5. A 4.50g nugget of pure gold absorbed 276J of heat. What was the final temperature of the gold if the initial temperature was 25.0oC. The specific heat of gold is 0.129J/goC. 6. A 155g sample of an unknown substance was heated from 25.0oC to 40.0oC. In the process, the substance absorbed 5696J of energy. What is the specific heat of the substance? 11. What is the specific heat of an unknown substance if a 2.50g sample releases 12.0 cal as its temperature changes from 25.0oC to 20.0oC?