THE MATERIAL WORLD

1. THE MATERIAL WORLD THE ORGANIZATION OF MATTER PROPERTIES OF MATTER FORMS OF ENERGY FLUIDS WAVES

2. THE ORGANIZATION OF MATTER KEY WORDS: • Matter • Pure Substance • Pure Elements • Pure Compounds • Atom and Molecule • Chemical Bond • Mixtures: Homogenous & Heterogeneous Mixtures • Solution: Solute & Solvent • Particle Model • Phases of Matter: Solid, Liquid Gas • Behavior of Matter

3. MATTER The air we breathe, the water we consume and the materials we use are all made up of matter. MATTER is anything that has volume or mass. • An ATOM is the smallest unit forming matter • A MOLECULE is a group of atoms attached by chemical bonds Example: A water molecule is formed by 2 hydrogen atoms bonding to one oxygen atom

4. PURE SUBSTANCES When all the particles in a substance are identical, we consider it a pure substance. There are two types of Pure Substances: • Pure Elements: Made up of only ONE type of atom. Gold (Au) is made of Au atoms only. • Pure Compounds: Formed by TWO or MORE different atoms combining. Water (H2O) is made up of two different atoms (2 hydrogen atoms and 1 oxygen atom) These 3 atoms have chemically bonded to create the molecule H2O (water).

5. MIXTURES Atoms and molecules can be mixed together without forming chemical bonds: this is called a mixture or solution. There are two types of Mixtures: • Homogenous: Made up of at least 2 substances that cannot be distinguished individually by the naked eye. • Heterogeneous:Made up of at least 2 substances that can be distinguished individually by the naked eye.

6. SOLUTIONS & SOLUBILITY Solution: A solution is a homogenous mixture in which it is impossible to distinguish its constituents (what it is made of). A solution is usually transparent. Solute:A substance that dissolves in another substance. Solvent: A substance that can dissolve a solute. Solubility: the amount of solute that can be dissolved in a given solvent. Water is typically the universal solvent. Colloid: A homogenous mixture in which at least 2 different substances can be seen under a magnifying instrument. Example: Blood is a colloid, under a microscope we can see both red and white blood cells. Milk is also a colloid.

7. HOMOGENOUS VS. HETEROGENEOUS MIXTURES HOMOGENOUS MIXTURE HETEROGENEOUS MIXTURE Composed of at least 2 pure substances Particles are not uniformly distributed Substances are not soluble in a solvent Contain 2 or more phases, which can be seen by the naked eye EX: Water and Sand • Composed of at least 2 pure substances • Particles are uniformly distributed • 1 or more substances (SOLUTES) are dissolved in a larger quantity of another substance (SOLVENT) • Contain 1 phase • Also called solutions • EX: Water and Salt

8. PARTICLE MODEL • A model is a visible representation of a situation that is abstract, hard to access or completely hidden. • Particles are naturally attracted to each other • The closer particles are to each other, the stronger the forces are. • Particles are in constant motion • Particles with a higher temperature move faster, on average, than particles with a lower temperature • The higher the temperature of a substance, the more energy its particles have.

9. PARTICLE MODEL (SOLIDS) • Particles of SOLIDS are very close together and therefore strongly attached to each other • The result is a very orderly arrangement of particles • Particles in a solid substance vibrate but do not actually move around

10. PARTICLE MODEL (LIQUIDS) • Particles in LIQUIDS are somewhat close together and therefore somewhat attached. • The arrangement of molecules is NOT very orderly • Particles in a liquid substance move around each other very slowly

11. PARTICLE MODEL (GASES) • Particles in GASES are widely spread out and weakly attached • The particles are arranged in a very disorderly fashion • Particles in a gaseous substance move rapidly in all directions

12. PROPERTIES OF MATTER KEY WORDS: • Physical Properties • Chemical Properties • Characteristic & Non Characteristic Properties • Density • Mass • Volume • Solubility • Concentration • Dilution • Ratio

13. PROPERTIES OF MATTER • A Property is information used to describe a substance. There are two types of Properties: 1.PHYSICAL PROPERTIES are those we can observe or measure without changing the nature of the substance. Examples: Boiling Point, Melting Point, Density and Solubility 2. CHEMICAL PROPERTIES describe how two properties react with each other, therefore you must modify the nature of the substance. Examples: Reaction to indicators (Litmus paper, fire alarm, pregnancy test)

14. CHARACTERISTIC VS. NON CHARACTERISTIC PROPERTIES A Non-Characteristic Property cannot help us tell one pure substance form another. Example: Saying the window is transparent will not help us understand what the glass is made of. It can be several things; glass, plastic, etc… A Characteristic Property helps us to identify precisely a pure substance. Example: Melting point and Density are examples of characteristic properties. Plastic and glass are examples of transparent solids. Their melting points and density, however, are very different.

15. DENSITY DENSITY: Measures the amount of matter in a given space. Every substance has its own specific density. (Ratio of Mass/Volume) • Example: The density of vegetable oil (0.92g/ml)is less than water (1.0 g/ml) and they cannot be mixed together; the oil simply floats on the surface of the water. • Example: The density of helium (gas) is lower than that of air, which explains why balloons filled with helium gas fly away.

16. CALCULATING DENSITYM Formula: p= M/V p V Mass is represented as g(grams) Volume is represented as ml (liquid) cm3 (solid) Example 1: A student measures the mass of an 8 cm3 block of brown sugar to be 12.9g. What is the density of the brown sugar? p=M/V 12.9g (mass) / 8cm3 (volume) = 1.61g/cm3 Example 2: A chef fills a 50ml container with 43.5g of cooking oil. What is the density of the oil? p=M/V 43.5g (mass) / 50ml (volume) = 0.87g/ml

17. CALCULATING DENSITY Example 3: Sam measures the mass of an empty cylinder to be 32.5g. Then, he adds 40 ml of the unknown liquid. Finally, he measures the mass of the cylinder with the liquid to be 82.9g. Calculate the density of the liquid. p=M/V 82.9g – 32.5g = 50.4g (mass of liquid) p = 50.4g (mass) / 40ml (volume) = 1.26g/ml Example 4: A piece of stone 12cm x 4cm x 4cm has a mass of 872g. What is the density of the stone? p=M/V Volume= length x width x height (12 x 4 x 4 = 192cm3) p = 872g (mass) / 192cm3 (volume) = 4.54g/cm3

18. SOLUBILITY There is a limit to the amount of solute you can add to a solvent. The maximum amount of solute in a solvent corresponds to its solubility. Unsaturated Solution: Contains less than the max amount of solute. Saturated Solution: Contains exactly the max amount of solute. If it contains more than the max amount of solute, a surplus solute will appear as a Precipitate. Water Soluble: A substance that can be dissolved in water. Lipo - Soluble: A substance that can be dissolved in lipids (fats).

19. CONCENTRATION The concentration of a solution corresponds to the quantity of dissolved solute in a given quantity of solution (Gatorade powder in water). Concentration increases when solute is added to the solution. A solutions concentration can be expressed in a variety of ways: • g/L (number of grams of solute per 1 liter of solution) • % M/V ( number of grams (g) of solute per 100ml of solution) • % V/V (number of milliliters (ml) of solute per 100ml of solution)

20. CALCULATING CONCENTRATION C = Mass/Volume • Solving for Concentration: Mass / Volume Mass • Solving for Mass: Concentration x Volume • Solving for Volume: Mass/Concentration Conc. Volume Mass = Solute (g) Volume = Total amount of Solution (L) Concentration = Amount of dissolved solute in a given amount of solution. CONVERSION TABLE

21. CALCULATING CONCENTRATION Example 1: You are mixing 2500mg of powder in a 4000ml solution. What is the concentration in g/L? Steps for Example 1: • Convert g/L 3. Identify variables 2. Write down formula 4. Divide Step by Step Answer: • 2500mg/1000 = 2.5g, 4000ml/1000 = 4 L • C= M/V M • M= 2.5g V= 4L C= ? C V • 2.5g / 4L = 0.625 g/L C= 0.625 g/L

22. CALCULATING CONCENTRATION Example 2: You have a mixture with a concentration of 12g/L in a 50 ml solution. What is the mass in grams? Steps for Example 2: • Convert g/L 3. Identify variables • Write down formula 4. Multiply Step by Step Answer: 1. 50 ml/ 1000 = 0.05L, 12g/L (Conversion not necessary) 2. C = M/V M 3. C= 12g/L V= 0.05L M= ? C V 4. 12g/L x 0.05L = 0.6g M= 0.6g

23. CALCULATING CONCENTRATION Example 3:Converting %M/V ,%V/V and ml to g/L A scientist is comparing the concentration of three different drinks. The concentrations of the drinks are listed below. Which drink has the lowest concentration? CONCENTRATION OF DRINKS Drink A: 0.65 g/ml Drink B: 12% M/V Drink C: 6.5g/100ml Step by Step Answer: Drink A: 1L = 1000ml, 0.65g x 1000ml = 650g/L Drink B: 12% M/V = 12g/ 100ml, 100ml x 10 = 1000ml or 1L, 12g x 10 = 120g/L ****(Same procedure for %V/V, replace g for ml) Drink C: 100ml x 10 = 100ml or 1L, 6.5g x10 = 65g/L

24. DILUTION Dilution is a technique that involves decreasing the concentration of a solution by adding solvent. ********INSERT PIC FROM P.16*************

25. DILUTION Formula:C1 V1=C2 V2 C= Concentration in g/L V = Volume in L or ml C1: Original concentration of solution V1: Original volume of solution C2: New concentration of diluted solution V2: New volume of diluted solution ** If concentration is represented as %M/V or %V/V, be sure to convert all variables to g/L before solving.**

26. CALCULATING DILUTION Example 1: You have a solution with a concentration of 60g/L. You add water to dilute the solution. The new concentration is 20 g/L in a 100ml solution. How much water in ml was added to the new solution? Steps for Examples 1, 2 & 3: • Write down formula 6. Isolate the unknown variable • Identify variables 7. Divide • Convert ml to L 8*. Convert L back to ml • Plug variables into formula 9*. Subtract V2-V1(to figure out • Multiply Brackets amount of added water) * Only use steps 8 & 9 when solving for V1* **See Example 3 for extra steps (converting %M/V to g/L)**

27. CALCULATING DILUTION Example 1: Step by Step Answer: • C1 V1 = C2 V2 • C1= 60g/L V1= ? C2= 20g/L V2= 100ml • 100ml / 1000 = 0.1L • (60g/L)(V1) = (20g/L)(0.1L) • 60 x V1 = 60V1 20 x 0.1 = 2 • 60V1 = 2 7. 2 ÷ 60 = 0.033L 8. 0.033L x 1000 = 33ml 9. V2-V1 100ml – 33ml = 67ml 67ml of water was added to the new solution.

28. CALCULATING DILUTION Example 2: OJ Delight juice has a concentration of 25g/L in a 300 ml solution. It was too concentrated. The company added 500ml of water to the juice. What is the new concentration? Step by Step Answer: • C1 V1 = C2 V2 • C1= 25g/L V1= 300ml C2= ? V2= 800ml (300ml+500ml) • 300ml / 1000 = 0.3L 800ml / 1000 = 0.8L • (25g/L)(0.3L) = (C2)(0.8L) • 25 x 0.3 = 7.5 C2 x 0.8 = C2 0.8 • 7.5 = C2 0.8 7. 7.5 ÷ 0.8 = 9.38g/L The new concentration of OJ Delight is 9.38g/L

29. CALCULATING DILUTION Example 3: A Scientist would like to dilute a 6.0% M/V solution. He takes the original 400ml of solution and adds water to prepare 2L of the new diluted solution. What is the concentration of the new diluted solution in % M/V? Extra steps for converting %M/V • 6.0% M/V = 6g / 100ml • Convert ml to L 100ml x 10 = 1L (1000ml) • Balance the ratio by also multiplying the 6g (6g x 10= 60g) • Concentration expressed in g/L = 60g/L • 60g/L is the same concentration as 6% M/V and 6g/100ml

30. CALCULATING DILUTION Example 3: Step by Step Answer: • C1 V1 = C2 V2 • C1= 6g/L V1= 400ml C2= ? V2= 2L • 400ml / 1000 = 0.4L • (6g/L)(0.4L) = (C2)(2L) • 6 x 0.4 = 2.4 C2 x 2 = C2 2 • 2.4 = C2 2 • 2.4 ÷ 2 = 1.2g/L • Convert back to %M/V( 1.2g/L ÷ 10 = 0.12g/100ml or 0.12%M/V) • The Concentration of the new diluted solution is 0.12 % M/V

31. FORMS OF ENERGY Key Words: • Energy • Thermal Energy • Radiant Energy • Chemical Energy • Mechanical Energy (Kinetic & Potential) • Energy Transfer • Energy Transformation • Physical Changes • Phase Changes (Sublimation, Deposition, Fusion, Freezing, Condensation, Vaporization) • Dissolution • Deformation • Chemical Changes (Synthesis, Decomposition, Oxidation, Precipitation) • Molecular Formula

32. ENERGY Energy is the capacity to do work or produce change. • When a person lifts his or her arm, we can see the arm move. This movement involves work. It’s because of energy that the work is produced. • When the snow melts, it changes from a solid to a liquid phase. This change is made possible because of the energy from the sun. • Energy is expressed in Joules (J). One J represents the energy needed to lift an object weighing 1 N (Newton) to a height of one meter. • 1 N = 100g.

33. FORMS OF ENERGY Thermal Energy is energy in the form of Heat that comes from the random motion of particles that make up a substance. Examples of the work or change that can be accomplished by Thermal Energy: • The melting of snow by the heat of the sun’s rays • The lifting of a hot air balloon by the burner’s heating of air • The production of water vapor in order to advance the train by the boiler of a steam engine

34. FORMS OF ENERGY Radiant Energy is energy in the form of Radiation that is contained in and transported by electromagnetic waves. Examples of the work or change that can be accomplished by Radiant Energy: • Producing images of the inside of the body using X – rays • Getting suntanned and sunburned by ultraviolet rays • Treating certain cancers with gamma rays (chemo therapy) • Producing light from a light bulb to brighten a room

35. FORMS OF ENERGY Chemical Energy is energy in reserve. On its own it doesn’t work or create change. To get the energy, the chemical bonds need to be reorganized. This means breaking the bonds to release energy they contain, and remaking bonds differently. Thus changing the molecular formula of the substance(s). Examples of the work or change that can be accomplished by Chemical Energy: • The operation of a car motor by the combustion of gasoline • The burning of a candle to produce light • Photosynthesis; energy from the sun being transformed and stored in glucose molecules • Cellular Respiration; helps living organisms use chemical energy stored in glucose molecules to grow, move and think

36. FORMS OF ENERGY Mechanical Energy is the energy that results from the speed of an object, its mass and its relationship to its surroundings. Potential vs. Kinetic Energy:A person is standing and holding a rock over a cliff. At this moment the rock has potential energy. As soon as the rock is released from the grasp of the person the rock is using kinetic energy until it hits the ground. If the rock is dropped from a higher position, mechanical energy will increase (relationship to surroundings) Speed vs. Mass: Let’s use the example of a car to understand how speed and mass can affect mechanical energy. If we compare the damage caused by a car accident at 100km vs. 30km, we can see that the mechanical energy of a car increases with speed. At the same speed, a large truck has more mechanical energy than a car. Greater mass = greater energy.

37. FORMS OF ENERGY Examples of the work or change that can be accomplished by Mechanical Energy: • The creation of a crater by an asteroid • The spinning of windmill blades by the wind • The powering of a hydroelectric power station by a waterfall

38. ENERGY TRANSFER vs. TRANSFORMATION Energy Transfer is the movement of energy from one place to another. Energy Transformation is the changing of energy from one form to another. *******INSERT PIC FROM PAGE 42**************

39. PHYSICAL vs. CHEMICAL CHANGES Physical Changes are changes that do not affect the nature or the characteristic properties of matter. Examples of Physical Changes: • Ice melting into liquid • Adding salt to soup • Folding a piece of paper Chemical Changes are changes that do affect the nature or characteristic properties of matter. Examples of Chemical Changes: • When baking soda and vinegar are mixed, we see gas being released • When garbage decomposes in landfills, it breaks down into simpler atoms or molecules • The oxidation of iron (the formation of rust) releases energy • Adding acidic substances to milk or cream produces a precipitate (curd), this is how cheese is made

40. PHYSICAL CHANGES There are 3 different types of Physical Changes we will discuss: 1.Phase Change is the transformation from one state or phase to another. There are 6 different phase changes: (Sublimation & Deposition, Fusion & Freezing, Condensation & Vaporization) 2. Dissolution is the creation of a solution by a solute dissolving in a solvent 3. Deformation means changing the shape of a material

41. PHASE CHANGE ******** INSERT PIC FROM PAGE 43******

42. DISSOLUTION Solute and solvent particles gap in-between each other. They do not have a chemical bond. ******Insert Pic From Page 47************

43. DEFORMATION Some deformations are reversible (stretching a coil spring) and some are permanent (tearing a sheet of paper). *******Insert Pic From Page 50*********

44. CHEMICAL CHANGES There are 4 different types of Chemical Changes we will discuss: 1. Synthesis is the formation of a complex molecule from atoms or simpler molecules (H2O). + = 2. Decomposition is the transformation of complex molecules into simpler molecules or into atoms (the opposite of Synthesis). = + 3. Oxidation is a chemical reaction involving oxygen or a substance that has similar properties to oxygen. 4. Precipitation is the formation of a solid (precipitate) that is less soluble, or not soluble, following the mixing of two solutions.

45. EXAMPLES OF CHEMICAL CHANGES ****** INSERT PICS OF SYNTHESIS p.54****** ******DECOMPOSITION (GARBAGE)****** ******* OXIDATION (RUST & BANANAS)***** ******PRECIPITATION(CHEESE CURDLING)*****

46. FLUIDS Key Words: • Liquid State • Gaseous State • Particle Model • Forces of Attraction • Compressible Fluid • Incompressible Fluid • Pressure • Force • Surface Area • Volume • Pressure Exerted • Particle Collisions • Depth & Density

47. WHAT IS A FLUID? A Fluid is a substance that has the capacity to flow and assume the form of the container into which it has been poured. • A liquid can be a fluid • A gas can be a fluid • A solid cannot be a fluid (Forms a pile instead of taking the natural shape of a container) ******INSERT PIC FROM PAGE 67 (BOTTOM)********

48. Fluids Through the Particle Model Particles in a liquid, such as water, are: • Somewhat close together • Held by relatively weak forces of attraction • Particles can slide over each other to take on the form of any container Particles of a gas, such as air, are: • Much farther apart from each other • Not bound together by any forces of attraction • They move in every direction and occupy all available space

49. Compressible vs. Incompressible Fluids Compressible Fluids • A compressible fluid is one that can be squeezed or compressed. • Its volume can change in response to a change in pressure. • The volume of a compressible fluid can change because the space between the particles in the fluid can vary. Incompressible Fluids • An incompressible fluid is a fluid whose volume cannot be varied. • Its volume cannot change noticeably because the particles in a liquid are already very close together. • The cannot be condensed any more.

50. Compressible vs. Incompressible Fluids *****INSERT PICS FROM P. 68 & 69*********