Chemistry Unit Two

1. Chemistry Unit Two Matter and Energy

2. Matter • Matter • anything that has a mass and takes up space. • Law of Conservation of Mass/Matter • Matter cannot be created or destroyed in an ordinary chemical reaction just rearranged to form different substances • Matter can be described using properties..

3. PROPERTIES CHEMICAL PHYSICAL EXTENSIVE INTENSIVE Types of Properties

4. Characteristics of Matter • Physical Properties • Characteristics of a substance that can be observed without the production of a new substance. • Examples: • Color,smell, taste, hardness, density, texture, melting/boiling/freezing points, magnetic attraction, solubility, electrical conductivity, temperature, state or phase

5. Two Types of Physical Properties • Extensive • Depends on the particular sample • examples: volume, mass, weight, shape, etc… • Intensive • Depends on the type of matter  NOT size of sample • examples: color, melting point, specific heat, density, appearance, etc…

6. Characteristics of Matter • Chemical Properties • describes how a substance reacts or fails to react with other substances to produce new substances. • Examples: • Oxidation, Corrosion, Hydrolysis, Combustion, Flammability, Reaction to Acid or Base.

7. Two Types of Changes • Physical Change • an alteration of a substance that only changes the physical properties of the substance. • Does not change the chemical composition of the matter!!

8. Characteristics of Matter • Chemical Change • an alteration of the chemical composition of a substance that results in the formation of a new substance • ALWAYS forms a new substance that has different physical and chemical properties than the original substance. • Also known as a chemical reaction.

9. Chemical and Physical change: • Practice 1 • Practice 2

10. Kinetic Theory • All matter is made of tiny particles in constant motion. • Potential Energy (PE) • energy due to the position or condition • at the atomic level: • the distance between the particles • closer= lower PE farther = higher PE • Kinetic Energy (KE) • energy due to motion • Faster=higher KE slower= lower KE

11. Phases of Matter Arranged in orderly pattern Yes Yes Touching, but not tightly packed No Yes Far apart and rarely touching No No

12. Phases of Matter Vibrational only Low Ice Very Low Vibrational & translational Low Moderate Water High High Vapor Move freely

13. Plasma • extraordinary state of matter • consists of high energy particles • electrons are stripped from their nuclei • examples: • fluorescent light • Stars • Lightning *Most Abundant State of Matter in the Universe!*

14. Phase Changes Changes of State • Adding or removing energy (heat) to a substance causes phase changes • The potential energy of the particles is increased or decreased • During a phase change, temperature does NOT change

15. Phase Changes • Melting • S Δ L (adding energy) • Freezing • L Δ S (removing energy) • Melting point & freezing point of a substance occur at the same temperature.

16. Phase Changes • Boiling • L Δ G (adding energy) • Evaporation • L Δ G (adding energy) • Condensation • G Δ L (removing energy) • Difference between boiling & evaporation: • Boilinga specific temp. below the surface • Evaporation any temp. at the surface

17. Phase Changes • Deposition • G Δ S (removing energy) • Examples: Snow, frost • Sublimation • S Δ G (adding energy) • Examples: solid CO2 (dry ice), solid air fresheners

18. Phase Change Graphs (T vs t) Liquid Melting Solid AB -heat Δ KE -move faster -temp.  -solid BC -heat Δ PE -get farther apart -temp. stay same -melting CD -heat Δ KE -move faster -temp.  -liquid

19. Phase Change Graph (T vs t) Gas Boiling DE -heat Δ PE -get farther apart -temp. stay same -boiling EF -heat Δ KE -move faster -temp.  -gas

20. Phase Change Graph (T vs t) A C B E D F CD -KE  -slows down -temp.  -Liquid AB -KE  -slows down -temp.  -Gas BC -PE  -closer together -temp. stays same -Condensation

21. Phase Change Graph (T vs t) A B C D E F DE -PE  -closer together -temp. stays same -Freezing EF -KE  -slows down -temp.  -Solid

22. Phase Change Graph (T vs t) Boiling Point Boiling Freezing Point & Freezing Melting Point Melting What is the boiling point? What is the melting point? What is the freezing point?

23. Phase Change Graph (T vs t) If melting & freezing points occur at the same temperature, how do you know which change is occurring? -depends on whether adding or removing energy

24. Phase Change Graph (T vs t) What is this substance? -Water How do you know? -Boiling & melting & freezing points of water (Intensive properties)

25. Heat Calculations • Heat (q) • Energy transferred from an object at a higher temperature to an object at a lower temperature. (heat lost = -heat gained) • q = mcT • q=mHfus • q=mHvap

26. Heat Calculations • A 10.0g sample of iron at 50.4oC is cooled to 25.0oC in 50.0g of water. Calculate the amount of heat lost by the iron. ciron= 0.449 J/goC *How much heat is gained by the water? • A 2.1g ice cube at –8.0oC melts completely and warms to 12.5oC. How much heat was required? Hfus ice = 334 J/g cice = 2.03 J/goC cwater = 4.18J/goC

27. Classification of Matter Matter Pure Substances Mixtures Elements Compounds Homogeneous Heterogeneous

28. Matter • Pure Substances • made of only one type of matter • Mixtures • a physical combination of two or more substances • no definite ratio of particles • Element • made of only one type of atom • cannot be broken down into simpler substances under normal laboratory conditions

29. Matter (cont’d) • Compound • Atoms of two or more elements, chemically combined in a definite ratio. • Homogeneous Mixtures • Atoms of two or more elements, physically combined in no definite ratio. • The same throughout. • Must be a SOLUTION • Heterogeneous Mixture • Atoms of two or more elements, • physically combined in no definite ratio. • Different throughout

30. Classifications of Mixtures • Solutions • Particles are very tiny, will not separate by filtering, will not settle out when allowed to stand, particles will not scatter light, (-) Tyndall effect. • Ex. Salt Water, Kool-Aid, Brass • Colloids • Particles are tiny, will not separate by filtering, will not settle out when allowed to stand, particles will scatter light, (+) Tyndall effect. • Ex. Milk, whipped cream, aerosols • Suspension • Particles visible with unaided eye, will separate when filtered, will settle out if allowed to stand, particles will scatter light, (+) Tyndall effect. • Ex. Muddy water, snow globe

31. Solutions • SOLUTION • a solute dissolved in a solvent. • The solvent is the part in greater quantity. • For example: In a solution of salt water, salt is the solute and water is the solvent. • ELECTROLYTE • a solution that conducts electricity in water or molten form. • Salt water will conduct electricity.(Electrolyte) • Sugar water will not.

32. Types of Solutions • Gas-Gas • Carbon dioxide, Nitrogen,Oxygen (air) • Liquid-Gas • Water Vapor in Air (moist air) • Gas-Liquid • Carbon dioxide in Water (soda water) • Liquid-Liquid • Acetic acid in Water (vinegar) • Solid-Liquid • Sodium chloride in Water (brine or salt water) • Solid-Solid • Copper in Silver (Sterling Silver)

33. Characteristics of Solutions • Homogeneous Mixture • Solute / solvent • Soluble- • Likes dissolve likes • Insoluble • Immiscible • Miscible

34. Solvation • When a solid solute is placed in a solvent, the solvent particles completely surround the surface of the solid solute. • If attractive forces between the solute particles and the solvent are greater than the attractive forces holding the the solute particles together, the solvent particles pull the solute particles apart and surround them.

35. - + + + - + + + - - - + - + + - + - + + - + - - + - + + + - + - - + + + + Process of Solvation H2O H = O = - - + - + NaCl Na = Cl = + -

36. Water- Universal Solvent • Polar molecule • Dipoles allow solvation of ions and polar molecules

37. Solvation Factors • Agitation • Increasing collisions and breaking up solute attraction • Increasing surface area of solute • Small pieces, more area for solvent to contact • Increasing temperature of solvent • Greater kinetic energy creates more collisions

38. Heat of Solution • Endothermic- • Solute and solvent particles break attractive forces holding them together • Exothermic- • When solute and solvent particles mix, particles now attract each other

39. Solubility • Refers to the maximum amount of solute that will dissolve in a given amount of solvent at a specified temperature and pressure. • g solute / 100 g solvent • Saturated vs Unsaturated vs Supersaturated

40. Concentration • How much solute is dissolved in a specific amount of solvent • Percent by mass • Percent by volume • Molarity: moles/Liter • Molality: moles solute/kilograms solvent

41. Molarity • # moles of solute / Liters of solution • Calculate the molarity of 1.60 L of a solution containing 1.55 g of dissolved KBr? • How many grams of CaCl2 would be dissolved in 1.0 L of a 0.10M solution of CaCl2?

42. Diluting Solutions • M1V1 = M2V2 • What volume of a 3.00 M KI stock solution would you use to make ).300 L of a 1.25 MKI solution?

43. Colligative properties of Solns • Physical properties of solutions that are affected by the number of particles but not the identity of dissolved solute particles • Vapor Pressure Lowering • Boiling point elevation • Freezing point depression