Physical and Chemical Changes Chapter 2

1. Physical and Chemical ChangesChapter 2

2. Changes in Matter • Physical Changes are changes to matter that do not result in a change of the fundamental components that make that substance • State Changes – boiling, melting, condensing • Chemical Changes involve a change in the fundamental components of the substance • Produce a new substance • Chemical reaction • Reactants  Products

3. Density • Density is a property of matter representing the mass per unit volume • For equal volumes, denser object has larger mass • For equal masses, denser object has small volume • Solids = g/cm3 • 1 cm3 = 1 mL • Liquids = g/mL • Density : solids > liquids >>> gases • Water: density = 1g/ml Iron: density = 7.86 g/cm3

4. Water • Freezes at 0°C • At 1 atm, solid at 0°C or below • Normal freezing point = normal melting point • Boils at 100°C • At 1 atm, liquid up to 100°C, then turns to steam • Normal boiling point • Boiling point increases as atmospheric pressure increases • Temperature stays constant during a state change • Relatively large amounts of energy needed to melt solid or boil liquid

5. Heating Curve • As heat added to solid, it first raises the temperature of the solid to the melting point • Then added heat goes into melting the solid • Temperature stays at the melting point • Heat of Fusion • As more heat added it raises the temperature of the liquid to the boiling point • Then added heat goes into boiling the liquid • Temperature stays at the boiling point • Heat of Vaporization • As more heat added it raises the temperature of the gas

6. The heating/cooling curve for water heated or cooled at a constant rate

7. Both liquid water and gaseous water contain H2O molecules

8. Representations of the gas, liquid, and solid state

9. Intramolecular (bonding) forces exist between the atoms in a molecule and hold the molecule together

10. Sublimation • Solids can change directly into gases without going through the liquid state • Dry ice is solid CO2; it never melts under normal conditions. Instead it sublimes to gaseous CO2. • The formation of frost is reverse of sublimation = vapor deposition

11. The Solution Process - Dissolving is not Melting • When ionic compounds dissolve in water they dissociate into ions • ions become surrounded by water molecules - hydrated • When solute particles are surrounded by solvent molecules we say they are solvated

12. When solid sodium chloride dissolves, the ions are dispersed randomly throughout the solution

13. Polar water molecules interact with the positive and negative ions of a salt

14. H O H H C O H H H H O H The Solution ProcessCovalent Molecules • Covalent molecules that are small and have “polar” groups tend to be soluble in water • The ability to H-bond with water enhances solubility

15. Attractive Forces and Properties • Larger attractive forces between molecules in pure substance means • higher boiling point • higher melting point (though also depends on crystal packing) • Like dissolves Like • Polar molecules dissolve in polar solvents • Water, alcohol • Molecules with O or N higher solubility in H2O due to H-bonding with H2O • Non-polar molecules dissolve in non-polar solvents • Oils and gasoline

16. Evaporation • Requires overcoming intermolecular attractions • Condensation is the reverse process • In a closed container, eventually the rate of evaporation and condensation are equal • Equilibrium • In open system, evaporation continues until all liquid evaporated

17. Behavior of a liquid in a closed container. The system is at equilibrium.

18. Equilibrium Liquid just poured into open container, little vapor Evaporation faster than Condensation Evaporation as fast as Condensation Vapor Pressure • Pressure exerted by a vapor in equilibrium with a liquid • Or solid • Increases with temperature • Larger intermolecular forces = Lower Vapor Pressure • Liquid boils when its Vapor Pressure = Atmospheric Pressure • Normal boiling point • Raising external pressure raises boiling point, & visa versa

20. Chemical Reactions • Reactions involve chemical changes in matter resulting in new substances • Reactions involve rearrangement and exchange of atoms to produce new molecules • Elements are not transmuted during a reaction Reactants  Products

21. Evidence of Chemical Reactions • a chemical change occurs when new substances are made • visual clues (permanent) • color change, precipitate formation, gas bubbles, flames, heat release, cooling, light • other clues • new odor, permanent new state

22. Bubbles of hydrogen and oxygen gas form when an electric current is used to decompose water

23. Hot and cold pack reactions

24. Chemical reactions

25. Chemical reactions

26. Chemical reactions

27. Chemical reactions

28. Chemical Equations • Shorthand way of describing a reaction • Provides information about the reaction • Formulas of reactants and products • States of reactants and products • Relative numbers of reactant and product molecules that are required • Can be used to determine weights of reactants used and of products that can be made

29. Conservation of Mass • Matter cannot be created or destroyed • In a chemical reaction, all the atoms present at the beginning are still present at the end • Therefore the total mass cannot change • Therefore the total mass of the reactants will be the same as the total mass of the products

30. Writing Equations • Use proper formulas for each reactant and product • proper equation should be balanced • obey Law of Conservation of Mass • all elements on reactants side also on product side • equal numbers of atoms of each element on reactant side as on product side • balanced equation shows the relationship between the relative numbers of molecules of reactants and products • can be used to determine mass relationships

31. Balancing by Inspection 1. Count atoms of each element • polyatomic ions may be counted as one “element” if it does not change in the reaction Al + FeSO4Al2(SO4)3 + Fe 1 SO4 3 • if an element appears in more than one compound on the same side, count each separately and add CO + O2 CO2 1 + 2 O 2

32. Balancing by Inspection 2. Pick an element to balance 3. Find Least Common Multiple and factors needed to make both sides equal 4. Use factors as coefficients in equation • if already a coefficient then multiply by new factor 5. Recount and Repeat until balanced

33. Examples • when magnesium metal burns in air it produces a white, powdery compound magnesium oxide • burning in air means reacting with O2 1. write the equation in words • identify the state of each chemical magnesium + oxygen magnesium oxide 2. write the equation in formulas • identify diatomic elements • identify polyatomic ions • determine formulas Mg + O2MgO

34. Examples • when magnesium metal burns in air it produces a white, powdery compound magnesium oxide • burning in air means reacting with O2 3. count the number of atoms of on each side • count polyatomic groups as one “element” if on both sides • split count of element if in more than one compound on one side Mg + O2MgO 1  Mg 1 2  O  1

35. Examples • when magnesium metal burns in air it produces a white, powdery compound magnesium oxide • burning in air means reacting with O2 4. pick an element to balance • avoid element in multiple compounds 5. find least common multiple of both sides & multiply each side by factor so it equals LCM Mg + O2MgO 1  Mg 1 1 x 2  O  1x 2

36. Examples • when magnesium metal burns in air it produces a white, powdery compound magnesium oxide • burning in air means reacting with O2 6. use factors as coefficients in front of compound containing the element • if coefficient already there, multiply them together Mg + O22MgO 1  Mg 1 1 x 2  O  1 x 2

37. Examples • when magnesium metal burns in air it produces a white, powdery compound magnesium oxide • burning in air means reacting with O2 7. Recount Mg + O22MgO 1  Mg 2 2  O  2 8. Repeat 2 Mg + O22MgO 2 x 1  Mg 2 2  O  2

38. O H H O + + C O O C H H H H O 1 C + 4 H + 2 O 1 C + 2 O + 2 H + O 1 C + 2 H + 3 O Combustion of Methane • methane gas burns to produce carbon dioxide gas and liquid water • whenever something burns it combines with O2(g) CH4 + O2 CO2 + H2O

39. O O O O H H H H + + + C C + H H O O O O H H 1 C + 4 H + 4 O 1 C + 4 H + 4 O Combustion of MethaneBalanced • to show the reaction obeys the Law of Conservation of Mass it must be balanced CH4 + 2 O2 CO2 + 2 H2O

40. Carbon Monoxide • Incomplete combustion of carbon compounds occurs when there is not enough oxygen to fully make CO2. 2 C4H10 + 13 O2 ----> 8 CO2 + 10 H2O 2 C4H10 + 9 O2 ----> 8 CO + 10 H2O • CO is called a by-product (not the intended product)

41. Carbon Monoxide • CO concentrations can be reported in the parts per million scale (ppm) • ppm tells how many particles out of a million are the compound in question • Air normally has 2 ppm CO • CO is a common pollutant from burning fossil fuels like gasoline or coal • Cars must have a catalytic converter, which helps ensure all C is converted to CO2

42. Carbon Monoxide • Effects of CO on humans • 200-800 ppm leads to headache, fatigue, dizziness in a few hours • Death would occur under these conditions if exposed for several hours • 1000 ppm exposure can lead to death in < 1 h • CO interferes with oxygen (O2) transport • The shape and size is about the same as O2 • Hemoglobin is tied up with CO

43. Reactions of Metals • Corrosion reactions can weaken structural metal • 4 Fe + 3 O2 ----> 2 Fe2O3 (rust) • Tarnishing of Silver • 2 Ag + H2S ----> Ag2S + H2 • Removing tarnish from silver • 2 Al + 3 Ag2S ----> 6 Ag + Al2S3

44. Nitrogen Reactions • Ammonia Fertilizer • N2 + 3 H2 ----> 2 NH3 • Liquid ammonia is easily stored and applied (-33 oC) • Ammonia contains high nitrogen content • Vehicle Airbags • 2 NaN3 ----> 2 Na + 3 N2 • Rapid formation of nitrogen gas inflates the airbag