Why Study Chemistry?

1. Why Study Chemistry? • Martin Jones, Christina Miller, “Chemistry in the Real World”, J.Chem.Ed., 78, 2001, 484-487. • Improving health care • Conserving natural resources • Protecting the environment • Food, clothing, shelter • “understand materials & properties/develop new materials • Always positives and negatives • Disposable Diapers

2. Why Study Chemistry? • Properties of matter (physical material of the universe), something unique • Chemistry – “atom & molecule” perspective • Composition & structure • Ethyl alcohol vs. Ethylene glycol • Cellulose & Amylose • Two realms – “macroscopic” & “submicroscopic”

4. Standard PS-3 • The student will demonstrate an understanding of various properties and classifications of matter. • Chemical/Physical Properties • Molecule/Atom • Pure Substance/Mixture • States of Matter • Phase Change

5. Standard PS-3 • 3.1: Distinguish chemical properties of matter (including reactivity) from physical properties of matter (including boiling point, freezing/melting point, density [with density calculations], solubility, viscosity, and conductivity. • Physical properties may be easier to start with…. substance remains unchanged

6. Physical Properties of Matter • No chemical change – substance unchanged • Phase Changes • Boiling Point (Condensation Point) • Melting Point (Freezing Point) • Density • Solubility • Viscosity • Electrical Conductivity

7. Physical Properties of Matter • Phase Changes (Changes of state): freezing condensation Solid  Liquid  Gas melting boiling • Temperature at which the “phase changes” -- Boiling Point -- Condensation Point -- Melting Point -- Freezing Point

8. Physical Properties of Matter • Density – “ratio” of mass to volume • Requires 2 measurements • Mass….. balance • Volume….. Depends on object (solid, liquid, gas) • Solid • Regular object (cube, cylinder)…physical measurements • Irregular object….displacement method • Liquid • Volumetric glassware • Gas • Volume of the container

11. Density of a Substance • Depends on the state of matter • Water

12. Densities of Some Substances Substance Density (g/cm3) • Air 0.001 • Balsa wood 0.16 • Ethanol 0.79 • Water 1.00 • Table Salt 2.16 • Iron 7.9 • Gold 19.32

13. Physical Properties of Matter • Solubility • Dissolves in a solvent  “Soluble” • Maximum amount (grams) that will dissolve in a “given amount” of solvent (100 mL, 1 liter,….) • “High” solubility to “low” solubility • 0.55 g lead iodide (PbI2) in 1 Liter of water • Barium sulfate • <0.01 mol/L “insoluble”

14. Physical Properties of Matter • Solubility • Solvent – most common is water • Again, no chemical change • Use caution as some solutes may undergo chemical reaction – “reactivity with water” • Sodium • Potassium

15. Physical Properties of Matter • Electrical Conductivity • Complete a circuit (current flows) • Conductors: allow current to flow • Insulators: no current flows • “High conductivity” – conductors • “Low conductivity” – insulators • Solutions conduct electricity depending on substance & solvent. If so…..”electrolyte”. • Electrolytes may be strong or weak.

16. Chemical Properties of Matter (including reactivity) • Chemical Change – matter undergoes a change forming a “new” substance • A “chemical reaction” has taken place. • Example “types” of reactions: • Burning (oxidation) • Rusting (oxidation) • Explosions (decomposition)

17. Chemical Properties of Matter (including reactivity) • Chemical property – often depend on the context • Potential chemical reactions a substance can undergo. • Capacity to oxidize (combine with oxygen) • New substance: original substance + oxygen

18. Chemical Properties of Matter (including reactivity) • Combustibility • A type of oxidation • Release heat and light • 2H2(g) + O2(g)  2H2O(g) • Capacity to corrode • A type of oxidation, usually metals • Metal + non-metal

19. Pure Substance • Distinct properties • Composition doesn’t vary • Examples: • Water • Table salt (sodium chloride) • Magnesium • Two (or more) pure substance  mixture • Homogeneous (Solutions) • Heterogeneous • Mixtures separated (different physical properties)

20. Separating Mixtures • Methods: • Solubility • Filtration • Distillation • Chromatography

22. Pure Substances • Pure Substance • Do Chemical Reactions  • Can’t be broken down  “Element” • Can be broken down  “Compound” • Aluminum  Element (100+ known) • Water  Compound • Pure Substance: either elements or compounds

24. Atoms & Molecules • Element: smallest unit is “atom” • Contain only ONE TYPE of atom • Misconception: element composed of “atoms” • Oxygen (ozone) • Carbon – diamond, graphite, fullerenes • Elements  Periodic Table • Compound: smallest unit is “molecule” • Two or more different TYPES of atoms

26. Pure Substances • Compounds: more than one type of atom • Molecular: covalently bonded (shared) • Ionic: opposite ions attracted • C12H22O11 “molecular formula” • NaCl  “formula” (or empirical formula) • Na+Cl- (more information)

27. Standard PS-3 • 3.5: Explain the effects of temperature, particle size, and agitation on the rate at which a solid dissolves in a liquid.

28. Standard PS-3 • Why do these have an effect? • Kinetic Theory • Assumptions: • Small particles • Constant motion (average) • Collisions occur

29. Standard PS-3 Solution Formation (molecular scale) • Solvent (water) • Constantly moving, rolling over but attracted • Solute (sugar) • Also constantly moving, not rolling over but attracted • Forming solution does not change substance

30. Standard PS-3 Solution Formation (molecular scale) • Surfaces in contact • Solute surrounded by solvent, disperse (Remove solvent  solute “re-attracts”

31. Standard PS-3 Solution Formation (molecular scale) • Temperature • Higher temperature, higher kinetic energy, more collisions between solvent & solute • Particle size • Smaller particles – more surface area, more collisions • Agitation • Mechanical force, more collisions • Ultrasonic cleaners

32. Standard PS-3 Solution Formation (molecular scale) • Temperature, Particle size, Agitation: • Affect “rate”, not whether it will dissolve • “Rate of Dissolving” and “Solubility – different concepts

33. Standard PS-3 • 3.6: Compare the properties of the four states of matter – solid, liquid, gas, and plasma – in terms of the arrangements and movement of particles.

34. States of Matter • Solid: • constantly moving, but can’t slip past • Can’t pour

35. States of Matter • Liquid: • constantly moving, but can slip past • Can pour

36. States of Matter • Gas: • constantly moving, particles far apart • Shape of the container

37. States of Matter • Plasma: • High temperature • Particles broken into ions, electrons • Most common form  “universal scale”

38. States of Matter • Caution: • At given temperature, solid particles NOT moving slower than liquid • Same temperature  Same kinetic energy • Phase changing – energy goes into overcoming “intermolecular” force”

39. Standard PS-3 • 3.7: Explain the processes of phase change in terms of temperature, heat transfer, and particle arrangement. • Heat: a type of energy • Temperature: measure of the average kinetic energy • Faster moving particles  higher kinetic energy  higher temperature

40. Standard PS-3 • Phase Change: (Kinetic Theory view) • Melting (Freezing): when enough energy added, overcome attractive forces • Boiling (Condensation): can occur at any temperature, if pressure change • Boiling occurs when the “vapor pressure” is the same as “ambient pressure”

41. Standard PS-3 • Phase Change: (Kinetic Theory view) • Temperature change is the “evidence that energy has been added, indicator of kinetic energy • Boiling: escape from the surface (bubbles) • Sublimation: melting point/boiling point very close

42. Standard PS-3 • Caution: • Heat (energy content) not equal to temperature • Pot of water & cup of water at the same temperature • Pot of water has more heat (energy)