reintroduction to Chemistry

1. reintroduction to Chemistry

2. Matter: Anything that has mass and volume States of Matter

3. Physical Changes • Boiling point • Melting point • Freezing point • Color • Density Properties that do not change the chemical nature of matter

4. Chemical ChangesChanges that do change the chemical nature of matter • Oxidation Rusting of iron = Iron Oxide (FeO2) • pH • Flammability Frying an egg

5. Physical & Chemical Changes

6. All Matter is composed of : • Atoms – Means “unable to be cut” • The smallest particle of an element that still has the characteristics of that element

7. Elements 96% of human mass is made up of C, H, O, N • Substances that can’t be broken down into simpler substances

8. Periodic Table of The Elements

9. Atoms are made up of smaller subatomic particles • Protons: positively charged (Located in the nucleus) • Neutrons: neutrally charged (Located in the nucleus) • Electrons: negatively charged (Located around the nucleus) Discovered by James Chadwick in 1932 Discovered by Ernest Rutherford in 1919 Discovered by J.J. Thomson in 1897

10. Each atom has the same number of protons and electrons Electrons located farther away from the nucleus have more energy Electrons have less mass than protons. 1 proton= 1870 electrons

11. Atomic number = number of protons Atomic mass = number of protons & neutrons Electron # = Proton # in Neutral elements

12. 2 He 4.003 Atomic number: Atomic mass: # of Protons # of Electrons # of Neutrons

13. Isotopes Neutron #’s change 1 Hydrogen: H 1 1 Deuterium: H 2 1 H Tritium 3

14. When elements combine to form substances with two or more atoms… Compounds are formed

15. Mixtures When one or more atoms are physically combined but NOT chemically combined (Individual atoms retain their own properties) Easily separated by non-chemical means

16. Solutions • Solute – Substance dissolved • Solvent – Dissolves substances (Water is the universal solvent) • Colloids – light does not penetrate • Suspensions

17. Solubility of Solutions • The amount of solute a solvent can dissolve Solubility affected by: • Temperature • Concentration

18. Intramolecular Bonds • Forces of attraction that hold atoms together within a molecule. • Stronger than intermolecular bonds • 2 types covalent ionic

19. Ions Ions are charged particles. Atoms have either lost or gained one or more electrons

20. Ionic bond Cation Anion

21. Redox Reactions Important in ATP synthesis Oxidation is a loss of electrons Reduction is a gain of electrons

22. Redox Reactions

23. Covalent Bond Occurs when atoms SHARE electrons Special Note Molecules are only formed with covalent bonds Diatomic molecule

25. Molecules • A group of atoms held together bycovalentbonds • Water, hydrogen gas, oxygen, carbon dioxide

26. Intermolecular Bonds • Intermolecular bonds refers to the forces of attraction that hold molecules together • Considerably weaker than intramolecular bonds • 3 types of intermolecular bonds Hydrogen bonds Van der Waals forces Molecule-ion attractions

28. Uneven distribution of electrons results in a polar molecule

29. – Hydrogenbonds + H – + H + –  – + Figure 3.2  – +  – + +  –  – +

30. Hydrogen Bond A bond formed in polar molecules

32. Cohesion & Adhesion: Due to Hydrogen Bonding Cohesion – The ability of hydrogen bonds to attract like molecules Adhesion – The ability of hydrogen bonds to attract different types of molecules

33. Why can these animals walk on water?

34. Capillary Action isDue to Adhesion

35. How does the crown get water?

36. Van der Waals forces are weak attractive forces that hold non-polar molecules together

37. molecule-ion attraction Hydration shell forms as water surrounds the dissolved ions

38. H2O The molecule that supports all of life

39. Molality vs Molarity • Molality – the # of moles solute which can be dissolved in 1 kg of solvent • Molarity – the # of moles solute which can be dissolved in 1 liter of soloution

40. Water Important biological properties • Cohesion • Adhesion • High specific heat • Heat of vaporization • Solvent of life • Insulation of bodies of water (floating ice) in winter

41. Heat & Temperature • Heat – measurement of the total amount of kinetic energy in matter • Temperature – measurement of the average kinetic energy of molecules Measurement of heat calorie (cal) = amount of heat required to raise the temperature of 1 g of water by 1°C - or – the amount of heat that 1 g of water releases when cooled by 1°C Calorie (kcal) = 1000 cal = amount of heat required to raise the temp of 1 kg of water by 1°C

42. Specific Heat • The amount of heat that must be absorbed or lost for 1 g of that substance to change its temperature by 1°C • Important because water does not change temperature quickly & large bodies of water can store much heat to warm the air.

43. Specific Heat of Water • Heat is absorbed when hydrogen bonds break • Heat is released when hydrogen bonds form

44. Evaporative Cooling • Due to the high specific heat, when water evaporated it removes much heat from the system, thereby cooling the organism

45. Latent Heat Latent Heat is the heat given up or absorbed by a substance as it changes state. It is called latent because it is not associated with a change in temperature

46. Heat of Vaporization The quantity of heat a liquid must absorb for 1 g of it to be converted from the liquid to the gaseous state

47. Hydrogen bond Liquid water Hydrogen bonds constantly break and re-form Ice Hydrogen bonds are stable Figure 3.5 The hydrogen bonds in iceAre more “ordered” than in liquid water, making ice less dense

48. Negative oxygen regions of polar water molecules are attracted to sodium cations (Na+). – Na+ + + – + – – Positive hydrogen regions of water molecules cling to chloride anions (Cl–). Na+ – + + Cl – Cl– + – – + – + – – Figure 3.6 The different regions of the polar water molecule can interact with ionic compounds called solutes and dissolve them

49. Acids & Bases Acids are formed by hydrogen cations Bases are formed by hydroxide anions

50. Acids • Donate protons (Hydrogen Ions) to water to form hydronium ions • pH 0-6.99 • Taste Sour • Turn litmus paper red • Strong acids completely dissociate to form ions