1. What is the chemical formula for water? Draw the structure of water. Write down all the types of bonding that you know of.

2. What allows this drop of water to hang there without falling?

3. Surface Tension Water Clings High Surface Tension – water molecules form hydrogen bonds with each other creating cohesion A water strider can walk on the surface of a pond. Capillary Action – water molecules from hydrogen bonds with other polar molecules creating adhesion Capillary action causes water to creep up a narrow glass tube and paper.

4. Hydrogen Bonding in Water Hydrogen bonding causes water to absorb a large amount of heat before its temperature increases appreciably and also causes it to lose large amounts of heat before its temperature decreases significantly. (Heat Capacity) High heat capacity allows organisms to maintain a constant body temperature. Hydrogen bonding causes liquid water to absorb a large amount of heat to become a vapour. Many organisms (including humans) dissipate body heat by evaporation of water from surfaces (skin-sweating; tongue-panting)

5. Chemical Context of Life  Matter (space & mass)  Element; compound  The atom  Atomic number (# of protons); mass number (protons + neutrons)  Isotopes (different # of neutrons); radioactive isotopes (nuclear decay)  Energy (ability to do work); energy levels (electron states of potential energy)

6. Chemical Bonding  Covalent  Double covalent  Nonpolar covalent  Polar covalent  Ionic  Hydrogen  van der Waals

7. Covalent Bonding  Sharing pair of valence electrons  Number of electrons required to complete an atom’s valence shell determines how many bonds will form  Ex: Hydrogen & oxygen bonding in water; methane

8. Covalent bonding Practice some examples of Covalent Bonding using Lewis structures. What other properties do covalent bonds have?

9. Polar/nonpolar covalent bonds  Electronegativity attraction for electrons  Nonpolar covalent electrons shared equallyEx: diatomic H and O  Polar covalent one atom more electronegative than the other (charged)Ex: water

10. Polar/nonpolar bonds Recall ► Covalent bonds may be polar or nonpolar Nonpolar covalent – electronegativity difference = 0 Polar covalent – electronegativity difference is greater than 0 but less than 1.7 Molecular polarity – dependant on both bond polarity and molecular shape Symmetrical molecules whose bonds are all polar are nonpolar. Asymmetrical molecules are nonpolar of all bonds are nonpolar, and they are polar if at least one bond is polar.

11. Ionic bonding  High electronegativity difference strips valence electrons away from another atom  Electron transfer creates ions (charged atoms)  Cation (positive ion); anion (negative ion)  Ex: Salts (sodium chloride)

12. Ionic bonds Practice drawing ionic compounds using lewis structures. What are some other properties of ionic bonds?

13. Hydrogen bonds  Hydrogen atom covalently bonded to one electronegative atom is also attracted to another electronegative atom (oxygen or nitrogen)

14. van der Waals interactions  Weak interactions between molecules or parts of molecules that are brought about by localized change fluctuations  Due to the fact that electrons are constantly in motion and at any given instant, ever-changing “hot spots” of negative or positive charge may develop

15. Water  Polar~ opposite ends, opposite charges  Cohesion~ H+ bonds holding molecules together  Adhesion~ H+ bonds holding molecules to another substance  Surface tension~ measurement of the difficulty to break or stretch the surface of a liquid  Specific heat~ amount of heat absorbed or lost to change temperature by 1oC  Heat of vaporization~ quantity of heat required to convert 1g from liquid to gas states  Density……….

16. Density  Less dense as solid than liquid  Due to hydrogen bonding  Crystalline lattice keeps molecules at a distance

17. Acid/Base & pH  Dissociation of water into a hydrogen ion and a hydroxide ion  Acid: increases the hydrogen concentration of a solution  Base: reduces the hydrogen ion concentration of a solution  pH: “power of hydrogen”  Buffers: substances that minimize H+ and OH- concentrations (accepts or donates H+ ions)

18. Two major parts of an atom Nucleus (not to scale) Electron Cloud

19. Three Major Sub-Atomic Particles Protons Neutrons Electrons

20. a single, relatively large particle with a positive charge that is found in the nucleus PROTON (p + )

21. THE PROTON p+p+ Fat (heavy) Positive (charge) Doesn’t move (lazy)

22. a single, relatively large particle with a neutral charge that is found in the nucleus NEUTRON ( N° )

23. THE NEUTRON N°N°  Fat (heavy) Neutral  (charge) Doesn’t move (lazy)

24. a single, very small particle with a negative charge that is found in a “cloud” around the nucleus ELECTRON (e - )

25. THE ELECTRON Skinny (very light) Negative  (charge) Moves a lot (runs around) e-e-

26. Review: Subatomic Particles e-e- N°N°  p+p+

27. Please complete the following table ProtonsNeutronsElectrons Where are they found? Mass Charge (attitude) Nucleus Electron Cloud Heavy Very Light Positive Neutral  Negative 

28. The total mass of all of the subatomic particles in an atom (but really # of protons and neutrons) ATOMIC MASS # (A)

29. the number of protons in an atom (assuming the atom is neutral, # of p + = # of e - ) ATOMIC NUMBER (Z)

30. Example: Sodium Na 11 22.99 Atomic # = # of protons Atomic Mass # = p + & N °

31. Another Notation Atomic # = # of protons Atomic Mass # = p + & N °

32. To calculate the number of neutrons, subtract the atomic number (smaller) from the atomic mass number (larger) A – Z = # of neutrons

33. Ex: How many neutrons does Sodium have? Mass # - Atomic # = #N° (You may need to round the atomic #) 23 - 11 = 12 N° Na 11 22.99

34. Atoms of the same element that differ in charge. (They have the same # of p +, but different # of e - ) ION

35. Positive Ions (cations) Negative Ions (anions) Na + (lost 1 e - ) Ca 2+ (lost 2 e - ) Al 3+ (lost 3 e - ) Pb 4+ (lost 4 e - ) H + (lost 1 e - ) Cl - (gain 1 e - ) O 2- (gain 2 e - ) P 3- (gain 3 e - ) S 2- (gain 2 e - ) OH - (gain 1 e - )

36. If an atom GAINS electrons, its overall charge becomes more negative. If it LOSES electrons, its charge becomes more positive

37. Atoms of the same element that differ in mass. (They have the same # of p +, but different # of N°) ISOTOPE

38. Isotopes are CHEMICALLY the SAME as atoms, but DIFFER PHYSICALLY because they have different masses.

39. A few examples of isotopes…

40. Complete the following table ProtonsNeutronsElectrons Na + Br w/ mass 84 O 2- with mass 13