Basic Chemistry

1. Basic Chemistry

2. basic unit of matter 92 recognized elements 25 essential for life 6 major elements in living organisms carbon hydrogen nitrogen oxygen phosphorus sulfur CHNOPS ( 98% ) Chemical Elements

3. atom - smallest unit of matter that retains the properties of an element • - carbon is an element • - may have multiple atoms of carbon (only one element) • Subatomic particles: • ParticleLocationCharge Mass • Proton nucleus positive 1 AMU • Neutron nucleus neutral 1 AMU • Electron orbiting negative negligible • nucleus

4. 12 C 6 Carbon Atomic Symbols and the Periodic Table MassNumber AtomicSymbol AtomicNumber atomic # is the # of protons (unique to each element) Atoms begin electrically neutral thus: positive and negative charges balance out # of protons = # of electrons Mass # equals the # of protons + # of neutrons Each element has a unique # of electrons # of electrons --> influences interactions w/ other atoms

5. Isotopes: • atoms typically have equal # of protons and neutrons • - except hydrogen which only has one proton • atoms MAY vary in the # of neutrons they have (Isotopes) • - if not equal with # of protons = more unstable • - radioactive isotopes • - acts the same as stable isotopes but can be ‘traced’ • - used in medical diagnosis and research • - spontaneously break down at a constant rate • - can be used in dating fossils

6. Periodic Table Vertical columns indicatenumber of electronsin outermost shell VIII I 1 H 1.008 2 He 4.003 1 VII III IV VI V II Horizontal periods indicatetotal numberof electron shells 3 Li 6.941 4 Be 9.012 5 B 10.81 6 C 12.01 7 N 14.01 8 O 16.00 9 F 19.00 10 Ne 20.18 2 11 Na 22.99 12 Mg 24.31 13 Al 26.98 14 Si 28.09 15 P 30.97 16 S 32.07 17 Cl 35.45 18 Ar 39.95 3 19 K 39.10 20 Ca 40.08 21 Ga 69.72 22 Ge 72.59 23 As 74.92 24 Se 78.96 25 Br 79.90 26 Kr 83.60 4

7. Electrons have energy: • - electrons are attracted to positively charged nucleus • - energy absorbed pushes the electron away • creates potential energy • - energy is released when electron moves closer • releases energy for use in chemical reactions

8. The Octet Rule forDistribution of Electrons • Bohr models show electron shells as concentric circles around nucleus • Each shell has two or more electron orbitals • Innermost shell has two orbitals • Others have 8 or multiples thereof • Atoms with fewer than 8 electrons in outermost shell are chemically reactive • If 3 or less – Tendency to donate electrons • If 5 or more – Tendency to receive electrons

9. Bohr Models of Atoms

10. Atoms interact and form chemical bonds dependent upon electron configuration. - electrons arranged in shells - # empty spaces in outer shell = the # of potential binding sites open circle - available binding sites

11. Molecules: • are formed when two or more atoms are bound together • made up of different elements or the same element • If all atoms in molecule are of the same element • Material is still an element • 2 oxygen atoms form 1 molecule oxygen gas • If at least one atom is from a different element • Material formed is a compound • 2 hydrogen & 1 oxygen form 1 water molecule • Characteristics dramatically different from constituent elements

12. Molecular / Chemical formula used to represent the # of atoms of each type of element in a molecule: CH4 - means one carbon and four hydrogen atoms together H2O - means two hydrogen and one oxygen • Chemical Bonds: • chemical bonds hold atoms together in a molecule & influence how molecules interact with one another

13. Types of Bonds between atoms/elements: • 1. Ionic bonds • transfer of electrons between two atoms • 2. Covalent bonds • sharing of electrons between two atoms • - equal sharing - non-polar covalent bonds • - unequal sharing - polar covalent bonds • 3. Hydrogen bonds • weak electrostatic charge - opposites attract • bond between polar moleculesor polar areas of a large molecule

14. Ionic Bonds: • reactivity - tendency to lose or gain electrons • full outer valence shell = stability IONS: Na+ / Cl- electrically charged atoms • transferring electron • formation of IONS

15. IONIC BOND: fig 2.7 • ions -- opposites attract • ionic bond - moderate strength bond • Na+ & Cl- form NaCl (sodium chloride) • fully charged atoms (lost or gained an electron) • strong (and opposite) electrical charges attract one another

16. double covalent bonds harder to break Covalent Bonds: fig 2.8 • sharing electrons • # of binding sites • dependent on outer electron shell • atoms share ONE pair of electrons • single covalent bond • atoms share TWO pairs of electrons • double covalent bonds

17. Non-polar Covalent Bonds: fig 2.8 • equal sharing of electron • equal pull on electron • molecule electrically neutral Polar covalent bonds: examples: H2, CH4 • covalent bonds share electrons • unequal sharing • electron spends more time around one of the atoms in bond • molecule has slightly negative side & slightly positive side fig 2.9 examples: H2O, NH3

18. fig 2.9b Hydrogen bonds: polar molecule • electronegative atom • (an atom with a slightly stronger pull on a shared electron) • attracts a HYDROGEN atom already w/in a polar covalent bond in a different molecule or another part of the same molecule excellent example: water

19. Ionic vs. Hydrogen bonds: stronger bond than hydrogen Hydrogen: two polar molecules or parts of molecules attracted to one another - slight electrical charge Ionic: ions (charged atoms) attracted to one another - strong electrical charge

20. Properties of Water: Water is a polar molecule due to polar-covalent bonds that are created. As such, water is held together by hydrogen bonds. This polarity and subsequent hydrogen bonding creates a number unique properties that are essential to life on Earth. 1. Its heat capacity 2. Its cohesive nature 3. Its reaction when forming a solid (ice) 4. Its use as a solvent 5. Its ability to ionize

21. as water is cooled…. bonds are formed first heat released as converted to solid as water is heated…. heat breaks the bonds between water molecules before water can vaporize! 1. Temperature Buffer -- Heat Storage water changes temperature slowly heat energy is absorbed to break bonds bonds formed before movement of molecule slows Temperature regulation in humans: • evaporative cooling

22. Water has a high heat capacity • - Temperature = rate of vibration of molecules • - Apply heat to liquid • - Molecules bounce faster • - Increases temperature • - But, when heat applied to water • - Hydrogen bonds restrain bouncing • - Temperature rises more slowly per unit heat • - Water at a given temp. has more heat than most liquids • Thermal inertia – resistance to temperature change • - More heat required to raise water one degree than most other liquids (1 calorie per gram) • - Also, more heat is extracted/released when lowering water one degree than most other liquids

23. High heat of vaporization • - To raise water from 98 to 99 ºC; ~1 calorie • - To raise water from 99 to 100 ºC; ~1 calorie • - However, large numbers of hydrogen bonds must be broken to evaporate water • - To raise water from 100 to 101 ºC; ~540 calories! • This is why sweating (and panting) cools • - Evaporative cooling is best when humidity is low because evaporation occurs rapidly • - Evaporative cooling works poorest when humidity is high because evaporation occurs slowly

24. Heat of fusion (melting) • - To raise ice from -2 to -1 ºC; ~1 calorie • - To raise water from -1 to 0 ºC; ~1 calorie • - To raise water from 0 to 1 ºC; ~80 calories! • This is why ice at 0 ºC keeps stuff cold MUCH longer than water at 1 ºC • This is why ice is used for cooling • - NOT because ice is cold • - But because it absorbs so much heat before it will warm by one degree

25. Heat Content of Waterat Various Temperatures

26. 2. Cohesion and Adhesion of Water • - Cohesion – Hydrogen bonds hold water molecules tightly together • - Adhesion – Hydrogen bonds form between water and other polar materials • - Hydrogen bonds -- breaking & reforming • - bonds last a few trillioniths of a second! • - always substantial % bonded to neighboring molecules

27. High Surface Tension • - Water molecules at surface hold more tightly than below surface • - Amounts to an invisible “skin” on water surface • - Allows small nonpolar objects (like water-strider) to sit on top of water

28. Allows water be drawn many meters up a tree in a tubular vessel

29. 3. Ice Formation - water molecules densest at 4°C -as temp drops bonds become more spaced & stable - density drops - frozen water less dense than liquid water Density of Waterat Various Temperatures

30. A Pond in Winter • Lakes/oceans don’t stay frozen over because ICE floats • Ice acts as an insulator on top of a frozen body of water • Otherwise, oceans and deep lakes would fill with ice from the bottom up • Melting ice draws heat from the environment • Turnover of nutrients / oxygen when lake thaws

31. 4. Water is the universal solvent • - A solvent (the most abundant part) and • - A solute (less abundant part) that is dissolved in the solvent • - Polar compounds readily dissolve; hydrophilic • - Nonpolar compounds dissolve only slightly; hydrophobic • - Ionic compounds dissociate in water • - Na+ • Attracted to negative (O) end of H2O • Each Na+ completely surrounded by H2O • - Cl- • Attracted to positive (H2) end of H2O • Each Cl- completely surrounded by H2O

32. break-up / surround polar molecules and ionized compounds Hydrophilic molecules: water-loving

33. 5. Water Ionizes • water molecules sometimes break • H2O <--> OH- + H+ • pH - measure of hydrogen ions (H+)concentration: • most chemical reactions w/in our bodies influenced by pH • most biological fluids act as buffers - neutralizing pH • basic solutions (high pH) - lower H+concentration • acidic solutions (low pH)- high H+concentration

34. Acids • Dissociate in water and release hydrogen ions (H+) • Sour to taste • Hydrochloric acid (stomach acid) is a gas with symbol HCl • In water, it dissociates into H+ and Cl- • Dissociation of HCl is almost total, therefore it is a strong acid • Bases: • Either take up hydrogen ions (H+) or release hydroxide ions (OH-) • Bitter to taste • Sodium hydroxide (drain cleaner) is a solid with symbol NaOH • In water, it dissociates into Na+ and OH- • Dissociation of NaOH is almost total, therefore it is a strong base • pH scale used to indicate acidity and alkalinity of a solution. • Values range from 0-14 • 0 to <7 = Acidic • 7 = Neutral • >7 to 14 = Basic (or alkaline)

35. The pH Scale

36. Health of organisms requires maintaining pH of body fluids within narrow limits • - Human blood normally 7.4 (slightly alkaline) • - Many foods and metabolic processes add or subtract H+ or OH- ions • - Reducing blood pH to 7.0 results in acidosis • - Increasing blood pH to 7.8 results in alkalosis • - Both life threatening situations • - Bicarbonate ion (-HCO3) in blood buffers pH to 7.4