1. Spring, 2012 Session 3 – General Chemistry Pt 1

2.  Definition of terms  Chemical formulas  Chemistry background  Reactions  Equilibrium and law of mass action  pH and acid base reactions  Coagulation  Water constituents

3.  Inorganic chemistry  Organic chemistry  Atomic structure – protons, neutrons, electrons  Atoms and molecules  Elements and compounds

4.  Shorthand notation to express which elements are in a molecule and in what ratio  For example, Fe 2 O 3 has 2 atoms of iron (Fe) for every 3 atoms of oxygen (O)  If atoms are ionized, they are combined to produce a neutrally charged molecule  Use your handout to explain what atoms are in each of the molecules on the following page and in what ratio

5. · CO · CaO · C 2 H 2 · CH 4 · Ca(OH) 2 · C 2 H 4 · CO 2 · SO 2 · C 2 H 6 · H 2 O · N 2 · H 3 (PO) 4 · HNO 3 · MgO · ZnO · H 2 SO 4 · HF · P 2 O 5 · NaOH · CaCO 3 · H 2 S

6.  Composed of chemical formulas  Show the proportions in which molecules or atoms react to form products  Reactants are shown on left hand side and products on right hand side  Numbers before chemical formula signify the number of molecules that go into the reaction  The same number of atoms of a species will be on both sides of the equals sign  For example, ZnS + 2HNO 3 = Zn(NO 3 ) 2 + H 2 S

7. For the reactions on the following slide explain:  what the reactants and products are  what the ratio these molecules are in Balance the following equations so that there is the same number of atoms of a species in the products as there is in the reactants

8.  CH 4 + O 2 = CO 2 + H 2 O  H 2 SO 4 + NaOH = Na 2 SO 4 + H 2 O  CH 3 SH + H 2 = CH 4 + H 2 S  C 6 H 6 + H 2 = C 6 H 14  C 7 H 16 = C 3 H 6 + C 4 H 8 + H 2  Fe + O 2 = Fe 2 O 3  FeS + O 2 = Fe 2 O 3 + SO 2  Ca(OH) 2 + H 2 SO 4 = CaSO 4 + H 2 O

9.  Chemical equations can also help you determine the amount of reactants to combine to produce a desired amount of products  Each molecule combines with the others in terms of its molecular weight  For example, one molecular weight of NaOH would combine with one molecular weight of HCl to obtain one molecular weight each of NaCl and H 2 O based upon the following chemical equation NaOH + HCl = NaCl + H 2 O

10.  Atomic and molecular weight (tabulated)  Atoms combine to form molecules  Obtain atomic weights or mass from chart (next slide) or from the following website http://www.chemicalelements.com/ http://www.chemicalelements.com/  A nice website with a lot of interactive information on chemistry is http://serendip.brynmawr.edu/sci_edu/chemsit es.html http://serendip.brynmawr.edu/sci_edu/chemsit es.html

12. Determine the molecular weight of a compound  Molecular weight is the sum of the: number of atoms times the atomic weight (mass) for each  The number of atoms is shown in the chemical formula – i.e. H 2 O has two atoms of hydrogen and one of oxygen  NaCl molecular weight is 1*23 + 1*35 = 58  Ferrous Sulfate, Fe 2 (SO 4 ) 2, molecular weight is: 2 * 55.8 + 2 * (1*32.1 + 4 * 16) = 303.8  Units are grams/gram mole or pounds/pound mole  Used to calculate amounts of chemicals to prepare solutions for analysis or for processes  Determine the molecular weights of the reactants and products in each of the reactions in the following slide. (You may need to balance the equations first.)

13.  CH 4 + O 2 = CO 2 + H 2 O  H 2 SO 4 + NaOH = Na 2 SO 4 + H 2 O  CH 3 SH + H 2 = CH 4 + H 2 S  C 6 H 6 + H 2 = C 6 H 14  C 3 H 8 + C 4 H 10 = C 7 H 16 + H 2  C 7 H 16 = C 3 H 8 + C 4 H 8  Fe + O 2 = Fe 2 O 3  FeS + O 2 = Fe 2 O 3 + SO 2  Ca(OH) 2 + H 2 SO 4 = CaSO 4 + H 2 O

14.  Precipitation (removal of phosphate)  Oxidation or reduction (CH 4 combustion)  Neutralization (addition of an acid to neutralize a base or vice versa)  Decomposition reactions – breaking one molecule into two or more products (like catalytic cracking)  Combination reactions – combining two molecules to make one or more products (like reforming)

15. Determine the types of the reactions on the following slide

16.  CH 4 + O 2 = CO 2 + H 2 O  H 2 SO 4 + NaOH = Na 2 SO 4 + H 2 O  CH 3 SH + H 2 = CH 4 + H 2 S  C 6 H 6 + H 2 = C 6 H 14  C 7 H 16 = C 3 H 6 + C 4 H 8 + H 2  Fe + O 2 = Fe 2 O 3  FeS + O 2 = Fe 2 O 3 + SO 2  Ca(OH) 2 + H 2 SO 4 = CaSO 4 + H 2 O

17.  Reactions need time, temperature and turbulence to take place effectively  Allow sufficient residence time  Reactions go faster at higher temperatures  Should provide good mixing  Examples are catalytic cracking, reforming, disinfection by chlorination, pH adjustment, pulp production, combustion

18.  Bonding concerns how the outer electrons in the atomic shell interact between two atoms to bind the two together  General classes of bonds are:  Ionic – electrons are lost or gained  Covalent – electrons are shared  Inorganic compounds generally have ionic bonds  Column 1 periodic table elements usually have a + 1 charge, column 2 a +2 charge  Column 17 elements usually have a – charge and column 18 have a 0 charge  Organic compounds (carbon bonded to hydrogen) generally have covalent bonds

20.  Molecules have a neutral charge  This means that the combination of positive and negative ions in a molecule must balance, i.e. for a O 2- ion we will need two H + ions to balance the charge to make one molecule of neutral water  Practice charge balancing the following reactions to make neutral products

21.  H + + O -2 =  H + + (SO 4 ) -2 =  H + + S -2 =  Ca +2 + (OH) -1  Fe +3 + O -2 =  H + + (PO 4 ) -3 =  Mg +2 + (SO 4 ) -2 =

22.  Tendency of distribution or reaction in two directions  Example is liquid vapor equilibrium in a distillation tower ( a species condensers from the vapor at the same rate as the same species vaporizes from the liquid)  This is physical equilibrium  Chemical equilibrium concerns reactions  Reactants are in equilibrium with products  Example of equilibrium: CO 2 + H 2 O ↔ H + + HCO 3 - the rate of reaction of carbon dioxide gas (CO 2 ) with water is in equilibrium with the rate of reaction of carbonate ion (HCO 3 - ) with H + ion to form carbon dioxide gas

23.  Law of mass action – adding more material to the left side of the reaction shifts the reaction to the right  Example;  HCO 3 - + Ca ++ + HO - ↔ CaCO 3 + H 2 O  Addition of base (HO - ) will cause more of the calcium (Ca ++ ) and bicarbonate (HCO 3 - ) to precipitate as CaCO 3  Phosphate (PO 4 -3 ) can be removed as precipitated calcium phosphate by adding calcium oxide (CaO). This is sometimes performed as tertiary treatment to rid water of nutrients (phosphates)

24.  pH = log [1/(H + )] (where H + concentration is in moles per liter)  pH of neutral water is 7  pH rises as the solution becomes more alkaline or basic  pH falls as the solution becomes more acidic  Can use a meter or pH paper to estimate pH  Restrictions on pH of discharged water  A fun website for pH is http://www.miamisci.org/ph/phlemon.html http://www.miamisci.org/ph/phlemon.html

25.  Colloids and coagulation  Colloids are very small particles that do not settle out (clay, organics, microscopic organisms)  Some are attracted to water, some not  Chemical coagulation is used to remove colloids  Coagulation is the use of chemicals to destablize colloids so that they can attract each other and settle  Coagulants include iron or aluminum sulfates  Flocculation is the process of attraction to form solids that settle  Polymers are often used as flocculation aids  Coagulation and flocculation are followed by settling

26.  Organic compounds present in wastewater  Alcohols (-OH), aldehydes (=O), carboxylic acids (-COOH), carbohydrates, fats, protein  Biodegradable organic content in wastewater is about 60 – 80%  Inorganic ions and compounds present in wastewater  Ions such as calcium, magnesium, phosphates, nitrates, carbonates  Gases such as hydrogen sulfide and oxygen