Chapter two: MOLECULES & SOLUTIONS

1. Chapter two: MOLECULES & SOLUTIONS

2. WHAT IS A MOLECULE? • Most atoms are found combined with other atoms in nature; exceptions: He, Au… • Molecule: 2 or more atoms that are chemically linked together behaving as an entity • An atom’s goal is to be as stable as possible; therefore an atom will react with other atoms so it may fill its valence shell • Octet (8) rule:elements will react with other atoms in order to acquire the same electron configuration as the noble gas closest to them; ex. F will gain an electron to have the same electron configuration as Ne; some elements like Li will follow the duet (2) rule when attempting to become stable & will follow He’s electron configuration • - H is a special case since it will tend to lose an electron at times but it may also gain en electron in other circumstances

5. 1.1 IONS • Ion: is a charged atom; becomes charged by losing or gaining an electron • Becoming an ion never changes the number of protons involved • Positive ion has lost 1 or several electrons, ex. Al3+ has lost 3 electrons; negative ion has gained 1 or several electrons, ex. Cl- has gained 1 electron • - Elements of a period will all have the same valence electron configuration; • ex. halogens, group VIIA will all tend to gain an electron to form an ion that looks like Br-

9. Polyatomic ions • Ex. SO42- (sulphate), OH- (hydroxide), CH3COO- (acetate), PO43- (phosphate), CO32- (carbonate), • HCO3- (bicarbonate) • Polyatomic ion: a group of linked atoms which are positively or negatively charged • Ex. Na2CO3 (sodium carbonate) Na+ + CO32- • Ca3(PO4)2 (calcium phosphate) Ca2+ + PO43- • - Fig.2.10, p.44 – Common polyatomic ions.

10. 1.2 THE NATURE OF CHEMICAL BONDS • - Chemical bond: when 2 or more atoms become linked due to the transfer or sharing of 1 or several electrons; this forms a new substance • 2 types of chemical bonds: • Ionic bonds: • often between a metal & a nonmetal; metals tend to lose electrons (becoming + ions) & nonmetals tend to gain electrons (becoming – ions); therefore the metal tends to give up electrons to the nonmetal which will readily accept them; the attraction between + & - ions forms ionic bonds; ex. NaCl, NaOH (sodium hydroxide), Na2SO4 (sodium sulphate), • KI (potassium iodide), LiBr (Lithium bromide), MgCl2 (magnesium chloride)

17. (2) Covalent bonds: a link between 2 or more atoms that happen when 1 or more electron pairs are shared; usu. between 2 nonmetals; covalent is from Latin & means ‘joint power’ • Ex. CO2, CH4, C6H12O6, H2O, O3, N2, C6H22O11, O2, F2….. • Several types of covalent bonds: • (1) Single bond: 1 pair of electrons is shared; shown by a single line linking atoms; ex. ethane C2H6, notice that ethane has 7 single bonds; draw a Lewis diagram to the right of the image to visualize what is happening

19. (2) Double bond: 2 pairs of electrons are shared; shown by drawing a double line across atoms; ex. ethane, C2H4, ethane has 1 double bond (& 4 single bonds); draw a Lewis diagram to the right of the image

20. (3) Triple bond: 3 pairs of electrons are shared; symbol: triple line across atoms; ex. ethylene, C2H2 has 1 triple bond (& 2 single bonds); draw a Lewis diagram

21. (4) Polar covalent bond: When the electrons of a covalent bond are unequally shared between the atoms the result is a covalent compound that has slight negative & positive charge in certain areas of its molecule; ex. H2O: the H atoms are slightly positive & the O atom is slightly negative; this means the negative charge is localized around the O atom because the electron distribution is asymmetric; meaning the negative charge migrates towards the O; O is therefore said to be electronegative, i.e., able to attract electrons more than the H atoms; it is because of this polarity that water can easily dissolve substances (NaCl + H2O: Na+ would be attracted O side & Cl- would be attracted to H side pulling NaCl apart & dissolving the salt into the water) & it is also because of this polarity that water molecules are attracted to each other forming a capillary action; polarity (or partial charge) is symbolized by the Greek letter delta

22. - HF (hydrogen fluoride) is another polar covalent compound (image right); where the H is slightly positive & F is slightly negative; the arrow shows where there is most electronegativity

23. 1.3 THE RULES OF CHEMICAL NOTATION & NOMENCLATURE • We use rules of notation & nomenclature to give the ratio of atoms within a compound & to name the compounds • - The following rules apply mainly tobinary molecules, molecules containing 2 atoms

24. The Rules of Notation (writing a molecular chemical formula) • Use the periodic table to determine the symbol for each element • (2) Determine the order of the elements • For a binary compound containing a metal & a nonmetal: the metal is 1st & the nonmetal is 2nd, ex. CaCl2 • - In all other cases, the following order is respected: B, Ge, Si, C, Sb, As, P, N, H, Te, Se, S, I, Br, Cl, O, & F, ex. You would write CO2 & not O2C since C is before O in this list • (3) Use subscriptswhen there is more than one atom; to determine the subscripts for ionic bonds the ‘strength’ of the ion must be considered; ex. a bond between Mg2+ & Cl- could not be MgCl since Mg has a ‘strength’ of 2 & Cl has a ‘strength’ of 1 instead MgCl2 would be correct; • Determine the formulas for the compounds these ions could form (ref. table 2.10, p.44): • 1. PO43- & Na+ 3. CrO42- & K+ • 2. Mg2+ & OH- 4. ClO3- & Ca2+

25. Determine the formulas for the compounds these ions could form (ref. table 2.10, p.44): • 1. PO43- & Na+ 3. CrO42- & K+ • 2. Mg2+ & OH- 4. ClO3- & Ca2+

26. The Rules of Nomenclature (naming molecules) Ionic Compounds (metal + nonmetal): The positive ion (metal) is named first & the negative ion (nonmetal) is named last Ex. HCl (H+ & Cl-) hydrogenchloride (2) The same rule applies for polyatomic ions: positive ion first, negative ion last Ex. NaOH (H+ & OH-) sodium hydroxide (3) Change the name of the second element so that it finishes in ‘-ide’; fig.2.17, p.50; KCl: potassium chloride& not potassium chlorine

27. Covalent Compounds (nonmetal + nonmetal): Name the first element following the order given in #2 of Rules of Notation; name the second element (2) If there are multiple atoms add the correct prefix (image right) -Ex. N2F4: dinitrogentetrafluoride; SF6: sulphur hexafluoride

28. The prefix mono is used only when 2 elements can form more than one compound together in order to reduce confusion; ex. COcarbon monoxide, since it also forms CO2,carbon dioxide • - Note: compounds can have more than 1 name, especially common compounds, for example, hydrogen chloride is also called hydrochloric acid which usually implies this substance is dissolved in water; same would be correct for HF hydrogen fluoride: hydrofluoric acid

30. Acids 7 Bases pH and acid-base indicators: Indicator: Substance which changes colour according to the environment in which it is situated (Acid or base). They don’t give a precise reading of the pH value of a solution but they remain useful. pH: Hydrogen potential. A level of acidity or alkalinity of a substance based on the concentration H+ and OH- ions. pH scale:

31. 1 2 3 4 5 6 7 8 9 10 11 12 13 14 x 10 x 102 On the pH scale, a jump of one corresponds to a power of ten. A solution having a pH of 3 is ten times more acidic than a pH of 4. Whereas a solution having a pH of 2 is one hundred times more acidic than a pH of 4. x 10

32. pH and concentration : From the concentration of an acid or base, we can deduce its pH level. The concentrations are based on the availability of the Hydride ion (H+1) All concentration of acids and bases, for now, are in negative powers of ten. Next year, you’ll see, we can have other types of concentrations. For example: 1. Acids: 0.001 mol/L = 1/1000 = 1-3 hence: pH is 3 0.0001 mol/L = 1/10000 = 1-4 hence: pH is 4 0.00001 mol/L = 1/100000 = 1-5 hence: pH is 5 2. Bases: 0.000000001 mol/L = 1/1000 000 000 = 1-9 hence: pH is 9 0.0000000001 mol/L = 1/10 000 000 000 = 1-10 hence: pH is 10 0.0000000001 mol/L = 1/100 000 000 000 = 1-11 hence: pH is 11

33. pH and concentration : Exercise: Find the pH of an HF solution knowing that 2.0 x 10-3 g are dissolved in 100 mL of solution.

34. ACID NEUTRAL H+(aq) OH-(aq) BASE

36. Ionic compounds: we learnt in the past that ionic compounds usually are substances that contain a metal bonded to a non-metal or group of non-metals. But really, ionic means that this substance once dissolved in water would dissociate into a positive ion (cation) and negative ion (anion). Since, acids, bases and salts all dissolve and dissociate in water, they are also considered ionic too. Also, substances that dissolve in water like acids, bases and salts can conduct electricity. They are called electrolytic. See next slide.

37. - - - - - - - - - - - + + + + + + + + + + + Cathode (-) Anode (+) Here we have salt (NaCl) which is ionic and therefore dissolves in water. It splits into ions. These ions conduct electricity because they are mobile.

38. Covalent molecule Covalent molecule Covalent molecule Covalent molecule Covalent molecule Covalent molecule Covalent molecule Covalent molecule Covalent molecule Covalent molecule - + Cathode (-) Anode (+) Here we have sugar (C12H22O12) which is covalent and therefore does not dissolves in water. It does not splits into ions, it stays as a molecule and therefore does not conduct electricity.

41. Acids Inorganic H-NM H-R where R is a radical group or a polyatomic ion that does not contain carbon Organic R-H where is a radical or a polyatomic ion that contains carbon Example: C2H5COOH →C2H5COO-1 + H+1 Examples: HCl(aq)→H+1 + Cl-1 HNO3(aq) →H+1 + NO3-1 H2SO4(aq) →H+1 + HSO4-1 HSO4-1(aq) →H+1 + SO4-2 Note that (aq) stands for aquaeous which means it is dissolved in water. Since it is dissolved in water, it dissociates to form ions.

42. Bases M-OH R-OH Where R is a radical or a polyatomic ion that does not include a carbon atom NaOH(aq)→ Na+1 + OH-1 NH4OH(aq) → NH4+1 + OH-1 Ba(OH)2(aq) → Ba+2 + 2 OH-1

43. Salts M-NM M-R NaCl(aq) → Na+1 + Cl-1 BaCl2(aq) → Ba+2 + 2 Cl-1 Na2CO3(aq) → 2 Na+1 + CO3-2 Ba(NO3)2(aq) → Ba+2 + 2NO3-1

44. (2) PROPERTIES OF SOLUTIONS • A mixture occurs when atoms & molecules combine together by physical means (no chemical reaction); recall that a mixture can be separated by using physical separation techniques (filtration, evaporation, sedimentation, decantation…) • A solution is a homogeneous mixture that has the same appearance even when seen through a magnifying instrument, i.e. phases cannot be seen; a solution is made up of a solute which dissolves into a solvent; ex. brass (Cu + Zn), air, sea water • We will look at 4 properties of solutions: solubility, concentration, electrical conductivity, & pH • An aqueous solution is a solution in which water is the solvent; ex. Seawater • - Ionic compounds & molecules with a degree of polarity, like sugar, can easily dissolve in water, whereas non-polar molecules like oil & gasoline cannot

45. Organization of matter Mixtures Pure substances Homogenous • Solutions • Alloys Heterogeneous Compounds Elements • Suspensions • Colloids • Covalent bonds • Share electrons • Metalloids-none-metal (M-NM) • H-NM-O • True bond • Transfer of electrons • M-NM • H and NM • Forms ions Acids Bases Salts

46. Note: Ion: a charged atom. It is not neutral. # of protons does not equal to the # of electrons Positive ion: cation Negative ion: anion Radical: polyatomic ion

47. 2.1 SOLUBILITY • Solubility: the maximum amount of solute that can be dissolved in a solvent at a certain T; point at which a solution is saturated; with added T more solid solute can dissolve • Gases show the opposite trend: the higher the T, the less a gas will dissolve in it (image right) • - Solubility is also dependent on the nature of the solute & solvent & the pressure (the higher the pressure the more gas can dissolve)

49. 2.2 CONCENTRATION • - The quantity of solute in a given solution • Dilution & Concentration • Concentration can be expressed in 4 ways: • g/L: grams solute contained in 1 litre of solution • (2) % m/V: grams solute per 100mL of solution • (3) % V/V: milliliters solute per 100mL of solution • (4) %m/m: grams solute per 100g of solution • Concentration can also be expressed in ppm (next section) • A formula is used to calculate concentration in g/L: • C = m/V • C = concentration in g/L • m = mass in g • V = volume in L

50. Concentration in PPM • Unit of concentration used when the amount of solute is very, very low in a solution • - Examples: used to express the concentration of pollutants, such as CO2,in the atmosphere, right or the concentration of minerals and elements found in spring water, next slide.