Grade 11: Chemistry Notes and Exam Review

Grade 11: Chemistry Notes and Exam Review Physical & Chemical Changes Physical Change Chemical Change Definition ● No new substance is produced ● Substance remains the sameeven with a change of state ● May require addition of energy ● Release of energy may occur ● Final substance is substantiallydifferent than initial substance ● New substance is alwaysproduced ● Energy is usually released butmay be required to get thechange going Properties ● Outside may look different ● Inside remains the same ● Particles may be rearranged ● Forces of attraction betweenparticles may be weaker orstronger ● A new substance is produced ● The particles of the newsubstance do not resemble thoseof the old substance ● Internally, the substanceproduced is different than the oldsubstances Examples ● Mixing sugar and water ● Ice melts into water ● Solid wax ==> Liquid wax ● Vinegar and baking soda mix toform carbon dioxide ● Hydrochloric acid reacts withmagnesium metal to formhydrogen gas Physical vs. Chemical Properties Physical properties: properties that use our senses or machines to identify a substance. Chemical properties: properties that describe how a substance reacts with another. Examples of Physical Properties: State: Solid, Liquid, Gas Malleability: how easily you can change the shape of something (eg. Aluminum foil). Solubility: how much of something can be dissolved in water. Viscosity: how easily a substance flows. Melting/Boiling Points NOMENCLATURE:

NAMING OF CHEMICAL COMPOUNDS International Union of Pure and Applied Chemistry (IUPAC) came up with naming conventions. We can divide the compounds into two main types: 1. Those which are true binary compounds (that is, contain only two types of elements). 2. Those which contain more than two types of elements. I. NAMING OF BINARY COMPOUNDS The names of all compounds containing only two elements end in IDE . Binary compounds may be subdivided into two types: 1. Those whose first element is a metal 2. Those whose first element is a non-metal In both cases the second element is a non-metal . 1. For binary compounds whose first element is a metal, we use the following system: [name of first element (metal)] [stem] + [IDE] The stem is merely an abbreviation for the name of the second element (the non-metal). Example: Naming the compound NaCl The symbol Na represents the metallic element sodium . The symbol Cl represents the non-metallic element chlorine , whose stem is chlor . Therefore, the name NaCl is sodium chlor+ide or sodium chloride. Likewise, for the following compounds: CaO is calcium oxide CaC 2 is calcium carbide AlN is aluminum nitride K 2 S is potassium sulfide Note that hydrogen is considered as a metal when it is written first in a binary compound. Valence Electrons: these are the outer most electrons in the last energy shell or level. ¨ Skipping the transition metals, we count from 1 to 8 to determine valence electrons. Ionization Energies: how much energy we need to remove one electron This decreases down the group, but increases across the period. The farther a valence electron is from the nucleus, the easier it is to remove it. If you go down a group (remember that in group 1, all have 1 valence electron but more and more total electrons) the ionization energy decreases). . BUT

As you add more valence electrons, it becomes harder to remove an electron because there is more negative chargeto be attracted to the positive nucleus. So the ionization energy increases across the period from left to right. http://web.archive.org/web/20130726231545/http://easystudy.ca/wp-content/uploads/2011/01/ionization-energy.png https://drive.google.com/file/d/1J-TTgJ46t5OlMNHLRD5Hr5_DHWhu1s-x/view?usp=share_link Atom Radii: the radius of an atom The atomic radius increases from top to bottom . This is because you are adding more and more electrons and but also more energy shells. BUT Across the period, atomic radii decrease because electrons are being added to the outer most shell (not creating newones). As more electrons are added, the positive nucleus tightens its grip like a belt and shrinks the atom. http://web.archive.org/web/20130726231545/http://easystudy.ca/wp-content/uploads/2011/01/atomic-radius.png Electron Affinity: This is the amount of energy created when an electron is added (on board) . This is the opposite of the ionization energy. There isn’t really a strong pattern, but it does tend to increase across the period from left to right .The halogens (group 7) have the highest electron affinity.

http://web.archive.org/web/20130726231545/http://easystudy.ca/wp-content/uploads/2011/01/electron-affinity.png Electronegativity: This is how much an atom wants to share another atom’s electrons. Think of it as how badly it wants anotherelectron. This increases across the period from left to right and decreases down a group . Electronegativity is used mainly to describe the types of chemical bonds an atom forms, not the atom itself. Why would group 7 have the highest electronegativities and group 1 the lowest? http://web.archive.org/web/20130726231545/http://easystudy.ca/wp-content/uploads/2011/01/electronegativity1.png Molecules: Chemical Bonding ¨ All atoms want to be stable. ¨ In order to be stable they need 8 electrons in their outer orbital. (Octet rule) ¨ When atoms from different elements temporarily have the same number of electrons we say they are isoelectronic . ¨ Atoms can either steal or share electrons. ¨ This is chemical bonding. ¨ Bonding: interaction between valence electrons of atoms. Property Ionic Compound Covalent Compound State at room temp. Crystalline solid Liquid, gas, solid Melting point High Low Electrical conductivity as aliquid Yes No Solubility in water Most have high solubility Most have low solubility Conducts electricity whendissolved in water(electrolyte) Yes Not usually Ionic bonds: one atom takes an electron and the other loses an electron. Example of groups that form ionic bonds: ¨ Halogens – 7 valence electrons

¨ Alkali metals – 1 valence electron ¨ Halogens will often “steal” an electron from an alkali metal. ¨ This forms a “bond” between the two atoms. Example: Recall: Diagrams that show ONLY valence electrons are called Lewis-Dot Diagrams. http://web.archive.org/web/20130827161507/http://easystudy.ca/wp-content/uploads/2011/01/NaCl1.png http://web.archive.org/web/20130827161507/http://easystudy.ca/wp-content/uploads/2011/01/NaCl2.png Why the positive and negative symbol? When sodium and chlorine combine they are make a compound that consists of positive sodium ions and a negative chlorine ions . Note: the negative Chlorine atom is now isoelectronic with Argon because both have 18 electrons. Conductivity ¨ Ionic compounds are good electrolytes ¨ Once dissolved in water or in liquid state, ions separate ¨ Can easily conduct charge (electricity) If there are more than one atom involved, for example CaF 2 , each fluorine will steal one electron from calcium leaving Ca 2+ and F - , F - Covalent Bonds: atoms “share” electrons so that all atoms achieve a stable octet. Example: Group 16 – 6 valence electrons Group 1 – 1 valence electron Oxygen (6 valence electrons) + 2 Hydrogens (1 valence electron each) = H 2 0 Covalent vs. Polar Covalent ¨ If the difference in electronegativity between both atoms is less than 1.7 but greater than 0.5 ¨ Both atoms share electrons, but unequally. ¨ The “shared” electrons will spend more time with one atom than another.

¨ One end of the molecule becomes more positive and one end more negative. http://web.archive.org/web/20130626230415/http://easystudy.ca/wp-content/uploads/2011/01/BohrDiagram.png Pure Covalent Bonds ¨ If a bond forms between two of the same kind of element both electronegativities will be the same. ¨ They will equally share electrons. ¨ No polar ends occur. Example: F 2 http://web.archive.org/web/20130626230415/http://easystudy.ca/wp-content/uploads/2011/01/F21.png Intermolecular and Intramolecular Forces In chemistry there are several different ways atoms and molecules can interact. These interactions take place within molecules and between molecules. Inter = between Intra = within Intermolecular Forces Intramolecular Forces Hydrogen Bonding Ionic Van der Waals/London Dispersion Polar Covalent Van der Waals/Dipole-Dipole Pure Covalent This course deals predominantly with Intramolecular Forces. For a detailed look at Intermolecular forces takeSCH4U. Intermolecular Forces are related to a substance’s PHYSICAL properties. Intramolecular Forces are related to a substance’s CHEMICAL properties. Intermolecular forces Relative strength of Intermolecular Forces:

● Intermolecular forces (dispersion forces, dipole-dipole interactions and hydrogen bonds) are much weakerthan intramolecular forces (covalent bonds, ionic bonds or metallic bonds) ● dispersion forces are the weakest intermolecular force ● dispersion forces < dipole-dipole interactions < hydrogen bonds Dispersion Forces (London Forces, Van der Waal’s Forces) ● are very weak forces of attraction between molecules resulting from: ● momentary dipoles occurring due to uneven electron distributions in neighbouring molecules as theyapproach one another ● The more electrons that are present in the molecule, the stronger the dispersion forces will be. ● Dispersion forces are the only type of intermolecular force operating between non-polar molecules, for example, dispersion forces operate between hydrogen (H 2 ) molecules, chlorine (Cl 2 ) molecules, carbon dioxide (CO 2 ) molecules, dinitrogen tetroxide (N 2 O 4 ) molecules and methane (CH 4 ) molecules. Dipole-dipole Interactions ● are stronger intermolecular forces than Dispersion forces ● occur between molecules that have permanent net dipoles ( polar molecules), for example, dipole-dipole interactions occur between SCl 2 molecules, PCl 3 molecules and CH 3 Cl molecules. Hydrogen Bonds § When hydrogen bonds to a very electronegative atom, it’s one electron spends most of its time with that new atom. § Leaves bare nucleus with proton (very positive) § This can form a temporary bonds between the hydrogen end of a molecule and the more negative end of anothermolecule. Nomenclature So far we have named compounds containing a non-metal and a metal, with only one possible valence. Part II Metals with Multiple Valences: The Stock and Old Method Formulas and Names of Binary Metal-NonmetalCompounds with More than One Valence: 1. IF the metal has more than one possible charge a. With the Stock Method you must indicate which charge in roman numerals (ie: FeCl 2 Iron (II) chloride). b. Alternatively the common (old) name may be used if the metal has more than one possible ion. Here use theLatin root and then add -ous for the lower charge. -ic for the higher charge.

i. FeCl 2 ferrous chloride ii. FeCl 3 ferric chloride i.e. Fe 2+ = ferrous, Fe 3+ = ferric (IC is bigger than OUS!) ¨ If given a formula, you must always look to the subscript on the non-metal to figure out the valence on the metal (reverse cross over). ¨ Be careful to check your periodic table for possible valences just in case that compound uses the lowest common denominator. Example: HgO Oxygen MUST have a valence 2-, therefore we know this formula was simplified from Hg 2 O 2 . HgO mercury(II)oxide More examples showing the two different systems: Compound Stock Method Common Name FeF 2 iron (II) fluoride ferrous fluoride FeF 3 iron (III)fluoride ferric fluoride HgBr 2 mercury (II) bromide mercuric bromide Polyatomic Compounds 1. Names of Polyatomic Ions a. Anions are negative, Cations are positive b. ammonium ion NH 41+ c. -ide ions i. CN 1- cyanide ii. OH 1- hydroxide d. Oxyanions i. -ate ate more oxygen. Formula Name NO 21- nitrite NO 31- nitrate ii. Sometimes oxyanions have an extra hydrogen Formula Name SO 42- Sulfate HSO 41- hydrogen sulfate (or bisulfate) iii. If more than two possibilities: Formula Name ClO 1- Hypochlorite ClO 21- Chlorite

ClO 31- Chlorate ClO 41- Perchlorate Naming compounds with polyatomic ions a. Positive charge species on left (using Stock method or common name) b. Negative charge species on right (using name of polyatomic ion) c. Use parentheses as needed Formula Ions Name BaSO 4 Ba 2+ and SO 42- barium sulfate Ca(NO 3 ) 2 Ca +2 and NO 31- calcium nitrate Ca(NO 2 ) 2 Ca +2 and NO 21- calcium nitrite Fe(NO 3 ) 2 Fe 2+ and NO 31- iron (II) nitrate orferrous nitrate Formulas and Names of Binary Nonmetal-Nonmetal Compounds 1. Systematic Nomenclature: a. For names start with element to the left side on the periodic table b. add -ide to the second element c. use Greek prefixes for number of atoms: mono, di, tri, tetra, penta, hexa, hepta, octa, nona, deca d. Example: i. CO carbon monoxide ii. CO 2 carbon dioxide iii. N 2 O 5 dinitrogen pentoxide 2. Common names: -ous and -ic ( -ic has greater charge, OR has fewer atoms). Examples: Formula Systematic Name Common Name NO nitrogen monoxide nitric oxide N 2 O dinitrogen monoxide nitrous oxide NO 2 nitrogen dioxide nitrogen peroxide N 2 O 5 dinitrogen pentoxide nitric anhydride N 2 O 3 dinitrogen trioxide nitrous anhydride 3. Hydrates (calcium sulphate dehydrate, cobalt (II) chloride hexahydrate) Oxyanions ● Oxyanions (negatively charged polyatomic ions which contain O) end in “-ate”. However, if there is morethan one oxyanion for a specific element then the endings are:

Two less oxygenthan the mostcommon startswith “hypo-” andends with “-ite” One lessoxygen thanthe mostcommon endswith “-ite” THE MOSTCOMMONOXYANION ENDSWITH “-ATE” One more oxygen thanthe most common startswith “per-” and ends with“-ate”< ClO - = hypochlorite ● ClO 2- =chlorite ● NO 2- = nitrite ● SO 32- =sulfite Most commonoxyanions withfour oxygens ● SO 42- = sulfate ● PO 43- = phosphate ● CrO 42- = chromate Most commonoxyanions withthree oxygens ● NO 3- = nitrate ● ClO 3- = chlorate ● CO 32- = carbonate ClO 4- = perchlorate ● Polyatomic anions (a negatively charged ion containing more than one type of element) often add ahydrogen atom; in this case, the anion’s name either adds “hydrogen-” or “bi-” to the beginning Example: CO 32- becomes HCO 3- “Carbonate” becomes either “Hydrogen Carbonate” or “Bicarbonate” Hydrogen as a Non-Metal If hydrogen is second in a binary compound, the compound ends in –ide. Example: NaH Sodium hydride Nomenclature: Acids 1. Hydro Acids: Hydro + halogen name + ic Formula Name HCl hydrochloric acid HF hydrofluoric acid 2. OxoAcids: polyatomic ion + acid. a. Recognize as polyatomic ions with a hydrogen at the beginning of the formula. b. Name with -ous and -ic suffix. (Works just like -ite and -ate suffix) c. -ic suffix is for acid with more oxygen atoms.

d. Examples Formula Name Source HNO 3 nitric acid nitric from nitrate HNO 2 nitrous acid nitrous from nitrite These acids have hydrogen as the positive ion and a radical containing oxygen as the negative ion. If the radicalends in ATE , take off the ate, put on IC and add the word acid. Examples : HNO 3 is nitric acid HClO 3 is chloric acid H 2 CrO 4 is chromic acid H 2 SO 4 is sulfuric acid H 3 PO 4 is phosphoric acid If the radical ends in ITE , take off the ite, put on OUS and add the word acid. Examples : HNO 2 is nitrous acid H 2 SO 3 is sulfurous acid HClO is hypochlorous acid Moles and Mass Recall 6.02 x 10 23 of anything represents one mole. The mass of one atom can be taken from the periodic table. 1 Carbon – 12 atom has a mass of 12 u. The mass of one mole of atoms is the same number, but in grams. 6.02 x 10 23 of carbon-12 atoms has a mass of 12 g. The molar mass is the mass of one mole of atoms. Mole Ratio

The coefficients in a balanced chemical equation can be used to determine therelative number of molecules, or moles of a compound involved in a chemicalreaction. Example: N 2 (g) + 3 H 2 (g) –> 2 NH 3 (g) 1 molecule of nitrogen (N 2 ) reacts with 3 molecules of hydrogen (H 2 ) to form 2 molecules of ammonia (NH 3 ) OR 1 mole of nitrogen (N 2 ) reacts with 3 moles of hydrogen (H 2 ) to form 2 moles of ammonia (NH 3 ). The coefficients in a balanced equation can be used to write a molar ratio. Molar ratios are conversion factors that can be used to relate: 1. moles of product formed from a certain number of moles of reactant 2. moles of reactant needed to form a certain number of moles of a product. 3. the number of moles of a particular reactant needed to completely react with acertain number of moles of a second reactant. For the following reaction: 4 NH 3 (g) + 5 O 2 (g) –> 4 NO (g) + 6 H 2 O (g) 4 moles NH 3 5 moles O 2 5 moles O 2 4 moles NO 4 moles NH 3 4 moles NO 4 moles NH 3 6 moles H 2 O 4 moles NO6 moles H 2 O 5 moles O 2 6 moles H 2 O THE MOLE - The mole is defined as a unit of measurement to measure the amount of matter in a substance - i.e. 16.1 mole of Carbon; means that the amount of carbon in the sample is 16.1 mol - 1 mol of substance contains (6.022)(10 23 ) particle (atoms, molecules, etc.) - (6.022)(10 23 ) is a special number, called Avogadro’s Number (Named after the man who discovered it

- Examples ● 1 mol of Al contains (6.022)(10 23 ) atoms ● 1 mol of S 8 contains (6.022)(10 23 ) molecules ● 1 mol of C 12 h 24 0 11 contains (6.022)(10 23 ) molecules ● 1 mole of NaCl contains (6.022)(10 23 ) Formula Units - The amount of matter in moles is directly related to the number of particles( atoms, molecules, etc.) in a sample of matter - Therefore we gather this equation: ● Number of moles= (Number of atoms, molecules, or Formula Units)/(Avogadro’s number) ● This is shortened to n=N/N A ● “n” is number of moles in mol ● “N” is Number of atoms, molecules, or Formula Units ● “N A ” is Avogadro’s Number (6.022)(10 23 ) - Carbon-12, each C-12 has a mass of 12 amu (atomic units) - A sample of carbon has weighted average atomic mass of 12.011 amu - 1 mole of any substance has a mass in grams equal to its weighted average atomic mass (given on periodic table, don’t forget to calculate for total atomic mass of the substance by adding up all the atomic masses of each atom(s)!!! ) - Therefore we can we can gather another equation: ● Number of moles= (Entire mass of substance)/(Total Atomic mass of entire substance) ● Shortened to n=m/MM ● “n” is moles in mol ● “m” is mass in grams(g) ● “MM” is the total atomic mass of substance, also known as molar mass, is measured in grams per mole(g/mol), and is found by adding up all the atomic masses for each atom(s) in a molecule - Tip : From n=N/N A and n=m/MM we can derive the formula N/N A = m/MM PERCENT COMPOSITION - Law of Definite proportions: the element in chemical compounds are always present in the same proportion, by mass - Percent composition are ratios by the masses of an element in a compound , this is represented by the formula: (mass of atom(s))/(Total mass of the substance) - In a chemical equation, the ratios are mole to mole

- i.e. 1NaCl —> 1Na + 1Cl, assume one mole of NaCl and find the percent composition of Cl ● 1 mol of NaCl ● 1 mol of Na ● 1 mol of Cl ● Therefore m Cl = (1 mol of Cl)( MM Cl ) and m Na = (1 mol of Na)(MM Na ) ● m Cl = 22.989768 g and m Na = 35.4527 g ● Therefore % Cl = (mCl)/(mNa + MCl) * 100% ● % Cl = 60.689252% —— don’t forget sig figs —–à = 60.6892% THE EMPIRICAL FORMULA - The simplest ratio of elements for a substance - Ie: CH 2 0; simplest ratio of elements for C 6 H 12 O 6 (Glucose) - It is found by finding the number of moles of each individual atom in a substance then dividing each of these by the smallest amount. These numbers are representative of the number of atoms there are in a single molecule orion for that atom. - Sample Question: A sample is found to contain 5.0 g of copper and 1.3 g of oxygen, and no other elements. Determine the empirical formula of this compound. ● What do we need? n Cu = ? and n O = ? ● n Cu = m Cu /MM Cu and n O = m O /MM O ● n Cu = 5.0g/(63.546g/mol) and n O = 1.3g/(15.9994g/mol) ● n Cu = 0.078683 mol and n O = 0.081253 mol ● Divide by small therefore n Cu / n Cu = 1 and n O / n Cu = 1.032667= 1 (this is because we cannot have decimals in a chemical formula) ● Since they are both 1 therefore the ratio of Cu to O is 1:1, therefore the Empirical Formula is Cu 1 O 1 or Cu Excess and Limiting Reagents In a chemical reaction, reactants that are not use up when the reaction is finished are called excess reagents . The reagent that is completely used up or reacted is called the limiting reagent , because its quantity limit the amount of products formed. Let us consider the reaction between sodium and chlorine. The reaction can be represented by the equation: 2 Na + Cl 2 = 2 NaCl, Thus, if you have 6 Na atoms, 3 Cl 2 molecules will be required. If there is an excess number of Cl 2 molecules, they will remain unreacted. We can also state that 6 moles of sodium will require 3 moles of Cl 2 gas. If there are more than 3 moles of Cl 2 gas, some will remain as an excess reagent, and the sodium is a limiting reagent. It limits the amount of the product that can be formed. Example 1

Calculate the number of moles of CO 2 formed in the combustion of ethane C 2 H 6 in a process when 35.0 mol of O 2 is consumed. Hint … The reaction is 2 C 2 H 6 + 7 O 2 = 4 CO 2 + 6 H 2 O 4 mol CO 2 35.0 mol O 2 ---------- = 20.0 mol CO 2 7 mol O 2 Example 2 Two moles of Mg and five moles of O 2 are placed in a reaction vessel, and then the Mg is ignited according to the reaction Mg + O 2 = MgO. Balance this equation and identify the limiting reagent in this experiment. Hint … The balanced reaction is, 2 Mg + O 2 = 2 MgO Thus, two moles of Mg require only ONE mole of O 2 . Four moles of oxygen will remain unreacted. Oxygen is the excess reagent, and Mg is the limiting reagent. Acids and bases in everyday life - The tartness of lemons and oranges comes from the weak acid citric acid. The acid is found widely in nature and in many consumer products (Charles D. winter) - The sting of ants is due to the weak acid formic acid, IDS and Bases in everyday life - Aspirin is a weak acid that has been used as an analgesic for over 100 years. (Charles D. Winter) - Glycine is representative of the amino acids that are the basis of proteins. The – co2h group is the acid portion of the molecule, and the –NH2 group is the basic portion (Charles D, winter) Acids and Bases in everyday life - Caffeine is a well known stimulant and a weak base (Charles D. Winter) - A sea slug excretes the strong acid in self defense. (sharksong/ M. Kazmers/Dembinski photo associates) Acids - Tastes sour - No characteristic feel

- Turns litmus paper red - Turns red phenolphthalein colourless - Conducts an electric current - Reacts with “active metals” to produce hydrogen gas - Reacts with carbonates to produce carbon dioxide gas Bases - Tastes bitter - Feels slippery - Turns litmus paper blue - Turns colourless phenolphthalein red - Conducts an electric current - Reacts with “fats/oils” to make soap The oldest theory is Arrhenius Theory - Arrhenius looked at the properties of acids. - Arrhenius said these properties were due to the production of H+ ions when acids dissociate in water. - HCl(g) à H+(aq) + Cl-(aq) - Arrhenius also looked at the properties of bases - Arrhenius said these properties were due to the production of OH- ions when bases are dissociated in water - NaOH(s) à Na+(aq) + OH- - Arrhenius’ Theory had a few questions that could not be answered - Why isn’t water included in the dissociation of an acid? - HCl(g) à H+(aq)+Cl (aq) - Where is the OH- when ammonia dissociates? - NH3(aq) à OH-?? - I think we are going to need a NEW theory ? Bronsted – Lowry Theory - Bronsted was Danish - Lowry was British

- Working at the same time, yet independent of one another, they came up with virtually the same theory and published at virtually in the same time - He fixed the NH3 theory because NH3–> NH4 + is possible A Bronsted Lowry ACID is a species - Which donates a proton in a proton transfer reaction - This was later re – written to read: - A species that is willing to have a proton removed - For the acid – base reaction: - HCl(aq) + KOH(aq) à H20+ KCl(aq) - Which species is the acid For the general acid – Base reaction - HA(aq) + B - (aq) à HB(aq) + A - (aq) - Which species is the acid A Bronsted – Lowry Base is a species which is capable of removing a proton from an acid - That is the species that takes the proton - For the acid – base reaction: - HBr(aq) + NaOH(aq) à H20(l) + NaBr(aq) - Which species is the base? Answer: NaOH - This reaction is a neutralization Why do we keep referring to protons?? - A “proton” is really just a hydrogen atom that has lost its election! (i.e. A hydrogen ion) - By definition, H+ is a proton Conjugate acid – Base pairs - When an acid loses its proton, the species that is left is called a conjugate base. - The conjugate base would now be ready to remove a proton from another acid. - When a base gains a proton from an acid, the species that is left is called conjugate acid. - The conjugate acid would now be ready to have a proton removed from another base. - Ex. HCl (aq)+H20àH30+(aq)+cl- - HCl(aq) is the acid

- H20(l) is the base - H30+(aq) is the conjugate acid - Cl-(aq) is the conjugate base Memorize 6 strong acids - HClO 4 - HI(aq) - HBr(aq) - H 2 SO 4 - HCl(aq) - HNO 3 - Why do we call them strong acids? - A strong acid will completely dissociate (100%) into its ions in aqueous solution. - H 2 SO 4 (aq) à 2H+(aq) + SO 4 2-(aq) Weak acids to know… - While any acid that is not on the list of strong acids is a weak acid because it will not dissociate completely, when dissolved in water. - The ones you should know are: - CH 3 COOH – acetic acid - CO 2 (aq) – aqueous carbon dioxide - SO 2 (aq) – Aqueous sulphur dioxide - HF (aq) – Hydrofluoric acid (strong bond) - HCN – hydrocyanic acid For a specific weak acid reaction: CH 3 COOH + H 2 O -> H 3 O + + CH 3 COO - - With a weak acid, like acetic acid, we have an equilibrium mixture of all four species. - Meaning that the rate of the forward reaction is equal to the rate of the reverse reaction - Weak acids do NOT dissociate 100% Memorize 6 strong bases - LiOH

- NaOH - KOH - CA(OH) 2 - SR(OH) 2 - Ba(OH) 2 - All group 1 hydroxides make strong bases - A strong base will completely dissociate into its ions in aqueous solution - Ca(OH) 2(aw) à Ca 2+(aq) +2OH - (aq) Weak bases to know - NH 3 – Ammonia - R-NH 2 – Amines (class of organic compound. R can be any element.) - A weak base will not dissociate fully For specific base acid reaction: NH 3 + H 2 O <-> OH - + NH 4+ - Much like a weak acid, a weak base will result in an equilibrium mixture of all four species. Meaning that the rate of ammonia dissolving is equal to the rate of ammonium hydroxide decomposing - Weak bases do NOT dissociate 100% Why do we write… NH 3 instead of H 3 N? CH3COOH instead of C 2 H 4 O 2 ? - The molecular formula is written for clarity. That is, clearly distinguish between acids and bases; and try to ensure that the proton being transferred is clearly identified. The Hydrogen is written separately so you can seewhere it gets pulled off. For a specific weak acid reaction: CH 3 COOH + H 2 O -> H 3 O + + CH 3 COO - - CH 3 COOH is the acid - H 2 0 is the base - H 3 O+ is the conjugate acid - CH 3 COO- is the conjugate base For specific weak base reaction:

NH 3 + H 2 O = OH - + NH 4+ - NH 3 is the base - H 2 O is the acid - OH - is the conjugate base - NH 4+ is the conjugate acid - NH 3 is the weak base. There are not many ions in solution. The strong smell is due to NH 3 (g) CH 3 COOH + H 2 O = H 3 0 + + CH 3 COO - NH 3 +H 2 O = OH - + NH 4+ AMPHOTERIC/AMPHIROTIC - According to the Bronsted –Lowry model… - A substance capable of acting as an acid or base in different chemical reactions… - That is, a substance that may donate or accept a proton. - Ex. Amphibian pH and the Autoionization of water - In any sample of water, there are some ions which result from the dissociation of water itself - H 2 O = H + + OH - - H 2 O + H 2 O = H 3 O + + OH - - Experiments show that in pure water at standard conditions (25degree, 101.3 kPa) - [H 3 O + ] = [H + ] = [OH - ] = 1.0 X 10 -7 M - Which is what we call a neutral solution - Water solutions in which - [H+]>[OH-] - Are acidic solutions - Water solutions in which - [OH-]>[H+] - Are basic solutions - You must keep in mind that is water solutions of acids and bases there are always OH- and H3O+ ions due to the self ionization of water.

- These ions from water do not add many ions to reasonable concentrations of acid and base solutions, but you must remember that they are there - For our purposes we will treat these concentrations as being negligible. - Since [H3O-], [H+], and [OH-] are so small, scientists use the concept of “p” when dealing with very small numbers, where “p” stands for “power” … on a logarithmic scale. - Ex. pH stands for “power of hydrogen” - pH =-log [H3O+] or =-log [H+] - pOH = -log [OH-] will measure concentration of hydroxide ion On a logarithmic scale, - Every increase by 1 pH represents a 10 fold increase in [H 3 O + ] GAS LAWS Standard Temperature and Pressure Standard Pressure (for theoretical purposes): 760 mm Hg = 760 torr = 1 atm = 101.3 kPa Standard Temperature: 0 celcius = 273 K Standard Ambient Temperature and Pressure (SATP) Real-life conditions: 25 degrees celcius and 100 kPa Doing Pressure Conversions I. between atmospheres and millimeters of mercury. One atm. equals 760.0 mm Hg, so there will be a multiplicationor division based on the direction of the change. Example #1 – Convert 0.875 atm to mmHg. Solution – multiply the atm value by 760.0 mmHg / atm. Notice that the atm values – one in the numerator and one in the denominator – cancel, leaving mmHg. Example #2 – Convert 745.0 mmHg to atm. Solution – divide the mmHg value by 760.0 mmHg / atm Notice that the mmHg values cancel and the atm, in the denominator of the denominator, moves to the numerator. II. between atmospheres and kilopascals. One atm equals 101.325 kPa, so there will be a multiplication or divisionbased on the direction of the change.

Example #3 – Convert 0.955 atm to kPa. Solution – multiply the atm value by 101.325 kPa / atm. Notice that the atm values – one in the numerator and one in the denominator – cancel, leaving kPa. Example #4 – Convert 98.35 kPa to atm. Solution – divide the kPa value by 101.325 kPa / atm. Notice that the kPa values cancel and the atm, in the denominator of the denominator, moves to the numerator. III. between millimeters of mercury and kilopascals. 760.0 mmHg equals 101.325 kPa, so both values will beinvolved. This situation is slighly unusual because most conversions involve a one, usually in the denominator. Theconversion examples above are examples of a one being involved. In this conversion, both 760.0 and 101.325 will be involved and the location of each (numerator or denominator) willdepend on the conversion. Example #5 – Convert 740.0 mmHg to kPa. Notice that the mmHg will cancel, since one is in the numerator and one is in the denominator, leaving kPa as the uniton the answer. Example #6 – Convert 99.25 kPa to mmHg. Notice that the kPa will cancel, since one is in the numerator and one is in the denominator, leaving mmHg as the uniton the answer. Ideal Gases ¨ In a perfect world gases would obey all these laws perfectly. Unfortunately, conditions are not always optimal. This is what we call ideal gases. ¨ But, we will use them to do theoretical calculations anyway because they are the best approximation we have. ¨ If we combine all the laws from the last lesson and add that to a little common sense we can get a new equation to work with. ¨ add that n (number of moles) is proportional to volume pv = nRT pressure x volume = number of moles x constant x temperature in Kelvin This is the ideal gas law. Example: What mass of neon gas should be introduced into an evacuated 0.88-L tube to produce a pressure of 90kPa at 30°C? Solution v = 0.88 L