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Lesson 1 - Working With Chemicals Safety is the number one concern WHMIS (Workplace Hazardous Materials Information System) was established to standardize information and symbols about chemicals in our lives o WHMIS informs workers about the chemicals in three ways 1. Controlled products must have informative labels, in both English and French on their containers 2. Each controlled product must have a MSDS (Materials Safety Data Sheet). 3. Workers who handle chemicals must complete an education program provided through their employer. http://www.brocku.ca/oehs/safety/whmis_symbols.pdf MSDS (Material Safety Data Sheet) was required to accompany every chemical bought and sold (p.7). http://www.lindecanada.com/en/msds/linde/Argon__Liquid_EN.pdf Classifying Matter Matter is defined as anything that has mass and volume Matter may be solid, liquid or gas. 1 Matter Mixture Pure Substances Combinations of matter that can be separated by physical means Matter that has a definite composition. Do not have a definite composition Heterogeneous Mixtures (Mechanical mixtures) Different components of the mixture are visible or layers can be seen. Gravel Homogeneous Mixtures Element Compound (Solutions) Cannot be chemically broken down into simpler substances Two or more elements that are chemically combined Different components are not visible Milk Gold Can be separated chemically into simpler substances Water (H2O) 2 Assignment: 1. State whether the following is a pure substance or a mixture: a) sea water b) iron c) bronze d) 14k gold e) table salt f) oxygen 2. State whether the following mixtures are homogeneous or heterogeneous: a) Oil-and-vinegar salad dressing b) Steel c) Cranberry juice d) Sugar dissolve in water e) Milk f) Antifreeze 3. State whether the following pure substances is an element or a compound: a) Copper b) Water, H2O c) Methane, CH4 d) Silver 4. Classify each of the following substances: a) Graphite, C b) Shampoo c) Coffee d) Motor oil 5. Name the following elements: 3 Au, S, Fe, Hg, W, Cu, At, K, Na, Pb, Zr, Mo, Ag, P, Ca, Cr, Ac, Ne, Fr, Sc, Ar, N, Mn, Be, Pt, Bi, Kr, C Hf, Th, Cs, Po, U, He, Y, Ir, In, Rn, Ce, Pu, Sb, Os, Assignment: a) p. 11 # 1 – 4 (copy question first) b) Find the number of protons, electrons, and neutrons for the following elements: Hf, Th, Pt, Po, Au, U, Bi, Cs, Pb, W Oxidizing material – rusting caused by oxygen (ex. iron) 4 Developing (History of) Atomic Theories 1. Dalton’s Atom (1766-1844) Dalton’s Atomic Theory (Pg 12) - All matter is made up of small particles called atoms - Atoms cannot be created, destroyed, or divided into smaller particles - All atoms of an element are identical in mass and size, but they are different in mass and size from the atoms of other elements - Compounds are formed when atoms of different elements combine in fixed proportions - Chemical reactions change the way atoms are grouped, but atoms themselves are not changed in reactions - “billiard ball” model http://www.rsc.org/chemsoc/timeline//pages/1803.html 2. J.J. Thomson (1856-1940) - English physicist - Atoms contain negatively charged electrons - Electrons are like raisins in a plum pudding or “raisin bun” model. 5 3. Ernest Rutherford (1871-1937) - Atom contains electrons and positively-charged particles - Atom composed of o A nucleus – a central region that is positively charged and contains most of the mass - protons are heavy positive particles within the nucleus o Electrons – particles with a negative charge and are very light (compared to protons). - Electrons circle around the nucleus o Empty space surrounding the nucleus is very large within which electrons move (planetary model). o Rutherford also proposed existence of the neutron to account for the mass difference between hydrogen and helium o Neutrons are heavy particles like protons but have no charge o Isotopes are atoms of the same element that differ in mass (but are chemically alike). (element with different number of neutrons) http://www.iun.edu/~cpanhd/C101webnotes/modern-atomic-theory/rutherford-model.html 4. Niels Bohr (1885-1962) - Electrons exist only in certain energy levels or orbits around the nucleus 6 - Only a certain number of electrons can exist in each energy level (orbit). 5. Modern Theory - Present day models of the atom are much more complex - Electron energy levels are divided into sublevels. - Neutrons and protons are made of even smaller particles called quarks. ATOMIC STRUCTURE Atom - the smallest part of an element (which retains the chemical and physical properties of the element). Atoms are made up of 3 sub-atomic particles 1. Electron (e- or e) -smallest particle in an atom -has a negative charge -located in the extra nuclear region of the atom 2. Proton (e+ or p) -Has a large mass -Has a positive charge -Located inside the nucleus 7 3. Neutron (n) -Same mass as a proton -Has a neutral charge (no charge) -Located inside the nucleus Nuclear Notation - Atomic number is the number of protons in the nucleus - The number of protons equals the number of electrons in a neutral atom (#p = #e) Atomic Mass Number is the total number of protons and neutrons in the nucleus - The mass number also identifies the particular isotope. - Number of neutrons = mass # - atomic # Atomic # = #p = #e Example: Find the number of protons, electrons and neutrons for iron and sodium. 8 Fe Atomic # = 26 Atomic mass = 55.85 p = 26 e = 26 n = 56 – 26 = 30 Note: when finding the number of neutrons we round the atomic mass to the nearest whole number. Na Atomic # = 11 Mass # = 22.99 p = 11 e = 11 n = 23 – 11 = 12 Assignment: Find the number of protons, electrons, and neutrons of the elements with atomic numbers 1 to 30 and 40-70. p. 11 #1-3 (copy questions or complete sentences) 9 Nuclear Notation Continued 1. One way to write isotopes of elements is: 12 6C where the top number is equal to the atomic mass, and the bottom number is equal to the atomic number. Atomic mass 12 6C Atomic number 13 6C #p = 6 #e = 6 #n = 12 – 6 = 6 #n = 13 – 6 = 7 2. Another notation used is: e.g. Lithium–7 or Li -7 where 7 is equal to the atomic mass. 10 From the table of elements we get the atomic number (which is 3). Thus, #p = 3 #e = 3 #n = 7 – 3 = 4 p.23 #5 to 8 p.24 # 1 to 4 p. 37 (b, c) sentences) (copy question first or complete p. 38 - define the key terms (first column) 11 Bohr’s Model of the Atom According to Bohr’s model electrons exist only in certain energy levels or orbits around the nucleus Only a certain number of electrons can exist in each energy level or orbit. The 1st orbit can hold a max. of The 2nd a max of 3rd 4th 5th 6th 2 electrons 8 8 18 18 32 When one orbit is filled the remaining electrons go to the next orbit – you cannot exceed the maximum allowed. We can draw the Bohr diagram for any element. It must have a nucleus showing the number of protons and neutrons and circles outside the nucleus showing the number of electrons. Reminder: # of protons = # of electrons = atomic # e.g. Draw the Bohr model for the following elements: 12 a) Lithium Step 1 – Look up the atomic number It’s 3. So, # of p = 3 # of e = 3 Step 2 – Look up the atomic mass. It’s 6.94 = 7 (round to the nearest whole #) Find the number of neutrons. Reminder: # of n = atomic mass – atomic # So, # of n = 7 – 3 = 4 Step 3 – Draw the diagram. #p=3 #e=3 #n=4 P=3 N=4 13 We can draw the orbits using a simplified version. e.g. Cobalt P = 27 N = 32 ___ 9 ___ 8 ___ 8 ___ 2 e Assignment: Draw the Bohr model of the atom for the elements: K, Cl, Mn, Zr, Ne, Ca, Sc, Ti, Fe, Al, Br, Ni, Cu, S, Ag, He, B, O, Na, Mg, Be, Ar, N, V, 1to10 and 15 to 30. And p. 38 #1, 5, 6 sentences) (copy question or complete 14 Periodic Table Periods are horizontal rows in which elements increase in atomic mass from left to right Groups or families are vertical rows made up of elements with similar properties. There are 4 special named groups. Group 1 – Alkali Metals Group 2 – Alkaline Earths Group 17 – Halogens Group 18 – Noble gases or Inert gases - ‘Staircase’ line separates metals from nonmetals - Metalloids border this line Francium is the most reactive metal. Fluorine is the most reactive non-metal. Valence Electrons - The outermost occupied energy level (orbit) of an atom is called its valence energy level. - The electrons in the valence energy level (electrons in the last orbit) are 15 called valence electrons. Electron dot diagrams or Lewis dot are useful ways to represent an atom. In an electron dot diagram, the electrons in the last orbit are shown as dots placed around the symbol. e.g. Li #p = 3 #n = 4 Bohr diagram e e There is 1 valence electron (that is, 1 electron in the last orbit). • Li So, the dot diagram will be e.g. Oxygen #p = 8 #e = 8 #n = 8 p=8 n=8 ____ 6 ____ 2e 16 There are 6 valence electrons. •• • O •• • Electron dot or Lewis dot diagram Assignment: Draw electron dot diagrams for the following elements: K, Cl, Mn, Zr, Ne, Ca, Sc, Ti, Fe, Al, Br, Ni, Cu, S, Ag, He, B, and p. 27 #9 to 12 p. 28 complete the table Science Test Friday Assignment: 17 Draw electron dot diagrams for: 1. Scandium 2. Fluorine 3. Beryllium 4. Vanadium 5. Gallium p. 38 – Define the key terms and p. 38 # 7, 8, 13 Quiz Draw electron dot diagrams for: 1. Sc 2. Na 3. Chromium 4. Ar Ions 18 - Ions are atoms, which have gained or lost electrons, in order to become more stable – it happens during chemical reactions. - Ions always have a charge - Positively charged ions have fewer electrons than protons – also called cations. - Most metals form cations – that means they lose electrons e.g. Li1+ 3 6.94 1+ Li loses an electron Li1+ Li Lithium - Negatively charged ions have more electrons than protons – also called anions. - Non-metals that form anions have a name ending in ‘ide’ e.g. chloride (Cl-), oxide (O-2 or O2-) 19 All non-metals gain electrons (that is, form anions). Compounds Compounds are formed when two or more elements are chemically combined. - Noble gases with their 8 valence electrons are very stable elements – they usually don’t form compounds. - Other atoms have different ways of becoming stable – they either gain or lose electrons when they form compounds. - **Metals give up electrons to other atoms, forming cations. - **Non-metals accept electrons, forming anions. - **Non-metals may share electrons with other atoms. e.g. Sulphur dioxide non-metal non-metal 20 Assignment p.36 #1-3, 5 p.38 #6, 9, 15,16 (copy question or complete sentences) Compounds There are two basic types: 1. Ionic 2. Molecular Ionic Compounds - Ionic compounds formed from just two elements are called binary ionic compounds - A metallic cation is joined to a non-metallic anion by an ionic bond. - Ions of an ionic compound are arranged in a regular repeating pattern called crystal lattice. Ionic compound – metal and non-metal joined chemically. 21 In ionic compounds electrons are traded. e.g. NaCl (p.30 – sketch Fig. 1.27 here) Molecular Compounds - Atoms which share electrons to become stable form molecular compounds (see p.32) - These groups of atoms are called molecules - Atoms in molecules are joined by covalent bonds. - All atoms in molecular compounds are nonmetals. Molecular compounds – non-metal and non-metal joined chemically. e.g. CO2 (p. 32 – sketch Fig. 1.29 here) Assignment: p.37 (i,j,k) p. 38 #13,15,16,17,18 (copy question) 22 Investigation 1-A Pg 33 Check Your Understanding Pg 36 Read Pg 37 Review Pg 38 23 Naming and Writing Binary Molecular Compounds - When two (binary) non-metallic atoms join by a covalent bond we have a molecular compound. e.g. Carbon dioxide Rules for naming 1. The first element in the compound is the one most left on the periodic table. 2. The suffix ‘ide’ is attached to the name of the second element. 3. Prefixes are used to indicate how many atoms of each type are present in one molecule of the compound. Prefixes: 1 = mono 2 = di 3 = tri 6 = hexa 7 = hepta 8 = octa MEMORIZE 24 4 = tetra 5 = penta 9 = nona 10 = deca No “mono” is used with the first element. e.g. Give the name or formula for each compound: NO2 – Nitrogen dioxide N2O – Dinitrogen monoxide N2O4 – Dinitrogen tetraoxide Nitrogen monoxide - NO Dinitrogen pentaoxide – N2O5 Carbon dioxide – CO2 Assignment: Name or give the formula: 1. Silicon dioxide 2. Sulphur monoxide 3. OF2 4. SiBr4 25 5. PH3 6. N2O 7. CO 8. NBr3 9. P2I3 10. SO3 11. N2O4 12. Tetraphosphorous hexaoxide 13. Dinitrogen tetraoxide 14. Heptasilicon monobromide 15. Octaboron decaiodide 16. B2O3 17. BrF7 18. N3O6 19. H2Cl5 20. Triselenium diastatide 21. Diarsenic pentaoxide 22. Sulphur trioxide 23. C3O2 24. C2H6 25. As3Br7 26. SO2 27. Selenium monoxide 28. Diboron trioxide 29. PF3 30. P2O5 31. P4O10 32. Arsenic trifluoride 26 33. 34. 35. BrF7 Hydrogen chloride N2O And p. 44 #1- 4, p. 62 #1 (copy question first) Binary Ionic Compounds - Are composed of ions of one metal element and ions of one non-metal element joined by ionic bonds Rules for naming 1. The first element in the name of the formula is the metal 27 2. The second element, the non-metal, is named as an ion. The suffix ‘ide’must be present. 3. No prefixes are used. e.g. Fe2O3 – Iron oxide CuS – Copper sulfide KCl – Potassium chloride p. 45 #5 - 7 p. 46 #9, 10 p. 47 #12 p. 48 #3, 5 (copy question) Writing Formulas for Binary Ionic Compounds In an ionic compound the total number of positive charges must equal the total negative charges – the compound must be electrically neutral. 28 This fact tells us how many of each atom is necessary to form a compound. e.g. sodium chloride Step 1 – use the table to find the charges on each ion (element) Na1+ Cl1- Step 2 – bring the two ions close together and see what the net charge is. Na+Cl- the two charges are equal so the formula is NaCl. Magnesium chloride Mg2+Cl1Question: how many of each ion is needed so that the molecule is neutral. Cl1- Mg2+ 29 Cl1- Therefore the formula is Chromium oxide 3+ Cr O MgCl2 Cr3+O2- 2- to balance the charges we use a shortcut method – charges are “traded” across. Cr2O3 Calcium oxide Ca2+O2- Ca2O2 CaO Multivalent Cations (metals) - Some atoms are able to form more then one cation. Ex. Ni2+ or Ni3+ - In the Stock system, the charge on the cation is written in brackets, as a Roman numeral after the name of the metal Example 30 Copper (II) oxide Cu2+O2- Tin (IV) fluoride Sn4+F1- SnF4 PbI2 CuO Lead (II) iodide Pb2+ I1- Cr2S3 Chromium (III) sulfide Cr3+S2Is this formula correct LiO Li1+O2- No – correct formula is Li2O p. 47 #11 p. 48 #4, 5 p. 49 #7, 9 31 Polyatomic Ions - Consist of two or more different atoms (covalently bonded) containing an overall charge. e.g. NO3- Found in the box at the top of the table. - All are negatively charged, except ammonium ion, and most names end in ‘ate’ - All act as non-metals except ammonium ion, NH4+, which acts as a metal in compounds. - The name of the cation (metal) is followed by the name of the anion (non-metal – negatively charged). - When writing formulas, brackets must surround the polyatomic ion (when more than one is present – i.e. subscript is not 1). Examples: 1. Potassium sulphate K2(SO4) K1+(SO4)2- “trade” charges or K2SO4 32 NH4NO3 Ammonium nitrate Al(NO3)3 Aluminum nitrate Sodium sulfate Na1+(SO4)2- Na2SO4 Na1+ SO42Na1+ Ammonium phosphate (NH4)1+(PO4)3- (NH4)3PO4 33 Gallium hydrogen carbonate Ga3+(HCO3)1- Ga(HCO3)3 Assignment: Practice Problems p. 52 #13-16 Practice Problems p. 53 #17-18 p. 55 #1-3 Properties of Ionic Compounds - In the solid state ionic compounds are crystalline Ionic compounds have fairly high melting points In the solid form they do not conduct electricity In the aqueous (dissolved in water) form ionic compounds are electrolytes – they conduct electricity 34 Properties of Molecular Compounds - Most molecular compounds have fairly low melting points – weak intermolecular bonds - Non-electrolytes – do not conduct electricity - When dissolved in water most do not conduct electricity (some do) p. 55 #4, 5 p. 80 #1 - 3, 6, 17,18, 20 (copy question) Special Compounds and Elements Special compounds – these compounds have special names, which do not follow the rules for naming. Water H2O Ozone Ammonia Hydrogen Peroxide Methanol Ethanol O3 NH3 H2O2 CH3OH C2H5OH or HOH 35 Sucrose Glucose Methane C12H22O11 C6H12O6 CH4 Diatomic and Polyatomic Elements If these elements are FREE, that is ALL ALONE, they are written as: H2 N2 O2 F2 Cl2 Br2 I2 At2 P4 S8 MEMORIZE For example, hydrogen has one electron and thus it wants to fill that orbit in order to become stable - so it will pair up with another hydrogen atom and they will share the two electrons – covalent bonding. H••H Thus, hydrogen when it’s not in a compound but all-alone is written as H2. 36 Assignment: p. 79 (a, b, c, e,) p. 80 – key terms p. 81 #25 Test tomorrow. Properties of Acids and Bases 1. Acids - a substance that reacts and releases hydrogen ions, H+(aq),in a water solution - taste sour - form colourless solutions - conducts electricity - formula starts with Hydrogen e.g. HCl H2SO4 - Hydrochloric acid Sulphuric acid 37 2. Bases - a substance that dissolves in water and releases hydroxide ions, OH- bitter tasting - feel slippery - form colourless solutions - conducts electricity e.g. NaOH - Sodium hydroxide Indicators and pH - an indicator is a chemical that changes a different colour in an acid vs a base - litmus is red in acids and blue in bases - phenolphthalein is colourless in acids but pink in bases - pH is a scale used to indicate the strength of the acid or base - pH scale ranges from 0 – 14, - pH of 7 is neutral – pure water - pH less than 7 – acid - pH greater than 7 – base 0 acid 7 base 14 38 p.80 #12,13 p.81 #23,25 (copy questions) Naming Acids All acids start with hydrogen. Acids have special names, which derive from the following rules. Chemical name Hydrogen _______ide becomes e.g. HCl Hydrogen chloride Acid name Hydro______ic acid Hydrochloric acid 39 H2S Hydrogen sulfide Hydrosulfuric acid Hydrogen _______ate _______ic acid H2SO4 Hydrogen sulfate HClO Hydrogen chlorate Hydrogen ______ite H2SO3 Hydrogen sulfite HClO2 Hydrogen chlorite Sulfuric acid Chloric acid _________ous acid Sulfurous acid Chlorous acid p. 70 #20 (a, b, c) #21 (a, b, c) #22,23 p. 71 #6 p. 79 (e, k) and p. 135 #26-28 40 Water - the shape of the water molecule is Oxygen end – slightly negative. 1050 Hydrogen end – slightly positive. - has two covalent bonds but the electrons shared in these bonds are not shared equally 41 - oxygen attracts the pairs of electrons closer to it - this creates an uneven distribution of charges or partial charges - the result is a polar molecule or dipole - the negative end or oxygen of one water attracts the positive end or hydrogen of another – hydrogen bond - hydrogen bonds are one kind of intermolecular force - intermolecular forces are attractions between molecules - intramolecular forces are attractions within molecules Properties of Water - the boiling point and melting points are higher in water than other similar substances – the need to break the hydrogen bonds - it requires a great deal of energy to raise the temperature of water – strong intermolecular forces - has a concave meniscus and shows capillary action – strong force of attraction between water and other molecules 42 - ice floats in liquid water – due to the rearrangement of the hydrogen bonds in the solid creating a greater volume and lower density - has a high surface tension – again due to the hydrogen bonds Chemical Reactions - chemical reactions occur when one or more substances change to form new substances - also called a chemical change - the substances that change are called the reactants - the substances formed are called products - evidence that a chemical change has occurred could involve one or more of the following o energy change – heat and/or light exothermic – release energy endothermic – absorb energy o odour change o colour change o formation of a gas – bubbling o formation of a solid – precipitate 43 Predicting Solubility - some ionic compounds are highly soluble in water while others have a very low solubility - we use a solubility table to help determine whether a substance is soluble or not – back of table. Step 1 – Locate one of the ions in the compound in the boxes across the top. Step 2 – Look for the other ion in the two vertical boxes below. If it is soluble write (aq) behind the compound to show that it is aqueous – it dissolves. If it is slightly (low) soluble show that it does not dissolve by writing (s) behind the compound so that it is solid. e.g. Determine if the following compounds are soluble or not by using the appropriate notation. NaCl(aq) Look for Na1+ or Cl1- across the top horizontal row. 44 PbI2(s) NH4OH(aq) CuSO4(aq) p.90 #1, 2 p. 93 #1- 4 p.128 # 9 Law of Conservation of Energy -Energy can be converted from one form to another, but the total energy of the universe remains constant (energy cannot be created nor destroyed). - Breaking chemical bonds is an endothermic process (energy is used). - Forming new chemical bonds is an exothermic process 1. When more energy is required to break bonds than is released when new bonds form, the reaction is endothermic e.g. energy + water hydrogen + oxygen 45 2. When less energy is required to break bonds than is released when new bonds form, the reaction is exothermic Example: hydrogen + oxygen → water + energy Lavoisier’s Law of Conservation of Mass During a chemical reaction, the total mass of the reacting substances (reactants) is always equal to the total mass of the resulting substances (products). Balanced Chemical Equations - a balanced chemical equation shows that atoms are conserved in a chemical reaction (that is, the numbers of each atom must be equal on both sides of the equation). - reactants are on the left side of the equation and products are on the right side 46 - coefficients are used to balance a chemical equation (tells us how many molecules or atoms are needed in the reaction). e.g. H2 + O2 → H2O (this is called a skeleton equation) 2H2 +1O2 → 2H2O H + H H H O → O (this is a balanced equation) O H H + O H H 47 http://funbasedlearning.com/chemistry/chemBalancer/ Assignment: Model Problem 1 p. 99 Model Problem 2 & 3 p. 100 p. 101 #5 p. 102 #5 (copy the EQUATIONS) Types of Chemical Reactions 1. Formation Reactions or Simple Composition - Two or more elements combine to form a new compound Element + Element → Compound X +Y → XY 48 The reactions must be balanced. e.g. Iron combines with oxygen to form iron (III) oxide. 4 Fe + 3 O2 → 2 Fe2O3 Copper reacts with chlorine to form copper (I) chloride. 2 Cu + 1Cl2 → 2 CuCl p. 114 #3 p. 115 #6 (Copy and balance) 2. Decomposition Reaction or Simple Decomposition One compound breaks down into two or more elements 49 compound → element + element + … XY → X +Y XYZ → X + Y + Z e.g. 2HCl → 1H2 + 1Cl2 2K2IO3 → 4K + 1I2 + 3O2 p. 127 (a-i,l,m) → complete sentences or copy the question. p. 128 #2, 4, 9 50 3. Single-Replacement Reactions one element takes the place of another element in a compound - many involve the reaction between a metal and a compound element + compound → new element + new compound A + BX → AX + B AX + Y AY + X → 51 Cu + 2AgNO3 → 2Ag + Cu(NO3)2 2NaBr + Cl2 → 2NaCl + Br2 Note: Metal replaces (switches with) a metal. Non-metal replaces a non-metal. Mg + CuSO4 → Cu + MgSO4 When a metal reacts with water, the water formula is written as HOH (first H “acts” as a metal). 2Na + 2HOH → H2 + 2NaOH OH1- (hydroxide ion) 52 4. Double-Replacement Reactions - Two different compounds react, forming two new compounds compound + compound → new + new compound compound AX + BY → BX + AY Note: metal switches with a metal and a non-metal with a non-metal. - a special kind of double-replacement reaction called neutralization is between an acid and a base NaOH + HCl → NaCl + HOH Base Acid Ba(OH)2 + Na2SO4 → 2NaOH + BaSO4 53 p. 114 #2,3,7 p.134 #21 p.136 #40 5. Hydrocarbon Combustion A hydrocarbon is an organic compound containing carbon and hydrogen (sometimes oxygen also) When hydrocarbons are burnt in a plentiful supply of oxygen complete combustion occurs - The two products are always carbon dioxide and water vapour Hydrocarbon + oxygen → carbon dioxide + water Hydrocarbon + O2 → CO2 + H2O 54 When hydrocarbons are burnt in a poor supply of oxygen incomplete combustion occurs. The products of this reaction are; carbon dioxide, water, carbon (soot) and carbon monoxide. Carbon monoxide is an odourless, colourless and highly toxic gas. - CO binds 200x more strongly to hemoglobin in the red blood cells than does O2 e.g. CH4 + 2O2 → CO2 + 2H2O 2C2H6 + 7O2 → 4CO2 + 6H2O p. 114 #1,4 p. 135 #27-29 Ammonia NH3 55 Ethane C2H6 Glucose C6H12O6 The Mole The mole is defined as the amount of substance that contains as many elementary entities (atoms, molecules, or formula units) as exactly 12 g of carbon-12, the most common isotope of carbon. One mole of a substance has been determined to contain 6.02 x 1023 elementary entities of a substance (atoms, molecules). This number is called Avogadro’s number. (Similar to dozen) dozen = 12 mole = 6.02 x 1023 Atomic Molar Mass 56 - is a weighted average of the mass of 1 mol of all of the naturally occurring isotopes of the element - listed for each element on the periodic table - example 1 mol of iron = 55.85 g/mol 1 mol of zinc = 65.39 g/mol - some elements exist as molecules such a nitrogen gas 1mol N2 = 2 x 14.01g/mol = 28.02g/mol Molar Mass of a Compound (M) - refers to the mass of 1 mol of any pure substance. - to find the molar mass of a compound use the chemical formula e.g. CO2 contains 1 carbon and 2 oxygen 1C = 1 x 12.01g/mol = 12.01 2 O = 2 x 16.00g/mol = 32.00 M = 44.01 g/mol 57 H2O Ca(OH)2 2 x 1.01 = 2.02 1 x 16.00 = 16.00 M = 18.02 g/mol 1 x 40.08 = 40.08 2 x 16.00 = 32.00 2 x 1.01 = 2.02 M = 74.10 g/mol p. 120 #9, 10 p. 123 #15, 16, 20 p. 125 #3 Calculate the molar mass of the following compounds: 1. PbI2 2. NH4OH 4. CaPO4 5. Mn(NO3)5 7. NH3 8. S2N4 10. C6H12O6 11. NH4HS 13. CoCl2 14. Cobalt(III) silicate 15. Potassium phosphate 16. Polonium (II) oxide 17. Mercury (II) sulfide 18. Fe2(OOCCOO)3 19. Zn(OH)2 20.Cu(NO2)2 21.Co2(Cr2O7)3 22. MgHPO4 3. CuSO4 6. Fe(OH)3 9. BaSO4 12. GaI3 58 Calculating mass of a sample (m) Molar mass (M) is equal to the mass of one mole of a compound. For example the molar mass of water is 18.02 g/mol. What if we have 2 moles of water? Then the mass of the water would be 2 x 18.02 = 36.04 g. We use the following formula: m = nM n = # of moles (mol) m = mass (g) M = molar mass (g/mol) How many grams are there in 3.5 moles of francium nitride? 59 Step 1 – Write the formula and find the molar mass. Fr3N 3 Fr – 3 x 223.00 = 669.00 1 N – 1 x 14.01 = 14.01 M = 683.01 g/mol Step 2 – List what’s given and apply the formula. n = 3.5 moles M = 683.01 g/mol m=? m = nM = (3.5)(683.01) m = 2390.54 g Mass of a substance to moles If the mass of the sample is given rearrange the formula for “n” n=m M e.g. 60 How many moles are there in a 16 g sample of carbon dioxide. CO2 1 x 12.01 = 12.01 2 x 16.00 = 32.00 M = 44.01 g/mol m = 16 g M = 44.01 g/mol n=? n = m/M = 16/44.01 n = 0.36 moles p. 122 #11-14 p. 123 #17,18 p. 125 #5,6 61 Moles summary 1. Molar mass (M) – must be calculated using the table. 2. Mass (m) – use the formula 3. Number of moles (n) – use m = nM n=m M p. 135 #29 - 32 p. 136 #41 - 43 62 Practice Problems Pg 122 Practice Problems Pg123 Check Your Understanding Pg 125 Read Pg 127 Chapter 3 Review Pg 128 Unit 1 Review Pg 134 TEST 63