Science Focus 10 Unit 1 Energy and Matter in Chemical Change pp 2-137 In this unit we will explore: Atoms, elements and compounds Ionic, molecular, acidic and basic compounds Naming chemical compounds The mole Water Chemical equations and reactions Chapter 1: Atoms, Elements and Compounds pp 2-39 1.1 Working with Chemicals p 6 Aboriginal peoples have been using chemical substances for thousands of years to preserve food, treat illness, build tools and decorate clothing. Many of these traditional processes are still used today. Many chemicals, even household products, have dangerous properties and must be handled properly. An MSDS, or Material Safety Data Sheet, lists important information including physical properties (ie melting and boiling points), chemical dangers, and how to store and dispose of the chemical. See figure 1.3 p7. The WHMIS or Workplace Hazardous Materials Information Systems informs users of the dangers of working with specific chemicals by informative labels on the chemical container, MSDS paperwork and worker training programs. See figure 1.4 p8 for examples of WHMIS symbols. Do BLM 1-1 (MSDS). Classifying Matter Matter is anything that has mass and occupies space. Mixtures can be mechanical (heterogeneous) where the separate parts are visible or solutions (homogeneous) where the different parts are not visible. Pure substances include elements (such as Na or Cl) or compounds (such as NaCl). Compounds can be chemically separated into simpler substances, elements cannot. See figure 1.5 p10. Do Practice Problems #1-4 p10 Do Check Your Understanding #1-3 p11 1.2 Developing Atomic Theories p12 The idea that matter is made up of small particles is rooted in work of scientists many centuries ago. One of the earliest theories is credited to the theorist Democritus (~400 BC), who stated that matter is made up of infinite tiny, indivisible, constantly moving units. This theory has evolved over recent history…. Historical Models of the Atom Model of atom Why model was proposed ● to account for the behaviour of substances when they interact with each other according to the law of conservation of mass (when substances react, the total mass of substance before and after are the same) Key features ● to account for the existence of small negative charges (cathode ray particles) associated with the atom gas discharge tube -see figure 1.9 p13 Thomson’s experiment – see figure 1.10 p15 ● All materials produce identical cathode ray particles, which are much smaller than even the smallest atom, hydrogen. ● All atoms contain electrons (small negative charges) that are distributed throughout a positively charged solid. Electrons are extremely tiny compared with the size of an atom. The Dalton Model “Billiard Ball” Dalton: 1766-1844 See figure 1.8 p13 The Thomson model “Raison Bun” Thomson: 1856-1940 See figure 1.11 p16 ● All matter is composed of small indivisible particles, called atoms, that can be neither created nor destroyed. ● All atoms of the same element are identical in mass and size, but different in mass and size from atoms of other elements. ● Atoms exist in an otherwise empty space and are in constant motion, during which they may collide to form new combinations (compounds). ● Chemical reactions change the way in which atoms are grouped, but they do not change the atoms themselves. Why model was rejected or modified ● contained inaccuracies regarding the relative masses of several atoms ● could not account for the electric nature of matter ● could not account for the production of high-frequency electromagnetic radiation ● could not explain the existence of alpha particles (heavy, positive charge) for radioactive materials Model of atom The Rutherford model “Solar System” Rutherford: 1871-1937 See figure 1.14 p18 The Bohr Model Bohr: 1885-1962 See figure 1.19 p20 Why model was proposed ● to account for the deflection and reflection of alpha particles from gold foil Rutherford’s experiment – see figure 1.12 p17 Explanation of the experiment – see figure 1.13 p17 Isotopes – see figure 1.15 p19 ● to account for the production of bright-line spectra in hot gases and absorption spectra in cold gases See figures 1.17, 1.18 p20 Key features ● At the centre of every atom, there is a small, positively charged nucleus. ● The nucleus accounts for the majority of the mass of the atom. ● Electrons are attracted to the nucleus and orbit in a cloud around the nucleus. A third subatomic particle, neutron, with no charge but a mass similar to a proton exists in the nucleus ● Electrons reside within stable energy levels (electron shells) at fixed radii relative to the nucleus of the atom. Transitions of electrons to higher energy levels require energy. Transitions to lower energy levels produce electromagnetic radiation (light and radio waves). ● Each energy level has a fixed maximum number of electrons that can reside in it. Why model was rejected or modified ● could not account for the emission spectra of elements ● included orbiting electrons that should constantly radiate energy and therefore collapse into the nucleus, but the radiation and collapse were not observed ● could not explain the existence of sublevels within higher energy levels, predicted by the shape of the periodic table ● could not fully explain the nature of electrons or the nucleus QUARKS Scientists now believe that neutrons and protons are made up of even smaller particles called quarks. They believe that matter is made up of dozens more of these subatomic particles. Features of a Simplified Modern Model of the Atom (figure 1.20 p22) a tiny, dense nucleus that is surrounded by electrons (e-) nucleus contains protons (p+) and neutrons (no), called nucleons (exception: H-1 nucleus contains one proton only) nucleus accounts for most of the mass of the atom e- exist at certain allowed energy levels p+ carry a positive charge, e- carry a negative charge and no carry no charge. See table 1.2 p 22. A neutral atom always has equal numbers of e- and p+ If a nucleus were the size of a baseball, the electron orbitals around the nucleus would take up the space of an arena! Nuclear Notation Elements can have 2 or more isotopes, where the number of neutrons is different. Scientists use the following notations to describe a specific isotope: 29 14 Si Silicon-29 or top number represents the mass = total number of protons + neutrons bottom represents the atomic number = total number of protons in the nucleus These two numbers can be used to determine the number of neutrons: mass number – atomic number = # neutrons so with Silicon-29, the number of neutrons is 29-14=15! Do Practice Problems #5-8 p23 + #4 p24 Do Check Your Understanding #2, 3, 5 p24 1.3 Electrons and the Formation of Compounds p25 During the 1800’s, a Russian chemist, Dmetri Mendeleev, examined 62 elements. He developed a table of these elements based upon the fact that they had repeating properties. The table was called a “periodic table” for this reason. He further predicted the existence and properties of unknown elements and left spaces on his periodic table for them. Examine our modern periodic table on p482. It displays the known elements in a format that follows various patterns and trends: Elements are arranged according to increasing atomic number (number of protons in the nucleus) Period – horizontal row Inner-transitional elements top period = lanthanoids (fits in period 6) bottom period = actinoids (fits in period 7) Group or family – vertical column, each element in a group has similar properties Group 1 = alkali metals (react with water) Group 2 = alkali earth metals (react with oxygen) Group 17 = halogens (very reactive) Group 18 = noble or inert gases (not reactive) Staircase line –bordered by metalloids, this line divides metals on the left from non-metals found on the right. See table 1.3 p25 to examine the properties of metals, non-metals and metalloids. Patterns of electron arrangement in periods & groups Recall that Neils Bohr inferred that electrons orbit the nucleus of the atom in fixed energy levels, and each level can only hold a certain maximum number of electrons. The first can hold 2 electrons, the second can hold 8, and the third can hold 8. We can use the patterns in periodic table to predict the number of energy levels of an elements and the number of electrons in each energy level. Examine figure 1.22B p26. period number shows the number of energy levels used in each element (eg. period 2 elements have 2 energy levels) group number describes how many electrons are found in the valence or outermost energy level (eg. lithium is in group 1 and has 1 valence electron) for groups 13 – 18, we use the last number to designate the number of valence electrons (eg. elements in group 16 have 6 valence electrons) electrons fill the first orbital before they can occupy the second, and fill the second before they can occupy the third when the valence level is full, it is referred to as a stable octet since there are 8 electrons occupying the orbital (unless it is the first level) The diagram representing the element beryllium looks like this: __2e-__ Try the diagram for fluorine: __2e-__ 4p+ Do Practice Problems #9-12 p27 and BLM 1-4. Electron Dot Diagrams, a.k.a. Lewis Dot Diagrams American chemist G.N. Lewis invented these structures to visualize and track electrons during bond formation. To draw: 1) write its chemical symbol 2) surround it by the number of dots that represent the atom’s valence electrons 3) if an atom has more than 4 valence electrons, the additional electrons are shown paired with others 4) note that elements in the same group will have identical dot diagrams since they have the same number of valence eDo BLM 1-6 and then BLM 1-5. Formation of Ions Any atom or group of atoms that either loses or gains eand thus carries either a positive or negative charge is called an ion. Cation - positive charge, has fewer e- than p+ - metallic atoms lose electrons, eg. Na+ - remember: cats have paws (pos) Anion - negative charge, has more e- than p+ - non-metallic atoms sometimes gain electrons from other atoms, eg. Cl- but they can also share electrons See p29 figure 1.24 and 1.25 for the 3 ways to represent the formation of ions. Formation of Ionic Compounds Ions do not form by themselves. Instead, as metallic and non-metallic atoms collide with one another, their valence electrons interact. The metal will lose its valence electrons (the cation) and an adjacent non-metal will gain those electrons (becoming the anion); this is a transfer of electrons. The two ions formed are opposite in charge and are greatly attracted to each other, forming a very strong ionic bond. The rearrangement of electrons allows each ion a full valence orbital (like its nearest noble gas) and therefore leads to a greater stability and lower energy level for the ionic compound. Important notes about ionic compounds: ions arrange themselves in a regular repeating pattern called a crystal lattice (see table salt + figure 1.28, p31) a binary ionic compound is formed from only 2 elements. Why does Na become Na1+ and Mg become Mg2+? Sodium (Na) has only 1 valence electron, so when it is lost, the resulting charge is 1+, however magnesium (Mg) has 2 valence electrons so its ion has a charge of 2+ Do BLM 1-7 to practice isotope and ion formation Formation of Molecular Compounds When non-metals react with one another, electrons are NOT transferred since both atoms tend to gain electrons to fill their valence energy level. Instead, the valence electrons are shared. These compounds are referred to as molecular and the sharing of electrons forms a bond called covalent (co = share, valent = valence shell) Important notes about molecular compounds: a diatomic molecule is composed of only 2 atoms and a polyatomic molecule is composed of many atoms (see and memorize table 1.5 p 31) molecules do NOT form a crystal lattice shape It is important to note that the physical and chemical properties of a compound is different from those of the individual elements that make up the compound. Do Investigation 1-A p 33 + Questions 1-6 p 35 Do Check Your Understanding #1,3,6,7 p36 Do Chapter 1 Review #2,4,5,7,10,11,12,15,16,17,19,20 p38 Chapter 2: Names, Formulas and Properties pp 40-81 2.1 Chemical Names and Formulas p 42 The IUPAC (International Union of Pure and Applied Chemistry) was founded in 1919 and has developed a method to name chemicals based on their composition. The systematic name of a compound allows you to write the chemical’s formula and predict some of its properties. Binary Molecular Compounds Naming Rules: see table 2.2 p 44 1st name: use the name of the element found furthest to the left on the periodic table 2nd name: add the suffix “ide” to the name of the second element (eg. oxygen becomes oxide) **attach a prefix to each name to indicate how many of each atom is present in one molecule of the compound (see and memorize the prefixes in table 2.1 p 43) **note that the prefix “mono” is not generally used for the first element **when hydrogen is the first element in the formula, the compound is molecular BUT prefixes are NOT used (see p54 – Hydrogen Compounds). More later. **memorize: diatomic elements (N2, O2, H2, F2, Cl2, Br2, I2, At2,) + P4 + S8 **memorize: O3 – ozone NH3 – ammonia CH4 – methane C6H12O6 – glucose CH3OH – methanol C12H22O11 - sucrose Do Practice Problems p44 C2H5OH – ethanol H2O2 – hydrogen peroxide Binary Ionic Compounds Naming Rules: 1st name: 2nd name: use the name of the metal element add the suffix “ide” to the name of the non-metal element **the chemical formula is the simplest whole ratio of each type of ion in the compound and it represents the smallest repeating unit (formula unit) within a crystal lattice (fig 1.28 p31) **prefixes are NOT used in naming ionic compounds Study Model Problem 1 and figure 2.3 p45 Do Practice Problems p45 Stock System: As you look at your periodic table, you may notice that many transitional elements have more than one possible ionic charge. Study Model Problem 2 and do Practice Problems #9,10 p46 A German chemist, Alfred Stock, developed a way to indicate which cation is in a compound. The Stock System uses a roman numeral to indicate the charge on the metal or cation. Study Model Problem 3 and do Practice Problems #11,12 p47 Do BLM 2-1 (and BLM 2-5 to practice the crossover method) Polyatomic Ions These ions consist of two or more different atoms joined together by covalent bonds. As a group, the bonded atoms have either a positive or negative charge. See the polyatomic ion table on the back of the periodic table, and become familiar with the common polyatomic ions in table 2.3 p51. Naming Rules: 1st name: name the cation 2nd name: name the anion **parentheses must be placed around the polyatomic ion if more than one are found in a chemical formula Study Model Problems 4,5 p51 and do Practice Problems p52 Getting to know the polyatomic ions: Many polyatomic anions contain oxygen and those that contain the same elements are part of a “family”. Anion name Per----ate ----ate ----ite Hypo----ite # of oxygen 1 more --------------1 less 2 less Do Practice Problems 17-19 p53 Example pernitrate nitrate nitrite hyponitrite Formula NO4NO3NO2NO1- Do BLM 2-3 (and Check Your Understanding p55 for practice) 2.2 Explaining Properties of Substances p56 Physical properties – observable and measurable qualities of a substance, such as melting and boiling points, conductivity, appearance, state (see table 2.7 p56) Chemical properties – reactive properties of a substance Property Ionic compound State at room temp. Melting point Solid (hard, brittle, Solid, liquid or gas crystal lattice shape) High Low Attraction between molecules Strong because of opposing charges Weak Conductivity when solid? No, molecules held rigidly in crystal lattice shape Yes, crystal lattice has broken down = electrolyte No Conductivity when dissolved in water or melted? Lab safety demo Read together and sign laboratory contract Review lab write-up guidelines Molecular compound Sometimes, but these compounds tend to be non-electrolytes Do Investigation 2-B p58 (use BLM 2-2) and Q #1-6 p59 Toxic Properties There are many chemicals found in society today with properties that produce pronounced physical and psychological effects. Some of these chemicals are regulated or even banned by government and others are not. Understanding chemistry can help people make informed decisions about personal use of such chemicals. Examine table 2.8 p 60 to familiarize yourself with effects of the chemicals ethanol and nicotine. Do Check Your Understanding #1,3,5 p62 2.3 Properties of Acids and Bases p 63 Both acids and bases have several properties in common. These chemicals can be both dangerous and useful. Acidic and basic solutions are electrolytes, so they contain freely moving ions and thus conduct electricity. See figure 2.7 p63 for examples of household acids and bases. Property Arrhenius definition (a Swedish chemist) Formula Electrolyte? Taste Feel Uses ACID releases H+ (hydrogen ions ) in a water solution HB(aq) ->H+(aq) + B-(aq) Yes Sour (like lemons) Will burn skin removes rust refine metal Less than 7 pH Indicators: Litmus Red Phenolphthalein Colorless BASE releases OH(hydroxide ions) in a water solution MOH(aq) ->M+(aq) + OH-(aq) Yes Bitter (like tonic water) Will burn skin, but bases feel slippery (soap) cleaning products medical drugs More than 7 Blue Pink For important acids and bases and their uses see tables 2.9 and 2.10 p64. Notes: pH indicators are made of chemicals that changes color in response to the concentration of hydrogen or hydroxide ions. pH meters and probes can also be used. These instruments measure the electrical property of the solutions, which relates to the solution’s pH. They are more precise but more expensive. When acids and bases react, they neutralize each other, or they lose their acid and base properties. The resulting solution is a salt dissolved into water. Strong acids react with metals to produce hydrogen gas. Strong acid = high conductivity, weak acid = weak conductivity Naming Acids: Name the compound as though it were an ionic compound. Use the following rules to complete the acidic name: 1. 2. 3. If the ionic name was: hydrogen _______ide The acidic name is: hydro_____ic acid eg. H2S = hydrogen sulfide = hydrosulfuric acid hydrogen _______ate ______ic acid eg. HClO3 = hydrogen chlorate = chloric acid hydrogen _______ ite ______ ous acid eg. H2SO3 = hydrogen sulfite = sulfurous acid Notes: Acids are aqueous solutions so the proper way to write the formula of an acid is with an (aq) subscript. H2S, hydrosulfuric acid, can also be called aqueous hydrogen sulfide. When trying to figure out the formula of an acid from its name, remember that you must have enough hydrogen ions to balance the negative charge of the anion. If the polyatomic ion that is part of the acid ends in a ---COO-, the hydrogen ion bonds at this end to become ---COOH (this is called a carboxylic acid). Do Practice Problems p70 Do Find Out Activity p 67 and BLM 2-6 Do Check Your Understanding #3,4,5,6,8 p71 2.4 Why Water is Weird p 72 Water is essential to life. The cells of most living organisms are 80-95% water, and virtually all chemical reactions for life occur in aqueous solutions. But water also has many other unique properties: Occurs in all 3 states on Earth Ice is less dense than liquid Its melting and boiling points are higher than similar substances (see table 2.14 p72) It has a very high surface tension Do Investigation 2-D p73 (use BLM 2-9 to record results) and do Q #1-3,6 p74 A Molecular View The above characteristics of water can be explained by several features of water. 1. Shape – bent (see figure 2.13 p75) 2. O-H bonds – each bond is made up of 2 electrons that are more strongly attracted by the oxygen atom so that it gains a slightly negative charge as the hydrogen gains a slightly positive charge, making it a polar or dipole molecule. These properties encourage individual water molecules to attract each other. The negative oxygen attracts the positive hydrogen and forms a hydrogen bond. This kind of bond is an intermolecular bond, one between molecules (versus intramolecular bonds, within molecules, like covalent bonds which are stronger). The hydrogen bond explains many of water’s unique properties: its melting and boiling points are high -more energy has to be added to water in order to break these intermolecular bonds and allow ice to melt or water to vaporize it takes much energy just to raise the temperature of water -the strong attraction between water molecules must be overcome to increase the average speed of the water molecules and thus heat the water it has a concave meniscus and shows capillary action -its polar molecules are attracted to the sides of the container when it is solid water, it floats in liquid water -hydrogen bonds force water molecules farther apart as it freezes and leaves the ice less dense than liquid water it has a high surface tension / cohesion -hydrogen bonds pull water molecules close together Do Check Your Understanding #1,2,4,6 p78 Do Chapter 2 Review #1,2,4,6,8,9,10,13,14,17,18,19,22,25 p80 Chapter 3: Chemical Reactions pp 82137 3.1 Recognizing and Describing Chemical Reactions p 84 A chemical reaction, or chemical change, has taken place when one or more substances have changed to form different substances. The reactants are those substances that undergo a chemical change and the products are the new substances formed. Evidence of a Chemical Reaction: (see table 3.1 p84) Odor change Color change Formation of a gas (bubbles) Formation of a solid (called a precipitate) Energy change such as: Temperature increase or decrease Light produced Sound produced Electricity produced *energy changes can also accompany physical changes, which are simply changes in state Do BLM 3-1 Solubility and Chemical Reactions Recall that when most ionic compounds are dissolved in water, the cations and anions separate from each other and move about freely. Also remember that some ionic compounds do not dissolve well in water. If two ionic compounds are mixed, their free cations and anions may react with each other to produce new substances. If the new compound does not dissolve in water, a precipitate is formed, and this indicated a reaction has occurred. Do Investigation 3-A p86 (use BLM 3-3 to record) and do Q#1-6 p87 Predicting Solubility Table 3.2 p88 outlines which ionic compounds are soluble in water and which have a low solubility. To determine whether a compound is soluble in water or not, follow these steps: 1. locate one of the ions found in the compound on the top row of the chart 2. look down that column of the table to find the second ion 3. if the 2nd ion is found in the “high solubility” category, the compound is soluble in water and should be given a subscript of (aq) (aqueous) OR If the 2nd ion is found in the “low solubility” category, the compound is NOT soluble in water and should be given the subscript (s) (solid) It will form a precipitate in water. Study Model Problem 1 p 89 Do Practice Problems 1-4 p90 and BLM 3-2 Chemical Reactions and Energy Changes Energy changes accompany any chemical reaction. Endothermic – energy is added to start the reaction Eg. Photosynthesis – energy from the sun is required to start the process of converting carbon dioxide and water into food (glucose) and oxygen in plants Exothermic – release of energy during the reaction Eg. Cellular respiration – the opposite reaction of photosynthesis, where glucose and oxygen are converted into carbon dioxide and useful energy in both plants and animals Do BLM 3-4 for reinforcement Law of Conservation of Energy This law states that energy can be converted from one form into another, but the total energy of the universe remains constant. Energy cannot be created or destroyed. Energy is required to break chemical bonds, and energy is released when chemical bonds are formed. Endothermic reaction – more energy is required to break the bonds of the reactants than is released when the products’ bonds are formed Exothermic reaction – less energy is required to break the bonds of the reactant than is released when product bonds are formed See figure 3.8 p 92. Do Check Your Understanding #1,3,4,5 p93 3.2 Representing Chemical Reactions p94 Closed system – an exchange of energy between the system and its surroundings occurs, but NOT an exchange of matter Example – a terrarium - the earth? Open system – an exchange of both energy and matter between a system and its surroundings Example – a flower pot Isolated system – exchange of neither energy nor matter between a system and its surroundings Example (theoretical) – a thermos Law of Conservation of Mass Developed by Antoine Lavoisier (1743-1794) This law states that during a chemical reaction, the total mass of the reacting substances (reactants) is always equal to the total mass of the resulting substances (products). Do Investigation 3-B p 95 and Q# 1-10 p96 Writing Balanced Chemical Equations Reaction Description: hydrogen gas reacts with oxygen gas to produce liquid water Word equation: Skeleton equation: hydrogen + oxygen -> water H2(g) + O2(g) -> H2O(l) Balanced equation: 2H2(g) + O2(g) ->2H2O(l) *shows atoms are conserved *coefficients (numbers in front of chemical formulas) show how many of each compound is there (note there is 1 molecule of oxygen gas, but 2 atoms of oxygen) Q: how many atoms of hydrogen are reacting? ______ *subscripts (letters below and to the right of each compound) show what state each compound is in How the balance equations: 1. When writing the correct skeleton equation, check your formulas to ensure they are correct. Include all states of matter for each reactant and product. 2. Balance the atom or polyatomic ion present in the greatest number. Find the lowest common multiple to obtain whole number coefficients to balance. 3. Repeat step 2 to balance each remaining atom/ion. HINTS: O2(g) – when O2(g) is present in the reaction, balance it last When single elements are present on one side of the equations, balance those elements last (ie. Na, Mg) As long as complex ions are found intact on both sides of the equation, treat the complex ion as one group 4. Check once again to make sure the equation is balanced by creating a table of atoms on each side of the equation. Study Model Problem 1-3 p99-100 Do Practice Problems 5-8 p101 Do Check Your Understanding #2-5 p102 Do Worksheets p2 (left side) 3.3 Types of Chemical Reactions p103 Do Find Out Activity p103 using BLM 3-5 to record. Do the appropriate sections of BLM 3-6 as equations are learned 1. Formation Reaction: two or more reactants combine to form a new product Eg. X + Y -> XY See figure 3.15 and 3.16 p104. 2. Decomposition Reaction: one compound breaks down into two or more simpler compounds/elements Eg. XY -> X + Y See figure 3.17 and 3.18 p 104. 3. Single-Replacement Reaction: one element takes the place of another element in a compound Eg.1. A + BX -> AX + B Eg.2. AX + Y -> AY + X *many of these reactions are between a metal and a compound (see p105) *sometimes the reaction is between a halogen and a halogencontaining compound (see p105) 4. Double-Replacement Reaction: the cations of two different compounds exchange places, forming two new compounds Eg. WX + YZ -> WZ + YX *many of these equations result in a precipitate forming or the formation of water (p105) 5. Reactions involving Carbon Compounds: The study of carbon-containing compounds is called organic chemistry. A hydrocarbon is a compound that contains only hydrogen and carbon atoms (eg. C2H6). These compounds are retrieved by refining crude oil and natural gas. Approximately 95% of these hydrocarbons are burned as fuels in exothermic combustion reactions to create thermal energy to warm buildings and provide energy for transportation. Complete Combustion: a hydrocarbon reacts with oxygen gas (or burns) to create carbon dioxide gas, water vapor and thermal energy. See figure 3.20 and 3.21 p108. Incomplete Combustion: when oxygen is in poor supply, the products are carbon dioxide gas, water vapor, carbon monoxide, carbon (soot) and less thermal energy than a complete combustion reaction. *carbon monoxide is a colorless, odorless, highly toxic gas that, when breathed in, strongly binds to red blood cells instead of oxygen and can lead to death Study the Tools of Science p109 and learn the chemical tests for hydrogen, oxygen, carbon dioxide and water Do Check Your Understanding #1-9 p114 Do the rest of your Worksheet 3.4 The Mole p116 Since atoms and many compounds are very tiny, scientists group them into an extremely large number called a mole. The mole: A mole is defined as the amount of substance that contains as many elementary entities (atoms or compounds) as exactly 12 g of carbon-12 (the most common isotope). One mole of a substance contains 6.022 141 99x1023 particles of that substance 6.02 x 1023 is called Avogadro’s number (see Did You Know p116) The unit mol is short for the German word “Molekulargewicht”, which is literally translated to, you guessed it, molecular weight! Atomic Molar Mass (M) this is a weighted average of 1 mol of all the naturally occurring isotopes of that element it is listed on the periodic table units are in g/mol Molar Mass of a Compound (M) this is the mass of 1 mol of any pure substance units are in g/mol it is determined by using the formula for that compound (see eg. p120) Do Practice Problems p 120 Do WS p.2,4 Review significant digits worksheet Converting Between Mass and Moles Scientists often convert between the mass (g) of a sample and the number of moles (mol). The following formula is used: n = _m_ M n = number of moles of substance m = mass of substance M = molar mass of substance Study Model Problem 1 and 2 p121 and 122 Do Practice Problems p 122 and 123 Do BLM 3-10 The Mole and the Law of Conservation of Mass Study the chemical reaction at the top of p124. The coefficients explain the number of molecules and/or the number of moles of each substance needed for the reaction to run. Now view table 3.5 at the bottom of the page. Notice how molecules, moles and mass are all related in the chemical equation. Also notice that the mass of the reactants is always equal to the mass of the products. Do Do Do Do BLM 3-10 + rest of Worksheet Check Your Understanding #2-6 p125 Chapter 3 Review #1-5,7,9-10,12(a-b),13(a-c),14-19 Unit 1 Review #1-38,39-43(a-c for each),46