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Energy and Matter in Chemical Change What is Chemistry? The study of matter; its properties, composition and the changes it undergoes. Why Chemistry is Cool It has the best jokes… Only the coolest people do chemistry… http://www.youtube.com/watch?v=bYtQ3_eh 094&feature=related A Walk Down WHMIS Compressed Gas Toxic Biohazardous Flammable Poisonous Oxidizing Corrosive Radioactive HHPS Danger H azardous H ousehold P roduct S ymbols Warning Caution MSDS M aterial S afety D ata S heets Science Lingo… • Scientists use an experiment to search for cause and effect relationships in nature. In other words, they design an experiment so that changes to one item cause something else to vary in a predictable way. • These changing quantities are called variables. A variable is any factor, trait, or condition that can exist in differing amounts or types. An experiment usually has three kinds of variables: independent, dependent, and controlled. • The independent (manipulated) variable is the one that is changed by the scientist. To insure a fair test, a good experiment has only one independent variable. As the scientist changes the independent variable, he or she observes what happens. • The scientist focuses his or her observations on the dependent (responding) variable to see how it responds to the change made to the independent variable. Let’s Look At An Example… • For example, if you open a faucet (the independent variable), the quantity of water flowing (dependent variable) changes in response--you observe that the water flow increases. The number of dependent variables in an experiment varies, but there is often more than one. • Controlled variables are quantities remain constant. For example, if we want to measure how much water flow increases when we open a faucet, it is important to make sure that the water pressure (the controlled variable) is held constant. That's because both the water pressure and the opening of a faucet have an impact on how much water flows. If we change both of them at the same time, we can't be sure how much of the change in water flow is because of the faucet opening and how much because of the water pressure. In other words, it would not be a fair test. Most experiments have more than one controlled variable. Some people refer to controlled variables as "constant variables." An Example… • A scientist is conducting an experiment to test if taking vitamin A could extend a person’s lifeexpectancy. Another Example… • A scientist wants to see if temperature affects the reaction rate of melting M&M’s. She sets up three beakers with different temperatures of water; hot, luke warm and ice cold. Summary • Independent (Manipulated) Variable: What we are changing • Dependent (Responding) Variable: The Response (what happens) • Controlled Variable: What remains the same • A hypothesis is an educated guess about how things work. • Your hypothesis should be something that you can actually test, what's called a testable hypothesis. • In other words, you need to be able to measure both "what you do" and "what will happen." Scientific Method • What is the scientific method? – It is a process that is used to find answers to questions about the world around us • Is there only one “scientific method”? – No there are several versions of the scientific method. Some versions have more steps, while others may have only a few. However, they all begin with the identification of a problem or a question to be answered based on observations of the world around us and provide an organized method for conducting and analyzing an experiment. Hypothesis • What is a hypothesis? – It is an educated guess based on observations and your knowledge of the topic. • What is data? Data – It is information gathered during an experiment Identify Form a Hypothesis Create an Experiment Perform an Experiment Analyze Data Is data accurate? Modify the Experiment Communicate the Results Scientists Aristotle • 400 BC • Fire, Earth, Water and Air Democritus • Believed that material was made up of small individual particles called “atoms” Alchemists Lavoisier • “Father of Modern Chemistry” • Came up with the Law of Conservation of Mass- “Mass is neither created nor destroyed” The other person who did a lot of significant scientific work but did not get credit because of her gender Dalton • In 1804 Dalton stated that atoms were – Tiny individual particles – All identical in a given element with identical properties – Two or more elements atoms will combine in fixed ratios – “Billiard Ball Model” Dalton’s Matter Model • All matter is made of atoms • All atoms of an element are identical in properties • Atoms of different elements can combine in specific ratios to form new substances Law of Definite Proportions • A compound always has the same ratio of atoms, regardless of how it is made. Example: water is always H2O, not HO, or HO2, or H2O2 30 J.J. Thompson • Discovered that atoms had negative particles (electron) contained inside the atom • “Raisin Bun Model” Nagoaka • Placed all the + charges in the centre of the atom and all the – charges in a ring around the centre. • “Saturn Ring” model Rutherford • Believed that electrons were outside the dense positive part of the atom (nucleus) • Used gold foil experiment Chadwick • Discovered the subatomic particles: – Protons: p+ – Electrons: e– Neutrons: n° Bohr • Thought that electrons were arranged in certain energy levels around a positive nucleus • Electrons in orbitals • “Bohr Model” de Broglie and Schrodinger • Electron cloud model • Electrons are not considered to be point charges but as a mist around the nucleus • Theory has evolved into the quantum mechanics model http://www.youtube.com/watch?v=7SjFJImg2Z8 The Latest and Greatest- Quantum Mechanics • The latest experimental evidence has disproved Bohr’s idea of fixed energy levels • Now, they visualize electrons not as a particle, but as a cloud of negative charge. Rather than following little race tracks around the nucleus, they occupy the whole space all at once at different energy levels • NOW GET YOUR HEAD AROUND THAT! Modern Atomic Structure • Current models are more complex and involve many subatomic particles such as quarks (protons and neutrons) 38 Your Task • Check & Reflect p 25 # 1,5-8 & 11 • Section Review page 27 # 5, 10, 11, 12, 15, 16 • Expect a scientist quiz next class Matter • Matter: Anything that has mass and takes up space. States of Matter • There are four states of matter: – Solid – Liquid – Gas – Plasma Particle Model of Matter 5 Main Points: 1) All substances are made of matter 2) All particles in a pure substance are the same. Different pure substances are made of different particles. 3) Particles are always in motion. The speed of the particles increases when temperature increases. 4) Particles have space between them 5) Particles may have attractive forces between them Organizing Matter Pure Substances • Made up of only one kind of matter and has a unique set of properties. • Can only be broken down by a chemical decomposition Elements • pure substance that cannot be separated into simpler substance by physical or chemical means. • Basic building blocks for all compounds • Each element has it’s own symbol Compounds Pure substance composed of two or more different elements joined by chemical bonds. – Made of elements in a specific ratio that is always the same – Has a chemical formula – Can only be separated by chemical means, not physically Mixtures • A combination of two or more pure substances that are not chemically combined. • substances held together by physical forces, not chemical • No chemical change takes place • Each item retains its properties in the mixture • They can be separated physically Chem4kids.com Heterogeneous Mixtures • Heterogeneous mixtures are also known as Mechanical Mixtures • In a mechanical mixture, the different substances that make up the mixture are visible. • Includes: Mechanical mixtures, suspensions, emulsions and colloids Homogeneous Mixtures • A Homogeneous Mixture is also known as a Solution • In a solution, the different substances that make it up are not separately visible. • One substance is dissolved in another. • Recall that a Solution is made up of a Solvent and a Solute Suspensions • A suspension is a cloudy mixture where tiny particles of one substance are held within another. • These particles can be separated out when the mixture is poured through filter paper. Colloids • A colloid is also a cloudy mixture, but the particles of the suspended substance are to small that they cannot be easily separated out from the substance. Test Your Understanding • A black solid with a constant melting point is heated to a high temperature, producing a gas and a shiny brown metal. The boiling point of the gas is -183°C and the melting point of the metal is 1085°C. Is the black solid an element, compound or mixture? Explain • Let’s break down the evidence: – Constant melting point (not a mixture) – But both a gas and a solid were formed. So it is not an element Compound Separation of Mixtures • Separating Mixtures to identify the ingredients can be accomplished through the following procedures. – Mechanical Mixtures: sifting or filtering – Solutions: distillation – Suspensions: filtration and centrifuging – Colloidal Solutions: centrifuging – Compounds: chromatography Phase Changes Physical vs. Chemical Change Physical Chemical http://cwx.prenhall.com/petrucci/chapter1/medialib/tutor/f20/0103.html Chemical Change • A chemical change always results in the formation of a different substance or substances. Physical Properties 2 Categories of Physical Properties: 1) Qualitative Properties ( sensed/ observed) • • • • • • • • Color Texture Smell Taste Malleability Texture Change in state Ductility Physical Properties 2) Quantitative Properties (measured) • • • • • • Melting point (mp) Boiling Point (bp) Density d= m/v (mass/volume) Solubility Viscosity (flow rate) Conductivity Chemical Reactions • Characteristics – Energy change (temperature, light, sound, electricity) – Odor Change (appearance or disappearance) – Color change – Formation of a gas – Formation of a solid (precipitate) http://www.youtube.com/watch?v=66kuhJkQCVM Test Your Understanding • A blue crystal is placed in water and after stirring it disappears and the water becomes blue. The liquid is then heated and the water evaporates and small blue crystals appear. Did a chemical reaction take place? = Answer • NO. The crystal dissolved and then the water evaporated. This was only a phase change not a chemical change. No new substance was created. And Now You Know! Test Your Understanding • If a marshmallow is cooked over a flame and becomes black on the outside, did a chemical reaction occur? Defend your answer = Answer • YES! A new color appears (black crust) and a new substance was produced (carbon crust) Change Frying and Egg Boiling Water Leaves changing in the Fall Rusting Iron Physical Chemical Evidence And One More Time… 1) Color change 2) Gas is formed 3) New odor is formed 4) Precipitate (solid) is formed 5) Temperature rise or drop Your Task • Create a set of notes summarizing the differences between chemical and physical changes/reactions • Check & Reflect p 25 # 1,5-8 & 11 • Section Review page 27 # 5, 10, 11, 12, 15, 16 The Periodic Table http://www.youtube.com/watch?v=lxJe7e5thkI&feature=relmfu http://www.youtube.com/watch?v=0YmypUfvJvY Elements • The number of protons in the nucleus and the distribution of electrons around it determines the type of element and the chemical and physical properties of the element. Eg. hydrogen is 1 p surrounded by 1e- bromine is 35 p and 45 n surrounded by 35 e- 72 Elements • Elements consist of three main classes – 1) metals – 2) non-metals – 3) metalloids • These classifications are based primarily on conductivity 73 Elements • Most elements occur naturally as single atoms (monoatomic) eg. Ne • Other elements occur naturally as combinations of two or more atoms (molecular elements) – diatomics eg. O2, N2, H2, Cl2, – polyatomics eg. P4, S8 74 Relationship of Chemical Symbols to Chemical Names • Alchemists used artistic symbols to represent the elements • Dalton updated the symbols but they were still cumbersome 75 Relationship of Chemical Symbols to Chemical Names • Baron Jons Jakob Berzelius (1779 -1844) developed the letter system in 1814 which was basically the same as the modern system except it used superscripts instead of subscripts. 76 International Union of Pure and Applied Chemistry -IUPAC • Founded in 1919 IUPAC maintains an international system for naming chemicals 77 International Union of Pure and Applied Chemistry -IUPAC • IUPAC develops rules, guidelines and standard conventions for the study of all aspects of Chemistry • IUPAC allows chemists to communicate clearly and precisely 78 Modern Atomic Symbols • The same in every language. • First letter upper case, second letter lower case (eg Co not CO) (if any) NOTE: Figure A2.3 p 30 Memorize element names and symbols of rows1- 4 [The elements most prevalent in living things are H, C, N, O, S, & Ca.] Worth of the Human Body link1 Worth of the Human Body link2 79 Periodic Table • John Newlands (1837-1898) • English scientist organized the 62 known elements on basis of atomic mass four years before Mendeleev (1864). H1 F8 Cl 15 Co/Ni 22 Br 29 Pd 36 I 42 Pt/Ir 50 Li 2 Na 9 K 16 Cu 23 Rb 30 Ag 37 Cs 44 Tl 53 Gl 3 Mg 10 Ca 17 Zn 25 Sr 31 Cd 34 Ba/V 45 Pb 54 Bo 4 Al 11 Cr 18 Y 24 Ce/La 33 U 40 Ta 46 Th 56 C5 Si 12 Ti 19 In 26 Zr 32 Sn 39 W 47 Hg 52 N6 P 13 Mn 20 As 27 Di/Mo 34 Sb 41 Nb 48 Bi 55 O7 S 14 Fe 21 Se 28 Ro/Ru 35 Te 43 Au 49 Os 51 80 Periodic Table - reference p 31 • Dimitri Mendeleev (1869) & Lothar Meyer (1870) individually arranged know elements in periodic repetition of characteristics • The success of this model was that it was predictive 81 The Periodic Table, Your Tool in Chemistry • The elements on the periodic table are organized according to their atomic number • There are about 115 elements known – Only 90 are naturally occurring – All elements are divided into one of three categories Metalloids Trends • The periodic table is arranged according to three • Main things: a) Increasing atomic mass b) Grouped in families because they have similar chemical properties c) Reactivity increases from L to R and from Top to Bottom • Each horizontal row is called a period – Grouped together because of energy levels – Numbered 1-7 • Each vertical column is a group or family – All members of a family have similar chemical properties – Numbered 1-18 Group (Medeleev used the term families) • Vertical columns, elements with similar chemical and physical properties (Alkali metals, alkali earths, halogens, & noble gases) - Group 1 & 2 are classified with increasing reactivity from top to bottom - Group 17 decreases in reactivity from top to bottom (fluorine is one of the most reactive substances known) (Note: Noble gases generally unreactive) 86 Period • Horizontal rows with elements arranged so that their properties repeat periodically across the table • Each new row represents repeating trends in reactivity 87 Groups or Families • Notable groups: - Alkali metals - Alkaline-earth metals - Noble gases - Halogens Alkali Metals • Soft, shiny, silver in colour • Very reactive with water • Compounds tend to be white solids that are soluble in water • They have 1 valence electron (e-) * Note that Hydrogen does not belong to the alkali family Alkaline Earth Metals • • • • Shiny, silver, but not as soft as alkali Compounds white, but less soluble than alkali Tend to react with oxygen 2 valence e- Halogens • Non-metals • Poisonous and react with alkali metals to form salts • Cl and F (g) Br (l), I (s) • 7 Valence e- Nobel Gases • Colorless gases • Very unreactive Other Groups Group 3-12: Transition Metals • They tend to be shiny, malleable and conduct thermal energy and electric current well Staircase line • Separates metals & non-metals • Metalloids border the staircase 94 Metalloids • Solids at room temperature • Have properties of metals and non-metals • Eight metalloids are: B, Si, Ge, As, Sb, Te, Po At border the staircase line of the periodic table • Important in semi-conductor industry 95 Lanthanides Lanthanides Elements in the first row that follow lanthanum are called lanthanides. These are also called rare-earth and inner-transition metals. These can be found naturally (though rarely, considering the name!) on earth. Actinides • Elements in the second row that follow actinium are called actinides. • These are all radioactive and some are not found in nature. Some of the elements with higher atomic numbers have only been made in labs Some Things to be Aware of… • Group 1 elements always donate 1 electron • Group2 elements always donate 2 electrons • Group 17 elements accept 1 electron • Group 16 elements accept 2 electrons Types of Elements Metals: – Solid at room temperature (except mercury) – Most are lustrous, malleable and ductile – Most will react with non-metals – They are good conductors of electricity – Exist to the left of the “staircase” Characteristics of Metals • 1. Conduction: Metals are good at conducting electricity. Silver (Ag) and copper (Cu) are some of the most efficient metals and are often used in electronics. 2. Reactivity: Metals are very reactive, some more than others, but most form compounds with other elements quite easily. Sodium (Na) and potassium (K) are some of the most reactive metals. 3. Chemical: Metals usually make positive ions when the compounds are dissolved in solution. Also, their metallic oxides make hydroxides (bases) (OH-), and not acids, when in solution. 4. Alloys: Metals are easily combined. Mixtures of many metallic elements are called alloys. Examples of alloys are steel and bronze. Alloys • solid solutions which usually have a metal as the solvent and a metal or a non-metal as the solute. Eg. Steel (Fe, C), solder (Sn, Pb/Ag), bronze (Cu, Sn), amalgam (Hg, Ag, Sn) 101 Types of Elements Non-metals: - Find them right of the “staircase” - Can exist as solid, liquid, or gas at room temp -Appearance varies, not very shiny -Poor conductors of electricity - Brittle/not ductile Element Song - YouTube Your Task • Periodic table worksheet • Read page 28-39 make sure to add notes for understanding on the following: – – – – – SATP Metals as elements Non-metals as elements, molecules that are elements Metalloids as elements The Modern Periodic Table and • • • • • • • • • Families or groups Periods The staircase The table key Alkali metals Alkaline earth metals Halogens Noble gases salts Atomic Theory Atom Song – YouTube Atoms • An atom is made up of three types of subatomic particles: – Protons p+ – Electrons e – – Neutrons n° These subatomic particles were discovered by… Atoms • The protons and neutrons are contained in the nucleus, while the electrons are outside the nucleus • Neutral element (no charge) – Electrons and protons are the same Organization Atomic Number • Elements are defined by the number of protons contained in the nucleus • The number of protons found in the nucleus of an atom is its atomic number Example: Hydrogen (H) - 1 p - 1H Gold (Au) - 79 p - 79Au 109 Isotopes • Isotopes: differing numbers of neutrons in an atom but the same number of protons – Results in a different atomic mass Note: • All three isotopes of hydrogen contain 1 proton in the nucleus • All three isotopes of hydrogen have 1 electron outside the nucleus • The three isotopes of hydrogen differ in the number of neutrons in the nucleus http://www.youtube.com/watch?v=BYX312koKps Isotopes PROTON AND ELECTRONS STAY THE SAME BUT THE NEUTRON # IS DIFFERENT 112 113 Example • Let’s consider the naturally occurring isotopes of chlorine: • 75.53% of all atoms are 3517 Cl isotopes; the atomic number is 17 and the atomic mass is 35 • 24.47% of all atoms are 3717Cl isotopes; the atomic number is 17 and the atomic mass is 37 • Note the atomic mass changes because the number of neutrons increases the mass Example: • An isotope with 45 proton and 48 neutrons would be what? Symbol • Ms. Godley remember to talk about how to symbolize isotopes Naming Isotopes • “atom name” – “mass number” Example: Oxygen-18 # of e = # of p = # of n = Symbol = Silver-200 # of e = # of p = # of n = Symbol = Copper-150 # of e = # of p = # of n = Symbol Isotopes • The average mass of each atom is what is recorded for each element in the Periodic Table. • All elements in the Periodic Table have isotopes that results in atomic numbers that are not whole numbers Energy Levels - aka Energy Shells • Electrons in atoms can only be at certain energy levels. [cf. stairs] • Transitions between energy levels involves either a gain ("up") or loss ("down") of energy. • Energy levels closest to the nucleus are always filled up first 120 Energy Levels The maximum number of electrons in each level is: Level 1 2 3 4 # of electrons 2 8 8 18 This level is filled 1st Bohr Diagrams *Lower levels must be filled before higher levels get any electrons. 121 Let’s Reintroduce Ourselves to Bohr Diagrams Drawing E Level Diagrams • Draw E level diagrams for various atoms on board and state how many valence electrons they have 123 Energy Levels & Chemical Properties • Why would sodium with 11 electrons have similar chemical properties as potassium with 19 electrons? • Draw E level diagrams for each! 124 A Little Chemistry Humor… • Why does hamburger have lower energy than steak? Because it's in the ground state. Valence e- & Octet Rule • The electrons occupying the outer- most E level are called valence electrons • Members of the same group (family) have the same number of valence e- 127 Valence e- & Octet Rule • Electrons are added as you move from left to right in a period • Notice that the period number is the same as the number of occupied energy levels • Notice also that the group number is the same as the number of electrons in the valence shell (Note: if > 10 remove the 1) 128 Valence e- & Octet Rule • Noble gases are stable because they have completely filled energy shells • Atoms bond in such a way as to create a stable octet (duet) [illustrate using NaCl] 129 Your Task • Draw energy diagrams and valence diagrams for 6 elements of your choice. You must show all of your work/steps Ions Let’s Get a Few Things Straight • Atoms = Proton # = electron # neutral charge Strive to be like Noble Gases Ions & Ionization - reference p 34 - 36 • Atoms "try" to become stable by having their outer energy levels filled with electrons (like noble gases). To do this they must either: a) gain or lose electrons (ionization), or b) share electrons with other atoms. Eg. • Na loses 1 e- Na+ ion. Cl gains 1 e- Cl- ion. [draw Li and F ion E diagrams] 133 Ions & Ionization • When atoms gain or lose e-, they become charged, forming ionic species or ions • Non-metal atoms tend to gain electrons to have a full outer energy level. Negative ions are called anions. [anion = a negative ion] Eg. Cl + 1e- Cl- 134 Ions & Ionization • Metals tend to lose electrons to have a full outer energy level. Positive ions are called cations. Eg. Na(s) Na+ + 1e- 135 • Summary: Ions – Ions are formed by the gain or loss of electrons – Anions: gain electrons (negative) – Cations: lose electrons (positive) – Non-metals gain electron(s) to become negative (anions) – Metals lose electron(s) to become positive (cations) – Members of the noble gas family do not ionize. Why???? Naming Ions • Metals do not change names when ionized – Example: Sodium is still called sodium in a compound. We simply add the word ion if it is alone • Non-metals have the last three letters of their names changed to “ide: – Example: Chlorine becomes chloride Multivalent Elements • Some atoms can exist as more than one ion type • On the periodic table the most common form of the ion is listed first Example: Cu2+ and Cu1+ 138 Recall that… • Ionic Compounds form when we transfer valence electrons • Molecular Compounds form when we share valence electrons Lewis Dot Diagrams • Only looks at Valence electrons Rules: 1) The element symbol represents the nucleus and inner-filled energy levels 2) A “dot”, “x”, or “O” represents valence electrons 3) There are 4 valence orbitals of the valence energy level 4) Each valence electron is distributed like compass points but must fill up 1 each before doubling up Lewis vs. Bohr Your Task • Check & Reflect p 39 #1 – 12 • Draw 5 Lewis dot models NOT in the examples and must be a representative element (#1-20) • Practice Questions in notes Compounds • Categories of Compounds: 1) Ionic (metal + non-metal) 2) Molecular (non-metal + non-metal) 3) Intermetallic (metal + metal combinations) Compounds 1) Ionic Compounds: metal/non-metal combination -Positive charge and negative charge combinations -They are all solid at SATP (standard ambient temperature and pressure) -Dissolve in water to form electrolytes Conduct electricity Take Electrons Ionic Compounds Identify the ionic compound(s) • Na3N • CS2 • SbBr3 • H2O End in “OH”- hydroxide Bases are Ionic Compounds • • • • Feel slippery Turns litmus paper blue Have a bitter taste Have a pH greater than 7.0 Also known as Covalent Compounds Compounds Share electrons 2) Molecular Compounds: non-metal and non-metal combinations (no charges!) - They can be solid, liquid or gas @ SATP -Do not conduct electricity Acids are Molecular Compounds • Unlike other molecular substances in water, acids DO conduct electricity • pH less than 7 • Taste sour H+ Hydrogen ion Recall… • Acids and bases will neutralize each other to form salts and water Compounds 3) Intermetallics: metal and metal combinations • Made of two or more different metal atoms Example: brass copper + zinc bronze copper + tin Compounds Ionic Compounds - Solid at SATP - aq (soluble in water) - Conduct electricity - pH greater than 7 - Red blue - Bitter, slippery Molecular Compounds - aq Conduct electricity pH less than 7 Blue red Sour Corrosive Chemical Bonds • Forces of attraction between atoms Ionic: Taken (stronger) Molecular: Shared 155 Types of Bonds Ionic Bonds • An ionic bond is formed by the attraction of oppositely charged atoms or groups of atoms. When an atom (or group of atoms) gains or loses one or more electrons, it forms an ion. Remember: -Cation = positive -Anion = negative Types of Bonds Covalent/ Molecular Bonds Let’s Draw a Lewis Dot to Explain!!!! • A covalent chemical bond results from the sharing of electrons between two atoms • A single covalent bond represent the sharing of two valence electrons (usually from two different atoms). • The Lewis structure below represents the covalent bond between two hydrogen atoms in a H2 molecule. Summary • Ionic Compounds form when we transfer valence electrons • Molecular Compounds form when we share valence electrons Remember… • Metals do not change names when ionized • Non-metals have the last three letters of their name changed to “ide” • Composed of metallic cations ( +charged) and a non metal anion ( - charge) • Form because of the attraction between + and charges Naming Compounds • Recall that compounds are pure substances which are combinations of two or more elements. • To form compounds, at least two atoms must form ions. One atom will lose an electron, while the other will gain the electron • Compounds can be classified by the type of chemical bonds they have: 1. Ionic 2. Molecular 160 http://www.youtube.com/watch?v=9Zl5Y4earhM Two types of Compounds: 1. Compounds which are combinations of metals and non-metals are called ionic compounds. • Number of electrons lost will always equal the number of electrons gained electrically neutral • Subscripts are used to represent the number of ions in a compound 2. Combinations of non-metals with other non-metals are called molecular (covalent) compounds. 161 Ionic Compounds • Collisions between non-metal atoms and metal atoms results in electron transfer, forming electrically charged ions. The attractions of oppositely charged ions are called ionic bonds. Eg. KCl 162 Ionic Compounds • These bonds form a 3-D crystal lattice arrangement of ions and the resulting compound is called an ionic compound. 163 Ionic Compounds Ionic compounds are; solids at room temperature, dissolve in water to some extent, and conduct electricity when in solution. 164 Types of Ions: • Monoatomic (simple) ions - Formed from single atoms. Eg. Ag+, Al3+, etc. • Polyatomic (complex) ions - Formed from groups of atoms & behave as a single unit. Eg. CH3COO-, PO43-, NO3-, SO42166 Ionic Compounds: Important Rule • IN AN IONIC COMPOUND ....… TOTAL POSITIVE CHARGE = TOTAL NEGATIVE CHARGE (Example: net charge = 0) Eg. Sodium chloride Na+ + Cl- -> NaCl Ca 2+ + F- -> CaF2 You need to balance the charges http://www.youtube.com/watch?v=h79HW83aoEw&feature=relmfu Naming Ionic Compounds reference p 40 - 46 NOMENCLATURE AND FORMULA WRITING • IUPAC rules for naming (same in any language) 1. Binary Ionic Compounds 2. Polyatomic Ionic Compounds http://www.youtube.com/watch?v=szhKhCrjUDI http://www.youtube.com/watch?v=NUVO_TKi6Vo&feature=relmfu 168 1. Ionic Compounds Nomenclature - Name of metal ion (cation) followed by name of non-metal (anion) ion suffix (ide). - Use lower case letters. - Name gives no indication of the number of each ion present. (Can be determined by balancing overall charge) Eg. calcium chloride - (CaCl2) 169 1. Ionic Compounds • If the metal can exist as ions with differing charges (eg. Cu), the ion charge is indicated with a Roman numeral in parentheses after the metal ions name. Eg. copper (I) chloride = CuCl copper (II) chloride = CuCl2 170 1. Ionic Compounds Formula Writing - Symbol of metal ion followed by symbol of non-metal ion. - Ion ratio is indicated by numerical subscripts (no subscript means one). Eg. NaCl, CuCl2, Fe2O3 1:1 1:2 2:3 171 1. Ionic Compounds Ratios are determined by balancing overall positive and negative charge Eg. potassium oxide = K2O aluminium sulfide = Al2S3 aluminium oxide 2 Al3+ + 3O2- Al2O3 172 Summary of Writing and Naming Ionic Compounds 1) Write the metal (cation) first using the element name 2) Write the nonmetal (anion) 2nd using the anion name and changing the last part to “ide” 3) Use the charges on each ion to determine the lowest whole # ratio that produces an overall charge of zero Multivalent Ionic Compounds • Ionic compounds where the metal can have multiple charges – Example: Multivalent Compounds • The charge that is listed first is the most common charge – Example: Cu2+ Cu+ Fe3+ Fe2+ • When there is more than 1 charge listed we MUST specify which charge was used with Roman Numerals Multivalent Compounds Rules: 1. Name the metal using the element name and follow with the roman numeral that corresponds to the charge 2. The negative charged ion will take the anion name NO ROMAN Numerals Your Task • Practice Problems 1 & 2 p 43-44 Polyatomic Ionic Compounds(Table of Polyatomics) • Poly “many” • Atomic “atoms” • Groupings of atoms that make their own special ion • Unlike a molecule which is neutral, these complex substances are like ions and carry a charge • CANNOT exist alone, they are neutralized by combining with another ion Polyatomic Ionic Compounds Nomenclature - Name of positive ion followed by name of negative ion. Eg. NH4Cl, NH4OH Parentheses may be required to keep numerical subscripts apart. Eg. Cu(NO3)2 181 Polyatomic Ions aka Complex Ions • Poly “many” • Atomic “atoms” • Groupings of atoms that make their own special ion • Unlike a molecule which is neutral, these complex substances are like ions and carry an electric charge • CAN NOT exist alone, they are neutralized by combining with another ion Polyatomic Ions • Naming Polyatomic Compounds – The positive ion is named first followed by the negative ion – Parentheses may be required to keep numerical subscripts apart Polyatomics –The positive ion is named first followed by the negative ion –The one with the lesser number of oxygen atoms has an “ite” ending while the one with the greater number of oxygen atoms has an “ate” ending • Nitrate vs. Nitrite Polyatomic Ions The most common polyatomic ions are: • • • • • • • • • • • • Ammonium Carbonate Chlorate Dihydrogen phosphate Hydrogen carbonate Hydroxide Nitrate Nitrite Permanganate Phosphate Sulfate Sulfite NH4+ CO32ClO3H2PO4HCO3OHNO3NO2MnO4PO4-3 SO4-2 SO3-2 And the best part about these is that you will need to memorize them all 2. Polyatomic Ionic Compounds Formula Writing - Positive ion first, negative ion last, using numerical subscripts to indicate ion ratio (found by balancing overall charge). Eg. iron (III) oxalate, ammonium chlorate, ruthenium (III) borate magnesium nitrate You get to copy these down in your notes at the bottom of the page!!! 186 Solubility • Ionic compounds usually dissolve in water(aq) BUT some form a precipitate (solid matter) • Can dissolve (soluble) = (aq) “aqueous” • Cannot dissolve (non-soluble)= (s) – leave it solid • We can use the solubility table to identify (aq) and (s) ionic compounds Solubility of Ions in Water • Group 1 metals and NH4+, H+, ClO3-, NO3- and ClO3- are all soluble in water • CH3COO- when in compounds is soluble in water except when added to Ag+ • Br-, Cl- and I- are soluble in water except when added to Ag+, Pb2+, and Cu+ • SO4-2 is soluble in water unless mixed with Ca2+, Sr2+, Ba2+, Ra2+, Pb2+, and Ag+ • OH- is soluble in water when combined with Group 1 elements, NH4+, Sr2+ or Ba2+ • PO4-3, SO3-2 and CO3-2 are soluble in water only when added to Group 1 elements and NH4+ A Few Examples… • Lead (II) iodide= • Sodium bromide= • Potassium sulfate= • Magnesium hydroxide= *If you DO NOT find the ions in the table assume it is a solid Ionic Compounds Review • Five properties that ionic compounds have in common –Have high melting points –Are solid at room temperature –Form crystal lattices –Will dissolve in water to a certain extent –When dissolved in water, the solution will conduct an electric current Your Task Practice Problems 3 & 4 p 46 State if the following are soluble in water • • • • • • • Sodium chloride Ag(NO3) Magnesium sulfate Calcium sulfate Be(OH)2 Na3(PO4) Lithium chlorate Name or write the formula for the following compounds. As well as state if they are soluble or insoluble • • • • • • Pb(SO4)2 Al(CN)3 Chromium (lll) hydroxide Cu3P Tin(ll) nitrite Fe(NO3)3 Molecular Compounds (aka covalent compounds) If no metal atoms are available to give electrons to non-metal atoms, the non-metal atoms share electrons to fill their outer energy levels. The bond which forms when electrons are shared is called a covalent bond. 195 Molecular Compounds (aka covalent compounds) Covalent (molecular) compounds exist as separate molecules, unlike the lattice of ionic compounds Eg. F2, H2O Covalent bonds are usually not as strong as ionic bonds (more easily broken). 196 Molecular Compounds (aka covalent compounds) Molecular compounds can be solids, liquids or gases at room temperature. Molecular compounds tend to be poor conductors of electricity 197 Electronegativity The strength with which an atom holds onto its outer (valence) electrons affects covalent bonding. On the periodic table, electronegativity increases as you go up and to the right. Fluorine is the most electronegative element. 198 Molecular Elements Recall that some elements exist as molecular elements when in pure form ie. diatomic (HOBrFINCl At) and polyatomic (P4 & S8) 199 Diatomic Molecular Elements • ALWAYS have 2 atoms • ALL ELEMENTS IN GROUP 17 ARE DIATOMICS • H2(g), N2(g), O2(g), F2(g), Cl2(g), Br2(l), I2(s), At2(s) Polyatomics (not the type in their own special table) • Phosphorous and Sulfur Molecular Compounds - Two different kinds of non-metal atoms joined by covalent bonds. Nomenclature - use lower case - (numerical prefix) + first name; (numerical prefix) + second name with (ide) suffix. (The prefix "mono" is usually omitted) Eg. N2O5, H2O, H2S 201 Prefixes • Mono 1 • Di 2 • Tri 3 • Tetra 4 • Penta 5 • Hexa 6 • Hepta 7 • Octa 8 • Nona 9 • Deca Binary Molecular Compounds Formula Writing - Write symbols in same order as in name - numerical prefix = numerical subscript (omitted when "one") eg. silicon tetrahydride oxygen dibromide carbon tetrafluoride 204 Molecular Compounds Memorize Common Molecular Names and Formulas; ammonia - NH3 glucose - C6H12O6 sucrose - C12H22O11 hydrogen peroxide - H2O2 hydrogen sulfide - H2S methane - CH4 Propane- C3H8(g) Methanol- CH3OH(l) Acetic Acid- CH3COOH(aq) Ethanol - C2H5OH(l) http://www.youtube.com/watch?v=2_W177RKt0U&feature=relmfu 205 Anomalous Properties of Water Reference p 60 Water is a polar molecule and exhibits some ‘unusual’ properties: • • • • adhesion & cohesion (surface tension) high c high mp & bp density 206 Naming Molecules that Contain Hydrogen • • • • • • • • • • • H2O(l) – water H2O2(l) – hydrogen peroxide NH3(g)- ammonia C12H22O11(s)- sucrose C6H12O6(s) – glucose H2S(g)- hydrogen sulfide CH4(g) – methane C3H8(g) – propane CH3OH(l)- methanol CH3COOH(aq) – acetic acid vinegar C2H5OH(l) - ethanol You need to memorize these ! Your Task Practise Problem 5 p 49 Check and Reflect p 50 # 1 -12 Supplemental p 77 # 28, 36 Check Answers p 495 Do Check and Reflect p 61 # 6 & 8 Acids and Bases Arrhenius • Arrhenius proposed that any substance that: – Produces hydrogen ions when dissolved in water = acid – Produces hydroxide ions when dissolved in water = base • We now refer to these as Arrhenius Acids and Bases Acids • • • • • • • • (aq), (s), (l), or (g) at room temp Soluble in water (to some degree) Solutions conduct electricity React with active metals to produce H2 (g). Taste sour Turns blue litmus paper red pH of <7 Identified by the presence of a H+ http://www.youtube.com/watch?v=70fot8t9zts&feature=relmfu Acids • Strong Acids: – Hydrochloric acid(HCl) – Nitric acid (HNO3 ) – Sulfuric acid (H2SO4 ) – Hydrobromic acid (HBr) – Hydroiodic acid (HI) – Perchloric acid (HClO4 ) It is in your best interest to become very familiar with these acids. This will save you significant amounts of time, especially on exams. • Strong acids will completely ionize conductive • Weak acids will not ionize low conductivity Bases • • • • • • • • Usually solids Usually soluble in water. Do not react with active metals. Solutions conduct electricity. Feel slippery, taste bitter. pH >7 Red litmus paper turns blue Identified by the presence of a OH213 Summary Property Acid Base Taste Sour Bitter Feel Wet Slippery Reaction with Mg Produce Hydrogen gas No reaction Indicator Turns blue litmus red Turns red litmus blue Conductive Yes Yes pH <7 >7 pH scale • Describes whether a substance is acidic, neutral or basic (alkaline) • Refers to the potency or strength of hydrogen ions present in the substance 215 pH Scale • Each number change represents a ten-fold change in the hydrogen ion concentration. – For example, a pH of 1 is 10 times stronger than a pH of 2 Indicators • Chemicals used to determine whether a solution is acidic or basic • Change color at different pH Example: litmus paper - bases -> turn red litmus blue - acids -> turn blue litmus paper red 217 Indicators Reference p 63 • Some other indicators • The Universal indicator is a mixture of several liquid indicators that turn different colors when exposed to hydrogen ions 218 Buffers • Any substance that keeps the pH of a solution nearly neutral even is a small amount of acid or base is added • Example: sodium hydrogencarbonate (NaHCO3) – is in the bloodstream. It neutralizes both acids and bases that enter the blood. Acid Nomenclature Acid Nomenclature - name as an ionic compound and then switch name according to acid naming table (next slide) or.... "aqueous" before the ionic name. No one really uses this nomenclature anymore. So you do not need to worry about it for the purposes of this course An acid formula should always be followed by the subscript (aq), meaning dissolved in water. 220 Acid Nomenclature • Hydrogen ion acts like a positive charge – Will be paired up with a negative charge Classical Naming System for Acids 1. If a substance ends with “ide” the acid becomes “hydro ic acid” - Hydrogen ide becomes hydro Example: See notes ic acid Classical Naming 2. If the substance ends with “ate” the acid becomes “ ic acid” - Hydrogen ate becomes ic acid Examples: See notes Classical Naming 3. If the substance ends with “ite” the acid becomes “ ous acid” -Hydrogen ite becomes ous acid Examples: See notes Acid Naming Table • hydrogen _ide becomes hydro_ic acid • hydrogen _ate becomes _ic acid • hydrogen _ite becomes _ous acid http://www.youtube.com/watch?v=70fot8t9zts 226 Or…. • H+ ide hydro Hydrochloric acid ic acid • H+ Sulfurous acid ite ous acid • H+ Nitric acid ate ic acid Exceptions • Hydrogen ions can hide at the end of the formula • Organic acids Formulas begin with carbon and end with H ions (hard to tell apart from bases) – Ethanoic acid (CH3COOH) acetic acid • • • • • • • • • • Acid Naming H2S HClO3 H2SO4 HNO3 H2SO3 HNO2 HClO Phosphoric acid Hydroiodic acid Hypochlorous acid Let’s name and write the formulas for these wonderful acids!!! Nomenclature of Bases • For this course we will consider bases to contain the hydroxide ion (OH) • Name as an ionic compound [Note: -COOH is an acid] 230 Naming Bases • The name of a base always ends in “hydroxide” • Some common Bases – Sodium hydroxide – Ammonium hydroxide – Calcium hydroxide – Magnesium hydroxide Neutralization Reaction Reference p 68 When an acid and a base are mixed they neutralize each other by producing water and a salt. (more later but be aware that salt doesn’t always mean NaCl) ACID H+ + BASE OH- = HOH Buffer - has the ability to neutralize either an acid or base solution 232 Task as a Class Check & Reflect p 69 # 3,4 & 7 Acid and Bases Naming Worksheets (included in your notes) Chemical Change • Is a process that involves recombining chemical bonds between atoms and energy flow. • New substances are formed. 234 Collision Reaction Theory • For a chemical reaction to take place: – Particles of the reactants before a rearrangement can occur – A minimum energy is particles – A certain orientation is particles for successful rearrangement must collide required of the required of the 235 Reactions Reactants - the substances which go into a chemical change. Products - the substances which come out of a chemical change. Reactants Products C6H12O6(s) + 6 O2(g) --> 6 CO2(g) + 6 H2O(g) 236 Rules for Chemical Equations • Use proper chemical symbols • Reactants on left, arrow pointing to products on right • Show states of matter (s, l, g, or aq) [Note: p 80] • Are balanced (matter is conserved) • include energy when possible More to come 237 Energy Changes Reference p 81 - 82 • In the formation of ionic compounds a considerable amount of energy (heat) may be absorbed (endothermic) or released (exothermic). • Spontaneous reaction - occurs on its own - when reactants are in contact. 238 Exothermic Reaction Reference p 81 LINK & LINK TOO • The energy not used to fuel the reaction (excess energy) is released to the surroundings. (energy written as product) Eg. combustion of methane CH4(g) + 2 O2(g) CO2(g) + 2 H2O(g) + E • energy stored in newly formed bonds is less than the energy available from the breaking of bonds in the reactants. Endothermic Reaction Reference p 81 link link too • Reactions which require energy from outside the system to drive the reaction. (energy written as reactant) Eg. Cold pack NH4Cl(s) + H2O(l) + E NH4+(aq) + Cl2 (aq) + H2O(l) • Energy required to form the new bonds is greater than the energy stored in the bonds of the reactants. 241 Biochemical Reactions Reference p 82 • Essential to all life on Earth Respiration C6H12O6(s) + 6 O2(g) --> 6 CO2(g) + 6 H2O(g) + E Photosynthesis 6 CO2(g) + 6 H2O(g) + E --> C6H12O6(s) + 6 O2(g) [also chemosynthesis] 242 Characteristics of Chemical Reactions Reference p 84 • Colour or density change • Gas produced (bubbles, new odour) • Energy change (heat, light, electricity) • A solid (precipitate) forms or dissolves 243 Conservation of Mass Reference p 84 • Antoine Lavoisier showed that the total number of each kind of atom remains the same in a reaction. That lady who did a lot of work but because of her gender doesn’t get any credit Father of Modern Chemistry Your Task • Check and Reflect p 85 #1-7,10,11 Recall… • Reactants- the substances which go into a chemical change • Products- the substances which come out of a chemical change Recall… • Word equations: use the “ chemical formula” • Formula Equations: use the “ chemical formulas” of reactants and products to represent a reaction Recall • Skeleton Equations: a formula equation that shows the identities of the substances involved in the reaction and which elements are present. Recall… • Balanced Equations: integers called “Coefficients” are used to show equal number of each element present on both sides Conservation of Mass • The amount of the reactants equals the amount of the products Law of Conservation of Mass & Balanced Equations • Chemical equations must be balanced this is important when writing chemical equations 251 Chemical Equation Basics 101 • Coefficient- the number in front of a symbol; represents the number of atoms or molecules - 3Na Coefficient • Subscript- the number behind the symbol; represents the amount of each atom - N2O5 Subscript Chemical Equations Basics 102 • Recall that the following elements NEVER exist as single atoms. They are represented as shown: N2 H2 O2 F2 Cl2 Br2 I2 At2 S8 P4 * Sometimes H2O is written as HOH Chemical Equations Reference p 86 - 88 • Adding coefficients to a skeleton equation, balances the equation to show the same number of atoms, of each type, on either side of the equation. • Coefficients in equations may be read as # of molecules, or # of moles. 254 Hints for Balancing Equations (NIT) 1. Make sure all reactants and products are written with correct formulas. 2. Balance each atom, one at a time, using numerical coefficients (leave O and H to the end). 2a. Balance complex ions as groups 3. Check if the coefficients can be reduced by a common divisor. 4. Sometimes writing water as HOH can make the solution easier to see. Remember • When counting elements don’t forget to look at the subscript and coefficient – P2O5 (has 2 phosphorous and 5 oxygen atoms) – 2P2O5 (has 4 phosphorous and 10 oxygen atoms) • Because there are 2 molecules (indicated by the coefficient) and 2 atoms in each molecule (indicated by the subscript) • NEVER CHANGE A SUBSCRIPT TO BALANCE AN EQUATION!!!!!! USE THE COEFFIECIENT INSTEAD! Let’s Do An Example Together • The decomposition of dinitrogen pentaoxide gas Step 1: Represent the reactant and products with formulas N2O5(s) N2(g) + O2(g) Notice that I included the states • Step 2: When there is an odd number of atoms of an element on one side and an even number on the other side, find a common multiple. • For oxygen the common multiple of 5 and 2 is 10 This requires you to understand what a multiple is • Step 3: Multiply each formula by a coefficient to represent the same number of atoms of oxygen on both side 2N2O5(s) N2(g) +5 O2(g) • Step 4: Add coefficients to the remaining formulas to balance the remaining atoms 2N2O5(s) 2N2(g) +5 O2(g) • Step 5: Check to make sure all atoms are balanced Left side Right Side 4 N and 10 O 4 N and 10 O Balance the following example equations: A. ___ Al(s) + ___ Br2(l) --> ___ AlBr3(s) B. _ C6H12O6(s)+ _ O2(g) --> _ CO2(g)+ _ H2O(g) C. Solid iron reacts with aqueous copper (II) chloride to form aqueous iron (II) chloride and solid copper 261 Further Assistance on Balancing Equations • http://www.youtube.com/watch?v=dQrV8Rdu ttU • http://www.youtube.com/watch?v=le5zr1kLE 4U • http://www.youtube.com/watch?v=3UeD32Q sKYM&feature=relmfu Your Task Example Problem p 89 Practice Problem p 89 # 1 Check & Reflect p 90 # 3,6 – 9 [W/S Balancing Chemical Equations – Supplemental] Five Reaction Types Reference p 91 - 106 1. Formation (synthesis) 2. Decomposition 3. Hydrocarbon Combustion 4. Single Replacement 5. Double Replacement 264 1. Formation (synthesis) Reference p 91-93 • Composition reactions • Two simple elements combine to form one complex compounds. • element + element --> compound – A + B AB – Reactant + reactant = product 2 K(s) + I2(s) --> 2 KI(s) + thermal E Example Problems p 92 & 93 Let’s Do This Together… • Example Problems p 92 & 93 2. Decomposition Reference p 94 • Complex compounds break down to form elements • Compound --> element + element You have to add energy to • ab a + b Eg. the system in order for water to separate 2 H2O(l) + electrical E --> 2 H2(g) + O2(g) http://www.youtube.com/watch?v=FyUvZ3ldtUU&feature=fvsr 3. Single Replacement Reference p 96 - 97 • Element + aqueous ionic compound --> new Element + new ionic compound • a + bc b +ac Cu(s) + 2 AgNO3(aq) -> Cu(NO3)2(aq) + 2 Ag(s) A + BC AC + B http://www.youtube.com/watch?v=WwH8I_K3yYM&feature=fvsr 4. Double Replacement Reference p 100 - 101 • Ionic compound + ionic compound new ionic compound + new ionic compound • ab + cd ad + bc Example Pb(NO3)2(aq) + 2 NaI(aq) -> 2 NaNO3(aq) + PbI2(s) Example Problem p 100 Practice Problem p 100 # 9 http://www.youtube.com/watch?v=7TtOUf91VSU&feature=fvsr 5. Hydrocarbon Combustion (oxidation reactions) Reference p 95 • A compound reacts with oxygen to form the most common oxide • If O2 is not limited -> produces H2O(g), CO2, and thermal/light energy Eg. burning propane C3H8(g) + 5 O2(g) -> 3 CO2(g) + 4 H2O(g) + E http://www.youtube.com/watch?v=CY9ldQjdgT8 Oxygen Reactions Oxygen is always present in these reactions! 1. Combustion Reactions 2. Corrosion Reactions: metal and oxygen always the reactants 3. Cellular Respiration: Glucose always one of the reactants Predicting Products Reference p 102 - 105 • Using knowledge gained regarding reaction types it is now possible to predict products of reactions. Your Task Example Problems p 100 - 105 Practice Problem p 103 #10 - 12 Check & Reflect p 106 # 1 – 8 [W/S Translating Formulae – Supplemental] What is a Mole? • We use “convenient numbers” in everyday situations • A “convenient number” is a simplified way to describe the number of things • It is often easier to group objects when counting them So What is a Mole? • A mole is a counting unit. It tells us how many objects there are. • A pair and a dozen are also counting units Pair means 2 objects Dozen means 12 objects Mole means 6.02x1023 objects So what is a Mole? • A term used to represent the number of objects • The amount of any substance containing 6.02x1023 particles – To define a mole, chemists measure carbon – And we also call 6.02x1023 Avogadro's Number Where does this Number come from? • Scientists took 12g of pure carbon-12 (composed of 6 protons+ 6 neutrons; mass number =12) and counted how many carbon atoms were in the sample • Guess how many there were? Avogadro’s Number -The Mole Reference p 107 • # of atoms in 12.0000 g of 12C is called a mole (mol) = 6.02 x 1023 Avogadro’s # • this number of atoms is observable and measurable. • coefficients in equations may be read as # of molecules, or # of moles. http://www.youtube.com/watch?v=xiVweBpjXJo&feature=related http://www.youtube.com/watch?v=xqw2BWdKl1Q&feature=relmfu http://www.youtube.com/watch?v=O7qjYRYxkso&feature=relmfu 283 Mole Ratio • The ratio of the coefficients gives us the mole ratio. 2 Al(s) + 3 Br2(l) 2 AlBr3(s) • In the balanced chemical equation above.. - if 3 mol of Br2 reacts, 2 mol of AlBr3 forms - if 6 mol of Br2 reacts, 4 mol of AlBr3 forms... etc. 284 Molar Mass Reference p 108 - 109 • We can now calculate the mass of one mole of a compound (molar mass ). • Use the atomic molar mass (M), in g/mol multiplied by the number of moles in the chemical formula then add the masses of all elements together Example Problem p 108 Eg. Find the molar mass of H2SO4(aq) hydrogen atomic mass = 1.01 g/mol # of moles in formula = 2 2 mol x 1.01 g/mol = 2.02 g sulfur atomic mass = 32.06 g/mol # of moles in formula = 1 1 mol x 32.06 g/mol = 32.06 g oxygen atomic mass = 16.00 g/mol # of moles in formula = 4 4 mol x 16.00 g/mol = 64.00 g ...molar mass of H2SO4 = 2.02 g + 32.06 g + 64.00 g = 98.08 g/mol 286 Your Task • PP p 108 # 13 – 16 Mass <--> Mole Conversions Chemists created the concept of molar mass to convert between mass and chemical amount These conversions can be performed using a formula (M=m/n) or the factor-label (unit analysis) method. 288 n= m/M Where: n= number of moles m= mass M= molar mass However, you are much too advance to use this formula and are capable of using unit analysis Molar mass is often used as a conversion factor because it contains the units grams and moles. Thus, it can be used to convert grams to moles and moles to grams. Example • How many moles are there in 20 g of carbon? • Find the number of moles of H2SO4 in a 100.0g sample. Mole to Mass Conversion • If we had a 4.60 mol amount of H2SO4, its mass would be........ m = (4.60 mol)(98.08 g/mol) m = 451 g 291 Mass to Mole Conversion • If we had a 0.58 g of H2SO4, its mole amount would be........ n = (0.58 g)/(98.08 g/mol) n = 56.8864 mol n = 5.9 x 10 -3 mol Example Problems p 109 Practice Problems p 108 - 109 # 13 - 20 Check & Reflect p 112 # 1 -12 292 Your Task Example Problems p 109 Practice Problems p 108 - 109 # 13 - 20 Check & Reflect p 112 # 1 -12 Nuclear Reactions (NIT) Chemical reactions involve the valence electrons and the bonds which they form (formation of new molecules). Nuclear reactions changes in the atom's nucleus, the formation of a different atom and thus a new element. 294 Nuclear Reactions (NIT) In nuclear reactions mass changes into electromagnetic radiation (E = mc2). Nuclear reactions involve up to 106 times more energy than chemical reactions. The high energy radiation given off is usually X-rays or gamma rays. 295 Nuclear Reactions (NIT) Radioactivity is the emission of particles and radiation from an atom's nucleus. Some isotopes of an atom's nucleus have an unstable ratio of protons to neutrons which makes them radioactive (radioisotopes). These atoms undergo radioactive decay to become more stable isotopes. Eg. Carbon- 14 14C 14N + radiation 296 Unit REVIEW • Use Section Reviews • p 27 • p 76 - 77 • p 113 and Unit Summary/Review • p 116 - 121 to help you prepare for the Unit Exam (selected answers on pg 498) But not really because that would be mean and inappropriate 297 1. 2. 3. 4. 5. 6. 7. 8. 9. Quiz Acetic Acid Hydrochloric acid Nitrous acid H2CO3 H2SO4 Magnesium hydroxide Chloric Acid Calcium hydroxide What is the molar mass of an unknown oil if 57.29g of the oil contains 0.15mol? 10.How many grams are in 5.2 moles of potassium sulfide?