Science 10 Chemistry Notes (Feb 2022) PDF

Signed Safety Contracts? Science 10 Chemistry Hand out notes/data sheets Look at the timeline A. The Atomic Theory – How we got to where we are now The scientific method in action. The scientific method consists of systematic observation, measurement, and exper...

Signed Safety Contracts? Science 10 Chemistry Hand out notes/data sheets Look at the timeline A. The Atomic Theory – How we got to where we are now The scientific method in action. The scientific method consists of systematic observation, measurement, and experiment, and the formulation, testing, and modification of hypotheses.  Watch This! https://www.youtube.com/watch? v=xazQRcSCRaY Dalton’s Atomic Theory – “Atoms” : 1808  all matter is composed of tiny, indivisible atoms particles called  identical properties atoms of an element have  atoms of different elements have different properties  atoms of two or more elements can combine ratio constant in to form new substances eg) H:0 ratio 2:1H2O  water H:0 ratio 2:2 H2O2 hydrogen peroxide J.J. Thompson – “raisin bun”: 1897 electrons credited with discovery of “raisin bun” “plum model or pudding” model  atom is asphere positive which is , negative with electrons embedded in it like raisins in a bun  most of themass is associated with the positive charge  Ernest Rutherford – “nucleus”: 1911  atoms have nucleus a positive which is massof the and has most  most of the atomempty is space occupied negatively charged electrons by the moving proposed the existence ofprotons Size Demonstration Neils Bohr – “planetary model”: 1913 circular orbits  electrons move in around the nucleus cannot exist between orbits Problems – electrons behave like waves and aren’t stuck in one place. Bohr Model James Chadwick – “neutrons”: 1932  showed that the nucleus must contain heavy neutral particles to account for all of the atom’s mass(neutrons) Schrodinger/de Broglie – “quantum model”: 19  quantum mechanical model electrons havedistinct energy levels  exact locations of electronsnot are defined , but the probable location region in of aspace can be predicted (electron cloud model) B. Atomic Structure  atom are the building blocks ALLofmatter s  consist of tiny a nucleus and a huge “cloud” region  nucleus makes99% up mass of the of an atom  cloud region makes up mostvolume of the of an atom Subatomic Particles 1. Protons (p+)  positive charge  found innucleus  determines the type of element 2. Neutron (n0)  no charge  found innucleus  hold nucleus together used to 3. Electron (e)  smallest subatomic particle  negative charge  “cloud” region found in  arranged in energy levels  maximum # of electrons in each level: Level 4 =we will only go up to 2 e Level 3 =8 e Level 2 =8 e Level 1 =2 e Atomic Mass atomic mass unit (amu) is the proton mass of or a neutron 1.7 =  1024 g net charge zero for all atoms is Information  on periodic table can help determine one or the other # e = # p+ in a neutral atom mass number is the sum of the protons and neutrons (rounded to the nearest whole number)  number of neutrons used to find the eg) lithium 3 atomic number = atomic mass =6.94 = 7 # protons =3 7– 3 = 4 # neutrons = Warm-Up How many p+, e-, and no in each? lithium, Pb, Tc Isotopes  same number of protons atoms that have the different but a number of neutrons  atomic mass on the periodic table is average an mass based on the percentage abundances naturallyof all occurring isotopes  Isotope Names (not in notes) hydrogen – 1 mass = 1 (number after the name) p+ = 1 (atomic number) no = 0 (atomic mass - #p = neutrons) carbon – 13 mass = 13 (number after the name) p+ = 6 (atomic number) no = 7 (13 – 6 = 7 neutrons) Examples (Not in your notes) copper - 64 64 # p + = 29 Cu # e - = 29 29 #n° =64 - 29 = 35 copper - 62 62 # p+ = 29 Cu # e- = 29 29 #n° =62 - 29 = 33 Trickier Question (Not in Notes) Name an atom with 10 neutrons Another Tricky Question In a particle accelerator, iron-56 and gold-197 nuclei are smashed together in an attempt to form a heavy atom. What atom will be formed? Is it likely to be a stable isotope? Activity How could you arrange the coloured shapes on the board? Discuss with a neighbour (2 minutes) Arrange shapes ????? C. The Periodic Table  the periodic table was developed by Dmitri Mendeleev in1800’s the mid http://www.chemicool.com/ 1. Atomic Number  number ofprotons in one atom of an element  increases from left to right and top to bottom 2. Properties  3 major categories: 1. metals good are conductors,strong, malleable(pound into thin sheet), ductile (can draw into a wire, bendable) have high luster; left side of stair case are found on mercury potassiu m copper 2. nonmetals poor are conductors, non-lustrous,weak, etc… opposite properties to right side metals; found on of staircase bromine sulphur iodine 3.metalloids show properties of along both metals and non-metals ; found staircase silicon boron arsenic  The Periodic Table Staircase separates Metals and non-metals 3. Groups  18 verticalcolumns are called groups or families  Group 1 (IA) Alkali Metals Reaction of Alkali Metals with Water – Links in D2L  Group 2 (IIA) Alkaline Earth Metals  Group 17 (VIIA)Halogens  Group 18 (VIIIA)Noble (Inert) Gases  Rare Earths Lanthanide Series (57-71)  Groups 3-12 (B series) Transition Metals The Element Song!  elements in each group share similar chemical properties (reactivity) intensity although changes  increases down group reactivity metals for up group andnonmetals for  group number indicates how many electronsoutermost are in the energy level 4. Periods  horizontal rows  show pattern a in reactivity which changes left from to right  each time you move to a new period you start the pattern over Back to the coloured paper What does the colour of the paper represent? What does the size of the paper represent? What will the next piece of paper in the sequence look like? Is there a missing piece of paper? Check: Find the element 1. 28 protons 2. Noble gas in the 2nd period 3. Its most common isotope has a mass of 16 4. A group 15 element lighter than sodium 5. Halogen closest to silver 6. 3rd period, 16th group 7. Element after Lanthanum 8. Lightest element with 3 electrons in its outer energy D. Electron Energy Level Diagram (EELD) ATOMS  nucleus –shows # p+ and n0  energy levels – shows # of e in each period level (# of levelselement = is in )  always fill thelowest energy level first (closest to nucleus) then move to the next level  valence electrons are the e in outermost same as energy group level ( #, ignoring the 1 in front of groups 13- 18 )  Recall  maximum # of electrons in each level: Level 1 1= 2 Level 2 1=2 3 45 6 78 Level 31 2 3 45 6 78 Level 41 2 3 45 6 78 We usually ignore the transition metals in high school! Except in AP chem! Examples sodium atomic #11 mass #22.99  23 # p+ =11 # e- =11 #n° = 23 – 11 = 12 Level 3 1 Level 2 e- 8 11 e- Level 1 2e-e- p+ =11 n° =12 Na nitroge n atomic # 7 mass #14.01  14 # p+ =7 # e- =7 #n° = 14 – 7 = 7 5 e- 7 e- 2 e- p+ =7 n° = 7 N Draw the EELD for the following: 1. Potassium atom 2. Chlorine atom potassiu atomic 19 # m mass #39.10  39 # p+ =19 # e- =19 #n° = 39 – 19 = 20 1 e- 8 19 e- e- 8 2e-e- p+ = 19 n° = 20 K chlorine 17 atomic # mass #35.45  35 # p+ =17 # e- =17 #n° = 35 – 17 = 18 7 e- 8 17 e- 2e-e- p+ = 17 n° = 18 Cl Warm-Up: Draw an EELD for an oxygen atom Ions  ions are atoms or groups of atoms that net have a charge  number of p+ and e are not equal  most atoms try to achieve the electron configuration noble gasof a  isoelectronic means having the same number of e as another atom or ion gains an electron eg) fluorine to be isoelectronic with neon neon atom  loses an electron eg) potassium to be isoelectronic argon with + potassium atom potassium ion argon atom  Cations  positively charged ions  lost electrons to obtain stable a electron configuration  METALS form cations  Naming: Just add ion eg) sodium loses one e to completely empty the last energy level to become a sodium ION  10 e and 11 p+ Na+ has  charge on a metal ion is the same as the group number 1,2,3,13,14 for groups (ignore the 1 in front of 13 Cations are metals.  Anions  negatively charged ions gained electrons stable to obtain a electron configuration (full energy level)  NONMETALS form anions  NAMING ALERT!! – drop the “gen”, “ur”, “ine” , “ourous”, or “ium” and add “IDE” eg) gains two e oxygen to completely fill the last energy level O2- has10 e and 8 p+ and is called ion oxide the  charge 18 minus on a non-metal is group number (not Roman Numerals) Negative Charges Note carbon doesn’t have a negative charge! The Dash is lacking a charge. In grade 12 you must memorize the charge on group F. Octet Rule DIFFERENT ORDER  stable atoms tend to be when the outer energyfull of is level electrons  the octet rule states that atoms bond in a way tofull havevalence a energy level (there are exceptions)  atoms will either share electrons, or gain or lose electrons in order to satisfy this octet rule  compounds are formed when two or more different elements bond together either by sharing or transferring electrons  EELD’s for Ions number of p+atomic = number number of e number = of p+ – charge number of n0 =atomic mass – atomic number (# of p+) Think about electrons like negative friends! If you get rid of a negative friend you become more POSITIVE! If you gain more negative friends you become more NEGATIVE! Examples 11 sodium ion atomic # mass #22.99  23 # p+ =11 # e- =11 – (+1) = 10 #n° = 23 – 11 = 12 Level 3EMPTY Lost 1 electron) Level 2 8 10 e- Level 1 2e-e- p+ =11 n° =12 Na+ Examples nitride ion atomic #7 mass #14.01  14 # p+ = 7 # e- = 7 – (-3) = 10 #n° =14 – 7 = 7 Level 2 8 10 e- Level 1 2e-e- p+ = 7 n° = 7 N3- Draw the EELD for the following: 1. potassium ion 2. chloride ion 3. calcium ion potassium ion 19 atomic # mass #39.10  39 # p+ =19 19 – (+1) = 18 # e- = #n° = 39 – 19 = 20 Lost to negative ion! 8 18 e- e- 8 2e-e- p+ = 19 n° = 20 K+  atomic #20 calcium ion mass #40.08  40 # p+ = 20 # e- = 20 – (+2) = 18 #n° =40 – 20 = 20 8 18 e- e- 8 2e-e- p+ = 20 n° = 20 Ca2+  17 atomic # chloride ion mass #35.45  35 # p+ =17 # e- =17 – (-1) = 18 #n° =35 – 17 = 18 8 e- 8 18 e- 2e-e- p+ = 17 n° = 18 Cl- How does titanium become an ion?  Titanium gains/loses ___ electrons Tricksy Warm-Up Thinky Question  A sulfur atom gains 3 protons, 2 electrons, and 5 neutrons. What could you call it?  Draw an EELD E. Electron Dot Diagrams  you can’t see atoms and electrons, draw therefore it is models convenient to chemical to show the structure and formation of bonds  an electron dot diagram is one such model  symbol for the element consists of the with dotsrepresenting the valence e  in each energy level, electrons are arranged in pairs in orbitals  when drawing the diagrams, look up the number of valence electrons, then place dots around the symbol clockwise for a maximum of four dots  if you have more electrons to place, go back to the top of the symbol and start pairing up the e     Na Mg  Al   Si             S   Cl    Ar   P         Warm-Up  The elements shown in these electron dot diagrams would be isoelectronic with krypton when they form ions. What are they?  G. Classification of Matter Where are we? Matter Pure Substance Mixtures Elements Compounds Heterogeneous Homogeneous Metals Ionic Colloids Suspension Mechanical Alloys Solutions Metalloid Molecular Mixture s Non-metals I. Elements – So pure! Order Change  metallic elements exist as single atoms (monatomic)  chemical formula is simply symbol the followed by the state at room temperature eg) sodium Na(s) mercury Hg( ) copper Cu(s) nonmetals  (not including noble gases ) do not exist as single atoms and areelements molecular called (diatomic, polyatomic)  chemical formula is the symbol with the subscript state and the at  room temperature memorize the subscripts (flagpole or hockey stick and puck)!!! H2 N2 O2 F2 P4 S8 Cl2 Br2 I2 DID YOU MEMORIZE THIS YET BECAUSE YOU MUST! Summary: Monatomic C(s), noble gases, all metals Diatomic H2(g), N2(g), O2(g), F2(g), Cl2(g), Br2(), I2(s) Polyatomic P4(s), S8(s) H. Properties  physical properties are see and measure properties you can eg) colour,state,boiling point  chemical properties are properties used to describe how react substances with other will elements eg) combustion, rusting,decomposition J. Molecular Compounds molecular compounds are formed when two or more nonmetals bond  together bonded by covalent bonds which is the sharing force of attraction between atoms that are electrons  properties: 1. do not conduct electricity when dissolved in water 2.dissolve in water to form either a neutral molecular solutionacidic solution or an 3.solids, liquids or gases at room temperature Naming Intro Activity Naming (when only two elements combined)  give theatom name first for element the prefix (with the if there more is one than name ) then for givethe thesecond element with “ide” ending prefix and include the  Note: if the first element is hydrogen not, do put a prefix (these are acids!) Prefixes 1 =mono 6 =hexa 2 = di 7 =hepta 3 =tri 8 =octa 4 =tetra 9 =nona 5 =penta 10 =deca Examples 1. CO(g) carbon monoxide 2. CO2(g) carbon dioxide 3. P4O10(s) tetraphosphorus decaoxide bromine heptahydride 4. BrH7(s) Writing Formulas  symbol simply write each followed by the subscript (from prefix) Examples 1. oxygen dibromide OBr2 2. diphosphorus pentasulphide P2S5 3. carbon tetraiodide CI4 4. phosphorus pentachloride PCl5  some molecular compounds have classical names…found on data sheet NH3 = ammonia H2O = water H2S =hydrogen sulphide HF, HCl, HBr, HI =no prefixes CH4 = methane CH3OH = methanol C2H6 = ethane C2H5OH = ethanol C6H12O6 = glucose C12H22O11 = sucrose O3 = ozone H2O2 = hydrogen peroxide Reminders  Safety contracts  Quiz corrections  Contacting me K. Ionic Compounds  ionic compounds are formed when transferred electrons are , allowing oppositely charged ions to bond together  ionic bond is the force of attraction between oppositely charged ions  properties of Ionic Compounds: 1. conduct electricity when dissolved in water 2.separate into ions when dissolved in water 3.crystalline solids at room temperature (non conductive as solids Naming Intro Activity Ionic compounds #1-3 1. Monovalent Ionic Compounds  monovalent means there oneischarge on the metal eg)Na+, Ca2+  metal + nonmetal Naming give the name for each ion…metals are normal, nonmetals “ide” have ending NO PREFIXES! eg) NaF sodium fluoride Na2S sodium sulfide the 2 means that two sodium ions are bonded with one sulphide ion… this doesn’t become a prefix!! Try These: 1. LiF lithium fluoride 2. KCl potassium chloride 3. BeS beryllium sulphide 4. Rb3P rubidium phosphide 5. MgF2 magnesium fluoride dinitrogen dioxide NOT IONIC 6. N2O2 cesium bromide 7. CsBr barium chloride 8. BaCl2 Naming Intro Activity #4 Writing Formulas  symbol for each ion look up the and write them listing metal the ion first balance the charges using subscript numbers ***in ionic compounds, the total positive charges must equal the total negative charges…the net charge is zero eg) sodium oxide Na2 O 1+ 2 1+  2 = 2+ 2  1 = 2  calcium phosphide Ca3 P2 2+ 3 2+  3 = 6+ 3  2 = 6 Try These: 1. magnesium chloride MgCl2 2. calcium chloride CaCl2 3. zinc sulphide ZnS 4. silver sulphide Ag2S 5. germanium oxide GeO2 6. calcium arsenide Ca3As2 7. magnesium nitride Mg3N2 Naming compounds intro activity #5-6  2. Multivalent Ionic Compounds  metal ions that have more than one possible charge are multivalent metals eg)Cu2+, Cu+, Fe3+, Fe2+ multivalent metal + anion  thefirst charge listedmost is thecommon ***NEW STEP** Check periodic table to see if the metal has more than 1 charge! Multivalent 4+ most common Monovalent charge. Naming  same rules as before the difference is you must include brackets charge containing the of the metal ion in Numerals (I,II,III,IV,V,VI,VII) Roman  figure out the charge on the metal by using how many negative charges there are in the nonmetal ions  Nonmetals are NEVER multivalent! eg) CuI copper (I) iodide TiBr4 titanium (IV) bromide Ti3P4 titanium (IV) phosphide Try These: 1. AuBr gold (I) bromide 2. CrCl2 chromium (II) chloride 3. Co2O3 cobalt (III) oxide 4. La2S3 lanthanum sulphide * NOT MULTIVALENT! plutonium (VI) nitride 5. PuN2 Writing Formulas  same rules  the charge on the metal is given to you in brackets the eg) iron (II) oxide FeO tin (II) chloride SnCl2 chromium (III) sulphide Cr2S3 Try These: 1. chromium (II) sulphide CrS NiCl3 2. nickel (III) chloride V3P4 3. vanadium (IV) AuI3 phosphide 4. gold (III) iodide Naming Ionic compounds #7-8  3. Mixed Ionic Compounds cation + polyatomic ion (complex ion) eg) PO43, SO42, HCO3 etc. ***NH4+ (ammonium ion) is the only positive complex ion…you’ll see it in the place of a metal Naming first ion  give the name for the then give the name for theion complex Try These: 1. KIO3 potassium iodate 2. NaCH3COO sodium acetate magnesium sulphite 3. MgSO3 ammonium nitrate 4. NH4NO3 calcium phosphate 5. Ca3(PO4)2 Writing Formulas same as before …look up symbol the for eachbalance ion then the charges  using subscripts if you must multiply one of the complex ions, brackets put around it first then write the subscript eg) aluminum phosphate AlPO4 aluminum chlorate Al(ClO3)3 calcium sulphite CaSO3 scandium acetate Sc(CH3COO)3 (NH4)2SO4 ammonium sulphate Summary: Ionic vs. Molecular Ionic Molecular  Cation + anion  all nonmetals Ionic bonds Covalent bonds  no prefixes  prefixes  Solids at room  solids, liquids or temp gases  solutions do not  solutions conduct conduct  solutions are  solutions are usually usually basic or acidic or neutral neutral Write the name or formula dinitrogen triphosphide SiO2 O2 magnesium nitrate FeCl3 4. Acids and Bases  matter can be subdivided into three groups based on its properties: 1. acids 2. bases 3. neutral substances  let’s look at the properties of acids and bases:  Acids Bases  always soluble in  usually soluble in H O 2  H 2O conduct electricity  conduct electricity  neutralize bases  neutralize acids  taste  taste sour bitter  reacts with metals  feel slippery to produce H2(g) Indicators Indicators   litmus -blue litmus - red  bromothymol blueyellow  bromothymol blue - - blue  phenolphthaleincolourles -  phenolphthalein – s bright pink  vinegar (acetic acid), lemon juice  Tums, ammonia Thinking question If you were trying to find out if something was an acid, and only had one type of litmus paper, which type would you prefer?  the pH scale was devised to indicate acidic how or basic a substance is stomach drain acid coffee water antacid cleane r 0 1 5 7 11 14 strong base strong acid Acids Neutral Bases  Naming and Writing Formulas for Bases  same rules as ionic compounds because most bases Strongare ionic bases hydroxide except contain ion (OH) !!!!ammonia!!! eg) NaOH sodium hydroxide KOH potassium hydroxide Ba(OH)2 barium hydroxide sodium hydrogen *NaHCO3 carbonate Ca(OH)2 calcium hydroxide CaCO3 *calcium carbonate NH3 *ammonia (molecular) Naming  acidsalways hydrogen contain firstaselement (almost always the )  acids are always aqueous (dissolved in water) (aq) eg) HCl(aq), H2SO4(aq), HNO3(s) not an Rules acid!!! 1. hydrogen______ide hydro____ic becomes acid ______ate ______ic 2. hydrogen______ite becomes ______ous acid 3. hydrogen becomes Step 1: Identify the nonmetal ion or polyatomic ion HNO2(aq) NO2- nitrite Step 2: Write ionic name hydrogen nitrite Step 3: Examine ending “ite” Step 4: Use rules: hydrogen nitrite becomes  nitrous acid Examples 1. HF(aq) hydrofluoric acid 2. H3BO3(aq) boric acid not an acid because of state 3. HCl(g) (g) acetic name -acid – (organic hydrogen acids H chloride 4.CH3COOH(aq) is gasat the end) Writing Formulas  use the naming acids rules in the opposite direction  come up with the“ionic” name then write the formula, balance “(aq)” the charges and add to the end eg) hydrogen sulphide H2S(aq) hydrosulphuric acid carbonic acidhydrogen carbonate H2CO3(aq) chlorous acidhydrogen chlorite HClO2(aq) What’s the difference between (aq) and (l)? L. States and Solubility acids – always (aq)  elements – can be (s), (l) or (g)… look up state periodic table on molecular compounds can – be (s), (l), or the question(g) usually tells you (or use common sense!) ionic compounds – either (s) or (aq)… look up solubility on the chart  Predicting Solubility 1. Break compound into its ions 2. Locate anion (-) in top row of table (if the anion isn’t there you can sometimes use the cation (group 1) and ammonium (NH4+)) 3. Look down column and find cation (+) OR GROUP element is in! In “HIGH SOLUBILITY” row  will dissolve and state is (aq) In “LOW SOLUBILITY” row  will form a precipitate and state is (s)  Solubility of Some Common Ionic Compounds in Water at 298.15 K Group 1 ClO3- Cl- PO43- Ion NH4+ NO3- CH3COO Br- SO42- S2- OH- SO32- H3O+ ClO4- - I- CO32- (H+) Solubility all all most most most Group 1 Group 1 Group 1 greater than or Group 2 NH4+ NH4+ equal to 0.1 NH4+ Sr2+ mol/L (very Ba2+ soluble) Tl+ (AQ) Solubility none none Ag+ Ag+ Ca2+ most most most less than or equal to 0.1 Hg+ Pb2+ Sr2+ mol/L Hg+ Ba2+ (slightly soluble) Cu+ Ra2+ Tl+ Ag+ SOLIDS Pb2+ Examples 1. LiCl( aq ) 2. AgCl(s ) 3. NaNO3aq ( ) 4. Ba(OH)2(aq ) s 5. BaSO4( ) aq 6. K2S( ) 7. Extra!! K2O ( ) Beware of Multivalent Metals! Is PbSO4 soluble? Is Pb(SO4)2 soluble? M. Law of Conservation of Matter  Law of Conservation of Matter states that matter cannot be created or destroyed, it only changes forms mass  of reactants = mass of products   when chemicals react they follow the Law of Conservation of Matter  equal numbers there must be of each element on both sides of the reaction  coefficients are used to increase the number of compounds present (balancing) coefficient 3 CaCl2 (aq)  The coefficient is a multiplier that multiplies the number of atoms by the coefficient Ca – 1 X 3 = 3 Ca atoms Cl – 2 X 3 = 6 Cl atoms More importantly the coefficient is the number of moles of a substance in a reaction. Counting Atoms  How many N and O in… NO NO2 2NO2 (NO2)2 3(NO2)2 Counting Practice! How many of each element are in the followin compounds? 1. NaCl 5. NH4CH3COO 2. BaBr2 6. 3 (NH4)2S 3.(NH4)3P 7. 2 CaCl2 4.Ba(OH)2 8. 8 PbI2 9. 4 Zn(CH3COO)2  Writing Chemical Reactions Word Equation – listed chemicals that are in the chemical reaction. You will have to balance each compound in order to balance the chemical equations. hydrogen + oxygen water Reactants “to produce” Products “reacts with”  Reactants Coefficients Products 1 Zn(s) + 2 HCl(aq) 1 ZnCl +1 H2(g) 2(aq) states states Writing Chemical Equations Skeleton Equation: Shows reactants and products, but not in balanced proportions H2 + O2 H2O Writing Chemical Equations Balanced Chemical Equation: Use coefficients in front of chemical formulas so there are the same number of atoms of each element on the reactants and products sides Must be whole #s (not fractions) ___ H2(g) + ___O2(g) ___H2O  Steps to Balance 1. Write skeleton equation from the word equation 2. List atoms present on both sides of equation 3. Count # of atoms on both sides 4. Use Coefficients to balance the number of atoms (so both sides are equal)  Helpful Hints  You CANNOT change any subscripts in chemical formula (can only add coefficients)  Keep atoms in polyatomic ions together if they appear intact on both sides of equation  Coefficients need to be in lowest form  DOUBLE check that your equation is balanced Example 1 ___N2(g) + ___H2(g) ___NH3(g)  Example 2 __Ca(NO3)2 ( ) + Ag ( )  AgNO3( ) + Ca ( )  Example 3 ___(NH4)3(PO4) ( ) + ___Ba(OH)2 ( )  _Ba3(PO4)2 ( ) + NH4OH ( ) Word equation to balanced chemical formula  Solid sodium reacts with liquid bromine to produce solid sodium bromide (have no fear…)  Example 2  Aqueous iron (III) nitrate reacts with aqueous sodium sulfide to produce solid iron (III) sulfide and aqueous sodium nitrate Warm-Up Write the balanced chemical reaction for: -Pure sulfur reacts with pure oxygen to produce sulfur dioxide gas -Sodium hydroxide reacts with phosphoric acid to produce sodium phosphate and water N. Chemical Reactions  chemical reactions can cause both physical and chemicalalways changes and involvenew the formation substanceof a Evidence 1. temperature change 2. colour change 3. solid (precipitate) produced 4. gas produced 5. Odour produced energychanges that occur with chemical reactions can be: 1. endothermic =energy is absorbed (enters) reactants + energy products 2 SO3(g) + 197.8 kJ 2 SO2(g) + O2(g) 2. exothermic = energy isreleased (exits) reactants products + energy Mg(s) + ½ O2(g) MgO(s) + 601.6 kJ O. Identifying Chemical Reactions 1. Hydrocarbon combustion  a compound containing hydrogen and carbon (a hydrocarbon) burns/combusts in the presence of O2(g)  always forms(products are) CO2(g) and H2O(g) C?H? + O2(g) CO2(g) + H2O(g) hydrocarbon eg) CH4(g) + 2 O2(g) CO2(g) + 2 H2O(g) 2. Formation  will always have more reactants than products  are usuallyexothermic  elements compound combine to form a element + element compound eg) 2 Mg(s) + O2(g) 2 MgO(s) 3. Decomposition  will always have more products than reactants  are usuallyendothermic  compounddecomposes into its elements compound element + element eg) 2 H2O(l) 2 H2(g) + O2(g) 4. Single Replacement  an element reacts with an ionic to form a different element and compound a different ionic compound element + compound element + compound eg) Cu(s) + 2 AgNO3(aq) 2 Ag(s) + Cu(NO3)2(aq) Cl2(g) + 2 NaBr(aq) Br2(l) + 2 NaCl(aq)  5. Double Replacement  two ionic compounds react to form two different ionic compounds compound + compound compound + compound eg) Pb(NO3)2(aq) + 2 KI(aq) PbI2(s) + 2 KNO3(aq) General Form of Reaction Types  Example 1 (Warm-Up) What is the reaction type? Complete the rxn ___C4H­10 (g) + _____O2(g)  ____CO2( ) + _____ H2O( )  Example 2 What is the reaction type? Complete the rxn ___ Cu(s) + ____ AgNO3(aq)  ____ Ag( ) + __ Cu(NO3)2( )  Example 3 Predict the reaction type using the general formulas, and identify products, their states and balance the chemical equation ___ Cl2(g) + ____ NaBr(aq)  Example 4 Predict the reaction type using the general formulas, and identify products, their states and balance the chemical equation ____ ZrI4(aq) + ____ Pb(NO3)2(aq)  Practice Write a complete, balanced reaction for the reaction between sulfur and yttrium. P. The Mole  the mass of a single atom is so small that we cannot easily measure it  the mole, n, is a concept that is used so that we can actually measure the mass of NUMBER elements and compounds…IT IS JUST A !!! Mathemagic Tricks How to change units You have 400 muffins. How many dozen muffins is that? Steps: 1.Write the given quantity 2.Multiply by a fraction 3.Numerator is units you want 4.Denominator is units you have You Try You have 65 dozen bags of chocolate chips (for obvious reasons). How many bags do you have? Isn’t this complicated? This method (called unit analysis) works for ALL types of mole calculations. 1. Avogadro’s Number 1 mole Avogadro’s = number (NA) = 6.022  1023 atoms, molecules that’s etc.000 000 000 000 000 000 602 000 atoms, molecules etc!  you can use Avogadro’s number to calculate the number of moles in a substance if you know the number of molecules n=# atoms NA Scientific Notation Number Scientific Notation 450 4.5 x 102 400 4 x 102 4000 4 x 103 5680 5.68 x 103 What is the rule to put a number in scientific notation? Practice Write these numbers in scientific notation 30 000 42 3 065 000 What are these numbers? 6.23 x 104 5.1 x 102 2.5 x 10-2 On your calculator The button for scientific notation usually looks like: E EE Exp To write Avogadro’s number on your calculator: 6.02 E 23 Example A diamond contains 50 atoms of carbon. How many dozen atoms of carbon are in the diamond? A diamond contains 5.0 x 1025 atoms of carbon. How many moles of carbon are in this diamond? How many seats are on 2.50 dozen bikes? How many tires? How many molecules are in 2.50 mol of carbon dioxide? How many oxygen atoms? 2. Atomic Molar Mass  the atomic masses given on the periodic table are an average of all thenaturally occurring isotopes of each element  molar mass is the mass of one mole of a substance  measured ing/mol  the atomic molar masses are given on the periodic table Examples: Calculate the molar mass of the following substances: Na(s) 1 x 22.99 g/mol = 22.99 g/mol O2(g) 2 x 16.00 g/mol = 32.00 g/mol CO2(g) 1 x 12.01 g/mol 12.01 g/mol = 2 x 16.00 g/mol = 32.00 g/mol 44.01 g/mol Al2O3(s) 2 x 26.98 g/mol 53.96 = g/mol 3 x 16.00 g/mol 48.00 = g/mol 101.96 g/mol 40.08 g/mol Ca(OH)2(s) 1 x 40.08 g/mol = 2 x 16.00 g/mol =32.00 g/mol 2 x 1.01 g/mol =2.02 g/mol 74.10 g/mol 3. Mole Calculations  now we can use number of moles and molar mass in a formula: n= m m = nM M where: n = number of moles in mol m = mass in g M = molar mass in g/mol Example 1 How many moles are in 200 g of table salt (NaCl)? m =200 g n = m M = 58.443 M g/mol = 200 g n=? 58.44 g/mol = 3.422… mol = 3.42 mol Example 2 How many grams are in 62.9 mol of lead (II) nitrate? Pb(NO3)2 n = 62.9 mol m = nM M = 331.23 = (62.9 mol)(331.23 g/mol g/mol) m=? = 20 834.367 g = 2.08  104 g Practice How many mol is 25 g of dinitrogen trisulfide? *What is the mass (in g) of a single molecule of dinitrogen trisulfide? To make a bike requires 2 wheels for each frame. So, what happens if you buy 1 ton of bike frames and 2 tons of wheels? Similarly, in chemical reactions, you cannot compare the amount of reactants and products using mass because… Mole Map 2 Li3N(aq) + 3Ca(NO3)2(aq)  1 Ca3N2(aq) + 6 LiNO3(aq) If 1 mol of lithium nitride reacts, how much lithium nitrate will be formed? If 3 mol of calcium nitride are produced, how much lithium nitrate was produced? You Try! For the reaction between iron (III) oxide and magnesium, if 4 mol of iron oxide is used, how many mol of magnesium will be required?