Chemistry Final Review

Science 10 Review CHEMISTRY:  WHMIS – Workplace Hazardous Materials Information System; symbols used to identify dangerous materials (xiii, 8)  Classifying Matter – matter is anything with mass and volume and could be either solid liquid or gas. Matter can be further divided into either a mixture or a pure substance. See page 10.  Atomic Theories –Early Chemists devised theories about the structure of an “atom”. o Dalton’s Atomic Theory (12):  All matter is made up of small particles called atoms  Atoms cannot be created, destroyed or divided into smaller parts  All atoms of the same element have the same mass but different elements have different masses  Put two elements together and you get a compound  Chemical reactions change the way atoms are grouped but atoms themselves do not change  Looks like a billiard ball (solid, uniform, sphere) o J.J. Thomson (15):  Thought atoms might be made up of smaller particles  Proved that negatively charged electrons were part of an atom  Viewed atom as a “raisin bun” o Rutherford (16):  Using alpha particle streams showed that positively charged protons and neutral neutrons existed within an atom’s nucleus.  A volume of empty space surrounded the nucleus 3 e Atomic model looks like a “solar system” o Bohr (19): 8 e Thought electrons were associated with certain energy levels Energy level Maximum # of electrons 2 e1 2 2 8 3 8 13 p+  What element is the drawing of? ______Al o Working Model of the Atom (22): Subatomic particle Proton Neutron electron  Charge Symbol 1+ 0 1- p+ n0 e- We say this particle has a mass of 0 2p+ 2n Nuclear Notation (22) using the periodic table o Atomic Number = number of protons o Mass Number = number of protons + number of neutrons o Number of neutrons = mass # - atomic #  How many neutrons in magnesium? ______12 Each proton has a mass of 1 so these protons add up to 2 Each neutron has a mass of 1 so these neutrons add up to 2 12 atomic number Mg 24.31 atomic mass  Periodic Table (25): o Organizes elements to help us predict their nonmetals properties metals o Three major sections:  Metals - left side  Non-metals – right side  Metalloids – border the “staircase line” that separates metals and non-metals  Read about each sections properties in the text book. o Elements arranged into periods (horizontal rows) and groups (vertical columns)  Periods – elements are in order of their atomic number, each period has the same number of energy levels for electrons. Eg. Elements with 3 energy levels (more than 10 electrons) are found in period (row) 3.  Groups – elements have similar properties as those found above and below them. Sometimes called “families” eg. Group 1 is the Alkali Metals (look on page 25 for the name of the rest). Main groups have the same number of valence electrons (electrons on the outer most shell).  Valence Electrons – If the outer shell of an atom is full this is called a “stable octet”. Atoms want their valence shell to be full so they gain or lose electrons.  Once the atom no longer has the same number of electrons as protons the atom has a charge. It is now called an ion instead. o Formation of Ions (29):  Positive Ions are called CATIONS – they have lost electrons (metals do this)  Negative Ions are called ANIONS – they have gained electrons (non-metals do this)  Ions form when atoms collide and their valence electrons interact. Since they both want to have full outer electrons shells like the nearest noble gas they negotiate electrons.  Compounds made of Ions are called Ionic Compounds – Positive + Negative o Held together by ionic bonds forming a crystal lattice o ELECTRICALLY NEUTRAL – the compound has no net charge!! o Naming Ions (44)  The 1st element in the name and the formula is the metal  The 2nd element, the non-metal named as an ion (add suffix “-ide” ) fluorine = fluoride  Eg. LiCl ____lithium chloride__________ KBr ____pottasium bromide___________ CaCl2 _____calcium chloride___________ o Formulas (45):  Use ion charges from you periodic table to decide the charge of each  Decide how many of each ion is required to create a neutral compound  If more than one, denote the number with a subscript beside the element symbol  Eg. Potassium Oxide = K+ and O2- we need two potassium to make neutral = K2O (s)  Eg. Potassium sulfide __________ (K2S) Beryllium fluoride ___________ (BeF2) o Multivalent Elements -The Stock System (46):  Some metals have more than one cation. The stock system determines the charge to use.  Cation ion charge is written in brackets as a roman numeral after the metal name  Eg. Ni2+ is nickel (II) and Ni3+ is nickel (III) o Polyatomic Ions (51):    A group of different atoms joined by a covalent bond with an overall charge or + or -. You will find the names, formulas and charges of these on the back of your periodic table To name use the name of the cation followed by the name of the anion – no suffix change  Eg. NO3- is called nitrate. It has an overall charge of 1-. When bonded with Zn2+ we get – Zn (NO3)2 or zinc nitrate. o Acids – read about the properties of acids on pages 63-66  An acid contains hydrogen as the cation (the first element in a formula)  Name first as an ionic compound and then apply the acid rules depending on the result  hydrogen ____ide = hydro____ic acid HCl = hydrogen chloride = hydrochloric acid  hydrogen_____ate = _________ic acid HClO3 = hydrogen chlorate = chloric acid  hydrogen_____ite = _______ous acid HClO2 = hydrogen chlorite = chlorous acid   Molecular Compounds (31):  Groups of atoms with no charge called molecules – not ions  Contains only non-metal atoms  Atoms of the same element can from bonds so elements can exist as molecules  Diatomic molecules (aka the Special Seven –MEMORIZE!!) O2(g) - H2(g) - N2(g) - F2(g) - Cl2(g) - Br2(g) - I2(g) & P4(s) - S8(s) o o Elements joined by covalent bonds – sharing electrons not exchanging them o Naming Molecular Compounds (43): 2 or more non-metals together  First element in the name is the one farther to the left on the periodic table  The suffix “-ide” is added to the name of the second element  Prefixes are used to tell how many atoms are of each type are in the molecule (prefixes can be found on the back of the periodic table) * mono is used only for the 2nd element *  Eg. CO2 = carbon dioxide, CCl4 = carbon tetrachloride  Some molecular compounds do not follow simple naming rules – MEMORIZE them H2O(l) water H2O2(1) hydrogen peroxide NH3(g) ammonia C12H22O11(s) sucrose CH4(g) methane C3H8(g) propane CH3OH(l) methanol C2H5OH(l) ethanol C6H12O6(s) glucose O3(g) ozone o Molecular and Ionic compounds have specific properties. Refer to section 2.2 of text. o For a summary of naming ionic and molecular compounds turn to table 2.6 on page 54 Chemical Reactions (85) o Occurs when one or more substances change to form different substances o Substances that undergo change are called reactants substance that result are called products  Reactants → Products: all chemical reactions involve a change in energy o Evidence a chemical reaction has occurred may include:  Energy change  Formation of Gas  Odor Change  Formation of a solid (Precipitate) in solution  Color Change o A precipitate is an insoluble solid formed out of solution. (88) We can predict if a cation and an anion will mix to form a precipitate using the Solubility chart on the back of your periodic table  High Solubility = (aq)  Low Solubility = (s)  Everything is soluble if NH4+, NO3-, or Group 1 is involved o Reactions release or absorb energy (90)  Release energy = exothermic reaction eg. Combustion of gasoline  Require energy = endothermic reaction eg. Photosynthesis needs the sun’s energy  The Law of Conservation of Energy states that energy cannot be create or destroyed only transferred. Breaking chemical bond = endothermic; forming new bonds = exothermic. o Writing Reactions (94) Reactants on left side of arrow o         Arrow means “produce” Zn(s) + AgNO3(aq) Zn(NO3)2(aq) + Ag(s) Products on right side of arrow Plus sign means “reacts with” o Balanced Equations (97):  There needs to be the same number of atoms on one side of the equation as the other  Called the Law of Conservation of Mass. We use Coefficients (big numbers in front of formulas) to balance atoms on either side. 1. Write out formulas for all compounds and elements, add arrow and plus signs if needed 2. Identify unbalanced atoms and polyatomic ions & add coefficients where necessary. 3. Check balancing at the end ___Li(s) + ___HOH(l) → ___LiOH(aq) + ___H2g) – Hydrogen appears by itself, balance that first ___Li(s) + _2_HOH(l) → ___LiOH(aq) + ___H2g) – Now there are two hydroxide molecules, balance LiOH ___Li(s) + _2_HOH(l) → _2_LiOH(aq) + ___H2g) – Now there are two Li atoms, balance Li (s) next _2_Li(s) + _2_HOH(l) → _2_LiOH(aq) + ___H2g) - Check and we’re done Types of Reactions (103) o Reactions are classified to help predict the products  Formation– two or more reactants combine to make a new compound (product) A+B→AB  Decomposition– a compound breaks into simpler compounds or elements  AB→A+B  Single Replacement – one element replaces another in a compound  A + BX → AX + B  Double Replacement – cations & anions exchange partners to form new compounds  AX + BY → AY + BX  These reactions may form a precipitate so check solubility chart  Combustion- adding oxygen to form most common oxide  X + O2 → XO?  Often involves hydrocarbons (compounds containing hydrogen and carbon)  Balance carbon first, then hydrogen and balance oxygen last. The Mole (116) o Avagadro’s number is 6.02 x 1023 – called a Mole (symbol mol) o The mass of 1 mol of all the isotope of a substance is called its atomic molar mass and is listed on the periodic table for each element. Eg. Iron – 55.85g, Sodium 22.99g o The molar mass (M) has the units g/mol is the mass of 1 mol of any pure substance  Molar mass of compound is found using formula – CO2 = 1(12.01g/mol)+ 2(16.00g/mol)  44.01 g/mol Converting between mass and molar mass (121) o n=m/M o n = amount (mol) o m = mass (g) o M = molar mass (g/mol) Coefficients of a balanced equation refer to the number of moles of each atom, molecule or formula unit (124) Atoms and Ions Review: Name Symbol # protons # electrons # neutrons Charge arsenic atom As 33 33 42 0 Chlorine atom Cl 17 17 18 0 Antimony ion Sb5+ 51 46 71 5+ Xenon atom Xe 54 54 77 0 Magnesium ion Mg2+ 12 10 12 2+ hydride H 1 2 1 -1 S2- 16 18 16 2- Sulphide ion Aluminum ion Al3+ 13 10 14 3+ Nitride ion N3- 7 10 7 3- - Naming Review - Use subscripts to indicate the state at room temperature 1. I, M, Name or A I sodium chloride Chemical Formula NaCl 2. I Calcium carbonate CaCO3 4. I sodium hydroxide NaOH 5. I Calcium oxide CaO 6. I Magnesium sulfate hepta hydrate MgSO47H2O 7. M carbon dioxide CO2 8. A acetic acid CH3COOH 9. M carbon C 10. I calcium sulfate CaSO4 11. I Sodium silicate Na2SiO3 12. I Calcium hydrogen carbonate Ca(HCO3)2 13. I magnesium hydroxide MgOH2 14. I potassium chloride KCl 15. I sodium thiosulfate pentahydrate Na2S2O3∙5H2O 16. I sodium hypochlorite NaClO 17. I Sodium carbonate Na2CO3 18. A Hydrochloric acid HCl(aq) 19. I potassium nitrate KNO3 20. I Copper sulfate penta hydrate CuSO45H2O magnesium oxide MgO 21. 22. I Potassium iodide KI 23. A Sulfuric acid H2SO4(aq) Calcium hydroxide Ca(OH)2 24. Balance the following equations. 1. ___K(s) + ___ Cl2(g)  ___ KCl(s) 2. ___Fe(s) + ___S8(s)  ___FeS(s) 3. ___H2O(l)  ___H2(g) + ___O2(g) 4. ___NaCl(s)  ___Na(s) + ___Cl2(g) 5. ___AsCl3(aq) + ___H2S(aq)  ___As2S3(s) + ___HCl(aq) 6. ___CuSO4 5H2O(s)  ___CuSO4(s) + ___H2O(aq) 7. ___Na(s) + ___O2(g)  ___Na2O(s) 8. ___H2S(aq) + ___KOH(aq)  ___HOH(l) + ___K2S(aq) 9. ___Fe(s) + ___H2O(g)  ___H2(g) + ___Fe3O4(s) 10.___Al(s) + ___H2SO4(aq)  ___H2(g) + ___Al2(SO4)3(aq) 1) 2) 3) 4) 5) 6) 7) 8) 9) 2,1,2 8,1,8 2,2,1 2,2,1 2,3,1,6 1,1,5 4,1,2 1,2,2,1 3,4,4,1 10) 2,3,3,1 11) 1,3,1,3 12) 1,2,2,1,1 13) 2,3,3,1 14) 1,2,2,1 15) 2,3,2,2 16) 1,5,3,4 17) 1,2,1,2 18) 1,3,3,1 11.___AlCl3(aq) + ___NaOH(aq)  ___Al(OH)3(s) + ___NaCl(aq) 12.___Na2CO3(aq) + ___HCl(aq)  ___NaCl(aq) + ___H2O(l) + ___CO2(g) 13.___Fe(s) + ___CuSO4(aq)  ___Cu(s) + ___Fe2(SO4)3(aq) 14.___H2SO4(aq) + ___KOH(aq)  ___HOH(l) + ___K2SO4(aq) 15.___ZnS(s) + ___O2(g)  ___ZnO(s) + ___SO2(g) Balance the following two reactions doing C, then H then O last. 16.___C3H8(g) + ___O2(g)  ___CO2(g) + ___H2O(g) 17.___CH4(g) + ___O2(g)  ___CO2(g) + ___H2O(g) 18.___H3PO4(aq) + ___NH4OH(aq)  ___HOH(l) + ___(NH4)3PO4(aq) Write and Balance the following Word Problems: 19. Exothermic reaction between aqueous sodium hydroxide and sulphuric acid. 2 NaOH(aq) + H2SO4 (aq)  Na2SO4 (aq) + 2 HOH (l) 20. Formation of a precipitate by reacting sodium carbonate and calcium chloride. Na2CO3 (aq) + CaCl2  2NaCl (aq) + CaCO3 (s) 21. Adding aqueous chlorine to aqueous sodium bromide. Cl2 (aq) + 2NaBr (aq)  2NaCl (aq) + Br2 (l) 22. Adding a drop of lead (II) nitrate to a sliver of zinc metal Pb(NO3)2 (aq) + Zn(s)  Zn(NO3)2 (aq) + Pb(s) 23. Adding a drop of tin nitrate to a sliver of zinc metal Sn(NO3)4 (aq) + 2 Zn(s)  2 Zn(NO3)2 (aq) + Sn(s) 24.There was a terrible accident in the lab! Someone took all the labels off and you need to help Ms. Williams re-label. The following is data collected from the unlabeled chemicals: Unknown A B C D E F State at SATP Solid Liquid Solid Liquid Solid Solid Solubility High High Low High Blue Litmus Test Blue Red Blue Blue Blue Blue Red Litmus Test Blue Red Red Red Red Red Conductivity High High High None High None Using your knowledge of ionic, molecular, acid and base properties and the data above, fill in the chart below and identify the unknown compounds. You have the following chemicals to choose from: C6H12O6, Ba(OH)2, CaCO3, HF, LiCl, CH3OH, Unknown Name A Barium hydroxide Formula (including state in the presence of water) Ba(OH)2 (aq) B Hydrofluoric acid HF(aq) C Lithium chloride LiCl (aq) D methanol CH3OH(l) E Calcium carbonate CaCO3 (s) F glucose C6H12O6 (s) or (aq)