Chemistry Booklet 2017 PDF

Unit-1 Energy and Matter in Chemical Change 1  Science 10 Unit A: Energy and Matter in Chemical Change (Nature of Science Emphasis) Key Concepts an...

Unit-1 Energy and Matter in Chemical Change 1  Science 10 Unit A: Energy and Matter in Chemical Change (Nature of Science Emphasis) Key Concepts and Learning Goals Focusing Questions: How has knowledge of the structure of matter led to other scientific advancements? How do elements combine? Can these combinations be classified and the products be predicted and quantified? Why do scientists classify chemical change, follow guidelines for nomenclature and represent chemical change with equations? Key Concepts:  how chemical substances meet human needs  evidence of chemical change  Workplace Hazardous Materials Information System  role and need for classification of chemical change (WHMIS) and safe practices  writing and balancing equations  International Union of Pure and Applied Chemistry  law of conservation of mass and the mole concept (IUPAC) nomenclature, ionic and molecular  Be able to identify responding, manipulated and compounds, acids and bases major controlling variables in an experiment Learning Goal 1: Describe the basic particles that make up matter, and investigate related technologies I Can: Knowledge level My Plan a) Outline the role of evidence in the 1-2-3-4-5-6-7-8-9- development of the atomic model 10 consisting of protons and neutrons (nucleons) and electrons i.e., Dalton, Thomson, Rutherford, Bohr b) Identify examples of chemistry-based 1-2-3-4-5-6-7-8-9- careers in the community (e.g., chemical 10 engineering, cosmetology, food processing) Learning Goal 2: Explain, using the periodic table, how elements combine to form compounds, and follow IUPAC guidelines for naming ionic compounds and simple molecular compounds I Can: Knowledge level My Plan a) Describe, define and illustrate WHMIS 1-2-3-4-5-6-7-8-9- guidelines, and symbols. Demonstrate 10 safe practices in the handling, storage and disposal of chemicals in the laboratory and at home b) Explain the importance of the IUPAC 1-2-3-4-5-6-7-8-9- system of naming compounds so that 10 scientists can communicate clearly and precisely c) Explain how and why elements combine 1-2-3-4-5-6-7-8-9- to form compounds in specific ratios 10 d) Predict formulas and write names for 1-2-3-4-5-6-7-8-9- ionic and molecular compounds and 10 common acids using IUPAC rules e) Classify ionic and molecular compounds, 1-2-3-4-5-6-7-8-9- acids and bases on the basis of their 10 properties; i.e., conductivity, pH, solubility, state f) Predict whether an ionic compound is 1-2-3-4-5-6-7-8-9- relatively soluble in water, using a 10 solubility chart g) Outline the issues related to personal and 1-2-3-4-5-6-7-8-9- societal use of potentially toxic or 10 hazardous compounds 2 Learning Goal 3: Identify and classify chemical changes, and write word and balanced chemical equations for significant chemical reactions, as applications of Lavoisier’s law of conservation of mass You Will: Knowledge level My Plan a) Provide examples of household, 1-2-3-4-5-6-7-8-9- commercial and industrial processes that 10 use chemical reactions to produce useful substances and energy b) Identify chemical reactions that are 1-2-3-4-5-6-7-8-9- significant in societies (e.g., reactions 10 that maintain living systems, such as photosynthesis and respiration; reactions that have an impact on the environment, such as combustion reactions and decomposition of waste materials) c) Describe the evidence for chemical 1-2-3-4-5-6-7-8-9- changes; i.e., energy change, formation of 10 a gas or precipitate, colour or odour change, change in temperature d) Differentiate between endothermic and 1-2-3-4-5-6-7-8-9- exothermic chemical reactions 10 e) Classify and identify categories of 1-2-3-4-5-6-7-8-9- chemical reactions; i.e., formation 10 (synthesis), decomposition, hydrocarbon combustion, single replacement, double replacement f) Translate word equations to balanced 1-2-3-4-5-6-7-8-9- chemical equations and translate balanced 10 chemical equations to word equations for chemical reactions that occur in living and nonliving systems g) Predict the products of formation 1-2-3-4-5-6-7-8-9- (synthesis) and decomposition, single and 10 double replacement, and hydrocarbon combustion chemical reactions, when given the reactants h) Define the mole as the amount of an 1-2-3-4-5-6-7-8-9- element containing 6.02 × 1023 atoms 10 (Avogadro’s number) and apply the concept to calculate quantities of substances made of other chemical species (e.g., determine the quantity of water that contains 6.02 × 1023 molecules of H2O) i) Interpret balanced chemical equations in 1-2-3-4-5-6-7-8-9- terms of moles of chemical species, and 10 relate the mole concept to the law of conservation of mass Other resources for Chemistry include: Vocabulary & Definitions for Unit: 3 Adapted from: Government of Alberta, Alberta Education Learner Assessment Branch, Copyright 2005 Unit 1 Vocabulary By the end of the unit you should be able to define all of the following terms in your own words. KEY TERMS 1) WHMIS 2) Polyatomic ion 3) Matter 4) Acid 5) Mixture 6) Base 7) Pure substance 8) Acid-Base Indicator 9) Element 10) pH Scale 11) compound 12) Hydrogen bond 13) heterogeneous mixture 14) Intermolecular force 15) homogeneous mixture 16) Intramolecular force 17) Nucleus 18) Chemical Reaction 19) Electron 20) Chemical Change 21) Proton 22) Reactants 23) Neutron 24) Products 25) Isotopes 26) Precipitate 27) Atomic number 28) Solubility 29) mass number 30) Law of Conservation of Energy 31) Periods 32) Endothermic 33) Groups 34) Exothermic 35) Valence electrons 36) Law of Conservation of Mass 37) electron dot diagram 38) Mole 39) Ion 40) Avogadro’s number 41) Cation 42) Atomic molar mass 43) Anion 44) Molar mass 45) ionic compound 46) molecular compound 47) Stable Octet 48) Valence energy level 49) Crystal lattice 50) Ionic bond 51) Covalent bond 4  Lesson 1: Understanding Matter Chemistry is the study of matter and the changes it undergoes. Matter: anything that has mass & takes up space Physical Properties: cause a physical change (colour, density, melting point, boiling point, state of matter) Chemical Properties: describe how one substance reacts with another A chemical property can only be identified once a substance has undergone a chemical reaction. Chemical Reaction: where matter changes to produce new substance(s) with properties different from those original materials. Elements: 1 type of atom, cannot be broken down into simpler substances Compounds: Made of 2 or more elements chemically united, Can be separated chemically into simpler elements Mechanical Mixture (Heterogeneous): Different components of the mixture are visible; composition is variable throughout the mixture Solution (Homogeneous): Different components are not visible; Composition is constant throughout the mixture Techniques to Separate Matter Mechanical: one or more components are picked out of the mixture either manually or by use of a magnet for magnetic substances. Settling: some heterogeneous mixtures can be separated by letting one of the components settle to the bottom. Spinning the mixture at a high speed (centrifuging) may be used to accelerate this process. Flotation: oil, detergents, or other chemicals are added to the heterogeneous mixture and air is blown through. The froth containing the desired component floats, and is skimmed off the surface. This technique is used to concentrate ores of zinc, copper, nickel, and lead, and to separate bitumen (tar) from sand. Filtration: a heterogeneous mixture, usually a solid in a liquid or gas, is passed through a screen or filter. The solid is trapped and separated from the liquid or gas. Extraction: the mixture is mixed with a solvent that dissolves one or more, but not all components. For example a coffee maker uses hot water to extract some of the components from ground coffee beans. Crystallization: a dissolved solid is separated from a solution by cooling or concentrating the solution to crystallize the solid. Fractional Distillation: a liquid mixture is boiled and one or more components are separated as they vaporize from the mixture. 5 Chromatography: a mixture is carried by a solvent through a stationary, porous medium such as a column of solids or a filter paper. Separation occurs because components of the mixture move at different rates in the porous medium. Assignment: 1. Answer the following question using pages 6 and 7 in your textbook. Identify historical examples of how humans worked with chemical substances to meet their basic needs. 6 2. Complete the Chart using page 10 in your textbook 3. Complete Practice Problems p10 #1-4 Lesson 2: Developing Atomic Theories (page 12-21) Using the information in the textbook complete the Theories of the Atom Assignment (on blue paper handed out in class). You can complete the information in the booklet as well, but you do not have to. Dalton’s Atomic Theory ( - )      J.J Thomson ( - ) What did he find? __________________________________________________________________________________________________ _________________________________________________________________________________________________ Draw and label the Thomsom model “ or “ 7 Rutherfords Atomic Model ( - ) “_________-_________” Nucleus- Electrons- Empty Space- Neutron- Isotpes- Niel Bohr ( - ) Energy Levels- Draw a diagram of Bohr’s atom: Energy Level Maximum # of electrons Atomic Theory Evolves: __________________________________________________________________________________________________ __________________________________________________________________________________________________ _______________________________________________ Working Model of the Atom (p22-23) The word atom comes from the Greek word meaning “uncuttable”. Through many years of experimentation, scientists have developed a model of the atom that we use today. We call this model the Modern Atomic Theory, which can be summarized in 3 main points: 1. Atoms consist of 3 subatomic particles: Particle Symbol Charge Relative Mass Location 8  Atoms of different elements have different numbers of _______________, _______________, and ________________. 2. The size of the nucleus is ________________ compared to the size of the atom, but the mass of the nucleus is __________________ to the mass of the entire atom. 3. An atom is electrically ____________ because the number of electrons is equal to the number of protons. Lesson 3: Periodic Table and Subatomic Particles The Periodic Table  The periodic table was invented by Dmitri Mendeleev in 1869  Mendeleev was able to create the entire Periodic Table based on the patterns he saw in only 56 known elements  He left gaps in the Periodic Table that were filled in as new elements were discovered Classifying the periodic table  Groups (Families):  Vertical columns  Elements in each group have similar  18 Groups on the Periodic Table chemical properties  Periods:  Horizontal rows  7 Periods on the Periodic Table Metals -are good conductors of electricity -form positive ions (charged particle) -are good conductors of heat -give away electrons -are solids at room temperature, except mercury 9 Non-Metals -are poor conductors of heat and electricity -form negative ions -depending on the element, could be at any state at room -accept electrons temperature Semi-Metals or Metalloids (Metals and nonmetals) -are elements that have properties that fall between metals and nonmetals -may or may not form ions -staircase elements Names of families that are expected to be known: Alkali Metals – Li, Na, K, Rb, Cs, Fr  Group 1  Very reactive  Silver coloured  React violently with water  Reactivity increases as you go down the group Alkaline Earth Metals – Be, Mg, Ca, Sr, Ba, Ra  Group 2  Quite reactive  React with oxygen to form oxides Transition Metals – families 3 through 12  Groups 3-12  Common metals  Contain the “coinage” metals Lanthanides  Period 6  Starts with the element Lanthanum Actinides  Period 7  Starts with the element Actinium Halogens – F, Cl, Br, I, At  Group 17 (VII)  Reactivity decreases as you go down the group  Solids, liquids or gases  Reactive with metal elements to form “salts”  Extremely reactive  React with hydrogen to form acids Nobel Gases – He, Ne, Ar, Kr, Xe, Rn  Group 18 (VIII)  Very low reactivity  All are colourless gases ***Note***: All elements above #93 are synthetic (ie. Can only be formed in a lab for a very short amount of time) What is an atom?  the building blocks of all substances  can be broken down into o Proton  located in the nucleus of the atom and have a positive charge.  the number of protons in an element is the same as its atomic number, and determines its properties o Electron  located outside the nucleus and have a negative charge  electrons can be gained, lost or shared 10  o Neutron  located in the nucleus but does not carry an electrical charge  helps stabilize the structure of an atom Particle Symbol Charge Mass (amu) Mass (g) Location Proton 1.67 x 10-24 Nucleus Neutron 1.67 x 10-24 Nucleus Electron 9.11 x 10-28 Space outside the Nucleus ***Note***: Since the mass of an electron is so small compared with a proton we consider it to be zero  An ATOM is always neutral in charge (protons=electrons)  Atomic Number: o Tells you the number of Protons found in an atom o Tells you the number of Electrons found in a neutral atom o ***NOTE***: Protons give an element its Personality AND Electrons give an element its Mood o All three subatomic particles have mass, but because electrons are so small we disregard the mass of electrons and assume: o # of protons +# of neutrons = mass number (the mass of one atom)  Proton = 1 atomic mass unit (AMU)  Neutron = 1 atomic mass unit (AMU) Nuclear or Isotope NotationExample: lithium Lithium Atomic # = _______________ # of protons = _____________ # of neutrons = ____________ # of electrons = ____________ Atomic mass = ______________  Atomic Mass:  The average mass of all of the isotopes of a particular atom  Reflects the percentage of abundance of each isotope as they are found in nature ISOTOPES  Atoms of the same element that contain the same number of Protons but different numbers of Neutrons  Nuclear notation is used to represent different isotopes 11   Example: Carbon can exist with masses of 12.00, 13.00 and 14.00 AMU. The isotopes of carbon can be represented by: Isotope Name Isotope # of Protons # of Neutrons Uses Notation Carbon-12 - Most common - Most abundant - Stable Carbon-13 - Used in MRI - Stable Carbon-14 - Radioactive - Carbon dating The isotopes differ only in the number of neutrons. All the isotopes of carbon have 6 protons. On your periodic table, the atomic mass that is listed is the average of all the isotopes. Assignment: Periodic Table (p25-29) Groups: Periods: Location in Classification Properties State Examples Periodic Table Metals Non-Metals 12  Metalloids What separates metals from non-metals on the periodic table? 13 Shade in the following chemical families on the periodic table below. Patterns of Electron Arrangement in Periods The periodic table represents patterns related to __________________________. What 2 questions can the periodic table help you answer? 1. 2. Name of Family Group # Elements in Group Properties 1 2 17 14  18 We can determine how many of each subatomic particle are present in an atom by the use of 2 other numbers that are unique to each atom: 1. Atomic Number: 2. Mass Number: number of neutrons =_____________-______________ Practice Problems p23 #5-8 Complete the following chart, using your periodic table and the information given: Atomic Mass # of # of Name Symbol # of Neutrons Number Number Protons Electrons Li 4 48 22 Gold 197 50 69 N 7 Hg 201 Lead 125 238 92 74 110 25 55 Potassium 20 31 15 15  10 10 35 45 27 13 Co 32 56 26 Carbon 6 127 53 Lesson 4: Drawing Atoms Atoms Atoms are stylistically drawn using three different diagrams: 1. Bohr Diagrams 2. Energy Level Diagrams 3. Lewis (Electron) Dot Diagrams Drawing Bohr Diagrams: 1. Find the number of Protons of your given element based on its Atomic Number 2. Find the number of Neutrons of your given element based on its Atomic Number and Mass Number 3. Find the number of Electrons of your given element based on its Atomic Number 4. Draw a circle representing the Nucleus of the Atom 5. Fill the Nucleus with the Number of Protons and Neutrons 6. Start to fill in the energy levels (orbitals) with the appropriate number of electrons 7. Always fill the lowest orbitals first ensuring that electrons are paired up ***NOTE***: Period # = Number of Orbitals Group # = Number of Valence Electrons (the electrons filling the outermost Orbital) Examples: oxygen atom lithium atom Drawing Energy Level Diagrams: 1. Find the number of Protons of your given element based on its Atomic Number 2. Find the number of Neutrons of your given element based on its Atomic Number and Mass Number 3. Find the number of Electrons of your given element based on its Atomic Number 16  4. Draw a circle representing the Nucleus of the Atom 5. Fill the Nucleus with the Number of Protons and Neutrons 6. Start to fill in the energy levels (orbitals) with the appropriate number of electrons 7. Always fill the lowest orbitals first ensuring that electrons are paired up ***NOTE***: Period # = Number of Orbitals Group # = Number of Valence Electrons (the electrons filling the outermost Orbital) Examples: magnesium atom carbon atom Drawing Electron Dot Diagram (Also known as Lewis Dot Diagrams) 1. Find the number of Valence Electrons of your given element based on its Group Number 2. Draw the chemical symbol for the atom 3. Draw the number of Valence electrons around the chemical symbol as dots ***NOTE***: Electrons should be drawn as pairs Examples: chlorine atom oxygen atom Assignment Using Electron Dot Diagrams to Represent Valence Electrons 1. Complete the following table. Number of Number of valence Name of element Period number Group number energy levels e- hydrogen 1 3 3 2 6 strontium 14 3 6 2 2. Complete the following table by drawing both the energy level diagram and electron dot diagram for each element. The first row is completed as an example. 17  Name of element Energy level diagram Electron dot diagram carbon oxygen lithium chlorine magnesium phosphorus 3. Draw the missing electron dot diagrams in the following table. Refer to the periodic table if necessary. 18  4. What feature of helium’s energy levels justifies placing its two valence electrons in a pair? (See the table above.) 5. Why do we use ELECTRON DOT DIAGRAMS? Lesson 5: Ions and Drawing Ions An ion is a charged particle (atom). Ions can either be positively charged (+) or negatively charged (-) This charge is due to a:  Loss of electrons (the atom becomes a positive (+) ion)  Gain of electrons (the atom becomes a negative (-) ion) Ions form when atoms join to form compounds. Atoms like to form a Stable Octet in their outer orbital when becoming an ion. (ie. 8 electrons filling the outer orbital as this is more stable) Atoms can accomplish this by either:  Giving away 1, 2, or 3 electrons to another atom  Accepting 1, 2, or 3 electrons from another atom Anions:  An anion is an ion with a negative charge (ie. The atom has gained electrons to become an ion)  Anions are always Non-Metals  Anions are always formed from atoms that have 5, 6 or 7 valence electrons  Anions will form when a Non-Metal gains electrons Cations:  A cation is an ion with a positive charge (ie. The atom has lost electrons to become an ion)  Cations are always Metals  Cations are always formed from atoms that have 1, 2 or 3 valence electrons  Cations will form when a Metal loses electrons Naming Ions Anions: When Non-Metal atoms become Anions the end of their name changes to “-ide”. If the anion is a polyatomic ion, name it from the table of polyatomic ions. Ex. chlorine becomes a “chloride” ion fluorine becomes a “fluoride” ion Cations: When Metal atoms become cations the name remains the same. If a polyatomic ion is the cation, name it from the table of polyatomic ions. Ex. magnesuim becomes a “magnesium” ion sodium becomes a “sodium” ion Drawing Diagrams for ions: Bohr Diagram: sodium ion calcium ion fluoride ion 19 Energy Level Diagram: Lewis Dot Diagram: sulfide ion magnesium ion nitride ion bromide ion barium ion Assignment: Each energy level of the periodic table can hold a certain number of _________. First level can hold a maximum ______ electrons. Second level can hold a maximum _____ electrons. Third level can hold a maximum _____ electrons. What is a “STABLE OCTET”? period # of an element= Patterns of Electron Arrangements in Groups What is valence energy? and valence electrons? For groups 1, 2 and 13-18 the ones place of group number will always equal the number of ___________ electrons. Do practice problems 9-12 on page 27. The Formation of Ions An ion is a ___________________ in which the number of protons does not equal the number of ________________. Ions have _______________ or _______________ charges associated with them called ionic charges. Atoms can form ions by either ____________ or _____________ electrons. WHY DOES THIS HAPPEN? 20 Metals tend to have 1, 2, or 3 electrons in their outermost energy level. They tend to ____________ these electrons to form ____________ ions. The names of ions of metal elements are _______________________________. Non-metals tend to have 5, 6, or 7 electrons in their outermost energy level. They tend to ___________ electrons to fill the energy level, and therefore form ions with a ____________ charge. When non-metals form ions, the name of the ion changes its ending to “_________”. Questions: 1. What kind of arrangement of electrons in the outer energy level does a stable ion have? 2. Draw Bohr diagrams for the stable ion formed by each of the following atoms: a) lithium b) fluorine c) sulfur d) calcium 3. Atoms and ions are described as isoelectronic if they have the same number of electrons. Name the noble gas that is isoelectronic with the following stable ions: a) Li+ b) F- c) Ca2+ d) S2- e) Br- f) Rb+ 4. Predict the names and charges of the ions that cesium, barium, and calcium might form. 5. Complete the following table by filling in the missing information about ions. Number of Name of Ion Symbol Number of Electrons Ion Charge Protons lithium ion 19 18 1+ Mg2+ chloride ion 9 1– O2– I– scandium ion 18 2– Se2– 7 10 21  Al3+ 10 4– calcium ion phosphide ion Lesson 6: Chemical Compounds and Ionic Compounds: Atoms form compounds so it may acquire a valence energy level like that of the closet noble gas. List the 3 ways this can occur: 1. 2. 3. There are two types of compounds we will be looking at: 1. Ionic Compounds- Composed of _______________. Binary Ionic Compound are made of ______ elements. They form when a ____________ reacts with a ____________. When the 2 elements react _____________ electrons from the metal are ______________ to the non-metal. This forms an _____________ bond. These bonds are very ____________ because they result from strong forces of _____________ between ____________ charged ions. What is a CRYSYTAL LATTICE? How does it relate to IONIC COMPOUNDS? 2. Molecular (Covalent) Compounds Made up of two or more non-metals that form molecules by ____________ electrons. Atoms in molecules are joined by __________ bonds. These bonds are different from __________ bonds. Neither atom wants to give up its ______________. Therefore there is no electron ___________. Instead they ______________ valence electrons in a ____________ bond. An atom can form enough ____________ bonds to complete its __________ energy level. Compounds are made up of atoms chemically bonded together. Compounds are represented by chemical formulas, which tell us how many and what type of each atom is present in the compound. Chemical formulas make use of chemical symbols and subscripts: H2O - 2 atoms of hydrogen - 1 atom of oxygen Chemical symbol (= element) Subscript (= # of atoms) 22 Ionic Compounds(p44-53) Compounds, like atoms, are electrically neutral, which means they have no overall charge. With ionic compounds, we are looking at charged species (ions) making up the compound. In order for the compound to be neutral, the charges on each ion must cancel out. This is important to keep in mind when writing the chemical formulas for ionic compounds. Binary Ionic Compounds are composed of 2 different ions. They are always a bond between a cation and an anion Ex. sodium chloride = NaCl Writing Formulas for Ionic Compounds Step 1: Write the symbols, with the metal first Step 2: Write the ionic charge above each symbol to indicate the stable ion that each element forms Step 3: Determine how many ions of each type you need so that the total ionic charge is ZERO Step 4: Write the formula using subscripts to indicate the number of each type of ion (1 as a subscript is not written) Examples: Write the formula of the compound formed from the following: lithium + fluorine magnesium + bromine calcium + phosphorus cesium + oxygen We can replace steps 2, 3, and 4 with the Criss-Cross Rule: write the ionic charges above the symbols and cross them down. Here, you need to ensure that the formula has the lowest number of ions that will produce an electrically neutral compound. Examples: a. lithium chloride c. scandium sulfide b. magnesium fluoride 23  d. calcium nitride Naming Ionic Compounds The name of the metal always comes first, followed by the name of the non-metal, ending in “ide”. If the metal can have more than one charge, indicate which charged state the metal is in using Roman numerals in brackets (we will deal with this type next lesson). Examples: a. BaI2(s) _____________________________________________ b. Ga2S3(s) _____________________________________________ Properties/Characteristics of an ionic compound  A complete transfer of electrons to form ions  Ionic bonds (attraction of opposite charges)  Form crystal lattices  Usually solid at room Temperature  Will conduct electricity in solution or liquified (electrolytes)  High melting and boiling point  Hard and brittle  No prefixes when naming, have to balance charge Assignment: Name the elements and how many atoms of each are in the following compounds: For each of the following, indicate… a) if it is a compound or an element b) how many different atoms are present c) how many atoms in total d) which atoms and how many of each Example: H2O a) Compound c) 3 b) 2 d) 2 hydrogen & 1 oxygen 1. H2SO4 2. KMnO4 3. H2 4. H2O2 5. C6H12O6 6. Mg 7. P4 8. NaCl Writing Names and Formulas of Binary Ionic Compounds In the spaces below, record your answers to the questions in Investigation 2-A: Writing Names and Formulas of Binary Ionic Compounds. 24  1. Circle the binary compounds in the following list. (a) HCl (b) SO3 (c) MgCO3 (d) hydrogen sulfide (e) sodium bicarbonate 2. (a) Which types of elements combine to form binary ionic compounds? (b) Which types of elements combine to form binary molecular compounds? 3. Identify whether each name or formula represents an ionic or molecular substance. (a) sodium sulfide (b) PCl3 (c) nitrogen dioxide (d) zinc oxide (e) MgI2 4. Complete the following table. Element Anion Name Symbol Name Symbol fluorine F fluoride F– chloride bromide oxide sulfide nitride 5. Complete the following table. Correct () or Correct formula and name Total charge on Total charge on Formula incorrect (X) of compound cation(s) anion(s) formula? (a) LiO 25  (b) MgO (c) K2S (d) AlBr3 (e) NaN3 6. Complete the following table to write the formula of each compound. Name of compound Cation Anion Formula (a) beryllium fluoride (b) sodium nitride (c) calcium sulfide (d) aluminum chloride (e) lithium oxide (f) magnesium nitride (g) gallium sulfide (h) barium bromide Ionic Compounds—Univalent Metal Ions 26  1.If the following pairs of elements were mixed and heated, they would combine into solid ionic compounds. Write the name and formula of each compound formed. Name Formula a) silver and iodine silver iodide AgI(s) b) magnesium and oxygen c) magnesium and bromine d) calcium and nitrogen e) zinc and selenium f) sodium and sulfur g) barium and phosphorus h) aluminium and fluorine i) potassium and chlorine j) silver and oxygen Lesson 7: Polyatomic and Multi-Valent Ionic Compounds Multi-Charged/Multi-Valent Ionic Compounds:  Some metals, found in the middle of the Periodic Table are able to form more than one kind of ion. To distinguish between the different ions, we write a Roman numeral in brackets after the name of the metal in the compound to indicate its ionic charge: Cu2+ Cu+ Sn2+ Sn4+ Copper (II) Copper (I) Tin (II) Tin (IV)  Multi-charged ions are always metals  Most often, these multi-charged or multi-valent elements are found in the “Transition Metals” portion of the periodic table 27  The most common charge is listed first on the periodic table  In order to name just the multi-charged ion you must use roman numerals to indicate the charge of the ion Ex. Ni+2 = Cu+ = Mn+4 =  The steps for writing the chemical formula of a Multi-Charged Ionic Compounds based on the name are as follows: 1. Determine the charge of each ion (***NOTE***: Since the first ion has the roman numeral you will know its charge without having to go to the periodic table) 2. Balance the charges by adding subscripts to the formula Example: a. copper (I) oxide b. nickel (III) sulfide  The steps for writing the name of Multi-Charged Ionic Compounds based on the chemical formula are as follows: 1. Write the name of the first ion in the chemical formula 2. Add roman numerals for the charge of the metal 3. Write the name of the second ion (the Non-Metal) with the “-ide” ending Examples: a. FeO(s) ________________________________ b. CoF (s)2 ________________________________ c. V N (s) 3 5 ________________________________ Complex / Polyatomic Ion: A cluster of atoms that behave as a single unit in a chemical compound. These Complex / Polyatomic ions are found in the table at the top of your periodic table. Ex. OH- = NO3- = CH3COO- = NH4+ = Since Complex / Polyatomic Ions are charged particles, they have the ability to form compounds with simple ions or with other Complex / Polyatomic Ions: Ex. Na2CO3 (NH4)2SO4 The steps for writing the chemical formula of Complex / Polyatomic Ionic Compounds based on the name are as follows: 1. Treat the Polyatomic Ion as one unit and find its charge in the Polyatomic Ions Table 2. Determine the charges for the Metal ion and the Polyatomic Ion 3. Balance the charges using subscripts (***NOTE***: You must use Brackets around the polyatomic Ion if there is more than one) Examples: a. iron (II) hydroxide b. barium sulfate 28  c. copper (II) chlorate d. ammonium silicate The steps for writing the name of Complex / Polyatomic Ionic Compounds based on the chemical formula are as follows: 1. Write the name of the cation (positive ion) first 2. Write the name of the anion (negative ion) second (***NOTE***: DO NOT CHANGE the ending of the anion to “-ide”) Examples: a. Ca(OH)2(s) b. Fe2(CrO4)3(s) Assignment: Ionic Compounds—Multivalent Metal Ions 1. If the following pairs of elements were mixed and heated, they would combine into solid ionic compounds. In this worksheet, use the most common ionic form of the multivalent metal ion. The most common form is listed first in the periodic table. For example, iron exists as both 2+ and 3+ ions, with iron(III) being the most common. Name Formula a) iron and sulfur iron(III) sulfide Fe2S3(s) b) copper and oxygen c) manganese and fluorine d) gold and nitrogen e) chromium and chlorine f) platinum and phosphorus g) nickel and oxygen h) cobalt and bromine i) tungsten and iodine j) manganese and sulfur Ionic Compounds—Polyatomic Ions The names and charges of polyatomic ions can be found in lists and need not be memorized. It is a good idea, however, to get to know the more common ones introduced in the practice below. Remember to form the name by combining the positive and negative ion: name = positive ion + negative ion 29  COMBINE IONS FORMULA NAME iron(II) & nitrate Fe2+ NO3− Fe(NO3)2(s) iron(II) nitrate aluminium & nitrate Al3+ NO3− Al(NO3)3(s) aluminium nitrate sodium & sulfate lead(IV) & sulfate magnesium & carbonate gold(III) & sulfite zinc & hydrogencarbonate ammonium & nitrate copper(I) & phosphate silver & hydroxide aluminium & hydroxide lead(II) & phosphate potassium & acetate manganese(V) & sulfate Ionic Compounds: Names and Formulas 1. Write the formulas for the following compounds: a) magnesium oxide _______________ k) copper (I) bromide _______________ b) sodium fluoride_______________ l) tin (II) iodide _______________ c) aluminum nitride_______________ m) iron (III) chloride _______________ d) potassium sulfide_______________ n) calcium phosphide _______________ e) lithium iodide _______________ o) lead (II) oxide _______________ f) calcium bromide_______________ p) lead (IV) fluoride _______________ g) beryllium oxide_______________ q) tin (IV) bromide _______________ 30  h) nickel chloride _______________ r) copper (II) sulfide ______________ i) magnesium nitride_____________ s) iron (II) oxide _______________ j) aluminum sulfide_______________ t) calcium nitride _______________ 2. Write the names for the following compounds: a) Li2O ____________________ k) PbS ____________________ b) AlCl3 ____________________ l) SnO2 ____________________ c) MgS ____________________ m) Na2S ____________________ d) CaO ____________________ n) Mg3P2 ____________________ e) KBr ____________________ o) NiO ____________________ f) BeF2 ____________________ p) CuI ____________________ g) Na3N ____________________ q) PbCl4 ____________________ h) Al2O3 ____________________ r) FeP ____________________ i) CuCl2 ____________________ s) CaF2 ____________________ j) FeBr3 ____________________ t) K3P ____________________ Lesson 8: Molecular Compounds MOLECULAR COMPOUNDS(p43-44) Molecular compounds are composed of ONLY NON-METALLIC atoms  Ex.  CO2 : carbon dioxide  CCl4 : carbon tetrachloride  The Non-Metals are held together by Covalent Bonds  Covalent Bonds involve the “Sharing” of electrons between the atoms involved in the compound  The simultaneous attraction of electrons between the nuclei of adjacent atoms results in the covalent bonds Made up of non-metals; no charged species are involved. When naming Binary Molecular covalent compounds we use a system of prefixes: Subscript 1 2 3 4 5 6 7 8 9 10 Prefix Mono Di Tri Tetra Penta Hexa Hepta Octa Nona Deca Writing Formulas for Covalent Compounds Step 1: Write the chemical symbol for the first element in the name Step 2: Write the appropriate subscript according to the prefix given in the name Step 3: Repeat steps 1 and 2 for the second element Practice: dihydrogen monoxide tetraphosphorus decaoxide 31  carbon dioxide nitrogen trichloride Naming Covalent Compounds Step 1: Look at the subscript on the first atom; write the proper prefix (mono is not used as a prefix on the first atom in the compound) Step 2: Write the name of the first atom as all one word with the prefix Step 3: Repeat steps 1 and 2 for the second atom, except here the ending of the name is “ide” Examples: PCl5 I3F7 S2O6 Some molecular compounds/elements must be MEMORIZED because they do not follow the prefix rule!! Name of Molecule Formula of Molecule propane C3H8(g) methane CH4(g) methanol CH3OH(l) ethane C2H6(g) ethanol (ethyl alcohol) C2H5OH(l) acetic acid (ethanoic acid) CH3COOH(aq) ammonia NH3(g) glucose C6H12O6(s) hydrogen sulfide H2S(g) ozone O3(g) sucrose C12H22O11(s) water H2O(l) hydrogen peroxide H2O2(l) There are a number of elements on the Periodic Table that are never found by themselves in nature. H2(g) hydrogen O2(g) oxygen N2(g) nitrogen F2(g) fluorine Cl2(g) chlorine Br2(l) bromine I2(s) iodine At2(s) astatine There are two more elements that are never found alone in nature, however, these elements are found as a quartet and octet respectively. P4(s) phosphorus S8(s) sulphur Properties of Molecular Compounds:  Share electrons  Covalent bonds  Non-Metal bonded to Non-Metal  Molecules exist as individual packages  s, l or g at room temperature  Usually will NOT conduct electricity in solution (non-electrolytes)  When naming use prefixes Assignment: 32 Answer the following questions: 1. How can you tell the difference between an ionic compound and a covalent compound? 2. Name the following compounds: a) CBr4 b) NI3 c) OF2 d) SiCl4 Binary Molecular Compounds A. Write the correct name for each compound below. Use prefixes to indicate the number of atoms of each element in the name of the molecular compound. (Remember: The prefix “mono-” is not used with the name of the first element.) 1 atom: mono- 3 atoms: tri- 5 atoms: penta- 7 atoms: hepta 9 atoms: nona- 2 atoms: di- 4 atoms: tetra- 6 atoms: hexa 8 atoms: octa- 10 atoms: deca- 1) BrCl3 2) BN 3) N2O3 4) NI3 5) SF6 6) XeF4 7) PCl3 8) CH4 9) PCl5 10) P2O5 11) S2Cl2 12) ICl2 13) NH3 14) P4O10 15) H2O 16) OF2 B. Write the correct formula for each compound below. Use subscripts to indicate the number of atoms of each element in the formula (never reduce). 1) chlorine monoxide 33 2) sulfur hexachloride 3) dinitrogen monoxide 4) nitrogen trifluoride 5) sulfur tetrachloride 6) xenon trioxide 7) carbon dioxide 8) boron trichloride 9) diphosphorus pentoxide 10) phosphorus trichloride 11) sulfur dioxide 12) bromine pentafluoride 13) disulfur dichloride 14) boron trifluoride 15) tetraarsenic decoxide 16) silicon tetrachloride Properties of Substances Use your text to help you fill in the following table: Property Ionic Substances Covalent Substances Melting Point Solubility in Water Conductivity as a Solid 34  Conductivity in Solution Bond Strength A substance that conducts electricity in solution is called an __________________, while a substance that does not conduct electricity in solution is called a ____________________. Lesson 9: Acids and Bases (p. 63-71) Properties of Acids and Bases: Acids Bases Tastes Sour Tastes Bitter Can burn skin Can burn skin Turns blue litmus paper red Turns red litmus paper blue Phenolphthalein remains colourless in an acid Phenolphthalein turns pink in a base Reacts with metal to for H2(g) N/A pH Scale: Acids: Most often contain Hydrogen Ex. HCl H2SO4 H3PO4 Can be solid, liquid, or gas at room temperature When an acid is dissolved in water it will “ionize” (ie. Split into its individual ions) Ex 1: HCl(g) H+(aq) + Cl-(aq) (represented as HCl(aq)) Ex 2: H2SO4(l) 2H+(aq) + SO42-(aq) (represented as H2SO4(aq)) IUPAC Naming: (International Union of Pure and Applied Chemistry) Since Acids are technically Ionic compounds they can be named just as any other ionic compound is named. The steps for giving the IUPAC name for acids are as follows: - Name the first ion in the formula (most often Hydrogen) 1. Name the second ion in the formula making sure that it has the proper ending (Ex. “-ide” , “- ate” , or “-ite”) 2. Put the word “aqueous” in front of the name Ex: HCl(aq) = HF (aq) = H3PO4(aq) = Classical Naming: 35  Since acids are a very special type of Ionic compound they get there own naming rules: After writing the IUPAC name for an acid, use the following rules to give the acid its Classical Name: hydrogen _______________ide → hydro_______________ic acid Ex. hydrogen chloride → hydrochloric acid (HCl(aq)) hydrogen _______________ate → _______________ic acid Ex. hydrogen borate → boric acid (H3BO3(aq)) hydrogen _______________ite → _______________ous acid Ex. hydrogen nitrite → nitrous acid (HNO2(aq)) Examples: Formula Ionic Name Name a. H ClO4(aq) b. H NO2(aq) c. H 2Se(aq) Exceptions to the rules Organic Compounds:  Organic compounds are made up of Carbon, Hydrogen and Oxygen primarily  When writing the formula for these organic acids they don’t have to start with hydrogen  Ex.  CH3COOH = acetic acid  C6H5COOH  Sulfur:  Add “-ur-” before the “-ic” or “-ous” in the Classic Name  hydrogen sulfide =  hydrogen sulfate =  hydrogen sulfite =  Phosphorous:  Add “-or-” before the “-ic” or “-ous” in the Classic Name  hydrogen phosphide =  hydrogen phosphate =  hydrogen phosphite = Bases:  Most often contain Hydroxide  Ex.  NaOH  Mg(OH)2  Fe(OH)3  When a base is dissolved in water it will “dissociate” (ie. Split into its individual ions)  Ex.  NaOH(s)  Na+(aq) + OH-(aq) (represented as NaOH(aq))  Fe(OH)3(s)  Fe+3(aq) + 3OH-(aq) (represented as NaOH(aq)) IUPAC Naming: (International Union of Pure and Applied Chemistry)  Since Bases are technically Ionic compounds they can be named just as any other ionic compound is named.  The steps for giving the IUPAC name for acids are as follows: 1. Name the first ion in the formula 2. Name the second ion in the formula (***NOTE***: most often Hydroxide) 3. Put the word “aqueous” in front of the name if indicated by the state  Ex.  NaOH(aq) = 36   Mg(OH)2(aq) =  Fe(OH)3(s) =  Since Bases are Ionic compounds you simply write the formula like you would for any other Ionic compound  Ex.  aqueous potassium hydroxide =  aqueous titanium (IV) hydroxide =  aluminum hydroxide = Assignment: Arrhenius’s Theory Svante Arrhenius defined an acid as _____________________________________________________ ____________________________________________________. In the space below, show the equation representing hydrogen chloride gas dissolving in water to form an acid: He defined a base as ________________________________________________________________ ___________________________________________________________. In the space below, write the equation representing sodium hydroxide dissolving in water to form a base: Complete the following table: Property Acids Bases Taste Feel pH Conductivity 2 Indicator Tests Common Examples Naming Acids A. Using classical rules name the following acids: 1. HCl(aq) 2. H2SO4(aq) 3. H3BO3(aq) 4. HNO2(aq) 5. HBr(aq) 37 6. H3PO4(aq) 7. H2CO3(aq) 8. H2CrO4(aq) 9. CH3COOH(aq) 10. HCN(aq) B. Using IUPAC rules write the formula for each of the following acids: 1. hydrofluoric acid 2. nitric acid 3. sulfurous acid 4. phosphorus acid 5. hypochlorous acid 6. chloric acid 7. hydroiodic acid 8. perchloric acid Classifying and Naming Compounds Ionic compounds begin with a metal or the ammonium ion. Molecular compounds contain only non-metals. Acids begin with H or end with COOH. 1. Classify each of the following as an ionic compound, a molecular compound, or an acid. Name each one. Type Name a) NaCl(s) b) N2O(g) c) HCl(aq) d) NH4Br(s) e) KOH(s) f) CH3COOH(aq) 38  g) XeF2(s) h) SCl3(g) i) NiCl3(g) j) H3PO4(aq) k) K2Cr2O7(s) l) NH4NO3(s) m) CH3OH(l) n) Fe2O3(s) 2. Classify each of the following as an ionic compound, a molecular compound, or an acid. Write the formula for each one. Type Formula a) solid sodium sulfate b) aqueous hydrogen nitrate c) gaseous sulfur trioxide d) gaseous dinitrogen trioxide e) solid manganese(IV) bromide f) solid ammonium phosphate g) aqueous hydrogen sulfate Activity: Qualitative Analysis A student was given 4 unknowns to analyze and identify. She knew that the substances were (in no particular order), a solution of potassium chloride, a solution of barium hydroxide, acetic acid, and a solution of sugar. As all substances were soluble in water, she tested each substance with litmus and for conductivity. Her results are summarized in the following table: 39  Chemical Name Formula Conductivity Effect on Litmus 1 high blue to red 2 none no change 3 high no change 4 high red to blue Using her results, complete the table by identifying and giving the formula for each solid. Lesson 10: Water (page 72-78) Water has many unique properties. 1. Water in solid state is less dense than water in liquid state. Important to life in lakes. 2.Water is polar.  water has a bent shape  the electrons are not shared evenly in O-H bonds  oxygen attracts the electrons more strongly  oxygen has a permanent partial negative charge δ-  hydrogen has a permanent partial positive charge δ+  water is a polar molecule: oxygen end is δ- and hydrogen end has δ+  polar molecules also called a dipole 3. Hydrogen bonds  water molecules attract one another bcs they are polar  hydrogen bonds are formed: an oxygen atom from one water molecule is attracted to a hydrogen atom in another molecule  hydrogen bonds are an intermolecular force: force between molecules  an intramolecular force: force within molecules (covalent bond) aka cohesive force  hydrogen bonds are the strongest of intermolecular forces BUT are much weaker than intramolecular forces aka adhesive force Properties of Water 1. The boiling point and melting point of water is higher than similar substances. Reason: it takes more energy to break hydrogen bonds 2. It requires a great deal of energy to increase the temperature of water. Reason: it takes more energy to speed molecules because of hydrogen bonding 3. Water has a concave meniscus and shows capillary action. Reason: the strong attraction between polar molecules (hydrogen bonding) 4. Ice floats in water. Reason: water expands when it freezes bcs water molecules arrange in pattern held by hydrogen bonds. 5. Water has high surface tension. Reason: hydrogen bonds pull molecules on surface into smallest possible area. 40 Assignment:Read p. 72-78 to answer the following questions. Part A: Water and its properties 1. Water is the only substance on Earth that exists in ____________________________________ in all ___________________________ states. Part B: Water: A Molecular View 1. Explain how the shape of a water molecule contributes to its unusual properties. Include a sketch of a water molecule in your answer. 2. Explain how the bonding within a water molecule contributes to its unusual properties. 3. Because of oxygen’s high electronegativity, electrons shared between an oxygen atom and a hydrogen atom spend most of their time around the ____________________________. As a result, there is a partial _____________________ charged around the oxygen while there is a partial _______________________ charge around the hydrogen. A water molecule is said to be _______________________________ or possesses ____________________________. Part C: Attraction between water molecules 1. Since water molecules are _______________________, the ____________________ hydrogen end of one molecule is attracted to the _________________ oxygen end of another water molecule. This results in the formation of _______________________ __________________. 2. An intermolecular force acts __________________________ molecules; an example of this type of force is a ______________________ _________________________. 3. An intramolecular force acts __________________________ molecules; an example of this type of force is a ______________________ _________________________. 4. In general, ________________________ bonds are much stronger than _____________________________ bonds. Part D: Explaining the Properties of Water 1. Why are the boiling point and melting point of water so much higher than those of other substances? 2. Why does water have a high specific heat capacity? 3. Why do we sweat after exercising? 41  4. Why does water have a concave meniscus? 5. Why does ice float in liquid water? 6. Why does water have a high surface tension? Lesson 11: Solubility Solubility:  The ability of a substance to dissolve in a particular solvent, such as water  When a substance is “insoluble”, it means that the substance cannot dissolve into a solvent  Substances that are insoluble are said to be “Precipitates” Ionic Compounds and Solubility:  When an Ionic compound is placed into water, the ions will split apart from each other  This is called dissociation  When an ionic compound dissociate, the crystal lattice breaks apart and the ions are able to move freely in the solvent  Ex. water NaCl(s)  Na+(aq) + Cl-(aq)  If two or more Ionic compounds are placed in water a Precipitate can form between two or more of the free ions  Ex. water Na2SO4 (s)  2Na+(aq) + SO42-(aq) water CaCl2(s)  Ca2+(aq) + 2Cl-(aq) Predicting Solubility  In order to predict whether or not an Ionic compound will be soluble in water or whether it will for a Precipitate you must use the solubility table.  In order to predict this solubility use the following steps: 1. Find the Anion in the top row of the Solubility Table 2. Look down the column of that Anion to find the Cation in the compound 3. If the Cation is found in the First row the compound will be “soluble” 4. If the Cation is found in the second row the compound will be “insoluble”  Example: Na2SO4  Example: CuCl Predicting Solubility of Two Compounds  When two or more Ionic compounds are mixed together in solution you need to predict if any of the ions will form a precipitate  Use the following steps to determine whether or not a precipitate will form: 1. Determine which ions well be present in solution 2. List all possible compounds that could be formed 42  3. Determine the insoluble compounds  Example: A chemist mixes a solution of lead nitrate, Pb(NO3)2(aq) and potassium iodide, KI(aq). Does a precipitate form? If so, what is the precipitate?  Example: A chemist mixes a solution of sodium sulfate, Na2SO4(aq) and barium chlorate, Ba(ClO3)2(aq). Does a precipitate form? If so, what is the precipitate? Solubility Assignment: Using a Solubility Table(p85-89) 1. Complete the following table. High or low Name Formula Cation Anion solubility? sodium chloride NaCl Na+ Cl– lithium iodide Mg(ClO3)2 strontium hydroxide BaCO3 2. Complete the following table. High or low Name Formula solubility? Al(OH)3 ammonium chloride K2S molybdenum(V) chlorate Pb(CH3COO)2 copper(II) iodide FeCO3 calcium sulfite Ba3(PO4)2 (s) 43  palladium(II) bromide HgI strontium sulfate 3. When in solutions of KI(aq) and AgNO3(aq) are mixed, a precipitate forms. a) What four types of ions are in solution when KI(aq) and AgNO3(aq) are first mixed? b) Other than KI(aq) and AgNOs(aq), what 2 possible compounds can form from these ions? c) Which one of these compounds is the precipitate? Briefy explain your Answer. Lesson 12: Chemical Reactions and Types of Reactions (p84-94) Chemical Reaction A chemical reaction (rxn) occurs when one or more substances change to form different substances. (chemical change) Evidence of a Chemical Reaction Energy Change All chemical rxns involve a change in Energy: temperature change (thermal E), emission of light , emission of sound, electrical energy Odour Change In some rxns odour change Colour Change In some rxns colour change Formation of a gas In some gases formed Eg. bubbles Formation of a solid (precipitate) in a solution In solution new substance no longer soluble forms solid (precipitate)  In order to know whether or not a chemical change has occurred it is best if two or more of those Evidences are apparent  Ex.  Odour change and gas formed  Colour change and heat produced  Etc.  However, the best indicators of whether or not a chemical reaction has occurred are as follows:  A new substance is formed.  Ex.  Hydrochloric acid mixed with Zinc will produce Hydrogen gas  The reaction can not be reversed  Ex.  Try to “uncook” an egg!!! Chemical Reactions & Energy Change  In ALL chemical reactions, energy needs to be either ABSORBED or RELEASED Exothermic Reaction: reactions that release energy, energy is a product 44  Endothermic Reaction: reactions that absorb energy, energy is a reactant Law of Conservation of Energy:  Energy can be converted from form to another, but the total energy of the universe remains constant. In order for this to be true, energy is linked to the formation or breakdown of chemical bonds  Breaking chemical bonds is endothermic Eg. energy + water → hydrogen + oxygen  Forming new chemical bonds is exothermic Eg. hydrogen + oxygen → water + energy More complex reactions cannot predict yet..... Law of Conservation of Mass: During a chemical reaction, the total mass of the reacting substances (the reactants) is always equal to the total mass of the resulting substances (the products)  Ex.  Reactant + Reactant  Product + Product Antoine Lavoisier (1743-1794):  Father of Modern Chemistry  Lavoisier developed the Law of Conservation of Mass based on a series of Closed System experiments using test tubes  Closed Systems:  does not allow exchange of matter between the system and its surroundings)  Open Systems:  allows the exchange of substances between the system and its surroundings Types of Chemical Reactions  Chemical reactions involve the mixing of two or more Reactants that forms one or more Products  Ex.  2Zn(s) + 2HCl(aq)  H2(g) + 2ZnCl(aq)  C6H12O6(s) + 6O2(g)  6CO2(g) + 6H2O(g) 1. Formation (Synthesis) Reaction o Reactions involve two Elements reacting to form one compound A + B → AB Eg. 2. Decomposition Reaction o involve one Compound decomposing to form two different Elements AB → A + B Eg. H2O(l) + Electricity  H2(g) + O2(g) 3. Single Replacement Reaction o involve one Compound reacting with one Element to form one new Compound and one New Element AB + C  CB + A Or DE + F  DF + E 45  Eg. 4. Double Replacement Reaction o involve two Compounds reacting to form two new Compounds AB + CD  AD + CB Eg. 5. Combustion Reaction o Combustion Reactions often are incredible reactions that we would think of as “Burning” but in reality they are when any substance reacts with oxygen Ex. o 4 Fe(s) + 3 O2(g)  2 Fe2O3(s) o 2 Mg(s) + O2(g)  2MgO(s) A. Hydrocarbon Combustion o Hydrocarbon Combustion reactions are a specific type of combustion reactions involve one hydrocarbon reacting with oxygen gas to form carbon dioxide and water General Formula: CxHyOz + O2(g)  CO2(g) + H2O(g) Ex. C6H12O6(s) + 6O2(g)  6CO2(g) + 6H2O(g) 6. Neutralization Reaction Neutralization reactions involve an Acid reacting with a Base to form Water and An Ionic Salt General Formula: HB + AOH  H2O(l) + AB  Ex.  HCl(aq) + NaOH(aq)  H2O(l) + NaCl(aq) Assignment: Physical and Chemical Changes Objects undergo many kinds of changes. Place a P in front of each physical change below and a C in front of each chemical change. ____ 1. rusting of iron ____ 13. etching glass with an acid ____ 2. breaking a tree branch ____ 14. formation of stalagmites in a cave ____ 3. cutting paper ____ 15. fertilizing a lawn ____ 4. action of yeast in breadmaking ____ 16. crushing ice in a blender ____ 5. souring of milk ____ 17. evaporation of water in a lake ____ 6. wadding up a sheet of paper ____ 18. eating food ____ 7. erasing a pencil mark ____ 19. burning gas in a car engine ____ 8. freezing water ____ 20. burning logs in a fireplace ____ 9. boiling water ____ 21. toasting marshmallows 46 ____ 10. salting the ice on a sidewalk ____ 22. adding bleach to a washer ____ 11. the action of baking powder ____ 23. slicing a block of cheese in baking a cake ____ 12. bending a metal wire ____ 24. letting ice cream melt Types of Reactions (p103-114) 1. Composition/Formation Reactions Definition: General Equation: Examples: 2. Decomposition Reactions Definition: General Equation: Examples: 3. Single Replacement Reactions Definition: General Equation: Examples: 4. Double Replacement Reactions Definition: General Equation: Examples: 5. Combustion Reactions Definition: General Equation: Examples: 47  Classifying Equations 1. Classify each reaction as a formation (F), decomposition (D), single replacement (SR), double replacement (DR), or combustion (C) reaction. Reaction Classification Li(s) + AlCl3(aq) Al (s) + LiCl(aq) NH3(g) N2(g) + H2(g) P4(s) + Cl2(g) PCl5(s) K(s) + Br2(l) KBr(s) C10H22(l) + O2(g) CO2(g) + H2O(g) NH4OH (aq) + H2CO3(aq) H2O(l) + (NH4)2CO3(aq) HNO3(aq) + Ba(OH)2(aq) H2O(l) + Ba(NO3)2(aq) H2O(l) H2(g) + O2(g) Cu(s) + Ag2CO3(s) CuCO3(s) + Ag(s) Al(s) + Cl2(g) AlCl3(s) Zn(s) + SnF4(aq) Sn(s) + ZnF2(aq) P4(s) + S8(g) P2S5(s) CuO(s) Cu(s) + O2(g) Na2CrO4(aq) + Cu(NO3)2(aq) NaNO3(aq) + CuCrO4(aq) Lesson 13: Writing and Balancing Chemical Reactions (p84-102) o A balanced equation is a recipe for the production of chemicals. The reactants are the ingredients. (Written on the left side of the equation) The products are the chemicals formed. (Written on the right side of the equation) o In chemical reactions, bonds are broken and formed. However, the atoms are not destroyed, they are only rearranged. Therefore, the number of atoms of each element cannot change before and after a reaction has taken place and chemical equations must be balanced. WRITING WORD EQUATIONS o Chemists use word equations to describe the many different chemical equations that can take place. Word equations describe the chemical change that has occurred. Word equations always are written in the same format. 48  o They contain three parts: Reactants, products and the arrow that separates them. o All reactants are found on the left hand side of the equation. These are the substances that are reacted together. o All products are on the right hand side of the equation. These are the new substances that are produced and have different properties than the substances that were reacted. An arrow points in the direction the reaction is going, from the reactants to the products. REACTANTS →PRODUCTS The arrow is read as PRODUCE OR PRODUCES OR FORMS, ETC.  Each reactant is separated from another reactant by a + sign REACTANT A + REACTANT B When reading the + sign is read as AND.  Each product is separated from another reactant by a + sign. PRODUCT A + PRODUCT B When reading the + sign is read as AND.  If given the sentence: Sodium hydroxide and hydrochloric acid produces sodium chloride (table salt) and water.  The word equation would be written: The skeleton chemical equation would be written: Balancing Chemical Equations  Antoine Lavoisier’s Law of Conservation of Mass has led to some interesting concepts in Chemistry  One major thing to consider is that the compounds in a chemical reaction have different numbers of the same atoms Eg. Cu(s) + O2(g) → CuO(s) * you need the same number of each atom on each side * you can only change the number in front  The rules of Balancing Chemical Reactions are as follows: 1. You need to have the same number of EACH atom on either side of the ARROW in the reaction 2. You can only add COEFFICIENTS to the front of each compound in the reaction (***NOTE***: Coefficients multiply each atom in the compound)  The following steps should be used to Balance Chemical Reactions: 1. Predict the products 2. Look for the most confusing looking compound in the chemical reaction (ie. The one with the most number of ions) 3. Balance one of these atoms by adding a Coefficient to the front of the compound (This will multiply the number of ions in the compound) 4. By adding a Coefficient you will also be multiplying the other ion in the compound. Next, add a Coefficient to the compound on the other side of the Arrow to Balance it. 5. Follow the above steps until all atoms are equal in number on either side of the Arrow Tips:  if there is a base (OH) change the water to HOH first  pick an element an element (usually not H or O) on one side and balance other  move on to another and do the same  double check, sometimes you will have to go back and change numbers  if can only have evens double the number 49   for states of ionic compounds in single and double replacement reactions check the table of solubility to determine if soluble (aq) or not (s). Ex 1. Cr2O3(s) → Ex 2. H2(g) + Fe2O3(s) → Ex 3. Cl2(g) + NaBr(aq) → Ex 4. Ca(OH)2(aq) + HCl(aq) → Ex 5. C2H4(g) + O2(g) → Assignment: Writing Formula Equations From Word Equations In the following exercises, recall that many of the non-metal elements exist as molecules, such as H2(g) or S8(s). 1. Rewrite the following sentences into: 1) word equations 2) skeleton equations. a) solid sodium metal reacts with chlorine gas to produce solid sodium chloride 1) Soild Sodium + Chlorine gas  Solid Sodium Chloride 2) Na(s) + Cl2(g)  NaCl(s) b) solid potassium metal reacts with oxygen gas to produce solid potassium oxide c) hydrogen gas reacts with oxygen gas to produce liquid water d) solid potassium chlorate decomposes into oxygen gas and solid potassium chloride e) solid aluminium oxide is decomposed into solid aluminium and oxygen gas f) mercury(II) sulfide is decomposed into liquid mercury and solid sulfur 50  g) aqueous cobalt(III) nitrate reacts with solid zinc to produce aqueous zinc nitrate and solid cobalt h) fluorine gas reacts with aqueous lead(IV) iodide to produce aqueous lead(IV) fluoride and solid iodine i) aqueous gold(III) bromide reacts with solid silver metal to produce solid silver bromide and solid gold metal j) aqueous sodium sulfate reacts with aqueous strontium hydroxide to produce aqueous sodium hydroxide and solid strontium sulfate k) aqueous thallium(I) hydroxide reacts with aqueous magnesium bromide to produce solid magnesium hydroxide and solid thallium bromide l) methane gas reacts with oxygen gas to produce carbon dioxide gas and water vapour m) magnesium metal reacts with liquid bromine to produce magnesium bromide n) solid sulfur and hydrogen gas react to produce hydrogen sulfide gas o) methane gas reacts with liquid water to produce liquid methanol and hydrogen Balancing Equations I Balance the following chemical equations: a) _____ Na(s) + _____ O2(g) à _____ Na2O(s) b) _____ Al(s) + _____ Cl2(g) à _____ AlCl3(s) c) _____ N2(g) + _____ O2(g) à _____ NO2(g) 51 d) _____ HI(g) à _____ H2(g) + _____ I2(s) e) _____ NH3(g) à _____ H2(g) + _____ N2(g) f) _____ Al2S3(s) à _____ Al(s) + _____ S8(s) g) _____ BN(s) + _____ Cl2(g) à _____ BCl3(g) + _____ N2(g) h) _____ SnF4(aq) + _____ Cr(s) à _____ CrF3 + _____ Sn(s) i) _____ Mg(s) + _____ HCl(aq) à _____ MgCl2(aq) + _____ H2(g) j) _____ (NH4)3PO4(aq) + _____ CaBr2(aq) à _____ Ca3(PO4)2(s) + _____ NH4Br(aq) k) _____ Pb(NO3)4(aq) + _____ K2Cr2O7(aq) à _____ Pb(Cr2O7)2(s) + _____ KNO3(aq) l) _____ AgClO4(aq) + _____ Na3PO4(aq) à _____ NaClO4(aq) + _____ Ag3PO4(s) m) _____ HCl(aq) + _____ Ca(OH)2(aq) à _____ CaCl2 + _____ H2O(l) n) _____ CH3COOH(aq) + _____ Ba(OH)2(aq) à _____ Ba(CH3COO)2(aq) + _____ H2O(l) o) _____ C3H8(g) + _____ O2(g) à _____ CO2(g) + _____ H2O(g) p) _____ C6H14(l) + _____ O2(g) à _____ CO2(g) + _____ H2O(g) Balancing Equations II Balance the following chemical equations: a) _____ Pb(s) + _____ O2(g) à _____ PbO(s) b) _____ N2(g) + _____ H2(g) à _____ NH3(g) c) _____ Na(s) + _____ H2O(l) à _____ NaOH(aq) + _____ H2(g) d) _____ C4H10(g) + _____ O2(g) à _____ CO2(g) + _____ H2O(g) e) _____ H3PO4(aq) + _____ KOH(aq) à _____ K3PO4(aq) + _____ H2O(l) f) _____ C5H12(l) + _____ O2(g) à _____ CO2(g) + _____ H2O(g) g) _____ Zn3N2(s) + _____ H2O(l) à _____ Zn(OH)2(s) + _____ NH3(g) h) _____ Fe2O3(s) + _____ H2(g) à _____ Fe(s) + _____ H2O(l) i) _____ Al(s) + _____ H2SO4(aq) à _____ H2(g) + _____ Al2(SO4)3(aq) 52 j) _____ CrS(s) + _____ O2(g) à _____ CrO(s) + _____ SO2(g) k) _____ HClO3(aq) + _____ HCl(aq) à _____ H2O(l) + _____ Cl2(g) l) _____ CaC2(s) + _____ AsBr3(aq) à _____ C(s) + _____ As(s) + _____ CaBr2(aq) m) _____ NH3(g) + _____ O2(g) à _____ NO(g) + _____ H2O(g) n) _____ HNO3(aq) + _____ NO(g) à _____ NO2(g) + _____ H2O(l) o) _____ Al(NO3)3(aq) + _____ NaOH(aq) à _____ NaNO3(aq) + _____ Al(OH)3(s) q) _____ NaIO3(s) à _____ NaI(s) + _____ O2(g) Putting It All Together For each of the following, classify the reaction, predict the product and then balance the reaction, including proper states where needed: a) CaCl2(aq) + K3PO4(aq) b) K(s) + H3PO4(aq) c) Cl2(g) + MgBr2(aq) d) Na2CO3(aq) + Fe(NO3)3(aq) e) Al2(SO4)3(aq) + Ba(OH)2(aq) f) CuO(s) g) K(s) + O2(g) h) C8H18(g) + O2(g) i) ZnS(s) + O2(g) j) Fe(s) + H2S(aq) k) Na2SO4(aq) + BaCl2(aq) l) P4(s) + Cl2(g) PCl5(s) m) AlCl3(s) 53  n) Na(s) + H2SO4(aq) o) Cu(s) + S8(s) p) Mg(s) + O2(g) q) CaBr2(s) r) C2H2(g) + O2(g) Lesson 14: Scientific Notation and Significant Digits: Scientific Notation  Scientific Notation is a way to write either really BIG numbers or really SMALL numbers in a shorter form. Ex. 30,567,000 = 0.0000004891 =  Scientific Notation always takes on the following form: Where: - N = is a number between 1 and 10 (***NOTE***: represents the actual number) - n = is the exponent and can either be POSITIVE (if the actual number is LARGE) or NEGATIVE (if the actual number is SMALL) whole number. (***NOTE***: represents the number of decimal places moved) In order to rewrite a number in Scientific Notation use the following steps: 1. To Find N: Place the decimal after the first real digit in the number 2. To Find n: Count the number of decimals that were moved from the original decimal position. - If you moved the decimal to the LEFT n will be POSITIVE. - If you moved the decimal to the RIGHT n will be NEGATIVE Ex. Express 568.762 in scientific notation Express 0.0000072 in scientific notation Express 504900000 in scientific notation Express 0.00005301 in scientific notation Significant Digits  The meaningful digits in a measured or calculated quantity.  Ex: Sometimes it doesn’t make sense to give all of the digits in a number  The Eiffel Tower is 324 m when in actual fact it is probably 324.54689 m. Significant Digits in Science There are two kinds of numbers in the world:  exact o There are exactly 12 eggs in a dozen o Most people have exactly 10 fingers and 10 toes  inexact numbers 54  o any measurement o See for yourself: what is the volume of the liquid in the graduated cylinder below? All measurements have digits that are KNOWN WITH CERTAINTY, plus ONE digit that is uncertain (estimated) – together they are known as the significant digits of the measurements. In order to determine the number of Significant Digits in a given number you must use the following rules. Rule 1: In numbers that do not contain zeros, ALL digits are significant Eg. Rule 2: All zeros BETWEEN significant (nonzero) digits are significant Eg. Rule 3: Zeros to the LEFT of the first nonzero (leading zeros) are NOT significant Eg Rule 4: All zeros to the RIGHT of the first nonzero (trailing zeros) ARE significant Eg. Rule 5: For number that do not contain decimal places, trailing zeros may or may not be significant. Eg. Use Scientific Notation  4.0 102 has 2 sig digs  4 102 has 1 sig digs Ex: Calculate the number of significant digits in each of the following numbers:  860.20 =  0.1003 =  0.00030 =  4.65 x 1081 =  501 = Significant Digits When Calculating: When multiplying or dividing numbers, your answer must be stated in the least number of significant digits given in a value in the question. 55 Multiplication The answer must be rounded off to 2 significant figures, since 1.6 only has 2 significant figures. Division The answer must be rounded off to 3 significant figures, since 45.2 has only 3 significant figures. Ex. Calculate the following and round to the appropriate number of significant digits. 40.24 x 3.567 = 0.0024 x 1.4 = 12 ÷ 0.56 = 500.0 x 0.19 = When adding or subtracting, your answer is given in the least number of decimal places given in a value in the question or the last digit retained is set by the first doubtful digit: EXAMPLES Addition Even though your calculator gives you the answer 8.0372, you must round off to 8.04. Your answer must only contain 1 doubtful number. Note that the doubtful digits are underlined. Subtraction Subtraction is interesting when concerned with significant figures. Even though both numbers involved in the subtraction have 5 significant figures, the answer only has 3 significant figures when rounded correctly. Remember, the answer must only have 1 doubtful digit. Ex. Calculate the following and round to the app