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Unit 2 Matter and Chemical Change Topic 1 Key Concepts WHMIS (Workplace Hazardous Materials Information System) Substances and their properties Elements, Compounds and Atomic Theory Topic 1 My Learning Outcomes When I have completed this section I will be able to : Identify and evaluate dangers of caustic materials and potentially explosive reactions. Investigate and describe properties of matter describe and apply different ways of classifying matter based on its composition and properties. Topic 1 Safety in the Science Classroom In any science activity, the safety of you, your classmates, and your teacher are of the utmost importance. There are two areas of special consideration: Understanding warning labels Following safety procedures Keep Safety in Mind Before you start any science activity: Follow all safety instructions from the teacher or from the text book. Identify possible hazards and report them immediately Show respect and concern for your own safety and the safety of others Read Toolbox 1: Safety in the Laboratory (Page 478) Organizing Matter Almost all matter exists as either a solid, liquid or gas. (A fourth state of matter is the plasma state - created when a large amount of energy is added to a gas) Physical Properties of Matter Matter Colour Lustre Melting Point Boiling Point Hardness Malleability Ductility Crystal Shape Solubility Density Conductivity Chemical Properties of Reaction with acids/bases ability to burn reaction with water behavior in air reaction to heat/light Pure Substances and Mixtures Observing Changes in Matter Changes in matter are classified as physical or chemical. Physical Change - matter changes state Chemical Change- two or more materials react to create new materials with different properties from the original substances. Evidence of Chemical Change A colour change Formation of an odour Formation of a gas (bubbles) Formation of a solid (precipitate) Release or absorption of energy (heat, light, sound) Topic 2.0 An understanding of the nature of matter has developed through observations over time. Evolving Theories of Matter Organizing the Elements The Periodic Table Today Topic 2 – Key Concepts Substances and their properties Elements, compounds and atomic theory Periodic table Topic 2 - Learning Outcomes When I have completed this section I will be able to: Distinguish between observation and theory and provide examples of how models are used to explain observations Demonstrate understanding of the origins of the periodic table and relate physical and chemical properties of the elements to their position on the periodic table Use the periodic table to Identify numbers of protons and electrons in each atom Describe the relationships between the structure of atoms in each group and the properties of the elements in that group. 2.1 Evolving Theories of Matter People try to make sense of the world around them by suggesting explanations (theories) to explain their observations. Over time these theories are modified as new evidence is discovered…. Ex. Human beings once thought the earth was flat and as we began to discover more about our planet the theories changed. Stone Age Chemists 8000B.C. In the Middle East Metals not discovered yet Learned to start and control fire Controlling fire led to producing glass and ceramic tile. Early Interest in Metals and Liquid Matter 6000 B.C. – 1000 B.C. Chemists investigated valuable materials Many were metals like gold & copper Chemists wanted to understand how copper could be controlled. Original discovery that heat made copper very useful was accidental – some fell in a fire. Later experiments (4500B.C.) led to creation of Bronze – made from copper and tin. Iron Age 1200 B.C. Hittites in Middle East discovered how to extract iron from rocks and turn it into a useful material. Led to combining iron & carbon to produce an even harder material – steel. Emerging Ideas About the Composition of Matter 500 B.C. - Idea that all matter is made up of particles started with Greek philosophers 400 B.C. – Democritus used word atomos to describe smallest particles that could not be broken down. (means indivisible) Believed each type of material was made up of different atomos These particles gave each material its own unique set of properties. 350 B.C. – Aristotle said everything was made of earth, air, fire, and water. He was so well known and respected that his ideas were accepted over Democritus for 2000 years. Chemistry Develops into a Science 1660’s (A.D.) Robert Boyle - matter is made up of tiny particles. 1770 – 1780’s – Antoine Laurent Lavoisier develops a system for naming chemicals “the father of modern chemistry” Atomic Theory Billiard Ball Model 1808 – John Dalton Suggests matter is made up of elements First person to define element Each element is composed of an atom Atoms cannot be broken into smaller particles Raisin Bun Model 1897 – J.J. Thomson First person to discover a subatomic particle Named them electrons Thomson discovered the electron in 1897 So, the atom was not the smallest unit of matter Since there are negative electrons in the atom, there must also be positive substances Thomson’s model of the atom: the positively charged substance fills the atom and the electrons are embedded throughout it The atom is like a “raisin bun” Rutherford Model 1898 – 1907 Earnest Rutherford (working in Canada) Atoms mainly empty space Core (center) was a tiny positively charged nucleus (1/10000th the size of the atom) Bohr’s Model 1913 Niels Bohr (while working with Rutherford) Electrons move in specific circular orbits – electron shells (not randomly) Electrons could jump shells by gaining or losing energy Bohr model refined by James Chadwick Discovered positive protons and neutral charged neutrons in the nucleus. Neutron and proton same mass but 1800 times more than an electron. Quantum Mechanics Model Describes electrons as existing in a charged cloud around the nucleus. 2.2 Organizing the Elements 1814 Jons Jacob Berzelius Suggested using letters rather than picture to represent elements.. The first letter (capitalized) of an element would be the symbol. For elements with same first letter (hydrogen & helium) a small second letter would be added. H – hydrogen & He for helium. 1869 – Dmitri Mendeleev Organized elements in a way that reflected patterns in the properties of elements. Was able to show that the properties vary periodically with increasing atomic mass Had gaps in his chart of elements but predicted that new elements would be discovered that would have the properties and atomic mass needed to fit. Within 16 years was proven right- gallium discovered. The Periodic Table Today Understanding the Periodic Table Horizontal rows (periods) Numbered 1-7 From left – right elements change from metals to non-metals Most reactive metals on the left Atomic # increases as we move to the right Groups (Families) Vertical columns Elements in the same family have similar properties Numbered 1 – 18 Special Names Alkali metals (group 1) – most reactive metals Alkaline Earth metals (group 2) very reactive Halogens (Group 17) most reactive non-metals Noble gases (group 18) – very stable, almost never react Element Information Element Symbols Some symbols are based on latin names of elements instead of the modern name Ex. Gold – Au (aurum) Lead – Pb(plumbum) Atomic Number Tells us # of protons Atoms of same element all have the same # of protons Ex. Oxygen – 8 protons Also tells us # of electrons Atomic Mass Tells us total mass of all protons and neutrons in an atom. (electrons are too tiny to effect mass of an atom) Not all atoms of same element have same atomic mass (different #’s of neutrons) Mass Number Is the sum of protons and neutrons in an atom Mass # - atomic # = neutrons Patterns of Information Metals Left ¾ of periodic table Shiny, malleable, ductile, conduct electricity Non-metals Right side of periodic table Solid or gas Dull, brittle elements Do not conduct electricity (carbon is exception) Metalloids Diagonal row of elements between metals & non-metals Have both metallic and non-metallic properties Topic 3 Compounds form according to a set of rules. Naming Compounds Ionic Compounds Molecular Compounds Topic 3 Key Concepts Periodic table Elements, compounds and atomic theory Chemical nomenclature Topic 3 - My Learning Outcomes When I have completed this section I will be able to: Distinguish between ionic and molecular compounds and describe examples of each. Read and interpret chemical formulas for compounds of two elements, and give the IUPAC name and common name for these compounds Identify examples of combining rations/number of atoms per molecule found in common materials and use ionic charges to predict combining rations in ionic compounds. 3.1 Naming Compounds Combining Elements to Make Compounds Looking around our homes we are surrounded by chemicals - from water to bleach. Each compound has a chemical name and a chemical formula. This formula identifies which elements and how much of each are in the compound. Sodium Chloride (aka. table salt) chemical formula NaCl. Sodium Bicarbonate (aka baking soda) NaHCO3 Naming Chemical Compounds Until 18th century there was no standard system. Names for chemicals differed from country to country. Ex. Hydrochloric acid and muriatic acid refer to the same thing. 1787 a French Chemist (Guyton de Morveau) created a naming system (nomenclature) for compounds The chemical name for each element in the compound is used in the compound name. Ex. Zinc and oxygen form zinc oxide. Since 1920, the International Union of Pure and Applied Chemistry (IUPAC) decides on the name for every chemical compound discovered. Interpreting Chemical Names and Formulas & Indicating Physical State. If you know the formula you can determine the chemical name; if you know the chemical name then you can determine the formula. Ex. H2O – the two indicates there are two atoms of hydrogen to go with every atom of oxygen in water. NaCl – one atom of sodium to every one atom of chlorine. Another common notation added to chemical compounds indicates the state at room temp. After the chemical formula a subscript (s) for solid, (l) for liquid, (g) for gas or an (aq) to show something dissolved in water. (aqueous solution) 3.2 Ionic Compounds Ionic Compounds are pure substances formed from the attraction of (+) & (-) charged particles called ions. (ex. Table salt formed from (+) sodium & (-) chloride ions. All ionic compounds Are solid at room temp. Conduct electricity when dissolved in water Formed from metals (+) and non-metals (-) Have high melting points Ion Charges An ion is an atom or group of atoms that has become electrically charged through the loss or gain of electrons. To indicate ions a (+) or (-) sign is placed to the upper right of the element symbol (Ex hydrogen ions are h+ and chlorine ions are Cl - ) Some ions also form when certain atoms of elements combine. (ex carbonate CO3 2-) These polyatomic ions act as one. Naming Ionic Compounds Two rules: The chemical name of the metal or (+) ion goes first followed by the name of the non-metal or (-) ion. The name of the non-metal (-) ion changes its ending to ide. (NaCl is not sodium chlorine but sodium chloride. o ** there is one exception to these rules – where negative ions are polyatomic ions the name remains unchanged. Some elements have more than one ion charge. To show which ion is being used a Roman Numeral is added. Ex. iron(III) oxide contains the Fe3+ ion. (Fe2O3) Ex. iron(II) oxide contains the Fe2+ ion (FeO) Using Ion Charges to Write Chemical Names Step 1 – print the metal elements symbol with its ion charge. Next to it print the non-metal elements symbol and ion charge. Ca2+ Cl 1 Step 2 – balance the ion charges. You must have an equal number of (+) charges & (-) charges. Ca2+ Cl 1- Cl 1 Step 3 – Write the formula showing the number of atoms of each element to balance the formula. ** Do Not Include The Ion Charge. CaCl2(s) Molecular Compounds Writing Formulas for Molecular Compounds Formed when non-metals combine. May be solids, liquids or gases Tend to be insulators, or poor conductors of electricity Have relatively low melting & boiling points Writing Formulas no ions are present and the ion charge is not used in the formula Hard to predict combinations Formulas show they type and # of atoms of each element. Ex. Ammonia NH3(g) Naming of Molecular Compounds Many molecular compounds known by their common name. Ex. Water & ammonia Names of these do not tell what elements they are made of All molecular compounds except those with hydrogen can be named using the following rules: The first element uses the element name The second element has the suffix (ends in) ide. More than one atom in the formula a prefix is used which specifies # of atoms. If the first element is only one atom – no prefix. Examples of Naming Molecular Compounds Topic 4 – Substances undergo a chemical change when they react to produce different substances 4.1 Chemical Reactions 4.2Conservation of Mass in Chemical Reactions 4.3Factors Affecting the Rate of a Chemical Reaction Topic 4 - Key Concepts Endothermic and exothermic reactions Reactants and products Conservation of mass Factors affecting reaction rates Topic 4 – My Learning Outcomes When I have completed this section I will be able to: Identify and critically evaluate if a new substance has been formed Observe and describe evidence of chemical change in reactions Distinguish between materials that react readily and those that do not Describe chemical reactions and represent these reactions by using word equations and chemical formulas and by constructing models of reactants and products. Chemical Reactions Writing and Identifying Chemical Reactions Chemical changes result from chemical reactions Starting materials – reactants New materials created – products Ex. Campfire – burning wood undergoes a combustion reaction. Wood and oxygen are the reactants. The products are carbon dioxide and water wood + oxygen carbon dioxide + water + heat Evidence of a chemical change (5 clues) A colour change A formation of an odour The formation of a gas The formation of a precipitate (solid) The release or absorption o energy (heat/light/sound) Common Types of Chemical Reactions Exothermic & Endothermic Reactions Chemical changes involving oxygen Combustion reactions (burning) Corrosion (rusting) Cellular respiration (body cells producing energy) Conservation of Mass in Chemical Reactions Products are formed when reactants undergo a chemical change Often look very different from reactants Conservation of Mass states that: Matter is not created or destroyed in a chemical reaction The mass of the products is always the same as the total mass of the reactants Ex. Combining 24.3g of Mg and 32.1 g of S creates 56.4g of MgS (magnesium sulfide) Closed System – no product escapes Open System – products escape and affect total product mass. Factors Affecting the Rate of a Chemical Reaction There are four factors that can affect the rate of a reaction. Presence of a catalyst Help reactions proceed faster – ex. Enzymes Concentration Greater the concentration, the faster the reaction Temperature More heat added to the reactants, the faster the reaction Surface Area The greater the surface area of the reactants, the faster the reaction Chapter 2 - Study Guide Topic 1 Classroom Safety Whmis Stands for Whmis symbols Organizing Matter Physical Properties State, colour, lustre, malleability, solubility, density, hardness, boiling/melting points, conductivity Chemical Properties Reaction with acids & waterbehavior in air, reaction to heat, ability to burn, Pure Substances Elements & compounds Mixtures Mechanical mixtures, solutions, suspensions & colloids Physical Changes Change state (Phase Change comp. doesn’t change) Evaporation, melting, sublimation, condensation, freezing,. Chemical Changes 2 or more materials react to create new materials 5 clues that a chemical reaction has occurred Metals, Non-metals, Metalloids Alkali Metals, Alkaline Earth Metals Halogens , Noble Gases Chapter 2 Study Guide Topic 2 2.1 Theories of Matter Timeline of discovery Democritus and Aristotle What were their ideas about matter Alchemists Lavoisier – naming chemicals Atomic Theory John Dalton – Billiard ball model JJ Thomson – Raisin Bun Model Nagaoka – Electrons in orbit Ernest Rutherford – famous experiment, his theory Neils Bohr – Bohr Model Electorn Shells James Chadwick Protons & neutrons Quantum Theory 2.2 Organizing the Elements Mendeleev and the Periodic Table Atomic Mass Definition Qhy is it important 2.3 The Periodic Table Periods – horizontal Groups (families) – Vertical Element Names – capital then lower case Atomic # vs Mass # Mass # = protons + neutrons Chapter 2 Study Guide Topic 3 Forming & Naming Compounds Chemical Names and Chemical Formulas Ionic Compounds Properties Solid at room temp. High melting points Good conductor Formed from metal (+) and non-metal (-) ions Rules for Naming Ionic Compounds Metal comes first and keeps name –Non-metal second – end of name changes to ide. Balance ion charge to determine ratio of atoms. Molecular Compounds Properties Solid, Liquid or gas at room temp Relatively low melting/boiling points Insulator (poor conductor) Formed from two non-metals Rules for Naming Molecular Compounds First element keeps element name Second element changes ending to ide If there is more than one atom a prefix is attached Exception to rule 3 when the first atom is only one we do not use the mono prefix. Chapter 2 Study Guide Topic 4 Chemical Reactions 5 Clues that a new substance has formed A colour change Formation of an odour Formation of a gas (bubbles) Formation of a solid (precipitate) Release or absorption of energy (heat, light, sound) Types of Reactions Endothermic & Exothermic Combustion Corrosion Cellular respiration Law of Conservation of Mass Reaction Rates Affected By Catalysts Concentration Temperature Surface Area