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Unit 2: Matter and Chemical Change Lesson 1: Properties of Matter All chemicals should be handled very carefully. WHMIS (Workplace Hazardous Materials Information System) has been developed to provide guidelines for handling, storage and disposal of reactive materials (chemicals). Some materials are CAUSTIC - will burn, corrode or destroy organic tissue MATTER: Matter is anything that has mass. All matter is described in terms of its properties PROPERTIES - are characteristics you can use to describe or identify different substances. Ex. Color, luster, state, taste, melting point, and behavior Properties can be classified as: 1. QUALITATIVE PROPERTIES - Those properties, which describe a quality of matter that has no numerical value (no number) - Usually involves one of the senses Ex. Taste, odor, texture, color, luster 2. QUANTITATIVE PROPERTIES -Those properties, which describe a quantity of matter - Have a number associated with the property Ex. Melting point of water is 00 C Freezing point, boiling point, number of legs, density etc. CHEMICAL PROPERTIES – those properties that involve the formation of a new substance. Chemical properties cannot be tested without destroying the substance. Ex. Rust is formed when iron reacts with oxygen Magnesium burns to produce a white powder Paper burns to produce carbon dioxide and water PHYSICAL PROPERTIES – do not involve the formation of a new substance. Physical properties can be tested without destroying the substance. Ex. Melting point – temperature at which something changes from a solid to a liquid Boiling point – temperature at which something changes from a liquid to a gas Malleability – can be pounded or rolled into sheets Ductility – can be stretched into a wire Density – the amount of mass in a given volume of a substance Conductivity – ability of a substance to conduct electricity Solubility – ability to be dissolved in another The Particle Model of Matter states that: 1. All matter is made up of extremely tiny particles. 2. Each pure substance has its own kind of particle, different from the particles of other pure substances. 3. Particles attract each other. 4. Particles are always moving. 5. Particles at a higher temperature move faster on average than particles at a lower temperature. All matter can be classified according to its state as a 1. Solid – has a definite shape and a definite volume - Particles are close together and the forces between the particles are strongest 2. Liquid - has a definite volume but takes the shape of its container - Are said to be fluid because they flow - Particles are further apart and the forces between the particles are weaker 3. Gas - has neither a definite shape nor a definite volume - Expands to fill the container - Gases are fluids because they also flow - Particles are furthest apart and the forces between the particles are very weak CHANGES OF STATE 1. 2. 3. 4. 5. Melting Freezing Boiling/Vaporizing Condensation Sublimation Classification of Matter Matter can also be classified according to its composition: Matter Pure Substances Mixtures Elements Compounds Mechanical Mixtures (heterogeneous) Suspension Colloids Solutions (homogeneous) PURE SUBSTANCES have properties that are always the same. Ex. melts Table salt is a white solid that at 801 °C and boils at 1465 °C. Water, vinegar, sugar etc. Elements -Pure substances that cannot be broken down into any simpler substances - most elements are solids but several are gases and two are liquids -All of the elements have been arranged on the Periodic Table according to certain properties. - Contains only one type of atom Ex. Silver, Oxygen, Iron, Carbon, Mercury Compounds - Pure substances that contain two or more different elements, combined in a definite fixed proportion. - Can be broken down chemically into different substances since it is made up of different kinds of atoms Ex. Water - 2 hydrogen + 1 oxygen ---> H2O Salt - 1 sodium + 1 chlorine ---> NaCl MIXTURES contain at least two different substances. Properties are variable Solutions -Homogeneous are a mixture in which one substance is dissolved in another. It is a uniform mixture that appears the same throughout. Ex. Salt water, apple juice, and air Mechanical - Heterogeneous are mixtures that do not appear the same throughout. The different components are visible. Ex. Soil, chocolate chip cookies, chicken noodle soup Suspensions – Heterogeneous mixture made of large particles that are uniformly mixed but will settle if left undisturbed Ex. Sand in water, powdered chalk in water Colloids - Heterogeneous mixture composed of fine particles evenly distributed throughout a second substance Ex. Hair gel Emulsions - Type of colloid in which liquids are dispersed in liquids - Many will separate quickly to form layers of the original liquids (oil and vinegar) Lesson 2: Chemical and Physical Changes Physical Changes - No new substance produced - Change of size, shape or state - Ex. Cutting, freezing, molding, boiling, dissolving Chemical Changes - Starting material is used up - New substance formed with different properties - Atoms are rearranged to form new molecules - Changes cannot be reversed - Ex. Cooking an egg, rusting, burning Evidence of a Chemical Change (a) Color Change (b) Gas Formed – bubbling (c) Solid material, called a precipitate is formed. - two solutions are combined and a solid is formed (d) Energy Change - energy is the ability to do work - Ex. Light, heat, mechanical, sound, electrical There are 2 types of energy change: 1. Endothermic – energy is required. - Energy is added to the starting materials - Ex. cooking 2. Exothermic – energy is released. - Ex. Burning Identify the following substances as pure substances (element / compound) or as a mixture (homogeneous / heterogeneous): 1. zinc 6. vinegar 2. carbon dioxide 7. tossed salad 3. orange juice 8. aluminum 4. nitrogen 9. kool-aid 5. sugar 10. windex Lesson 3: History of the Atomic Theory Aristotle: (350 BC) – Greek philosopher - Believed that everything was made of 1. Earth (dry and cold) 2. Air (wet and hot) 3. Fire (dry and hot) 4. Water (wet and cold) Robert Boyle: (1660’s) – England Recognized that elements could be combined to form compounds Lavoisier: (1770-1780) – France 1. Defined elements as pure substances that cannot be decomposed (broken down into simpler substances) 2. Developed a system for naming chemicals, so that all scientists could use the same words 3. Identified 23 pure substances as elements 4. Discovered that in a chemical change, “the mass of the new substances is always the same as the mass of the original substances” – LAW OF CONSERVATION OF MASS.+ John Dalton: (1808) – England Atomic Theory: 1. All matter is made up of small particles called atoms 2. Atoms cannot be created, destroyed or divided into smaller particles 3. All atoms of the same element are identical in mass and size. Atoms of one element are different in mass and size from the atoms of other elements 4. Compounds are created when atoms of different elements link together in definite proportions. Dalton’s Theory led to the current definitions: Element – a pure substance made up of one type of particle, or atom. Compounds – pure substances that are made up of 2 or more elements chemically combined together. Compounds can be broken down into elements again by chemical means. J.J. Thompson: (1897) England - Raisin bun model (plum pudding) - Atom is a sphere, which is positive, with negative electrons embedded in it like raisins in a bun Ernest Rutherford: (1911) McGill University, Canada - Atoms have a nucleus which is positive - Most of the atom is empty space occupied by the moving negatively charged electrons - Proposed the existence of protons in a nucleus Neils Bohr: (1913) – Danish -Electrons move in circular orbits around the nucleus - Like a miniature solar system James Chadwick: (1932) - showed that the nucleus must contain heavy neutral particles to account for all of the atom's mass - proposed the existence of neutrons Lesson 4: Element Symbols - All elements have been given an atomic symbol (a) A single capital letter – O – oxygen (b) Capital letter & a lower case letter – Co – cobalt (c) Capital letter & 2 lower case letters – Uun – ununnilium - In the 1860’s Dmitri Mendeleev, a Russian chemist arranged the elements in order of increasing ATOMIC MASS and created the PERIODIC TABLE - ATOMIC MASS is the average mass of an atom of an element Ex. Oxygen = 16.00 g/mol - Mendeleev found that the properties of the elements repeated at definite, or periodic intervals (ex. Lithium, sodium and potassium have similar properties so he placed them in the same family or vertical row) - He left blanks in the table where he predicted elements should be and predicted what their properties would be, based on where they were on his table - After the development of atomic theory, the periodic table was rearranged in order of increasing ATOMIC NUMBER - ATOMIC NUMBER is the number of protons an element has in its nucleus Ex. Fluorine – atomic number = 9, therefore it must have 9 protons in its nucleus. The Periodic Table contains a lot of information about the different atoms. For example: Atomic number 4 Name Symbol Be Beryllium 9.01 Atomic Mass The horizontal rows on the periodic table are called Periods. The vertical rows on the table are called Groups or Families. Elements in the same family have similar properties (behave in a similar manner) There are 18 Groups or Families. The key on the periodic table will indicate the state of each element. Ex. White box – solid Grey box – gas Black box - liquid - All elements can be classified as metals, non- metals or metalloids depending on their properties Metals - Found to the left of the staircase line - 80% of all elements - Lustrous (shiny) - Ductile (stretched into wire) - Malleable (hammered/shaped) - Conduct electricity - All solids, except mercury liq. - Ex. Sodium, iron Non-metals - located to the right of of the staircase line - 20 % of all elements - dull - non-ductile - brittle - non-conductors - Mostly gases, some metals, 1 Ex. Oxygen, bromine Metalloids - these elements have properties of both metals and non-metals Ex. Silicon – shiny like a metal, poor conductor like a non-metal There are 4 special named groups in the table: Group 1 – Alkali Metals - Most reactive metals - Never found in pure form in nature Ex. Lithium, sodium, potassium Group 2 – Alkaline Earth Metals - React fairly vigorously with some substances Ex. Magnesium, calcium, barium Group 17 – Halogens - Most reactive non-metals Ex. Fluorine, Chlorine Group 18 – Noble Gases - Most non-reactive elements - Used to be called “Inert” gases until 1963 when a Canadian chemist at UBC, made some of them react - Different noble gases produce different colors Ex. Argon – blue Helium – yellow-white Lesson 5: ATOMIC STRUCTURE Atom - the smallest part of an element (which retains the chemical and physical properties of the element). Atoms are made up of 3 subatomic particles 1. Electron (e) -Smallest particle in an atom -Has a negative charge -Located in the extra nuclear region of the atom - outside the nucleus 2. Proton (p) -Has a large mass -Has a positive charge -Located inside the nucleus 3. Neutron (n) -Same mass as a proton -Has a neutral charge (no charge) -Located inside the nucleus Nuclear Notation - Atomic number is the number of protons in the nucleus - The number of protons equals the number of electrons in a neutral atom (#p = #e) - Atomic Mass Number is the total number of protons and neutrons in the nucleus Number of neutrons = mass # - atomic # And, Atomic # = #p = #e Example: Find the number of protons, electrons and neutrons for iron and sodium. Fe Atomic # = 26 Atomic mass = 55.85 = 56 (round the mass) Therefore: # of p = 26 # of e = 26 # of n = 56 – 26 = 30 Note: when finding the number of neutrons we round the atomic mass to the nearest whole number. Na Atomic # = 11 Mass # = 22.99 = 23 p = 11 e = 11 n = 23 – 11 = 12 Au Atomic # = 79 Atomic mass = 196.96 = 197 p = 79 e = 79 n = 197 – 79 = 118 Lesson 6: Bohr’s Model of the Atom - Bohr’s model states that electrons can be found only in certain energy levels or orbits around the nucleus - He also stated that only a certain maximum number of electrons are allowed in each orbit. Orbit # 1st 2nd 3rd 4th 5th 6th Max. # of electrons 2 8 8 18 18 32 When one orbit is filled the remaining electrons go to the next orbit – you cannot exceed the maximum allowed. We can draw the Bohr diagram for any element. It must have a nucleus showing the number of protons and neutrons and circles outside the nucleus showing the number of electrons. Reminder: # of protons = # of electrons = atomic # e.g. Draw the Bohr model for the following elements: a) Lithium Step 1 – Look up the atomic number It’s 3. So, # of p = 3 # of e = 3 Step 2 – Look up the atomic mass. It’s 6.94 = 7 (round to the nearest whole #) Find the number of neutrons. Reminder: # of n = atomic mass – atomic # So, # of n = 7 – 3 = 4 Step 3 – Draw the diagram. 1st orbit #p=3 #e=3 #n=4 Atomic # =3 Atomic Mass=7 P=3 n=4 The Bohr model diagram can be simplified – we can use lines instead of circles – this is called Electron Energy Level Representation. Electron Energy Level Representations (EELR) Ex. Zinc - Atomic # = 30 Mass # = 65 Therefore, #p = 30 #e = 30 #n = 65 – 30 = 35 12 8 8 2e P=30 n =35 These must add up to 30. Lesson 7: Molecules and Compounds Molecule – a particle formed when two or more atoms combine * can be 2 atoms which are the same ex. H2, or O2 OR * can be 2 or more different atoms ex. CuS, NaCl, CO Compound is a pure substance made of 2 or more different elements - Compounds can be broken down into simpler substances - ELECTROLYSIS – use of electricity to separate a chemical compound into its elements Ex. Water –broken down into hydrogen and oxygen - CHEMICAL BONDS hold elements together - If elements SHARE electrons to form a bond it is called a MOLECULAR BOND -- MOLECULAR ELEMENT or MOLECULAR COMPOUND – Non-metals share electrons to form molecular compounds. E.g. CO2 If atoms transfer electrons from one atom to another to form a bond it is called an IONIC BOND -- IONIC COMPOUND – metals transfer electrons to non-metals to form ionic compounds. E.g. NaCl - Sodium chloride - CHEMICAL FORMULAS use symbols and numbers - If only one atom is represented, no numbers are used - if there is more than one of that type of atom present a small number written below the line is used to tell us the number of that type of atom. This is called a SUBSCRIPT. E.g. H2O - One water molecule is made up of 2 atoms of Hydrogen and 1 Oxygen. NaCl - sodium chloride 1 – sodium 1 – chlorine H2 - hydrogen C12H22O11 - sucrose 12- carbon 22 – hydrogen 11 – oxygen Cu(NO3)2 - copper nitrate 1 – copper 2 – nitrogen 6 – oxygen 2 – hydrogen All pure substances can be identified in two ways: * Element or compound * Atom or molecule NaCl H2 C6H12O6 Zn Cu(NO3)2 - compound/molecule - element/molecule - compound/molecule - element/atom - compounds/molecule Lesson 8: Molecular Elements and Compounds Diatomic Elements – Molecules made of 2 atoms of the same element - All of the halogens, plus oxygen, hydrogen and nitrogen are diatomic elements H2, O2, F2 Cl2, N2, Br2 Molecular Compounds - Formed when 2 non-metals share electrons - Most molecular compounds have low melting and boiling points; therefore they are found as solids, liquids and gases at room temperature - They are poor conductors Naming Binary Molecular Compounds IUPAC – International Union of Pure and Applied Chemistry – Determines how compounds are named followed by chemists around the world Step 1. Write the entire name of the first element Step 2. Change the ending on the name of the second element to – ide Step 3. Use a prefix to indicate the number of each type of atom in the formula. Prefixes are: 1 – mono (only used for the second element) CO – carbon monoxide 2 – di 3 – tri 4 – tetra 5 – penta 6 – hexa 7 – hepta 8 – octa 9 – nona 10 – deca Ex. P2O5 – SiO SCl2 - diphosphorous pentaoxide silicon monoxide sulphur dichloride NO2 – Nitrogen dioxide N2O – Dinitrogen monoxide N2O4 – Dinitrogen tetraoxide Writing Molecular Formulas 1. Write the symbol for the elements in the same order as they appear in the name. 2. Use subscripts to indicate the numbers of each type of atom. Ex. Carbon tetrabromide - CBr4 Triarsenic hexasulphide - As3S6 We use small symbols in parentheses after the formula for each compound to indicate the state of matter (s) - solid - NaCl(s) (l) - liquid - H2O(l) (g) - gas - CO2 (g) Prefixes: 1 = mono 2 = di 3 = tri 4 = tetra 5 = penta 6 = hexa 7 = hepta 8 = octa 9 = nona 10 = deca MEMORIZE Assignment 11: Name_________________________ Name or give the formula: 1. Silicon dioxide ____________________________________ 2. Sulphur monoxide_________________________________ 3. OF2 _________________________________________ 4. SiBr4 ______________________________________ 5. PH3 _______________________________________ 6. N2O _______________________________________ 7. CO ________________________________________ 8. NBr3 _______________________________________ 9. P2I3 _________________________________________ 10. SO3 _________________________________________ 11. N2O4 _____________________________________________ 12. Tetraphosphorous hexaoxide _________________________________ 13. Dinitrogen tetraoxide _______________________________________ 14. Heptasilicon monobromide _____________________________________ 15. Octaboron decaiodide _________________________________________ 16. B2O3 ______________________________________________ 17. BrF7 ______________________________________________ 18. N3O6 _____________________________________________ 19. H2Cl5 _____________________________________________ 20. Triselenium diastatide _________________________________ 21. Diarsenic pentaoxide _____________________________________ 22. Sulphur trioxide _________________________________________ 23. C3O2 ___________________________________________________ 24. C2H6 ___________________________________________________ 25. As3Br7 __________________________________________________ 26. SO2 _____________________________________________________ 27. Selenium monoxide __________________________________________ 28. Diboron trioxide _____________________________________________ 29. PF3 ________________________________________________________ 30. P2O5 __________________________________________ __________________________________________________ 31. P4O10 ________________________________________ 32. Arsenic trifluoride ______________________________ 33. BrF7 __________________________________________ Lesson 9: Ionic Compounds - Atoms that gain or lose electrons to become stable are called IONS. - If they gain an electron they have more negative charges than positive charges so they have a slight negative charge - Non-metals gain electrons. - If they lose an electron they have less negative charges than positive charges so they have a slight positive charge - Metals lose electrons. Ionic Compounds - Made up of a metal bonded to a non-metal – electrons transfer from the metal to the non-metal. - All solids at room temperature - Separate into positive and negative ions when they dissolve in water - The ions conduct electricity Naming Ionic Compounds 1. The name includes both elements in the compound, with the name of the metallic element first. 2. The non-metallic element is second. Its ending is changed to – ide. 3. No prefixes used in naming. Ex. CaCl2 - calcium chloride (1 calcium/2 chlorine) Na2S - sodium sulphide (2 sodium/1 sulphur) Fe2O3 - iron(iii) oxide (rust) Lesson 10: Chemical Reactions Chemical Reactions: formation of a new substance Chemical bonds are broken and new bonds formed Reactants Starting materials Products become End materials Reactants: Any substance that is used up in the reaction Products: Any substance that is produced in the reaction Word equation: -Gives the names of all the reactants (separated by a + sign) -Arrow points to the names of all the products (separ. by + sign) Write Word Equations for the following reactions: (a) When sodium reacts with chlorine, sodium chloride (salt) is produced. Sodium + Chlorine Sodium chloride (b) Hydrogen gas and zinc chloride are produced when a piece of zinc metal is dropped into a beaker of hydrochloric acid. Zinc + Hydrochloric acid Hydrogen + Zinc chloride (c) Potassium iodide is decomposed to produce potassium metal and iodine. Potassium iodide Potassium + Iodine Law of Conservation of Mass: In a chemical reaction the total mass of the reactants is always equal to the total mass of the products. In a chemical reaction mass is neither gained nor lost. Molecules may be broken apart and new ones may be formed, but the atoms in the products are the same ones that were in the reactants Ex. Start with 100 grams of reactants - end up with 100 grams of product. Start out with 20 Hydrogen and 10 oxygen - end up with 20 hydrogen and 10 oxygen Using numbers called COEFFICIENTS in front of the elements and compounds in the reaction balances chemical reactions. 2 H2O --- 2 H2 + 1 O2 Reaction Rate – A measure of how fast a reaction occurs - Some reactions are naturally fast others are slow We can influence the rate in several ways: 1. Temperature – The higher the temperature the faster the rate - Molecules move faster at higher temperatures - Molecules collide more often and form new substances more quickly - The faster the rate, the less time needed for the reaction 2. Concentration - Refers to the amount of solute present in a specific amount of solution - the higher the concentration, the faster the reaction 3. Surface Area - For 2 substances to react, they must come into close contact - the greater the surface area, the more contact the two substances have therefore the faster the reaction - Can increase surface area by grinding up a chemical 4. Stirring - Increases the chances of collisions, therefore speeds up the reaction rate Lesson 11: Catalysts and Inhibitors Catalyst - A substance that speeds up the rate of a reaction (without being changed itself). Ex. saliva – acts as a catalyst to break down starch Inhibitors - Substances that slow down chemical reactions Ex. Added to some foods and medicines to slow down their decomposition Corrosion Corrosion is a chemical reaction. It is the “eating away” of a metal as it reacts with other substances in the environment. Corrosion of iron is called rusting 4 Fe(s) + 3 O2(g) iron + oxygen 2 Fe2O3(s) produces rust Different metals corrode at different rates. Iron corrodes quickly. Gold does not corrode at all. Aluminum and copper corrode only on the surface. Corroded materials lose their strength. Rusting is sped up by high temperature, surface area, and the presence of water, air, salt or acid. If the metal is totally protected from air or water, rusting cannot occur. Rust Protection 1. Paint. 2. Rust Check – spray with an oil film to keep air and water away 3. Galvanization - coat it with zinc - Zinc is more resistant to corrosion Ex. Galvanized nails Electroplating – covering a metal with another metal by using electrolysis -Ex. Bumpers are coated with a thin layer of chromium - chromium improves the hardness, stability, and appearance. 4. Combustion (Burning) - Chemical reaction that occurs when oxygen reacts with a substance to form a new substance - Combustion is exothermic - Oxygen is always one of the reactants – no oxygen no combustion (no fire). - Carbon dioxide and water vapor are the products of combustion when one of the reactants contains “carbon” Ex. Methane - CH4 CH4 + 2 O2 Methane + oxygen CO2 + 2 H2O Carbon dioxide + water Identification Tests for the Products of Combustion 1. Test for carbon dioxide - Bubble the gas through limewater solution - If the limewater turns milky, the gas is carbon dioxide 2. Test of water - Touch the cobalt (II) chloride paper to the liquid - If the paper turns pink, the liquid is water