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Unit 1: Chemistry (P. 134-279)

Unit 1: Chemistry (P. 134-279). Patterns and Compounds Periodic Table, Naming, Balancing Equations Chemical Reactions Energy, 4 Types, Combustion Acids and Bases Properties, pH, Reactions Chemical Reactions in the Environment Factors affecting Rates, Chemicals and Us.

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Unit 1: Chemistry (P. 134-279)

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  1. Unit 1: Chemistry (P. 134-279) • Patterns and Compounds • Periodic Table, Naming, Balancing Equations • Chemical Reactions • Energy, 4 Types, Combustion • Acids and Bases • Properties, pH, Reactions • Chemical Reactions in the Environment • Factors affecting Rates, Chemicals and Us

  2. Classification of Matter

  3. Classification of Matter • PURE SUBSTANCE:A substance with a fixed composition and constant properties • ELEMENT: A substance that cannot be broken down into simpler substances by chemical means. Atoms are the simplest particles that cannot be broken down by chemical means • COMPOUND: A substance that is made up of two or more different atoms (molecules). These substances can be broken down only by chemical means. • MIXTURE: A mixture consists of two or more kinds of matter, each keeping its own characteristic properties. • SOLUTIONS: A mixture that is homogeneous. If the solution is a liquid or gas, it is transparent. • MECHANICAL MIXTURE: A heterogeneous mixture with parts that are visibly distinguishable

  4. MetalsMetalloidsNonmetals

  5. Metals vs. Non Metals • Shiny • Ductile • Malleable • Conducts Electricity • Conducts Heat • Dull • Brittle • Does Not Conduct Electricity • Does Not Conduct Heat

  6. Elements and the Periodic Table:A Review • Every Element has a unique • Name • Symbol • Atomic Mass Number (A) • Represents the number of protons + number of neutrons • Atomic Number (Z) • Represents the number of protons and the number of electrons in a neutral atom • The number of neutrons can be calculated by subtracting A - Z

  7. Atomic Number Atomic Mass Number of electrons, protons 2 Atomic Number He 4 Number of Protons and Neutrons Atomic Mass Number of Neutrons = Atomic Mass – Atomic Number

  8. Example

  9. Grouping of Elements • Elements are subdivided into: • “groups” or "families" (vertical columns) • and “periods” (horizontal rows) • Metals elements are on the left • Non-metal elements on the right • separated by a dark "staircase line". • Elements bordering this division line exhibit some properties of both metals and non-metals and are called metalloids. • Copy table 5.1 on page 140 into your notes.

  10. Group Names Alkali Metals Alkaline Earth Metals Group 3 Carbon Group Nitrogen Group Oxygen Group Halogens Noble Gases

  11. Bohr-Rutherford Diagrams • The following information is required: • 1. Number of Electrons: • Is the same as the number of protons in a neutral atom. The electrons are organized into shells in the following order. • up to 2 electrons in the first shell • up to 8 electrons in the second shell • up to 18 electrons in the third shell • up to 32 electrons in the fourth shell • 2. Number of Protons • Is the same number as the atomic number • 3. Number of Neutrons • Can be determined by subtracting the atomic mass from the atomic number

  12. Bohr Rutherford Diagram Nucleus with protons and Neutrons 18p 36n Electron Orbits (shells) with a 2,8,8, pattern

  13. Predicting Chemical Reactivity • Elements with 8 electrons in their outer energy level appear to have a special significance. Elements with this arrangement do not react easily and are considered stable. • All noble gases (Neon, Krypton, Xenon, Radon) have 8 electrons in their outer energy level and are very non-reactive elements (Helium is a special gas that is very stable with 2 electrons in its first level). • All elements want to be stable and therefore want to gain or lose electrons in order to achieve a stable 8 configuration (Stable Octet). • Whenever an atom gains or loses electrons they become negative or positive and they are called ions.

  14. Two main factors determine chemical activity(reactivity) • 1) The number of electrons in the outer energy level • i) Elements with 1-3 electrons on outer level lose electrons (become positive) • ii) Elements with 5-7 electrons in outer level gain electrons (become negative) • iii) Elements with 4 electrons in outer level are special (tend to become positive) • 2) The number of energy levels • As the number of energy levels increase, the attraction between those electrons in the outermost energy level and the positive nucleus decrease

  15. Ions: To gain or lose an Electron • Positively Charged: Cations • When a neutral atom gives up one or more electrons, the positively charged ion that results is called a Cation. • For example: • Negatively Charged: Anions • When a neutral atom gains one or more electrons, the negatively charged ion that results is called an Anion. • For example:

  16. Electron Dot Diagrams • A Bohr-Rutherford diagram represents an atom and all its electrons. • A simpler way to represent atoms and ions of atom is with electron dot diagrams • Electron Dot Diagrams show only the outer energy level (valence shell) of an atom. Only these electrons are represented because they are responsible for an atom’s chemical properties. For example:

  17. Lewis Dot / Electron Dot diagrams N C

  18. Chemical Bonds: Forming Compounds • Most substances on earth do not exist as elements, they are composed of two or more different elements joined together to make compounds. • When two atoms collide, valence electrons on each atom interact. A chemical bond forms between them if the new arrangement of their valence electrons have less energy than their previous arrangement. • For many atoms that new arrangement of their electrons will be that of their closest noble gas. • Atoms may acquire a valence shell like that of its closest noble gas in one of three ways: • 1. An atom may give up electrons and forma ion • 2. An atom may gain electrons and form an ion • 3. An atom may share electrons

  19. Ionic Compounds Substances held together by ionic bonds are called Ionic compounds e.g. NaCl, KCl. Ionic Bonds occur because of the attraction of cations and anion for each other. Electrons are transferred between the atoms during bond formation. Properties include: • High melting point (i.e. strong bonds) • Conduct electricity when dissolved in water or molten • Form crystal lattice structures • Soluble in water

  20. Molecular Compounds • Substances that are composed of molecules are called molecular compounds. Many non-metals form compounds with other non-metals. When this occurs there is no transfer of electrons between the two atoms instead they share electrons forming a covalent bond. • Although bond between atoms are strong, bonds between molecules are weak. eg. Moth crystals, nitrogen gas etc. • Properties Include: • Low melting and boiling points • Often have an odour • Don’t conduct heat • Don’t conduct electricity (non-electrolytes) • Diatomic molecules (e.g. O2, F2 etc.) are also the • result of covalent bonds.

  21. Chemical Naming and Formulas • Binary Ionic • Transition Metals • Stock versus Classical • Polyatomic Ions • Binary Molecular

  22. General Rules • The Metal is always written first • The nonmetal suffix in a compound is either “ide” or “ate” • Every compound must be electrically neutral • All Positive charges must equal Negative charges

  23. Binary Compounds: Formula to Name • Composed of two Elements • One metal and one nonmetal • Write the name of the metal first unchanged • Write the name of the nonmetal second • Change the ending to an “ide” • LiCl  Lithium Chloride • MgI2 Magnesium iodide

  24. Binary Compounds: Name to Formula • Write the symbol for each element with the metal written first • Find the ionic charge for each element • Cross the number value of the charge and place it as the subscript of the other element • Reduce the values to lowest ratio • Magnesium Oxide  Mg2+ O2- Mg2O2 MgO

  25. Transition Metals: Groups 3-12+Name to formula • Almost all are able to form more than one cation • When writing the formula the charge of the metal cation will be indicated by roman numerals after the metal • Lead (III) chloride  PbCl3 • Iron (II) oxide  FeO

  26. Transition Metals: Groups 3-12+Formula to Name • Finding the charge on the metal can be done two ways • Reverse Cross-Over Method • The subscript of the nonmetal becomes the charge of the metal • Sometimes the charge is misleading • Charge Balancing • Charge = Subcript of the nonmetal multiplied by the charge of the nonmetal divided by the subscript of the metal

  27. Chemical Equations and Reactions A chemical equation is a description of a chemical reaction using chemical symbols, not words Steps: 1) The reactants are written first 2) The products are written second 3) The state for each atom is indicated (g) gas, (s) solid, (l) liquid, (aq) aqueous 4) The reactants and products are separated by an "arrow" (  ) e.g. Word Equation Hydrogen gas plus chlorine gas produces hydrogen chlorine gas e.g. Chemical Equation H2(g) + Cl2(g)  HCl(g)

  28. Balanced and Unbalanced Chemical Equations The Law of Conservation of Mass states: Matter cannot be created or destroyed; it can only be changed from one form to another. Therefore, the number of atoms in the reactants must equal the number of atoms in the products An unbalanced or skeleton equation does not follow the Law of Conservation of Mass. The number of atoms on the left side (reactants) does not equal the atoms on the right side (products) e.g. H2(g) + Cl2(g)  HCl(g) 4 atoms (2 H, 2 Cl) 2 atoms(1 H, 1 Cl)

  29. A balanced chemical equation follows the Law of Conservation of Mass. The number of atoms on the left side (reactants) equals the atoms on the right side (products) e.g. 1H2(g) + 1Cl2(g)  2HCl(g) 4 atoms (2 H, 2 Cl) 4 atoms(2 H, 2 Cl)

  30. Writing Balanced Chemical Equations 1. Write the chemical formula for each reactant and product followed by the state of each: solid (s); liquid (l); gas (g); aqueous(aq) 2. Adjust the numbers of molecules until there are the same number of atoms of each type on both sides of the equation. This balances the mass of both the reactants and products. 3. Usually, balancing is easiest when hydrogen and oxygen atoms are left until the end NOTE: Do not change the subscript in a formula to balance an equation. Changing these numbers changes the molecular structure of the molecule.

  31. Energy Changes and Chemical Reactions Chemical reactions, physical changes of state and dissolving processes often involve energy changes. Exothermic Processes: Processes that release energy (e.g. heat and light) and increase the temperature of the surroundings. Endothermic Processes: Processes that absorb energy and decrease the temperature of the surroundings.

  32. Factors Affecting Chemical Reaction Rate The Rate of Reaction is defined as: The time it takes for a given product to form, or for given amounts of reactant to react. Reaction rate is determined by: i. Measuring how fast reactants are used up. ii. Measuring how fast the products are formed. Factors affecting Reaction Rate 1. Concentration and Reaction Rate  Concentration (amount of substance in a given volume)  Rate 2. Surface Area and Reaction Rate  Surface Area (area exposed)  Rate 3. Temperature and Reaction rate  Temperature  Rate

  33. 4. Catalysts and Reaction Rates A Catalyst is defined as: A substance that speeds up the rate of a chemical reaction without being used up in the reaction. Catalyst lower the energy required to break the bonds that hold substances together. Examples include: enzymes (biological catalysts), platinum, rhodium and palladium (chemical catalyst used in catalytic converters)

  34. Types of Chemical Reactions There are four basic patterns that most chemical reactions follow: 1) Synthesis Reactions This type of reaction fits the general pattern: A + B  AB e.g. N2(g) + 3H2(g)  2NH3(g) CaO(s) + H2O(l)  Ca(OH)2 A synthesis reaction involves the formation of a new compound from simpler elements or compounds Combustion reactions (involving the reaction with O2) are examples of Synthesis Reactions e.g. Cu(s) + O2(g)  2CuO(s) Mg(s) + O2(g)  2MgO(s)

  35. 2) Decomposition Reactions These type of reactions are opposite to direct combinations. They fit the general pattern: AB  A + B e.g.CuCO3(s)  CuO(s) + CO2(g) 2KClO(s)  2KCl(s) + 3O2(g) A decomposition reaction involves the breaking down of a compound into simpler compounds or elements

  36. 3) Single Displacement Reactions A single displacement or substitution reaction fits the general pattern of: A + BC  AC + B This type of reaction involves a change in partners. One element displaces or knocks off another element in a compound.. e.g. Zn(s) + 2HCl(aq)  ZnCl2(aq) + H2(g) 3C(s) + Fe2O3(s)  3CO(g) + 2Fe(s)

  37. 4) Double Displacement Reactions A double displacement reaction fits the following general pattern: AB + CD  AD + CB This type of reaction involves a change of both partners. The cation (positive element or polyatomic ion) of one compound changes place with the cation of the second compound. e.g. Na2S(aq) + ZnCl2 (aq)  ZnS(s) + 2NaCl(aq) AgNO3(aq) + KBr(aq)  AgBr(s) + KNO3(aq) SF4(s) + 2H2O(l)  SO2(g) + 4HF(aq)

  38. Carbon Chemistry Organic Chemistry: The study of carbon containing compounds and their properties e.g. hydrocarbons When hydrocarbons (contain carbon and hydrogen) are burned in enough oxygen complete combustion occurs. Hydrocarbon + oxygen gas  carbon dioxide + water + E (good supply) If hydrocarbons are burned in a poor supply of oxygen, incomplete combustion occurs. Hydrocarbon + oxygen gas  carbon dioxide + water + E (poor supply) + carbon monoxide + residue

  39. Classification of Substances by Their Behaviour The process of grouping substances according to common properties is called classification. Previously we have classified substances according to: i) State (e.g. solid, liquid or gas) ii) Composition (e.g. pure substances, mixtures etc.) Matter can also be classified by chemical behaviour. Acids and bases make up two classes of compounds that have been classified by their chemical behaviour.

  40. Acids and Bases Acids: An acid is a compound that dissolves in water to produce hydrogen ions (H +) in solution. e.g. HCl Bases: A base is a compound that dissolves in water to produce hydroxide ions in solution (OH -) e.g. NaOH Copy Table 7.3 “Acids and Bases: A Summary” found on page 230 in your text.

  41. Preparation of Common Acids A common way to prepare an acid is to react a nonmetal oxide with water. An oxide is an element combined with only oxygen e.g. sulphur trioxide + water  sulphuric acid carbon dioxide + water  carbonic acid Some common acids in the laboratory include: i) sulfuric acid ( H2S04 ) ii) nitric acid (HNO3) iii) hydrochloric acid (HCl) iv) acetic acid, (CH3COOH) Other common acids include: i) acetylsalicylic acid (aspirin) ii) ascorbic acid (vitamin C) iii) carbonic acid (carbonated soft drinks)

  42. Preparation of Common Bases A common way to prepare a base is to react a metal oxide with water. e.g. sodium oxide + water  sodium hydroxide calcium oxide + water  calcium hydroxide Some common bases in the laboratory include: i) Sodium hydroxide (NaOH) ii) Calcium hydroxide (Ca(OH)2) iii) Potassium hydroxide (KOH) iv) Magnesium hydroxide (Mg(OH)2)

  43. Indicators An indicator is a chemical that changes colour as the concentration of H+ (aq) and OH- (aq) changes. e.g. i) Litmus: • blue litmus turns red in acid • red litmus turns blue in base ii) Phenolphthalein • turns pink in base Indicators can be made from flowers, fruits, vegetables, leaves (e.g red cabbage, tea etc.) Synthetic Indicators are more easy to use than natural indicators because they: • last longer than natural indicators • can be produced in large quantities e.g. bromothymol blue (BTB) phenolphthalein methyl orange methylene blue

  44. The pH Scale The pH scale describes the "strength of the hydrogen ion (H+)". The scale is numbered from 0 to 14 • acids have a pH less than 7 [H+] > [OH-] • bases have a pH more than 7 [H+] < [OH-] • neutral substances have a pH of 7 [H+]= [OH-] The change in 1 pH unit represents a tenfold increase in the concentration of hydrogen ions in solution. e.g. A pH of 2 is 10 x's stronger than a pH of 3 A pH of 2 is stronger than a pH of 5 A pH of 2 is _ stronger than a pH of 7 pH can be estimated using pH paper or measured using a pH meter (measures electric properties)

  45. The Strength Of Acids And Bases The strength of an acid or base is dependant on two factors: 1. Concentration The concentration of an acid or base is the amount of the pure substance dissolved in 1 L of water. 2. Ionization When acids and bases are dissolved in water, they ionize (break apart into charged particles). The term “Percent Ionization” refers to the number of molecules that will ionize for every 100 molecules that dissolve. e.g. HCl + H2O  H3O++ Cl- Solutions that form ions in water are called electrolytes. Electrolytes conduct electricity. The higher the concentration of ions the stronger the electrolyte.

  46. The Strength of Acids Strong acids: ionize completely in water e.g H2SO4 Weak acids: ionize partially in water e.g. CH3COOH The Strength of Bases Strong Bases: ionize completely in water e.g NaOH Weak Bases: ionize partially in water e.g. NH3

  47. Neutralization Neutralization occurs when hydroxide ions (base) and hydrogen ions (acid) are mixed to make water. The general word equation is: Acid + Base  Water + Salt e.g hydrochloric + sodium  water + sodium chloride acid hydroxide (HCl) (aq) + ( NaOH)(aq)  ( H2O)(l) + ( NaCl) (aq) After neutralization, the solution no longer has a high concentration of either ion.

  48. Soaps and Detergents What makes up soap ? 1. fatty acid (lipid) 2. strong base (NaOH) The word equation is: fat + base  soap + glycerol Soap curds cling as scum to whatever it comes into contact with, and does not rinse away easily. This problem led to the development of synthetic detergents called syndets. Advantages include: 1. good at removing dirt 2. more soluble in water 3. prevented dirt from collecting back onto clothes 4. did not form a curd 5. mild to hands and fine fabrics 6. less expensive (made from plant oils and animal fats)

  49. How soap cleans A soap or detergent molecule consists of two ends: 1. Hydrophillic (water loving) The end with the sodium ion is attracted to water and becomes soluble 2. Hydrophobic (water hating) Hydrocarbon end is attracted to insoluble dirt (grease) on clothes etc. For example:

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