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Chemistry

Chemistry. Chapters 5-8. What is chemistry?. Chemistry = The study of matter, its properties, and how it changes. Matter = Anything that has mass + takes up space. Pure substances = all particles that make up a

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Chemistry

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  1. Chemistry Chapters 5-8

  2. What is chemistry? • Chemistry = The study of matter, its properties, and how it changes. • Matter = Anything that has mass + takes up space. • Pure substances = all particles that make up a substance are the same and have constant properties. (Ex: pure water is colorless, freezes at 0 Celsius etc) • Elements = pure substances that can’t be broken down into simpler substances. (Ex: any of the individual “boxes” on the periodic table.)

  3. The Periodic Table: • Where are the metals/non-metals found? • Which elements are solid, liquid, gases? • Where/what are the metalloids? • Why is Hydrogen separated from the others? • Where are the transition metals? • Where do the names/symbols of the elements come from? • What’s up with those really weird named ones? • Why are their blanks and questions marks on some? • What do the numbers in the boxes mean? • Why is chemistry so awesome and exciting?

  4. Use the periodic table at the back of your book as a guide: • On your blank periodic table, label the following: • 1) Metals-color the boxes green 2) Non-metals-color the boxes orange • 3) Metalloids-color the boxes purple 4) Transition metals-label columns/groups 3-12 • 5) Alkali metals-label column/group 1 6) Alkaline earth metals-label column/group 2 • 7) Halogens-label column/group 17 8) Noble gases-label column/group 18 • 9) Solids –symbols in pencil or black 10) Liquids-write the symbols in blue • 11) Gases-write the symbols in red 12) Hydrogen-(color the box yellow) • *First letter of element symbol is capitalized; others lower case. • *Make sure your labels are clear and your work is neat. • *Put your name on it and hand in to me when finished.

  5. The Periodic Table: • The periodic table is a logical arrangement of the elements based on their properties. It allows us to explain and predict the physical and chemical properties of the elements based on their position in the table. • Families (a.k.a. groups) = elements grouped into the same vertical column, they share similar physical and chemical properties. (Ex: alkali metals (column #1), halogens (column 17 etc). • Periods = elements in the same horizontal row (they don’t share similar properties)

  6. Elements and atomic structure (pg 185): • Protons = heavy, positively charged particles found in the nucleus of atoms • Neutrons = heavy particles with no charge, found in the nucleus of atoms • Electrons = negatively charged particles with almost no mass, “circle” the nucleus at different energy levels • Bohr diagrams = show the arrangement of electrons in orbits

  7. Atomic number = the # of protons an element has • # of protons = the # of electrons for an element • Finding the number of neutrons: • A) First, round the atomic mass to nearest whole # • B) Then, subtract the # of protons (atomic number) • Electrons “circle” atoms at different levels (orbits). The first orbit has a maximum of two electrons, the second has a maximum of 8, the third also has a maximum of 8. This is the 2, 8, 8 rule.

  8. Use this information to complete the handouts provided.

  9. Practice/review questions: • Atomic Structure Worksheet (page 18 of package) • Bohr diagrams worksheet (page 19 of package) • Understanding chemical reactions (pages 20-25) • Periodic table crossword puzzle • Element bingo

  10. Compounds = substances made of 2 or more different elements. (Ex: CO2, H2O, NaCl etc) • Physical change = change in the size or form of a substance that does not change its chemical properties. (Ex: water freezing, ice melting etc.) • Chemical change = change into a new substance. (Ex: burning something.)

  11. How do elements form compounds? • Electrons in the outermost orbit of atoms have the most energy and are involved in bonding. These are called valence electrons. • Atoms “like” to have their outermost orbit full in order to become stable. In order to achieve this, an atom will either lose or gain valence electrons. • Metals lose (give away) electrons while non-metals gain (accept) electrons when they come into contact with another element. • When atoms no longer have an equal number of protons and electrons, they have either a positive or negative charge and are called ions. • *With Mr. Jessome’s help, label your periodic table with the ion charges for families 1, 2, and 13-18. • Page 191: #’s 1 and 2

  12. Ionic compounds: • Made up of metals and non metals joined together. • Ex: NaCl, Fe2O3. Their names have two parts: • 1) Name of the metal • 2) Name of the non-metal with the ending changed to “ide” • Ex: NaCl is named: Sodium chloride, Fe2O3 is named: Iron oxide • The metal has a positive charge; non-metal has a negative charge. • Attractive force of these opposites hold the compound together. The opposite charges must be equal (balance each other out). • Ex: MgCl2 • There are two Cl atoms because Mg has a charge of 2+ and Cl only has a charge of 1-, so two are needed because 2 x 1- = 2- • *Also, see page 193 for these steps, or 194 for the criss-cross method.

  13. Multi-valent metals: • “Multi” means “multiple” or “more than one”. • “valent” refers to valence electrons and therefore its charge. • Multi-valent metals can have more than one possible ion charge. • Example: Iron (Fe) has two possible charges; 3+ or 2+ • We assume the common ion charge is used (listed on top/first), unless Roman numerals in the name tell us otherwise. • Ex: The common ion charge for Iron is 3+, so to indicate a charge of 2+, you would name the ion like this: Iron (ll) oxide • Roman numeral (ll) indicates 2+ ion, instead of normally used 3+

  14. Practice questions: • Handout: Bohr diagrams, ions transfer etc • Ionic compounds practice sheet • Page 195: 1-10

  15. 1. Write the chemical formulas for the following ionic compounds: • a) Silver (I) iodide b) Gold (III) chloride c) Tin sulfide • d) Calcium nitride e) Sodium bromide f) Iron (III) oxide • g) Lithium fluoride h) Magnesium phosphidei) Potassium chloride j) Lead phosphide • 2. Write the chemical name for the following ionic compounds: • a) CoSe b) HgF2 c) Na3P • d) Ba3As2 e) CaBr2 f) Al2O3 • g) CuF h) MgCl2i) SnS2 j) FeBr2

  16. Practice questions: • Page 214, #’s: 1 (a, b, c) 2 (a, b, c, d, e, f) 3 (a, b, c, d, e, f) 6- all 9 - all

  17. Chemistry Bingo (see page 213): • Alkali metals • Alkaline earth metals • Bohr diagram • Chemical change • Chemical family • Chemistry • Combining capacity • Compound • Element • Halogens • Ion • Ionic charge • Ionic compound • Matter • Molecular compound • Electron • Neutron • Noble gases • Organic compound • Periodic table • Physical change • Polyatomic ion • Product • Proton • Pure substance • Reactant • Valence

  18. Polyatomic ions: • “Poly” = “more than one” or “several”. Ions are atoms with a charge. • Polyatomic ions = “several atoms together with an overall charge” • Examples: Hydroxide: OH, which has a charge of 1- • Carbonate: CO3, which has a charge of 2- • The same rules apply when writing formulas and names. When multiplying atoms to balance the opposite charges, brackets around the polyatomic part show that several atoms are being multiplied. • Ex: Magnesium hydroxide is written as Mg(OH)2 to show that both the oxygen and hydrogen atoms are being multiplied.

  19. Polyatomic ions: • Page 198: #’s 1, 3, 4 • Page 214: # 10 • Page 330: #’s 4, 12, 13

  20. Molecular compounds: • Made up of non-metals joined together. • Share their electrons equally in the form of covalent bonds. • Co = “equal”, valent = “outer electrons” • (*Ionic compounds have unequal sharing and ionic bonds) • Covalent bonds makes them stable, and neutral in charge (no ions). • In our everyday lives, so we have many common names for them. • Ex: H2O = water, NH3 = ammonia • Also scientific names: carbon monoxide(CO), carbon dioxide(CO2) • Scientific names: • A) end in “ide” • B) prefixes tell us how many of each atom make up the compound.

  21. Naming Molecular Compounds: • Prefixes and common names to know: • 1= Mono 2= Di 3= Tri 4= Tetra • 5 = Penta 6= Hexa 7= Hepta 8= Octa • Water = H2O Methane = CH4 Ammonia = NH3 • Ozone = O3 H2, F2, I2, Cl2, O2 = (element name) gas • Make sure you identify them as non-metals (molecular compounds). • Use the prefixes to create the name. • Using the name, you can easily write the formula (by knowing prefixes). • These compounds do not always have charges that balance like ionic compounds do-so learn the prefixes and common names!

  22. Practice questions: • Page 204: #’s 1,2,5 • Page 215: # 11 • Page 330: # 14

  23. Answers to the questions: • Page 204: • 1: Ionic = metal and non-metal, molecular = non-metals • 2. a) Non-metal atoms, b) share their electrons, c) covalent bonds • 5. A) Carbon tetrabromide B) Nickel triiodide C) oxygen diflouride D) Tin tetrachloride

  24. Answers to the questions: • Page 215: 11. A) CO • B) NI3 • C) Sulfur dichloride • D) Carbon tetrachloride • Page 330: 14. A) PCl3 • B) SO3 • C) Carbon tetrabromide • D) Nitrogen dioxide

  25. Try these: • a) NF3 b) NO c) NO2 d) B2O3 e) N2O • f) N2O4 g) PCl3 h) PCl5i) SF6 j) CO2 k) Phosphorus triiodide l) sulfur dichloride m) xenon trioxide n) Dinitrogentetraflouride o) sulfur tetraflouride p) boron trichloride q) Phosphorus pentafluoride

  26. More practice: • Complete the practice worksheet on molecular compounds. • Complete the review assignment with both ionic and molecular compound naming practice.

  27. Counting atoms: • In order to write proper chemical equations we must first learn to correctly count the numbers of atoms being shown in compounds. • (Follow along with Mr. Jessome to complete the worksheet: “How To Count Atoms Review”, and then try the “Counting Atoms Worksheet” on your own).

  28. Types of Chemical equations: A) Word equations - Show the reactants and products: Methane + oxygen gas  carbon dioxide + water B) Skeleton equations - Reactants and products are replaced with chemical formulas: CH4 + O2 CO2 + H2O C)Balanced chemical equations -Skeleton equations are “fixed” to show equal #’s of atoms on each side of equation: 1CH4 + 2O2 1CO2 + 2H2O

  29. Chemical equations: • Rules: • 1) You can only add coefficients to make the number of atoms balance. You can’t add subscripts, and cannot change the chemical formulas. • 2) It’s usually easiest to start with the most complicated looking group (most atoms). Leave the individual elements for last (such as O2, H2 etc). • 3) Use lowest possible whole number coefficients.

  30. Chemical equations: • Try examples on pages: • 219: #’s 2, 3 and 4 • 229: #’s 1, 2, 3 and 4 (a + b) • 252: #’s 4, 5 and 6

  31. The Law of Conservation of Mass: • “In a chemical reaction, the total mass of reactants is always equal to the total mass of the products”. • In chemical reactions, atoms are not created or destroyed; they are simply rearranged. The atoms present before a reaction occurs are still there after a reaction has occurred. • This is why we balance chemical equations.

  32. Evidences of Chemical Changes/Reactions: • A) A new color appears • B) Heat or light is given off • C) A gas is given off • D) A precipitate (solid) is formed *You will see such things happen when we do the lab

  33. Types of chemical reactions: • A) Synthesis (a.k.a. combination) reactions: • The combining of smaller molecules into larger molecules. • General formula: A + B  AB • Example: H2 + O2  H2O • B) Decomposition reactions: • The splitting of a larger molecule into smaller molecules. • General formula: AB  A + B • Example: H2O  H2 + O2

  34. C) Combustion reactions: • The burning of something, in the presence of oxygen • General formula: Fuel + oxygen  oxides + energy • Example: C3H8 + O2  CO2 + H2O + heat • (Propane plus oxygen produces carbon dioxide, water vapor and heat) • D) Single displacement reactions: • An element and a compound as reactants. One element displaces/replaces another element from a compound. • General formula: A + BC  AC + B (A and B switched) • Example: Br2 + CaI2  CaBr2 + I2

  35. Types of chemical reactions: • E) Double displacement reactions: • Elements in different compounds displacing/replacing another each other to form new compounds. • General formula: AB + CD  CB + AD (A and C switch) • Example: PbCl2 + KI  KCl + PbI2 • Page 252: #4 • Page 247: # 4 • Page 331: #16 (and if you are bored, #17)

  36. Answers to the questions: • Page 252, #4: • A) Combustion B) Single displacement • C) Decomposition D) Synthesis, combustion • E) Double displacement F) Synthesis • Page 247, #4: • A) Synthesis B) Single displacement • C) Double displacement D) Decomposition • E) Decomposition F) Single displacement • G) Synthesis

  37. Page 331, #16: • A) Single displacement B) Decomposition • C) Double displacement D) Combustion • E) Synthesis

  38. What to know for the test: • Chemistry vocabulary: (pure substances, elements, physical vs chemical changes etc –our bingo words from your notes) • Periodic table (metals vs non-metals, alkali metals, halogens etc-labeled periodic table) • Bohr diagrams (How to find the # of electrons, protons, neutrons for atoms) • Ionic compounds (naming them and writing formulas-includes polyatomic ions and multi-valent metals) • Molecular compounds (naming them and writing formulas-know your prefixes and common names) • Balancing chemical equations • *Not on the test: evidence of chemical reactions, 5 types of reactions (we will have a quiz on these next week, after some demos and a lab)

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