Unit 3: Chemical Processes

1. Unit 3: Chemical Processes Section 1: Chemicals in Action

2. Chemistry – the study of matter, its properties, and its changes or transformations • Matter – anything that has mass and takes up space • Matter has both physical and chemical properties

3. Matter • Matter is either classified as a pure substance or as a mixture • Pure substance – one in which all the particles that make up the substance are the same • E.g. Pure Water • Pure substances can be further classified as elements or compounds

4. Elements are pure substances that cannot be broken down into simpler substances (made of only one type of atom) • E.g. oxygen, hydrogen, … • Compounds are pure substances that are made of two or more different elements in a fixed proportion • E.g. carbon dioxide, sodium chloride, …

5. Properties of Matter • Matter has both physical and chemical properties • Physical property – is a characteristic of a substance (color, hardness, odour, solubility, state, melting and boiling point, size) • A physical change is a change in the size or form of a substance, which does not change the chemical properties of the substance

6. Chemical Property – is a characteristic behaviour that occurs when a substance changes to a new substance • A chemical change is also called a chemical reaction. It involves changing one substance into one or more different substances with different properties • The starting materials in a chemical reaction are called the reactants and the new materials are called the products.

7. Chemical tests – are distinctive chemical reactions that can be used to identify an unknown substance

8. WHMIS, HHPS, and MSDS • WHMIS, HHPS, and MSDS are all symbols that are used to identify hazardous chemicals and the precautions that should be used when working with these chemicals. • WHMIS – Workplace Hazardous Materials Information System • HHPS – Hazardous Household Product Symbols • MSDS – Materials Safety Data Sheet

9. Homework • Questions 1, 2, 4, and 5 page 175.

10. Elements and the Periodic Table • Periodic Table – a structured arrangement of elements • Elements on the periodic table are grouped into chemical families – groups of elements in the same vertical column of the periodic table. They tend to share similar physical and chemical properties.

11. Alkali metals – group 1 elements • Alkaline earth metals – group 2 elements • Noble gases – group 18 elements • Halogens – group 17 elements

12. Elements and Atomic Structure • Atoms are composed of three subatomic particles (protons, neutrons, electrons) • Protons – positively charge particles found in the nucleus of the atom. Each proton has a mass of 1 • Neutrons – neutral particles found in the nucleus of the atom. Each neutron has a mass of 1 • Electrons – negatively charged particles with almost no mass and can be found in orbits surrounding the nucleus

13. The number of electrons in an atom is equal to the number of protons. • Electrons that are in the outer orbit (valence shell) have greater energy levels and a greater probability of being involved in a chemical reaction. These electrons are called valence electrons.

14. Bohr – Rutherford Diagrams • Each orbit around the nucleus of an atom contain a specific number of electrons. • 1st orbit = 2 electrons • 2nd orbit = 8 electrons • 3rd orbit = 8 electrons • The electrons fill the orbits in order. Each orbit must be filled completely before electrons can be placed in a new orbit.

15. Sample Bohr – Rutherford Diagrams • Create a diagram for Nitrogen 1. Determine the number of electrons (7) 2. Draw the first orbit with the maximum number of electrons allowed (2) 3. Draw the second orbit with the remaining electrons (5) 4. Draw the nucleus with the protons and neutrons 7p 7n

16. You Try It (1) • Create a Bohr – Rutherford diagram for Magnesium

17. Answer • Determine the number of electrons, 12 • Draw the first orbit (2) • Draw the second orbit (8) • Draw the third orbit (2) • Draw the nucleus with the protons and neutrons 12p 12n

18. You Try It (2) • Create a diagram for Sodium

19. Answer • Determine the number of electrons (11) • Draw the first orbital (2) • Draw the second orbital (8) • Draw the third orbital (1) • Draw the nucleus with the protons and neutrons.

20. You Try It (3) • Create a diagram for Neon

21. Answer • Determine the number of electrons (10) • Draw the first orbit (2) • Draw the second orbit (8) • Draw the nucleus with the protons and neutrons. 10p 10n

22. The noble gases (column 18) are considered to be fairly stable. • Look at the arrangement of the electrons in the noble gases by drawing a Bohr – Rutherford diagram for He, Ne, and Ar. What do you notice about the arrangement of the electrons?

23. In some compounds, electrons are transferred from one atom to another to create a stable electron arrangement (like the noble gases) • Create a Bohr – Rutherford Diagram for Lithium.

24. Notice that lithium has two electrons in the first orbit and one in its outer orbit (valence shell). • If lithium loses one electron it will have the same electron arrangement as helium. • However the charge on the atom is no longer neutral. It has 3 protons (positive charges) and 2 electrons (negative charges) giving an overall charge of 1+ P = 3 N = 4

25. An ion is a charged atom in which the number of electrons is different from the number of protons • The term ionic charge is used to describe the overall charge an ion has. • Example Li which has lost an electron is written as Li+.

26. Create a Bohr – Rutherford diagram for calcium. • How many electrons would calcium have to lose in order to become stable like a noble gas? • What would the ionic charge be on Ca?

27. Metals tend to lose their electrons and nonmetals tend to gain electrons.

28. YOU DO IT • What would most likely to happen to the electrons in fluorine if it was to become stable like a noble gas? What would be the ionic charge on fluorine?

29. Fluorine would gain one electrons (nonmetal). • Fluorine would have a net ionic charge of 1- and the atom would be represented as F- (fluoride ion) • The ending “ide” is added to nonmetal ions

30. YOU DO IT • What would most likely to happen to the electrons in sulfur if it was to become stable like a noble gas? What would be the ionic charge on sulfur?

31. Sulfur would gain two electrons • S2- (sulfide ion)

32. Homework • 1. Create a Bohr – Rutherford diagram for: • Boron • Chlorine • Nitrogen • Beryllium • 2a. Create a Bohr – Rutherford diagram for the stable ion formed by each of the above atoms • b. State the ionic charge on each of the ions, and write the name of the atom in ionic form.

33. How Elements Form Compounds • Compounds are made by combining elements together. • Ionic Compound – is formed when positive and negative ions join together. The ions combine in a fashion that generally leaves the ionic compound with a neutral charge. • Molecular Compound – are formed when nonmetals combine with other nonmetals.

34. Chemical Formula • A chemical formula is a combination of symbols that represent a compound • Example: magnesium chloride (MgCl2) describes a compound with one magnesium ion to two chloride ions

35. Homework • Answer questions 2 and 3 on page 189. For part b construct a Bohr – Rutherford Diagram instead of a Bohr diagram.

36. Ionic Compounds • Metals and nonmetals combine to form compounds by sharing electrons (metals lose electrons to form positive ions and nonmetals gain electrons to form negative ions) • Example: Aluminum chloride (AlCl3) Al = 3+ and Cl = 1- (3+) + 3(1-) = 0

37. Writing Formulas For Ionic Compounds • What is the formula for the ionic compound formed by calcium and iodine? • Steps: • Write the symbols, with the metal first. Ca I • Write the ionic charge above each symbol to indicate the stable ion that each element forms. 2+ 1- Ca I • Determine how many ions of each type you need so that the total ionic charge is zero. One Ca2+ ion will balance the charge of two I- ions. • Write the formula using subscripts to indicate the number of ions of each type CaI2

38. YOU TRY IT • What is the formula for the ionic compound formed by aluminum and sulfur?

39. ANSWER • Al2S3

40. YOU TRY IT 2 • What is the formula for the ionic compound formed by aluminum and sulfur?

41. ANSWER • Al2S3

42. YOU TRY IT 3 • What is the formula for the ionic compound formed by nickel and oxygen?

43. ANSWER • NiO

44. Naming Ionic Compounds • The name of the metal is first followed by the name of the nonmetal. However, the ending of the name for the nonmetal changes to “ide” • Example: Calcium and Iodine become Calcium iodide • Example: Aluminum and Sulfur become Aluminum sulfide

45. Sc 3+ Ti 4+ 3+ V 5+ 4+ Cr 3+ 2+ Mn 2+ 4+ Fe 3+ 2+ Co 2+ 3+ Ni 2+ 3+ Cu 2+ + Zn 2+ Pd 2+ 4+ Ag + Cd 2+ Sn 4+ 2+ Sb 3+ 5+ Pt 4+ 2+ Au 3+ + Hg 2+ + Tl + 3+ Pb 2+ 4+ Bi 3+ 5+ Po 2+ 4+ Transition Metal Charges

46. Ce 3+ Pr 3+ Nd 3+ Pm 3+ Sm 3+ Eu 3+ 2+ Gd 3+ Tb 3+ Dy 3+ Ho 3+ Er 3+ Tm 3+ Yb 3+ Lu 3+ Ac 3+ Th 4+ Pa 5+ U 6+ 4+ Np 5+ Pu 4+ 6+ Am 3+ Cm 3+ Bk 3+ Cf 3+

47. Examples • Name or write the chemical formula for the following molecules. • 1) Tin (IV) oxide • 2)PbO • Platinum (II) sulfide • Ag3N • Answers • 1) SnO2 • Lead (II) oxide • PtS • Silver (I) nitride

48. Homework • Read page 195 • Questions 1, 3 – 9 on page 195

49. Polyatomic Compounds • Are formed when metals combine with polyatomic ions • Polyatomic Ions – groups of atoms that tend to stay together and carry an overall ionic charge

50. Writing Formulas for Polyatomic Compounds • Rules • Write the symbols of the metal and of the polyatomic group. • Write the ionic charges • Choose the number of ions to balance the charge. • Write the formula using subscripts