The Structure of Matter

1. The Structure of Matter Chapter 4

2. Chapter Preview • 4.1 Compounds and Molecules • What Are Compounds? • Models of Compounds • How Does Structure Affect Properties? • 4.2 Ionic and Covalent Bonding: • What Holds Bonded Atoms Together? • Ionic Bonds • Metallic Bonds • Covalent Bonds • Polyatomic Ions • 4.3 Compound Names and Formulas • Naming Ionic Compounds • Writing Formulas for Ionic Compounds • Naming Covalent Compounds • Chemical Formulas for Covalent Compounds • 4.4 Organic and Biochemical Compounds • Organic Compounds • Polymers • Biochemical Compounds

3. Section 4.1 Compounds and Molecules • Objectives: • Distinguish between compounds and mixtures. • Relate the chemical formula of a compound to the relative numbers of atoms or ions present in the compound. • Use models to visualize a compound’s chemical structure. • Describe how the chemical structure of a compound affects its properties.

4. What are Compounds? • A compound is made up of two or more elements. • Ex : Table Salt made up of Na and CL (NaCl) Chemical bonds distinguish compounds from mixtures. Chemical bonds: the attractive force that holds atoms or ions together.

5. Compounds • Chemical bonds distinguish compounds from mixtures. • Each substance in a mixture keeps its own properties. • A compound always has the same chemical formula. Chemical formula shows the types and numbers of atoms or ions making up the simples unit of the compound. H20

6. Chemical structure shows the bonding within a compound • Chemical structure: the arrangement of bonded atoms or ions with a substance.

7. Two terms are used to specify the positions of atoms relative to one another in a compound. • Bond length: the arrangement of bonded atoms or ions within a substance. • Bond angle: the angle formed by two bonds to the same atom.

8. Models can represent physical events: • Scientific model: is a representation of an object or event that can be studied to understand the real object or event.

9. Three types of Models use to visualize Compounds • Some models give you an idea of bond lengths and angles. • Types of Models: • Structural formulas: shows the structure of compounds.

10. Space-filling modes show the space occupied by atoms • The problem with this model is that it is harder to “see” bond lengths and angles.

11. Ball and stick model • This model makes it easy to see the bonds and the angles they form in a compound.

12. Structure Affects Properties • Compounds with network structures are strong solids. • Some networks are made of bonded ions. • These strong attractions cause high melting and boiling points. • Good conductors of electricity. Table Salt

13. Some compounds are made of molecules • Salt and sugar are both white solids you can eat, but their structures are very different.

14. Section 4.1 Summary • Atoms or ions in compounds are joined by chemical bonds. • A compound’s chemical formula shows which atoms or ions it is made of. • A model represents a compound’s structure visually. • Substances with network structures are usually strong solids with high melting and boiling points.

15. Section 4.2Ionic and Covalent Bonding

16. Section 4.2 Objectives • Explain why atoms sometimes join to form bonds. • Explain why some atoms transfer their valence electrons to form ionic bonds, while other atoms share valence electrons to form covalent bonds. • Differentiate between ionic, covalent, and metallic bonds. • Compare the properties of substances with different types of bonds.

17. What Holds bonded Atoms Together? • Atoms bond with other atoms when their valence electrons interact. • Atoms with full outermost energy levels are non- reactive than atoms with only partially filled outermost energy levels.

18. Stable atom vs non stable atom: Full Valence : Non Reactive Partially Filled : More Reactive

19. Chemical Bonding In Action • Atoms want to join to form bonds so that they can fill the outermost energy levels.

20. Bonds: can bend and strtch without breaking • Bonds behave more like flexible springs than like sticks. • Three types of Bonds : • Ionic Bonds • Metallic Bonds • Covalent Bonds

21. Ionic Bonds • Ionic Bond – a bond formed by the attraction between oppositely charged ions. 1. Atoms of metal elements form positive ions (cat- ion). 2. Atoms of non metal elements form negative ions (anion).

22. Ionic Bond Example 3. Some atoms do not share electrons: • They transfer electrons where one atom gains and one loses. This process is known as ionization. • Ions have filled valence energy level • They result in (+) and (–) ions. • Opposite charged ions attract and form ionic bonds ex:NaCl

23. Metallic Bonds • Metallic Bond - bond formed from attraction between positive charged metal ions and the electrons around them. • Metal ions are surrounded by what are known as delocalized electrons. Delocalized electrons are valenceelectrons from metal atoms. They move from one place to another and are not associated with a particular metal atom.

24. Covalent Bonds • Covalent Bond – bond formed when atoms share 1 or more pairs of valence electrons. • They are often formed between nonmetal atoms and can be solid, liquid, or gas. • They have a low melting point mostly below 300°C except silicon dioxide and any other compounds with network structures.

25. Atoms do not always share electrons equally: • An unequal sharing of valence electrons forms a polar covalent bond. • An equal sharing of valence electrons forms a non-polar covalent bond.

26. Ionic VS. Covalent Bonds

27. Atoms do not always share electrons equally. • Polyatomic ion: an ion made of two or more atoms that are covalently bonded and that act like a single ion. • There are many polyatomic ions:

28. Section 4.2 Summary Report • Atoms bond when their valence electrons interact. • Cations and anions attract each other to form ionic bonds. • When ionic compounds are melted or dissolved in water, moving ions can conduct electricity. • Atoms in metals are joined by metallic bonds. • Metals conduct electricity because electrons can move from atom to atom. • Covalent bonds form when atoms share electron pairs. Electrons may be shared equally or unequally. • Polyatomic ions are covalently bonded atoms that have either lost or gained electrons. Their behavior resembles that of simple ions.

29. Section 4.3 Compound Names and Formulas

30. Section 4.3 Objectives: • Name simple ionic and covalent compounds. • Predicts the charge of a transition metal cation in an ionic compounds. • Write chemical formulas for simple ionic compounds. • Distinguish a covalent compound’s empirical formula from its molecular formula.

31. Naming Ionic Compounds • Remember: Ionic Bond = the attraction between a positive and a negative ion (left side attracted to the right side of the periodic table). • Use Table 4-4 on page 123 and Table 4-5 on page 124 to show the charges for common cations and anions. • Remember: Anions (right side) change their name. Ex. Fluorine to Fluoride ion.

32. Naming Ionic Compounds • So….if you are given CsBr, by looking at the table you will see that Cs is Cesium and Br is Bromide therefore the compound’s name is… Cesium Bromide • Ok now CaCl2…. Calcium Chloride Li3N Lithium Nitride

33. Naming Ionic Compounds • Now that you can see that all you have to do is look at the table to name them, how about Metallic compounds? • Look at Table 4-6 page 124 • Notice that Roman Numerals are used. • Roman Numerals represent the charge of the transition metal.

34. Naming Metal Cations • So if you are given Fe2O3 • What is Fe? • Iron • What is O? • Oxygen that changes to Oxide • So put them together and you have Iron Oxide, but this is not it! • What happened to the numbers?

35. Naming Metal Cations • Notice that is was Fe2O3 If you see that the charge for Iron is 3+ Then the name of the compound will now be Iron (III) Oxide. This only applies to Metal Cations in using Roman Numerals Charge of Iron Charge of Oxygen

36. Writing Formulas for Ionic Compounds • By using the tables you can determine charges of ions. • How do we write chemical formulas? • Lets take aluminum fluoride • Aluminum from the table is Al3+ • Fluoride from the table is F1- • So now you see Al3+ F1- • But this is not the formula

37. Writing formulas for Ionic Compounds • Al3+ F1- What is the least common multiple of 3 and 1? • The least common multiple of 3 and 1 is 3 • Your objective is to make the charges neutral • So take Al3+ F1- • Your formula is now AlF3 3 x -1 = -3

38. Writing formulas for Ionic Compounds • Now try lithium oxide • Lithium Li+1 Oxide O2- • Together … Li+1 O2- • The least common multiple is 2 • Li+1 O2- • Formula is now Li2O 2 x 1 = 2

39. Writing Formulas for Ionic Compounds • How about titanium (III) nitride? • Here we have a metal cation so the charge of titanium will be the roman numeral given. • So…Titanium is Ti3+ and nitride is N3- • Together…Ti3+N3- • Notice that 3 is already the least common multiple so the formula will be TiN

40. Naming Covalent Comounds • Covalent compounds, like SiO2 and CO2, are named using different rules than those used to name ionic compounds. • Numerical prefixes are used to name covalent compounds of two elements.

41. Naming Covalent Compounds • Remember that a covalent compound is a bond when atoms share one or more pairs of electrons. • Use Table 4-7 on page 126 • This table will be used to determine the prefix of the element when we name it. • Given N2O4 • How many Nitrogen atoms are there? • Correct…2 so look at the table and what is the prefix for 2?

42. Naming Covalent Compounds • The prefix for 2 is Di • Now how many oxygen are there? • Correct… 4 so what is the prefix for 4? • The prefix for 4 is tetra • So given N2O4 using the prefixes we have • Dinitrogen Tetroxide

43. Naming Covalent Compounds • How about BF3? • Boron Trifluoride • Now As2O5 • Diarsenic pentoxide

44. Section 4.3 Review • To name an ionic compound, first name the cation and then the anion. • If an element can form cations with different charges, the cation name must include the ion’s charge. The charge is written as a Roman Numeral in parentheses. • Prefixes are used to name covalent compounds made of two different elements. • An empirical formula tells the relative numbers of atoms of each element in a compound. • A molecular formula tells the actual numbers of atoms in one molecule of a compounds. • Covalent compounds have both empirical and molecular formulas.

45. Section 4.4Organic and Biochemical Compounds • Objectives: • Describe how carbon atoms bond covalently to form organic compounds. • Identify the names and structures of groups of simple organic compounds and polymers. • Identify what the polymers essential for life are made of.

46. Section 4.4 Organic and Biochemical Compounds • In chemistry, the organic is used to describe certain compounds. • Organic compound: any covalently bonded compound that contains carbon. • Organic compounds contain carbon and, almost always, hydrogen. • Examples of Organic Compounds: • Acetylsalicylic acid, C9H8O4 • Sorbitol: C6H14O4 • Aspartame : C14H18N2O5

47. Carbon atoms form four covalent bonds in organic compounds. • When a compound is made of only carbon and hydrogen atoms it is called a hydrocarbon. • The simples hydrocarbon is Methane CH4 • Methane gas is formed when living matte, such as plants, decay, so it is often found in swamps and marshes. • Carbon atoms have four (4) valence electrons to use for bonding.

48. Types of Hydrocarbon • Alkanes are hydrocarbons that have only single covalent bonds. • For example; • Methane (CH4) : • Ehtane (C2H6) : • Propane: (C3H8) :

49. Arrangements of carbon atoms in alkanes can vary: • The carbon atoms in methane, ethane, propane and butane all lie up in a row this is called normal alkane.

50. Alkanes chains may be branched or unbranched, and they can even form rings.