Structure of Matter

1. Structure of Matter Chapter 5

2. Students Will Be Able To: • Describe how a compound differs from its component elements. • Explain what a chemical formula represents. • State a reason why chemical bonding occurs. • Determine an atom’s number of valence electrons. • Identify the importance of the octet rule.

3. Chemical Compound • Most matter is in this form • A compound is a substance formed by the chemical combination of two or more elements Examples: • NaCl 1 Sodium + 1 Chloride = Salt • H2O 2 Hydrogen + 1 Oxygen = Water • CH4 1 Carbon + 4 Hydrogen = Methane

4. Chemical Formula • Is a representation of a compound • It shows the types and numbers of atoms or ions making up the simplest unit of the compound • Element symbol tells the type of element • Subscript tells you how many atoms • Examples: • NaCl Salt • H2O Water • CH4 Methane

5. Valence Electrons • Are the outermost electrons in an orbital • Can have between 1 and 8 • Groups have the same number of valence electrons • Group 1 has 1 valence electron • Group 2 has 2 valence electrons • Group 13 has 3 valence electrons • Group 14 has 4 valence electrons • Group 15 has 5 valence electrons • Group 16 has 6 valence electrons • Group 17 has 7 valence electrons • Group 18 has 8 valence electrons • Groups 3-12 varies

6. Valence Electrons • Atoms with partially filled energy levels can lose, gain, or share electrons to obtain a stable outer energy level • Are able to do this by combining with other atoms that also have unfilled outer energy levels • When atoms gain, lose, or share electrons an attraction forms between the atoms and pulls them together to form a chemical compound • Attraction is referred to as a chemical bond

8. Octet Rule • Is a tendency to have either empty outer energy levels or full outer energy levelsof eight electrons

10. Electron Dot Symbols • Are dots placed around atomic symbol to indicate the number of valence electrons • Group 1A atoms (Na, K, etc) have a single dot • Group 2A atoms (Mg, Ca, etc) have 2 dots • Helium is an exception

11. Students Will Be Able To: • Define an ion. • Differentiate between cations and anions. • Describe the process of ionic bonding. • Describe the properties of ions.

12. Ions • Atoms are electrically neutral due to the number of protons equaling the number of electrons • By gaining or losing electrons an atom can be converted into a charged particle called an ion

13. Types of Ions • Cations • Occurs when there is a loss of one or more electrons • Has a positive charge • Anions • Occurs when there is a gain of one or more electrons • Has a negative charge

14. A Way to Remember • CATIONS have more protons than electrons • ANIONS have more electrons than protons • Remember…… • CATS have PAWS • CATions are PAWSitive

15. Cations • Its symbol is written by adding a positive charge as a superscript to the symbol for the element • Example: • Na (sodium) loses an electron to make the sodium cations (Na+)

16. Anions • Its symbol is written by adding a negative charge as a superscript t the symbol for the element • Example: • Cl (chlorine) gains an electron to make the chloride anion (Cl-)

19. Ionic Bonding • Formed when one or more electrons are transferred from one atom to another • Lose electron- positive charge (cation) • Are often metals • Gain electron- negative charge (anion) • Are often nonmetals • Less energy is required to give up one valence electron than to give up seven valence electrons

21. Ionic Bonds • Opposite electrical charges are attracted to one another • Example: • When sodium combines with chlorine, sodium transfers electron to chlorine forming Na+ and Cl- ions • The oppositely charged Na+ and Cl- ions are held together by an ionic bond, making an ionic compound

22. Ionic Compound Structure • They are usually crystalline solids • They vary in size and charge • They have high melting and boiling points • Due to strong attraction between ions

24. Dissolving Capabilities • An ionic compound will dissolve in water if the attraction between water and the ions is greater than the attraction between the ions in an ionic compound • An ionic compound will not dissolve in water if the attraction between water and the ions is not greater than the attraction between the ions in an ionic compound

26. Students Will Be Able To: • Identify binary compounds. • Name cations, anions, and ionic compounds. • Write chemical formulas for ionic compounds, such that an overall neutral charge is maintained. • Explain why polyatomic ions and their salts are named.

27. Binary Ionic Compounds • Consists of two elements joined together • One element is a cation and the other is an anion

28. Naming Cations • Main group metal cations (Groups 1A, 2A, and 3A) are named by identifying the metal, followed by the word “ion” • Examples: • K+ = Potassium ion • Mg2+ = Magnesium ion • Al3+ = Aluminum ion

29. Naming Cations • Transition metals can often form more than one type of cation • Two methods used • Old name • Adding a different ending the metal’s name • New name • Write out metal’s name and then put in parentheses the charge number (written in Roman numeral)

30. Naming Cations • Chromium • Can have 2 charges • Cr2+ • Chromous ion • Chromium (II) ion • Cr3+ • Chromic ion • Chromium (III) ion

31. Naming Anions • Main group nonmetal anions (Groups VA, VIA, and VIIA) are named by identifying the nonmetal and changing the ending to “ide” followed by the word “ion” • Examples: • Cl- = chloride ion • O2- = oxide ion • P3- = phosphide ion

32. Monoatomic Ions • Mono = one • Atomic = atom • Monoatomic = one atom • Consists of only one element

33. Naming Ionic Compounds • Ionic compounds are named by citing first the cation and then the anion with a space between the words • Examples: • NaBr = sodium bromide • MgSO4 = magnesium sulfate • SnCl2 = Tin (II) chloride • SnCl4 = Tin (IV) chloride • Al2O3 = Aluminum oxide • **Subscript represents the number of atoms**

34. Ionic Formulas • Formulas of an ionic compound show the lowest possible ration of atoms in the compound

35. Writing Ionic Formulas • Ionic compounds have a balance of positive and negative charges • Sodium has a positive 1 charge and chlorine has a negative 1 charge • Written as NaCl • Sometimes the charges are not balanced • Therefore, the elements must be combined in a way to allow the charges to balance • Calcium has a positive 2 charge and chlorine has a negative 1 charge • It will take 2 chlorine atoms to produce a negative 2 charge so the charges can balance • Written as CaCl2

36. Polyatomic Ions • Poly = many • Atomic = atoms • Polyatomic = many atoms • Is a charged group of 2 or more atoms bonded together • Parentheses group the atoms of a polyatomic ion • Example • The chemical formula for ammonium sulfate is written as (NH4)2SO4, not N2H8SO4

37. Polyatomic Ions • Ammonium ion, NH4+, is composed of 1 nitrogen and 4 hydrogen atoms • Together is has 10 electrons due to the + charge • Nitrogen has 7 electrons and 4 hydrogen atoms have 4 4 electrons

38. Polyatomic Ions • Some polyatomic anion names relate to their oxygen content • An -ate ending is used to name an ion with more oxygen • Examples: sulfate (SO42–), nitrate (NO3–), chlorate (ClO3–) • An -ite ending is used to name an ion with less oxygen • Examples: sulfite (SO32–), nitrite (NO2–), chlorite (ClO2–)

39. Polyatomic Ions • The presence of hydrogen is often indicated by an ion’s name starting with hydrogen • The prefixes mono- anddi-arealso used • HPO42- is monohydrogen phosphate • H2PO4- is dihydrogen phosphate • The prefix thio- means “replace an oxygen with a sulfur • K2S2O3 is potassium thiosulfate

40. Naming Polyatomic Ions • Follow these steps when naming an ionic compound that contains one or more polyatomic ions: • Name the cation • Recall that a cation is simply the name of the element • Name the anion • Recall that salts are electrically neutral • Name the salt • Recall that the name of a salt is just the names of the cation and anion

41. Students Will Be Able To: • Describe covalent bonds. • Identify the particles produced by covalent bonding. • Name covalent compounds • Write formulas for covalent compounds.

42. Covalent Bonds • Atoms share valence electrons to become stable • Known as molecular compounds • Valence electrons are considered as part of the outer energy level for both atoms • No gaining or losing takes place • Occurs between nonmetals ONLY

43. Covalent Bonds • Formed when electrons are shared from one atom to another; molecule results • 2 electrons shared- single bond • 4 electrons shared- double bond • 6 electrons shared- triple bond

46. Naming Covalent Bonds • The first element is always named first • The second element ends in –ide • Prefixes are used to indicate the number of atoms each type of element in the compound

47. Naming Covalent Bonds • The system of prefixes is used to show the number of atoms of each element in the molecule

48. Naming Covalent Bonds • Prefixes can be used to show the number of each type of atom in diphosphorus pentasulfide

49. Metallic Bonding • A metallic bond is a bond formed by the attraction between positively charged metal ions and the electrons around them • Electrons move freely between metal atoms • This model explains why metals • Conduct electricity • Conduct heat • Are flexible

50. Students Will Be Able To: • Distinguish between a nonpolar and polar covalent bond. • Differentiate among single, double, and triple covalent bonds. • Draw Lewis structures for single covalent compounds and polyatomic ions.