Writing Chemical Formulas and Naming Chemical Compounds

1. Writing Chemical Formulas and Naming Chemical Compounds Part 1: Writing Chemical Compounds

2. Chemical Formula • short hand method of indicating a ratio of atoms in a compound. • identifies the atoms in a compound • the less electronegative element is listed first in the formula • for covalent compounds, the formula tells you how many atoms of each type in the molecule.

3. What does this look like?

4. Valence or Oxidation Numbers • describes how many electrons from an atom are used in bond formation • if its an ionic bond • indicates how many electrons are donated ( + ) • indicates how many electrons are received ( - ) • if it’s a covalent bond it indicates how many electrons are contributed to the bond as if the electrons were completely removed or gained.

5. Oxidation Numbers • An oxidation number is the charge an atom would have if the electron pair that is shared between two atoms belonged entirely to the more electronegative atom.

6. Rules for Oxidation Numbers The following rules will help to assign oxidation numbers. The rules are listed in Priority Sequence. • free atoms (0) Ex. Al(s) • atoms bound to eachother. (0) Ex. Cl2, H2 ..... • monoatomic ions have the same oxidation number as the charge number Ex. Cl1- (-1), Mg2+ (+2), ....... • F is always (-1) • O is almost always (-2) -except in peroxides • H is almost always (+1) -except for metallic hydrides • Assign oxidation numbers to the most electronegative atom first

7. Rules Continued  Group I (+1) Group II (+2) Al (+3) Group VII (-1) Group VI (-2) When forming binary compounds withmetals. Group V (-3) nonmetals only In general, if two atoms form an ionic bond, the valence tells you the charges on the ions that are formed. If a covalent bond is formed, the valence tells you how many electrons the atoms contribute to the covalent bond.

9. Writing Chemical Formulas Using Valences Use the Zero Sum Rule (for neutral compounds only) • algebraic sum of oxidation numbers is zero • for charged polyatomic ions, the sum of oxidation numbers equals the charge of the ion. • Example: KF • Each potassium ion has a charge of +1 and each fluorine ion has a charge of -1. Because there is one of each ion in the formula, the sum of the valences is zero.

10. Polyatomic Ions • ions made of more than two atoms • are charged molecules • very strong covalent bonds keep these ions together and react as a single inseparable ion.

11. Need to Memorize... Ammonium NH4+ Acetate C2H3O2- Hydroxide OH- Nitrate NO3- Carbonate CO32- Chlorate ClO3- Sulfate SO42- Phosphate PO43- Cyanide CN-

12. Writing Chemical Formulas

13. Writing Chemical Formulas

14. Writing Chemical Formulas

15. Writing Chemical Formulas and Naming Chemical Compounds Part 2: Naming Chemical Compounds

16. Chemical Nomenclature • Each chemical compound has been given a name. Some have a trivial or common name, H2O is known as water. • These names have been used for centuries. However all chemicals have been given a name using Systematic Nomenclature. • uniquely describes it s compound • the name is derived from its chemical formula • from a name a formula can be determined • ex. NaCl common name: salt; (sodium chloride)H2O common name: water (hydrogen oxide)

17. Positive Monoatomic Ions • the name of the positive monoatomic ion is the same as the element name • Ex. Li+ -lithium  Stock System: if the metal has more than one oxidation number, a Roman Numeral is used: Sn2+ -Tin(II) Sn4+ Tin(IV) Old Method: the old name along with the suffix -ous or -ic is used to indicate the oxidation number. Sn2+ -stannous Sn4+ -stannic Memory Aid: -ous -indicates the lower oxidation number -ic -indicates the higher oxidation number

18. Negative Monoatomic Ions • add -ide to the end of the element name. • Ex. chlorine (Cl) becomes chloride (Cl- )

19. Binary Compounds Containing a Metal and a Nonmetal • consist of only 2 different elements • the element with the more positive oxidation number is written first. • Ex. NaCl and not ClNa

20. Binary Compounds Containing Hydrogen and Another Element • compounds of hydrogen and nonmetals from group VI and VII are called: • hydrogen _______ide • ex. HBr is call hydrogen bromide • compounds of hydrogen and a metal, the metal is written first and are called: • ________________ hydride • ex. NaH is called sodium hydride

21. Binary Compounds Containing 2 Nonmetals • this method uses prefixes to indicate the number of nonmetal elements present __________name _______name ex. As2S3 is called diarsenic trisulfide NO is called nitrogen monoxide N2O5 is called dinitrogen pentoxide

22. Naming Chemical Compounds Containing More than Two Elements Polyatomic ion + nonmetal ion • name the positive portion first then the negative ion • negative monoatonic ions end in –ide • ex. NH4Cl is called ammonium chloride Cation + polyatomic ion • oxygen containing polyatomic ions end in either –ate or –ite. • an –ite polyatomic ion contains one less oxygen than does an-ate polyatomic ion. (Memory aide: memorize the –ate form of ions) • ex. SO42- is sulfate and . SO32- is sulfite

23. Group VII Polyatomic Ions have Additional Forms • ClO-hypochlorite • ClO2- chlorite • ClO3- chlorate • ClO4-perchlorate

24. Nomenclature of Acids • Acid molecules contain hydrogen atoms that are easily removed when dissolved in water. These ionizable hydrogen atoms appear at the beginning of the chemical formula. Binary Acids • hydrogen and a nonmetal element from group VI or VII is called: hydrogen ______ide. • when this compound is dissolved in water, it becomes an acid called: hydro________ic acid ex. HCl (g) + H2O (l) → HCl (aq) hydrogen chloride hydrochloric acid

25. Oxyacids • hydrogen and a negative polyatomic ion is called: hydrogen polyatomic ion ex. H2SO4 (g) hydrogen sulfate H2SO3 (g) hydrogen sulfite • when these compounds are dissolved in water are called: ex. H2SO4 (aq) sulfuric acid H2SO3 (aq) sulfurous acid Note: In acid form –ate polyatomic ions become -ic acid -ite polyatomic ions become -ous acid

26. Intermolecular Forces

27. Proof for existence of intermolecular forces

28. Types of Intermolecular forces

29. Examples to explain Van der waals’ forces

30. Examples to explain dipole-dipole forces

31. Hydrogen bonding • Hδ+— Fδ- ---------------Hδ+— Fδ- • hydrogen bond • Considerably stronger than other intermolecular forces. • Affects the physical properties of the compounds in which it exists.

32. Examples of H-Bonding

33. Chemical Reactions and Collision Theory

34. Collision Theory A reaction will only be successful if it has the correct orientation and energy. Baseball Model: Baseball bat is Reactant A, baseball is Reactant B. The reaction is successful if the batter hits a homerun!

35. The scenarios… • Pitcher throws a fast ball, batter swings and misses. • Pitcher throws off-speed and the batter just makes contact with the ball. It’s a foul. • Pitcher throws a curve ball, batter swings and just hits the ball sending it to right field. • Pitcher throws a fast ball, batter swings and the ball goes flying high in the air. HOMERUN! When did the reaction occur?

36. Try with molecules Reaction: H2 + I2 2 HI +  H2 I2 HI How would you orient the molecules to get an effective collision capable of producing two hydrogen iodide molecules?

37. Types of Reactions • There are five types of chemical reactions we will talk about: • Synthesis reactions • Decomposition reactions • Single displacement reactions • Double Displacement reactions • Combustion reactions • You need to be able to identify the type of reaction and predict the product(s)

38. Steps to Writing Reactions • Some steps for doing reactions • Identify the type of reaction • Predict the product(s) using the type of reaction as a model • Balance it Don’t forget about the diatomic elements! (BrINClHOF) For example, Oxygen is O2 as anelement. In a compound, it can’t be a diatomic element because it’s not an element anymore, it’s a compound!

39. 1. Synthesis reactions • Synthesis reactions occur when two substances (generallyelements) combine and form a compound. (Sometimes these are called combination or addition reactions.) reactant + reactant  1 product • Basically: A + B  AB • Example: 2H2 + O2  2H2O • Example: C+ O2  CO2

40. Practice • Predict the products. Write and balance the following synthesis reaction equations. • Sodium metal reacts with chlorine gas 2 Na(s) + Cl2(g)  2 NaCl (s) • Solid Magnesium reacts with fluorine gas Mg(s) + F2(g)  MgF2 (s) • Aluminum metal reacts with fluorine gas 2 Al(s) + 3 F2(g)  2 AlF3 (s)

41. 2. Decomposition Reactions • Decomposition reactions occur when a compound breaks up into the elements or in a few to simpler compounds • 1 Reactant  Product + Product • In general: AB  A + B • Example: 2 H2O  2H2 + O2 • Example: 2 HgO  2Hg + O2

42. Decomposition Exceptions • Carbonates and chlorates are special case decomposition reactions that do not go to the elements. • Carbonates (CO32-) decompose to carbon dioxide and a metal oxide • Example: CaCO3  CO2 + CaO • Chlorates (ClO3-) decompose to oxygen gas and a metal chloride • Example: 2 Al(ClO3)3  2 AlCl3 + 9 O2 • There are other special cases, but we will not explore those in Chemistry 11

43. Practice • Predict the products. Then, write and balance the following decomposition reaction equations: • Solid Lead (IV) oxide decomposes PbO2(s)  Pb (s) + O2(g) • Aluminum nitride decomposes 2 AlN(s)  2 Al (s) + N2(g)

44. Practice Identify the type of reaction for each of the following synthesis or decomposition reactions, and write the balanced equation: N2(g) + O2(g) BaCO3(s)  Co(s)+ S(s)  NH3(g) + H2CO3(aq)  NI3(s)  Nitrogen monoxide (make Co be +3)

45. 3. Single Displacement Reactions • Single Displacement Reactions occur when one element replaces another in a compound. • A metal can replace a metal (+) OR a nonmetal can replace a nonmetal (-). • element + compound product + product A + BC  AC + B (if A is a metal)OR A + BC  BA + C (if A is a nonmetal) (remember the cation always goes first!) When H2O splits into ions, it splits into H+ and OH- (not H+ and O-2 !!)

46. Single Replacement Reactions • Write and balance the following single replacement reaction equation: • Zinc metal reacts with aqueous hydrochloric acid Zn(s) + HCl(aq) ZnCl2 + H2(g) Note: Zinc replaces the hydrogen ion in the reaction 2

47. Single Replacement Reactions • Sodium chloride solid reacts with fluorine gas NaCl(s) + F2(g)  NaF(s) + Cl2(g) Note that fluorine replaces chlorine in the compound • Aluminum metal reacts with aqueous copper (II) nitrate Al(s)+ Cu(NO3)2(aq)

48. 4. Double Replacement Reactions • Double Replacement Reactions occur when a metal replaces a metal in a compound and a nonmetal replaces a nonmetal in a compound • Compound + compound  product + product • AB + CD  AD + CB

49. Double Replacement Reactions • Think about it like “foil”ing in algebra, first and last ions go together + inside ions go together • Example: AgNO3(aq) + NaCl(s)  AgCl(s) + NaNO3(aq) • Another example: K2SO4(aq) + Ba(NO3)2(aq)  KNO3(aq) + BaSO4(s)

50. Practice • Predict the products. Balance the equation • HCl(aq) + AgNO3(aq)  • CaCl2(aq) + Na3PO4(aq)  • Pb(NO3)2(aq) + BaCl2(aq)  • FeCl3(aq) + NaOH(aq)  • H2SO4(aq) + NaOH(aq)  • KOH(aq) + CuSO4(aq) 