1. Warm Up Net Ionic Equation
A solution of ammonia is added to a dilute solution of acetic acid

2. Unit 1: Chemistry Fundamentals
Zumdahl: Chapter 2

3. Atomic Theory Foundations
Law of Conservation of Mass – mass is neither created or destroyed during a chemical or physical change Law of Definite Proportions – a compound contains the same proportions by mass regardless of the size of the sample Example: NaCl – always 39.34% Na & 60.66% Cl

4. Atomic Theory Foundations
Law of Multiple Proportions – if two or more different compounds are composed of the same two elements, then the ratio of the masses of the second element combined with a certain mass of the first element is always a ratio of small whole numbers Example CO2 and CO

5. Dalton’s Atomic theory
All matter is composed of atoms
Atoms of an element have the same size, mass and properties; atoms of a different element have different sizes, masses and properties
Atoms cannot be divided, created or destroyed
Atoms of different elements combine in simple whole number ratios
Chemical reactions combine, separate, or rearrange atoms

6. Modern Atomic Theory Atoms can be divided
Atoms of the same element can have different masses
All else remains the same

7. Comparing Theories Dalton See notes Solid sphere Thompson
Plum pudding model
Electrons scattered thru positively charged cloud
Rutherford
Concept of the nucleus
Positively charged

8. Isotopes
II.A.2(c) – compare the characteristics of isotopes of the same element
IV.B.2(d) – calculate the weighted average atomic mass of an element from isotopic abundance, given the mass of each contributor

9. Def: atoms of the same element that have different masses
Isotopes
Def: atoms of the same element that have different masses

10. protium – 1 proton in nucleus deuterium – 1 proton; 1 neutron
Example: hydrogen
protium – 1 proton in nucleus
deuterium – 1 proton; 1 neutron
tritium – 1 proton; 2 neutrons
*Nuclide – general term for any isotope

11. Writing Isotopes
Example – Uranium 235

12. Average Mass Number
Def: the weighted average of the atomic masses of the naturally occurring isotopes of an element Like calculating a “weighted” grade (decimal % of each isotope x mass of that isotope)

13. Sample Calculation

14. The Periodic Table
IV.B.2(c) – use the periodic table to determine the atomic number, atomic mass, mass number, and number of protons, electrons and neutrons in isotopes of elements

15. Reading the Periodic Table
Atomic Number
# of protons in nucleus
Element Symbol
Element Name
Atomic Weight
Electron Configuration 3 Li
Lithium
6.941
[He]2s1

16. Calculating Protons, Electrons & Neutrons
IV.B.2(c) – use the periodic table to determine the atomic number, atomic mass, mass number, and number of protons, electrons and neutrons in isotopes of elements

17. Mass Number
Def: the number of protons and neutrons in the nucleus of an isotope mass # - atomic # = # of neutrons Example – oxygen Mass # (16) – atomic # (8) = # of neutrons (8)

18. Formulas
II.A.2(a) – use the IUPAC symbols of the most commonly referenced elements
III.A.1(a) – distinguish between chemical symbols, empirical formulas, molecular formulas and structural formulas
III.A.1(b) – interpret the information conveyed by chemical formulas for numbers of atoms of each element represented
III.A.(d) – provide the interconversion of molecular formulas, structural formulas, and names, including common binary and ternary acids

19. Chemical Symbols One or two letter symbol used to represent an element
H – hydrogen
Ca – Calcium
Sn – tin (stannum)
Cl – chlorine

20. Molecular Formulas Represent the elements in a compound
Subscripts tell how many atoms of that element are in a compound
Example – sulfuric acid
H2SO4
2 – hydrogen
1 – sulfur
4 – oxygen

21. Empirical Formulas
consists of the symbols for the elements combined in a compound, with subscripts showing the smallest whole-number ratio of the different atoms in the compound

22. Structural Formulas
ethane
ascorbic acid (vitamin C)
cholesterol

23. Counting by Weighing

24. Percent Composition
III.A.1(c) – calculate the percent composition of a substance given its formula or masses of each component element in a sample

25. % Composition Sample Problems

26. NOMENCLATURE Zumdahl – Ch 2
III.A.1(d) - provide the interconversion of molecular formulas, structural formulas, and names, including common binary and ternary acids
NOMENCLATURE

27. Oxidation Numbers
Pure elements in their standard states have an oxidation number of zero
Fluorine has an oxidation of -1 in all compounds
Oxygen has an oxidation number of -2 except when in a peroxide when its oxidation number is -1
Hydrogen has an oxidation number of +1 except when bonded to a metal
The algebraic sum of the oxidation numbers of all atoms in a neutral compound is zero
The algebraic sum of the oxidation numbers of all atoms in a polyatomic ion is equal to the charge of the ion
These rules apply to covalently bonded atoms

28. Oxidation Numbers Cont’d
Group 1 = +1
Group 2 = +2
Transition Metals = varies
Group 15 = varies (-3 usually)
Group 16 = varies (-2 usually)
Group 17 = varies (-1 usually)

29. General Rules for Naming Chemical Compounds
Cation (pronounced cat-ion)
is named first
has a positive charge/oxidation number
name does NOT change
Anion (pronounced an-ion)
is named second
has a negative charge/oxidation number
ending is “-ide”

30. Binary Ionic Compounds Type I
Metal is from Group 1 or 2
charges/oxidation numbers = 0
Examples:
NaCl – Na+1 Cl-1 = sodium chloride
KF – K+1 F-1 = potassium fluoride
CaBr2 – Ca+2 Br-1 = calcium bromide
Li2O – Li+1 O-2 = lithium oxide

31. Binary Ionic Compounds Type II
Metal is a transition metal
Many metals form more than one type of positive ion
Iron: Fe2+ or Fe3+
Copper: Cu1+ or Cu2+
Lead: Pb2+ or Pb4+
Tin: Sn2+ or Sn4+
Roman numeral indicates charge on the metal ion (NOT the number of atoms)
Look at 2nd element to determine charge on transition metal

32. Examples of Type II
EX: FeCl2 – iron(II) chloride FeCl3 – iron(III) chloride CuO – copper(II) oxide Cu2O – copper(I) oxide SnF4 – tin(IV) fluoride SnO2 – tin(IV) oxide SnO – tin(II) oxide **NOTE: sum of oxidation numbers = 0

33. Exceptions to the Rule
Aluminum – always Al3+ Cadmium – always Cd2+ Zinc – always Zn2+ Silver – always Ag+ No Roman numerals needed

34. Binary Covalent Compounds or Binary Molecular Compounds
Formed between two non-metals
first element name doesn’t change
Second element ends with “-ide”
Don’t use “mono-” on first name
Use prefixes to indicate number of atoms of element

35. Prefixes Mono– 1 Di– 2 Tri- 3 Tetra- 4 Penta- 5 Hexa- 6
Hepta- 7 Octa- 8 Nona- 9 Deca- 10 Undeca – 11 Dodeca - 12

36. Examples of Binary Covalent Compounds
CO – carbon monoxide CO2 – carbon dioxide CCl4 – carbon tetrachloride SiO2 – silicon dioxide SeBr2 – selenium dibromide BF3 – boron trifluoride P2O4 – diphosphorous tetroxide P4O10 – tetraphosphorous decoxide

37. Polyatomic Ions Group of atoms with a shared charge
III.A.1(c) – use the names, formulas, and charges of commonly referenced polyatomic ionsl

38. Polyatomic Rule #1 Common form ends with –ate
Example: nitrate, chlorate, sulfate

39. Polyatomic Rule #2
If oxygen decreases by one, O-1, changes ending to “-ite”
Example: nitrite, chlorite, sulfite

40. Polyatomic Rule #3 If oxygen decreases by two, O-2, add prefix “hypo-”
Keep ending “-ite”
Example: hypochlorite

41. Polyatomic Rule #4 If oxygen increases by one, O+1, add prefix “per-”
Keep ending “-ate”
Examples: perchlorate, permanganate

42. Polyatomic Ions Chlorate Bromate Iodate Nitrate Permanganate Carbonate
Silicate
Selenate
Phosphate
Arsenate
ClO3-
BrO3-
IO3-
NO3-
MnO4-
CO32-
SiO32-
SeO32-
PO43-
AsO43-

43. More Polyatomic Ions Acetate Hydroxide Bicarbonate Bisulfate Ammonium
Thiocyanate
Thiosulfate
Oxalate
Chromate
Dichromate
C2H3O2- OH- HCO3- HSO4- NH4+ SCN- S2O32- C2O42- CrO42- Cr2O72-

44. Binary Acids
III.A.1(d) - provide the interconversion of molecular formulas, structural formulas, and names, including common binary and ternary acids

45. Binary Acids Hydrogen takes place of metal ion
Use prefix hydro- in place of hydrogen
Use suffix -ic on the root
Examples –
HCl – hydrochloric acid
HF – hydroflouric acid
H2S – hydrosulfuric acid

46. Oxyacids
III.A.1(d) - provide the interconversion of molecular formulas, structural formulas, and names, including common binary and ternary acids

47. Oxyacids contain hydrogen, a non-metal, and oxygen
usually contain 3 or 4 oxygen atoms
add suffix -ic to the stem
Examples –
HBrO3 – bromic acid
H2CO3 – carbonic acid
HClO3 – chloric acid
HNO3 – nitric acid
H3PO4 – phosphoric acid
H2SO4 – sulfuric acid

48. Oxyacids Cont’d Common name ends with -ate
One extra oxygen add per- to stem name
Example –
HClO3 – chloric acid
HClO4 – perchloric acid
One less oxygen change -ic ending to -ous
HClO2 – chlorous acid

49. Oxyacids cont’d
Two fewer oxygens change -ic ending to -ous AND add prefix hypo-
Examples -
HClO3 – chloric acid
HClO – hypochlorous acid