Chapter 2 part 1

1. Chapter 2 part 1

2. Law of conservation of mass • Dalton • Mass is neither created or destroyed in chemical rxns.

3. Law of definite proportions • Dalton • In a given compound the relative number and kind of atoms is constant. • Example: Water = 2 H and 1 O

4. Dalton’s Atomic Theory • each element is composed of atoms • All atoms of a given element are identical • A chemical compound is the result of the combination of atoms of 2+ different elements • Atoms of an element are not changed into different types of atoms by chemical rxns (MOVIE)

5. Millikan • Discovered the charge and mass of an electron 1.6 X 10 -19 C (coulomb)

6. Radioactivity • Three types of radiation • Alpha α +2 charge • Gamma γ no charge • Beta β -1 charge

7. Rutherford • Most of the mass of an atom and all of its positive charge resides in a very small dense region (nucleus) • Most of the volume of the atom is empty space around the nucleus

9. Atomic Structure • Proton • Symbol p+ • Charge = 1.602 x 10-19 C (coulomb) • Mass = 1.0073 amu • Electron • Symbol = e- • Charge = -1.602 x 10-19 C • Mass = 5.486 x 10-4 amu • Neutron • Symbol = n0 • no charge • Mass = 1.0087 amu

10. Units = amu or g/mol

11. Complete Chemical Symbol A Z X A = Atomic mass ( p+ n) Z = Atomic number ( p) X = element Symbol

12. QuestionHow many neutrons, protons and electrons are in carbon? C 12 6 C = Carbon 12 = Atomic Mass ( protons + neutrons) 6 = atomic Number (protons)

13. Answer 12 = p + n 6 = protons 12 – 6 = 6 neutrons In an atom with out a charge the number of protons is equal to the number of neutrons Thus: 6 electrons

14. Question • Write the complete chemical symbol for the following elements. • Magnesium • Sodium • Tungsten

15. Determining Sub Atomic Particles 12 6 C C = Carbon 12 = Atomic Mass/weight ( protons + neutrons) 6 = atomic Number (protons)

16. Isotopes • Isotopes • Atoms of a given element that differ in number of neutrons and thus in mass.

17. Isotopes Cont. • When writing isotopes the atomic number ( # of p+) stays the same, but the Atomic Mass (p+ + n0) changes due to the addition and subtraction of neutrons.

18. Formulas • Chemical Indicates actual numbers and type of atoms in a molecule • H2O2 C2H4 • Empirical gives only relative number of atoms of each type • HO CH2 • Structural individual bonds are shown, indicated by lines

19. average atomic mass • Amu = Average Atomic Mass Unit • The average atomic mass (weight) of an element is equal to the sum of the products of each isotope’s mass (in amu) multiplied by it’s relative abundance.

20. EXAMPLE OF AVERAGE ATOMIC MASS PROBLEM • Naturally occurring chlorine is 75.53% Cl-35 which has an atomic mass of 34.969 amu, and 24.47% Cl-37, which has an atomic mass of 36.966 amu. • Calculate the average atomic mass of chlorine.

21. EXAMPLE OF AVERAGE ATOMIC MASS PROBLEM (CONT) • Average atomic Mass = [ (%/100) (Atomic Mass) ] • Average atomic mass = (0.7553) (34.969 amu) + (0.2447) (36.966 amu) • = 26.41 amu + 9.045 amu • = 35.46 amu • NOTE: The average atomic mass of an element is closest in value to the atomic mass of the most abundant isotope.

22. Halogens Alkali metals TRANSITION METALS Alkaline Earth Metal Metalloids Metals Non Metals

23. DIATOMIC SEVEN

24. Nomenclature Chapter 2 part 2 Check out these videos for more help or if you are absent Naming molecular compounds Writing formulas for molecular compounds Naming ionic compounds video Writing formulas for Ionic compounds

25. 2.7 Ions and ionic compounds • When negative electrons are removed or added to an atom the charge of that atom changes from its neutral state to a charged state ( + or - ) • Ion: charged particle

26. a cation is a particle that carries a positive electrical charge. The cation gets this positive charge from losing negatively charged electrons.

27. Anions are ions that carry a negative electrical charge. Anions get their negative charge by gaining one or more electrons

28. Trick Cl- Na+

29. Example Na = atomic number 11 atomic mass = 23 11 protons, 12 neutrons, 11 electrons Na- = 11 protons, 12 neutrons, 12 electrons We added an additional negative charge (e-) creating a negatively charged particle. 11 + -12 = -1 charge on Na

30. Reverse example Na = atomic number 11 atomic mass = 23 11 protons, 12 neutrons, 11 electrons Na+ = 11 protons, 12 neutrons, 10 electrons We subtracted a negative charge (e-) creating a positively charged particle. 11 + -10 = +1 charge on Na

31. Question • How many protons and electrons does Se2- ion posses? • How many protons and electrons does Cr3+ ion posses? • What kind of ions are these molecules and why?

32. Answer • Se2- : • anion • 34 protons (atomic number) • 36 electrons = -34 + -2 = -36 • Cr3+ : • Cation • 24 protons (atomic number) • 21 electrons = 24 – 3 = 21

33. Nomenclature (Naming) • As of 2007 there are 31,000,000 known compounds. • Your are options are A: memorize all 31,000,000 names B: Learn how to name them memorize about 50 things that will allow you to name all 31,000,000.

34. Polyatomic Ions • Atoms joined as a molecule, but they have a net positive or negative charge. • Example: NO3-, SO42- I will give you a list. You need to try to memorize them all… yes all.

35. Polyatomic Ion Rap

36. Cation and Anions to Memorize +/-3 +/-4 Write these on your binder periodic table

37. Putting the pieces together Na+ Cl- Mg2+ 2Cl- NaCl is an ionic compound MgCl2 is an ionic compound

38. Ionic Compounds • Contain both positively and negatively charged ions. • In general ionic compounds are made of metals and nonmetals.

39. Covalent compounds • 2 negatively charged elements • 2 non metals

42. Question • Write the ionic compound for: • Magnesium and Nitrogen • Magnesium and NO3-

43. Answer • Mg3N2 • Mg(NO3)2 • Count the total number of each type of atom in the molecules above.

44. Homework • Pg 71 • #’s 37,40,45,47,48

45. 2.8 Naming ionic compounds • PINK SHEET