1. Atoms and Molecules By Doba Jackson, Ph.D.

2. Outline of Concepts in this chapter Law of Constant Composition (sec. 2-4) –Mass Percentages Dalton’s Atomic Theory (sec. 2-5) –Postulates Atomic and Molecular Mass (sec. 2-7, 2-8) –Atomic Mass Unit (amu) Protons, Neutrons, and Electrons (sec 2-9)

3. Dalton’s Atomic Theory 1.Elements are made up of tiny particles called atoms. 2.Each element is characterized by the mass of its atoms. 3.Chemical combination of elements to make different chemical compounds called molecules. 4.Chemical reactions only rearrange the way that atoms are combined in molecules; the atoms themselves don’t change.

4. Law of Constant Composition 2Hg2HgO Chemical Formula O2O2 + Chemical Equation 1)All matter is made of a specific composition of elements from the periodic table. 2)The mass percentage of each element in a pure molecule is conserved.

5. Law of Constant Composition Chemical Formula 2Hg2HgOO2O2 + Chemical Equation HgO Mass percentage of mercury Mass percentage of oxygen

6. Suppose we analyze 2.83 g of Lead Sulfide, PbS (consists of lead and sulfur), and find that it consist of 2.45 g of lead (Pb) and.380 g of sulfur (S). Calculate the mass percentages of lead (Pb) and sulfur (S) in the sample. Example % Lead (Pb) % Sulfur (S)

7. Definitions that lead to Dalton’s Atomic Theory Law of Constant Composition: The relative amount of each element in a particular compound is always the same regardless of the source of the compound, or how the compound was prepared. Law of Multiple Proportions: This law says that when a given element, (X), combines with another element, (Y), then the ratio of masses of X and Y are small whole numbers. Law of Conservation of Mass: For any closed system, the mass of the system will remain constant over time

8. Law of Multiple Proportions Law of Multiple Proportions: Elements can combine in different ways to form different compounds, with mass ratios that are small whole-number multiples of each other. Compound Grams of oxygen Grams of nitrogen Ratio Nitric Oxide 8.10 grams7.60 grams1.07 : 1.00 NO Nitrous Oxide 16.3 grams8.1 grams2.01 : 1.00 NO 2

9. Law of Multiple Proportions Law of Multiple Proportions: Elements can combine in different ways to form different compounds, with mass ratios that are small whole-number multiples of each other.

10. Law of Conservation of Mass: Mass is neither created nor destroyed in chemical reactions. Aqueous solutions of mercury(II) nitrate and potassium iodide will react to form a precipitate of mercury(II) iodide and aqueous potassium iodide.

11. Law of Conservation of Mass: Mass is neither created nor destroyed in chemical reactions. Hg I 2 (s) + 2KNO 3 (aq)Hg(NO 3 ) 2 (aq) + 2K I(aq) 4.55 g + 2.02 g = 6.57 g 3.25 g + 3.32 g = 6.57 g

12. Dalton’s Atomic Theory 1.Elements are made up of tiny particles called atoms. 2.Each element is characterized by the mass of its atoms. 3.Chemical combination of elements to make different chemical compounds called molecules. 4.Chemical reactions only rearrange the way that atoms are combined in molecules; the atoms themselves don’t change.

13. Molecules are groups of atoms joined together

14. Each element has a unique individual mass (atomic mass) Atomic mass: The mass of a single element Molecular mass: The mass of a single molecule

15. Atomic masses can be used to find mass percentages Find the mass percentages of lead (Pb) and sulfur (S) in lead sulfide (PbS). Use the periodic table. Molecular mass of PbS Atomic mass of Pb Atomic mass of S =+ 207.2 amu + 32.07 amu = 239.3 amu % Lead (Pb) % Sulfur (S)

16. What is the atom made of? Cathode-Ray Tubes: J. J. Thomson (1856–1940) proposed that cathode rays must consist of tiny, negatively charged particles which we now call electrons. JJ Thompson reasoned that these cathode rays come from negatively charged particle called electrons.

17. Subatomic particles of the atom

18. Atomic Structure: Protons and Neutrons Rutherford’s Scattering Experiment Ernest Rutherford (1871–1937) directed a beam of alpha particles at a thin gold foil. A very small number, (about 1 in every 20,000) were deflected at an angle.

19. Rutherford’s Scattering Experiment Rutherford proposing that an atom must be almost entirely empty space and have its mass concentrated in a tiny central core that he called the nucleus.

21. A comparison of subatomic particles The mass of the atom is primarily in the nucleus.

22. A comparison of subatomic particles The charge on a proton is opposite the charge of the electron

23. Atomic Numbers, Mass Numbers, Isotopes Atomic Number (Z): Number of protons in an atom’s nucleus. Equivalent to the number of electrons around the atom’s nucleus. Mass Number (A): The sum of the number of protons and neutrons in an atom’s nucleus. This is the atomic mass in units of amu (1 amu = 1.661 x 10 -24 g). Isotopes: Atoms with identical atomic numbers but different mass numbers.

24. How to write an atoms atomic numbers, mass numbers for different Isotopes Carbon-14 C 14 6 Atomic number Mass number Carbon-12 C 12 6 Atomic number Mass number 6 protons 6 electrons 8 neutrons 6 protons 6 electrons 6 neutrons

25. Atomic Mass, Atomic Number The mass of 1 atom of carbon-12 is defined to be 12 amu. Atomic Mass: The weighted average of the isotopic masses of the element’s naturally occurring isotopes.

26. Why is the atomic mass of the element carbon 12.01 amu? Carbon-12:98.89 % natural abundance12.0000 amu Carbon-13:1.11 % natural abundance13.0034 amu Carbon-14: 1 x10 -12 % natural abundance 14.0032 amu = 12.01 amu = (12.00 amu)(0.9889) + (13.0034 amu)(0.0111) = 11.87 amu + 0.144 amu Average mass of a Carbon atom Atomic masses on the periodic table are weighted averages of all naturally occurring isotopes