1. Chem 1151: Ch. 2 Atoms and Molecules

2. “It’s a theory not a fact” Hypothesis: A possible explanation. Theory: A proposed explanation that is consistent with all available evidence that has been repeatedly tested. A theory evolves from interpretation of data (facts). “Theory” is frequently misunderstood by the general public. Gravity (Newton and Einstein) is “just a theory”, but I predict that if you jump from the roof of your house, you will fall. The scientific method is a sequence of steps for systematically analyzing the natural world.

3. The origin of atoms Following the big bang  Expansion of space  Cooling  Formation of fundamental particles  Formation of low mass nuclei (H and He)  Star formation  Fusion reactions forming elements up to Iron-56 http://www.lbl.gov/abc/wallchart/chapters/10/0.html

4. Structure of the Atom Mass (g)Mass (u) Proton (p+)1.67 x 10 -24 1 Neutron (n)1.67 x 10 -24 1 Electron (e-)9.07 x 10 -28 1/1836 Most of the mass is actually in the nucleus Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7 th Edition, 2011

5. 4 Fundamental Forces in the Universe http://hyperphysics.phy-astr.gsu.edu/hbase/forces/funfor.html#c4 These forces generally considered manifestations of the same force.  Unified.

6. Synthesis of Matter Nucleosynthesis: Protons and neutrons join to form nuclei. Fusion: Multiple nuclei join to form heavier nucleus. Formation of heavier elements: Two protons collide – Releases positron, neutrino Nucleus with proton and neutron collides with another proton – releases gamma ray Two He-3 atoms collide – Produces He-4 – Releases two protons http://en.wikipedia.org/wiki/Image:FusionintheSun.png

7. Abundance of Elements in the Universe Mg K Cd http://en.wikipedia.org/wiki/File:SolarSystemAbundances.png Fuse 3 4 He atoms together  12 C, basic building block of all organic compounds  life. Elements with FW or AMU multiples of 4 are more abundant: (C, O, Si, Ca, etc.) Fusion Process continues up through Fe.

8. Periodic Table of the Elements All matter in our universe categorized in periodic table. – Based on atomic number (number of protons). Arranged in columns (groups) and rows (periods). – Groups have similar properties. – Periods correspond to filling of quantum shells by electrons.

9. Elements from Group 7A chlorine bromine iodine Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7 th Edition, 2011

10. Elemental Notation http://hyperphysics.phy-astr.gsu.edu/HBASE/pertab/pertab.html

11. Applications of Atomic and Mass Numbers – On the periodic table, the atomic number is written as a whole number above the symbol F. – In the written description, fluorine is said to have 9 protons (the atomic number is the number of protons). – In the symbol, the number 9 is written in the atomic number or Z (lower left) position. – Note: The periodic table does not show the mass number for an individual atom. It lists an average mass number for a collection of atoms! Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7 th Edition, 2011

12. Elemental Notation Q: How to represent lead-208? Q: How many p+, e-, n?82, 82, 126 Q: How to represent element X with 4 p+ and 5 n.

13. Isotopes Isotopes are atoms that have the same number of protons in the nucleus but different numbers of neutrons. That is, they have the same atomic number but different mass numbers. Because they have the same number of protons in the nucleus, all isotopes of the same element have the same number of electrons outside the nucleus. Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7 th Edition, 2011

14. Isotope Symbols Isotopes are represented by the symbol, where Z is the atomic number, A is the mass number, and E is the elemental symbol. Isotopes are also represented by the notation: Name-A, where Name is the name of the element and A is the mass number of the isotope. An example of this isotope notation is magnesium-26. This represents an isotope of magnesium that has a mass number of 26.

15. Relative Masses and Mass Units The extremely small size of atoms and molecules makes it inconvenient to use their actual masses for measurements or calculations. Relative masses are used instead. Relative masses are comparisons of actual masses to each other. For example, if an object had twice the mass of another object, their relative masses would be 2 to 1. An atomic mass unit is a unit used to express the relative masses of atoms. One atomic mass unit is equal to 1/12 the mass of a carbon-12 atom. A carbon-12 atom has a relative mass of 12 u because carbon- 12 has 6 protons and 6 neutrons. Mass (g)Mass (u) Proton (p+)1.67 x 10 -24 1 Neutron (n)1.67 x 10 -24 1 Electron (e-)9.07 x 10 -28 1/1836 Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7 th Edition, 2011

16. Determining Mass However, carbon has an actual mass listed of 12.011 u, not 12 u. Why are these values different? C 6 12.011 Mass (g)Mass (u) Proton (p+)1.67 x 10 -24 1 Neutron (n)1.67 x 10 -24 1 Electron (e-)9.07 x 10 -28 1/1836 12 Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7 th Edition, 2011

17. How Isotopes Determine Atomic Weight The atomic weight of an element is the relative mass of an average atom of the element expressed in atomic mass units. Many elements have more than 1 isotope (e.g. – 12 C, 13 C, 14 C). Abundance of isotopes are not evenly distributed. Weighted atomic mass of Carbon ( 12 C, 13 C only) = (0.98882*12u) + (0.01108 * 13.300335u) = 12.011u. C 6 12.011 12 C 6 12.011 12 AMU12 u13.300335 u

18. Determining Atomic Weight A specific example of the use of the equation is shown below for the element boron that consists of 19.78% boron-10 with a mass of 10.01 u and 80.22% boron-11 with a mass of 11.01u. This calculated value is seen to agree with the value given in the periodic table. Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7 th Edition, 2011

19. Molecular Weight The relative mass of a molecule in atomic mass units is called the molecular weight of the molecule. Because molecules are made up of atoms, the molecular weight of a molecule is obtained by adding together the atomic weights of all the atoms in the molecule. The formula for a molecule of water is H 2 O. This means one molecule of water contains two atoms of hydrogen, H, and one atom of oxygen, O. The molecular weight of water is then the sum of two atomic weights of H and one atomic weight of O: MW = 2(at. wt. H) + 1(at. wt. O) MW = 2(1.01 u) + 1(16.00 u) = 18.02 u Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7 th Edition, 2011

20. Molecular Weight The clear liquid is carbon disulfide, CS 2. It is composed of carbon (left) and sulfur (right). What is the molecular weight for carbon disulfide? Answer: MW = 1(atomic weight C) + 2(atomic weight S) 12.01 u + 2(32.07 u) = 76.15 u Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7 th Edition, 2011

21. The Mole THE MOLE CONCEPT APPLIED TO ELEMENTS – The number of atoms in one mole of any element is called Avogadro's number and is equal to 6.022x10 23. – A one-mole sample of any element will contain the same number of atoms as a one-mole sample of any other element. – One mole of any element is a sample of the element with a mass in grams that is numerically equal to the atomic weight of the element. EXAMPLES OF THE MOLE CONCEPT – 1 mole Na = 22.99 g Na = 6.022x10 23 Na atoms – 1 mole Ca = 40.08 g Ca = 6.022x10 23 Ca atoms – 1 mole S = 32.07 g S = 6.022x10 23 S atoms Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7 th Edition, 2011

22. The Mole THE MOLE CONCEPT APPLIED TO COMPOUNDS – The number of molecules in one mole of any compound is called Avogadro's number and is numerically equal to 6.022x10 23. – A one-mole sample of any compound will contain the same number of molecules as a one-mole sample of any other compound. – One mole of any compound is a sample of the compound with a mass in grams equal to the molecular weight of the compound. EXAMPLES OF THE MOLE CONCEPT – 1 mole H 2 O = 18.02 g H 2 O = 6.022x10 23 H 2 O molecules – 1 mole CO 2 = 44.01 g CO 2 = 6.022x10 23 CO 2 molecules – 1 mole NH 3 = 17.03 g NH 3 = 6.022x10 23 NH 3 molecules Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7 th Edition, 2011

23. What these numbers mean? Mass of 1 atom of Mg = 4.037 x 10 -23 g Mg atomic weight = 24.31 u How many atom of Mg in24.31 g Mg? 24.31 g Mg x 1 atom Mg = 6.022 x 10 23 atoms Mg 4.037 x 10 -23 g Mass of 1 atom of C = 1.994 x 10 -23 g C atomic weight = 12.01 u 12.01 g C x 1 atom C = 6.022 x 10 23 atoms C 1.994 x 10 -23 g The number of atoms of an element, in a mass equivalent to it’s atomic weight, is equal to Avogadro’s number (mole). or 1 u = 1 g/mole

24. Relationships: Mass, Moles, Molecular Weight 1 mol C atoms = 6.022 x 10 23 atoms C 6.022 x 10 23 atoms C = 12.01 g C 1 mol C atoms = 12.01 g C C has atomic weight of 12.01 u or 12.01 g/mol 1 mol S atoms = 6.022 x 10 23 atoms S 6.022 x 10 23 atoms S = 32.1 g S 1 mol S atoms = 32.1 g S S has atomic weight of 32.1 u or 32.1 g/mol

25. Compound (Molecular) Formulas The compound (molecular) chemical formula represents the numerical relationships that exist between atoms in a compound. This also applies to moles. Compound formula: all elements and number of each in a compound Examples: Urea1C, 4H, 2N, 1O Hydrofluoric acid 1H, 1F Sodium bicarbonate2H, 1C, 3O Sodium Azide 1Na, 3N 1 mol of H 2 SO 4 contains 2 mol of H 1 mol of S 4 mol of O

26. Mole Calculations (continued) The mole concept applied earlier to molecules can be applied to the individual atoms that are contained in the molecules. An example of this for the compound CO 2 is: 1 mole CO 2 molecules = 1 mole C atoms + 2 moles O atoms 44.01 g CO 2 = 12.01 g C + 32.00 g O 6.022x10 23 CO 2 molecules = 6.022x10 23 C atoms + (2) 6.022x10 23 O atoms Any two of these nine quantities can be used to provide factors for use in solving numerical problems. Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7 th Edition, 2011