1. Chapter 3 Notes

2. The Atom: From Idea to Theory Historical Background  In approximately 400 BC, Democritus (Greek) coins the term“atom” (means indivisible). Before that matter was thought to be one continuous piece - called the continuous theory of matter. Democritus creates the discontinuous theory of matter. His theory gets buried for thousands of years  18th century - experimental evidence appears to support the idea of atoms.

3. Law of Conservation of Mass Antoine Lavosier (French) -1700’s The number of each kind of atoms on the reactant side must equal the number of each kind of atoms on the product side A + B + C ABC

4. Law of Multiple Proportions John Dalton (English) - 1803 The mass of one element combines with masses of other elements simple in whole number ratios. Water (H 2 O) is always: 11.2% H; 88.8% O Sugar (C 6 H 12 0 6 ) is always: 42.1% C; 6.5% H; 51.4% O

5. Law of Multiple Proportions Cont’d Wt. of H Wt. of O H+OH2OH2O216 H+OH2O2H2O2 232 The ratio of O in H 2 O 2 to O in H 2 O = _______2 to 1______________

6. Dalton’s Atomic Theory 1. Everything is made of atoms. 2. Atoms of the same element are identical. (NOT) 3. Atoms can not be broken down, created, or destroyed. (NOT) 4. Atoms combine in simple whole number ratios to form chemical compounds. 2:1 1:1 5. A chemical reaction is the combining, separation, or rearrangement of atoms. C (s) + O 2 (g) -------->CO 2 (g)

7. 3.2 The Structure of the Atom Updating Atomic Theory  1870’s - English physicist William Crookes - studied the behavior of gases in vacuum tubes(Crookes tubes - forerunner of picture tubes in TVs). Crookes’ theory was that some kind of radiation or particles were traveling from the cathode across the tube. He named them electrons  20 years later, J.J. Thomson (English) repeated those experiments and devised new ones. Thomson used a variety of materials, so he figured cathode ray particles must be fundamental to all atoms. 1897 - discovery of the positive charge Plum Pudding Model

8. The Structure of the Atom Cont’d oCharge and Mass of the electron oThomson and Milliken (oil drop experiment) worked together to discover the charge and mass of the electron  charge = 1.592 × 10 −19 coulombs this is the smallest charge ever detected  mass = 9.1093821545 × 10 −31 g this weight is pretty insignificant

9. The Structure of the Atom Cont’d 1909 - Gold Foil Experiment (Rutherford - New Zealand) Nuclei are composed of ‘nucleons’: protons and neutrons Top: Expected results: alpha (+) particles passing through the plum pudding model of the atom undisturbed. Bottom: Observed results: a small portion of the particles were deflected, indicating a small, concentrated positive charge.

10. Important Subatomic Particles a.m.u.Mass, kgChargeLocatio n Job Proton (p + ) 11.67265×10 -27 +1 nucleus ID Neutron (n ° ) 11.67495×10 -27 0nucleusStabilize atom Electron (e - ) 09.10953×10 -31 cloudsBonding

11. Important Subatomic Particles Cont’d Electrostatic force - pulls nuclei apart: protons and neutrons Strong Nuclear Force- force holds nuclei together

12. Weighing and Counting Atoms We look to the periodic table to give us information about the number of particles are in atoms and also to help us count atoms in a sample. Counting Atoms  Atomic Number (Z) Number of protons in the nucleus Uniquely labels each element  Mass Number (M)  Number of protons + neutrons in the nucleus

13. Weighing and Counting Atoms Counting electrons  Atoms Same number of electrons and protons  Ions – lost or gained electrons  Ionic charge (q) = #protons - #electrons Positive ions are cations Negative ions are anions

14. Weighing and Counting Atoms If the mass # comes from the p + and n 0 [each with masses of exactly 1], why don’t the atomic weights/masses of the all elements turn out to be whole numbers? Because the atomic weights/masses on the P- Table are the “weighted averages,” of the naturally occurring isotopes of the element. (remember: ignore the mass of the e -, it’s too small to care about.

15. Review of Formulas atomic # (Z) - (always a whole number, smaller number on the periodic table) = # of protons in the nucleus - also indicates the # of electrons if the element is not charged atomic mass – the average mass of all of the isotopes of an element – is a number with a decimal – is always the larger number on the periodic table. mass number (A) - sum of the protons and neutrons in a nucleus this number is rounded from atomic mass due to the fact that there are isotopes # neutrons = A - Z example - # of neutrons in Li = 6.941-3 = 3.941 rounds to 4 Ion – a charged atom. Atoms become charged by gaining electrons (become a negative charge) or losing electrons (become a positive charge)

16. Lets Practice! p+p+ e-e- n°n° Atomic # = (# of p + ) Mass # = (p + + n 0 ) C666612 Ca20 40 U92 14692238 Cl17 181735 Mg12 24 14 C668614 S -2 161816 32 Na +1 1110121123

17. Isotopes Two atoms of the same element (same # of p + ) but with different weights (different # of n 0 ) Average Atomic Mass (“weighted average”) Definition - The average weight of the natural isotopes of an element in their natural abundance. History lesson - originally H was the basis of all atomic masses and was given the mass of 1.0. Later, chemists changed the standard to oxygen being 16.000 (which left H = 1.008). In 1961, chemists agreed that 12 - C is the standard upon which all other masses are based. 1/12 of the mass of 1 atom of 12 - C = 1 amu

18. Isotope Calculations Carbon consists of two isotopes: 98.90% is C-12 (12.0000 amu). The rest is C-12 (13.0034 amu). Calculate the average atomic mass of carbon to 5 significant figures. 12.011 amu Chlorine consists of two natural isotopes, 35Cl (34.96885) at 75.53% abundance and 37Cl (36.96590) at 24.47% abundance. Calculate the average atomic mass of Chlorine. 35.46 Antimony consists of two natural isotopes 57.25% is 121Sb (120.9038). Calculate the % and mass of the other isotope if the average atomic mass is 121.8.

19. The Mole, Avogadro’s number and Molar Mass The Mole  Atoms are tiny, so we count them in “bunches”.  A mole is a “bunch of atoms”.  The Mole (definition) -The amount of a compound or element that contains 6.02 x 10 23 particles of that substance. 1 mole = 1 gram formula mass = 6.02 x 10 23 particles

20. The Mole, Avogadro’s number and Molar Mass Molar Mass  Molar Mass - the sum of the atomic masses of all atoms in a formula  Round to the nearest tenth! (measured in amu or grams)  ex - H 2 H 2 OCa(OH) 2 2.0g 18.0 g 74.1 g

21. The Mole, Avogadro’s number and Molar Mass Molar mass is a term that can be used for atoms, molecules (covalent compounds or elements) and formula units (ionic compounds) Official names may also be:  Formula mass (ionic compounds)  Molecular mass (covalent compounds and diatomic elements)  Atomic weight, Atomic mass, grams formula weight, etc.

22. The Mole, Avogadro’s number and Molar Mass Examples:1 mole Na = 6.02 x 10 23 atoms = 23.0 g 1 mole O 2 = 6.02 x 102 3 molecules = 1 mole HCl = 6.02 x 10 23 molecules = 1 mole NaCl =6.02 x 10 23 formula units=

23. Mole Map Liters Atoms, molecules, particles Grams Mole 22.4 L 6.022 x10 23 Molar Mass

24. Examples 2 steppers  convert 13.8 g Li to moles  convert 2.0 moles Ne to g  convert 3.0 moles of Be to atoms  convert 44.8 L of O 2 to moles

25. Examples 3 and 4 steppers  convert 1.2 x 1024 atoms of Magnesium to grams  convert 128 g of O 2 to molecules of O 2  convert 128 g of O 2 to atoms of oxygen  Convert 100. g of Ar to liters of Ar