Ch 3: Atoms Problem Set Ch 3: page 89-90 1-3, 7-9, 12-13, 17-20, 22-24 Chapter 3 problem set: Problem Set Ch 3: page 89-90 1-3, 7-9, 12-13, 17-20, 22-24

3.1 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.

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

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 (H2O) is always: 11.2% H; 88.8% O Sugar (C6H1206) is always: 42.1% C; 6.5% H; 51.4% O Law of Multiple Proportions Video

Law of Multiple Proportions – John Dalton (English) - 1803 Ex1: wt. of H wt. of O H + O  H2O 2 16 H + 0  H2O2 2 32 The ratio of O in H2O2 to O in H2O = 32/16 = 2:1 (small whole numbers

Dalton’s Atomic Theory

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 cathode rays .

3.2 The Structure of the Atom 20 years later, J.J. Thomson (English) repeated those experiments and devised new ones. Cathode Ray Tube Thomson used a variety of materials, so he figured cathode ray particles must be fundamental to all atoms. 1897 - discovery of the electron.

3.2 The Structure of the Atom Thomson and Milliken (oil drop experiment) worked together (their data, not them) to discover the charge and mass of the electron

3.2 The Structure of the Atom Charge and Mass of the electron - Thomson and Milliken (oil drop experiment) worked together to discover the charge and mass of the electron Oil Drop charge = 1.602 x 10-19 coulomb this is the smallest charge ever detected mass = 9/109 x 10-28 g this weight is pretty insignificant

3.2 The Structure of the Atom 1909 - Gold Foil Experiment (Rutherford - New Zealand) Nuclei are composed of ‘nucleons’: protons and neutrons Alpha particles from Polonium (in the lead box) were released towards a thin sheet of gold foil. Most of the particles went through and were seen on the detector screen. 1 in 20,000 alpha particles bounced back.

Rutherford’s Conclusion Concluded: 1 – the positive portion of the atom is in the middle 2 – most of the atom is empty 3 – most of the mass is in the middle 4 – electrons orbit the nucleus Analogy: if an atom is the size of the Linc, then the nucleus is the size of a tennis ball floating in the middle of the stadium.

Table: Subatomic particles important in chemistry. http://curriculum.media.pearsoncmg.com/curriculum/science/cg_wsmw_chem_12/UntamedSci/CHEM-UT-04/player.html

Table: Subatomic particles important in chemistry. Nuclear Forces Forces in Atoms

3.3 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 nucleons Atomic Number (Z) Atomic # Number of protons in the nucleus Uniquely labels each element Mass Number (M) Mass # Number of protons + neutrons in the nucleus

Counting nucleons

Counting electrons Atoms Ions Same number of electrons and protons Ionic charge (q) = #protons - #electrons Positive ions are cations Negative ions are anions

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)

Lots of Practice!!! p+ e- n° Atomic # = (# of p+) Mass # = (p+ + n0) C 6 12 Ca 20 40 U 92 146 238 Cl 17 18 35 Mg 24 14C 8 14 S-2 16 32 Na+1 11 10 23

Isotopes Isotopes Isotope Video Two atoms of the same element (same # of p+) but with different masses (different # of n0)

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 12C is the standard upon which all other masses are based. 1/12 of the mass of 1 atom of 12C = 1 amu

Carbon consists of two isotopes: 98. 90% is C-12 (12. 0000 amu) Carbon consists of two isotopes: 98.90% is C-12 (12.0000 amu). The rest is C-13 (13.0034 amu). Calculate the average atomic mass of carbon to 5 significant figures. (.9890)(12.0000)+(.0110)(13.0034)=x 11.8680+.1430=12.011

Ex1: Chlorine consists of two natural isotopes, 35Cl (34. 96885) at 75 Ex1: 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. (.7553)(34.96885)+(.2447)(36.96590)=x 26.41+9.045=35.46 Ex2: 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. (.5725)( 120.9038)+(.4275)(x) =121.8 69.22+.4275x=121.8 -69.22- 69.22.4275x= 52.59x=123.0

The Mole, Avogadro’s number and Molar Mass The Mole Mole Video 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 1023 particles of that substance. Avogadro 1 mole = 1 gram formula mass = 6.02 x 1023 particles

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 - H2 H2O Ca(OH)2 2.0g 18.0 g 74.1 g

Official names may also be: 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. Molar Mass

Examples: 1 mole Na = 6.02 x 10 23atoms = 23.0g 1 mole O2 = 6.02 x 1023 molecules=32.0g 1 mole HCl = 6.02 x 1023 molecules = 36.5 g 1 mole NaCl =6.02 x 1023 formula units = 58.5 g

Mole Relationships 1 mole 6.02 x 1023 atom/molecule P-Table for gram 22.4 liter (at STP