Post on 02-Jan-2016
transcript
Chapter 3 Notes
The Atom: From Idea to Theory
Historical BackgroundIn 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 Cont’d
Wt. of H
Wt. of O
H + O H2O 2 16
H + O H2O2 2 32
The ratio of O in H2O2 to O in H2O = _______2 to 1______________
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) + O2 (g) --------> CO2 (g)
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
The Structure of the Atom Cont’d
o Charge and Mass of the electron o Thomson 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
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.
Important Subatomic Particles
a.m.u. Mass, kg Charge Location
Job
Proton (p+)
1 1.67265×10-27 +1 nucleus
ID
Neutron (n°)
1 1.67495×10-27 0 nucleus Stabilize atom
Electron (e-)
0 9.10953×10-31 -1 clouds Bonding
Important Subatomic Particles Cont’d
Electrostatic force - pulls nuclei apart: protons and neutrons
Strong Nuclear Force- force holds nuclei together
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 nucleusUniquely labels each element
Mass Number (M) Number of protons + neutrons in the nucleus
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 cationsNegative ions are anions
Weighing and Counting Atoms
If the mass # comes from the p+ and n0 [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.
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)
Lets Practice!
p+ e- n° Atomic # =(# of p+)
Mass # =(p+ + n0)
C 6 6 6 6 12
Ca 20 20 20 20 40
U 92 92 146 92 238
Cl 17 17 18 17 35
Mg 12 12 12 12 24
14C 6 6 8 6 14
S-2 16 18 16 16 32
Na+1 11 10 12 11 23
Isotopes
Two atoms of the same element (same # of p+) but with different weights (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 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
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.
The Mole, Avogadro’s number and Molar Mass
The MoleAtoms 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.
1 mole = 1 gram formula mass = 6.02 x 1023 particles
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 - H2 H2O Ca(OH)2
2.0g 18.0 g 74.1 g
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.
The Mole, Avogadro’s number and Molar Mass
Examples: 1 mole Na = 6.02 x 1023 atoms = 23.0 g
1 mole O2 = 6.02 x 1023 molecules =
1 mole HCl = 6.02 x 1023 molecules =
1 mole NaCl =6.02 x 1023 formula units=
Mole Map
Liters
Atoms, molecules, particles
Grams
Mole
22.4 L
6.022 x10 23Molar Mass
Examples
2 steppersconvert 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 O2 to moles
Examples
3 and 4 steppersconvert 1.2 x 1024 atoms of Magnesium to grams
convert 128 g of O2 to molecules of O2
convert 128 g of O2 to atoms of oxygen
Convert 100. g of Ar to liters of Ar