Ch 4: Atoms

Scientists (key experiments and contributions) Democritus Lavosier Dalton Crookes Thomson Miliken Rutherford Counting P,N,E Atoms, Ions, Isotope Cation vs anion Define Ion and Isotope Average Atomic Mass Problems

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 + 2B + C —> AB2C

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

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

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 (video) 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 electron. Plum Pudding Model https://www.youtube.com/watch?v=O9Goyscbazk

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 charge = 1.602 x 10-19 coulomb this is the smallest charge ever detected mass = 9.11 x 10-28 g this weight is pretty insignificant

3. 2 The Structure of the Atom https://www. youtube. com/watch 1909 - Gold Foil Experiment (Rutherford - New Zealand) Nuclei are composed of ‘nucleons’: protons and neutrons

A) Oil Drop B) Found Nucleus C) Law of conservation of mass D) Plum Pudding Model E) Coined the term cathode ray F) Term atom G) Law of multiple proportions H) Cathode Ray Tube experiments I) Gold Foil Experiment J) Charge and Mass of electron 1) Lavosier 2) Democritus 3) Milliken 4) Rutherford 5) Dalton 6) Thomson 7) Crookes

Table: Subatomic particles important in chemistry.

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. Atomic Number (Z) -Number of protons in the nucleus - Uniquely labels each element Mass Number (M) - Number of protons + neutrons in the nucleus

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)

Practice- How many Protons? 23 16 7

Practice- How many Electrons? 23 16 7

Practice- How many Neutrons? 23 16 7

Lets try a few together

Lots of Practice!!! p+ e- n° Atomic # = (# of p+) Mass # = (p+ + n0) C Ca U Cl Mg 14C S-2 Na+1

Protons Neutrons Electrons Sr Sr +2 O -2 H+ B +3

Protons Neutrons Electrons 38 50 36 8 10 1 5 6 2 Sr Sr +2 O -2 H+ B +3

Complete Practice Problems

Isotopes Isotopes Two atoms of the same element (same # of p+) but with different weights (different # of n0) H-1, H-2,H-3 C-12, C-13, C-14 U-235, U-238

How do you determine your grades in this class? 60% tests: 90 40% other: 100

Average Atomic Mass (“weighted average”) Definition - The average weight of the natural isotopes of an element in their natural abundance.

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=35.46amu 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 69.2174255+.4275x=121.8 .4275x= 52.5825745 =123.0amu

Practice Worksheet Problems

Beanium Lab

Scientists (key experiments and contributions) Democritus Lavosier Dalton Crookes Thomson Miliken Rutherford Counting P,N,E Atoms, Ions, Isotope Cation vs anion Define Ion and Isotope Average Atomic Mass Problems