Brief Timeline of Atomic Theory

Democritus 400BC Greek philosopher

Hard Particle (Cannonball)Theory Proposed that they world was made up of tiny, indivisible particles moving through a void of empty space “atom” comes from the Greek word “atomos”, meaning indivisible (cannot be divided)

John Dalton 1808 AD First modern atomic theory

Daltons Atomic Theory All matter is composed of tiny, indivisible particles called atoms All atoms of an element are identical Atoms of different elements are all different Atoms combine in simple ratios to form compounds

J.J. Thomson 1897-1904 “Plum Pudding Model” Cathode Ray tube experiment demo

demo

Cathode Ray Tube Thompson showed that cathode rays (electrons) were composed of negatively charged particles that separated from the gas atoms inside the tube Significant because: this meant that atoms are not hard, indivisible particles. Atoms are composed of smaller “subatomic” particles

Thomson’s Plum Pudding Model The atom was a hard sphere that was positively charged with negatively charged electrons that “dotted” the atom like raisins in plum pudding

The discovery of radioactivity Henri Becquerel 1896 Discovered that uranium ore released rays that could expose photographic film

The discovery of radioactivity Marie & Pierre Curie Extracted 2 new elements from uranium (U)ore: radium (Ra) and polonium (Po) Marie Curie

Magnetic Field Experiment Gold Foil Experiment (1911) Ernest Rutherford Magnetic Field Experiment Gold Foil Experiment (1911) Was able to separate radioactive rays into 2 types: alpha (α) & beta (Β) Determined that α rays were composed of helium nuclei (He +2 charge) Lead to discovery of the nucleus, as a positively charged center of atom, containing the mass Most of the atom is negatively charged empty space, electrons are outside the nucleus

Magnetic Field Experiment

Gold Foil Experiment

Gold Foil Experiment

Gold Foil Experiment

Gold Foil Experiment

Rutherford’s Atomic Model

Rutherford’s “Nuclear Model” Most of the atom is negatively charged empty space, surrounding a small, positively charged nucleus, containing most of the mass of the atom

Modern Theory of Atomic Structure Developed by Niels Bohr, based on the science of nuclear physics Bohr determined that an element's position on the periodic table was related to its electron configuration.

Electron configuration  Electron configuration – shows how many electrons are in each energy level or “ring” Ex: Carbon 2-4

Bohr’s Planetary Atomic Model Niels Bohr (1922) Determined that electrons rotate around the nucleus in discrete paths or rings

Planetary Model of Atomic Structure

Wave-Mechanical Model Current (modern) theory of atomic structure Moseley used x-ray analysis to calculate an integer for each element: these integers are the atomic numbers

Wave-Mechanical Model There is a tiny, dense positively charged nucleus at the center of a huge negatively charged electron cloud

Wave-Mechanical Model

Orbital Region of probability of finding an electron

“The whole point:” The modern model of the atom is the result of many investigations that have been revised over a long period of time by many scientists Atomic theory song

Place the models of atomic structure in order from earliest to the modern theory:

Basic Atomic Structure The nucleus occupies less than 0.01% of the total volume of an atom but accounts for 99.97% of its mass. Thus most of an atom is EMPTY SPACE where the ELECTRONS are found, this is called an ELECTRON CLOUD. One atomic mass unit is 1/12TH THE MASS OF A CARBON-12 ATOM. This is the standard by which the masses of all other elements are determined. It is abbreviated “u”.

Start here with lecture notes!!

Subatomic Particles Proton p+ In Nucleus +1 1 Neutron n0 In nucleus Symbol Location Charge Mass (amu) Mass (g) Proton p+ In Nucleus +1 1 Neutron n0 In nucleus Electron e- Outside nucleus -1

Nucleons: Def: particles found in the nucleus : protons & neutrons Number of Nucleons = Atomic Mass

Nuclear Charge: Def: equal to the number of protons (b/c only protons have charge, neutrons are neutral) For this atom the nuclear charge would be +6, because there are 6 protons

protons neutrons protons neutral protons electrons 6 6 6

2713Al 3517Cl 11H 20782Pb Signature Atomic Number Mass Number Nuclear charge # of PROTONS # of NEUTRONS # OF ELECTRONS 2713Al 3517Cl 11H 20782Pb

Signature Atomic Number Mass Number Nuclear charge # of PROTONS # of NEUTRONS # OF ELECTRONS 2713Al 13 27 +13 14 3517Cl 17 35 +17 18 11H 1 +1 20782Pb 82 207 +87 87 125

The only number that never changes for an element is ATOMIC NUMBER !!

24?N 64?Cu Atomic number:___________ # protons: ______________   Atomic number:___________ # protons: ______________ # electrons: _____________ Mass number: ___________ # neutrons: _____________ Name-mass:_______________________________ 64?Cu

147N 3115P 2713Al 10847Ag Atomic number:___________   Atomic number:___________ # protons: ______________ # electrons: _____________ Mass number: ___________ # neutrons: _____________ Name-mass:______________________________ _ 3115P Name-mass:_______________________________ 2713Al 10847Ag

Atomic Structure Practice: Determining Subatomic Particles 1.) Complete the following definitions: a. ATOM b. ATOMIC NUMBER – equals the number of ________ in an atom c. MASS NUMBER – equals the number of ________ + _________ d. Charge of a PROTON - ________ Mass of a PROTON- ________ e. Charge of a NEUTRON - ______ Mass of a NEUTRON -_______ f. Charge of an ELECTRON- _____ Mass of an ELECTRON- ______

WORD BANK: proton, neutron, electron, nucleus, electron cloud

For the atom shown above: f. Number of Protons= _______________ g. Atomic Number =________________ h. Number of Neutrons= _____________ i. Mass Number = _______________ Using the ATOMIC NUMBER, identify the element name___________ and SYMBOL _____

Atomic Structure 1 Name Symbol Atomic Number Mass Number Charge # of Protons # of Neutrons # of electrons   199F Helium-4 11 23 Nitrogen-14 14 32 16 6429Cu  Manganese-60  6025Mn 25 60  25  35 Barium-137  13756Ba 56  137  81 56  Iodine-131  13153I 53 131 78 53   Iodine-127  12753I 127  74

Atomic Structure 1 Name Symbol Atomic Number Mass Number Charge # of Protons # of Neutrons # of electrons   199F Helium-4 11 23 Nitrogen-14 14 32 16 6429Cu 25 35 81 56 53 131 74

**Shade the columns representing the nucleons light blue Phosphorus-32   146C Potassium-39 16 8 5626Fe 18 40 29 35 79 197 2412Mg **Shade the columns representing the nucleons light blue

Changes in number of subatomic particles Isotopes Ions Change in number of neutrons Same atomic number, different mass Same number protons, different number neutrons Change in number of electrons A cation is positive ion, results from loss of electrons, reducing radius An anion is negative ion, results from gain of electrons, increasing radius

Name-mass:_______________________________   Atomic number:___________ # protons: ______________ # electrons: _____________ Mass number: ___________ # neutrons: _____________ Name-mass:_______________________________ 2311? ??Li ??I

ISOTOPE Forms of the same element having different mass due to different number of neutrons. Indicated by “element name-mass”

158O 168O Name: _______________ Mass: ________________ Protons: ______________ Neutrons: _____________ Name: _______________ Mass: ________________ Protons: ______________ Neutrons: _____________

Practice: Name Symbol Atomic # Mass # # Protons # Neutrons # Electrons   235U 238U Carbon-12 Carbon-13

The mass on the periodic table is the MOST ABUNDANT mass! ** You can estimate which isotope is found in the highest abundance as the one with a mass closest to the mass listed on the periodic table Example: Chlorine-35 mass 34.969g Chlorine-37 mass 36.966g Look on the periodic table for the mass of chlorine____________________________ The more abundant isotope has a mass closer to the mass given on the periodic table_____________

Practice: Which isotope of silicon would be found in the highest percentage? 2814Si, mass 27.977 2914Si, mass 28.976 3014Si, mass 29.974   Why? _______________________________________ _______________________________________

Atomic Structure 2 Isotopic Notation Number of protons   Number of protons Number of neutrons Number of electrons Mass number 1.Oxygen-16 O-16 16O 2.Oxygen-18 3. Ar-40 4. 18 5. 16 32 6. 34S 7. 19 20 8. 41 9. Iron- 10. 57Fe 11. 26 12. Ne-20 13. 10 22 14.Hydrogen- 1 15. H-2 16. 3H

2.) Calculate the weighted average of the following naturally occurring isotopes. SHOW ALL WORK! a.) 95.50%7Li & 7.50% 6Li d.) 99.63%14N & 0.37%15N   b.)80.20%11B & 19.80%10B e.) 78.9%24Mg, 10.00%25Mg, & 11.01%26Mg c.)95.02%32S, 0.75%33S, & 4.21%34S f.) 92.23%28Si, 4.67%29Si, & 3.10%30Si

IONS A charged part of an atom, resulting from the loss or gain of electrons VALENCE electrons: outermost electrons, the last number in an electron configuration KERNEL electrons: all electrons except valance electrons

Electron configuration  Electron configuration – shows how many electrons are in each energy level or “ring” Ex: Carbon 2-4

Electron configuration of sodium:

2 diagrams of atomic structure: Bohr diagrams Lewis electron dot diagrams Bohr realized that the rows on the periodic table corresponded to the number of shells of electrons Lewis realized that the groups/families on the periodic table correspond to the number of valence electrons This model shows the nucleus, indicating the number of protons and neutrons, surrounded by rings, representing each energy level This model shows the element symbol surrounded by dots, representing the valence electrons. You must place one dot at each (3, 6,9,12 o’clock) location before “doubling up” (exception: Helium)   189F electron configuration 2-7 F electron configuration 2-7

Bohr Atomic Structures 1 18 Bohr Atomic Structures tables to fill in the electron 4   configurations, as shown, then draw the Bohr Atomic Structure for each element 1-20. 2 13 14 15 16 17 7 9 11 12 19 20 3 5 6 8 10 2-1 2-2 2-3 2-4 23 24 27 28 31 32 35 40 2-8-1 2-8-2 2-8-3 2-8-4 39 Rules: 1.) Show placement 2.) The nucleus is 3.) Indicate the number of ALL electrons represented by a center of electrons in each circle showing the energy level, by writing *use atomic # # of protons & the the number on each ring. OR the entire # of neutrons 2-8-8-1 2-8-8-2 electron configuration ** closest to nucleus is 1st

2 Main Types of Ions: anion A negative ion Ex: Cl-, O-2 cation A positive ion Ex: Na+, Al+3

The octet rule Atoms will gain or lose electrons in order to have a full valence shell of 8 electrons. Exception: Helium can have a maximum of 2 valance electrons

When an atom gains 1 or more electrons It becomes a negative ion and it’s radius increases. A negative ion is an anion.

When an atom loses 1 or more electrons It becomes a positive ion and it’s radius decreases. A positive ion is a cation.

Change ending of element to “-ide”   CATION ANION Definition positive ion negative ion Results from Loss of electron(s) Gain of electron(s) Indicated by (+) charge (-) charge What happens to radius??? Gets smaller Gets bigger Na Na+ Naming “Element name-ion” Change ending of element to “-ide” Lewis Dot Structure [Na] + .. [:.F.:]-

How to predict if an element will form an anion or cation: The “electron clock”: 8/0 7 1 6 2 5 3 4 # valance electrons

Atomic Structure 3: Predicting Ions Element Electron configuration Lewis dot structure of atom Lose or gain electrons? How many electrons lost or gained? Ionic Charge ** Lewis dot structure of ion Radius increase or decrease? F 2-7 gain 1 -1 increase Mg 2-8-2 lose 2 +2 decrease O   Al N Fr C

Electron configuration Lewis dot structure of atom Element Electron configuration Lewis dot structure of atom Lose or gain electrons? How many electrons lost or gained? Ionic Charge ** Lewis dot structure of ion Radius increase or decrease?   2-8-8-1 2-8-7 2-8-18-18-8-2 2-8-6 2-8-5 2-3 **In the “ionic charge” column only: shade the cation charges red and the anion charges blue

Atomic Structure 4 3517Cl 2311Na 94Be 6530Zn 147N 3216S 2010Ne 12753I   # of Protons # of Neutrons # of Electrons Nuclear Charge Bohr Diagram of Atom Lewis Dot of Atom Predict Ionic Charge Lewis Dot of Ion Name of Ion ex 3517Cl 17 18 +17 Cl -1 Chloride 1 2311Na 2 94Be 3 6530Zn 4 147N 5 3216S 6 2010Ne 7 12753I 8 10847Ag 9 7031Ga 10 126C

Atomic Spectra

Radiant Energy Energy that travels through space as electromagnetic waves at the speed of light

Electromagnetic Spectrum Includes all types of radiant energy from gamma rays (hi E) to radiowaves (lo E) Visible light is only a small portion of the spectrum

1 photon = 1 quantum Quanta: tiny packets of energy released or absorbed by objects *Einstein and Plank determined that energy is released or absorbed in a continuous flow of small packets or quantum/photons

Release or Absorption of Energy:

Higher energy levels (excited state) Electrons absorb energy when jumping to Electrons release energy when falling to Lower energy levels (ground state)

Bohr used the emission spectrum as proof of planetary model But his model only works for hydrogen because he didn’t account for electrons moving between energy levels

Spectral Lines Characteristic wavelengths (λ) of photons of energy released as electrons fall from hi to lo energy

Spectral lines demo: Salt of Element Color of Flame Strontium Chloride   Strontium Chloride Barium Chloride Copper (II) Chloride Lithium Chloride Potassium Chloride Identity Unknown Element Unknown Mixture

Emission Spectrum:

Each element has it’s own characteristic spectrum:

Compare H & He: hydrogen helium

Because electrons do move between energy levels, emitting “spectral lines”, we had to change our view of atomic structure:

Excited State Electron Configurations Occurs when elements absorb energy and jump to a higher energy level. ** it will not look like it is written on periodic table, be sure they add to the correct number! Ground state: 2-8-1 Excited state : 2-7-2

“Crib Sheet” #p+ = atomic number *#n0 = mass-atomic number #e- = #p+ - charge (use the sign of the charge) Isotope: same #p+, different #no OR same atomic number, different mass To calculate weighted average: (%/100 x atomic mass) + (%/100 X atomic mass) + ….. *Ion: same # p+, different #e- Charge= #p+ - #e-

Atomic Structure Review p. 17 11 9 43 92 118 13 4 9. Br 10. C 11. Sn 12. Zn 13. Cl 14. 40 15. 16

Atomic Structure Review p. 17 17) =(.789x24)+(.10x25)+(.1101x26) = 18.936 + 2.5 + 2.8626 = 24.2986 = 24.30g 16.) =(.925x7) + (.0750x6) =6.475 + .45 =6.925 =6.93g

Atomic Structure Review p. 18 18.) 2-8-1 19.) Na 20.) 2-7-2 21.) 19 22.) 1 23.) Y 24.) Ar 25.) Not possible 27.) as electrons fall from excited state to ground state energy is released as radiant energy (spectral lines). 28.) you can ID the gas element using spectral line analysis. 29.) electrons are negatively charged particles. B has 5 e-, its e- config. is 2-3, with 2 e- in the 1st energy level and 3 e- in the 2nd (valence) level

Atomic Structure Review MC?s 1.) 2 13.) 1 2.) 4 14.) 4 3.) 1 15.) 1 4.) 1 16.) 3 5.) 3 17.) 4 6.) 1 18.) 3 7.) 4 19.) 2 8.) 2 20.) 2 9.) 4 21.) 3 10.) 3 11.) 4 12.) 3 pg 19-20 1.) 4 13.) 2 2.) 3 14.) 4 3.) 2 15.) 3 4.) 3 16.) 3 5.) 2 17.) 2 6.) 3 18.) 1 7.) 1 19.) 1 8.) 2 20.) 4 9.) 3 21.) 4 10.) 1 22.) 2 11.) 3 23.) 3 12.) 1 pg 21-22

Atomic Structure Review p. 23 1.) 19p, 20n, 18e 2.) 9p,10n,10e 3.) 5p,6n,2e 4.) 15p,16n,18e 5.) 16p,16n,18e 6.) 14p,14n,10e 7.) 7p,7n,10e 8.) 20p,20n,20e 9.) 37p,48n,36e 10.) 53p,75n,54e 11.) 30p,35n,28e 12.) 6p,6n,10e