Atomic Structure
What is a theory? An organized system of accepted knowledge that applies in a variety of circumstances to explain a specific set of phenomena
Early Theories Democritus: 4 B.C.: atom He thought that you could cut something until it could no longer be divided and he called this indivisible particle atomos.
Dalton: 1766-1844 >All elements composed of tiny particles called atoms >Atoms of same element are identical; atoms of different elements are different >Atoms of different elements can physically mix together or chemically combine to form compounds >Chemical reactions cannot change atoms of one type of element to another
J.J. Thomson: 1856-1940 >discovered electrons in 1897 >used a cathode ray tube >the ray produced was deflected by an electrical field (showed that atoms had particles with (-) charge)
Summary- Atomic Models Dalton s Sphere
Thomson s Plum Pudding
Rutherford: 1871-1937 >Gold Foil Experiment >Discovered the nucleus
Rutherford s Gold Foil Experiment Experiment Shot positively charged alpha particles at gold foil Results 1) Most particles passed through the foil 2) A few were deflected
Rutherford s Gold Foil Experiment Conclusions: (Important to Remember) 1) The atom has a small, dense, positively charged core (nucleus). 2) The rest of the atom is empty space.
Rutherford s Nuclear Model
Modern Theories Bohr planetary model electrons arranged in concentric circular patterns paths or orbits around nucleus (energy level)
Bohr Model of the Atom
Wave-Mechanical Model Electron Cloud Model Orbitals are the area of highest probability where an electron will be found. Orbitals have a variety of shapes and names (s, p, d, f)
Example: Wave Mechanical Model An orbital predicts the probability of finding an electron in a given area.
Wave-Mechanical Model
Summary- Atomic Models Dalton s Sphere
Thomson s Plum Pudding
Rutherford s Nuclear Model
Bohr s Planetary Model
Wave-Mechanical Model
1 amu = 1/12 th mass of a carbon-12 atom Subatomic Particles Name Symbol Charge Mass Proton (located in nucleus nucleon) Neutron (located in nucleus nucleon) Electron (located outside the nucleus in orbitals) p + +1 1 amu n 0 0 1 amu e - -1 1/1836 amu **Note: amu = atomic mass unit
Atomic Number Equal to the number of protons Every element has its own atomic number See Periodic Table 6 C
Mass Number Equal to the sum of the protons and the neutrons (whole number) Can be written as carbon-12 12 C
To find: # of protons look up atomic number on Periodic Table
To find: # of electrons in a neutral atom, it is equal to the number of protons
To find: # of neutrons if protons + neutrons = mass then, # of neutrons = mass # - # protons
Practice Element Atomic # Mass # # of protons Ca # of neutrons # of electrons Mg Na He
Do Now 1. What is the nuclear charge of a sulfur atom? 2. If an atom has 5 protons and 6 neutrons how many electrons does it have? What would the mass of this atom be? 3. What must all atoms of the same element have in common?
. List the relative mass, charge and location of all three subatomic particles? 5. How many protons, neutrons and electrons does an atom of Lithium-7 have? 6. What conclusions about the atom where discovered as a result of the Gold Foil Experiment?
. List the relative mass, charge and location of all three subatomic particles? 5. How many protons, neutrons and electrons does an atom of Lithium-7 have? 6. What conclusions about the atom where discovered as a result of the Gold Foil Experiment?
Ions Defined as charged particles Ions are formed when the number of electrons changes. If a (+) ion is formed, electrons are lost (called cations). If a (-) ion is formed, electrons are gained (called anions).
Examples Ca 2+ A Ca atom has 20 protons and 20 electrons. A Ca 2+ ion has lost two electrons to have 18.
Examples Cl - A Cl atom has 17 protons and 17 electrons. A Cl - ion has gained one electron to have 18.
Practice Element Zn Fe 3+ F I - Li + Atomic # 26 9 53 3 Mass # 30 65 56 19 127 7 p n e 30 26 9 53 3 35 30 10 74 4 30 23 9 54 2
Isotopes Definition: Atoms that have the same atomic number (same # of protons) but a different mass number (different # of neutrons)
Isotopic Symbols Must write isotopic symbol to show mass 2 Isotopes will have the same atomic # (bottom) and a different mass # (top) Mass # Atomic # X
Write the isotopic symbol for: Carbon-14 (write a symbol for a different isotope of carbon) 14 C 6
Write the isotopic symbol for: Oxygen-17 (write a symbol for a different isotope of carbon) 17 O 8
Write the isotopic symbol for: Chlorine-37 (write a symbol for a different isotope of carbon) 37 Cl 17
3 Common Isotopes of Hydrogen Name Symbol #p #e #n Mass Protium 1 H 1 1 0 1 1 Deuterium 2 H 1 1 1 2 Tritium 3 H 1 1 2 3 1 1
Why is atomic mass not a whole number? The atomic mass on the periodic table is a weighted average of the isotopes of the elements. The weighted atomic mass takes into account the relative abundances of all the naturally occurring isotopes.
How do you calculate a weighted average? To calculate the weighted average you convert each percentage to a decimal by moving it 2 places left. Multiply the decimal by the mass for each isotope and add them all up. Or you can multiply the percent abundance (without moving the decimal) by the mass for each isotope add them all up and divide by 100 See Examples
Example of a general weighted average Your grade in chemistry 70% exams 85 10% quizzes 100 10% labs 95 10% HW/CW 80 (0.70)85 + (0.10)100 + (0.10)95 + (0.10)80 = 87
Example 1: Determine weighted atomic mass Boron-10 19.78% 10.013 amu Boron-11 80.22% 11.009 amu (0.1978) 10.013 + (.8022) 11.009 = 10.812 amu
Example 2 Determine weighted atomic mass Potassium-39 93.12% 38.964 amu Potassium-41 6.88% 40.962 amu (0.9312) 38.964 + (0.0688) 40.962 = 39.101 amu
Do Now 1. How many total electrons does an Al +3 ion have? 2. If a neutral atom has 10 neutrons and 8 electrons how many protons does it have? 3. How does an atom of Lithium-7 differ from an atom of Beryllium-9 4. Compare a Na +1 ion to a Na atom? 5. How are 14 N, 15 N and 16 N different and the same?
Bohr models How do electrons orbit the nucleus? Each principal energy level is a fixed distance from the nucleus can hold a specific number of electrons has a definite amount of energy
The greater the distance from the nucleus the greater the energy of the electrons in it. The orbits are called principal energy levels or shells.
Increasing distance from nucleus Increasing energy Energy levels or shells energy level number of e- 1 2 2 8 3 18 4 32
Bohr models: examples nucleus--- # protons And neutrons P+ n 0 2 e- 8 e- -energy levels and total number of electrons Electron configuration: element s symbol and number of electrons in each orbit; LOOK UNDER ATOMIC NUMBER ON PT of E
TRY THESE Mg 12 p+ 12 n 0 2 e- 8 e- 2 e- Electron configuration (bottom left corner on PT): Mg 2-8-2 H Na F C
answers H 1p+ 1 e- Na H 1 Na 2-8-1 11 p+ 12 n 0 2 e- 8e-1e- F 9 p+ 10 n 0 2 e- 7 e- F 2-7 C 2-4 C 6 p+ 6 n 0 2 e- 4 e-
Lewis Dot Diagrams Valence shell: outer most shell of an atom that contains electrons Valence electrons: electrons that occupy the valence shell (last number in electron configuration) Electron dot diagrams or Lewis dot diagrams: show only the valence shell of the atom Ex: Lewis dot for nitrogen: N
TRY THESE O F C Ne I K
Ions For ions: remember that ions have gained or lost electrons. Use periodic table to find charge of ion (see table) For dot diagrams of Ions (+)ions indicate charge no dots around the symbol (-)ions use brackets, charge, and always 8 dots around the symbol
Dot Diagrams for Ions Ca Ca +2 Cl [ Cl ] -1
Ground State vs. Excited State When all electrons in an atom occupy the lowest available orbitals, it is said to be in the ground state. When electron(s) absorb energy, they have the ability to jump to higher energy levels. The excited state is when electrons have absorbed energy and no longer occupy the lowest available energy levels.
Possible Excited States Na (ground state) Na (possible excited state) Na (another possible excited state) Na (another possible excited state) 2-8-1 2-7-2 2-6-3 2-5-4
Absorption When an electron jumps to a higher energy level it absorbs energy. The excited state is a temporary state. Excited State e- (i.e. energy level 2) Ground State (i.e. Energy level 1)
Emission The electron then falls back down to the ground state, emitting energy. The energy is in the form of light.
This radiant energy has a characteristic color and wavelength that can be determined. Every electron transition produces a specific wavelength of light and all transitions for an element blend together. This light can be separated through a prism into its various wavelength components. Every element has its own unique bright line spectrum that can be used to help identify the presence of that element. Ex: elements in a star, forensic analysis, flame tests, spectroscopy
Light and Atomic Spectra (bright line spectra) Electromagnetic spectrum consists of light that exists as waves.
Sunlight and prisms Sunlight produces a continuous range of wavelengths and frequencies that can be separated into all the colors of the rainbow. R O Y G B I V.
Atomic emission spectra produce narrow lines of color called bright line spectra. Each line corresponds to an exact wavelength.
Experiments Flame Tests Flame Tests demonstrates the emission spectrum of a substance. Completed by heating elements to high temperatures so they may enter excited state. Characteristic color will be emitted as excited electrons return to ground state. Used to determine metal ion presence in unknown substance.
Experiments Spectroscopy Spectroscopy used to view the bright line spectra for given gases. Completed by viewing a gas tube through which an electric current is passed. Use an instrument called a spectroscope, contains a prism to separate emitted light into line spectra.