Atomic Structure and the Periodic Table of the Elements Flashcards
Dalton’s five basic ideas about atoms
(helium nucleus 4/2 He) Positively charged, 2+
Ex. 230/90 Th –> 4/2 He + 226/88 Ra
- All matter is made up of very small, discrete particles called atoms 2. All atoms of an element are alike in weight, and this weight is different from that of any other kind of atom. 3. Atoms cannot be subdivided, created, or destroyed 4. Atoms of different elements combine in simple whole-number ratios to form chemical compounds. 5. In chemical reactions, atoms are combined, separated, or rearranged
Niels Bohr
Published a theory explaining the line spectrum of hydrogen. He proposed a PLANETARY MODEL that quantized the energy of electrons to specific orbits. Louis de Broglie and others showed that quantum theory described a more PROBABILISTIC MODEL of where the electrons could be found that resulted in the theory of orbitals.
J. J. Thomson
Discovered the electron and ratio of the electrical charge the electron to its mass Using a cathode ray - beam was deflected because it was magnetics and electrical
Robert Millikan
Experiment determined the mass of an electron - Oil drop experiment
Ernest Rutherford
Gold foil experiment using alpha particles confirmed that there was mostly empty space between the nucleus and electron
James Chadwick
Discovered the neutron
Bohr Model of the Atom
Electron distribution to principal energy levels has the formula: 2n^2 Principal Energy Level –> Max. Number of Electrons (2n^2) 1 –> 2 2 –> 8 3 –> 18 4 –> 32 5 –> 50
Nucleons
Proton and neutrons are in the nucleus. These particles are known as nucleons.
atomic number
The number of protons in the nucleus
Henry Moseley
First determined the atomic number of the elements through the use of X-ray
Mass number
sum of protons and neutrons in the nucleus
Average atomic mass
The weighted average of the atomic masses of the naturally occurring isotopes of an element
Calculating Average Atomic Mass
69.17% copper-63 –> atomic mass of 62.919 30.83% copper-65 –> atomic mass of 64.93 (0.6917 x 62.919) + (0.3083 x 64.93) = ANSWER
Valence electrons
The electrons found on the outer-most energy level
Lewis Structure
A Lewis strucutre shows the atomic symbol to represent the nucleus and inner shell electrons. It shows dots to represent the valence electrons.
Change in energy (DELTA E)
DELTA E = Efinal - Einitial
Photons
discrete radiant energy
Atomic Spectra Chart
Lyman series: When an electron cascades from a level higher than the first level down to n = 1 (ultraviolet range) Balmer series: n = 2 (visible range) Paschen series: n = 3
Spectra
(spectroscope) distinct color line
Frequency
= velocity of light/ wavelength
Mass spectroscopy
Mass spectroscope separates isotopes of the same element based on differences in their mass
Max Planck
Quantum theory of light that light has both particle like properties and wavelike characteristics.
Louis de Broglie
Particles can also have wavelike characteristics
Uncertainty principle
Werner Heisenber It is impossible to know both the precise location and precise velocity of a subatomic particle at the same time
Wave Mechanism Modle
Orbitals - the orbital is a three-dimensional region around the nucleus that indicates the probable location of an electron but gives no information about its pathway (NOT RELATED to the Bohr orbits at all)
Principal quantum number (n)
1, 2, 3, 4, 5, etc. The values of n = 1, 2, 3… Average distance of the orbital from the nucleus. (principle energy level)
Angular momentum (l) quantum number
s, p, d, f (in order of increasing energy) The value of l can be = 0, 1, … n-1 l = 0: spherical shaped s orbital l = 1: dumbbell-shaped p orbital l = 2: five orbital orientation d orbital This number refers to the shape of the orbital Limited by the principal quantum number
Magnetic quantum number (ml)
Number of spatial orientations of orbitals s = 1 space-oriented orbital p = 3 space-oreinted orbitals d =5 space-oriented orbitals f = 7 space-oriented orbitals Can equal: -l…. 0 …. l
Spin quantum number (ms)
+ spin - sin The value of m = +1/2 or -1/2
Pauli Exclusion Principle:
Now two electrons can have the same four quantum numbers
p orbitals
have a dumbbell shape, oriented on the x, y, and z-axes, and can hold a total of 2 electrons each, making a total of 6
s orbitals
spherical and can hold 2 electrons
d orbitals
Have 5 orientations and can hold 2 electrons each, making a total of 10
Limits of quantum numbers
n = 1, 2, 3,… l = 0, 1, … (n-1) ml = -l, … 0 … +l
Aufbau Principle
An electron occupies the lowest energy orbital that can receive it. Then they will pair up when they are all filled.
Hund’s Rule of Maximum Multiplicity
An electron occupies the lowest energy orbital that can receive it. –> Then they will pair up when they are all filled.
Transitional element characteristic properties
They often form color compounds They can have a variety of oxidation states At least one of their compounds has an incompleted electron subshell They are often good catalysts They are silvery blue and room temperature (except for copper and gold) They are solids at room temp. (except mercury) They form complex ions They are often paramagnetic
Dimitri I. Mendeleev
is given credit for the first periodic table. It was based on placement by properties.
Henry Moseley
Changed the basis of the periodic law from atomic weight to atomic number
Periodic Table Trends
Acid-forming properties increase from left to right on the table Base-froming properties are high on the left side and decrease to the right The atomic radii of elements decrease from left to right across a period Metallic properties are greatest on the left side of the table and decrease to the right Nonmetallic properties are greatest on the right side of the table and decrease to the left
Atomic and Ionic Radii
Atomic radii decrease from left to right across a period in the Periodic Table (until the noble gases) Tomic radii increases from top to bottom in a group or family Atomic vs. Ionic: positive vs. negative charges - negative means ionic radius is larger - positive means ionic radius is smaller
Electronegativity
is a number that measures the relative strength with which the atoms of the element attract valence electrons in a chemical bond. Value less than 2 = metal Decreases down a group and increases across a period Lower the electronegativity the more electropositive an element is. Most electronegative = F Least = Fr (francium)
First ionization Energy
The energy that is supplied to one outer electron to remove it from its atom Lowest ionization elements are found with the least electronegative elements
Alpha Particle
(helum nucleus 4/2He) Positively charged, 2+
Ex. 230/90Th –> 4/2He + 226/88Ra
- Ejection reduces the atomic number by 2. the atomic weight by 4 amu.
- High energy, relative velocity
- Range: about 5 cm in air
- Shielding needed: stopped by the thickness of a sheet of paper, skin
- Interactions: produces about 100,000 ionizations per centimeter; repelled by the positivelely charged nucleus; attracts electrons, but does not capture them until its speed is much reduced. 6.
Beta Particle
(fast electron) Negatively charged, 1-
Ex. 234/91 Pa –> 234/91 U + 0/-1 e
- Ejected when a neutron decays into a proton and an electron
- High velocity, low energy
- Range: about 12 m
- Shielding needed: stopped by 1 cm of aluminum or thickness of average book
- Interactions: weak because of high velocity, but produces about 100 ionizations per centimeter.
Gamma Radiation
(electromagnetic radiation identical with light; high energy) No charge
- Beta particles and gamma rays are usually emitted together; after a beta is emitted, a gamma ray follows
- Arrangement in nculeus is unknown. Same velocity as visible light.
- Range: no specific range
- Shielding needed about 13 cm of lead
- Interations: weak of itself; gives energy to electrons, which then perform the ionization.
Methods of detection of alfpha, beta, and gamma rays
- Photographic plate
- Scintillation counter
- Greiger counter
Half-life
Radon
Transmutation
conversion of an element to a new element
