
Dmitri Mendeleev
Chemistry Review 3
History of the Periodic Table of the Elements 10/10/03
Greek Thinkers- 400BC
Used words for atoms and elements
1. Fire
2. Earth
3. Water
4. Air
Johann Dobereiner- 1817
Arranged elements in groups of three (triads)
1. Cl, Br & I
2. Li, Na & K
John Newlands- 1864
Octaves- properties repeat every 8th element
Lother Meyer- 1864
Organized atoms by atomic weight
Published abbreviated periodic table. 1868 constructed an extended version.
Dmitri Mendeleev- 1869
Arranged 63 known elements in increasing atomic mass
Showed gaps where elements would be found
Predicted physical and chemical traits of unknown elements
William Ramsey- 1898
Placed argon between chlorine and potassium
Created the zero/18 group; inert/noble gases
Henry Mosely- 1913
Rearranged table by atomic number, not mass
First modern form of periodic table
Specialized in x-rays and radiation. Died in 1915 in a battle
Glen Seaborg- 1940
Discovered all the transuranic elements (numbers 90- 102)
Placed the actinide series below the lanthanide
10/27 Periodic Trends
Atomic Radii- measured by ½ the distance between two nuclei
A Atomic Radii Trends
1. Increases going down a group
2. Energy levels increase, more electrons
3. Decreases going across a period
4. Adding 1p+ and 1e-
5. Shielding affect
B Ionization Energy- Energy required to remove an e- from a gaseous atom
C Ionization Trends
1. Decreases down a group
2. Electrons are held with less force in outer energy levels
3. Increase moving left to right across a period
4. Electrons are held with more force
D Electron Affinity
1. The energy that accompanies the addition of an electron
a. Cl(g) + e- à Cl-(g)
2. Becomes more positive down a group
3. The electron is added at increasing distances from the nucleus
4. More negative going across a period
5. There are some exceptions to this trend
6. A positive number is endothermic
7. A negative number is exothermic
E Electronegativity
1. Linus Pauling- 2 nobels
2. “The ability of an atom in a molecule to attract shared electrons to itself”
Book notes Section 4-3
- Groups are vertical, and indicate a change in energy levels
- Periods are horizontal
Trend- predictable change in a particular direction
Bond Radius- ½ the distance between nuclei of atoms in each molecule- known with great precision, will be used in class. 2/3 van der Waal’s radius.
Van der Waal’s radius- not precise; measures radii of atoms in separate molecules?
- Atomic radius increases as you move down a group. This is because there are more energy levels around the nucleus
Electron shielding- electrons closer to the nucleus stop the outside electrons from receiving the same positive charge of attraction. Therefore, outer electrons don’t experience the same force of attraction.
- The nuclear charge increases as you move down a group (more protons, more charge), but the charge acting on the outer electrons is constant (more electrons, more electron shielding, it all evens out)
- Atomic radius decreases as you move across a period. This is because the nuclear charge increases, and electron shielding is constant (because no additional energy levels are added) so the electrons are sucked closer to the nucleus
Ionization energy- energy used to remove electrons. When there is only one electron in the outer energy level (group 1 on the periodic table), the removal of an electron erases a whole energy level, making the ion smaller than the neutral atom.
- Energy needed to ionize increases across a period and decreases down a group. It is the opposite of atomic size. This is because a smaller atom indicates a higher attraction between the nucleus and the electrons. Obviously, it takes more energy to break up a stronger attraction.
Electron affinity- the ability of an atom to hold an electron. Neutral atoms can attract electrons if their outer energy levels aren’t full. With a full outer energy level, the nucleus is well shielded from outside electrons (hence the reluctance of the noble gases to react). If not, an electron can still feel a positive pull and jumps into an empty orbital.
- Energy is usually given off when an atom attracts an electron.
- Electron affinity is more difficult to measure than ionization energy
- Electron affinity becomes more negative across a period and decreases from top to bottom in a group. Electron affinity follows ionization in its trends on the periodic table. This is because the ability to hold an electron directly corresponds with the energy needed to take that electron away.
- Electron affinity trends aren’t as regular as those for ionization energy; there are some exceptions
Melting and boiling points
- The highest melting and boiling points occur when an orbital is half full. This is due to stronger bonding between the individual atoms of the element (remember the electron- sharing bonds we learned about in physical science? When the outer energy levels are about half full, the atom is sharing quite a few of its electrons with the other atoms around it. When the outer energy level is almost full or almost empty, only a few electrons are being shared. It’s like the difference between a hug and a handshake). It takes more energy (higher temperatures) to break up the stronger bonds.
- This trend is more visible in the higher energy levels. It’s hard to see a trend in period 1, for example, which only contains hydrogen and helium.
Section 4-4
Natural and synthetic elements
- There are 93 naturally occurring elements. Three (technetium, promethium, and neptunium) aren’t found on Earth, but have been detected in the light spectra of stars (remember that lab with the defraction grading lens?). Chemists have synthesized (made) other elements.
- Naturally occurring elements are created in the center of stars.
- Big Bang theory- 12-16 billion years ago, a particle that could’ve fit on the head of a pin exploded, creating tiny particles (hydrogen atoms)(Ahem, God, Ahem, Ahem)
Nuclear reaction- a reaction involving protons and neutrons in the nuclei of atoms.
- Nuclear reactions caused hydrogen atoms to fuse, creating helium.
- Nuclear fusion reactions only occur in extreme heat (centers of stars) where the kinetic energy is great enough to overcome the repulsion between positively charged particles (protons).
- A tiny bit of mass is lost (converted into energy) during these fusions.
- Einstein’s E=mc squared describes mass- to energy conversion quantitatively (with numbers). C is the speed of light.
- As hydrogen is used up, the star shrinks and the temperature rises (due to the energy that came from the fusions)
- Under these even hotter temperatures, hydrogen can fuse into other elements with higher atomic number, etc.
- Iron and nickel are the most massive atoms formed by this process.
- When a star uses all its H and He, it collapses and expands as a red giant, tearing some atoms apart. These atoms are absorbed by other nuclei, creating elements with atomic numbers higher than Fe and Ni.
- Super-hot red giants expel some matter into space
- Supernovas explode, and their matter creates a new 2nd generation of stars. These already have some heavy metals along with Hydrogen and Helium. The sun is a 4th or 5th generation star.
Artificial elements
Transmutation- a nuclear reaction that cannot be achieved by ordinary chemical means
- Alchemists obsessed over transmuting cheaper elements into gold
- Ernest Rutherford produced protons while studying the passage of high-speed alpha particles through water vapor. He found long thin tracks, which turned out to be protons (nuclei of H). Guessed that atomic nuclei in air were disintegrating upon being struck by helium nuclei. They turned into hydrogen nuclei plus the nuclei of some other atom.
- W.D. Harkins (US) and P.M.S. (he musta gotten beat up a lot in grade school, eh?) Blackett (UK) studied this further and confirmed Rutherford’s hypothesis.
Cyclotron- invented 1930 E.O. Lawrence (US). Increases energy of charged particles, enabling them to fuse creating much higher atomic numbers. Used to create synthetic elements.
Synchotron- can’t accelerate as many types of elements as a cyclotron, but can reach enormous energies. Used to create superheavy metals.
Superheavy elements- elements more massive than the actinides
- The US, Germany, Russia and Sweden have the largest research teams when it comes to superheavy metals.
- Scientists predict that they will be able to make atoms stable up to an atomic number of 114.
1. Master group information
2. Grouping
3. Metals, Nonmetals and Metalloids
4. Electron Configuration
5. Periodic Trends, Electronegativity, Ionization, atomic radii
6. History of the Periodic Table