Periodic Trends

Properties of the elements exhibit trends in the periodic table. These properties include: atomic radius (or size), ionization energy, electron affinity, electronegativity, metallic and nonmetallic property.

ATOMIC SIZE

Atomic size refers to how big or how small an atom is. The size of an atom is determined by the distance of the valence electron (the outermost electron) from the nucleus of the atom. The shorter this distance is, the smaller is the atomic size. This distance is further dependent on other factors such as

a. the number of energy levels in an atom (the more is the number of energy levels, the bigger is the atom);

b. shielding effect (refers to the number of inner electrons that "shields" the attraction between the nucleus and the valence electrons-- the higher the shielding effect, the less is the attraction, therefore the bigger will be the size of the atom);

c. effective nuclear charge (the number of protons found inside the nucleus of the atom). The higher is the nuclear charge, the greater is the ability of the nucleus to pull electrons towards itself thus, decreasing the size of the atom.

The following "periodic table" shows a graphical representation of the sizes of the atoms down a group and across a period.

The absence of representations for those of the metals indicates that there is difficulty in establishing a clear trend for atomic size of these transition elements. Atomic size of the noble gases, Group 8A, are estimates from those of nearby elements.

IONIZATION ENERGY

When an atom loses an electron, it becomes a positive ion (cation). The removal of an electron from an atom requires an absorption of ionization energy. Ionization energy is the energy required to "knock-off" the most loosely held electron (the valence electron) from an atom.

The more tightly held is the electron to the nucleus (atomic size is small), the higher is the energy required to remove it. Thus a smaller atom will require higher ionization energy than a bigger one.

The following graph shows the ionization energies of the first 20 elements.

A more comprehensive survey of the ionization energies of the main group elements is shown below.

ELECTRON AFFINITY

Electron affinity is the energy released when an electron is added to a neutral atom to form a negative ion. If the electron affinity is low, the electron is weakly bound; if the electron affinity is high, the electron is strongly bonded. Generally, electron affinity increases from left to right across the period because of increase in nuclear charge and decrease in atomic size. This causes the incoming electron to experience a greater pull of the nucleus thus giving a higher electron affinity.

Electron affinity decreases down a group because of increasing number of energy levels (and atomic size). This causes the incoming electron not to experience much attraction to the nucleus thus giving a lower electron affinity.

The electron affinities of completely filled atoms (Group 8A) is almost zero. An atom does not accept an electron in its outermost energy level if it already has a stable configuration.

ELECTRONEGATIVITY

When two elements are joined in a chemical bond, the element that attracts the shared electrons more strongly is more electronegative (usually nonmetallic elements). Elements with low electronegativities (the metallic elements) are said to be electropositive. It is important to note that electronegativities are properties of atoms that are chemically bound to each other; there is no way of measuring the electronegativity of an isolated atom.

Electronegativity values increases across a period in the periodic table. This is because elements are becoming more nonmetallic (atomic size is decreasing) from left to right. Electronegativity decreases down a group because of increasing atomic size. Elements found below are more metallic thus they are more electropositive than those found above.

METALLIC and NONMETALLIC PROPERTIES

Metals are very good electron donors. In terms of atomic size, the bigger an atom is, the more metallic it is. Thus, across a period in the periodic table metallic property decreases. Down a group, where atomic size increases, metallic property also increases.

Nonmetals, on the other hand, are very good electron acceptor. The smaller is the size of the atom the easier for it to attract (or receive) an electron from an external source. Therefore across a period in the periodic table, nonmetallic property increases. Down a group, nonmetallic property decreases.

24-Item Test on Atomic Structure and the Periodic Table

This test covers the topics on atoms and the periodic table. Before starting the test, you will be asked to enter your name (please include your course, year and section, i.e. LEMUEL OWEN H. CABANA BSN1A). A quiz with no name will be scored zero. You will also need a periodic table and a calculator in answering most of the questions in the test.

Kindly print the score report ("certificate") found at the end of the quiz. Submit such to me on our next lecture class meeting.

Good luck!

link: http://www.proprofs.com/quiz-school/story.php?title=atomic-structure--periodic-table

Chemistry: A Volatile History (a BBC Four Documentary)

As a Chemistry teacher, watching this narrative feels amazingly like watching a big-budget Hollywood film :) This superbly pieced-together documentary and Chemistry lesson in one show beautifully explains the progression of the human understanding of the elements. (If only my chemistry lessons were this good at school, all my students would have become scientists! LOL)

This three-part documentary, produced by BBC Four, features Jim Al-Khalili, a professor of Theoretical Physics, tracing the extraordinary story of how the elements were discovered and mapped. In the process, he uncovers tales of success and heartaches in the story of well-known chemists' battle to control the elements which, in many ways, have helped build our modern world.

Episode 1: Discovering the Elements

Episode 2: The Order of the Elements

Episode 3: The Power of the Elements

Episode 2 is particularly interesting for me since I am presently discussing the Periodic Table of the Elements in my classes.

Many thanks to AtheistMedia.com for uploading the complete set of the documentary on youtube.

Episode 2: The Order of the Elements Part 1

Episode 2: The Order of the Elements Part 2

Episode 2: The Order of the Elements Part 3

Episode 2: The Order of the Elements Part 4

Episode 2: The Order of the Elements Part 5

Episode 2: The Order of the Elements Part 6

Rutherford's Gold-Foil Experiment: The Nuclear Model of the Atom

In 1909, Hans Geiger and Ernest Marsden, two researchers in Ernest Rutherford's laboratory at the University of Manchester, fired a beam of alpha particles at a thin gold metal foil. An alphaparticle, identified and named a decade earlier by Rutherford, is one of the types of radiation given off by radioactive elements such as uranium. Because these particles are fast-moving and positively charged (they're now known to be high-speed Helium nuclei), Rutherford reasoned they will serve as a good probe of the atomic structure of matter.

The experiment was done to validate the prevailing atomic model then, the plum-pudding model (some refer to it as the raisin-bread model) championed by J.J. Thomson. According to Rutherford, if Thomson's atomic model is correct, each high-velocity alpha particle will just pass straight through the thin gold foil suffering, at most, minor deflections because of the weak influence of the spread-out ball of positive charge and the effect of the electrons being negligible.

What Geiger and Marsden actually observed was stunning. Most of the alpha particles did indeed traveled straight through the foil with little or no deviation. But a small fraction (about 1 in 10,000) rebounded, ending up on the same side of the foil as the incoming beam. A few were returned almost along the same tracks as they went in. Rutherford upon hearing of these rebounds, described it as the most incredible event of his life. It was, he said, "as if you fired a 15-inch shell at a piece of tissue paper and it came back and hit you." Such huge deflections could mean only one thing: some of the alpha particles had run into massive concentrations of positive charge and, since like charges repel, had been hurled straight back by them. Thomson's plum-pudding model of the atom failed to explain these observations.

After months of studying, in 1911, Rutherford, presented a new model of the atom in which all of the positive charge is crammed inside a tiny, massive nucleus about ten thousand times smaller than the atom as a whole. That's equivalent to a marble in the middle of a football stadium. The much lighter electrons, he assumed, are positioned outside the nucleus. Everyone was amazed at the implication of Rutherford’s proposal---the atoms of which planets, people, objects, and everything else are made consisted almost entirely of empty space.

Rutherford's nuclear model of the atom was considered as one giant step forward in understanding nature at the microscopic level. But even as it closed the issue on the alpha particle experiment, it threw open another one. Since the nucleus and its retinue of electrons are oppositely charged, and therefore attract one another, there didn't seem anything to stop the electrons from being pulled immediately into the nucleus. Rutherford countered by saying that the atom was like a miniature solar system: the electrons circled the nucleus in wide orbits just as planets orbit the sun. This is the picture of atoms that most of us still carry around in our heads. It's an appealing, easy-to-grasp image – one that's inspired many a logo of the atomic age. Yet theorists were well aware of its shortcomings right from the start.

The fatal flaw in Rutherford's model is that it contains charges that are accelerating. The charges are on the electrons and the acceleration is due to the electrons always changing direction as they move around their orbits. (Things accelerate when they change speed and/or direction.) Since Maxwell's time, scientists had known that accelerating charges radiate energy. What was to stop the orbiting electrons in Rutherford's atom quickly (in fact, in about one hundred-millionth of a second) losing all their energy and spiraling into the nucleus?

The answer came from a young Dane, Niels Bohr, who joined the team at Manchester for a six-month spell in 1912, shortly after Rutherford went public with his new vision of the atom. Bohr played a hunch. He knew about Planck's quantum. He knew there was no way to save an electron inside an atom from plummeting into the nucleus if it could give off energy continuously. And so he said simply that electrons inside atoms can’t radiate continuously. They can only radiate in lumps, and these lumps are the same as Planck's quanta. For a given type of atom, say hydrogen, there's a limited number of stable orbits that an electron can occupy. Each of these orbits corresponds to a whole multiple of the basic quantum. As long as an electron is in one of these orbits, its energy, contrary to whatever classical physics might say, stays the same. If it jumps from an outer (higher energy) orbit to an inner (lower energy) orbit, the energy difference between the two is given off as a quantum of light. Once the electron reaches the lowest energy orbit, it can't fall any further and is safe from the clutches of the nucleus.

GenChem Midterm Lab Activities

7 is here...

8 is here...

9 is here...

you know what to do... :)