Atomic radius trends on periodic table. Ionization energy trends. Ionization energy: period trend. First and second ionization energy. Electron affinity: period trend. Electronegativity and bonding. Periodic trends and Coulomb's law. Worked example: Identifying an element from successive ionization energies.
Ionization energy: group trend. Current timeTotal duration Google Classroom Facebook Twitter. Video transcript Voiceover: Let's think a little bit about the notion of atomic size or atomic radius in this video.
At first thought, you might think well this might be a fairly straight-forward thing. If I'm trying to calculate the radius of some type of circular object I'm just thinking about well what's the distance between the center of that circular object and the edge of it. So the length of this line right over here.
That would be the radius. And so a lot of people when they conceptualize an atom they imagine a positive nucleus with the protons in the center right over here then they imagine the electrons on these fixed orbits around that nucleus so they might imagine some electrons in this orbit right over here, just kind of orbiting around and then there might be a few more on this orbit out here orbiting around, orbiting around out here.
And you might say, "well okay, that's easy to figure out the atomic radius. I just figure out the distance between the nucleus and the outermost electron and we could call that the radius. Electrons are not in orbits the way that planets are in orbit around the sun and we've talked about this in previous videos. They are in orbitals which are really just probability distributions of where the electrons can be, but they're not that well defined.
So, you might have an orbital, and I'm just showing you in 2 dimensions. It would actually be in 3 dimensions, where maybe there's a high probability that the electrons where I'm drawing it in kind of this more shaded in green. But there's some probability that the electrons are in this area right over here and some probability that the electrons are in this area over here, and let's say even a lower probability that the electrons are over this, like this over here.
And so you might say, well at a moment the electron's there. The outermost electron we'd say is there. You might say well that's the radius. But in the next moment, there's some probability it might be likely that it ends up here.
Explanation: The following diagram shows the periodic trends of atomic radius of the representative elements main group elements for the first six periods.
Related questions Why do periodic trends exist for electronegativity? Why does atomic size increase down a group? What do periodic trends of reactivity occur with the halogens? How can I determine atomic size of ions? The atomic radii of transition metals do not decrease significantly across a row.
Atomic radii are measured in picometers one picometer is equal to one trillionth of a meter. Hydrogen H has the smallest average atomic radius at about 25 pm, while caesium Cs has the largest average radius at about pm. There are two main atomic radius trends.
One atomic radius trend occurs as you move left to right across the periodic table moving within a period , and the other trend occurs when you move from the top of the periodic table down moving within a group. Below is a periodic table with arrows showing how atomic radii change to help you understand and visualize each atomic radius trend. At the end of this section is a chart with the estimated empirical atomic radius for each element.
The first atomic radius periodic trend is that atomic size decreases as you move left to right across a period. Within a period of elements, each new electron is added to the same shell. When an electron is added, a new proton is also added to the nucleus, which gives the nucleus a stronger positive charge and a greater nuclear attraction. Comparing carbon C with an atomic number of 6 and fluorine F with an atomic number of 9, we can tell that, based on atomic radius trends, a carbon atom will have a larger radius than a fluorine atom since the three additional protons the fluorine has will pull its electrons closer to the nucleus and shrink the fluorine's radius.
The second atomic radius periodic trend is that atomic radii increase as you move downwards in a group in the periodic table. For each group you move down, the atom gets an additional electron shell.
Each new shell is further away from the nucleus of the atom, which increases the atomic radius. While you may think the valence electrons those in the outermost shell would be attracted to the nucleus, electron shielding prevents that from happening.
Electron shielding refers to a decreased attraction between outer electrons and the nucleus of an atom whenever the atom has more than one electron shell. As an example, potassium K has a larger average atomic radius pm than sodium Na does pm. The potassium atom has an extra electron shell compared to the sodium atom, which means its valence electrons are further from the nucleus, giving potassium a larger atomic radius. The two atomic radius trends we discussed above are true for the majority of the periodic table of elements.
However, there are a few exceptions to these trends. One exception is the noble gases. Because electronegativity is a qualitative property, there is no standardized method for calculating electronegativity. However, the most common scale for quantifying electronegativity is the Pauling scale Table A2 , named after the chemist Linus Pauling. The numbers assigned by the Pauling scale are dimensionless due to the qualitative nature of electronegativity.
Electronegativity values for each element can be found on certain periodic tables. An example is provided below. Electronegativity measures an atom's tendency to attract and form bonds with electrons. This property exists due to the electronic configuration of atoms. Most atoms follow the octet rule having the valence, or outer, shell comprise of 8 electrons. Because elements on the left side of the periodic table have less than a half-full valence shell, the energy required to gain electrons is significantly higher compared with the energy required to lose electrons.
As a result, the elements on the left side of the periodic table generally lose electrons when forming bonds. Conversely, elements on the right side of the periodic table are more energy-efficient in gaining electrons to create a complete valence shell of 8 electrons. The nature of electronegativity is effectively described thus: the more inclined an atom is to gain electrons, the more likely that atom will pull electrons toward itself.
According to these two general trends, the most electronegative element is fluorine , with 3. Ionization energy is the energy required to remove an electron from a neutral atom in its gaseous phase. Conceptually, ionization energy is the opposite of electronegativity. The lower this energy is, the more readily the atom becomes a cation. Therefore, the higher this energy is, the more unlikely it is the atom becomes a cation. Generally, elements on the right side of the periodic table have a higher ionization energy because their valence shell is nearly filled.
Elements on the left side of the periodic table have low ionization energies because of their willingness to lose electrons and become cations. Thus, ionization energy increases from left to right on the periodic table. Another factor that affects ionization energy is electron shielding. Electron shielding describes the ability of an atom's inner electrons to shield its positively-charged nucleus from its valence electrons. When moving to the right of a period, the number of electrons increases and the strength of shielding increases.
As a result, it is easier for valence shell electrons to ionize, and thus the ionization energy decreases down a group. Electron shielding is also known as screening. Some elements have several ionization energies; these varying energies are referred to as the first ionization energy, the second ionization energy, third ionization energy, etc.
The first ionization energy is the energy requiredto remove the outermost, or highest, energy electron, the second ionization energy is the energy required to remove any subsequent high-energy electron from a gaseous cation, etc. Below are the chemical equations describing the first and second ionization energies:. Generally, any subsequent ionization energies 2nd, 3rd, etc. Ionization energies decrease as atomic radii increase. The relationship is given by the following equation:.
As the name suggests, electron affinity is the ability of an atom to accept an electron. Unlike electronegativity, electron affinity is a quantitative measurement of the energy change that occurs when an electron is added to a neutral gas atom. The more negative the electron affinity value, the higher an atom's affinity for electrons.
Electron affinity generally decreases down a group of elements because each atom is larger than the atom above it this is the atomic radius trend, discussed below. This means that an added electron is further away from the atom's nucleus compared with its position in the smaller atom.
With a larger distance between the negatively-charged electron and the positively-charged nucleus, the force of attraction is relatively weaker. Therefore, electron affinity decreases. Moving from left to right across a period, atoms become smaller as the forces of attraction become stronger.
This causes the electron to move closer to the nucleus, thus increasing the electron affinity from left to right across a period. The atomic radius is one-half the distance between the nuclei of two atoms just like a radius is half the diameter of a circle.
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