Urgent Chemistry Help-Trends

I need quick help
What trends are present across periods and down groups on the periodic table and why for the following:

-Electrical Conductivity
-Combining Power/Valency
-Melting Point
-Boiling Point

I really need help, I've looked all over google but couldn't find anything. i'll be a very happy person if you can help me with this, Needed by Tuesday!

Well.

That's odd. What grade are you in? I'm only in Gr. 10, so I may or may not be able to help.

Really the only thing I know is about the valences.
if you look at hydrogen, you will notice that it has one electron it's outer orbit. The same goes for lithium, sodium, potassium, and all the others in the first group.

The next group has two as a valence, then the third group has three (not counting transition metals I think), then next four, and so on, and so on. Helium is placed in the last group because it has two in the ONLY orbit, the max for that orbit, which makes it stable, like the rest in that group, which have 8.

I wish I could help you with the other stuff, but I don't know PM me if you're confused about anything

I can't remember anything about conductivity, so I'll just help you for the other 3.

Valence electrons increase from left to right, and top to bottom. This is basically the defining feature of the atom (unless it's an ion).

The melting point and boiling point go together. Increase the boiling point, and you generally decrease the freezing point. Think about this one - when you're boiling something, you're ripping the molecules away from each other. Bigger atoms (and bigger molecules) are much harder to break apart than smaller ones. Basically the bigger the atom (meaning the lower ones, or the ones toward the right) the higher the boiling point, and the lower the freezing point. In terms of the boiling/freezing point of molecules, it's slightly different, but I don't think that's what you're being taught yet.

Another trend you might want to consider is electronegativity. This basically determines the strength of a bond. Generally, the elements on the top right are the most electronegative (but I'm not sure about the noble gasses on the far right) and the ones at the bottom left are the least. Because of this, you can get a strong molecule from binding an atom from the bottom left with one on the top right - NaCl, salt, is a good example.

Hope that helps. I'm in college, but chemistry is my weakness...

Originally Posted by Jazz Man

That's odd. What grade are you in? I'm only in Gr. 10, so I may or may not be able to help.

11 Prelim.

We had to find the trends and explain why of the following:
-Electrical Conductivity
-Atomic Radius
-Ionisation Energy
-Melting Point
-Boiling Point
-Combining Power/Valency
-Elctronegativity
-Reactivity

The melting point and boiling point go together. Increase the boiling point, and you generally decrease the freezing point. Think about this one - when you're boiling something, you're ripping the molecules away from each other. Bigger atoms (and bigger molecules) are much harder to break apart than smaller ones. Basically the bigger the atom (meaning the lower ones, or the ones toward the right) the higher the boiling point, and the lower the freezing point. In terms of the boiling/freezing point of molecules, it's slightly different, but I don't think that's what you're being taught yet.

Yer I get that but do you know the trends of them going acroos the period and down a group.

Wow, I can't believe I've had this stuff already. I'm grade 9! I just wish I had my notes back....thankfully I have my book though.

What Mad Hatter is saying is that as you go right across a period, the boiling point raises and the freezing point decreases. Going down a family (group, whatever) is the same.

Going right across a period (not counting the noble gases or transition metals) you get more and more reactive, so they combine more easily. As you go down a column, I'm not so sure. Since all families have the same number of valence electrons, I don't know if they combine any easier. I'd say it stays the same, but what do I know?

Electrical conductivity is easier than you think. Metals conduct electricity, and nonmetals don't. So as you go to the right in a period (metals to nonmetals) the elements are less conductive. Going down a group, I think they change but I'm not sure how. If I remember correctly (and I might not) I think the atom is less conductive as you go down, too. Not sure on that one, you can reason it out, though.

Hope some of that helped.

A visual aid for trends in atomic radii:

at the moment all i can give you is that as you go down a group, the atomic radii gets larger as an extra 'shell' of electrons is added.

i got a whole stash of notes on my AS Chem, from last year, ill put some up if ya need it

Thanks but HR the answer you gave me is will get me 0 or 1 mark, but you did help none the less . As for atomic radius I don't need that information considering I didn't ask that at the start of the thread, but thanks for your concern.

The properties of the elements exhibit trends. These trends can be predicted using the periodic table and can be explained and understood by analyzing the electron configurations of the elements. Elements tend to gain or lose valence electrons to achieve stable octet formation. Stable octets are seen in the inert gases, or noble gases, of Group VII of the periodic table. In addition to this activity, there are two other important trends. First, electrons are added one at a time moving from left to right across a period. As this happens, the electrons of the outermost shell experience increasingly strong nuclear attraction, so the electrons become closer to the nucleus and more tightly bound to it. Second, moving down a column in the periodic table, the outermost electrons become less tightly bound to the nucleus. This happens because the number of filled principal energy levels (which shield the outermost electrons from attraction to the nucleus) increases downward within each group. These trends explain the periodicity observed in the elemental properties of atomic radius, ionization energy, electron affinity, and electronegativity.

Atomic Radius

The atomic radius of an element is half of the distance between the centers of two atoms of that element that are just touching each other. Generally, the atomic radius decreases across a period from left to right and increases down a given group. The atoms with the largest atomic radii are located in Group I and at the bottom of groups.

Moving from left to right across a period, electrons are added one at a time to the outer energy shell. Electrons within a shell cannot shield each other from the attraction to protons. Since the number of protons is also increasing, the effective nuclear charge increases across a period. This causes the atomic radius to decrease.

Moving down a group in the periodic table, the number of electrons and filled electron shells increases, but the number of valence electrons remains the same. The outermost electrons in a group are exposed to the same effective nuclear charge, but electrons are found farther from the nucleus as the number of filled energy shells increases. Therefore, the atomic radii increase.

Ionization Energy

The ionization energy, or ionization potential, is the energy required to completely remove an electron from a gaseous atom or ion. The closer and more tightly bound an electron is to the nucleus, the more difficult it will be to remove, and the higher its ionization energy will be. The first ionization energy is the energy required to remove one electron from the parent atom. The second ionization energy is the energy required to remove a second valence electron from the univalent ion to form the divalent ion, and so on. Successive ionization energies increase. The second ionization energy is always greater than the first ionization energy. Ionization energies increase moving from left to right across a period (decreasing atomic radius). Ionization energy decreases moving down a group (increasing atomic radius). Group I elements have low ionization energies because the loss of an electron forms a stable octet.

Electron Affinity

Electron affinity reflects the ability of an atom to accept an electron. It is the energy change that occurs when an electron is added to a gaseous atom. Atoms with stronger effective nuclear charge have greater electron affinity. Some generalizations can be made about the electron affinities of certain groups in the periodic table. The Group IIA elements, the alkaline earths, have low electron affinity values. These elements are relatively stable because they have filled s subshells. Group VIIA elements, the halogens, have high electron affinities because the addition of an electron to an atom results in a completely filled shell. Group VIII elements, noble gases, have electron affinities near zero, since each atom possesses a stable octet and will not accept an electron readily. Elements of other groups have low electron affinities.

Electronegativity

Electronegativity is a measure of the attraction of an atom for the electrons in a chemical bond. The higher the electronegativity of an atom, the greater its attraction for bonding electrons. Electronegativity is related to ionization energy. Electrons with low ionization energies have low electronegativities because their nuclei do not exert a strong attractive force on electrons. Elements with high ionization energies have high electronegativities due to the strong pull exerted on electrons by the nucleus. In a group, the electronegativity decreases as atomic number increases, as a result of increased distance between the valence electron and nucleus (greater atomic radius). An example of an electropositive (i.e., low electronegativity) element is cesium; an example of a highly electronegative element is fluorine.

Summary of Periodic Table Trends

Moving Left --> Right

* Atomic Radius Decreases
* Ionization Energy Increases
* Electronegativity Increases

Moving Top --> Bottom

* Atomic Radius Increases
* Ionization Energy Decreases
* Electronegativity Decreases

--> From About.com

~GoL

Originally Posted by GðÐ-Õƒ-Lîñk

The properties of the elements exhibit trends. These trends can be predicted using the periodic table and can be explained and understood by analyzing the electron configurations of the elements. Elements tend to gain or lose valence electrons to achieve stable octet formation. Stable octets are seen in the inert gases, or noble gases, of Group VII of the periodic table. In addition to this activity, there are two other important trends. First, electrons are added one at a time moving from left to right across a period. As this happens, the electrons of the outermost shell experience increasingly strong nuclear attraction, so the electrons become closer to the nucleus and more tightly bound to it. Second, moving down a column in the periodic table, the outermost electrons become less tightly bound to the nucleus. This happens because the number of filled principal energy levels (which shield the outermost electrons from attraction to the nucleus) increases downward within each group. These trends explain the periodicity observed in the elemental properties of atomic radius, ionization energy, electron affinity, and electronegativity.

Atomic Radius

The atomic radius of an element is half of the distance between the centers of two atoms of that element that are just touching each other. Generally, the atomic radius decreases across a period from left to right and increases down a given group. The atoms with the largest atomic radii are located in Group I and at the bottom of groups.

Moving from left to right across a period, electrons are added one at a time to the outer energy shell. Electrons within a shell cannot shield each other from the attraction to protons. Since the number of protons is also increasing, the effective nuclear charge increases across a period. This causes the atomic radius to decrease.

Moving down a group in the periodic table, the number of electrons and filled electron shells increases, but the number of valence electrons remains the same. The outermost electrons in a group are exposed to the same effective nuclear charge, but electrons are found farther from the nucleus as the number of filled energy shells increases. Therefore, the atomic radii increase.

Ionization Energy

The ionization energy, or ionization potential, is the energy required to completely remove an electron from a gaseous atom or ion. The closer and more tightly bound an electron is to the nucleus, the more difficult it will be to remove, and the higher its ionization energy will be. The first ionization energy is the energy required to remove one electron from the parent atom. The second ionization energy is the energy required to remove a second valence electron from the univalent ion to form the divalent ion, and so on. Successive ionization energies increase. The second ionization energy is always greater than the first ionization energy. Ionization energies increase moving from left to right across a period (decreasing atomic radius). Ionization energy decreases moving down a group (increasing atomic radius). Group I elements have low ionization energies because the loss of an electron forms a stable octet.

Electron Affinity

Electron affinity reflects the ability of an atom to accept an electron. It is the energy change that occurs when an electron is added to a gaseous atom. Atoms with stronger effective nuclear charge have greater electron affinity. Some generalizations can be made about the electron affinities of certain groups in the periodic table. The Group IIA elements, the alkaline earths, have low electron affinity values. These elements are relatively stable because they have filled s subshells. Group VIIA elements, the halogens, have high electron affinities because the addition of an electron to an atom results in a completely filled shell. Group VIII elements, noble gases, have electron affinities near zero, since each atom possesses a stable octet and will not accept an electron readily. Elements of other groups have low electron affinities.

Electronegativity

Electronegativity is a measure of the attraction of an atom for the electrons in a chemical bond. The higher the electronegativity of an atom, the greater its attraction for bonding electrons. Electronegativity is related to ionization energy. Electrons with low ionization energies have low electronegativities because their nuclei do not exert a strong attractive force on electrons. Elements with high ionization energies have high electronegativities due to the strong pull exerted on electrons by the nucleus. In a group, the electronegativity decreases as atomic number increases, as a result of increased distance between the valence electron and nucleus (greater atomic radius). An example of an electropositive (i.e., low electronegativity) element is cesium; an example of a highly electronegative element is fluorine.

Summary of Periodic Table Trends

Moving Left --> Right

* Atomic Radius Decreases
* Ionization Energy Increases
* Electronegativity Increases

Moving Top --> Bottom

* Atomic Radius Increases
* Ionization Energy Decreases
* Electronegativity Decreases

~GoL

Copy and Paste? Yer Not very impressive....

Anyway you can lock this mods if you want, I've handed the assessment in now. Thanks guys for your help.

:embrsd: I forgot to list the creator... My bad. It was from About.com... You see, I can't view the computer for more than 20 minutes on weekdays, so I did what I can under the rush of time.