Periodic Table Trends

The trends need to know are as follows:
1. Metallic properties
2. Atomic Radius
3. Ionization energy
4. Electronegativity
5. Reactivity
6. Ion Charge
7. Melting/Boiling Point
8. Density

If you were to look carefully at many of the properties of the elements, you would notice something besides the similarity of the properties within the groups. You would notice that many of these properties change in a fairly regular fashion that is dependent on the position of the element in the periodic table.  As you compare elements from left to right across the periodic table, you will notice a trend or regular change in a number of properties. The same thing happens if you go up and down on the periodic table and compare the properties of the elements. 

Atomic Radii

1) As you move down a group, atomic radius increases.

WHY? - The number of energy levels increases as you move down a group as the number of electrons increases.  Each subsequent energy level is further from the nucleus than the last.  Therefore, the atomic radius increases as the group and energy levels increase.

2) As you move across a period, atomic radius decreases.

WHY? - As you go across a period, electrons are added to the same energy level.  At the same time, protons are being added to the nucleus.  The concentration of more protons in the nucleus creates a "higher effective nuclear charge."  In other words, there is a stronger force of attraction pulling the electrons closer to the nucleus resulting in a smaller atomic radius.
 First Ionization Energy

Definition:  The energy required to remove the outermost (highest energy) electron from a neutral atom in its ground state.

1) As you move down a group, first ionization energy decreases.

WHY?
Electrons are further from the nucleus and thus easier to remove the outermost one. "SHIELDING" - Inner electrons at lower energy levels essentially block the protons' force of attraction toward the nucleus.  It therefore becomes easier to remove the outer electron

2) As you move across a period, first ionization energy increases.

WHY? - As you move across a period, the atomic radius decreases, that is, the atom is smaller.  The outer electrons are closer to the nucleus and more strongly attracted to the center.  Therefore, it becomes more difficult to remove the outermost electron.

Exceptions to First Ionization Energy Trends

1) Xs2 > Xp1  e.g. 4Be > 5B WHY? - The energy of an electron in an Xp orbital is greater than the energy of an electron in its respective Xs orbital.  Therefore, it requires less energy to remove the first electron in a p orbital than it is to remove one from a filled s orbital.  2) Xp3 > Xp4  e.g.  7N > 8O WHY? - After the separate degenerate orbitals have been filled with single electrons, the fourth electron must be paired.  The electron-electron repulsion makes it easier to remove the outermost, paired electron. (See Hund's Rule)

Metals

Common characteristics:

Metallic luster (shine)

Generally solids at room temperature

Malleable

Ductile

Conduct heat and electricity

Exist as extended planes of atoms

Combine with other metals to form alloys which have metallic characteristics

Form positive ions, e.g.  Na⁺, Mg²⁺, and Al³⁺

Nonmetals

Common characteristics:

Rarely have metallic luster (shine)

Generally gases at room temperature

Neither malleable nor ductile

Poor conductors of heat and electricity

Usually exist as molecules in thier elemental form

Combine with other nonmetals to form covalent

Generally form negative ions, e.g.  Cl^-, SO₄^2-, and N^3-

The differences in the characteristics of metals and nonmetals can be explained by the following:

Metals have relatively few electrons in their valence shells.

Metals have lower ionization energies than nonmetals.

Metals have smaller electron affinities than nonmetals.

Metals have larger atoms than nonmetals.

1) As you move across a period, metallic character decreases and nonmetallic character increases.

2) As you move down a group, metallic character increases and nonmetallic character decreases.

Semimetals (Metalloids)

A class of 8 elements that have properties of both metals and nonmetals.

B

Si

Ge

As

Sb

Te

Po

At

Reactivity
Metals and non-metals show different
trends.
The most reactive metal is Francium; the
most reactive non-metal is Fluorine

Ion Charge
elements ion charges depend on their
group (column).


Melting Point
Elements in the center of the table of the
highest melting point
Noble gases have the lowest melting
points
Starting from the left and moving right,
melting point increases (until the middle of
the table)

DENSITY
metals will tend to have higher density because of their positive charge. And radioactive elements also are also very dense. But you can say that as you go down the table, atoms are more dense.



Trend of Ionization:http://www.youtube.com/watch?v=ywqg9PorTAw
General Trend:http://www.youtube.com/watch?v=CHHy2ex0dw4
http://www.youtube.com/watch?v=h7XWqwgZII0&feature=related

Periodic Table Families

PERIODIC TABLE FAMILIES!

Families. We all have them. Sometimes, they may drive us crazy. But love them or hate them, we're stuck with the ones we've got, so we have to stick with them for what they're worth.

Just like us, the periodic table can be grouped into certain families as well! These families are:



Alkali metals: Shown as the yellow strip on the far left, alkali metals display properties such as high density levels, a single valence electron, low ionization energies, and a large atomic radii.

Alkali earth metals: Shown as the bright, bright blue strip next to the alkali metals, alkali earth metals have two electrons in their outer shells, low electronegativities, and slightly smaller atomic radii than the alkali metals.

Transition metals: These metals form the huge block between the alkali earth metals and the metalloids. They have high melting and boiling points, are malleable, and are good conductors of electricity.

Metalloids: These form a staircase consisting of boron, silicon, germanium, arsenic, antimony, and tellurium. Polonium is still under debate regarding its status as a metalloid. Metalloid properties vary widely. They have differing boiling and melting points, and are good semi-conductors.

Halogens: Halogens, as shown by the strip of pale yellow, are highly reactive elements. They display high levels of electronegativity, and all have seven valence electrons.

Noble gases: The orange family on the far right, noble gases tend to be very stable because they have full outer shells. Noble gases have low boiling points and rarely lose or gain electrons. They display low electronegativity.

Rare earths: These are the two sky blue rows at the very bottom, the Lanthanide series and the Actinide series. These metals are good conductors, are usually silvery in colour, and have high densities.



And now here's a nice poem by Michael Carungi!

Carbon, the Champion

Carbon is the element that is the basis
Of all organic life in all places
It can bond with elements, that's a fact
Over ten million compounds to be exact
When united with a substance like air
Carbon dioxide is created, which is used to prepare
The growth of plants. And when Carbon is combined
With Hydrogen, fuels are made. Diamonds and Graphite are a kind
Of Carbon, where Diamonds are the hardest ever known
And a softer substance than Graphite is unknown
Carbon also has the highest melting grade
And Carbon cannot be artificially made.
Yes, Carbon is truly a great thing
And it should be declared the elemental king.

UNTIL NEXT TIME!

Valence Electron Predictions

Valence Electrons!

As you may recall from previous science courses, valence electrons are the electrons that are on the outermost energy level of an atom.

Valence electrons are the ones that take part in chemical reactions. The more unstable an atom is, the more likely it will gain or lose electrons to become stable.

Take a neutral sodium atom, for example. It has eleven electrons, two on the first energy level, eight on the second, and one on the last one. Sodium in this case has an open shell, because its outermost shell is not completely filled.

How how sad it is, to be that lonely little valence electron hanging out all by itself! Poor little particle! Won't anyone love him for who he is? He just wants to be LOVED! He doesn't need much at all, really, just someone out there to care for him, and he'll care for them in return too! Um...*cough* I mean...okay, anyway...moving on...

Noble gases do not gain or lose electrons easily. This is because they typically have closed shells, which mean that their outermost shells are filled up.

Take argon, for instance. See how all the available spaces are taken up by electrons?

Writing Core Notation to Figure out the # of Valence Electrons

Here's a trick for writing electron configurations that can save you a lot of time. Let's say you're writing the configuration for an ion of manganese - Mn ²⁺ .

A neutral atom of Mn has 25 electrons, and Mn²⁺  will have 23, since it lost two to gain a +2 charge. The electron configuration of Mn²⁺  will be 1s²2s²2p⁶3s²3p⁶4s²3d³

Now look closely at the noble gases. Notice how the first part of the configuration of argon (1s²2s²2p⁶3s²3p⁶) is the same as Mn²⁺. In this case, we can substitute the first part of Mn²⁺   that contains all the same configuration as Ar for the notation of Ar itself!

So now we have [Ar]4s²3d³.

To find all the valence electrons, count up the # of electrons that was NOT included in the core. Do NOT count filled d- or f- shells.

Mn²⁺  has 5 valence electrons. Amazing!

Looking for further help? Check out this informative website!

Complete Electron Configurations

History of Periodic Table

DID you know that there has been a prodigious amount of change in the way periodic table arranges?

In the Beginning

the first scientific discovery of an element occurred in 1649 when Hennig Brand discovered phosphorous. During the next 200 years, a vast body of knowledge concerning the properties of elements and their compounds was acquired by chemists. By 1869, a total of 63 elements had been discovered. As the number of known elements grew, scientists began to recognize patterns in properties and began to develop classification schemes.

Law of Triads

In 1817 Johann Dobereiner noticed that the atomic weight of strontium fell midway between the weights of calcium and barium, elements possessing similar chemical properties. In 1829, after discovering the halogen triad composed of chlorine, bromine, and iodine and the alkali metal triad of lithium, sodium and potassium he proposed that nature contained triads of elements the middle element had properties that were an average of the other two members when ordered by the atomic weight (the Law of Triads).

Law of Octaves

John Newlands, an English chemist, wrote a paper in 1863 which classified the 56 established elements into 11 groups based on similar physical properties, noting that many pairs of similar elements existed which differed by some multiple of eight in atomic weight. In 1864 Newlands published his version of the periodic table and proposed the Law of Octaves (by analogy with the seven intervals of the musical scale). This law stated that any given element will exhibit analogous behavior to the eighth element following it in the table.

Who Is The Father of the Periodic Table?

There has been some disagreement about who deserves credit for being the "father" of the periodic table, the German Lothar Meyer (pictured here) or the Russian Dmitri Mendeleev. Both chemists produced remarkably similar results at the same time working independently of one another. This consisted of about half of the known elements listed in order of their atomic weight and demonstrated periodic valence changes as a function of atomic weight. Unfortunately for Meyer, Mendeleev's table became available to the scientific community via publication (1869) before Meyer's appeared (1870)!

Discovery of the Noble Gases

In 1895 Lord Rayleigh reported the discovery of a new gaseous element named argon which proved to be chemically inert. This element did not fit any of the known periodic groups. In 1898, William Ramsey suggested that argon be placed into the periodic table between chlorine and potassium in a family with helium, despite the fact that argon's atomic weight was greater than that of potassium. This group was termed the "zero" group due to the zero valency of the elements. Ramsey accurately predicted the future discovery and properties neon.

Atomic Structure and the Periodic Table

Although Mendeleev's table demonstrated the periodic nature of the elements, it remained for the discoveries of scientists of the 20th Century to explain why the properties of the elements recur periodically.

In 1911 Ernest Rutherford published studies of the scattering of alpha particles by heavy atom nuclei which led to the determination of nuclear charge. He demonstrated that the nuclear charge on a nucleus was proportional to the atomic weight of the element.

 This charge, later termed the atomic number, could be used to number the elements within the periodic table. In 1913, Henry Moseley (see a picture) published the results of his measurements of the wavelengths of the x-ray spectral lines of a number of elements which showed that the ordering of the wavelengths of the x-ray emissions of the elements coincided with the ordering of the elements by atomic number. With the discovery of isotopes of the elements, it became apparent that atomic weight was not the significant player in the periodic law as Mendeleev, Meyers and others had proposed, but rather, the properties of the elements varied periodically with atomic number.


PERIODIC LAW: properties of chemical elements recur periodically when the elements are arranged from lowest to highest atomic #s.The question of why the periodic law exists was answered as scientists developed an understanding of the electronic structure of the elements beginning with Niels Bohr's studies of the organization of electrons into shells through G.N. Lewis' (see a picture) discoveries of bonding electron pairs.

The Modern Periodic Table

The last major changes to the periodic table resulted from Glenn Seaborg's work in the middle of the 20th Century. Starting with his discovery of plutonium in 1940, he discovered all the transuranic elements from 94 to 102. He reconfigured the periodic table by placing the actinide series below the lanthanide series. In 1951, Seaborg was awarded the Nobel Prize in chemistry for his work. Element 106 has been named seaborgium (Sg) in his honor.


If you are truly passionate about this, you can also check out these very educational video by BBC!
They summarizes the history from the first discovery to the first organizing to the modern periodic table we know now...

http://www.youtube.com/watch?v=nsbXp64YPRQhttp://www.youtube.com/watch?v=25lprEvoFJ8&feature=related
http://www.youtube.com/watch?v=25lprEvoFJ8&feature=related
http://www.youtube.com/watch?v=hHL80A93lCA&feature=related

Electron Configeration

Electronic Configeration = orbitals electron occupy + # of electrons in each orbital.
History: Bohr's proposal-electrons moves from one orbital to another when emit or absorb energy.
Energy level= amount of energy an electron can possess( n)
Ground State: all e- in lowest possible energy level
Excited State: 1 or more e- in energy level other than lowest available level

( orbital: region occupied by an electron in a specific energy level.-> S, P, D, F
( shell: set of all orbitals w same n-value. (n) = energy level
( subshell: set of orbitals of same type.

Comparing Energy level of Hydrogen to Polyelectronic atom


H (1S1)   1 e-
He ( 1S 2)  2 e-
Be (1S2 2S2)  4 e-

Orbital:
n=1 : only s-type possible
n=2: s, p-types are possible
n=3: s, p, d-types are possible
n=4: s, p, d, f-types are possible

Subshell:
s-subshell include 1 s-orbital
p-subshell include 3 p-orbitals
d-subshell include 5 d-orbitals
f-subshell include 7 f-orbitals



The diagram above shows how s, p, d, f are distrbuted on the periodic table for each element, and we can use this diagram to understand and write the electronic configeration.

TWO ESSENTIAL RULES:

1. electrons MUST be added to lowest energy orbital 1st.
2. maximum of 2 e- in each orbital:
 2 e- in s-type subshell, 6 e- in p-type subshell, 10 e- in d-type subshell, 14 e- in f-type

E-Configeration for Neutral atoms
like Na, P, Ca, Fe, ......

1.how many e-?
2.each e- has 1 upward and 1 downward arrow

ex. Silicon, Si has 14 e-.

Look at the diagram below and starts with the lowest energy level, then keep adding until it fulfills the all # of e-.

Si ( 1s2 2s2 2p6 3s2 3p2)

Exercise:

a)Mn b) Co c) Rb d) P

Answer:
a) 1s2 2s2 2p6 3s2 3p6 4s2 3d5
b) 1s2 2s2 2p6 3s2 3p6 4s2 3d7
c) 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s1
d) 1s2 2s2 2p6 3s2 3p3

Core notation: simplifies the work by replacing a set of configeration with the nearest noble gas with less atomic number.

try those with the same exercise above, did you get the answer?
a) [Ar] 4s2 3d5
b) [Ar] 4s2 3d7
c) [Kr] 5s1
d) [Ne] 3s2 3p3

E.Configeration for ions

negative: add e- (same as charge) to last UNFILLED subshell, where neutral atom left off.

Ex. P 3-
1s2 2s2 2p6 3s2 3p3 -> 1s2 2s2 2p6 3s2 3p6

positive: remove e- from largest n-value, remove p-electron before s-electron before d-electron.

Ex. Sn 2+
[Kr]5s2 4d10 5p2 -> [Kr] 5s2 4d10

I believe simple rule following like this won't preplex all of you!

Thanx for comin!

April 18, 2011

Hello, hello, hello! What a beautiful, nice sunny day it is! Today we had no rain at all with the sun shining above! So let's get straight down to what we learned today..... Today it was basically a review of what we learned in grade 9, ahhh grade 9 seems so long ago, gosh, ok right getting off topic here..... today we learned about......

Atomic Structure!

- The subatomic particles are: protons, neutrons and electrons.
 ˜Proton:
Symbol = (we use the P)
Relative Mass = 1
Electric Charge = +1
Location in the Atom = nucleus

 ˜Neutron:

Symbol =
Relative Mass = slightly bigger than 1
Electric Charge = 0
Location in the Atom = nucleus

 ˜Electron:

Symbol =  (we use the E)
Relative Mass = 0
Electric Charge = -1
Location in the Atom = cloud surronding the nucleus

- In a Neutral atom, # of Protons = # of Electrons

Atomic Number (Z): The proton number
- The Atomic Number (Z) is the number of protons found in the nucleus of an atom.
- Atoms have no overall electric charge (i.e. charge of atom is zero).
- Atomic Number = # of protons = # of electrons

Ions:
- Atoms that have gained or lost electrons are called ions.
- An ion is an electrically charged atom (or groups of atoms).
- Negatively-charged ion = anion
- Positively-charged ion = cation
- For Ions; # of electrons = protons - charge

Mass Number (A):
- Is the total number of protons and neutrons or atomic mass number.
- Atomic Number = the # of protons
- Atomic Mass = # of protons + # of neutrons
- Number of neutrons = Mass Number - Atomic Number
- Mass Number = # of protons + # of neutrons

Atomic Mass:
- The average mass of an element's isotopes.
- The atomic mass is very close (not exact!) to the mass number.
- # of neutrons = Mass Number - Atomic Number
                         Atomic Mass - Atomic Number
- The atomic mass is an average! The mass number is usually calculated by rounding the atomic mass to the nearest whole number.

Isotopes:
- Are atomic species having the same atomic number (protons) but different atomic masses/mass numbers (neutrons). Basically, same # of protons and electrons, but different # of neutrons.

Whew! Alot of information but pretty simple to comprehend, which is a very good thing :) So now let's give you some important visuals!










Hope those are helpful, and now here are some worksheets & extra information for you:
1) http://www.sciencejoywagon.com/chemzone/02atomic-structure/
2) http://misterguch.brinkster.net/propertyworksheets.html
3) http://docs.google.com/viewer?a=v&q=cache:Svl1HS3qzfsJ:misterguch.brinkster.net/001_021.doc+atomic+structure+worksheet&hl=en&gl=ca&pid=bl&srcid=ADGEESj_WYrtiEN2lX55CTrVQlm1hc7YmKY1mVbgw1DoKBPwfd7ZkvmnBtHXM9MonGdgMHr7rYPbP_zspbkgYZtxil2dk5ld9UqPcFm79Mukr5YOnQD-v3EXpsaki74QhGzie3QH_Cry&sig=AHIEtbTL6KdKPbx4-wFOQzdx-a1ZI3Et9g&pli=1
4) http://cmsweb1.loudoun.k12.va.us/52820831134912597/lib/52820831134912597/Atoms%20and%20Atomic%20Theory/Homework/ws.atomic.20structure.20set.pdf
5) http://www.allaboutcircuits.com/worksheets/atomic.html
6) http://staff.fcps.net/jswango/unit2/atomic_structure/Basic%20Atomic%20Structure%20Worksheet.pdf
7) http://chemistry.about.com/library/weekly/blatomquiz2.htm
8) http://www.softschools.com/quiz_time/chemistry/atomic_structure/theme81.html

Last but not least, here are some youtube videos:
1) http://www.youtube.com/watch?v=7ohfJ9ku8gc&feature=fvwrel
2) http://www.youtube.com/watch?v=WWxnZK_g5ug
3) http://www.youtube.com/watch?v=Z6Y4Ffod1hQ&feature=related

Thanks for reading & have a great long weekend! Wooooooooooo!!! :)

[History of the Atom]

|We are starting a new unit today called the Atomic Theory.|

|This is the start of atom and where ideas of atoms emerged...|

|Why do you think ancient philosophers were so interested in the phenomenon of atoms? |

|Perhaps they were curious about the foundation of all matters and that was when they came up with atom...|

• Aristotle became the first to propose that all matter consists of four elements either

|EARTH|  |WATER| |FIRE| OR |AIR|
Aristotle: http://www.iep.utm.edu/aristotl/

• Alchemist Brought Au to Earth by changing the mechanics of metal!
• The Four Element theory lasted for about 2000 years!
• Democritus proposed that atom = indivisible particles though this was not a testable theory and merely a model.=[ There was also no discovery of nucleus, proton or electrons. Thus, it cannot be used to explain chemical reactions

Alchemist full metal image

• Lavoisier started 1st version of Law of Conservation of Mass & Law of Definite Proportions.

Law of Definite Proportion means that the % of an element on a compound will never change!

• Proust proposed:  even if compound are broken down, the ratio in the compound would still apply to the products.
• Dalton provided a vision of an atom = solid, sphere and his experiments and discoveries were based on the Law of Conservation of Mass
• Dalton's Law #1: Elements made up of atoms
• Dalton's Law #2: All atoms of a element are identical.
• Dalton's Law #3: atoms are distinguished by relative weight
• Dalton's Law #4: Atom(of different element) + Atom(of different element) = chemical Compound, with same # of same types of atom
• Dalton's Law #5: Atoms are indivisible, non-destructable and they cannot be created.

A new revoluntion started when another chemist proposed his hypothesis stating that there are both positive and negative particles inside an atom.

Thomson gave new characteristics to an atom: they are solid spheres with negative particles embedded (electrons) and positive particles (protons). He figured out the mass of electron, 9.10938188 x 10 -28 grams.



Adding to Thomson's discovery, RUtherford explained that all the positive particles gathers in the dense centre of an atom. They electrons are outside of the dense positive center. His model is called the planetary model.

this is the famous GOLD LEAF experiment provides evidence for the positive centre of an atom.

Niels Bohr contributed to this RACE of ATOM most notably for discovering the Electrons in specific energy level or shells outside the nucleus.

BOHR diagram of atoms 

After evolutionary change in understanding of atom, our knowledge of an atom today is the smallest
particle of an element that reatains the properties of element. There are 3 kinds of particles called subatomic particles: e, p, n


http://www.youtube.com/watch?v=h0UllT0DE7s&feature=related
This is an exellent video presenting the historical timeline of the atomic theory!
ENJOY~