Electrons and the Periodic Table:

Transcription

1 Electrons and the Periodic Table: Development of the Periodic Table: The development of the periodic table started in the 1800 s. * By the end of the 1700 s, scientists had identified about 30 elements. This amount doubled in the next 100 hundred years. * In the early 1800 s, a German by the name of Dobereiner grouped the existing elements into triads (groups of 3) according to similar properties. * In 1865, John Newlands (English), a huge music fan, arranged the now 62 elements by atomic mass and found that properties repeated themselves every 8 elements. He called them octaves. * In 1869, Dmitri Mendeleev (Russian) and Lothar Meyer (German) published nearly identical arrangements of elements according to atomic mass. But Mendeleev was better at communicating his ideas and gets all the credit to this day. * Mendeleev created the table because he was a teacher and wanted an easy way for his students to memorize the elements and their properties. * Like Newlands, Mendeleev found that when elements were arranged according to increasing atomic mass, properties repeated themselves every 8 elements. * This is called the PERIODIC LAW (when elements are arranged according to atomic number, properties repeat periodically). We now use atomic number because it s more convenient and it works better. * The true genius of Mendeleev s work was that he was able to predict the properties of elements that weren t even discovered yet. He left spaces in his table and identified what properties they would have. Once these elements were discovered, it was found that he was right almost every time.

2 * All the creators of the periodic table were European. What impact do you think this fact had on the creation of the periodic table (it s appearance and even it s existence)? The Current Periodic Table: * Elements with similar properties are placed in groups or families. These are the columns in the table. These groups are number 1 through 18, but some of them also have special names: Group 1: Group 2: Group 17: Group 18: Groups 3 to 12: * Elements are also arranged into rows and it is across these rows that properties repeat. For example, as you travel across any given row, the size of the atom decreases. At the start of the next row, the atom is large and gets smaller as it goes across the next row, etc. * There is also a special region at the bottom of the table where the Lanthanides and Actinides are. They are not really a group, but have unique chemistry. * The rows in the periodic table are called periods because they show trends in many element properties.

3 * The table is also split in 2 by the staircase: To the left of the staircase are the METALS: To the right of the staircase are the NON-METALS: Those touching the staircase are considered to be METALLOIDS/SEMI-METALS: ASSIGNMENT: -Fill in the Blank Periodic Table with atomic numbers -Identitfy (with colours, patterns, etc) the following on the periodic table: -Actinides -alkali metals -alkaline earth metals -halogens -transition metals -lanthanides -metalliods -metals -noble gases -nonmetals -Be sure to include a legend!

4 Electron Configuration: * As it turns out, the electrons are the most important subatomic particle when trying to understand and element s reactivity and other properties. * Electron configuration allows us to get a vague idea about how electrons are arranged in the space around the nucleus. Much of it was figured out be Neils Bohr when he was studying the electron of the hydrogen atom. * Where are electrons exactly? * When we talk about an electrons position, we use probabilities. We talk about regions of space where there is a good chance we will find an electron at any given moment. Orbitals are: There are 4 different types of orbitals. They tell us the shape of the expected path of an electron found in that orbital. s- orbitals: p-orbitals:

5 d-orbitals: f-orbitals: * There are also shells: * The way that electrons are arranged in an atom is called its ELECTRON CONFIGURATION. The Rules of Electron Configuration: 1/ ELECTRONS ARE ALWAYS PLACED FROM LOWER ENERGY TO HIGHER ENERGY POSITIONS. The lowest energy position is nearest to the nucleus (the lowest energy position for you is on the ground). The further away from nucleus we get, the more energy it takes to get there. 2/ ONLY TWO ELECTRONS CAN FIT IN AN ORBITAL. This is because of electron-electron repulsions (remember they are both negatively charged). In order for the two to avoid each other, they spin in different directions. 3/ HUND S RULE. I call this rule the anti-social electrons on a bus rule. When electrons have multiple orbitals of equal energy available (ex: p-orbitals), they fill the orbitals singly first, then double up.

6 We are now ready to try some. We will be filling orbitals with the electrons that an atom has. We represent orbitals with circles and electrons with half arrows. Representing the electron configuration this way is called an ORBITAL DIAGRAM. Ex: the sodium atom The Electron Configuration for sodium is: Ex: Oxygen The Electron Configuration for oxygen is:

7 Ex: Co 2+ The Electron Configuration for this cobalt cation is: Try One: P 3-

8 Try another: the argon atom * What do you notice about the 2 examples you tried? * Substances with the same number of electrons will automatically have the same electron configuration. Substances with the same electron configuration are called ISOELECTRONIC. * Oddly enough there is an easier way to do this. It avoids the orbital diagrams. We can get the electron configuration straight from the periodic table. - SEE HANDOUT -

9 * All of the electron configurations we have done so far have been the long version. We also have an abbreviated version. Just like in the written English language (can t, won t, it s), we use a symbol to represent missing parts. We use the NEAREST, PREVIOUS NOBLE GAS to sum up all the electrons before that. **** YOU WILL BE EXPECTED TO BE ABLE TO SHOW THE ORBITAL DIAGRAM AND LONG & ABBREVIATED VERSIONS OF ELECTRON CONFIGURATION. **** Valence Electrons: * Not all of the electrons are really important in the element s chemistry. Only the electrons in the outermost shell (farthest away from the nucleus) is involved in the chemical reaction of an atom. * Electrons that are in the outermost shell of an atom are called VALENCE ELECTRONS. The ones in all the shells below are called CORE ELECTRONS. * Remember that that when we did electron configuration, the outermost shell would be the one with the highest shell number. Ex: How many valence electrons are in Bromine? Ex: How many electrons are there in S 2-? Try One: How many valence electrons are there in Xenon?

10 Lewis Dot Structures: * are symbols that SHOW the number of valence electrons that an element has. Electrons are shown with dots and THE ELECTRONS STAY AS FAR AWAY FROM EACH OTHER. THEY PAIR UP AFTER THEY FILL THE SINGLE SPACES. Ex: Chlorine Ex: Carbon: Try these: Argon, Chromium, Iron, Phophorus, Sulphur, Boron The Octet Rule:

11 * Valence electrons determine (although not completely) the reactivity of an element. The Octet Rule determines how reactive the valence electrons. OCTET RULE: Which elements are the only ones that have a full outer shell of electrons? What do we know about these elements? Ex: Boron Ex: Lithium Ex: Bromine Ex: Carbon Do these: (will it gain or lose, how many, what will be the charge?) sodium, magnesium, aluminum, silicon, phosphorus, sulfur, chlorine, argon

12 For each of the elements you just did, which other elements in the periodic table would gain/lose the same and, therefore, have the same charge? Periodic Trends: * Certain trends occur across periods/rows and then repeat periodically. We have already seen the trend for charges and we know it has to do with the octet rule (gain/lose). * We will also be looking at Ionisation Energy, Atomic Radius, and Electronegativity. 1/ Ionisation Energy (IE): * IONISATION ENERGY is the energy it takes to rip an electron from an atom (difficult to do = high energy & vice versa) I want you to graph the following data, draw a line (point to point), outline the line, and label each dot with the element symbol. Element Ionisation Energy (kj)

13 H (1) 1312 He (2) 2372 Li (3) 519 Be (4) 900 B (5) 799 C (6) 1088 N (7) 1406 O (8) 1314 F (9) 1682 Ne (10) 2080 Na (11) 498 Mg (12) 736 Al (13) 577 Si (14) 787 P (15) 1063 S (16) 1000 Cl (17) 1255 Ar (18) 1519 K (19) 418 Ca (20) 590 What trend do we see when we go across a row/period? Why?

14 What trend do we see when going down a group/family? Why? There s a blip when going from group 15 to group 16. Why? There s also a blip when going from group 2 to 13. Why?

15 2/ Atomic Radius/Size: * As we go across a row/period, atomic size decreases. Why? * As we go down a group/family, atomic size increases. Why? ALSO:

16 * The anion of an atom is larger than the atom. Why? * The cation of an atom is smaller than the atom. Why? 3/ Electronegativity (EN): * ELECTRONEGATIVITY is a measure of an element s ability to bring electrons to itself in a chemical bond. * When we go across a row/period, electronegativity increases. Why? * When we go down a group/family, electronegativity decreases. Why?