Follow it … or else! Section 1

By 1860 more than 60 elements had been discovered. Chemists had to memorize the properties of the elements AND the properties of the compounds they formed. Also, there was no method for determining the atomic mass or number of atoms of an element.

 In September 1860, a group of chemists assembled at the First International Congress of Chemists in Karlsruhe, Germany

 At this meeting Italian chemist Stanislao Cannizzaro presented a method to accurately measure the masses of atoms  Thanks to him, chemists were able to come up with agreed upon values for atomic mass

 Russian chemist  He was writing a book on the elements and wanted to organize them by their properties.

 He did this by placing the name and atomic mass of each element together on a card and listed the physical and chemical properties of each element on the card

 Mendeleev noticed when elements are arranged by increasing atomic mass certain similarities of their chemical properties appeared at regular intervals

 This repeating pattern is known as periodic  Created the first periodic table in 1869

Many chemists accepted Mendeleev’s periodic table and this earned him credit as the discoverer of the periodic law. However… 1. Why could most of the elements be arranged in the order of increasing atomic mass but a few could not? 2. What was the reason for chemical periodicity?

Henry Moseley analyzed the spectra of different metals. He realized they fit into patterns better when they were arranged by increasing number of protons.

 Discovered the elements fit into patterns better when they were arranged by increasing atomic number Led to the modern definition of…  Periodic Law- The physical and chemical properties of the elements are periodic functions of their atomic numbers

When elements are arranged in order of increasing atomic number, elements with similar properties appear at regular intervals.

 Since Mendeleev, chemists have discovered and synthesized new elements  Periodic Table- the arrangement of the elements by their atomic numbers so elements with similar properties fall in the same group

 1894 English physicist John William Strutt and Scottish chemist Sir William Ramsay discovered Argon  1895 showed Helium existed on Earth  1898 discovered Kr and Xe

 Final noble gas, Radon, Rn, was discovered in 1900 by Friedrich Ernst Dorn.  Found in Group 18

 The 14 elements with atomic numbers from 58 (cerium, Ce) to 71 (lutetium, Lu).  Also called the rare earth elements  Used as catalysts in oil refinery, television industry (used to make the red color), streetlights, searchlights, stadium lights, coloring ceramics and glass, camera lenses and binoculars

 14 elements with atomic numbers from 90 (thorium, Th) to 103 (lawrencium, Lr).  Used in cancer therapy, making atom bombs, used to power heart pacemakers, used as power sources to transmit messages from the Moon to Earth.  The lanthanides and actinides belong in periods 6 & 7

 Who was the first person to attempt to organize the elements by their properties?  What is the definition of periodic?  How are elements now arranged in the periodic table of elements?  Where are the noble gases found?  Why are periods 6 & 7 placed below the periodic table?

Section 2 Electron Configuration and the Periodic Table

 Write the electron-configuration notation for He, Ne, and Ar  Now look at the highest energy level, what do you notice?

 Since the highest energy level of the noble gases is completely full the noble gases are stable and therefore unreactive

 Elements are arranged vertically b/c of similar chemical properties.  They are also arranged horizontally in rows (periods)

The length of the period (# of elements in each period) can be determined by the # of electrons it will hold. Wow!

For example: In the first period, the 1s sublevel is being filled… The 1s sublevel can hold 2 electrons… Therefore the first period has two elements!

 The highest energy level number of an element’s electron configuration notation also tells you what period it is in!

 Lithium Electron Notation 1s²2s¹ Highest Energy Level: 2 Period it is located in: 2

 What period is the element with the configuration [Xe] 6s² in?

The periodic table is also divided into s,p,d, and f sublevel blocks. The name of each block is determined by whether an s, p, d, or f sublevel is being filled in successive elements of that block.

Nitrogen # electrons: 7 Configuration: 1s²2s²2p³ What block is Nitrogen located in? p block

 Chemically reactive metals  Group 1 is more reactive than Group 2 This is b/c they only contain 1 electron in the 1s sublevel. One electron is easier to lose than 2 electrons. Group configuration: ns¹

 Group 1 elements are called the alkali metals  Silvery appearance and soft enough to cut with a knife  Not found in nature b/c so reactive

 Elements of Group 2 are called alkaline-earth metals.  Harder, denser, and stronger than Group 1 elements  Not found in nature; less reactive than Group 1 elements though

 Group configuration: ns²

 Without looking at the periodic table, give the group, period, and block where the element with the configuration [Rn] 7s¹ is located.

 Hydrogen and Helium

 Located in Group 1 but shares none of the same properties as the other Group 1 elements  Hydrogen’s properties do not resemble those of any group.

 Has an ns² configuration like Group 2 but belongs to Group 18  This is due to its full sublevel making it very unreactive.

Without looking at the table, identify the group, block, and period where [Xe]6s² is found.

 Group 3 configuration: (n-1)d¹ns²  Group 12 configuration: (n-1)d^10ns²

 There are deviations from orderly “d” sublevel filling in Groups 4-11  Elements in these groups do not necessarily have identical electron configurations.

 Notice: The sum of the outer “s” and “d” electrons is = to the group number.

 Platinum Notation: [Xe]4f^145d^96s¹ Add up 9 and 1 = 10 Platinum is located in Group 10

 Transition Metals  They have metallic properties  They are good conductors of electricity and have a high luster  Less reactive (do exist as free elements in nature)

 Group 13 group configuration: ns²np¹  Group 14 group configuration: ns²np²  Group 18 group configuration: ns²np ⁶

 Main-group elements Except….. Group 17 elements are halogens These are very reactive

 They involve filling the 4f sublevel  Shiny metals  Similar reactivity to Group 2 elements

Periodic Trends

 One-half the distance between the nuclei of identical atoms that are bonded together  p. 150 Fig. 12

 Periodic Trend As you move from left to right across Period 2, the atomic radius decreases

 The reason is because of the increasing positive charge of the nucleus.  The electrons are pulled closer to the positively charged nucleus

 Generally, increases as you go down a group  There are exceptions (Group 13 elements)

 Of the elements magnesium, chlorine, sodium, and phosphorus which has the largest atomic radius?

 Energy required to remove an electron from an atom

A + energy  A+ + e- A: elementA+: Ion e- : electron

 Ion- atom or group of atoms that has a positive and negative charge - A+ from the previous example  Ionization- process that results in the formation of an ion

 In general, the energy will increase across each period

 In general, they will decrease down the groups

 Which of the following elements belongs in s block and which one belongs in the p block: Element A: Ionization energy 419 kJ/mol Element B: Ionization energy 1000 kJ/mol

S Block: Element A P Block: Element B

 Electrons used in bonding  Can be lost, gained, or shared  Located in incompletely filled main-energy levels

 Measure of the ability of an atom in a compound to attract electrons from another atom in the same compound.  Causes uneven distribution of charge

 Increase across each period  There are exceptions

 Decrease down a group or stay about the same  Some noble gases will not form compounds therefore they have no electronegativity value assigned to them

 Energy change occurring when an electron is gained by a neutral atom

 Positive Ion is called a cation *Formation of a cation leads to a decrease in atomic radius

 Negative Ion is called an anion *Formation of an anion leads to an increase in atomic radius

 Cationic radii decrease across a period  Anionic radii decrease across the periods in Groups 15-18

 Gradual increase in ionic radii down a group