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Ch 5 Sec 5-1 Around 1850, more than 60 elements had been discovered. Chemists had to learn the properties of these elements as well as the compounds they formed, which was very difficult. At that time, there was no accurate way to determine the atomic mass or the # of atoms of an element in a particular compound. Different chemists used different atomic masses for the same element, which made it impossible to compare results. In 1860, a group of chemists met at the 1st International Congress of Chemists in Germany. • They wanted to settle the issue of atomic mass and other matters. • Italian Chemist – Stanislao Cannizzaro – presented a convincing method for accurately measuring the relative mass of an atom. • They agreed on his method as a standard and initiated a search for relationships between the atomic mass and other properties of the elements. Dmitri Mendeleev (Russian teacher) heard about the new method to calculate the atomic masses, he decided to include the new values in a chemistry textbook he was writing. Arranged note cards with the name of the element, atomic mass, physical and chemical properties. Started to arrange the note cards in various ways – he looked for trends. Organized them in increasing atomic mass – then the phy. & chem. properties repeated periodically. (See pg 124 for the 1st periodic table.) ** Mendeleev left holes in the table but he predicted the existence & properties of 3 elements that would fit in the spaces. Therefore, he is credited with the periodic law. A few questions remained: 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 the chemical periodicity? The first question wasn’t answered until about 40 years later after the first published table. Henry Moseley (1887 – 1915) – worked with Rutherford. • He examined the spectra of 38 different metals. • Discovered a previously unrecognized pattern. • The elements fit better if they were arranged by the # of p+. • This led to the modern definition of the atomic mass and that the periodic table should be based on the atomic # not the atomic mass. Periodic Law – The physical and chemical properties of the elements are periodic functions of their atomic #. The periodic table continues to undergo slight changes, but it is now arranged so that elements with similar properties fall in the same column. Sec 5-2 In the periodic table, we have special names for different sections. 1. 2. 3. 4. 5. 6. 7. 8. 9. Alkali Metals – group 1 (except Hydrogen) Alkaline – Earth – group 2 Transition Metals – groups 3 – 12 Main group elements – groups 13 – 18 Halogens – group 17 Noble gases – group 18 Lanthanides Actinides Metalloids Alkali Metals – • Silvery appearance • Soft – could cut with a knife • So reactive they are not found as free elements in nature • React violently with most non-metals • React violently with water to produce hydrogen gas • Usually stored in kerosene Alkaline – earth metals – • Harder, denser, stronger, and have higher melting points than the Alkali metals. • Less reactive than the Alkali metals, but are still reactive enough that they are not found free in nature. Transition Metals •As all metals, they are good conductors of electricity and heat. They also have a high luster. •They are typically less reactive than the alkali and alkali-earth metals. Therefore some are found in nature as pure elements. Gold for example. Main group elements Properties of the elements vary greatly. Includes all of the non-metals except H. Contain all the metalloids Halogens Most reactive non-metals React vigorously with most metals to form “salts”. Gases at room temperature Noble Gases Unreactive because they contain 8 outershell e-. Section 5 - 3 Electron Configuration and Periodic Properties Atomic Radii One way to express an atom’s radius is to measure the distance between two identical atoms that are chemically bonded together, then divide the distance by two. Atomic Radius— ½ the distance between the nuclei of identical atoms that are bonded together What would the atomic radius be of the Chlorine atom below: 198 pm Periodic Trends As you travel from left to right across the periodic table, the atomic radii decreases Caused by the increased pull of the e- by the nucleus Group Trends As your travel from the top of the periodic table to the bottom, the atomic radii increases Periodic Table of Atomic Radii Decrease in Atomic Radii Increase in Atomic Radii Examples Ex 1: Of the elements magnesium, Mg, chlorine, Cl, sodium, Na, and phosphorus, P, which has the largest atomic radius? Explain. Ex 2: Of the elements calcium, Ca, beryllium, Be, barium, Ba, and strontium, Sr, which has the largest atomic radius? Explain. Ionization Energy Ionization Energy—the amount of energy needed to remove an e Can be expressed as: A + energy A+ + eAtom Ion with a 1+ charge Ex: Na + energy Na+ + eCa + energy Ca2+ + e- Period Trends Group 1 metals have the lowest ionization energythey lose e- the easiest Group 18 noble gases have a very high ionization energy the do not lose e- very easily Increase across a period is caused by the increase in nuclear charge In general, nonmetals have a higher ionization energy than metals Group Trends As you travel down the periodic table the e- in the outer shell get further away from the nucleus Can be removed much more easily Valence e- can very easily be taken off of the outer shell. However, after that it takes an immense amount of energy Examples Consider two main-group elements A and B. Element A has a first ionization energy of 419 kJ/mol. Element B has a first ionization energy of 1000 kJ/mol. For each element, decide if it is more likely to be in the s block or p block. Which element is more likely to form a positive ion? Examples Consider the four hypothetical main-group elements Q, R, T, and X with the outer electron configurations indication below. Then answer the following questions: Q: 3s2 3p5 a) b) c) d) R: 3s1 T: 5s2 4d10 5p5 X: 5s2 4d10 5p1 Identify the block location of each main-group element. Which of these elements are in the same period? Which are in the same group? Which elements would you expect to have the highest first ionization energy? Which would you expect to have the lowest? Which of the elements is most likely to for a 1+ ion? Metals (on the left) need to lose e- to form an octect, or eight e- in their outer shell. Makes them (+) charged—Called a cation Ex: Li+1, Na+1, and K+1 all need to lose 1 e- to be happy Ex: Be2+, Mg2+ and Ca2+ all need to lose 2 e- to be happy Metals are such good conductors because they want to lose e- so easily Nonmetals don’t want to lose any e-, they would rather gain. Electron Affinity Electron Affinity—the energy that occurs when an e- is added to an atom. Adding an e- to an atom would make the atom (-) charged— this is called an anion Ex: F-1, Cl-1, and Br-1 all need to gain 1 e- to be happy Ex: O-2, S-2, and Se-2 all need to gain 2 e Cations, like Na+1 will combing with anions, like Cl-1 so both are happy Will form NaCl (table salt) Ionic Radii Cation—Positive ion Formed by the loss of one or more e Results in the decrease in atomic radius Anion—Negative ion Formed by the addition of one or more e Results in the increase in atomic radius Period Trends—Ionic Radii Metals on left form cations Radii decreases across a period b/c e- cloud shrinks Nonmetals at upper right form anions Group Trends—Ionic Radii Gradual increase in Ionic Radii as you travel down a group Valence Electrons The electrons that interact to form chemical compounds are those in the highest energy levels Valence Electrons—The electrons available to be lost, gained, or shared in the formation of chemical compounds Located in incompletely filled main-energy levels Ex: When Na loses an electron it becomes Na+. The e- that is lost is known as a valence e Occupy the s and p orbitals Examples How many valence e- does the following have: 1) 2) 3) 4) 5) 6) Mg Cs C O Ar Al Electronegativity Valence e- hold chemical compounds together The negative charge is concentrated closer to one atom than another Linus Pauling—devised an electronegativity scale Electronegativity—a measure of the ability if an atom in a chemical compound to attract electrons Most electronegative element is Fluorine Given an electronegativity value of 4 Period Trends Tend to increase across each period Alkali and Alkaline-Earth Metals are the least electronegative Nitrogen, Oxygen, and halogens are the most electronegative Atoms attract e- strongly in compounds Group Trends • Tend to decrease down a group • Noble Gases are unique in that some are unreactive and have no electronegativity Example Among the elements Ga, Br, and Ca which has the highest electronegativity. Why?