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Unit 13 Chemical Periodicity A Little Bit of History In the early 1800s, German chemist J.W. Dobereiner observed that several elements could be classified into sets of three which he called triads. • (Li, Na, K), (Ca, Sr, Ba), (Cl, Br, I) • The elements within each triad had similar properties. • Several properties of the middle elements were averages of the other two. In 1865, English chemist J.A.R. Newlands observed that when the then 62 known elements were arranged in order of increasing atomic mass, every eighth element had similar properties. • This was known as the Law of Octives • The first was light the eighth; Na was like Li • The second was like the ninth, Mg was like Be • He knew nothing of the noble gases In 1869, Russian chemist Dmitri Mendeleev and German chemist Lothar Meyer published nearly identical schemes for classifying elements. • Mendeleev got the credit because he published his first and he was more successful at demonstrating it value. • Mendeleev eventually produces the first working periodic table of the elements. • He arranged his table so that elements in the same column had similar properties. • He switched the order of three pairs of elements so as to keep elements in columns with similar properties • He switched the order of three pairs of elements so as to keep elements in columns with similar properties. • This was contrary to all other chemists who arranged elements according to increasing atomic mass. • He boldly pronounced that perhaps the calculated atomic masses for those mixed up elements should be recalculated • Mendeleev also had the insight to predict the existence and some of the properties of three new elements (these elements were missing from his periodic table. In 1913, English chemist Henry Moseley determined the atomic number of the elements. • He already know that protons are positively charged particles that exist inside of nuclei. • He used x-rays to determine that each element had a certain number of positive particles in its nucleus. • He said these positive particles must be the protons discovered by Goldstein. • This discovery led to the term atomic number. The modern periodic law states that when elements are arranged in order of increasing atomic number, their physical and chemical properties show a periodic pattern. Mosley’s discovery proved Mendeleev to be correct. Different kinds of periodic tables may present different kinds of information or present information in different ways. • The shape of the periodic table comes from the periodic law. • Elements with similar properties are aligned in vertical columns called groups or families. • The horizontal rows are called periods. • Some of the groups have family names. (Alkali metals, alkaline earth metals, halogens, and the noble gases.) • Elements can be classified as metals, nonmetals, and semimetals. • Metals are typically solids at room temperature. • characteristic shine, conduct heat and electricity, malleable (hammered into sheets), ductile (drawn into wires), lose e-‘s to become cations • Nonmetals can be solids, liquids, or gases • do not have luster (dull), poor conductors of heat and electricity, brittle. • quite a variation in physical properties (some are colored, others are colorless; some are soft solids while others form hard solids); gain e-‘s to become anions • Semimetals have some properties of metals and some of nonmetals, or they have properties that are intermediate. The outermost “s” and “p” electrons, which are largely responsible for an atom’s chemical behavior, are called valence electrons. • The elements in a group have similar properties because they have the same number of valence electrons (which means similar electron configurations). • To save space in writing electron configurations and to focus attention on valence electrons, chemists often use abbreviated electron configurations. • An atom’s inner electrons are represented by the symbol for the nearest noble gas with a lower atomic number. (This is called the noble gas inner core) Octet rule is a very important concept to understanding bonding in chemistry. • Octet refers to the eight outer electrons of an atom. • An atom with all of its “s” and “p” electrons is extremely stable. • The noble gases have those eight outer electrons and will not bond with another atom. • Metals will lose electrons to back up to have the same number of electrons as the previous energy level’s noble gas.(they become cations) • Nonmetals will gain electrons to have the same number of electrons as the noble gas at the end of the row. (they become anions) • Sometimes nonmetals share electrons to gain the octet The key to understanding the shape of the periodic table is to examine the elements’ electron configuration. (The four main sections of the periodic table correspond to the four sublevels, s p d f.) • The s-block elements are groups 1 and 2. • The p-block elements are groups 3-8. • The d-block elements are the elements whose last electrons fill the d sublevel. • The f-block elements are the elements whose last electrons fill the f sublevel. The shape of the periodic table is a result of the way the electrons fill the s, p, d, and f orbitals of the different energy levels. • The four blocks have other names as well. • The elements in the s- and p- blocks are called the representative elements. • The elements in the d-block are called the transition metals. • Those in the f-block are known as the inner transition metals. Periodic Trends Many properties of the elements change in a predictable way as you move through the periodic table. These systematic variations are called periodic trends. The periodic trends that you must understand are… • Atomic radius • Ionization energy • Electronegativity • Shielding Effect Atomic Radius The atomic radius is the distance from the center of an atom’s nucleus to its outermost electron. • Atoms get larger going down a group. (ie. size increases with increasing atomic #) • Atoms get smaller moving from left to right across each period. Why? – As a rule, atoms that have more positive charge in their nuclei exert a stronger pull on the electrons in a given principal quantum number. – A stronger attractive force shrinks the electrons’ orbitals and makes the atom smaller. • As you learned in Unit 5, an atom can gain or lose electrons to form an ion. • When an atom loses electrons it becomes smaller. • Loss of electrons not only vacates the atom’s largest orbitals, it also reduces the repulsive force between the remaining electrons, allowing them to be pulled closer to the nucleus. • When an atom gains electrons it becomes larger. • A fluoride ion is larger than the fluorine atom mainly because of the greater number of electrons, which increases the electric repulsive forces among them. • The increased repulsions spread out the electrons, thus making the ion larger than the atom. • The atoms of a particular element typically form only certain ions. • The elements on the left side of the table form positive ions. (cations) • The elements on the right side of the table form negative ions. (anions) • The elements in the eighth column do not form ions. (noble gases) • An atom’s ionization energy is the energy needed to remove one of its electrons. • You can think of ionization energy as a reflection of how strongly an atom holds onto its outermost electron. • Atoms with high ionization energies hold onto their electrons very tightly, whereas atoms with low ionization energies are more likely to lose one or more of their outermost electrons and gain a positive charge. • Ionization energies generally increase as you move from left to right across a period. • Metals have low ionization energies while the noble gases have the highest ionization energies. • There will be peaks and valleys due to filled sublevels or partially-filled sublevels. • Ionization energies generally decrease as you move down a group. • The energy required to remove the first electron from an isolated atom is called the first ionization energy. • The successive ionization energies are the energies required to remove electrons beyond the first electron. • For each element you can find one very large increase between a different pair of ionization energies. • This happens whenever the inner core of electrons are attacked. Electronegativity • An atom’s electronegativity reflects its ability to attract electrons in a chemical bond. • Electronegativity increases as you go from left to right across the period. • Electronegativity decreases as you go down a group. – Metals have low electronegativity while nonmetals have high electronegativity. – Fluorine has the highest and francium has the lowest. Shielding Effect • Shielding effect is the result of full energy levels separating the outermost electrons from the nucleus. • An element like lithium only has one energy level separating its outer electron from the nucleus. (therefore low shielding effect) • But, an element like francium has many energy levels separating its outer electrons from the nucleus. (therefore high shielding effect) • This shielding effect is the reason why ionization energy and electronegativity decreases as you go down a group. • Shielding effect increases as you go down a group. • Shielding effect remains the same as you go across a period.