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Enriched Chemistry Chapter 5 – The Periodic Law Geneva High School Section 5-1 History of the Periodic Table • By 1860, more than 60 elements had been discovered, yet chemists had no way of organizing them or agreeing on basic properties of the element. Mendeleev’s periodic table grouped elements by their properties. • Mendeleev placed the name of each known element on a card, with the atomic mass and the observed phys/chem properties • He then arranged them in various configurations, looking for patterns or trends • He noticed that when the elements were arranged in order of increasing atomic mass, certain similarities appeared at regular intervals. • He published his periodic table in 1869 • His P.T. left several empty spaces, but he predicted the properties of the elements that would fill three of these spaces. • By 1886, all three elements had been discovered. • The success of his predictions persuaded most chemists to accept his periodic table and earned him credit as the discoverer of the periodic table. • periodic – a regular, repeating pattern Eka-aluminum (now known as gallium) Three truths and a lie about Mendeleev (You do not need to copy this!) Which of the following is NOT true? A. Mendeleev only cut his hair once a year. B. He had two cats, which he named after the first two elements he discovered. C. He was in charge of measuring the amount of alcohol in vodka in Russia. D. There is a crater named after him on the moon. Henry Mosely arranged elements by their atomic numbers. Mosely discovered that the elements fit better when arranged in increasing order by their atomic number What elements does this effect? Mendeleev’s principle is known as the periodic law: the physical and chemical properties of the elements are periodic functions of their atomic numbers. This means that their properties repeat themselves based on their atomic numbers Modern periodic tables arrange the elements by both atomic number and properties. • The P.T. has changed over time, and new elements have been discovered and synthesized • synthesized – man made • However, they can all be placed in a group of other elements with similar properties • The periodic table is an arrangement of the elements in order of their atomic numbers so that elements with similar properties fall in the same column, or group. Periodic Table Structure Rows are called periods Columns are called groups or families Significant periodic table discoveries The Noble Gases • In the late 1800s, the noble gases were discovered and placed in their own group • They had escaped detection because of their lack of reactivity The Lanthanides and Actinides • Lanthanides - 14 elements (58-71) that are all very similar in physical and chemical properties which made them difficult to isolate • Actinides – 14 elements (90 – 103) • To save space, they are set off below the main portion of the periodic table Periodicity • The differences between the atomic numbers of successive elements in a group are 8, 8, 18, 18, 32 • The reason for periodicity is explained by the arrangement of the electrons around the nucleus. Chapter 5 Section 2 – Electron Configuration and the Periodic Table (aka Groups of the Periodic Table) Generally, the electron configuration of an atom’s highest occupied energy level governs the atom’s chemical properties. Sodium atom – notice This little guy is the only electron that matters in terms of reactivity of Francium! how much closer the outer shell electron is to the positively charged nucleus. Main Idea – The period of an element is determined by its electron configuration. • Elements are arranged vertically in groups (or families) that share similar chemical props • They are also organized into rows, or periods • The period of an element can be determined from the element’s electron configuration • ex: [Ar]3d104s24p3 • The P.T. can be divided into four blocks, the s, p, d, and f blocks The s-Block Elements: Groups 1 and 2 • Elements of the s-block are chemically reactive • Group 1 elements are known as the alkali metals • • • • 1 valence electron (v.e.) silvery appearance and soft never found in nature as free elements combine vigorously with nonmetals • Group 2 elements are known as the alkaline earth metals • • • • 2 v.e. harder, denser, and stronger than the alkali metals less reactive than alkali metals never found as free elements https://www.youtube.com/watch?v=cqeVEFFzz7E Hydrogen and Helium • despite having a 1s1 config, hydrogen does not share the same properties as group 1 elements • helium is non-reactive because its outer energy is filled with 2 electrons The d-Block elements: groups 3-12 (aka the Transition Metals) • The d-block elements are metals w/typical metallic props and are often referred to as transition elements • good conductors of electricity and have a high luster • less reactive than groups 1 and 2 The p-Block elements: groups 13-18 • The p-block elements together with the s-block elements are called the main-group elements • props of p-block elements vary greatly • includes all of the nonmetals except H and He, and the six metalloids • The families are known by the first element in each family, so the Boron Family, Carbon Family, Nitrogen Family and Oxygen Family. • The elements of Group 17 are known as the halogens • most reactive nonmetals, reacting vigorously with most metals to form salts • this is due to their 7 v.e., one electron short of the stable noble-gas configuration • rarely found as free elements in nature • Fluorine is the most reactive nonmetal. Metalloids • Also called semiconducting elements • Mostly brittle solids with some properties of metals and some of nonmetals • They will conduct electricity, but a higher voltage is necessary than with true conductors • Used in transistors, LEDs, photovoltaic cells Chapter 5 Section 3 – Electron Configuration and Periodic Properties Main Idea – Atomic radii are related to electron configuration. • The boundaries of an atom are fuzzy, making it difficult to describe its size. • One way to express an atom’s radius is to measure the distance between the nuclei of two identical atoms that are bonded together and then divide by two, giving the atomic radius. Atomic radius – defined as one-half the distance between the nuclei of identical atoms that are bonded together. Atomic radii…( cont) • Periodic trends Atomic radii decrease as you go across a period. This is caused by the increasing charge of the nucleus, which gains a proton from one element to the next. The attraction of the nucleus is somewhat offset by repulsion among the increased number of electrons. • Group trends The atomic radii of the maingroup elements increase down a group. • Main idea – Removing electrons from atoms to form ions requires energy. A + energy A+ + eElement A+ represents an ion, which is an atom or group of bonded atoms that has a positive or negative charge. Any process that results in the formation of an ion is referred to as ionization. To compare the ease with which atoms of different elements give up electrons, chemists compare ionization energies. the energy required to remove one electron from a neutral atom. Ionization Energy • Period Trends In general, ionization energies of the main-group elements increase across each period. This is caused by an increasing nuclear charge. • Group Trends Among the main-group elements, ionization energies general decrease down the groups. This is because as you go down a group, there are more energy levels, and the electrons are further from the nucleus, making them easier to remove. Also, the electrons between the nucleus and the outer electrons have a shielding effect. • Removing Electrons from Positive Ions Electrons can be removed from positive ions as well as neutral atoms. The energies for removal of additional electrons from an atom are referred to as the second ionization energy, third ionization, and so on. The energy required is always greater as you remove more electrons. Each successive electron removed from an ion feels an increasingly stronger nuclear charge. The biggest ionization energies occur when an ion assumes a noblegas configuration. (see figure 3.6 on page 147) Main idea – Adding electrons to atoms to form ions also requires energy. • The energy change that occurs when an electron is acquired by a neutral atom is called the atom’s electron affinity. • Most atoms release energy when they acquire an electron. A+ e- A- + energy • On the other hand, some atoms must be “forced” to gain an electron by the addition of energy. A + e -- + energy A– The vast majority of elements release energy when they acquire an electron. Main Idea – When atoms become ions, their radii change • The formation of a cation by the loss of one or more electrons always leads to a decrease in atomic radius. • The formation of an anion by the addition of one or more electrons always leads to an increase in atomic radius. Main Idea – Only the outer electrons are involved in forming compounds. • Chemical compounds form because electrons are lost, gained, or shared between atoms. • The electrons available to be lost, gained or shared in the formation of chemical compounds are called valence electrons. Group number valence electrons 1 1 2 2 13 3 14 4 15 5 16 6 17 7 18 8 Main Idea – Atoms have different abilities to capture electrons. • Linus Pauling, one of America’s most famous chemists, devised a scale that reflects the tendency of an atom to attract electrons. • Electronegativity is a measure of the ability of an atom in a chemical compound to attract electrons from another atom in the compound. • The most electronegative element, fluorine, is arbitrarily assigned a value of 4. • This concept is vitally important in many biochemical reactions, such as photosynthesis and aerobic respiration. • Period trends – Electronegativities tend to increase across each period. • Group trends – electronegativities tend to either decrease down a group or remain about the same.