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Chapter 5

Chapter 5. Periodic Law. History of the Periodic Table. Section 1. Mendeleev and Chemical Periodicity. When the Russian chemist Dmitri Mendeleev heard about the new atomic masses he decided to include the new values in a chemistry textbook that he was writing

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Chapter 5

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  1. Chapter 5 Periodic Law

  2. History of the Periodic Table Section 1

  3. Mendeleev and Chemical Periodicity • When the Russian chemist Dmitri Mendeleev heard about the new atomic masses he decided to include the new values in a chemistry textbook that he was writing • Mendeleev hoped to organize the elements according to their properties

  4. He placed the name of each known element on a card, together with the atomic mass of the element and a list of its observed physical and chemical properties • He then arranged the cards according to various properties and looked for trends or patterns

  5. Mendeleev noticed that when the elements were arranged in order of increasing atomic mass, certain similarities in their chemical properties appeared at regular intervals • Such a repeating pattern is referred to as periodic

  6. Mendeleev’s procedure left several empty spaces in his periodic table • In 1871, he predicted the existence and properties of the elements that would fill three of the spaces • By 1886, all three elements had been discovered • scandium, Sc, gallium, Ga, and germanium, Ge • Their properties are very similar to those predicted by Mendeleev

  7. Success of Mendeleev’s predictions persuaded most chemists to accept his periodic table and earned him credit as the discoverer of the periodic law • Two 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 chemical periodicity?

  8. Moseley and the Periodic Law • In 1911  English scientist Henry Moseley examined the spectra of 38 different metals • Moseley discovered a previously unknown pattern

  9. The elements in the periodic table fit into patterns better when they were arranged in increasing order according to nuclear charge, or the number of protons in the nucleus • Moseley’s work led to both the modern definition of atomic number and the recognition that atomic number, not atomic mass, is the basis for the organization of the periodic table

  10. Electron Configuration and the Periodic Table Section 2

  11. The Modern Periodic Table • “Periodic” - Repeating patterns • Listed in order of increasing number of protons (atomic #) • Properties of elements repeat • Periodic Law- “the physical and chemical properties of the elements are periodic functions of their atomic numbers.”

  12. Periods and Blocks of the Periodic Table • Elements are arranged vertically in the periodic table in groups that share similar chemical properties • They are also organized horizontally in rows, or periods • The length of each period is determined by the number of electrons that can occupy the sublevels being filled in that period

  13. main group elements

  14. Metals • Most solids (Hg is liquid) • Luster – shiny. • Ductile – drawn into thin wires. • Malleable – hammered into sheets. • Conductors of heat and electricity. • Include transition metals – “bridge” between elements on left & right of table

  15. Non-Metals • Properties are generally opposite of metals • Poor conductors of heat and electricity • Low boiling points • Many are gases at room temperature • Solid, non-metals are brittle (break easily) • Chemical properties vary

  16. Metalloids • stair-step pattern • Have properties similar to metals and non-metals • Ability to conduct heat and electricity varies with temp • Better than non-metals but not metals • semiconductors

  17. Group 1 – The Alkali Metals • The elements of Group 1 of the periodic table (lithium, sodium, potassium, rubidium, cesium, and francium) are known as the alkali metals • In their pure state, all of the alkali metals have a silvery appearance and are soft enough to cut with a knife

  18. Because they are so reactive, alkali metals are not found in nature as free elements • They combine strongly with most nonmetals • And they react strongly with water to produce hydrogen gas and aqueous solutions of substances known as alkalis • Because of their extreme reactivity with air or moisture, alkali metals are usually stored in kerosene or oil

  19. Group 2 – The Alkaline-Earth Metals • The elements of Group 2 of the periodic table are called the alkaline-earth metals • Atoms of alkaline-earth metals contain apair of electrons in their outermost s sublevel

  20. Group 2 metals are harder, denser, and stronger than the alkali metals • They also have higher melting points • Less reactive than the alkali metals, but also too reactive to be found in nature as free elements

  21. Transition Elements • Good conductors of electricity and have a high luster • They are typically less reactive than the alkali metals and the alkaline-earth metals • Some are so unreactive that they do not easily form compounds, existing in nature as free elements

  22. Mercury Tungsten Vanadium

  23. uranium Rare Earth Elements • Lanthanide series (period 6) • Actinide Series (period 7) • Some radioactive • Separated from table to make easy to read/print • silver, silvery-white, or gray metals. • Conduct electricity

  24. Halogen Family (“salt-former”) -7 Valence Electrons -most active nonmetals -never found pure in nature -react with alkali metals easily (forms salts) -F most active halogen

  25. bromine Halogens cont… • F compounds in toothpaste • Cl kills bacteria • I keeps thyroid gland working properly

  26. The Noble Gases (Inert Gases) Neon - non-reactive • outermost e- shell is full (8 VE) • In “neon” lights -in earth’s atmosphere (less than 1%)

  27. Section 5.3 Electron Configuration and Periodic Properties

  28. Periodic Trends • In periodic table, there is a DECREASE in atomic radii across the periods from left to right • Caused by increasing positive charge of nucleus (more protons = more positive charge)

  29. Group Trends • Radii of elements decrease as you go UP a group • Electrons occupy sublevels in consecutively higher main energy levels (located further away from nucleus) • In general, the atomic radii of the main-groups elements decrease from the bottom to the top of a group

  30. 2. Ionization Energy • Electrons can be removed from an atom if enough energy is supplied • Using A as a symbol for an atom of ANY element, the process can be expressed as follows: A + energy  A+ + e-

  31. A+ represents an ion of element A with single positive charge (a 1+ion) • Ion an atom or group of bonded atoms that have a positive or negative charge • Ionization  any process that results in the formation of an ion

  32. Period Trends • In general, ionization energies of the main-group elements INCREASE across each period • Caused by increasing nuclear charge • Higher charge more strongly attracts electrons in same energy level

  33. Group Trends • Ionization energies generally INCREASE going UP a group • Electrons going down in group are in higher energy levels, so further away from the nucleus • Removed more easily • Also more electrons between outermost electrons and the nucleus (shields them from attraction to positive nucleus)

  34. What are Valence electrons? • outermost e-’s • Responsible for chem props • Elements in same group… same # of VE • ALL atoms want full outer energy level (usually 8 VE) • To get full outer energy level, some elements: • lose e- (metals) • gain e- (non-metals) • share electrons (some non-metals & metalloids)

  35. Main-group elements – valence e- are in outermost s and p sublevels • Inner e- held too tightly by nucleus to be involved in compound formation

  36. 6. Electronegativity • Valence e- hold atoms together in compounds • In many compounds, negative charge centered around one atom more than another • Uneven distribution of charge has effect on compound’s properties • Useful to have measurement of how strongly one atom attracts e- of another atom

  37. Electronegativity measure of the ability of an atom in a chemical compound to attract electrons • Most e-neg element (fluorine) – randomly assigned value of 4 • Other values calculated in relation to F

  38. Period Trends • e-negs tend to INCREASE across each period • There are exceptions (of course) • Alkali and alkaline-earth metals are least e-neg • In compounds, their atoms have low attraction for e-

  39. Group Trends • Electronegativities tend to INCREASE going UP a group or stay the same • At higher energy levels electrons being added are further away from the nucleus • Therefore, less attraction to the nucleus • Also more electrons between outermost electrons and the nucleus (shields them from attraction to positive nucleus)

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