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PERIODIC TABLE

PERIODIC TABLE. The most awesome chemistry tool ever!. Unit Objectives:. Understand the history of who built up the periodic table, how they did it, and what law was made

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PERIODIC TABLE

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  1. PERIODIC TABLE The most awesome chemistry tool ever!

  2. Unit Objectives: • Understand the history of who built up the periodic table, how they did it, and what law was made • Become familiar with structure of periodic table, how the e- is related to the structure of the table, and properties of chemicals based on their location • Discover trends related to electron configuration and periodic properties; use electron configuration to predict location on periodic table.

  3. History • Explainthe roles of Mendeleev and Moseley in the development of the periodic table. • Describe the modern periodic table. • Explain how the periodic law can be used to predict the physical and chemical properties of elements. • Describe how the elements belonging to a group of the periodic table are interrelated in terms of atomic number.

  4. Mendeleev and Periodicity • In the late 1800s, scientist Mendeleev noticed: • when the elements were arranged in order of increasing atomic mass, certain similarities in their chemical properties appeared at regular intervals and formed a pattern. • Repeating patterns are referred to as periodic. • Something that occurs periodically occurs at regular/fixed intervals

  5. Mendeleev and Periodicity • Not all elements had been discovered, so when arranging the table by atomic mass, • He left spaces for future elements to be discovered and placed into the table based on their mass • He even predicted what their properties would likely be

  6. Mendeleev and Periodicity • Where Mendeleev’s work left off there were some questions: • Why could most elements be arranged in the order of increasing atomic mass and other not? • What was the reason behind chemical periodicity?

  7. Moseley and the Periodic Law • Moseley continued research on the periodic table and was able to propose an answer to the first question. • Elements follow a clearer pattern when arranged by their nuclear charge (# of protons) instead of their nuclear mass! • Why do you think this is? • Hint: remember Hydrogen-3 and Helium-3 and their similarities and differences?

  8. Moseley and the Periodic Law • A big example of what Moseley did: he realized that Tellurium (which has an atomic mass of about 127.6 amu) should be before Iodine (which has an atomic mass of 126.9 amu) on the periodic table because Tellurium has the atomic number of 52 and Iodine has the atomic number of 53.

  9. Moseley and the Periodic Law • Moseley led to the editing of Mendeleev’s principle of chemical periodicity: • Periodic law: the physical and chemical properties of the elements are periodic functions of their atomic numbers. • (When elements are arranged by increasing atomic number, similar properties occur in elements at regular intervals)

  10. Modern Periodic Table • After all the elements had been discovered (some synthesized as well) our modern periodic table was been established • Periodic table: an arrangement of the elements in order of their atomic numbers so that elements with similar properties fall into the same column or “group”.

  11. Specific Group: Noble Gases • Discovered in the late 1800s, the noble-gases were slowly discovered, often due to indirect observation • Helium was discovered due to its emission spectrum in sunlight • Argon was not-reactive and was deduced when mass of an unknown gas could not be accounted for when all other gases had been reacted out.

  12. Lanthanides and Actinides have their own block • Lanthanides: • The elements between Cerium (Atomic # 58) and Lutenium (Atomic # 71) did not follow the properties of the elements around them • Whereas Lanthanum (atomic # 57) and Hafnium (atomic # 72) do fit in their groups • Actinides: • The elements between atomic #s: 90 and 103 • Many are synthetic and they do not follow the properties of the metals around them in the periodic table

  13. Periodicity • We will begin to explore the similar properties groups of the periodic table exhibit. • Groups to begin thinking about: • Alkali Metals (s-block) • Alkaline Earth Metals (s-block) • Main-group elements (p-block • Halogens (p-block) • Noble Gases (p-block) • Transition metals (d-block)

  14. Electron Configuration and Periodic properties Objectives: • Explainthe relationship between electrons in sublevels (shapes) and period length in the periodic table. • Locate and name the four blocks of the periodic table. Explain the reasons for these names. • Discuss the relationship between group configurations and group numbers. • Describe the locations in the periodic table and the general properties of the alkali metals, the alkaline-earth metals, the halogens, and the noble gases.

  15. Periods and blocks of the periodic table • Periods are the horizontal rows of the periodic table. • Elements are arranged vertically in the periodic table in groups that share similar chemical properties. • The four blocks: the s, p, d, and fblocks • each block corresponds to the electron sublevel being filled in that block.

  16. Periods & Blocks of the Periodic Table

  17. S-block elements • The elements of Group 1 of the periodic table are known as the alkali metals. • lithium, sodium, potassium, rubidium, cesium, and francium • In their pure state, all of the alkali metals have a silvery appearance and are soft enough to cut with a knife. • The elements of Group 2 of the periodic table are called the alkaline-earth metals. • beryllium, magnesium, calcium, strontium, barium, and radium • Group 2 metals are less reactive than the alkali metals, but are still too reactive to be found in nature in pure form.

  18. Hydrogen and Helium • Hydrogen is NOT a metal; • Its properties are very different from the metals in the alkali metal group • Helium fits in with the Noble Gases as it has a full valence shell and is not at reactive • Alkaline earth metals have a full s-sublevel, but their whole valence shell is not full (their p-sublevel is empty) • A valence shell is the highest energy level that electrons occupy. Elements are more stable/less reactive when their valence shells are full of electrons

  19. The d-block elements (Groups 3-12) • Transition metals: the d-block elements are metals with typical metallic properties • Ductility • Malleability • Conductivity • Having luster • Groups and periods within the transition metals have varied properties and it is therefore hard to group metals by individual properties

  20. p-block elements (group 13-18) • Main-group elements: the elements in the p-block of varying properties • All non-metals are in the p-block (aside from hydrogen) • All metalloids (boron, silicon, germanium, arsenic, antimony, and tellurium) are in the p-block

  21. Relationships between Group #, blocks, and electron config.

  22. Halogens • Elements in group 17 are the halogens • Fluorine, chlorine, bromine, iodine, and astatine • They are highly reactive in similar ways • All form salts with group 1 and 2 metals • They are volatile ( gas forms easily)

  23. Metalloids • Metalloids have properties of both metals and non-metals • boron, silicon, germanium, arsenic, antimony, and tellurium • Semiconducting, brittle, harder than most metals, have an opalescent luster • Relatively reactive (Bi only found in elemental form)

  24. f-block elements Lanthanides and Actinides • Lanthanides are shiny metals and have similar reactivity to group 2 metals • Their position reflects the fact that they involve the filling of the 4f sublevel • Actinides: • Only 4 are found in nature (thorium through neptunium)

  25. Alkali Metals Alkaline Earth Metals Other Metals Metalloids Inner transition Metals Other Non-Metals Halogens Noble-Gases

  26. Do Now: • List properties of the following elemental groups: • Table 1: Halogens • Table 2: Noble Gases • Table 3: Alkali Metals • Table 4: Alkaline Earth Metals • Table 5: Transition metals • Table 6: Actinides • Share with your table • Describe what you know about an element by looking at its position in the periodic table. • Table 1: Br • Table 2: Ar • Table 3: K • Table 4: Ca • Table 5: Cu • Table 6: Fe • Identify any noticeable trends.

  27. Do Now If you were absent Friday: • How did Mendeleev organize the periodic table? • Why didn’t this always work? • How did Moseley organize the periodic table? • What are the groups of the periodic table? • Include at least 1 property for each group If you were here Friday: • Please take out your practice problem sheet and complete the problems for Atomic radius and ionization energy using your graphic organizer • Look over your SRQ Answers • Help your neighbor with their do now if they were absent Friday

  28. Atomic Radii (the distance between the nucleus & the end of the cloud) • One half the distance between the nuclei of identical atoms that are bonded together Atomic radii: 200 pm Bond length 400 pm

  29. Atomic Radii • Increases in size from right to left • Increases in size from up to down

  30. Ionization Energy (how to knock an electron off!) • An ion is an atom that • has more or less electrons than it has protons. • is either positively (cation) or negatively (anion) charged • Ionization is anything that gives matter a charge • Ionization Energy (IE) is the amount of energy needed to remove an electron from an atom

  31. Ionization Energy Continued • Ionization often occurs because • The valence electrons are more stable in the ionic configuration • To complete a valence shell • Gaining an electron can cause a neutral atom to lose or gain energy depending on the atom • A metal gaining an electron would absorb energy and become less stable • A non-metal gaining an electron would release energy and become more stable

  32. Ionization Energy (how to knock an electron off!) • Ionization energy trends: • Increases from left to right • Increases from down to up

  33. Additional Ionization energies • Removing a second electron is the 2nd Ionization Energy, and it typically greater than the first ionization energy • With each additional electron removed, the Ionization Energy will increase.

  34. Electron Affinity • The change in energy associated with the addition of an electron to a neutral atom • Green means no electron affinity • Yellow means high electron affinity (trends by groups)

  35. Ionic Radii • Positive ion = cation • Negative ion = anion

  36. Ionic Radii

  37. Valence Electrons • Valence electrons are available to be lost, gained, or shared during chemical reactions and compound formation.

  38. Electronegativity • Electronegativity how much an atom in a compound attracts electrons from another atom in the compound • (how much the atom is greedy for electrons)

  39. Periodic properties of thed-block and f-block • Atomic radii decrease from left to right • Ionization energy increases from left to right • Electronegativity slightly increase from left to right

  40. -1 -2 -3 +1 +2 +3 Common Elemental Ions 8.2

  41. Why? • A lot of these trends are related to electron shielding • Electron shielding is how the inner electrons effect how the outer electrons act due to charge different and repulsion

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