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Ch 7. The Periodic Table. Objectives. SWBAT Define the periodic properties Apply periodic trends to the periodic table. The Periodic Table. J.W. Dobereiner The elements in the triad has similar chemical properties.
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Ch 7 The Periodic Table
Objectives SWBAT Define the periodic properties Apply periodic trends to the periodic table
The Periodic Table • J.W. Dobereiner • The elements in the triad has similar chemical properties. • Several elements could be classified into groups of three called triads. • Jar Newlands • Arranged the elements in order of increasing mass.
The Periodic Table • Dmitri Mendeleev • Organized the elements into the Periodic Table • He started by sorting elements by increasing mass. • He saw a repetition of pattern. • Same column → same properties • Predicted the existence of new elements • Henry Moseley • The atomic number is based on the amount of positive charges in the nucleus.
The Periodic Law • When elements are arranged in order of increasing atomic number, their physical and chemical properties show a periodic pattern. • Elements with similar properties are in vertical columns called Groups or Families. • Horizontal rows are called Periods.
European notation American system IUPAC (Inter. Union of Pure & Applied Chemistry) The elements in a group have similar properties because they have valence electrons in similar configurations Roman numerals IA, IIA, IIIA 1-18 (no letters) Labeling
Valence Electrons • Elements in the same column contain the same number of valence electrons (electrons in their outer-shell orbitals).
Atomic Radius (in a group) • Measure the atomic radius from the center of the nucleus to the outermost electron. Atom size increases going down a group. • Reason: The principal quantum number of the outermost (valence) electrons increases.
Atomic Radius (in a period) Atomic size decreases going across the period. • Reason: Same principal quantum number going across the period. More protons are added going across the period. The protons have a stronger pull on the electrons. The strong attractive forces between the protons and the outermost (valence) electrons shrinks the orbitals and makes the atoms smaller. Generally speaking, effective nuclear charge is the charge felt by the valence electrons after you have taken into account the number of shielding electrons that surround the nucleus. Effective nuclear charge increases (which decreases shielding).
Atomic Radius • http://www.shodor.org/chemviz/ionization/students/background.html
Ionic Size (when you create an ion) • An ion is created when an atom gains or loses an electron. • LOSE AN ELECTRON (create a positive ion) • Size becomes smaller • Reason: Loss of an electron vacates outer orbitals and reduces the repulsive forces between electrons. • GAIN AN ELECTRON (create a negative ion) • Size becomes larger • Reason: There are a greater number of electrons. There is a greater repulsion force between electrons. ***Elements in a group form ions of the same charge*** • Left side of periodic table forms POSITIVE ions. • Right side of the periodic table form NEGATIVE ions.
Ionization Energy • This is the energy needed to remove the outermost electron of an atom. Li(g) → Li+(g)+ e- ionization energy 8.64 x 10-19 J/atom • HIGH ionization energy means the atom hold onto the electron tightly. • LOW ionization energy means the atom holds onto the electron loosely. • Since an atom is very small, scientists use a larger unit of measure called the mole. Therefore, ionization energy is measured in J/mol
Ionization Energy Trend • Ionization energy decreases as you move down a group. Reason: the electrons being removed are, in general, farther from the nucleus. As "n" increases, the size of the orbital increases, and the electron is easier to remove. (NOTE: This trend is the opposite of the atomic radius trend)
Ionization Energy Trend Ionization energy increases as you move from left to right on the periodic table. Reason: electrons added in the same principal quantum level do not completely shield the increasing nuclear charge caused by the added protons. The electrons in the same principal quantum level are generally more strongly bound when moving left to right across the periodic table
Ionization Energy • http://www.shodor.org/chemviz/ionization/students/background.html
Successive IE’s • The energy required to remove a second electron from an atom is called its second IE, and so on…
Electron Affinity (attraction) • This is the energy change that occurs when a gaseous atom gains an extra electron. Ne(g) + e- → Ne-1(g) electron affinity= 29 kJ/mole • If the electron affinity is a negative number, the atom releases energy. • Normally, non-metals have a more negative electron affinity than metals. The exception is the noble gases.
Electron Affinity • Electron affinity generally becomes increasingly negative moving from left to right. • (exception: the addition of an electron to a noble gas would require the electron to reside in a new, higher-energy subshell. Occupying a higher-energy subshell is energetically unfavorable, so the electron affinity is positive, meaning that the ion will not form)
Electron Affinity • Electron affinity does not change greatly as we move down a group. Electron affinity should become more positive (less energy released). Reason: Moving down a group the average distance between the added electron and the nucleus steadily increases, causing the electron-nucleus attraction to decrease. The orbital that holds the outermost electron is increasingly spread out, however, proceeding down the group, reduces the electron-electron repulsions. A lower electron-nucleus attraction is thus counterbalanced by lower electron-electron repulsions.
Electronegativity • This is the ability for an atom to attract an electron in a chemical bond. • The scale ranges from 0.7 to 4.0 • There are no units for this number.
METALS/NON-METALS/METALLOIDS • The more an element exhibits the physical and chemical properties of metals, the greater its metallic character. • Metallic character generally increases going down a column and decreases going from left to right across a period
METALS • Metals conduct heat and electricity. They are malleable (can be pounded into thin sheets) and ductile (can be drawn into wire). • Metals are solids at room temperature except Hg (which is a liquid)
METALS • Metals tend to have low ionization energies and are consequently oxidized (lose electrons) when they undergo chemical reaction. • Many transition metals have the ability to form more than one positive ion.
Reactions with Metals • Metal oxide + water → metal hydroxide • Metal oxide + acid → salt + water • These reactions will be helpful to answer your Reaction Prediction questions.
Try Writing the Reaction Sodium oxide is reacted with water. Barium oxide is reacted with HCl.
Reaction Answers Sodium oxide is reacted with water. Na2O + H2O NaOH Barium oxide is reacted with HCl. BaO + HCl BaCl2 + H2O
Non-Metals • Non-metals are not lustrous and are poor conductors of heat and electricity. • Non-metals commonly gain enough electrons to fill their outer p sub-shell completely, giving a noble gas electron configuration. • Compounds composed entirely of nonmetals are called molecular substances.
Non-Metal Reactions Nonmetal oxide + water → acid Nonmetal oxide + base → salt + water • These reactions will be helpful to answer your Reaction Prediction questions.
Try these Reactions. Carbon dioxide reacts with water. Tetraphosphorous decaoxide reacts with water. Sulfur trioxide reacts with barium hydroxide. Tetraphosphorous decaoxide reacts with sodium hydroxide.
Reaction Answers Nonmetal oxide + water → acid CO2 (g) + H2O(l) ---> H2CO3 (aq) P4O10 (s) + 6 H2O(l) ---> 4 H3PO4 (aq) Nonmetal oxide + base → salt + water SO3(g) + Ba(OH)2(aq) → BaSO4(aq) + H2O(l) P4O10(s) + 12 NaOH(aq) → 4 Na3PO4(aq) + 6 H2O(l)
Metalloids • Metalloids have properties between those of metals and nonmetals. • Metalloids are also called semi-metals.
Group Names • Know the group (family) names: • Alkali metals • Alkaline earth metals • Halogens • Noble gases
Alkali Metals • Soft • Metallic luster • High thermal and electrical conductivity • Low densities • Low melting points • Exist in nature as compounds
Alkaline Earth Metals • Solids • Harder than alkali metals • More dense than alkali metals • Higher melting points than alkali metals
Homework Problems • 4, 5, 9, 11, 13-17, 20, 23, 24, 27, 28, 32, 33, 35, 38, 40, 39, 41, 43, 45, 49, 54 • Do 40 before 39