1. Ch 7 The Periodic Table The Periodic Table

2. Objectives SWBAT SWBAT Define the periodic properties Define the periodic properties Apply periodic trends to the periodic table Apply periodic trends to the periodic table

3. The Periodic Table J.W. Dobereiner J.W. Dobereiner The elements in the triad has similar chemical properties. The elements in the triad has similar chemical properties. Several elements could be classified into groups of three called triads. Several elements could be classified into groups of three called triads. Jar Newlands Jar Newlands Arranged the elements in order of increasing mass. Arranged the elements in order of increasing mass.

4. The Periodic Table Dmitri Mendeleev Dmitri Mendeleev Organized the elements into the Periodic Table Organized the elements into the Periodic Table He started by sorting elements by increasing mass. He started by sorting elements by increasing mass. He saw a repetition of pattern. He saw a repetition of pattern. Same column → same properties Same column → same properties Predicted the existence of new elements Predicted the existence of new elements Henry Moseley Henry Moseley The atomic number is based on the amount of positive charges in the nucleus. The atomic number is based on the amount of positive charges in the nucleus.

5. The Periodic Law o When elements are arranged in order of increasing atomic number, their physical and chemical properties show a periodic pattern. o Elements with similar properties are in vertical columns called Groups or Families. o Horizontal rows are called Periods.

6. Labeling  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)

7. Valence Electrons Elements in the same column contain the same number of valence electrons (electrons in their outer-shell orbitals). Elements in the same column contain the same number of valence electrons (electrons in their outer-shell orbitals).

8. Atomic Radius (in a group) Measure the atomic radius from the center of the nucleus to the outermost electron. 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. Reason: The principal quantum number of the outermost (valence) electrons increases.

9. Atomic Radius (in a period) Atomic size decreases going across the period. Reason: Same principal quantum number 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).

10. Atomic Radius http://www.shodor.org/chemviz/ionization/students/background.html http://www.shodor.org/chemviz/ionization/students/background.html

11. Ionic Size (when you create an ion) An ion is created when an atom gains or loses an electron. An ion is created when an atom gains or loses an electron. 1) 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. 2) 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.

12. Ionization Energy This is the energy needed to remove the outermost electron of an atom. 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 1) HIGH ionization energy means the atom hold onto the electron tightly. 2) 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

13. Ionization Energy Trend 1) 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)

14. Ionization Energy Trend 1) 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

16. Ionization Energy http://www.shodor.org/chemviz/ionization/students/background.html http://www.shodor.org/chemviz/ionization/students/background.html

17. Successive IE’s The energy required to remove a second electron from an atom is called its second IE, and so on… The energy required to remove a second electron from an atom is called its second IE, and so on…

18. Electron Affinity (attraction) This is the energy change that occurs when a gaseous atom gains an extra electron. 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. 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. Normally, non-metals have a more negative electron affinity than metals. The exception is the noble gases.

19. Electron Affinity Electron affinity generally becomes increasingly negative moving from left to right. 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) (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)

20. Electron Affinity Electron affinity does not change greatly as we move down a group. Electron affinity should become more positive (less energy released). 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.

21. http://www.shodor.org/chemviz/ionization/students/background.html http://www.shodor.org/chemviz/ionization/students/background.html

22. Electronegativity o This is the ability for an atom to attract an electron in a chemical bond. o The scale ranges from 0.7 to 4.0 o There are no units for this number.

23. METALS/NON-METALS/METALLOIDS The more an element exhibits the physical and chemical properties of metals, the greater its metallic character. 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 Metallic character generally increases going down a column and decreases going from left to right across a period

24. MetalsNon-Metals Have a shiny luster; various colors, although most are silvery Do not have a luster; various colors Solids are malleable and ductile Solids are usually brittle; some are hard, and some are soft Good conductors of heat and electricity Poor conductors of heat and electricity Most metal oxides are ionic solids that are basic Most non-metallic oxides are molecular substances that form acidic solutions Tend for form cations in aqueous solutions Tend to form anions or oxyanions in aqueous solution

25. METALS Metals conduct heat and electricity. They are malleable (can be pounded into thin sheets) and ductile (can be drawn into wire). 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 are solids at room temperature except Hg (which is a liquid)

26. METALS Metals tend to have low ionization energies and are consequently oxidized (lose electrons) when they undergo chemical reaction. 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. Many transition metals have the ability to form more than one positive ion.

27. Reactions with Metals Metal oxide + water → metal hydroxide Metal oxide + water → metal hydroxide Metal oxide + acid → salt + water Metal oxide + acid → salt + water These reactions will be helpful to answer your Reaction Prediction questions. These reactions will be helpful to answer your Reaction Prediction questions.

28. Try Writing the Reaction Sodium oxide is reacted with water. Barium oxide is reacted with HCl.

29. Reaction Answers Sodium oxide is reacted with water. Na 2 O + H 2 O  NaOH Barium oxide is reacted with HCl. BaO + HCl  BaCl 2 + H 2 O

30. Non-Metals Non-metals are not lustrous and are poor conductors of heat and electricity. 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. 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. Compounds composed entirely of nonmetals are called molecular substances.

31. Non-Metal Reactions Nonmetal oxide + water → acid Nonmetal oxide + base → salt + water These reactions will be helpful to answer your Reaction Prediction questions. These reactions will be helpful to answer your Reaction Prediction questions.

32. 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.

33. Reaction Answers Nonmetal oxide + water → acid CO 2 (g) + H 2 O(l) ---> H 2 CO 3 (aq) P 4 O 10 (s) + 6 H 2 O(l) ---> 4 H 3 PO 4 (aq) Nonmetal oxide + base → salt + water SO 3(g) + Ba(OH) 2(aq) → BaSO 4(aq) + H 2 O (l) SO 3(g) + Ba(OH) 2(aq) → BaSO 4(aq) + H 2 O (l) P 4 O 10(s) + 12 NaOH (aq) → 4 Na 3 PO 4(aq) + 6 H 2 O (l) P 4 O 10(s) + 12 NaOH (aq) → 4 Na 3 PO 4(aq) + 6 H 2 O (l)

34. Metalloids Metalloids have properties between those of metals and nonmetals. Metalloids have properties between those of metals and nonmetals. Metalloids are also called semi-metals. Metalloids are also called semi-metals.

35. Group Names Know the group (family) names: Know the group (family) names: Alkali metals Alkali metals Alkaline earth metals Alkaline earth metals Halogens Halogens Noble gases Noble gases

36. Alkali Metals Soft Soft Metallic luster Metallic luster High thermal and electrical conductivity High thermal and electrical conductivity Low densities Low densities Low melting points Low melting points Exist in nature as compounds Exist in nature as compounds

37. Alkaline Earth Metals Solids Solids Harder than alkali metals Harder than alkali metals More dense than alkali metals More dense than alkali metals Higher melting points than alkali metals Higher melting points than alkali metals

38. Homework Problems 4, 5, 9, 11, 13-17, 20, 23, 24, 27, 28, 32, 33, 35, 38, 40, 39, 41, 43, 45, 49, 54 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 Do 40 before 39