Download to read offline
More Related Content
Similar to Periodic-Trends-Dacillo-Bron-Lacson.pptx
Periodic-Trends-Dacillo-Bron-Lacson.pptx
- 1. PERIODIC TRENDS Prepared by:Maureen Bron & Lorraine Lacson
- 2. PERIODIC TRENDS ■ Sincethe atomic number directly represents the number of protons, the properties of elements within a group and across a period will vary. The properties of elements that are affected are metallic and nonmetallic properties, atomic radius, ionization energy, electron affinity, ionic size, and electronegativity. ■ Periodic trends, arising from the arrangement of the periodic table, provide chemists with an invaluable tool to quickly predict an element's properties. These trends exist because of the similar atomic structure of the elements within their respective group families or periods, and because of the periodic nature of the elements.
- 4. METALLIC PROPERTY ■ Elementswith three or less electrons in the outer energy level are classified as metals, while elements with live or more electrons in the outer energy level are classified as nonmetals. Elements adjacent to the ladder like line exhibit some of the properties of both metals and nonmetals. These elements are called metalloids. Metals have high densities, high melting temperatures, luster, and are good conductors of heat and electricity. Metals are usually solid and their oxides are often soluble in acidic solutions. They can be shaped into wires (ductilel or hammered into flat shapes (malleable). All metals are temperature except mercury, which is liquid.
- 5. ■ Nonmetals areoften in gaseous or liquid form. They have low densities, low melting temperatures, and are poor conductors of heat and electricity. Their oxides are usually soluble in basic solutions. Nonmetals are not very shiny, malleable, or ductile. Some common examples of nonmetals are C. N. CI. S. and P.
- 7. •Metallic character refersto the level of reactivity of a metal. •Non-metallic character relates to the tendency to accept electrons during chemical reactions. •Move left across period and down the group: increase metallic character (heading towards alkali and alkaline metals) •Move right across period and up the group: decrease metallic character (heading towards nonmetals like noble gases)
- 9. Atomic Radius ■ Atomicradius or Atomic Radii is the total distance from the nucleus of an atom to the outermost orbital of its electron. ■ Unlike a ball, an atom has fuzzy edges. ■ The radius of an atom can only be found by measuring the distance between the nuclei of two touching atoms, and then dividing that distance by two.
- 11. Period: in general,as we go across a period from left to right, the atomic radius decreases –Effective nuclear charge increases, therefore the valence electrons are drawn closer to the nucleus, decreasing the size of the atom Family: in general, as we down a group from top to bottom, the atomic radius increases –Orbital sizes increase in successive principal quantum levels
- 13. CONCEPT CHECK Which shouldbe the larger atom? Why? ■ O or N ■ K or Ca ■ Cl or F ■ Be or Na ■ Li or Mg
- 14. Ionization energies ■ Whena neutral atom loses or gains an electron, the product particle is an ion, and the gain or loss process is termed ionization. ■ Energy is required to remove one or more electrons from a neutral atom, and this is referred to as ionization energy or ionization potential.
- 15. GROUP TREND: Invertical groups, ionization energy decreases from top to bottom. This is because electrons are farther from the nucleus & filled levels cause a shielding effect. SHIELDING EFFECT: Inner electrons shield outer electrons from the positive nucleus. This means outer electrons are not held as tightly. PERIOD TREND: Ionization energy tends to increase as you move from left to right toward the noble gases. This is because metals tend to lose electrons & nonmetals tend to gain electrons. All of them want to as stable as the noble gases.
- 17. ■ The smallerthe atomic radius, the more closely the electrons are held by the nucleus, and thus the higher the ionization energy. ■ So, in general, ionization energy increases from the bottom-left to the upper-right of the periodic table.
- 19. The minimum amountof energy required to dislodge the least firmly attached electron from an atom in the gaseous state is called the first ionization energy, the second ionization energy is the amount of energy required to remove the second electron from the gaseous atom; and so forth.
- 20. Which atom wouldbe harder to remove an electron from?
- 21. IONIZATION ENERGY •X +energy → X+ + e– X(g) → X+ (g) + e– • First electron removed is IE1 • Can remove more than one electron (IE2 , IE3 ,…) Ex. Magnesium Mg → Mg+ + e– IE1 = 735 kJ/mol Mg+ → Mg2+ + e– IE2 = 1445 kJ/mol Mg2+ → Mg3+ + e– IE3 = 7730 kJ/mol* *Core electrons are bound much more tightly than valence electrons
- 23. Problems The following seriesof problems reviews general understanding of the aforementioned material. 1. Based on the periodic trends for ionization energy, which element has the highest ionization energy? A. Fluorine (F) B. Nitrogen (N) C. Helium (He) 2.) Nitrogen has a larger atomic radius than oxygen. A.) True B.) False 3.) Which has more metallic character, Lead (Pb) or Tin (Sn)?