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Organizing the Elements - In a self-service store, the products are grouped according to similar characteristics. With a logical classification system, finding and comparing products is easy. You will learn how elements are arranged in the periodic table and what that arrangement reveals about the elements.

Searching For an Organizing Principle How did chemists begin to organize the known elements?

● Chemists used the _____________ of elements to sort them into groups.

● Chlorine, bromine, and iodine have very similar _____________ properties. Mendeleev’s

Periodic Table - How did Mendeleev organize his periodic table?

● Mendeleev arranged the elements in his periodic table in order of increasing __________________________.

● The periodic table can be used to predict the _____________ of undiscovered elements.

In 1913, Henry Moseley used X-ray experiments to determine the atomic numbers of elements, proving they were more important than atomic mass. Before Moseley, elements were arranged mainly by increasing atomic mass, which sometimes caused errors in their placement. Moseley proved that elements should be organized by atomic number, leading to a more accurate periodic table. His work resolved issues with elements like iodine and xenon and established the modern periodic law, which organizes elements by increasing atomic number. Moseley’s contributions were essential to the accurate periodic table we use today.

The Periodic Law - How is the modern periodic table organized?

● In the modern periodic table, elements are arranged in order of increasing _________________________.

The periodic law: When elements are arranged in order of increasing atomic number, there is a periodic repetition of their physical and chemical properties.

● The properties of the elements within a period change as you move across a period from left to right.

● The pattern of properties within a period repeats as you move from one period to the next.

What are three broad classes of elements?

Three classes of elements are metals, nonmetals, and metalloids.

Across a period, the properties of elements become less metallic and more nonmetallic.

Metals, Metalloids, and Nonmetals in the Periodic Table Metals

● Metals are good conductors of heat and electricity

● In general, nonmetals are poor conductors of heat and electricity

● 80% of elements are metals.

● Most nonmetals are gases at room temperature

● Metals have a high luster, are ductile, and are malleable.

● A few nonmetals are solids, such as sulfur and phosphorus. ●

One nonmetal, bromine, is a dark-red liquid.

Metalloids

● A metalloid generally has properties that are similar to those of metals and nonmetals.

● The behavior of a metalloid can be controlled by changing conditions.

Each square in a periodic table also contains information. You will learn what types of information are usually listed in a periodic table.

Squares in the Periodic Table - What type of information can be displayed in a periodic table? The periodic table displays the symbols and names of the elements, along with information about the structure of their atoms. The background colors in the squares are used to distinguish groups of elements. ● The Group 1A elements are called alkali metals. ● The Group 2A elements are called alkaline earth metals. ● The nonmetals of Group 7A are called halogens. Electron Configurations in Groups ● How can elements be classified based on their electron configurations? ● Elements can be sorted into noble gases, representative elements, transition metals, or inner transition metals based on their electron configurations. The Noble Gases The noble gases are the elements in Group 8A of the periodic table. The electron configurations for the first four noble gases in Group 8A have full s and p orbitals. The Representative Elements Elements in groups 1A through 7A are often referred to as representative elements because they display a wide range of physical and chemical properties. ● The s and p sublevels of the highest occupied energy level are not filled. ● The group number equals the number of electrons in the highest occupied energy level. ● In atoms of the Group 1A elements, there is only one electron in the highest occupied energy level. In atoms of the Group 4A elements, there are four electrons in the highest occupied energy level. Transition Elements There are two types of transition elements—transition metals and inner transition metals. They are classified based on their electron configurations. In atoms of a transition metal, the highest occupied s sublevel and a nearby d sublevel contain electrons. In atoms of an inner transition metal, the highest occupied s sublevel and a nearby f sublevel generally contain electrons. 2 6.3 Periodic Trends The periodic table was designed around the electron structure of atoms, the last accupied sublevel with electrons, the # of valence electrons, information on elements with similar properties, and the activities of elements. (This refers to how easily an element gains or loses electrons) Most active metal = Francium Most active Nonmetal = Fluorine Periodicity in Properties Coulombic Forces: Many properties are based on the attractive forces between (+) and (-) parts of the atom and the distance that separates these charges. 1. Coulombic Attractions a. A force of attraction exists between bodies that have opposite charge i. As the amount of charge increases, the force of attraction ii.As the distance of separation of charge increases decreases, the force of attraction increases. Shielding Effect The shielding effect is caused by the repulsion between kernel electrons and valence electrons. 1. Kernel Electrons: The electrons not in the highest energy level 2. Valence Electrons: The electrons in the highest energy level 3. Shielding: The electrons that lie between the nucleus and the electrons in highest energy level shield the outer electrons from the effect of the nuclear charge. Trends in Atomic Size What are the trends among the elements for atomic size? ● The atomic radius is one half of the distance between the nuclei of two atoms of the same element when the atoms are joined. Group and Periodic Trends in Atomic Size In general, atomic size increases from top to bottom within a group and decreases from left to right across a period. Trends in Atomic Size Atomic Radius ● Trend of Atomic Radius o Going down a group Atomic Radius INCREASES ▪ Reasons: ● Larger distance from nucleus to valence electrons ● Shielding is increased o Going across a period Atomic Radius DECREASES ▪ Reasons: ● Greater nuclear charges 3 Covalent Atomic Radius = ½ the Van der Waals Radius = ½ the Atomic radius in metals = ½ the distance between the nuclei of a distance between the nuclei of distance between two nuclei in a identical atoms that have formed no crystalline metallic form of the covalent bond. In the Cl2 molecule, bond. Two Helium atoms are shown element. Two Magnesium atoms are the covalent atomic radius of chlorine shown here. (160pm) is half the distance between the nuclei here. (31pm) of the bonded hydrogen atoms. (198pm / 2 = 99pm) Ions Some compounds are composed of particles called ions. ● An ion is an atom or group of atoms that has a positive or negative charge. ● A cation is an ion with a positive charge. ● An anion is an ion with a negative charge. How do ions form? - Positive and negative ions form when electrons are transferred between atoms. Ionic Radius - Ionic Radius is the radius of (+) or (-) ions of an element, (+) ions are Cations, (-) ions are Anions Neutral atoms radii > (+) ions radii Mg = 0.136 nm Mg+2 = 0.065 nm (-) ions radii > Neutral atoms radii N-3 = 0.171 nm N = 0.070 nm Trends in Ionic Size During reactions between metals and nonmetals, metal atoms tend to lose electrons, and nonmetal atoms tend to gain electrons. The transfer has a predictable effect on the size of the ions that form. Cations are always smaller than the atoms from which they form. Anions are always larger than the atoms from which they form. Trends in Ionization Energy The energy required to remove an electron from an atom is called ionization energy. ● The energy required to remove the first electron from an atom is called the first ionization energy. ● The energy required to remove an electron from an ion with a 1+ charge is called the second ionization energy. Group and Periodic Trends in Ionization Energy First ionization energy tends to decrease from top to bottom within a group and increase from left to right across a period. Neutral Atom + Ionization energy 🡪 Ion+1 + e1st Ionization energy 🡪 removes 1st electron 2nd Ionization energy 🡪 removes 2nd electron 3rd Ionization energy 🡪 removes 3rd electron For Nitrogen 1st Ionization energy = 1.4 (Kj/mol * 10-3) 2nd Ionization energy = 2.9 (Kj/mol * 10-3) 3rd Ionization energy = 4.6 (Kj/mol * 10-3) 1. Trends: a. Group – Going down a group the ionization energy DECREASES i. Reasons: 1. Atoms are larger, more shielding of intervening electrons 2. Distance b. Period – Going across a period the ionization energy INCREASES i. Reasons: 1. Greater nuclear charge, but atoms are no larger in size 2. No additional shielding of energy levels 4 5 Trends in Electronegativity Electronegativity - Electronegativity is the measure of the ability of an atom to attract electrons in a chemical bond In general, electronegativity values decrease from top to bottom within a group. Ionic bonds e- are transferred between atoms Covalent Bonds 1 pair of e- is shared between 2 atoms Nonpolar = e- shared evenly H : H Nonpolar covalent molecule Polar = e- are not shared evenly C -:-> O Polar covalent Molecule Trend of Electronegativity a. Going down a group Electronegativity DECREASES i. Reasons: 1. Larger distance from nucleus to electron 2. Shielding is increased b. Going across a period Electronegativity INCREASES i. Reasons: 1. Greater nuclear charges 0 If the electronegativity difference is: Nonpolar - Covalent 0.3 Polar - Covalent 1.7 Ionic . 🡪 |-------------------------------------|---------------------------------|---------------------------------| Classify the following compounds as Covalent or Ionic bonds: a. HCl b. H2O c. Cl2 d. NaCl Summary of Trends What is the underlying cause of periodic trends? 6 Periodicity Review 1. 2. 3. 4. 5. 6. 7. 8. 9. 10. 11. 12. 13. 14. 15. 16. 17. 18. 19. 20. 21. Mendeleev Mosley Periodic Law Horizontal Rows are called Vertical Columns are called Period numbers tell the # of __________________ for all elements within the period Group numbers tell the number of _____________ for all elements within the group Alkali Metals 22. Semimetals = Metalloids Alkaline Earth Metals 23. Shielding Effect Transition Metals 24. Atomic Radius Semimetals (metalloids) 25. Ionization Energy Nonmetals 26. Electron Affinity Halogens 27. Ionic Radius Noble Gases a. Which is the biggest N-3, O-2, or F-1? Lanthanide Series b. Which is the smallest Mg+2 or Na+1? Actinide Series 28. Electronegativity Most Active Metal 29. Covalent Bonds Most Active Nonmetal 30. Ionic bonds Isoelectronic Species 31. Nonpolar Properties of Metals 32. Polar Properties of Nonmetals Nonpolar – Covalent Polar – Covalent Ionic 0 0.3 1.7 🡪 |-----------------------------------|------------------------------|-----------------| Does this property increase or decrease going down a Group? Does this property increase or decrease going across a Period? Reason Atomic Radius # of Valence Electrons # of Energy Levels Ionic Radius Ionization Energy Electronegativity Metallic Properties Nonmetallic Properties 7