Periodic Table Trends

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1 Name Date Period Periodic Table Trends (Ionization Energy and Electronegativity) Ionization Energy The required to an electron from a gaseous atom or ion. Period Trend: As the atomic number increases, ionization energy. WHY? increases; Electrons are held more to the nucleus. Also, the outer energy level is closer to being an. Example 1: Circle the element with the higher 1 st ionization energy. Ca or Br Group Trend: As the atomic number increases, ionization energy. WHY? Outer electrons are in energy levels. Farther from the nucleus, which means more ; therefore, the electrons are more removed. Example 2: Circle the element with the higher 1 st ionization energy. Na or Cs Practice 1: Which has a higher ionization energy? Mg or Ba Which has a higher ionization energy? Al or Si Which has a higher ionization energy? Mg or Cl 2 nd Ionization Energy The required to the electron. (There can also be 3 rd, 4 th and so forth ionization energies. Example 3: Aluminum Al(g) Al + (g) + e - I 1 = 580 kj/mol Al + (g) Al 2+ (g) + e - I 2 = 1815 kj/mol Al 2+ (g) Al 3+ (g) + e - I 3 = 2740 kj/mol Al 3+ (g) Al 4+ (g) + e - I 4 = kj/mol Why is there an increase in successive ionization energies?

2 Example 4: Consider atoms with the following electron configurations: 1s 2 2s 2 2p 6 1s 2 2s 2 2p 6 3s 1 1s 2 2s 2 2p 6 3s 2 Which atom has the largest 1 st ionization energy? Which one has the smallest 2 nd ionization energy? Practice 2: The successive ionization energies for an unknown element are I 1 = 786 kj/mol I 2 = 1,577 kj/mol I 3 = 3,232 kj/mol I 4 = 4,355 kj/mol I 5 = 16,091 kj/mol To which family in the periodic table does the unknown element most likely belong? Explain. Electronegativity The ability of an atom to electrons to itself in a chemical bond. Period Trend: As the atomic number increases, electronegativity. WHY? The atoms are getting in size with more protons in the nucleus, so they have more ability to other elements electrons and pull them toward themselves. Example 5: Circle the element with the higher electronegativity. Be or O Group Trend: As the atomic number increases, electronegativity. WHY? Elements at the of a group have electrons held by the nucleus. Elements at the bottom of a group are and have more (and ) between the nucleus and outer/valence electrons. Example 2: Circle the element with the higher electronegativity. N or Sb

3 Name Date Period CHEMICAL BONDING describes the that hold adjacent atoms together in a compound. 3 General Types of Bonds 1. IONIC BONDS Form when one or more electrons are from one atom to another, creating and ions. a. Properties of Ionic Bonds i. High and points. ii. Crystalline ( ) when dried and are. iii. Often. iv. Conduct electric current (electricity) in the form and forms. 2. COVALENT BONDS Involves of valence electrons between atoms. a. Properties of covalent Bonds i. melting and boiling points. ii. electric conductor in any form. iii. Exist as whole, not ions. iv. Most are in water. 3. METALLIC BONDS Form in atoms (positive metal ions with electrons). The force holding the metal together is the electrostatic attraction among and. a. Properties of Metallic Bonds i. Good - Electrons flow in metals, conducting electrical signals. ii. Good - Free flowing electrons transmit. iii. melting and boiling points. iv.,, and.

4 b. Alloys A material composed of or metals. Examples: 1. Brass made of and Karat Gold made of and. 3. Bronze made of and. Chemical reactions result in the, or of valence electrons. So, only the valence electrons are involved in. Lewis Electron Dot Symbols A useful way to represent electrons in the valence shell of an atom. The symbol of the element represents the atomic nucleus together with electrons. electrons are represented by and are placed one-by-one around the element symbol. Draw the Lewis electron dot symbols for each element in period 2. How do you draw Lewis electron dot symbols for IONS? To draw the Lewis electron dot symbol of cations: 1. the same number of electrons as the charge. 2. Draw the symbol for the element with no dots. 3. Place around the structure. 4. Write the of the ion the brackets. Practice 1: Draw the Lewis electron dot structures for these cations. Calcium Ion Aluminum Ion Sodium Ion To draw the Lewis electron dot symbol of anions: 1. the same number of electrons as the charge. 2. Draw the new electron arrangement. 3. Place around the structure. 4. Write the of the ion the brackets. Practice 2: Draw the Lewis electron dot structures for these anions. Nitride Ion Sulfide Ion Bromide Ion

5 IONIC BONDING Ionic Bonds form when one or more valence electrons are from one atom to another, creating and ions. To draw Lewis structures for ionic compounds: 1. Write the correct formula for the compound. 2. Draw the Lewis electron dot symbol for the ion(s). 3. Draw the Lewis electron dot symbol for the ion(s) to the RIGHT of the positive ion. Example 1: Draw the Lewis structure for lithium oxide. Practice 3: Draw the Lewis structure for the following ionic compounds: Sodium phosphide Magnesium nitride COVALENT BONDING Covalent bonding involves the of valence electrons between atoms. One pair of electrons is represented by a dash ( ) A pair of electrons shared ( - ) are represented by a pair of around their atom ( ) Two atoms can share more than one pair of valence electrons Double bonds ( ) and Triple bonds ( ) How to draw Lewis Dot Structures for Covalent Compounds: 1) Determine the arrangement of atoms within a molecule. The central atom is usually the electronegative atom. Hydrogen is a atom because it typically bonds to only one other atom. 2) Determine the total number of electrons in a molecule or ion. In a neutral molecule, this number will be the sum of the valence electrons for each atom. a. For an anion, the number of electrons equal to the negative charge. b. For a cation, the number of electrons equal to the positive charge. 3) Place one pair of electrons between each pair of atoms to form a bond. Count the number of valence electrons in the molecule. Subtract 2 electrons from the total valence electrons for every bond you drew. 4) Use any remaining pairs as pairs around each atom (except hydrogen) so that each terminal atom is surrounded by electrons. If, after this is done, there are electrons left over, assign them to the atom. (If the central atom is an element in the third or higher period, it can have more than eight electrons.)

6 5) If the central atom has than 8 electrons at this point, change one or more of the lone pairs on the terminal atoms into a bonding pair between the central atom and terminal atom to form a ( or ) bond. a. As a general rule, double or triple bonds are most often encountered when both atoms are from the following list: Carbon, Nitrogen or Oxygen Using the steps above, draw the Lewis structures for the following covalent compounds. Example 2: Phosphorus trichloride Example 3: Carbon monoxide Example 4: Silicon dioxide Practice 4: Arsenic tribromide Practice 5: Carbon tetrafluoride Practice 6: Water

7 EXCEPTIONS to the OCTECT RULE: 1) Incomplete Octets A central atom with than electrons in its outer energy level. a. Incomplete octets are pretty rare and generally are only found in some, and compounds. b. Boron and aluminum form compounds in which they have valence electrons, rather than the usual 8 as predicted by the octet rule. c. Beryllium will form compounds in which it only has valence electrons. d. DO NOT double bond to satisfy their octets. 2) Expanded Octets A central atom with than electrons. a. These structures are only possible when the principle quantum number is greater than or equal to n = because their orbitals are available for bonding. Some can have up to 12 electrons surrounding the central atom! Example 5: Xenon tetrafluoride Example 6: Aluminum chloride Example 7: Sulfur hexafluoride Practice 7: Iodine pentachloride Practice 8: Bromine pentafluoride Practice 9: Boron trifluoride

8 Lewis Structures for POLYATOMIC IONS: Use the rules for drawing Lewis structures for covalent compounds. When determining the total number of valence electrons: o o For an anion, the number of electrons equal to the negative charge. For a cation, the number of electrons equal to the positive charge. Enclose the entire structure in and write the of the ion outside the brackets as a superscript. Example 8: Ammonium Ion Example 9: Sulfate Ion Practice 10: Carbonate Ion Practice 11: Perchlorate Ion Resonance Structures The possible structures of a molecule for which more than one Lewis structure can be written. Example 10: Ozone Formula: O 3 Practice 12: Carbonate Ion Formula: CO 3 2-

9 Valence Shell Electron Pair Repulsion (VSEPR) Theory Based on the idea that the bond and non-bond (lone) electron pairs in the valence shell of an element each other and seek to be as far apart as possible. It is this repulsion that causes the molecule or ion to have a particular. VSEPR Vocabulary: 1) Electron Pair Geometry The geometry of the on the central atom. (Count the number of atoms and non-bonding pairs of electrons around the central atom to determine the electron pair geometry. 2) Molecular Geometry The arrangement of the and electrons that represent the of the molecule. VSEPR CHART: Example # of atoms bonded to the central atom # of lone pairs of electrons on the central atom Electron Pair Geometry Molecular Geometry (Shape)

10 Draw the following Lewis structures to determine the electron pair geometry and the molecular geometry (shape) of the molecule or ion. Example 11: PF 3 Electron Pair Geometry: Molecular Geometry: Example 12: NO 2 1- Electron Pair Geometry: Molecular Geometry: Practice 13: H 2O Electron Pair Geometry: Molecular Geometry: Practice 14: BCl 3 Electron Pair Geometry: Molecular Geometry:

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