Study Guide for Chapters 7 & 8, Study Guides, Projects, Research of Chemistry

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Study Guide 7 & 8

Chapter 7 - Periodic Properties of the Elements Main Topic : I. Mendeleev and the Periodic Table II. Periodicity & Atomic Trend III. Effective Nuclear Charge IV. Atom & Ion Size V.Isoelectronic Series VI. Ionization Energy VII. Electron Affinity VIII. General Trend IX. Electron Configuration X. Group trends XI. Metal, Nonmetals, & Metalloids Key Concepts :

• Mendeleev: Mendeleev’s table is special because not only did it predict elements that had yet been discover but also made fairly accurate prediction of their properties
• Atomic Number / Henry Moseley: developed the atomic number which is now used to organized the periodic table after Ernest Rutherford discovered the nuclear atom.
• Periodicity are patterns of properties based on atomic number Atomic Trends: 1) Sizes of atoms and ions 2) Ionization energy 3) Electron affinity 4) Some group chemical property trends

• Effective Nuclear Charge = net positive charge pulling electrons toward nucleus ↳ Effective Nuclear Charge increases as you move to right or down
• Atom Size: is increasing the more you to the left or down ↳ Decreases from left to right [ Period = Z eff increases ] ↳ Increases from top to bottom [ Group = n increases ]
• Ion Size: Determined by inter atomic distance ionic compounds ↳ Depends on:
  1. Nuclear charge
  2. Number of electros
  3. Orbital which electrons reside in ↳ Cations : Smaller than parent & Repulsions = Reduced ↳ Anions : Larger than parent & Repulsions = Increased
• Isoelectronic Series = Ions that have the same number of electrons ↳ Ion size decreases when nuclear charge increases Valence Electrons are tightly bound to more positive nucleus
• Ionization Energy = min. Energy required to remove an electron ↳ The first ionization energy is that energy required to remove the first electron. ↳ The second ionization energy is that energy required to remove the second electron. Exceptions: enters new sub-level & first electron to pair in sublevel
• Electron Affinity = is the energy change accompanying the addition of an electron to a gaseous atom/how likely to gain another electron: → More toward halogens
• General Trends
• Group Trends = element in groups have similar properties
• Metal: forms cations PROPERTIES:
  1. Shiny luster

ion size, isoelectronic series, ionization energy, electron affinity, electron configuration, group trends, and the properties of metals, nonmetals, and metalloids. The chapter begins with an overview of Mendeleev's contribution to the development of the periodic table and how it uniquely predicted elements yet to be discovered, along with their properties. It then proceeds to discuss atomic trends, which are patterns of properties based on atomic numbers. These trends include the sizes of atoms and ions, ionization energy, electron affinity, and some group chemical property trends. The concept of effective nuclear charge is introduced, which is the net positive charge pulling electrons toward the nucleus. This charge is shown to increase as you move to the right or down on the periodic table. The size of atoms is another key topic covered in this chapter. It's explained that atom size tends to increase the further you move to the left or down on the periodic table. Conversely, atom size decreases from left to right, primarily because the effective nuclear charge increases. Ion size is carefully analyzed, with factors determining it including interatomic distance in ionic compounds, nuclear charge, the number of electrons, and the orbital in which electrons reside. The chapter further explains that cations are smaller than their parent atoms, while anions are larger. The chapter also explores the concept of isoelectronic series, referring to ions that have the same number of electrons. It highlights that ion size decreases when the nuclear charge increases, as valence electrons are tightly bound to a more positive nucleus. Ionization energy, defined as the minimum energy required to remove an electron, is discussed in detail. The chapter distinguishes between the first and second ionization energy, noting that each corresponds to the energy required to remove the first and second electron, respectively. Electron affinity, the energy change accompanying the addition of an electron to a gaseous atom, is also covered. The chapter further examines general trends, group trends, and electron configuration. The properties of metals, nonmetals, and metalloids form a significant part of the discussion. Metals are characterized by their ability to form cations, and their properties include having a shiny luster, conducting heat and electricity, being malleable and ductile, being solid at room temperature (except for mercury), and having low ionization energies. Nonmetals, on the other hand, form anions, and their properties vary depending on the particular element. However, generally, solid nonmetals are dull, brittle, and poor conductors. Metalloids, which form ionic compounds, possess properties of both metals and nonmetals and are electronic semiconductors.

The chapter concludes by presenting the formula for effective nuclear charge (Z eff = Z - S), where Z represents attractions and S represents shielding. It also discusses the theories and rules related to Van der Waals radius and covalent radius. Chapter 8 - Basic Concepts of Chemical Bonding Main Topic : I. Chemical Bonds II. Lewis Symbols & Structures III. Octet Rule IV. Lattice Engery V. Pairs & Bonds VI. Polarity VII. Electronegativity VII. Formal Charge VIII. Resonance IX. Bond Enthalpies Key Concepts :

• [I] Ionic Bonds: ↳ Electrostatic attraction between ions ; Born-Haber Cycle Metals & Non-metal; Electron transfer; Very exothermic Brittle, High melting point, crystalline, cleave along smooth lines Energies: atoms = endo ; cation = endo & IE ; anion = exo & EA ; solid = exo
• Covalent Bonds: ↳ Sharing electrons ; “Only true bond”

• Bond Enthalpy = the enthalpy associated with breaking 1 mole of a particular bond in a gaseous substance Formulas : Formal Charge = VE - 1/2 (B=bonded)- (N = Non-bonded) Lattice Energy = E el = K Q1 Q2 / d Bond Enthalpies = ΔHrxn = Σ [Broken Bonds] - Σ [Bond Created] Theories & Rules : ⟡ Octet Rule : Each element wants to be a Noble Gas/Have 8 (e-) [Oct = 8] ↳ EXCEPTIONS: an incomplete octet, (2) odd-electron molecules, and (3) an expanded octet. ; 1A, 2A, & 3A less than Octet & 3rd period onward more than Octet Watch Exceptions Video Exceptions 1) If the ion has an odd number of (e-) 2) If the ion has less that an octet = H,Be,B,Al 3) If the ion has more than eight valence = PF Summary : Chapter 8 covers different types of bonds and related concepts.

Ionic bonds are formed through the electrostatic attraction between ions. This type of bond is typically found between metals and non-metals where electron transfer occurs. The reaction is usually very exothermic. Characteristics of ionic bonds include being brittle, having high melting points, and cleaving along smooth lines. Covalent bonds, on the other hand, are formed by sharing electrons to achieve noble gas configuration. The electrostatic interactions within covalent bonds can be described as attractions between electrons and nuclei, and repulsion both between electrons and between nuclei. Metallic bonds involve free electrons which hold metal atoms together, often referred to as a 'sea of electrons'. Lattice Energy is a concept that refers to the energy required to convert one mole of an ionic solid into a gaseous ionic constituent. This is a key concept in understanding the strength and stability of ionic compounds. Bond polarity is a measure of how equally or unequally the electrons in a covalent bond are shared. If one atom has a greater electron-attracting force, the bond is polar. If both atoms have equal electron-attracting forces, the bond is non-polar. The Octet Rule is a fundamental concept in chemistry, stating that each element aims to have eight electrons in its outer shell to achieve stability. However, there are exceptions to this rule, such as ions with odd numbers of electrons, ions with less than an octet, and ions with more than eight valence electrons. Study Guide Key: (e-) = Electron IE = Ionization energy EA = Electron Affinity LE = Lattice Engery Trends: ←↑→↓

Lattice Energy = ↑→ Electronegative = ↑→ Ionization energy = ↓→