CHEM 1030 Cagg Chapter 4 Notes
CHEM 1030 Cagg Chapter 4 Notes Chem 1030
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This 4 page Class Notes was uploaded by Amy Notetaker on Tuesday February 23, 2016. The Class Notes belongs to Chem 1030 at Auburn University taught by Brett A Cagg in Spring 2016. Since its upload, it has received 42 views. For similar materials see Fundamental Chemistry I in Chemistry at Auburn University.
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Date Created: 02/23/16
Lecture / Book Notes: Chapter 4 (2/17/2016 & 2/22/2016) CHEM 1030 Cagg Highlighted: Vocab ----- Highlighted: Formula/Numbers Section 4.1 v The Development of the Periodic Table • The elements on the periodic table are arranged in order of increasing atomic number, and not by mass as scientists originally thought. • Periodicity: the tendency to be periodic or to be able to reoccur. Section 4.2 v The Modern Periodic Table • The outermost electrons determine what categories the elements are divided into. - The main group elements - The noble gasses - The transition elements/the transition metals - The lanthanides - The actinides • The main group elements: also known as the representative elements, which are in groups 1A-7A. • All noble gasses except helium have a full p subshell. • The transition metals either have a full d subshell, or are willing to lose electrons in order to achieve incomplete d subshells. - The transition metals are in groups 1B and 3B-8B. • Lanthanides and actinides have incompletely filled f shells, so they are sometimes referred to as f-block elements. • Valence electrons: the outermost electron shell, which determines how an atom behaves with other atoms. • Since noble gasses have completely filled “ns” and “np” shells, they keep to themselves and don’t often mingle with other elements, because they are already stable. • Transition metals are unique from the other elements because they have certain properties due to incompletely filled d subshells. Section 4.3 v Effective Nuclear Charge • Nuclear charge: the number of protons in an atom’s nucleus. • Effective nuclear charge: the magnitude of positive charge that an atom experiences. - In hydrogen, the nuclear charge and effective nuclear charge are the same. • Shielding: when electrons are attracted to the nucleus, but repelled by one another. • Core electrons: the most effective electrons at shielding. • As you move right across the 2 period, the nuclear charge increases by 1 with every new element. The effective nuclear charge only increases by 0.64. • As you move down a column of the periodic table, the effective nuclear charge changes less than the nuclear charge. The nuclear charge is large when you move down. Section 4.4 v Periodic Trends in Properties of Elements • Most chemical and physical properties depend on the nuclear charge. v Atomic Radius • The atomic radius: the distance between the nucleus of the atom, and its valence shell. This also has 2 other definitions: - Metallic radius: half of the distance between 2 nuclei of identical metal atoms that are right next to each other. - Covalent radius: half of the distance between 2 chemically bonded nuclei that are right next to each other. • As you move left to right across a period, the atomic radius decreases. - When you move from left to right, the effective nuclear charge increases because more electrons are added into the valence shell, making the atomic radius more negative and therefore decrease. • As you move top to bottom within a group, the atomic radius increases. - This is because the outermost shell has an increasing value of “n”, which means that it is further from the nucleus, so the radius is bigger. v Ionization Energy • Ionization energy: the minimum energy required to remove an electron during the gas phase. • This is expressed using kJ/mol. • Cation: an atom that has lost an electron and has become more positive. • Ionization increases as effective nuclear charge increases. • In a shell, electrons with higher levels of ℓ are easier to remove due to more energy and because they are loosely held by the nucleus. • 2A elements require a removal of an electron from an s orbital for ionization to occur. • 3A elements require a removal of an electron from a p orbital for ionization to occur. • 5A and 6A elements require a removal of an electron from a p orbital for ionization to occur. - 6A requires the removal of a PAIR. • Coulomb’s law states that the attractive force between the effective nuclear charge and valence electrons gets weaker as the distance between them increases. • It is harder to remove an electron from a cation than from an atom. • It takes less energy to remove valence electrons and more energy to remove core electrons. - This is due to the core electrons being closer to the nucleus and because they experience a greater effective nuclear charge. • The process of ionization is a chemical process since the identity of whatever involved changes. v Electron Affinity • Electron affinity: the energy released when a gas phase atom accepts an electron. - Positive electron affinity shows that a process is energetically favorable. - The larger and more positive the electron affinity is, the more favorable the process will be. • Anion: an atom that has gained an electron and has become more negative. • Electron affinity increases when moving right to left across a period, because it becomes easier to add a negatively charged electron when the positive charge of the nucleus increases. • The electron affinity of group 2A elements is lower than group 1A elements. • The electron affinity of group 5A elements is lower than group 4A elements. • More than one electron can be added to an atom, as it can also be taken away from an atom. v Metallic Character • Metals - Shiny, lustrous, malleable, and ductile - Are good conductors of heat and electricity - Commonly form cations, due to having low ionization energies • Nonmetals - Vary in color and don’t have a shiny appearance - Are brittle - Are poor conductors of heat and electricity - Commonly form anions, due to having high ionization energies • Metallic character increases from top to bottom in a group. • Metallic character decreases from left to right in a period. • Metalloids: elements that have combined properties of metals and nonmetals ▯▯×▯▯ • Coulomb’s law formula for periodic trends: F α ▯▯ - “F” is the force between 2 objects - "Q▯and Q "▯are the 2 objects - "d” is the distance between the 2 objects - Opposite sign charges indicate an attractive force. - Same sign charges indicate a repulsive force. • Diagonal relationship: the similarities between the pairs of elements in different groups and periods. - This is due to similar charge densities of cations. Section 4.5 • Elements in 1A and 2A have the ion charges +1 and +2. • Elements in 6A and 7A have the ion charges -2 and -1. v Ions of Main Group Elements • Group 8A elements have the highest ionization energy. • Isoelectronic: identical electron configuration. • When it comes to writing the configuration for ions formed by main group elements, you: - Write normal configuration of the atom. - Add or remove the number of electrons needed. v Ions in d- Block Elements • An atom always loses an electron from the highest n value shell. Section 4.6 • An atom’s radius increase/decreases when it gains or loses an electron. • Ionic radius: the radius of a cation or an anion that affects the chemical property of the substance that it deals with. v Comparing Ionic Radius with Atomic Radius • When an atom becomes a cation, its radius decreases. - This is because an electron is taken away from an atom. • A significant decrease can be seen when all of the valence electrons are removed. • When an atom becomes an anion, its radius increases. - This is because an electron is added to an atom. v Isoelectronic Series • Isoelectronic series: when 2 or more species have the same electron configuration with different nuclear charges. • Species that has the smallest nuclear charge has the largest radius. • Species that has the largest nuclear charge has the smallest radius.
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