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GEN CHEM I chap 8

nerve transmission 1. movement of ions across cell membranes is the basis for the transmission of nerve signals.
nerve transmission and what ions 1. Na+ and K+ ions are pumped across membranes in opposite directions through ion channels. 2. Na+ pumped out of cell and K+ pumped in. 3.The ions channels can differentiate Na+ from K+ by their difference in size.
periodic properties ions size and other properties of atoms are periodic properties whose values can be predicted based on the element's position on the Periodic Table.
Group 1A or 1A metals tend to: lose one electron to form cations with a 1+ charge.
K lies directly belwo Na which means: 1.It indicates that it has more protons, neutrons, and electrons than sodium. (higher atomic number.) 2. Although a higher atomic number does not always result in a larger ion (or atom), it does in the case of potassium.
1 pm= ? m 1 pm=10^-12 m
the pumps and channels within the cell membranes are so sensitive... ... that they can distinguish between the sizes of these two ions and selectibely allow only one or the other to pass
Mendeleev Order elements by atomic mass Saw a repeating pattern of properties
Meendeleev and Quantum mechanics 1. Mendeleev’s Periodic Law allows us to predict what the properties of an element will be based on its position on the table 2.It doesn’t explain why the pattern exists Quantum Mechanics is a theory that explains why the periodic trends in the properties exist and knowing Why allows us to predict What
1.Electron Configurations 2. Ground state 1. An ELECTRON CONFIGURATION for an atom shows the particular orbitals that are iccupied for that atom. 2. GROUND STATE- or lowest energy state- electron configuration for a hydrogen atom. 1s^-1 --The 's' is the Orbital --the -1 is the number of electrons in the orbital
The Property of Electron Spin Spin is a fundamental property of all electrons All electrons have the same amount of spin The orientation of the electron spin is quantized, it can only be in one direction or its opposite spin up or spin down The electron’s spin adds a fourth quantum number to the description of electrons in an atom, called the Spin Quantum Number, ms not in the Schrödinger equation
Spin Quantum Number, ms, and Orbital Diagrams ms can have values of +½ or −½ Orbital Diagrams use a square to represent each orbital and a half-arrow to represent each electron in the orbital By convention, a half-arrow pointing up is used to represent an electron in an orbital with spin up Spins must cancel in an orbital paired
We often represent an orbital as a square and the electrons in that orbital as arrows the direction of the arrow represents the spin of the electron one electron is one arrow. two arrows is two arrows. A blank square is unoccupied,
Pauli Exclusion Principle No two electrons in an atom may have the same set of four quantum numbers Therefore no orbital may have more than two electrons, and they must have with opposite spins Knowing the number orbitals in a sublevel allows us to determine the maximum number of electrons in the sublevel
Cont. Pauli Exclusion Principles: s sublevel has 1 orbital, therefore it can hold 2 electrons p sublevel has 3 orbitals, therefore it can hold 6 electrons d sublevel has 5 orbitals, therefore it can hold 10 electrons f sublevel has 7 orbitals, therefore it can hold 14 electrons
Shielding & Effective Nuclear Charge 1. Each electron in a multielectron atom experiences both the attraction to the nucleus and repulsion by other electrons in the atom 2. These repulsions cause the electron to have a net reduced attraction to the nucleus – it is shielded from the nucleus 3. The total amount of attraction that an electron feels for the nucleus is called the effective nuclear charge of the electron
Penetration 1. The closer an electron is to the nucleus, the more attraction it experiences 2. The better an outer electron is at PENETRATING through the electron cloud of inner electrons, the more attraction it will have for the nucleus 3. The degree of penetration is related to the orbital’s radial distribution function in particular, the distance the maxima of the function are from the nucleus
n electron far from the nucleus is partly shielded by the electrons in th 1s orbital, reducing the effective net nuclear charge that it experiences.
An electron that penetrates the electron cloud of the 1s orbital experiences more of the nuclear charge.
Penetration and Shielding 1. The radial distribution function shows that the 2s orbital penetrates more deeply into the 1s orbital than does the 2p 2. The weaker penetration of the 2p sublevel means that electrons in the 2p sublevel experience more repulsive force, they are more shielded from the attractive force of the nucleus 3. The deeper penetration of the 2s electrons means electrons in the 2s sublevel experience a greater attractive force to the nucleus and are not shielded as effectively
Effect of Penetration and Shielding 1. Penetration causes the energies of sublevels in the same principal level to not be degenerate 2. In the fourth and fifth principal levels, the effects of penetration become so important that the s orbital lies lower in energy than the d orbitals of the previous principal level 3. The energy separations between one set of orbitals and the next become smaller beyond the 4s the ordering can therefore vary among elements causing variations in the electron configurations of the transition metals and their ions
picture on slide 29 Notice the following: because of penetration, sublevels within an energy level are not degenerate 2. penetration of the 4th and higher energy levels is so strong that their s sublevel is lower in energy than the d sublevel of the previous energy level 3. the energy difference between levels becomes smaller for higher energy levels (and can cause anomalous electron configurations for certain elements)
ELECTRON CONFIGURATION CONTINUEED: Electrons generally occupy the lowest energy orbitals available.
ELECTRON SPIN nd SUBLEVEL ENERGY SPLITTING CONTINUED 1. it affects the number of electrons allowed in one orbital. 2. determines the order of orbital filling within a level.
the orbital diagram the electron configuration of hydrogen (1s^1) can be represented in a slightly different way by an orital diagram. 2. the arrow pointing up or down represents the electron spin.
Spin 1. Spin is a basic property of electrons. One electron does not have more or less spin than another, all electrons have the same amount of spin. 2. The orientation of the electron's spin is quantized, with only two possibilities that we can call spin up and down.
quantize is the procedure of constraining something from a relatively large or continuous set of values to a relatively small discrete set.
the spin quantum number (ms) the values of (ms) are +1/2 (spin up) or -1/2 (spin down).
the hydrogen atom electron configuration: Helium: with its one electron spin up. Helium: is the first element on the periodic table that contains two electrons. The two electrons occupy the 1s orbital. How do the electrongs spin? Addressed in Pauli exclusion principle
quantum number one of four interrelated numbers that determine the shape and energy of orbitals, as specified by a solution of the Schrodinger equation.
Pauli exclusion principle: no two electrons in an atom can have the same four quantum numbers. **It emplies that each orbital can have a maximum of only two electrons with opposing spins in Helium.
Paulie exclusion principle: Helium Helium: 1. 2 electrons occupy the same orbital and have 3 identical quantum numbers (n, l, m1) BECAUSE they ar in the same orbital. 2. Therefore they must have different quantum numbers. 3. Electron configuration: 1s^2
In a hydrogen atom,the energy dpends only on: depends only on 'n'
degenerate orbitals Degenerate orbitals are orbitals at identical energy levels
multielectron atom are not degenerate. 1.Their energy depends on the value of 'l' (it's an L). 2.The energies of the sublevels are split.
Value of 'l' The lower the value of 'l' within a level, the lower the energy of the corresponding orbitals.
In order to understand why the sublevels split we muse examine: 3 key concepts associated with the energy of an electron in the vicinity of a nucleus. 1. Coulomb's law. 2. Shielding. 3. Penetration
Quick overview of Coulomb's law, shielding, and penetration: 1. Coulomb's law: describes the interactions between charged particles. 2. Shielding: describes how one electron can shield another electron from the full charge of the nucleus. 3. Penetration: describes how one atomic orbital can overlap spatially with another, thus penetrating into a region that is close to the nucleus (and therefore less shielded from the nuclear charge).
Periodic Law: 1.When the elements are arranged in order of increasing atomic mass, certain sets of properties recur periodically 2. Put elements with similar properties in the same column. Used pattern to predict properties of undiscovered elements Where atomic mass order did not fit other properties, he re-ordered by other properties Te & I
Coulomb’s Law Coulomb’s Law describes the attractions and repulsions between charged particles For like charges, the potential energy (E) is positive and decreases as the particles get farther apart as r increases. For opposite charges, the potential energy is negative and becomes more negative as the particles get closer together The strength of the interaction increases as the size of the charges increases electrons are more strongly attracted to a nucleus with a 2+ charge than a nucleus with a 1+ charge