Review of Unit 3- Energy and Bonding
Chapter 10: Energy
Chapter 11: Modern Atomic Theory
10.2 The Flow of Energy
10.3 Energy and Chemical Reactions
10.1 Energy, Temperature, and Heat
10.4 Using Energy in the Real World
Energy is the ability to do work or produce heat.
The law of conservation of energy states that energy can be converted from one form to another but can be neither created nor destroyed.
Temperature is a measure of the random motions of the components of a substance.
Heat can be defined as a flow of energy due to a temperature difference.
When a process results in the evolution of heat, it is said to be exothermic; that is, energy flows out of the system.
Processes that absorb energy from the surroundings are said to be endothermic.
The internal energy, E, of a system can be defined most precisely as the sum of the kinetic and potential energies of all “particles” in the system.
The calorie is defined as the amount of energy (heat) required to raise the temperature of 1 g of water by 1°C.
The amount of energy required to change the temperature of one gram of a substance by one Celsius degree is called its specific heat capacity or, more commonly, its specific heat.
For a reaction occurring under conditions of constant pressure, the change in enthalpy (Delta H) is equal to the energy that flows as heat.
Hess's Law: In going from a particular set of reactants to a particular set of products, the change in enthalpy is the same whether the reaction takes place in one step or in a series of steps.
The sign of Delta H indicates the direction of the heat flow at constant pressure.
The Greenhouse Effect: Molecules in the atmosphere, principally H2O and CO2, strongly absorb infrared radiation and radiate it back toward the earth. A net amount of thermal energy is retained by the earth’s atmosphere, causing the earth to be much warmer than it would be without its atmosphere.
Energy spread means that in a given process, concentrated energy is dispersed widely.
Entropy (designated by the letter S) is a measure of disorder or randomness.
11.1 Atoms and Energy
11.3 Atomic Orbitals
11.4 Electron Configurations and Atomic Properties
11.2 The Hydrogen Atom
The wavelength (symbolized by the Greek letter lambda) is the distance between two consecutive wave peaks
The frequency of the wave (symbolized by the Greek letter nu) indicates how many wave peaks pass a certain point per a given time period.
When atoms are excited, they release energy by emitting light. The energy of the photon corresponds exactly to the energy change experienced by the emitting atom.
All hydrogen atoms have the same set of discrete energy levels. We say the energy levels of hydrogen are quantized. That is, only certain values are allowed.
In the Bohr atom, the energy levels in the hydrogen atom represented certain allowed circular orbits.
The wave mechanical model gives no information about when the electron occupies a certain point in space or how it moves, it just provides the probabilities of where the electron is located. In fact, we have good reasons to believe that we can never know the details of electron motion, no matter how sophisticated our models may become.
The probability map for the hydrogen electron on the previous page is called an orbital. Although the probability of finding the electron decreases at greater distances from the nucleus, the probability of finding it at even great distances from the nucleus never becomes exactly zero.
We call the discrete energy levels principal energy levels and label them with whole numbers. Next we find that each of these levels is divided into sublevels.
Pauli exclusion principle: An atomic orbital can hold a maximum of two electrons, and those two electrons must have opposite spins.
Principal energy level 2 has two sublevels. These sublevels are labeled 2s and 2p. The 2s sublevel consists of one orbital (called the 2s), and the 2p sublevel consists of three orbitals (called 2px, 2py, and 2pz).
For example, level 3 has three sublevels, which we label 3s, 3p, and 3d. The 3s sublevel contains a single 3s orbital, a spherical orbital larger than 1s and 2s. Sublevel 3p contains three orbitals: 3px, 3py, and 3pz, which are shaped like the 2p orbitals except that they are larger. The 3d sublevel contains five 3d orbitals with the shapes and labels shown in the Textbook.
Level 4 has four sublevels labeled 4s, 4p, 4d, and 4f. The 4s sublevel has a single 4s orbital. The 4p sublevel contains three orbitals (4px, 4py, and 4pz). The 4d sublevel has five 4d orbitals. The 4f sublevel has seven 4f orbitals.
The 4s, 4p, and 4d orbitals have the same shapes as the earlier s, p, and d orbitals, respectively, but are larger. We will not be concerned here with the shapes of the f orbitals.
In its ground state, hydrogen has its lone electron in the 1s orbital. This is commonly represented in two ways. First, we say that hydrogen has the electron arrangement, or electron configuration, 1s1.
Valence electrons are the electrons in the outermost (highest) principal energy level of an atom.
Elements with the same valence electron arrangement show very similar chemical behavior.
Chapter 12: Chemical Bonding
Atoms get larger as we go down a group on the periodic table and that they get smaller as we go from left to right across a period.
12.1 Characteristics of Chemical Bonds
12.2 Characteristics of Ions and Ionic Compounds
12.3 Lewis Structures
12.4 Structures of Molecules
A bond is a force that holds groups of two or more atoms together and makes them function as a unit.
An ionic compound results when a metal reacts with a nonmetal.
The type of bonding where electrons are shared by nuclei is called covalent bonding.
A polar covalent bond is formed when the atoms are not so different that electrons are completely transferred but are different enough so that unequal sharing of electrons results,
Electronegativity is the relative ability of an atom in a molecule to attract shared electrons to itself.
The polarity of a bond depends on the difference between the electronegativity values of the atoms forming the bond.
When representative metals and nonmetals react, they transfer electrons in such a way that both the cation and the anion have noble gas electron configurations.
Chemical compounds are always electrically neutral.
A cation is always smaller than the parent atom, and an anion is always larger than the parent atom.
The Lewis structure is a representation of a molecule that shows how the valence electrons are arranged among the atoms in the molecule.
In writing Lewis structures, we include only the valence electrons.
A single bond involves two atoms sharing one electron pair. A double bond involves two atoms sharing two pairs of electrons.
A molecule shows resonance when more than one Lewis structure can be drawn for the molecule. In such a case, we call the various Lewis structures resonance structures.
The bonding and nonbonding electron pairs (lone pairs) around a given atom are positioned as far apart as possible.
Whenever two pairs of electrons are present around an atom, they should always be placed at an angle of 180° to each other to give a linear arrangement.
Whenever three pairs of electrons are present around an atom, they should always be placed at the corners of a triangle (in a plane at an angle of 120° to each other).
Whenever four pairs of electrons are present around an atom, they should always be placed at the corners of a tetrahedron (the tetrahedral arrangement).
When using the VSEPR model to predict the molecular geometry of a molecule, a double bond is counted the same as a single electron pair.