Chem 110
potential energy
In terms of potential energy, chemistry is completely dominated by electrostatic forces, where the relative positioning of charged particles (nuclei, electrons, or ions) decides their energy. We already know that the force of interaction between two charged particles (with charges Q1 and Q2) is conveyed by Coulomb's law (E01-1-1). From basic physics, we also know that work (w) is equal to force (F) multiplied by distance (r), and that energy is "equivalent" to work (w). Thus, potential electrostatic energy (Eel) can be expressed by equation E01-4-2, with the charges Q1 and Q2 in coulombs (C). The proportionality constant is called Coulomb's constant, k = 1/4πε0 = 8.99 × 109 J·m/C2 where ε0 (= 8.86 × 10−12 F/m) is vacuum permittivity.
Enthalpy is the heat exchanged with the surroundings under constant pressure
Let's assume we have two flasks that start in the same state, and contain the same amount of dry ice (solid CO2) at room temperature. One flask is closed with a stopper and one has a deflated plastic bag over the mouth of the flask. Let's further assume that the same amount of heat (q) enters each of these systems at the same rate. Room temperature (25 °C) is higher than the temperature of the dry ice (−78 °C), and the solid will change into a gas as it warms up. The internal energy of both systems will increase as the samples warm up, but will the two systems end up with the same amount of internal energy? The one with the stopper remains at constant volume because it cannot expand; the pressure rises (F01-5-1a). The system with the bag remains at constant pressure (at least initially) because its volume can expand (F01-5-1b), while performing work against the ambient pressure. Figure F01-5-1. (a) The system absorbs heat from the surroundings at constant volume. No work is done by the system. (b) The system absorbs heat from the surroundings at constant pressure (until the bag is fully inflated). The work done by the system on the surroundings, against ambient pressure (P), is w = PΔV. At any point in time (before the bag fully inflates) after an equal amount of heat is absorbed by both systems, which one will have the higher internal energy? The system that absorbs heat at constant P (F01-5-1b) does work on the surroundings that lowers its internal energy. The system that absorbs heat at constant V does no work and therefore will have a higher internal energy. Anyone watching the flasks will be waiting for the cork to pop off the flask in Figure F01-5-1a. And it does! Eventually, when the bag is fully filled, it breaks as well under the increasing pressure of CO2. In chemical and physical changes occurring under constant pressure, it is indeed quite common that the only work done by the system is due to volume change (ΔV). This P-V work term is particularly important when the transformation leads to a significant change in the volume of a system, as often happens for processes involving gases. Since P and V are state functions, P-V work is also a state function, and under constant pressure, when only heat and P-V work contribute to energy changes, we can write (E01-5-1): ΔE=qp+w=ΔH−P⋅ΔV The heat exchanged with surroundings under such conditions is called enthalpy (qp = ΔH). Like internal energy, enthalpy is a state function. It is a very useful quantity as it is relatively easy to measure (we will learn how in later Lessons), and for many reactions in liquids or solids where change in volume is small or zero, it directly matches the changes of the internal energy of the system. Internal energy and enthalpy are both extensive properties, as they depend on the quantity of the matter in the system. Transfer of heat to and from the system during chemical and physical processes is so common that chemists use special names to indicate the direction of the heat flow. Processes where ΔH > 0 (i.e., the heat flows into the system) are called endothermic (endo means "into"), while processes where ΔH < 0 (with heat transfer to the surroundings) are called exothermic (exo means "out of").
Chapter 1: Lesson 1: Atomic Structure
Atoms are composed of protons and neutrons held in the nucleus surrounded by an electron cloud the basic building blocks of matter are atoms Atomic sizes and masses are very small. Atoms are 100 to 500 pm in diameter (1 to 5 Å where 1 Å = 100 pm = 10-10 m), with the heaviest having masses on the order of 10-22 g. They are built from even smaller subatomic particles, only three of which have bearing on chemical behavior: the proton, the neutron, and the electron. Protons and neutrons reside in the extremely small nucleus of the atom (ca. 10-14 m) and account for virtually all of the mass of the atom. The vast majority of volume of the atom, on the other hand, is essentially empty space occupied by a "cloud" of rapidly moving electrons, which contribute negligibly to the mass of the atom.
Ions
Atoms gain or lose electrons to form anions or cations which are attracted to each other by electrostatic forces Atoms may gain or lose electrons to become ions An ion with a net negative charge (excess of electrons) is called an anion (AN-ion) and an ion with a net positive charge (deficiency of electrons) is called a cation (CAT-ion). Simple examples of such ions include NH4+ (ammonium ion), HO- (hydroxide ion), or SO42- (sulfate). In such ions, the total number of electrons does not match the total number of protons in all nuclei. The numerical excess or deficit of electrons defines their overall charge. Atoms combine to form molecules or extensive solids Atoms may also combine with other atoms without forming ions. Atoms of the same element or of two or more elements may bond together to form molecules or extensive solids. For example, the simplest element, hydrogen, exists as a diatomic molecule, H2. The most stable form of oxygen is diatomic oxygen, O2, but another form is the triatomic version O3, which is called ozone. These different elemental forms, called allotropes, have different structures and different physical and chemical properties. Formation of extended solids made of many atoms that share electrons is common for the metallic and metalloid elements.
Electrons
Electrons and protons are held together by an attractive electrostatic force The electrons do not fly away from the atom because they are attracted to the protons in the nucleus by an electrostatic force that is proportional to the magnitude of the charges (Q1 and Q2) on the interacting particles and inversely proportional to the square of the distance (r) between them. Coulomb's law (k is proportionality constant): An electron has a negative charge of -1.602 × 10-19 coulombs (C). The charge of the proton is equal in magnitude to that of an electron, but has the opposite sign (+1.602 × 10-19 C). Neutrons, as the name indicates, have no charge (are neutral). Atoms, as a whole, have no net charge, as the number of electrons is equal to the number of protons. atomic mass unit (amu or u): 1.66054 × 10-24 g; exactly 1/12 of the mass of carbon-12 which contains 6 protons and 6 neutrons in its nucleus
Molecules
Molecules constructed with more than one type of atom are called molecular compounds. For example, hydrogen and all halogens (F, Cl, Br, I) typically form just one bond, oxygen tends to form two bonds, nitrogen prefers to form three bonds, and carbon does four. The number of bonds formed is called the valence of the atom. Atoms can be "connected" via single or multiple (double or triple) bonds. Carbon is able to form four bonds to other carbon atoms or to atoms of many other elements, resulting in a large variety of diverse and occasionally very complex structures. Indeed, life on Earth is based on compounds of carbon — these are called organic molecules.
The kinetic energy of particle motion and electrostatic potential energy dominate chemical processes
Potential energy can be thought of as "stored" energy and can manifest itself in many forms, such as electrostatic, nuclear, or gravitational energy. These types of energy are directly linked to the corresponding fundamental physical forces. On the other hand, kinetic energy (Ek) is determined only by the mass (m) and velocity (v) of a moving object kinetic energy equation
Energy is the capacity to do work and transfer heat
The SI unit of energy is the joule, (1 J = kg·m2/s2). Since energy is an extensive property (i.e. it depends on the amount of substance present), kilojoules (1 kJ = 103 × J) are often used on the molar scale (kJ/mol). Alternatively, calories (1 cal = 4.184 J) or kilocalories (kcal) are employed. There is also a nutritional Calorie unit (with a capital C), equal to 1 kcal.
The change in internal energy of a system results from the exchange of heat or workwith the surroundings
The internal energy of a system (E) is the total energy associated with the system, the sum of all sources of kinetic and potential energy. In most situations, chemists do not care about "absolute" internal energy of a sample, but instead are more interested in the internal energy changes (ΔE) directly connected to a physical or a chemical process under consideration (Figure F01-4-1). Since energy cannot be created or lost, chemists define the system and the surroundings and observe the energy flow between the two to determine the energy changes from the initial state to the final state. The system is defined as a collection of particles of interest, examples of which could be an atom emitting light energy, or a mole of molecules undergoing a chemical reaction. The surroundings are defined as everything else. The energy flow is always referenced with respect to the system, in a way analogous to the balance in a bank account. Thus, if the energy of the system is lowered (by transferring some of it to the surroundings) the energy change has a negative sign (energy has been "withdrawn" from the system). On the other hand, if energy is added to the system (from the surroundings) the energy change has a positive sign (energy has been "deposited" into the system).
The internal energy of a system is a function of state
The internal energy of the system does not depend on the path or method (i.e., the mechanism) used to get to the current state. It is a state function; it depends only on the existent state of the system (its temperature, volume and pressure). Let's look at a simple example in Figure F01-4-4. The internal energy is lowest for ice, and highest for hot water, but the internal energy of the samples at 25 °C is the same, regardless whether prepared by cooling the hot water or by heating and melting the ice. In a simple analogy, the overall change in altitude when traveling from State College, PA (370 m above sea level) to Boulder, CO (1650 m above sea level) is always 1280 m, regardless of the road chosen for our trip. In this example, the altitude behaves as a state function. Figure F01-4-4. Samples of identical masses of H2O in different states: ice at 0 °C, water at 25 °C, and water at 99 °C, close to a boiling point. The internal energy of the samples at 25 °C is the same.
is the interactions attractive or repulsive
We have now discussed the two major ways in which energy is transferred between a system and its surroundings, heat (q) and work (w). Adding heat (+q) to the system and doing work on the system (+w) both increase the overall energy of the system. ΔE=q+w Conversely, removing heat from the system (-q) and having the system do the work (-w) lowers the overall energy of the system. The sign convention followed in such energy transfer processes are illustrated in Figure 01-4-3.
is it a state functions
We use capital letters to signify state functions like P (pressure), T (temperature), E (internal energy), and H (enthalpy). However, heat and work (q and w) are not state functions. Depending on how we set up a process, we can release lots of heat and do little or no work, or we can force the process to do more work and release less heat. Consider the combustion of a gallon of gasoline. We could light a gallon of gasoline on fire and it will generate a great deal of heat as it burns. However, if we burn the gasoline in a car engine we can use the released energy to power the car. Some heat will still be produced, but a substantial fraction of the energy will be used to do work. In each case the total amount of energy released by the burning of the gasoline will be the same, but the relative amount of heat and work produced is different (ΔE = q + w).
Isotopes
are atoms with the same atomic number but a different number of neutrons When an element occurs in nature as a mixture of isotopes, its atomic weight (AW) is the average of the masses of all isotopes present. AW: E [(istotope mass) X (fractional natural abundance)] The atomic weight of carbon: For example, naturally occurring carbon is composed of 98.93% of the carbon-12 isotope with atomic mass of 12 amu (exactly ) and 1.07% of the carbon-13 isotope with atomic mass of 13.00335. AW(C) = (0.9893)(12amu)+(0.0107)(13.00335amu) = 12.01amu
Atomic components
protons (p+), charge (C): +1.602X10^-19 Charge (au): +1 Mass (kg): 1.6726X10^27 Mass (amu): 1.0073 neutrons (n^0) charge (C): 0 charge (au): 0 mass (kg): 1.6749 × 10−27 mass (amu): 1.0087 electrons (e^-) charge (C): −1.602 × 10−19 charge (au): −1 mass (kg): 9.1094 × 10−31 mass (amu): 5.486 × 10−4
