Chemistry Essentials For Dummies
there will probably be changes within the year that diverge from this but take this as a general reference
Chapter 1: Matter and Energy: Exploring the Stuff of Chemistry
Knowing the States of Matter and Their Changes
Matter is anything that has mass and occupies space. It can exist in one of three classic states: solid, liquid, and gas. When a substance goes from one state of matter to another, the process is called a change of state, or phase change. Some rather interesting things occur during this process, which I explain in this section.
Solids, liquids, and gases
Particles of matter behave differently depending on whether they’re part of a solid, liquid, or gas. As figure 2-1 shows, the particles may be organized or clumped, close or spread out. In this section, you look at the solid, liquid, and gaseous states of matter.
Solids
At the macroscopic level, the level at which you directly observe with your senses, a solid has a definite shape and occupies a definite volume. Think of an ice cube in a glass — it’s a solid. You can easily weigh the ice cube and measure its volume.
At the microscopic level (where items are so small that people cant directly observe them), the particles that make up the solid are very close together and aren’t moving around very much (see Figure 2-1a). That’s because in many solids, the particles are pulled into a rigid, organized structure of repeating patterns called a crystal lattice. The particles in the crystal lattice are still moving but barely — it’s more of a slight vibration. Depending on the particles, this crystal lattice may be of different shapes.
Liquids
Unlike solids, liquids have no definite shape; however, they do have a definite volume, just like solids do. The particles in liquids are much farther apart than the particles in solids, and they’re also moving around much more (see Figure 2-1b).
Even though the particles are farther apart, some particles in liquids may still be near to each other, clumped together in small groups. The attractive forces among the particles aren’t as strong as they are in solids, which is why liquids don’t have a definite shape. However, these attractive forces are strong enough to keep the substance confined in one large mass — a liquid — instead of going all over the place.
Gases
A gas has no definite shape and no definite volume. In a gas, particles are much farther apart than they are in solids or liquids (see Figure 2-1c), and they’re moving relatively independent of each other. Because of the distance between the particles and the independent motion of each of them, the gas expands to fill the area that contains it (and thus it has no definite shape).
Condensing and freezing
If you cool a gaseous or liquid substance, you can watch the changes of state, or phase changes, that occur. Here are the phase changes that happen as substances lose energy:
✓ Condensation: When a substance condenses, it goes from a gas to a liquid state. Gas particles have a high amount of energy, but as they’re cooled, that energy decreases. The attractive forces now have a chance to draw the particles closer together, forming a liquid. The particles are now in clumps, as is characteristic of particles in a liquid state.
✓ Freezing: A substance freezes when it goes from a liquid to a solid. As energy is removed by cooling, the particles in a liquid start to align themselves, and a solid forms. The temperature at which this occurs is called the freezing point (fp) of the substance.
You can summarize the process of water changing from a gas to a solid in this way:
H2O(g) \rightarrow H2O(l) \rightarrow H2O(s)
Here, the (l) stands for liquid, the (g) stands for gas, and (s) stands for solid.
Melting and boiling
As a substance heats, it can change from a solid to a liquid to a gas. For water, you represent the change like this:
H2O(s) \rightarrow H2O(l) \rightarrow H2O(g)
This section explains melting and boiling, the changes of state that occur as a substance gains energy.
From solid to liquid
When a substance melts, it goes from a solid to a liquid state. Here’s what happens: If you start with a solid, such as ice, and take temperature readings while heating it, you find that the temperature of the solid begins to rise as the heat causes particles to vibrate faster and faster in the crystal lattice.
After a while, some of the particles move so fast that they break free of the lattice, and the crystal lattice (which keeps a solid solid) breaks apart. The solid begins to go form a solid sate of a liquid state — a process called melting. The temperature at which melting occurs is called the melting point (mp) of the substance. The melting point for ice is 32°F, or 0°C.
During changes of state, such as melting, the temperature remains constant — even though a liquid contains more energy than a solid. So if you watch the temperature of ice as it melts, you see that the temperature remains steady at 0°C until all the ice has melted. The melting point (solid to liquid) is the same as the freezing point (liquid to solid).
From liquid to gas
The process by which a substance moves from the liquid state to the gaseous state is called boiling.
If you heat a liquid, such as a pot of cool water, the temperature of the liquid rises and the particles move faster and faster as they absorb the heat. The temperature rises until the liquid reaches the next change of state — boiling. As the particles heat up and move faster and faster, they begin to break the attractive forces between each other and move freely as a gas, such as steam, the gaseous form of water.
Skipping liquids: Sublimation.
Most substances go through the logical progression from solid to liquid to gas as they’re heated (or vice versa as they’re cooled). But a few substances go directly from the solid to the gaseous state without ever becoming a liquid. Scientists call this process sublimation. Dry ice — solid carbon dioxide, written as CO2(s) — is the classic example of sublimation. You can see dry ice pieces becoming smaller as the solid begins to turn into a gas, but no liquid forms during this phase change.
The process of sublimation of dry ice is represented as
CO2(s) \rightarrow CO2(g)
Besides dry ice, mothballs and certain solid air fresheners also go through the process of sublimation. The reverse of sublimation is deposition — going directly from a gaseous state to a solid state.
Pure substances
A pure substance, like salt or sugar, has a definite and constant composition or makeup. A pure substance can either be an element or a compound, but the composition of a pure substance doesn’t vary.
Elements
An element is composed of a single kind of atom. An atom is the smallest particle of an element that still has all the properties of the element. For instance, if you slice and slice a chunk of the element gold until only one tiny particle is left that can’t be chopped anymore without losing the properties that make gold gold, then you have an atom. (I discuss properties later in the section “Nice Properties You’ve Got There.”)
The atoms in an element all have the same number of protons. Protons are subatomic particles — particles of an atom. (Chapter 2 covers the three major subatomic particles in great, gory detail.) The important thing to remember right now is that elements are the building blocks of matter. They’re represented in the periodic table, which you explore in Chapter 3.
Compounds
A compound is composed of two or more elements in a specific ratio. For example, water (H2O) is a compound made up of two elements, hydrogen (H) and oxygen (O). These elements are combined in a very specific way — in a ratio of two hydrogen atoms to one oxygen atom (hence H2O). A lot of compounds contain hydrogen and oxygen, but only one has that special 2-to-1 ratio called water.
A compound has physical and chemical properties different from the elements that make it up. For instance, even though water is made up of hydrogen and oxygen, water’s properties are a unique combination of the two elements.
Chemists can’t easily separate the components of a compound: They have to resort to some type of chemical reaction.
Throwing mixtures into the mix.
Mixtures are physical combinations of pure substances that have no definite or constant composition; the composition of a mixture varies according to whoever prepares the mixture. Each component of the mixture retains its own set of physical and chemical characteristics.
Chemists can easily separate the different parts of a mixture by physical means, such as filtration. For example, suppose you have a mixture of salt and sand, and you want to purify the sand by removing the salt. You can do this by adding water, dissolving the salt, and then filtering the mixture. You then end up with pure sand.
Mixtures can be either homogeneous or heterogeneous:
✓ Homogeneous mixtures: Sometimes called solutions, homogeneous mixtures are relatively uniform in composition. Every portion of the mixture is like every other portion. If you dissolve sugar in water and mix it really well, your mixture is basically the same no matter where you sample it. I cover solutions in Chapter 10.
✓ Heterogeneous mixtures: The composition of heterogeneous mixtures varies from position to position within the sample. For instance, if you put some sugar in a jar, add some sand, and then give the jar a couple of shakes, your mixture doesn’t have the same composition throughout the jar. Because the sand is heavier, there’s probably more sand at the bottom of the jar and more sugar at the top.
Measuring Matter
Scientists often make measurements, which may include such things as mass, volume, and temperature. If each nation had its own measurement system, communication among scientists would be tremendously hampered, so scientists adopted a worldwide measurement system to ensure they can speak the same language.
The SI system (from the French Système international) is a worldwide measurement system based on the older metric system. SI is a decimal system with basic units for things like mass, length, and volume and prefixes that modify the basic units. For example, here are some very useful SI prefixes:
kilo- (k) means 1,000
centi- (c) means 0.01
milli- (m) means 0.001
So a kilogram (kg) is 1,000 grams, and a kilometer (km) is 1,000 meters. A milligram (mg) is 0.001 grams — or you can say that there are 1,000 milligrams in a gram.
Here are some basic SI units ad how they compare to the English units common in the U.S.:
✓ Length: The basic unit of length in the SI system is the meter (m). A meter is a little longer than a yard; 1.094 yards are in a meter. The most useful SI/English conversion for length is 2.54 centimeters = 1 inch
✓ Mass: The basic unit of mass in the SI system for chemists is the gram (g). And the most useful conversion for mass is 454 grams = 1 pound
✓ Volume: The basic unit for volume in the SI system is the liter (L). The most useful conversion is 0.946 liter = 1 quart
Suppose you want to find the weight of a 5.0-lb. bag of potatoes in kilograms. The setup would look like this:
\frac{5.0\operatorname{lb}s}{1}\cdot\frac{454g}{1\operatorname{lb}}\cdot\frac{1\operatorname{kg}}{1000g}=2.3\operatorname{kg}
Nice Properties You’ve Got There
When chemists study chemical substances, they examine two types of properties:
✓ Chemical properties: These properties enable a substance to change into a brand-new substance, and they describe how a substance reacts with other substances. Does a substance change into something completely new when water is added — like sodium metal changes to sodium hydroxide? Does the substance burn in air?
✓ Physical properties: These properties describe the physical characteristics of a substance. The mass, volume, and color of a substance are physical properties, and so is its ability to conduct electricity. Physical properties can be extensive or intensive.
Extensive properties, such as mass and volume, depend on the amount of matter present
Intensive properties, such as color and density, don’t depend on the amount of matter present. A large chunk of gold, for example, is the same color as a small chunk of gold.
Intensive properties are especially useful for chemists because intensive properties can be used to identify a substance. For example, knowing the differences between the density of quartz and diamond allows a jeweler to check out that engagement ring quickly and easily.
Density (d) is the ratio of the mass (m) and volume (v) of a substance. Mathematically, it looks like this:
d = m/v
Usually, mass is described in grams (g) and volume is described in milliliters (mL), so density is g/mL. Because the volumes of liquids vary somewhat with temperature, chemists usually specify the temperature at which they made a density measurement. Most reference books report densities at 20°C, for example, is 1 g/mL.
You may sometimes see density reported as g/cm3 or g/cc, both of which mean grams per cubic centimeter. These units are the same as g/mL.
Calculating density is pretty straightforward. You measure the mass of an object by using a balance or scale, determine the object’s volume, and then divide the mass by volume.
With an irregular solid, like a rock, you can measure the volume by using the Archimedes principle. The Archimedes principle states that the volume of a solid is equal to the volume of water it displaces. Simply read the volume of water in a container, submerge the solid object, and read the volume level again. The difference is the volume of the object.
Energy Types
Matter is one of the two components of the universe. Energy is the other. Energy is the ability to do work.
Energy can take several forms, such as heat energy, light energy, electrical energy, and mechanical energy. But two general categories are especially important to chemists, kinetic energy and potential energy
Kinetic energy
Kinetic energy is energy of motion. A baseball flying through the air toward a batter has a large amount of kinetic energy — just ask anyone who’s ever been hit with a baseball.
Chemists sometimes study moving particles, especially gases, because the kinetic energy of these particles helps determine whether a particular reaction may take place. As particles collide, kinetic energy may be transferred from one particle to another, causing chemical reactions.
Kinetic energy can be converted into other types of energy. In a hydroelectric dam, the kinetic energy of the falling water is converted into electrical energy. In fact, a scientific law — the law of conservation of energy — states that in ordinary chemical reactions (or physical processes), energy is neither created nor destroyed, but it can be converted from one to another.
Potential energy
Potential energy is stored energy. Objects may have potential energy stored in terms of their position. A ball up in a tree has potential energy due to its height. If a ball were to fall, that potential energy would be converted to kinetic energy.
Potential energy due to position isn’t the only type of potential energy. Chemists are far more interested in the energy stored (potential energy) in chemical bonds, which are the forces that hold atoms together in compounds.
Human bodies store energy in chemical bonds. When you need that energy, your body can break those bonds and release it. The same is true of the fuels people commonly use to heat their homes and run their automobiles. Energy is stored in these fuels — gasoline, for example — and is released when chemical reactions take place.
Temperature and Heat
When you measure, say, the air temperature in your backyard, you’re really measuring the average kinetic energy (the energy of motion) of the gas particles in your backyard. The faster those particles are moving, the higher the temperature is.
The temperature reading from your thermometer is related to the average kinetic energy of the particles. Not all particles are moving at the same speed. Some are going very fast, and some are going relatively slow, but most are moving at a speed between two extremes.
If you’re in the U.S., you probably use the Fahrenheit scale to measure temperatures, but most scientists use either the Celsius (°C) or Kelvin (K) temperature scale. (Remember: There’s no degree symbol associated with K.) Water boils at 100°C (373 K) and freezes at 0°C (273 K)
Here’s how to do some temperature conversions:
Fahrenheit to Celsius: °C = ⁵⁄₉(°F - 32)
Celsius to Fahrenheit: °F = ⁹⁄₅(°C) + 32
Celsius to Kelvin: K + °C + 273
Heat is not the same as temperature. When you measure the temperature of something, you’re measuring the average kinetic energy of the individual particles. Heat, on the other hand, is the amount of energy that goes from one substance to another.
The unit of heat in the SI system is the joule (J). Most people still use the metric unit of heat, the calorie (cal). Here’s the relationship between the two:
1 calorie = 4.184 joules
The calorie is a fairly small amount of heat: The amount it takes to raise the temperature of 1 gram of water 1°C. I often use the kilocalorie (kcal), which is 1,000 calories, as a convenient unit of heat. If you burn a large kitchen match completely, it produces about 1 kcal.
Chapter 2: What’s In an Atom?
Subatomic Particles
The atom is the smallest part of matter that represents a particular element. For quite a while, the atom was thought to be the smallest part of matter that could exist. But in the latter part of the 19th century and early part of the 20th, scientists discovered that atoms are composed of certain subatomic particles that no matter what the element, the same subatomic particles make up the atom. The number of the various subatomic particles is the only thing that varies.
Scientists now recognize that there are many subatomic particles. But to be successful in chemistry, you really only need to be concerned with three major subatomic particles:
Protons
Neutrons
Electrons
Table 2-1 summarizes the characteristics of these three subatomic particles. The masses of the subatomic particles are listed in two ways: grams and amu, which stands for atomic mass units. Expressing mass in amu is much easier than using the gram equivalent.
Atomic mass units are based on something called the carbon-12 scale, a worldwide standard that’s been adopted for atomic weights. By international agreement, a carbon atom that contains six protons and six neutrons has an atomic weight of exactly 12 amu, so 1 amu is defined as ¹⁄₁₂ of this carbon atom. Because the masses in grams of protons and neutrons are almost exactly the same, both protons and neutrons ae said to have a mass of 1 amu. Notice that the mass of an electron is much smaller than that of either a proton or neutron. It takes almost 2,000 electrons to equal the mass of a single proton.
Table 2-1 also shows the electrical charge associated with each subatomic particle. Matter can be electrically charged in one of two ways: positive or negative. The proton carries one unit of positive charge, the electron carries one unit of negative charge, and the neutron has no charge — it’s neutral.
Scientists have discovered through observation that objects with like charges, whether positive or negative, repel each other, and objects with unlike charges attract each other.
The atom itself has no charge. It’s neutral. (Well, actually, certain atoms can gain or lose electrons and acquire a charge, as I explain in the later section “Ions: Varying electrons.” Atoms that gain a charge, either positive or negative, are called ions.) So how can an atom be neutral if it contains positively charged protons and negatively charged electrons? The answer is that there are equal numbers of protons and electrons — equal numbers of positive and negative charges — so they cancel each other out.
The last column in Table 2-1 lists the location of the three subatomic particles. Protons and neutrons are located in the nucleus, a dense central core in the middle of the atom, and the electrons are located outside the nucleus (for details, see “Locating Those Electrons?” later in this chapter.)
Centering on the Nucleus
In 1911, Ernest Rutherford discovered that atoms have a nucleus — a center — containing protons. Scientists later discovered that the nucleus also houses the neutron.
The nucleus is very, very small and very, very dense when compared to the rest of the atom. Typically, atoms have diameters that measure around 10-10 meters (that’s small!). Nuclei are around 10-15 meters in diameter (that’s really small!). If the Superdome in New Orleans represented a hydrogen atom, the nucleus would be about the size of a pea.
The protons of an atom are all crammed together inside the nucleus. Now you may be thinking, “Okay, each proton carries a positive charge, and like charges repel each other. So if all the protons are repelling each other, why doesn’t the nucleus simply fly apart?” It’s the Force, Luke. Forces in the nucleus counteract this repulsion and and hold the nucleus together. Physicists call these forces nuclear glue. (Note: Sometimes this “glue” isn’t strong enough, and the nucleus does break apart. This process is called radioactivity, and I cover it in Chapter 4.)
Not only is the nucleus very small, but it also contains most of the mass of the atom. In fact, for all practical purposes, the mass of the atom is the sum of the masses of the protons and neutrons. (I ignore the minute mass of the electrons unless I’m doing very, very precise calculations.")
The sum of the number of protons plus the number of neutrons in an atom is called the mass number. And the number of neutrons in a particular atom is given a special name, the atomic number. Chemists commonly use the symbolization in Figure 2-1 to represent these amounts for a particular element
As Figure 2-1 shows, chemists use the placeholder X to represent the chemical symbol. You can find an element’s chemical symbol on the periodic table of elements. The placeholder Z represents the atomic number — the number of protons in the nucleus. And A represents the mass number, the sum of the number of protons plus neutrons. The mass number is listed in amu.
For example, you can represent a uranium atom that has 92 protons and a mass number of 238 as in Figure 2-2
You can find the number of neutrons in an atom by subtracting the atomic number (number of protons) from the mass number (protons plus neutrons). For instance, you know that uranium has an atomic number of 92 and a mass number of 238. So if you want to know the number of neutrons in uranium, all you have to do is subtract the atomic number (92 protons) from the mass number (238 protons plus neutrons). The answer shows that uranium has 146 neutrons.
But how many electrons does uranium have? Because the atom is neutral (it has no electrical charge), there must be equal numbers of positive and negative charges inside it, or equal numbers of protons and electrons. So there are 92 electrons in each uranium atom.
You can find both the element symbol and its atomic number on the periodic table, but the mass number for a particular element is not shown there. What is shown is the average element, taking into account the percentages of each found in nature. See the later section “Isotopes: Varying neutrons” for details on other forms of an element.
Locating Those Electrons
Many of the important topics in chemistry, such as chemical bonding, the shape of molecules, and so on, are based on where electrons in an atom are located. Simply saying that the electros are located outside the nucleus isn’t good enough; chemists need to have a much better idea of their location, so this section helps you figure out where you can find those pesky electrons.
The quantum mechanical model
Early models of the atom had electrons going around the nucleus in a random fashion. But as scientists discovered more about the atom, they found that this representation probably wasn’t accurate. Today, scientists use the quantum mechanical model, a highly mathematical model, to represent the structure of the atom.
This model is based on quantum theory, which says that matter also has properties associated with waves. According to quantum theory, it’s possible to know an electron’s exact position and momentum (speed and direction, multiplied by mass) at the same time. This is known as the uncertainty principle. So scientists had to develop the concept of orbitals (sometimes called electron clouds), volumes of space in which an electron is likely present. In other words, certainty was replaced with probability.
The quantum mechanical model of the atom uses complex shapes of orbitals. Without resorting to a lot of math (you’re welcome), this section shows you some aspects of this newest model of the atom.
Scientists introduced four numbers, called quantum numbers, to describe the characteristics of electrons and their orbitals. You’ll notice that they were named by top-rate techno-geeks.
Principal quantum number n
Angular momentum quantum number l
Magnetic quantum number ml
Spin quantum number ms
Table 2-2 summarizes the four quantum numbers. When they’re all put together, theoretical chemists have a pretty good description of the characteristics of a particular electron.
The principal quantum number n
The principal quantum number n describes the average distance of the orbital from the nucleus — and the energy of the electron in an atom. It can have only positive integer (whole number) values: 1, 2, 3, 4, and so on. The larger the value of n, the nigher the energy and the larger the orbital, or electron shell.
The angular momentum quantum number l
The angular momentum quantum l describes the shape of the orbital, and the shape is limited by the principal quantum number n: The angular momentum quantum number l can have positive integer values from 0 to n - 1. For example, if the n value is 3, three values are allowed for l: 0, 1, and 2.
The value of l defines the shape of the orbital, and the value of n defines the size.
Orbitals that have the same value of n but different values of l are called subshells. These subshells are given different letters to help chemists distinguish them from each other. Table 2-3 shows the letters corresponding to the different values of l.
When chemists describe one particular subshell in an atom, they can use both the n value and the subshell letter — 2p, 3d, and so on. Normally, a subshell value of 4 is the largest needed to describe a particular subshell. If chemists ever need a larger value, they can create subshell numbers and letters.
Figure 2-3 shows the shapes of the s, p, and d orbitals. In figure 2-3a, there are two s orbitals — one for energy level 1 (1s) and the other for energy level 2 (2s). S orbitals are spherical with the nucleus at the center. Notice that the 2s orbital is larger in diameter than the 1s orbital. In large atoms, the 1s orbital is nestled inside the 2s, just like the 2p is nestled inside the 3p.
Figure 2-3b shows the shapes of the p orbitals, and Figure 2-3c shows the shapes of the d orbitals. Notice that the shapes get progressively more complex.
The magnetic quantum number ml
The magnetic quantum number ml describes how the various orbitals are oriented in space. The value of ml depends on the value of l. The values allowed are integers from -l to 0 to +l. For example, if the value of l = 1 (p orbital — see Table 3-4), you can write three values for ml: -1, 0, and +1. This means that there are three different p subshells for a particular orbital. The subshells have the same energy but different orientations in space.
Figure 2-3b shows how the p orbitals are oriented in space. Notice that the three p orbitals correspond ml values of -1, 0, and +1, oriented along the x, y, and z axes.
The spin quantum number ms
The fourth and final quantum number is the spin quantum number ms. This one describes the direction the electron is spinning in a magnetic field, either clockwise or counter-clockwise. Only two values are allowed for ms: +½ or -½. For each subshell, there can only be two electrons, one with a spin of +½ and another with a spin of -½.
Putting the quantum numbers together
Table 2-4 summarizes the quantum numbers available for the first two energy levels
Table 2-4 shows that in energy level 1 (n = 1), there’s only an s orbital. There’s no p orbital because an l value of 1 (p orbital) is not allowed. And notice that there can be only two electrons in that 1s orbital (ms of +½ and -½). In fact, there can only be two electrons in any s orbital, whether it’s 1s or 5s.
Each time you move higher in a major energy level, you add another orbital type. So when you move from energy level 1 to energy level 2 (n = 2), there can be both s and p orbitals. If you write out the quantum numbers for energy level 3, you see s, p, and d orbitals.
Notice also that there are three subshells (ml) for the 2p orbital (see Figure 2-3b) and that each holds a maximum of two electrons. The three 2p subshells can hold a maximum of six electrons.
Energy level diagrams
Chemists find quantum numbers useful when they’re looking at chemical reactions and bonding (and those are things many chemists like to study). But they find two other representations for electrons — energy level diagrams and electron configurations — more useful and easier to work with.
Chemists use both of these things to represent which energy level, subshell, and orbital are occupied by electrons in any particular atom. Chemists use this information to predict what type of bonding will occur with a particular element and to show exactly which electrons are being used. These representations are also useful in showing why certain elements behave in similar ways.
In this section