By Fatskills Exam Guides Team — the exam nerds behind 28,500+ quizzes and 2.1M practice questions across 500+ global exams.
Atomic Structure Atoms are made up of three subatomic particles: protons, neutrons, and electrons. The protons have a positive charge and are located in the nucleus of the atom. Neutrons have a neutral charge and are also located in the nucleus. Electrons have a negative charge, are the smallest of the three particles, and are located in orbitals that surround the nucleus.
An atom is neutral if it has an equal number of electrons and protons. If an atom does not have an equal number of electrons and protons, it is an ion. When there are fewer electrons than protons, leaving it with a positive charge, it is termed a cation. When there are more electrons than protons, leaving a negative charge, it is an anion.
There are many levels of orbitals that surround the nucleus and house electrons. Each orbital has a distinct shape and three quantum numbers that characterize it: n, l, and ml. The n is the principal quantum number and is always a positive integer, or whole number. As n increases, the size of the orbital increases and the electrons are less tightly bound to the nucleus. The angular momentum quantum number is represented by l and defines the shape of the orbital.
There are four numerical values of I that correspond to four letter representations of the orbital: 0 or s, 1 or p, 2 or d, and 3 or f. Lastly, ml represents the magnetic quantum number and describes the orientation of the orbital in space. It has a value in between -l and l, including zero. Electrons exist outside the nucleus in energy levels, and with each increasing period, there is an additional energy level. Hydrogen and helium are the only elements that have only one energy level, which is an s orbital that can only hold two electrons. Their relative electron configurations are 1s1 and 1s2, respectively. Each orbital holds two electrons with opposite spins.
The second energy level can hold a total of eight electrons: two in the s orbital, and six in the p orbital, of which there are three, since electrons pair up: pxpypz. All alkali metals and alkali Earth metals have an electron configuration that ends in either s1 or s2. For example, magnesium ends in 3ss and cesium ends in 6s1. Groups 3A to 8A all end in the Xpy configuration, with X representing the energy level and y representing how for inland it is. For example, oxygen is 2p4. The third energy level has s, p, and d orbitals, and together can hold 18 electrons, since d can hold 10. D orbitals start with the transition metals, and even though it appears that they are in the fourth energy level, they actually are on the third. The energy level for the whole d block is actually the period number minus one.
Example: Mn: 1s22s22p63s23p64s23d5 The f block, which can hold 14 electrons, has a similar pattern for its preceding number, except two are subtracted from the period number. Structure Atoms are the most basic portion of an element that still retains its properties. All of the elements known to man are catalogued in the periodic table, a chart of elements arranged by increasing atomic number. The atomic number refers to the number of protons in an atom’s nucleus. It can be found either in the upper left-hand corner of the box or directly above an element’s chemical symbol on the periodic table. For example, the atomic number for hydrogen (H) is 1. The term “atomic mass” refers to the sum of protons and neutrons in an atom’s nucleus. The atomic mass can be found beneath an element’s abbreviation on the periodic table. For example, the average atomic mass of hydrogen (H) is 1.008. Because protons have a positive charge and neutrons have a neutral charge, an atom’s nucleus typically has a positive electrical charge. Electrons orbiting the nucleus have a negative charge. As a result, elements with equal numbers of protons and electrons have no net charge. Atomic Radii Atomic radius refers to the size of an atom. Going down a group, because each level represents an energy level, atomic radii increase. For example, when comparing lithium and potassium, potassium will have a greater atomic radius because it has two more energy levels. Contrary to what might be predicted, when going from left to right on the periodic table across a period, atomic radius actually decreases. This is due to proton pull. As each element gains a proton in its nucleus, it is like gaining a positive magnet in the middle. The more magnets in the middle, the stronger they will suck in any outer electrons.
Comparing carbon and fluorine, fluorine has a smaller atomic radius due to more (stronger) proton pull of the nucleus.
Ionic radius is a bit different. When sodium forms a cation, it loses an electron, losing an entire energy level. Thus, the ionic radii of cations are smaller than their elemental form. Anions, however, have a larger radius than their elemental form. This is because an extra electron is like adding a negative magnet to seven other negative magnets, resulting in a repulsive force that pushes them apart.
When atoms collide, in solutions, for example, they touch each other and then ricochet apart. The closest they come to each other is the combined length of each of their radii. This radius is known as van der Waals radius for each atom. When atoms bond with each other, however, they must come closer together than they would if they were just colliding. Their attraction also draws them closer together. The bonding atomic radius, or the covalent radius, is equal to one-half of the distance between the two atoms when they are bonded together. Within each group in the periodic table, the covalent radius increases from top to bottom. Within each period, the covalent radius decreases from left to right. Ions Atoms that have gained or lost electrons wind up having a net electrical charge and are termed ions. The following are the main ions related to human health:
For the majority of light atoms, the number of protons is similar to the number of neutrons. An isotope is a variation of an element having the same number of protons, but a different number of neutrons. For example, all isotopes of carbon (C) have six protons. Nevertheless, C-12 has six neutrons, C-13 has seven neutrons, and C-14 has eight neutrons. Some isotopes are radioactive and result in nuclear decay. Not all radioactive isotopes are harmful, and some are even useful to scientists and physicians. For example, C-14 is radioactive and can be used in the process of radiocarbon dating, which can be used to determine the age of organic remains. A radioactive isotope of gold (Au-198) can be utilized to treat ovarian, prostate, and brain cancer. Periodicity Periodicity refers to the repeating patterns, or trends, in the properties of elements. The atomic number and atomic structure are the key determinants of the properties of elements. During the mid-1800s, the Russian chemist Dmitri Mendeleev utilized the principal of periodicity to arrange elements in a manner similar to the modern periodic table. Mendeleev’s periodic table was arranged in rows according to increasing atomic mass and in columns according to similar chemical behavior. The modern periodic table is arranged in order of increasing atomic number, which is defined as the number of protons in an atom’s nucleus. Elements near each other are more similar than elements that are distant on the periodic table. Periodic Table The periodic table is a chart of all 118 known elements. The elements are organized according to their quantity of protons, also known as their atomic number, their electron configurations, and their chemical properties. The rows are called periods and the columns are called groups.
Groups have similar chemical behavior. For example, Group 8A is the noble gases, and because their outer electron shell is full, they are non-reactive. The closer the element is to having a full set of valence electrons, the more reactive it is. Group 1, the alkali metals, and Group 7, the halogens, are both highly reactive for that reason. The alkali metals form cations and lose their lone electron while halogens pick up an electron. In each box in the period table, an element’s symbol is the abbreviation in the center and its full name is located directly below that. The number in the top left corner is the atomic number and the atomic mass of the element is the number underneath. The atomic mass of the element is noted in atomic mass units, or amu, which represents the number of protons and neutrons combined. The number of protons defines the element, but the mass number can be different due to the existence of elements with a different number of neutrons, also called isotopes. The amu shown on the periodic table is a weighted mass (based on abundance) of all known isotopes of a particular element. Ionization Energy and Electron Affinity Energy is required to remove an electron from an atom or ion while it is in its ground state, usually as a gas. Ionization energy applies to all elements; it is just greater for those on the right side of the table. The ionization energy of an atom or ion is the minimum amount of energy required to do so. The greater the ionization energy, the harder it is to remove the electron. The energy required for the first electron is noted as I1, the second as I2, and then continues consecutively numbered for each successive electron. With each electron that is removed, the ionization energy of the element increases. I1 increases in elements across the periods and moving down the columns of the period table. In contrast to ionization energy, energy is released from an atom when an electron is added. This negative change in energy is known as electron affinity. Ionization energies are always positive, because atoms require more energy to let go of an electron than to add one. Looking at the periodic table, elements on the left, the alkali metals and alkali earth metals, have a low ionization energy. It takes little energy to peel off of their electron, because they have the least (weakest) proton pull in their period. Elements on the right, like halogens (gases), have an extremely high ionization energy, because they have a very strong proton pull pulling electrons in, so it would require even more energy to peel off an electron. Noble gases have an extremely high ionization energy. Going down a group, as energy levels increase, proton pull becomes weaker, so ionization energy decreases. Molecular Weight Elements can exist alone or combine to form compounds and molecules. Using the atomic mass, or atomic weights, that are noted on the periodic table, the molecular weight of a substance can be found by multiplying the subscript of each element by the atomic mass of that element and then adding together the weights of all of the elements in the molecule together.
For example, to find the molecular weight of one glucose molecule, C6H12O6, the atomic mass of carbon is multiplied by six, the atomic mass of hydrogen is multiplied by twelve, the atomic mass of oxygen is multiplied by six and then the three resulting numbers are added together. The answer is one glucose molecule has a molecular weight of 180 amu. Chemical Bonding Nomenclature and Formulas Chemical nomenclature refers to how compounds and substances are named, and the formulas are how they are written. Inorganic substances do not contain carbon, while organic ones do. Some common compounds have “household names,” like water, ammonia, and table salt in addition to their chemical formulas, whereas others are strictly referred to by their official nomenclature. There are a variety of conventions or standards by which chemists derive the names for chemical compounds. These systematic rules help create unity and establish conventions worldwide so that a given substance is referred to the same way around the world. The International Union of Pure and Applied Chemistry (IUPAC) has developed most of the chemical nomenclature used around the world and has published “color books” that serve as chemical dictionaries in a sense, listing the names of compounds based on their type.
For example, the Blue Book contains the names for organic compounds, while the Red Book lists inorganic ones.
For inorganic compounds, nomenclature may be established based on different factors such as the substance’s composition, arrangement, and type of bonding.
For example, ionic compounds, which are composed of a metal and non-metal bonded together (e.g., NaCl, sodium chloride, and Ca3N2, calcium nitride) adhere to the IUPAC’s standard conventions established for this type of compound. The monoatomic cation, which is positive, is listed first in the name and formula and retains the element’s name. In NaCl, for example, the cation is sodium, so it is written and said first in the compound and is simply called “sodium.” The monoatomic anion, chlorine, with a negative charge is written second and is given the suffix -ide, becoming ”chloride.” Therefore, compound NaCl, table salt, is called “sodium chloride.” Bonding Chemical bonding occurs between two or more atoms that are joined together. There are three types of chemical bonds: ionic, covalent, and metallic. The characteristics of the different bonds are determined by how electrons behave in a compound. Lewis structures were developed to help visualize the electrons in molecules; they are a method of writing a compound structure formula and including its electron composition. A Lewis symbol for an element consists of the element symbol and a dot for each valence electron. The dots are located on all four sides of the symbol, with a maximum of two dots per side, and eight dots, or electrons, total. The octet rule states that atoms tend to gain, lose, or share electrons until they have a total of eight valence electrons. Ionic Bonds Ionic bonds are formed from the electrostatic attractions between oppositely charged atoms. They result from the transfer of electrons from a metal on the left side of the periodic table to a nonmetal on the right side. The metallic substance often has low ionization energy and will transfer an electron easily to the nonmetal, which has a high electron affinity.
An example of this is the compound NaCl, which is sodium chloride, or table salt, where the Na atom transfers an electron to the Cl atom. Due to strong bonding, ionic compounds have several distinct characteristics. They have high melting and boiling points, and are brittle and crystalline. They are arranged in rigid, well-defined structures, which allow them to break apart along smooth, flat surfaces. The formation of ionic bonds is a reaction that is exothermic. In the opposite scenario, the energy it takes to break up a one mole quantity of an ionic compound is referred to as lattice energy, which is generally endothermic.
The Lewis structure for NaCl is written as follows: Covalent Bonds Covalent bonds are formed when two atoms share electrons, instead of transferring them as in ionic compounds. The atoms in covalent compounds have a balance of attraction and repulsion between their protons and electrons, which keeps them bonded together. Two atoms can be joined by single, double, or even triple covalent bonds. As the number of electrons that are shared increases, the length of the bond decreases. Covalent substances have low melting and boiling points, and are poor conductors of heat and electricity.
The Lewis structure for Cl2 is written as follows: Metallic Bonds Metallic bonds are formed by electrons that move freely through metal. They are the product of the force of attraction between electrons and metal ions. The electrons are shared by many metal cations and act like glue that holds the metallic substance together, similar to the attraction between oppositely-charged atoms in ionic substances, except the electrons are more fluid and float around the bonded metals and form a sea of electrons. Metallic compounds have characteristic properties that include strength, conduction of heat and electricity, and malleability. They can conduct electricity by passing energy through the freely moving electrons, creating a current. These compounds also have high melting and boiling points. Lewis structures are not common for metallic structures because of the free-roaming ability of the electrons. Hydrogen Bonding Hydrogen bonds are temporary and weak. They typically occur between two partial, opposite electrical charges. For example, hydrogen bonds form when a hydrogen (H) atom is in the vicinity of nitrogen (N), fluorine (F), or oxygen (O) atoms. These partial electrical charges are called dipoles, and are caused by the unequal sharing of electrons between covalent bonds. Water is the most prevalent molecule that forms hydrogen bonds. Hydrogen bonds contribute to the adhesiveness and cohesiveness properties of molecules like water. Adhesiveness confers glue-like properties to molecules which ensure they stick or connect more easily with other molecules, similar to wetting a suction cup before sticking it to a surface. Cohesiveness refers to a molecule’s ability to form hydrogen bonds with itself. For example, the cohesiveness of water is the reason why it has a high boiling point, which is a physical property. Reactions and Reaction Mechanisms Chemical reactions are represented by chemical equations. The equations help to explain how the molecules change during the reaction. For example, when hydrogen gas (H2) combines with oxygen gas (O2), two molecules of water are formed.
The equation is written as follows, where the “+” sign means reacts with and the “→” means produces: 2 H2 + O2 → 2 H2O
Two hydrogen molecules react with an oxygen molecule to produce two water molecules. In all chemical equations, the quantity of each element on the reactant side of the equation should equal the quantity of the same element on the product side of the equation due to the law of conservation of matter. If this is true, the equation is described as balanced.
To figure out how many of each element there is on each side of the equation, the coefficient of the element should be multiplied by the subscript next to the element. Coefficients and subscripts are noted for quantities larger than one.
The coefficient is the number located directly to the left of the element. The subscript is the small-sized number directly to the right of the element. In the equation above, on the left side, the coefficient of the hydrogen is two and the subscript is also two, which makes a total of four hydrogen atoms. Using the same method, there are two oxygen atoms. On the right side, the coefficient two is multiplied by the subscript in each element of the water molecule, making four hydrogen atoms and two oxygen atoms. This equation is balanced because there are four hydrogen atoms and two oxygen atoms on each side. The states of the reactants and products can also be written in the equation: gas (g), liquid (l), solid (s), and dissolved in water (aq). If they are included, they are noted in parentheses on the right side of each molecule in the equation. Types of Chemical Reactions Chemical reactions are characterized by a chemical change in which the starting substances, or reactants, differ from the substances formed, or products. Chemical reactions may involve a change in color, the production of gas, the formation of a precipitate, or changes in heat content.
The following are the five basic types of chemical reactions:
Oxidation-Reduction Reactions Oxidation-reduction reactions, or redox reactions, are chemical reactions in which electrons are transferred from one compound to another. A helpful mnemonic device to remember the basic properties of redox reactions is “LEO the lion says GER”, in which “LEO” means Lose Electrons Oxidation and “GER” stands for Gain Electrons Reduction. A reducing agent is an electron donor and an oxidizing agent is an electron acceptor. An example of an oxidation-reduction reaction is the electrochemical cell, which is the basis for batteries. The electrochemical cell comprises two different cells, each equipped with a separate conductor, and a salt bridge. The salt bridge isolates reactants, but maintains the electric current. The cathode is the electrode that is reduced, and the anode is the electrode that is oxidized. Anions and cations carry electrical current within the cell and electrons carry current within the electrodes. Balancing Chemical Reactions Chemical reactions are conveyed using chemical equations. As mentioned, chemical equations must be balanced with equivalent numbers of atoms for each type of element on each side of the equation. Antoine Lavoisier, a French chemist, was the first to propose the Law of Conservation of Mass for the purpose of balancing a chemical equation. The law states, “Matter is neither created nor destroyed during a chemical reaction.” The reactants are located on the left side of the arrow, while the products are located on the right side of the arrow. Coefficients are the numbers in front of the chemical formulas. Subscripts are the numbers to the lower right of chemical symbols in a formula. To tally atoms, one should multiply the formula’s coefficient by the subscript of each chemical symbol.
For example, the chemical equation 2 H2 + O2 2H2O is balanced. For H, the coefficient of 2 multiplied by the subscript 2 = 4 hydrogen atoms. For O, the coefficient of 1 multiplied by the subscript 2 = 2 oxygen atoms. Coefficients and subscripts of 1 are understood and never written. When known, the form of the substance is noted with (g) for gas, (s) for solid, (l) for liquid, or (aq) for aqueous. Balancing Redox Reactions Keep track of oxidation states or oxidation numbers to ensure the chemical equation is balanced. Oxidation numbers are assigned to each atom in a neutral substance or ion. For ions made up of a single atom, the oxidation number is equal to the charge of the ion. For atoms in their original elemental form, the oxidation number is always zero. Each hydrogen atom in an H2 molecule, for example, has an oxidation number of zero. The sum of the oxidation numbers in a molecule should be equal to the overall charge of the molecule. If the molecule is a positively-charged ion, the sum of the oxidation number should be equal to overall positive charge of the molecule. In ionic compounds that have a cation and anion joined, the sum of the oxidation numbers should equal zero. All chemical equations must have the same number of elements on each side of the equation to be balanced. Redox reactions have an extra step of counting the electrons on both sides of the equation to be balanced. Separating redox reactions into oxidation reactions and reduction reactions is a simple way to account for all of the electrons involved. The individual equations are known as half-reactions. The number of electrons lost in the oxidation reaction must be equal to the number of electrons gained in the reduction reaction for the redox reaction to be balanced.
The oxidation of tin (Sn) by iron (Fe) can be balanced by the following half-reactions: Oxidation: Sn2+ → Sn4+ + 2e- Reduction: 2Fe3+ + 2e- → 2Fe2+ Complete redox reaction: Sn2+ + 2Fe3+ → Sn4+ + 2Fe2+ Equilibrium and Reaction Rates Reaction Rates The rate of a reaction is the measure of the change in concentration of the reactants or products over a certain period of time. Many factors affect how quickly or slowly a reaction occurs, such as concentration, pressure, or temperature. As the concentration of a reactant increases, the rate of the reaction also increases, because the frequency of collisions between elements increases. High-pressure situations for reactants that are gases cause the gas to compress and increase the frequency of gas molecule collisions, similar to solutions with higher concentrations. Reactions rates are then increased with the higher frequency of gas molecule collisions. Higher temperatures usually increase the rate of the reaction, adding more energy to the system with heat, and increasing the frequency of molecular collisions. Equilibrium Equilibrium is described as the state of a system when no net changes occur. Chemical equilibrium occurs when opposing reactions occur at equal rates. In other words, the rate of reactants forming products is equal to the rate of the products breaking down into the reactants—the concentration of reactants and products in the system does not change. Although the concentrations are not changing, the forward and reverse reactions are likely still occurring. This type of equilibrium is called a dynamic equilibrium. In situations where all reactions have ceased, a static equilibrium is reached. Chemical equilibriums are also described as homogeneous or heterogeneous. Homogeneous equilibrium involves substances that are all in the same phase, while heterogeneous equilibrium means the substances are in different phases when equilibrium is reached. When a reaction reaches equilibrium, the conditions of the equilibrium are described by the following equation, based on the chemical equation aA + bB ↔ cC + dD: This equation describes the law of mass action. It explains how the reactants and products react during dynamic equilibrium. Kc is the equilibrium constant and it is obtained when molarity values are put into the equation for the reactants and products. It is important to note that Kc is only dependent on the stoichiometry of the equation. If Kc is greater than 1, the equilibrium occurs when there are more products generated; the equilibrium lies to the right. If Kc is less than 1, the equilibrium occurs when there are more reactants generated and the equilibrium is to the left. Similar to finding Kc, the quantity of reactants and products, as well as the direction of the reaction, can be determined at any point of time by finding Q, the reaction quotient. Qc is substituted for the Kc in the equation above. If Q is less than K, the concentration of the reactants is too large and the concentration of the products is too small, so the reaction must move from left to right to achieve equilibrium. If Q is equal to K, the system is at equilibrium. If Q is greater than K, the concentration of the products is too large and the concentration of the reactants is too small; the reaction must move from right to left to reach equilibrium. Catalysts Catalysts are substances that accelerate the speed of a chemical reaction. A catalyst remains unchanged throughout the course of a chemical reaction. In most cases, only small amounts of a catalyst are needed. Catalysts increase the rate of a chemical reaction by providing an alternate path that requires less activation energy. Activation energy refers to the amount of energy necessary for the initiation of a chemical reaction. Catalysts can be homogeneous or heterogeneous. Catalysts in the same phase of matter as its reactants are homogeneous, while catalysts in a different phase than the reactants are heterogeneous. It is important to remember catalysts are selective. They do not accelerate the speed of all chemical reactions, but catalysts do accelerate specific chemical reactions. Enzymes Enzymes are a class of catalysts instrumental in biochemical reactions, and in most, if not all, instances, they are proteins. Like all catalysts, enzymes increase the rate of a chemical reaction by providing an alternate path that requires less activation energy. Enzymes catalyze thousands of chemical reactions in the human body. Enzymes possess an active site, which is the part of the molecule that binds the reacting molecule, or substrate. The “lock and key” analogy is used to describe the substrate key fitting precisely into the active site of the enzyme lock to form an enzyme-substrate complex. Many enzymes work in tandem with cofactors or coenzymes to catalyze chemical reactions. Cofactors can be either inorganic (not containing carbon) or organic (containing carbon). Organic cofactors can be either coenzymes or prosthetic groups tightly bound to an enzyme. Coenzymes transport chemical groups from one enzyme to another. Within a cell, coenzymes are continuously regenerating, and their concentrations are held at a steady state. Several factors—including temperature, pH, and concentrations of the enzyme and substrate—can affect the catalytic activity of an enzyme. For humans, the optimal temperature for peak enzyme activity is approximately body temperature at 98.6 0F, while the optimal pH for peak enzyme activity is approximately 7-8. Increasing the concentrations of either the enzyme or substrate will also increase the rate of reaction, up to a certain point. The activity of enzymes can be regulated. One common type of enzyme regulation is termed feedback inhibition, which involves the product of the pathway inhibiting the catalytic activity of the enzyme involved in its manufacture. Stoichiometry Stoichiometry investigates the quantities of chemicals that are consumed and produced in chemical reactions. Chemical equations are made up of reactants and products; stoichiometry helps elucidate how the changes from reactants to products occur, as well as how to ensure the equation is balanced. Limiting Reactants Chemical reactions are limited by the amount of starting material—or reactants—available to drive the process forward. The reactant that is present in the smallest quantity in a reaction is called the limiting reactant. The limiting reactant is completely consumed by the end of the reaction. The other reactants are called excess reactants.
For example, gasoline is used in a combustion reaction to make a car move and is the limiting reactant of the reaction. If the gasoline runs out, the combustion reaction can no longer take place, and the car stops. Reaction Yield The quantity of product that should be produced after using up all of the limiting reactant can be calculated, and is called the theoretical yield of the reaction. Since the reactants do not always act as they should, the actual amount of resulting product is called the actual yield. The actual yield is divided by the theoretical yield and then multiplied by 100 to find the percent yield for the reaction. Solution Stoichiometry Solution stoichiometry deals with quantities of solutes in chemical reactions that occur in solutions. The quantity of a solute in a solution can be calculated by multiplying the molarity of the solution by the volume. Similar to chemical equations involving simple elements, the number of moles of the elements that make up the solute should be equivalent on both sides of the equation. When the concentration of a particular solute in a solution is unknown, a titration is used to determine that concentration. In a titration, the solution with the unknown solute is combined with a standard solution, which is a solution with a known solute concentration. The point at which the unknown solute has completely reacted with the known solute is called the equivalence point. Using the known information about the standard solution, including the concentration and volume, and the volume of the unknown solution, the concentration of the unknown solute is determined in a balanced equation.
For example, in the case of combining acids and bases, the equivalence point is reached when the resulting solution is neutral. HCl, an acid, combines with NaOH, a base, to form water, which is neutral, and a solution of Cl- ions and Na+ ions. Before the equivalence point, there is an unequal number of cations and anions and the solution is not neutral. Kinetic Theory The Kinetic Theory of Matter states that matter is composed of a large number of small particles (specifically, atoms and molecules) that are in constant motion. The distance between the separations in these particles determines the state of the matter: solid, liquid, or gas. In gases, the particles have a large separation and no attractive forces. In liquids, there is moderate separation between particles and some attractive forces to form a loose shape. Solids have almost no separation between their particles, which gives them a defined, set shape. The constant movement of particles causes them to bump into each other, thus allowing the particles to transfer energy among one another. This bumping and transferring of energy helps explain the transfer of heat and the relationship between pressure, volume, and temperature.
According to kinetic molecular theory: - Gas particles have Brownian, or random, motion (consistent with the second law of thermodynamics). - Gas particles travel in a straight line until they hit another object and change direction. - The volume of a gas particle is virtually zero. - There are no attractive or repulsive forces between gas particles, and there are no unaccounted chemical or physical interactions between particles. - There are elastic collisions between the gas and its container and other particles; particles do not slow down, and they are constantly moving. - Movement, or kinetic energy, is a function of temperature. As temperature increases, kinetic energy increases. Average kinetic energy is equal to temperature in Kelvin.
Collision theory is closely related to kinetic molecular theory. Based on the fact that as temperature increases, movement increases, the premise of collision theory is that reactions happen faster as temperature increases due to the increases in reactant collisions with each other and with catalysts. Of note, from a biological perspective, there is a limit to this temperature increase because, if it is a reaction involving biological protein catalysts (enzymes), the enzyme will eventually denature, “cook,” and lose functionality. States of Matter Chemistry is the study of matter, including its properties and behavior. Matter is the material that the universe is made of; it is any object that occupies space and has mass. Despite the diversity of items found in the universe, matter is comprised of only about 100 substances, known as elements. Elements cannot be broken down into simpler substances. Hydrogen and oxygen are two examples of elements. When different elements join together, they form compounds. Water is a compound made from hydrogen and oxygen. Atoms and molecules are among the smallest forms of matter. Matter can be found in three different states: gas, liquid, or solid.
Gas is a state that does not have a fixed volume or shape. It can expand to fill large containers or compress to fill smaller ones. Gas molecules are far apart from each other and float around at high speeds. They collide with each other and the container they are in, filling the container uniformly. When the gas is compressed, the space between the molecules decreases and the frequency of collisions between them increases. A liquid has an exact volume. It molds to the shape of the container that holds it. Liquid molecules are close together but still move rapidly. They cannot be compressed and slide over each other easily when liquids are poured. Solids have a definitive shape and volume. Similar to liquids, solids cannot be compressed. The molecules are packed together tightly in a specific arrangement that does not allow for much movement. The physical and chemical properties of matter can help distinguish different substances. Physical properties include color, odor, density, and hardness. These are properties that can be observed without changing the substance’s identity or composition. When a substance undergoes a physical change, its physical appearance changes but its composition remains the same.
Chemical properties are those that describe the way a substance might change to form another substance. Examples of chemical properties are flammability, toxicity, and ability to oxidize. These properties are observed by changing the environment of the substance and seeing how the substance reacts. A substance’s composition is changed when it undergoes a chemical change. Many properties of matter can be measured quantitatively, meaning the measurement is associated with a number. When a property of matter is represented by a number, it is important to include the unit of measure, otherwise the number is meaningless. For example, saying a pencil measures 10 is meaningless. It could be referring to 10 of something very short or 10 of something very long. The correct measurement notation would be 10 centimeters, because a centimeter has a designated length. Other examples of properties of matter that can be measured quantitatively are mass, time, and temperature, among others. Gas Laws The Ideal Gas Law states that pressure, volume, and temperature are all related through the equation: PV = nRT, where P is pressure, V is volume, n is the amount of the substance in moles, R is the gas constant, and T is temperature. Through this relationship, volume and pressure are both proportional to temperature, but pressure is inversely proportional to volume. Therefore, if the equation is balanced and the volume decreases in the system, pressure needs to increase proportionally to keep both sides of the equation balanced. In contrast, if the equation is unbalanced and the pressure increases, then the temperature would also increase, since pressure and temperature are directly proportional. Causes and Effects of Changes in State When environmental changes occur, such as temperature or pressure changes, one state of matter can convert to another. States of matter are able to undergo phase transitions. Vaporization refers to the transformation of a solid or liquid into a gas. There are two types of vaporization—evaporation and boiling. Evaporation is a surface phenomenon and involves the conversion of a liquid into a gas below the boiling temperature at a given pressure. Evaporation is also an important component of the water cycle. Boiling occurs below the surface and involves the conversion of liquid into a gas at or above the boiling temperature. Condensation represents the conversion of a gas into a liquid. It is the reverse of evaporation and an important process in the water cycle. It is a crucial component of distillation.
There is one other state of matter called plasma, which is seen in lightning, television screens, and neon lights. Plasma is most commonly converted from the gas state at extremely high temperatures.
The amount of energy needed to change matter from one state to another is labeled by the terms for phase changes. For example, the temperature needed to supply enough energy for matter to change from a liquid to a gas is called the heat of vaporization. When heat is added to matter in order to cause a change in state, there will be an increase in temperature until the matter is about to change its state. During its transition, all of the added heat is used by the matter to change its state, so there is no increase in temperature. Once the transition is complete, then the added heat will again yield an increase in temperature. Each state of matter is considered to be a phase, and changes between phases are represented by phase diagrams. These diagrams show the effects of changes in pressure and temperature on matter. The states of matter fall into areas on these charts called heating curves. Solutions Concentration (pH) One mole is the amount of matter contained in 6.02 x 1023 of any object, such as atoms, ions, or molecules. It is a useful unit of measure for items in large quantities. This number is also known as Avogadro’s number. One mole of 12C atoms is equivalent to 6.02 x 1023 12C atoms. Avogadro’s number is often written as an inverse mole, or as 6.02 x 1023/mol. Molar Mass A mole is always the same number, equivalent to Avogadro’s number. The molar mass of a substance is the mass in grams of one mole of molecules of that substance. It is numerically equivalent to the molecular weight of the substance. The molecular weight of glucose (C6H12O6) is 180 amu. Therefore, the molar mass of glucose is 180 grams per mole, written as 180 g/mol. In other words, one mole of glucose, or 6.02 x 1023 molecules of glucose, has a mass of 180 grams. Two substances with different molecular weights will have two different molar masses. Compared to glucose, O2 has a molecular weight of 32 amu, which is less than that of glucose. So, the molar mass of O2 is 32 g/mol. One mole of O2 has an equivalent number of molecules as one mole of glucose, but it weighs less. A simple calculation can determine how many moles are in a certain number of grams of a substance. The amount of substance in grams is divided by the molar mass of that substance and the result is the number of moles of the substance. Similarly, to convert the number of moles of a substance to the number of grams, multiply the number of moles by the molar mass of the substance. The result is the number of grams of the substance. It is also possible to calculate the number of molecules in a certain number of grams of substance using Avogadro’s number and the molar mass. The number of known grams is divided by the molar mass and then multiplied by Avogadro’s number. The result is the number of molecules in the starting number of grams. Molarity Molarity is the concentration of a solution. It is based on the number of moles of solute in one liter of solution and is written as the capital letter M. A 1.0 molar solution, or 1.0 M solution, has one mole of solute per liter of solution. The molarity of a solution can be determined by calculating the number of moles of the solute and dividing it by the volume of the solution in liters. The resulting number is the mol/L or M for molarity of the solution. Ionic solutions can also be described by molarity values. Since ionic compounds dissolve in solution, the chemical formula of the compound can be used to determine the relative concentrations of the ions in the solution. For example, in a 1.0 M solution of NaCl, there is 1.0 M Na+ ions and 1.0 M Cl- ions. In a 1.0 M solution of Na2SO4, there are two Na+ ions (2.0 M) for each individual SO42- ion (1.0 M). pH Scale pH refers to the power or potential of hydrogen atoms and is used as a scale for a substance’s acidity. In chemistry, pH represents the hydrogen ion concentration (written as [H+]) in an aqueous, or watery, solution. The hydrogen ion concentration, [H+], is measured in moles of H+ per liter of solution. The pH scale is a logarithmic scale used to quantify how acidic or basic a substance is. pH is the negative logarithm of the hydrogen ion concentration: pH = -log [H+]. A one-unit change in pH correlates with a ten-fold change in hydrogen ion concentration. The pH scale typically ranges from 0 to 14, although it is possible to have pHs outside of this range. Pure water has a pH of 7, which is considered neutral. pH values less than 7 are considered acidic, while pH values greater than 7 are considered basic, or alkaline. Generally speaking, an acid is a substance capable of donating hydrogen ions, while a base is a substance capable of accepting a hydrogen ion. A buffer is a molecule that can act as either a hydrogen ion donor or acceptor. Buffers are crucial in the blood and body fluids, and prevent the body’s pH from fluctuating into dangerous territory. pH can be measured using a pH meter, test paper, or indicator sticks. Solubility A solution is a homogenous mixture of more than one substance. A solute is another substance that can be dissolved into a substance called a solvent. If only a small amount of solute is dissolved in a solvent, the solution formed is said to be diluted. If a large amount of solute is dissolved into the solvent, then the solution is said to be concentrated. For example, water from a typical, unfiltered household tap is diluted because it contains other minerals in very small amounts. If more solute is being added to a solvent, but not dissolving, the solution is called saturated.
For example, when hummingbirds eat sugar-water from feeders, they prefer it as sweet as possible. When trying to dissolve enough sugar (solute) into the water (solvent), there will be a point where the sugar crystals will no longer dissolve into the solution and will remain as whole pieces floating in the water. At this point, the solution is considered saturated and cannot accept more sugar. This level, at which a solvent cannot accept and dissolve any more solute, is called its saturation point. In some cases, it is possible to force more solute to be dissolved into a solvent, but this will result in crystallization. The state of a solution on the verge of crystallization, or in the process of crystallization, is called a supersaturated solution. This can also occur in a solution that seems stable, but if it is disturbed, the change can begin the crystallization process.
Although the terms dilute, concentrated, saturated, and supersaturated give qualitative descriptions of solutions, a more precise quantitative description needs to be established for the use of chemicals. This holds true especially for mixing strong acids or bases. The method for calculating the concentration of a solution is done through finding its molarity. In some instances, such as environmental reporting, molarity is measured in parts per million (ppm). Parts per million, is the number of milligrams of a substance dissolved in one liter of water.
To find the molarity, or the amount of solute per unit volume of solution, for a solution, the following formula is used: In this formula, c is the molarity (or unit moles of solute per volume of solution), n is the amount of solute measured in moles, and V is the volume of the solution, measured in liters.
Example: What is the molarity of a solution made by dissolving 2.0 grams of NaCl into enough water to make 100 mL of solution? To solve this, the number of moles of NaCl needs to be calculated: First, to find the mass of NaCl, the mass of each of the molecule’s atoms is added together as follows: 23.0g (Na) + 35.5g (Cl) = 58.8g NaCl Next, the given mass of the substance is multiplied by one mole per total mass of the substance: 2.0g NaCl × (1 mol NaCl/58.5g NaCl) = 0.034 mol NaCl Finally, the moles are divided by the number of liters of the solution to find the molarity: (0.034 mol NaCl)/(0.100L) = 0.34 M NaCl
To prepare a solution of a different concentration, the mass solute must be calculated from the molarity of the solution. This is done via the following process: Example: How would you prepare 600.0 mL of 1.20 M solution of sodium chloride? To solve this, the given information needs to be set up: 1.20 M NaCl = 1.20 mol NaCl/1.00 L of solution 0.600 L solution × (1.20 mol NaCl/1.00 L of solution) = 0.72 moles NaCl 0.72 moles NaCl × (58.5g NaCl/1 mol NaCl) = 42.12 g NaCl This means that one must dissolve 42.12 g NaCl in enough water to make 600.0 L of solution. Factors Affecting the Solubility of Substances and the Dissolving Process Certain factors can affect the rate in dissolving processes. These include temperature, pressure, particle size, and agitation (stirring). As mentioned, the ideal gas law states that PV = nRT, where P equals pressure, V equals volume, and T equals temperature.
If the pressure, volume, or temperature are affected in a system, it will affect the entire system. Specifically, if there is an increase in temperature, there will be an increase in the dissolving rate. An increase in the pressure can also increase the dissolving rate. Particle size and agitation can also influence the dissolving rate, since all of these factors contribute to the breaking of intermolecular forces that hold solute particles together. Once these forces are broken, the solute particles can link to particles in the solvent, thus dissolving the solute. A solubility curve shows the relationship between the mass of solute that a solvent holds at a given temperature. If a reading is on the solubility curve, the solvent is full (saturated) and cannot hold anymore solute. If a reading is above the curve, the solvent is unstable (supersaturated) from holding more solute than it should. If a reading is below the curve, the solvent is unsaturated and could hold more solute. If a solvent has different electronegativities, or partial charges, it is considered to be polar. Water is an example of a polar solvent. If a solvent has similar electronegativities, or lacking partial charges, it is considered to be non-polar. Benzene is an example of a non-polar solvent.
Polarity status is important when attempting to dissolve solutes. The phrase “like dissolves like” is the key to remembering what will happen when attempting to dissolve a solute in a solvent. A polar solute will dissolve in a like, or polar solvent. Similarly, a non-polar solute will dissolve in a non-polar solvent. When a reaction produces a solid, the solid is called a precipitate. A precipitation reaction can be used for removing a salt (an ionic compound that results from a neutralization reaction) from a solvent, such as water. For water, this process is called ionization. When a solute is added to a solvent to lower the freezing point of the solvent, it is called freezing point depression. This is a useful process, especially when applied in colder temperatures. For example, the addition of salt to ice in winter allows the ice to melt at a much lower temperature, thus creating safer road conditions for driving. Unfortunately, the freezing point depression from salt can only lower the melting point of ice so far and is ineffectual when temperatures are too low. This same process, with a mix of ethylene glycol and water, is also used to keep the radiator fluid (antifreeze) in an automobile from freezing during the winter. Acids-Base Theories Acids and bases are defined in many different ways. An acid can be described as a substance that increases the concentration of H+ ions when it is dissolved in water, as a proton donor in a chemical equation, or as an electron-pair acceptor. A base can be a substance that increases the concentration of OH- ions when it is dissolved in water, accepts a proton in a chemical reaction, or is an electron-pair donor. Water can act as either an acid or a base. When mixed with an acid, water can accept a proton and become an H3O+ ion. When mixed with a base, water can donate a proton and become an OH- ion. Sometimes water molecules donate and accept protons from each other; this process is called autoionization. The chemical equation is written as follows: H2O + H2O → OH- + H3O+. Strength of Acids and Bases Acids and bases are characterized as strong, weak, or somewhere in between. Strong acids and bases completely or almost completely ionize in aqueous solution. The chemical reaction is driven completely forward, to the right side of the equation, where the acidic or basic ions are formed. Weak acids and bases do not completely disassociate in aqueous solution. They only partially ionize, and the solution becomes a mixture of the acid or base, water, and the acidic or basic ions. Strong acids are complemented by weak bases, and vice versa. A conjugate acid is an ion that forms when its base pair gains a proton. For example, the conjugate acid NH4+ is formed from the base NH3. The conjugate base that pairs with an acid is the ion that is formed when an acid loses a proton. NO2- is the conjugate base of the acid HNO2. Nuclear Chemistry: Radioisotopes Nuclear chemistry is the study of reactions in which the nuclei of atoms are transformed and their identities are changed. These reactions can involve large changes in energy—much larger than the energy changes that occur when chemical bonds between atoms are made or broken. Nuclear chemistry is also used to create electricity. Nuclear reactions are described by nuclear equations, which have different notations than regular chemical equations. Nuclear equations are written as follows, with the superscript being the mass number and the subscript being the atomic number for each element: The equation describes the spontaneous decomposition of uranium into thorium and helium via alpha decay. When this happens, the process is referred to as nuclear decay. Similar to chemical equations, nuclear equations must be balanced on each side; the sum of the mass numbers and the sum of the atomic numbers should be equal on both sides of the equation. In some cases, the nucleus of an atom is unstable and constantly emits particles due to this instability. These atoms are described as radioactive and the isotopes are referred to as radioisotopes. There are three types of radioactive decay that occur most frequently: alpha (α), beta (β), and gamma (γ). Alpha radiation is emitted when a nucleus releases a stream of alpha particles, which are helium-4 nuclei. Beta radiation occurs when a stream of high-speed electrons is emitted by an unstable nucleus. The beta particles are often noted as β–. Gamma radiation occurs when the nucleus emits high-energy photons. In gamma radiation, the atomic number and the mass remain the same for the unstable nucleus. This type of radiation represents a rearrangement of an unstable nucleus into a more stable one and often accompanies other types of radioactive emission. Radioactive decay is often described in terms of its half-life, which is the time that it takes for half of the radioactive substance to react. For example, the radioisotope strontium-90 has a half-life of 28.8 years. If there are 10 grams of strontium-90 to start with, after 28.8 years, there would be 5 grams left. There are two distinct types of nuclear reactions: fission and fusion reactions. Both involve a large energy release. In fission reactions, a large atom is split into two or more smaller atoms. The nucleus absorbs slow-moving neutrons, resulting in a larger nucleus that is unstable. The unstable nucleus then undergoes fission. Nuclear power plants depend on nuclear fission reactions for energy. Fusion reactions involve the combination of two or more lighter atoms into a larger atom. Fusion reactions do not occur in Earth’s nature due to the extreme temperature and pressure conditions required to make them happen. Fusion products are generally not radioactive. Fusion reactions are responsible for the energy that is created by the Sun.
Join 4M+ learners. Unlock unlimited quizzes, wrong-answer tracking, flashcards + reminders, study guides, and 1-on-1 challenges.