Atoms and the Periodic Table PDF

I. ATOMS AND THE PERIODIC TABLE The Atomic Theory Dalton’s work marked the beginning of the modern era of chemistry. The hypotheses about the nature of matter on which Dalton’s atomic theory is based can be summarized as follows: 1. Elements are composed of extremely small particles calle...

I. ATOMS AND THE PERIODIC TABLE The Atomic Theory Dalton’s work marked the beginning of the modern era of chemistry. The hypotheses about the nature of matter on which Dalton’s atomic theory is based can be summarized as follows: 1. Elements are composed of extremely small particles called atoms. 2. All atoms of a given element are identical, having the same size, mass, and chemical properties. The atoms of one element are different from the atoms of all other elements. 3. Compounds are composed of atoms of more than one element. In any compound, the ratio of the numbers of atoms of any two of the elements present is either an integer or a simple fraction. 4. A chemical reaction involves only the separation, combination, or rearrangement of atoms; it does not result in their creation or destruction. Atoms - 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 and 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. Three major subatomic particles: 1. Proton- positively charged particles (Plum-pudding Model) 2. Electrons- negatively charged particles (Cathode Ray Experiment) 3. Neutrons- electrically neutral particles having a mass slightly greater than that of protons Atomic Number (Z) - the number of protons in the nucleus of each atom of an element. Mass Number (A) - the total number of neutrons and protons present in the nucleus of an atom of an element. For example, if the mass number of a particular boron atom is 12 and the atomic number is 5 (indicating 5 protons in the nucleus), then the number of neutrons is 12 – 5 = 7. Note that all three quantities (atomic number, number of neutrons, and mass number) must be positive integers, or whole numbers. Isotopes - Atoms that have the same atomic number (Z) but different mass numbers (A) are called isotopes - All neon atoms have 10 protons in their nuclei, and most have 10 neutrons as well. A very few neon atoms, however, have 11 neutrons and some have 12. We can represent these three different types of neon atoms as Nuclear Stability One of Becquerel’s students, Marie Curie, suggested the name radioactivity to describe this spontaneous emission of particles and/or radiation. Since then, any element that spontaneously emits radiation is said to be radioactive. Three types of rays are produced by the decay, or breakdown, of radioactive substances such as uranium. Two of the three are deflected by oppositely charged metal plates. Alpha (a) rays consist of positively charged particles, called a particles, and therefore are deflected by the positively charged plate. Beta (b) rays, or b particles, are electrons and are deflected by the negatively charged plate. The third type of radioactive radiation consists of high-energy rays called gamma (g) rays. Like X rays, g rays have no charge and are not affected by an external field. Average Atomic Mass The sum of the masses of its isotopes, each multiplied by its mass abundance (the decimal associated with percent of atoms that are a given isotopes) Example: The mass spectrum of a particular sample of carbon shows that 98.93% of the carbon atoms are carbon-12 with a mass of exactly 12 u; the rest are carbon-13 atoms with a mass of 13.0033548378 u. Therefore: Mole In the SI system the mole (mol) is the amount of a substance that contains as many elementary entities (atoms, molecules, or other particles) as there are atoms in exactly 12 g (or 0.012 kg) of the carbon-12 isotope. The actual number of atoms in 12 g of carbon-12 is determined experimentally. This number is called Avogadro’s number (N A ), in honor of the Italian scientist Amedeo Avogadro. The currently accepted value is Molar Mass Molar mass ( m ), defined as the mass (in grams or kilograms) of 1 mole of units (such as atoms or molecules) of a substance. Example We know the molar mass of carbon-12 is 12.00 g and there are 6.022 x 1023 carbon-12 atoms in 1 mole of the substance; therefore, the mass of one carbon-12 atom is given by II. QUANTUM THEORY AND THE ELECTRONIC STRUCTURE OF ATOMS Quantum Theory - Max Plank (1900)- discovered that atoms and molecules emit energy only in certain discrete quantities, or quanta - Classical physics places no limitations on the amount of energy a system may possess, whereas quantum theory limits this energy to a set of specific values. The difference between any two allowed energies of a system also has a specific value, called a quantum of energy. This means that when the energy increases from one allowed value to another, it increases by a tiny jump, or quantum. - Physicists had always assumed that energy is continuous and that any amount of energy could be released in a radiation process. Nature of Light - Heinrich Hertz (1888)- that when light strikes the surface of certain metals, electrons are ejected. This phenomenon is called the photoelectric effect and the electrons emitted through this process are called photoelectrons. The salient feature of the photoelectric effect is that electron emission only occurs when the frequency of the incident light exceeds a particular threshold value When this condition is met,  the number of electrons emitted depends on the intensity of the incident light, but  the kinetic energies of the emitted electrons depend on the frequency of the light. - Albert Einstein (1905)- proposed that electromagnetic radiation has particle-like qualities and that “particles” of light, subsequently called photons by G. N. Lewis, have a characteristic energy given by Ephoton = hv, where v is the frequency of light and the following value for the constant h. We now call it Planck’s constant, and it has the value h = 6.62607 x 10-34 J s Properties of Waves Wave- a vibrating disturbance by which energy is transmitted. a. Wavelength (lambda) is the distance between identical points on successive waves. b. Frequency (nu)- the number of waves that pass through a particular point in 1 second.  c. Amplitude- is the vertical distance from the midline of a wave to the peak or trough. Bohr’s Theory Of The Hydrogen Atoms Neil Bohr (1913)- postulated the following explain the line spectrum of the hydrogen atom 1. The electron moves about the nucleus with speed u in one of a fixed set of circular orbits; as long as the electron remains in a given orbit, its energy is constant and no energy is emitted. Thus, each orbit is characterized by a fixed radius, r, and a fixed energy, E. 2. The electron’s angular momentum, , is an integer multiple of , that is , with n = 1, 2, 3, and so on. 3. An atom emits energy as a photon when the electron falls from an orbit of higher energy and larger radius to an orbit of lower energy and smaller radius. Bohr was able to derive equations for the energies and radii of the allowed orbits: RH is the numerical constant, called the Rydberg constant, with a value of RH = 2.17868 x 10-18 J According to the equation:  The energy of the hydrogen atom is quantized. By this we mean the energy is restricted to specific values: - RH , -( RH /4), -( RH /9), -( RH /16) and so on.  All the allowed energy values are negative. The theory that leads to equation employs the convention that the energy of the electron is defined to be zero when the electron is free of the nucleus, that is, when it is infinitely far away from the nucleus. Physically, n= corresponds to the situation in which the electron is free of the nucleus. Quantum Mechanics Werner Heinsenberg- formulated what is now known as the Heisenberg uncertainty principle: it is impossible to know simultaneously both the momentum p (defined as mass times velocity) and the position of a particle with certainty. * Applying the Heisenberg uncertainty principle to the hydrogen atom, we see that in reality the electron does not orbit the nucleus in a well-defined path, as Bohr thought. If it did, we could determine precisely both the position of the electron (from its location on a particular orbit) and its momentum (from its kinetic energy) at the same time, a violation of the uncertainty principle. Erwin Schrödinger (1926)- formulated an equation that describes the behavior and energies of submicroscopic particles in general, an equation analogous to Newton’s laws of motion for macroscopic objects. - specifies the possible energy states the electron can occupy in a hydrogen atom and identifies the corresponding wave functions. These energy states and wave functions are characterized by a set of quantum numbers, with which we can construct a comprehensive model of the hydrogen atom. Quantum Numbers describe the distribution of electrons in hydrogen and other atoms. These numbers are derived from the mathematical solution of the Schrödinger equation for the hydrogen atom. 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 (wholenumber) values: 1, 2, 3, 4, and so on. The larger the value of n, the higher the energy and the larger the orbital, or electron shell. The angular momentum quantum number l The angular momentum quantum number 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. 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, 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. 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 counterclockwise. Only two values are allowed for ms: +1⁄2 or –1⁄2. For each subshell, there can be only two electrons, one with a spin of +1⁄2 and another with a spin of –1⁄2. Atomic Orbitals s orbitals All s orbitals are spherical in shape but differ in size, which increases as the principal quantum number increases. Although the details of electron density variation within each boundary surface are lost, there is no serious disadvantage. For us the most important features of atomic orbitals are their shapes and relative sizes, which are adequately represented by boundary surface diagrams. p orbitals It should be clear that the p orbitals start with the principal quantum number n = 2. If n = 1, then the angular momentum quantum number / can assume only the value of zero; therefore, there is only a 1 s orbital. As we saw earlier, when / = 1, the magnetic quantum number m/ can have values of -1, 0, 1. Starting with n = 2 and / = 1, we therefore have three 2 p orbitals: 2 p x , 2 p y , and 2 p z. d Orbitals and Other Higher-Energy Orbitals When / = 2, there are five values of m/, which correspond to five d orbitals. The lowest value of n for a d orbital is 3. Because / can never be greater than n - 1, when n = 3 and / = 2, we have five 3 d orbitals (3dxy, 3dyz, 3dxz, 3dx2 2 y2, and 3dz2). Orbitals having higher energy than d orbitals are labeled f , g ,... and so on. Electronic Configurations The electron configuration of an atom is a designation of how electrons are distributed among various orbitals in principal shells and subshells. Rules for Assigning Electrons to Orbitals 1. Electrons occupy orbitals in a way that minimizes the energy of the atom. Aufbau principle, the total energy of an atom depends not only on the orbital energies but also on the electronic repulsions that arise from placing electrons in particular orbitals. With only a few exceptions, the order in which orbitals fill is 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p 2. Only two electrons may occupy the same orbital, and these electrons must have opposite spins. Pauli exclusion principle, this principle states that no two electrons in an atom can have the same set of four quantum numbers. If two electrons in an atom should have the same n, /, and m/ values (that is, these two electrons are in the same atomic orbital), then they must have different values of ms. In other words, only two electrons may occupy the same atomic orbital, and these electrons must have opposite spins. 3. When orbitals of identical energy (degenerate orbitals) are available, electrons initially occupy these orbitals singly and with parallel spins. Hund’s Rule, a simplified statement of Hund’s rule is that, for a given configuration, the arrangement having the maximum number of parallel spins is lower in energy than any other arrangement arising from the same configuration. This behavior can be rationalized as follows. Here are a couple of electron configurations you can use to check your conversions from energy level diagrams: Chlorine (Cl): 1s22s22p63s23p5 Iron (Fe): 1s22s22p63s23p64s23d6 Electron Configurations and the Periodic Table  The group 1 atoms (alkali metals) have one outer-shell (valence) electron in an s orbital, that is,  The group 17 atoms (halogens) have seven outer-shell (valence) electrons, in the configuration ns2np5.  The group 18 atoms (noble gases)—with the exception of helium, which has only two electrons—have outermost shells with eight electrons, in the configuration ns2np6. four blocks of elements according to the subshells being filled: s block. The s orbital of highest principal quantum number (n) fills. The s block consists of groups 1 and 2 (plus He in group 18). p block. The p orbitals of highest quantum number (n) fill. The p block consists of groups 13, 14, 15, 16, 17, and 18 (except He). d block. The d orbitals of the electronic shell n - 1 (the next to outermost) fill. The d block includes groups 3, 4, 5, 6, 7, 8, 9, 10, 11, and 12. f block. The f orbitals of the electronic shell n - 2 fill. The f-block elements are the lanthanides and the actinides. III. PERIODIC TRENDS OF THE ELEMENT Modern periodic Table In the mid-1800s, Dmitri Mendeleev, a Russian chemist, noticed a repeating pattern of chemical properties in the elements that were known at the time. Mendeleev arranged the elements in order of increasing atomic to form something that fairly closely resembles the modern periodic table. He was even able to predict the properties of some of the then-unknown elements. Later, the elements were rearranged in order of increasing atomic number, the number of protons in the nucleus of the atom. Arrangement of Elements Metals you can see a stair-stepped line starting at boron (B), atomic number 5, and going all the way down to polonium (Po), atomic number 84. Except for germanium (Ge) and antimony (Sb), all the elements to the left of that line can be classified as metals. These metals have properties that you normally associate with the metals you encounter in everyday life. They’re solid at room temperature (with the exception of mercury, Hg, a liquid), shiny, good conductors of electricity and heat, ductile (they can be drawn into thin wires), and malleable (they can be easily hammered into very thin sheets). All these metals tend to lose electrons easily. As you can see, the vast majority of the elements on the periodic table are classified as metals. Non Metals Except for the elements that border the stair-stepped line the elements to the right of the line, along with hydrogen, are classified as nonmetals. Nonmetals have properties opposite those of the metals. The nonmetals are brittle, aren’t malleable or ductile, and are poor conductors of both heat and electricity. They tend to gain electrons in chemical reactions. Some nonmetals are liquids at room temperature. Metalloids The elements that border the stair-stepped line in the periodic table are classified as metalloids. The metalloids, or semimetals, have properties that are somewhat of a cross between metals and nonmetals. They tend to be economically important because of their unique conductivity properties (they only partially conduct electricity), which make them valuable in the semiconductor and computer chip industry. (The term Silicon Valley doesn’t refer to a valley covered in sand; silicon, one of the metalloids, is used in making computer chips.) Arrangement by Families and Periods Periods: The seven horizontal rows are called periods. The periods are numbered 1 through 7 on the left-hand side of the table. Within each period, the atomic numbers increase from left to right. Families: The vertical columns are called groups, or families. The families may be labeled at the top of the columns in one of two ways. The older method uses roman numerals and letters. The newer method simply uses the numbers 1 through 18. Effective Nuclear Charge The effective nuclear charge (Zeff ) is the nuclear charge felt by an electron when both the actual nuclear charge (Z) and the repulsive effects (shielding) of the other electrons are taken into account. In general, Zeff is given by where sigma is called the shielding constant (also called the screening constant). The shielding constant is greater than zero but smaller than Z. Atomic Properties Atomic Radius. which is one-half the distance between the two nuclei in two adjacent metal atoms or in a diatomic molecule. Within a group we find that atomic radius increases with atomic number. Ionic radius. the radius of a cation or an anion - Cations are smaller than the atoms from which they are formed. - For isoelectronic cations, the more positive the ionic charge, the smaller the ionic radius. - Anions are larger than the atoms from which they are formed. For isoelectronic anions, the more negative the charge, the larger the ionic radius. Ionization Energy - The ionization energy, Ei, is the quantity of energy a gaseous atom must absorb to be able to expel an electron. The electron that is lost is the one that is highest in energy, and therefore, is most loosely held. - With relatively few exceptions, ionization energies increase from left to right across a period and decrease from top to bottom within a group. - Ionization energies decrease as atomic radii increase. Electron Affinity - Electron affinity, Eea, can be defined as the enthalpy change, that occurs when an atom in the gas phase gains an electron. According to this definition, the electron affinity of fluorine is a negative quantity. - The overall trend is an increase in the tendency to accept electrons (electron affinity values become more positive) from left to right across a period. The electron affinities of metals are generally lower than those of nonmetals. The values vary little within a given group. Electronegativity - The electronegativity (EN) of an element is a measure of the relative tendency of an atom to attract electrons to itself when it is chemically combined with another atoms - Elements with high electronegativities (non metals) often gain electrons to form an ions. Elements with low electronegativities (metals) often lose electrons to form cations.  - Electronegativities usually increases from left to right across periods and decrease from to bottom within groups. Magnetic Properties - In a diamagnetic atom or ion, all electrons are paired and the individual magnetic effects cancel out. A diamagnetic species is weakly repelled by a magnetic field. - A paramagnetic atom or ion has unpaired electrons, and the individual magnetic effects do not cancel out. The unpaired electrons possess a magnetic moment that causes the atom or ion to be attracted to an external magnetic field. The more unpaired electrons present, the stronger is this attraction Polarizability - The polarizability of an atom provides a measure of the extent to which its electron cloud can be distorted, for example, by the application of an externally applied electric field or by the approach of another atom, molecule, or ion. It is often expressed in units of volume. The polarizability of an atom depends on how diffuse or spread out its electron cloud is, and in general, polarizability increases with the size of the atom. - polarizability decreases from left to right across a period and increases from top to bottom within a group. IV. IONIC AND COVALENT COMPOUNDS Lewis Theory 1. Electrons, especially those of the outermost (valence) electronic shell, play a fundamental role in chemical bonding. 2. In some cases, electrons are transferred from one atom to another. Positive and negative ions are formed and attract each other through electrostatic forces called ionic bonds. 3. In other cases, one or more pairs of electrons are shared between atoms. A bond formed by the sharing of electrons between atoms is called a covalent bond. 4. Electrons are transferred or shared in such a way that each atom acquires an especially stable electron configuration. Usually this is a noble gas configuration, one with eight outer-shell electrons, or an octet. Lewis Symbol A Lewis symbol consists of a chemical symbol to represent the nucleus and core (inner-shell) electrons of an atom, together with dots placed around the symbol to represent the valence (outer-shell) electrons. Thus, the Lewis symbol for silicon, which has the electron configuration [Ne]3s23p2, is Lewis Structure Lewis structure is a combination of Lewis symbols that represents either the transfer or the sharing of electrons in a chemical bond. IONIC COMPOUND AND BONDING Ionic bonding is the complete transfer of valence electron(s) between atoms. It is a type of chemical bond that generates two oppositely charged ions. In ionic bonds, the metal loses electrons to become a positively charged cation, whereas the nonmetal accepts those electrons to become a negatively charged anion. Ionic bonds require an electron donor, often a metal, and an electron acceptor, a nonmetal. Ionic bonding is observed because metals have few electrons in their outer-most orbitals. By losing those electrons, these metals can achieve noble gas configuration and satisfy the octet rule. Similarly, nonmetals that have close to 8 electrons in their valence shells tend to readily accept electrons to achieve noble gas configuration. In ionic bonding, more than 1 electron can be donated or received to satisfy the octet rule. The charges on the anion and cation correspond to the number of electrons donated or received. In ionic bonds, the net charge of the compound must be zero. This sodium molecule donates the lone electron in its valence orbital in order to achieve octet configuration. This creates a positively charged cation due to the loss of electron. This chlorine atom receives one electron to achieve its octet configuration, which creates a negatively charged anion. Lattice Energy The strength of the bond between the ions of opposite charge in an ionic compound therefore depends on the charges on the ions and the distance between the centers of the ions when they pack to form a crystal. An estimate of the strength of the bonds in an ionic compound can be obtained by measuring the lattice energy of the compound, which is the energy given off when oppositely charged ions in the gas phase come together to form a solid. Example: The lattice energy of NaCl is the energy given off when Na+ and Cl- ions in the gas phase come together to form the lattice of alternating Na + and Cl- ions in the NaCl crystal shown in the figure below. The lattice energies of ionic compounds are relatively large. The lattice energy of NaCl, for example, is 787.3 kJ/mol, which is only slightly less than the energy given off when natural gas burns. NAMING IONS AND IONIC COMPOUNDS - Ionic compounds are neutral compounds made up of positively charged ions called cations and negatively charged ions called anions. - The Stock system, an ion’s positive charge is indicated by a roman numeral in parentheses after the element name, followed by the word ion. Thus, Fe2+ is called the iron(II) ion, while Fe3+ is called the iron(III) ion. This system is used only for elements that form more than one common positive ion. We do not call the Na+ ion the sodium(I) ion because (I) is unnecessary. Sodium forms only a 1+ ion, so there is no ambiguity about the name sodium ion. - The common system, is not conventional but is still prevalent and used in the health sciences. This system recognizes that many metals have two common cations. The common system uses two suffixes (-ic and -ous) that are appended to the stem of the element name. The -ic suffix represents the greater of the two cation charges, and the -ous suffix represents the lower one. - The name of a monatomic anion consists of the stem of the element name, the suffix -ide, and then the word ion. Thus, as we have already seen, Cl− is “chlor-” + “-ide ion,” or the chloride ion. Similarly, O2− is the oxide ion, Se2− is the selenide ion, and so forth. - POLYATOMIC IONS A polyatomic ion is an ion composed of more than one atom. The ammonium ion consists of one nitrogen atom and four hydrogen atoms. Together, they comprise a single ion with a 1+ charge and a formula of NH4+. The carbonate ion consists of one carbon atom and three oxygen atoms and carries an overall charge of 2−. The formula of the carbonate ion is CO32−. The atoms of a polyatomic ion are tightly bonded together and so the entire ion behaves as a single unit. Covalent Bonding Covalent bonding is the sharing of electrons between atoms. This type of bonding occurs between two atoms of the same element or of elements close to each other in the periodic table. This bonding occurs primarily between nonmetals; however, it can also be observed between nonmetals and metals. If atoms have similar electronegativities (the same affinity for electrons), covalent bonds are most likely to occur. Because both atoms have the same affinity for electrons and neither has a tendency to donate them, they share electrons in order to achieve octet configuration and become more stable. Naming Molecular Compounds Naming binary (two-element) molecular compounds is similar to naming simple ionic compounds. The first element in the formula is simply listed using the name of the element. The second element is named by taking the stem of the element name and adding the suffix -ide. A system of numerical prefixes is used to specify the number of atoms in a molecule. Note: - Generally, the less-electronegative element is written first in the formula, though there are a few exceptions. Carbon is always first in a formula and hydrogen is after nitrogen in a formula such as NH3. The order of common nonmetals in binary compound formulas is C, P, N, H, S, I, Br, Cl, O, F. - The a or o at the end of a prefix is usually dropped from the name when the name of the element begins with a vowel. As an example, four oxygen atoms, is tetroxide instead of tetraoxide. - The prefix is "mono"is not added to the first element’s name if there is only one atom of the first element in a molecule. Percent Composition of Compounds  V. REPRESENTING MOLECULES Octet Rule - Octet rule, formulated by Lewis: An atom other than hydrogen tends to form bonds until it is surrounded by eight valence electrons. - Single bond: In a single bond, two atoms are held together by one electron pair. - Double bond: two atoms share two pairs of electrons, the covalent bond is called a double bond. - A triple bond arises when two atoms share three pairs of electrons, as in the nitrogen molecule (N2): Electronegativity - A property that helps us distinguish a nonpolar covalent bond from a polar covalent bond - Polar covalent bond: A polar covalent bond exists when atoms with different electronegativities share electrons in a covalent bond. Consider the hydrogen chloride (HCl) molecule. Each atom in HCl requires one more electron to form an inert gas electron configuration. - Nonpolar covalent bond: Nonpolar covalent bonds are a type of bond that occurs when two atoms share a pair of electrons with each other. These shared electrons glue two or more atoms together to form a molecule. An example of a nonpolar covalent bond is the bond between two hydrogen atoms because they equally share the electrons. Summary Scheme for Drawing Lewis Structure Formal Charge Formal charge is the electrical charge difference between the valence electrons in an isolated atom and the number of electrons assigned to that atom in a Lewis structure. To assign the number of electrons on an atom in a Lewis structure, we proceed as follows: - All the atom’s nonbonding electrons are assigned to the atom. Ex. Ozone molecule (O3) - We break the bond(s) between the atom and other atom(s) and assign half of the bonding electrons to the atom. The formal charge on each atom in O3 can now be calculated according to the following scheme: Resonance - - Resonance structure: one of two or more Lewis structures for a single molecule that cannot be represented accurately by only one Lewis structure. The double-headed arrow indicates that the structures are resonance structures. - Resonance: the use of two or more Lewis structures to represent a particular molecule. Exception to the Octet Rule Here are three general exceptions to the octet rule: 1. Molecules, such as NO, with an odd number of electrons 2. Molecules in which one or more atoms possess more than eight electrons, such as SF6 whose Lewis structure must accommodate a total of 48 valence electrons [6 + (6 × 7) = 48] 3. Molecules such as BCl3, in which one or more atoms possess less than eight electrons. The boron atom has only six valence electrons, while each chlorine atom has eight. A reasonable solution might be to use a lone pair from one of the chlorine atoms to form a B-to-Cl double bond: REFERENCES:  Whitten, K., et.al (2004). General Chemistry. 7th Ed.  Petrucci, R., et.al (2017). General Chemistry: Principles and Modern Applications. 11th Ed.  Moore, J. (2010). Chemistry Essentials for Dummies  Chang, R. (2010). Chemistry. 10th Ed.  www.khanacademy.com  www.chemlibretext.org  www.chemed.chem.purdue.edu  www.sciencedirect.com  www.study.com