Predict bond type from electron configuration and electronegativity and link it to properties

Learning goal: Use valence electron configurations and the octet rule to predict the ions an atom forms and how ionic compounds arise; use electronegativity differences to place a bond on the continuum from nonpolar covalent through polar covalent to ionic; describe metallic bonding; explain why bond forming releases energy and bond breaking absorbs it; and connect each bonding type to the properties of the material.

Why atoms bond at all. Atoms of most elements are more stable bonded to other atoms than alone. Stable means lower in energy: when two atoms bond, energy is released to the surroundings and the bonded system sits at lower energy than the separated atoms. The reverse is also true. Bond energy is the energy required to break a particular bond in one mole of gaseous molecules, and breaking a bond always requires an input of energy, so bond breaking is endothermic and bond forming is exothermic. That pair of statements is the foundation for every energy argument later in the course. A common misconception says that breaking bonds releases energy, because fuels are said to store energy in their bonds. What actually happens when a fuel burns is that weaker bonds break and stronger bonds form, and the net result is a release; the release comes from the forming, not the breaking.

Valence electrons and the octet rule. Valence electrons are the highest-energy electrons in an atom, those in its outermost occupied level, and they decide bonding. Elements in the first three periods, apart from hydrogen and helium, tend to gain, lose or share electrons until they have eight valence electrons, the configuration of a noble gas. This octet rule predicts the most stable ion for main-group elements: sodium, [Ne]3s¹, loses one electron to become Na⁺ with the neon configuration; magnesium, [Ne]3s², loses two to become Mg²⁺; chlorine, [Ne]3s²3p⁵, gains one to become Cl⁻ with the argon configuration; oxygen gains two to become O²⁻. Hydrogen aims for the two-electron configuration of helium instead. The rule is a strong guide for main-group elements but has exceptions even in the first three periods (boron in BF₃ has only six valence electrons, and phosphorus and sulfur can hold more than eight, as in PCl₅ and SF₆); it is a rule about stability, not a law.

Transfer or share, decided by position. Whether atoms transfer electrons or share them depends on where they sit on the periodic table. Metals on the left have low ionization energies and low electronegativities; they lose electrons easily. Nonmetals on the right have high electronegativities; they hold their own electrons and attract others. Between a metal and a nonmetal, transfer is favored and an ionic bond results. Between two nonmetals, neither gives up electrons, so they share, forming a covalent bond. Between metal atoms, a third arrangement occurs, metallic bonding.

Ionic bonding. When sodium meets chlorine, each sodium atom transfers its 3s electron to a chlorine atom. Writing it in configurations: Na [Ne]3s¹ + Cl [Ne]3s²3p⁵ → Na⁺ [Ne] + Cl⁻ [Ar]. The oppositely charged ions attract in every direction, so they do not form pairs; they pack into a three-dimensional ionic lattice in which each Na⁺ is surrounded by six Cl⁻ and each Cl⁻ by six Na⁺. A separated pair of ions in the gas phase is far less stable than the ions in the lattice, because in the lattice each ion is attracted by many neighbors; building the lattice releases a large amount of energy. The strong, all-directional attractions explain the properties: ionic compounds are hard, brittle crystalline solids with high melting points; they do not conduct as solids, because the ions cannot move, but conduct when melted or dissolved, when the ions are free; and many dissolve in water.

Covalent bonding and the electronegativity continuum. In a covalent bond two atoms share a pair of electrons that is attracted to both nuclei. A single bond shares one pair, a double bond two pairs and a triple bond three pairs; between the same two elements, more shared pairs mean a stronger and shorter bond, so a C≡C triple bond is stronger and shorter than a C=C double bond, which is stronger and shorter than a C–C single bond. Sharing is rarely equal. Electronegativity difference, the gap between the two atoms' values, predicts where a bond falls on a continuum. When the difference is essentially zero, as in Cl–Cl or H–H, the electrons are shared equally and the bond is nonpolar covalent. As the difference grows the shared electrons spend more time near the more electronegative atom, giving that end a partial negative charge and the other end a partial positive charge: a polar covalent bond with a dipole. When the difference is very large, as in Na–Cl at 2.23, the electron is effectively transferred and the bond is ionic. A common guide treats differences below about 0.4 as nonpolar, from about 0.4 to about 1.7 as polar covalent and above about 1.7 as ionic; these cutoffs are conventions, and a bond such as H–F at 1.78 is best described as very polar. Bond polarity also depends on bond length, since the same charge separation over a longer distance makes a larger dipole.

Bond polarity is not molecular polarity. A molecule can contain polar bonds and yet be nonpolar overall, if the bond dipoles are arranged symmetrically so that they cancel. Carbon dioxide has two polar C=O bonds (difference 0.89) pointing in exactly opposite directions along a straight line, so the molecule has no net dipole. Water has two polar O–H bonds (difference 1.24) at an angle, so the dipoles add and water is polar. To decide molecular polarity you need the shape, which the next lesson supplies; here the rule is that identical polar bonds arranged symmetrically cancel and unsymmetrical or unlike bonds do not.

Metallic bonding. In a metal the valence electrons are not held by individual atoms. The metal atoms form a lattice of positive ions through which the valence electrons move freely, a shared pool often called a sea of electrons. The lattice is held together by the attraction between the positive ions and the mobile electrons. This one picture explains the metallic properties: electrical conductivity, because the electrons move when a voltage is applied; thermal conductivity, because the mobile electrons carry energy quickly; and malleability and ductility, because layers of ions can slide past one another without the bond breaking, since the electron sea simply flows around the new positions. In a covalent bond the electrons are localized between two particular atoms; in a metallic bond they belong to the whole lattice.

Carbon and the shapes of large molecules. Carbon has four valence electrons and forms four covalent bonds, with other carbon atoms and especially with hydrogen, oxygen, nitrogen and sulfur. Chains, rings and branching networks result, and these skeletons appear in fossil fuels, in synthetic polymers such as polyethylene and nylon, and in the large molecules of living things. A polymer is a very long molecule made by joining many small molecules, monomers, end to end; polymerization is the reaction that links them. Ethene monomers join to form polyethylene; cellulose in plant cell walls and proteins are natural polymers; polyester and polyethylene are synthetic. The properties of a polymer depend on the monomers used and on the length and structure of the chain: longer, less branched chains pack more closely and give stiffer, stronger materials. Detailed biological structures are for later courses.

One substance, several bond types. Most compounds contain more than one type of bond. Sodium sulfate contains ionic bonds between Na⁺ and sulfate ions and polar covalent bonds inside each sulfate ion. Naming the bonds in a substance means naming all of them.

A worked model: classify from properties. Four white or gray solids are tested. Solid A melts easily in a warm-water bath, does not dissolve in water and does not conduct: molecules held by weak attractions, so covalent bonding, and probably nonpolar since water does not dissolve it. Solid B does not melt in the bath, dissolves in water, and the solution conducts strongly while the dry solid does not: ionic. Solid C conducts as a solid and can be bent: metallic. Solid D dissolves in water but its solution does not conduct and it does not melt in the bath: polar covalent molecules, such as sugar, which dissolve without forming ions. The reasoning always names the particles (ions, molecules, metal ions plus electrons) and the attraction between them.

Investigation: sort household solids by bonding type. Materials: table salt, granulated sugar, a paraffin candle shaving, a strip of aluminum foil, distilled water, vegetable oil, small cups, a battery-powered conductivity tester (a 3 V cell, a bulb or LED with a resistor, and two clip leads), a bowl of hot tap water about 60 °C, goggles. Steps: test each dry solid for conduction; place a little of each in a sealed bag floating in the hot water for three minutes and record whether it melts; stir a small amount of each into water and into oil and record dissolves or not; test each water mixture for conduction. Record all four results for each solid in one table. A result that supports the claim: salt is the only one whose solution conducts and it does not melt in the bath (ionic); foil conducts dry and does not melt (metallic); wax softens or melts (paraffin melts at roughly 50 to 65 °C, so in a 60 °C bath it may only soften), does not dissolve in water and never conducts (nonpolar covalent); sugar dissolves in water without conducting and does not melt in the bath (polar covalent molecules). What the model cannot show: a bulb tester cannot distinguish weak from strong conduction, and a 60 °C bath cannot tell a melting point of 200 °C from one of 800 °C, so the classification is coarse. No open flame is used; do not heat sugar or wax over a burner without a supervising adult.

Worked example 1 — form an ionic compound from configurations. Magnesium is [Ne]3s²; chlorine is [Ne]3s²3p⁵. Magnesium reaches an octet by losing two electrons; each chlorine needs only one. So one magnesium atom transfers one electron to each of two chlorine atoms: Mg → Mg²⁺ [Ne] + 2e⁻ and 2Cl + 2e⁻ → 2Cl⁻ [Ar]. The ratio is one Mg²⁺ to two Cl⁻, giving MgCl₂, and the ions assemble into a lattice, releasing energy. The formula is predicted by the configurations, not memorized.

Worked example 2 — place bonds on the continuum. Using electronegativities Na 0.93, H 2.20, C 2.55, N 3.04, O 3.44, Cl 3.16: C–H differs by 0.35, nonpolar in practice; N–H differs by 0.84, polar covalent with the partial negative charge on nitrogen; O–H differs by 1.24, strongly polar with the partial negative charge on oxygen; Na–Cl differs by 2.23, ionic. Arranged from least to most polar: C–H < N–H < O–H < Na–Cl. The number ranks the bonds; the sign of the dipole follows from which atom is more electronegative.

Worked example 3 — polar bonds, nonpolar molecule. Carbon tetrachloride, CCl₄, has four C–Cl bonds each with a difference of 0.61, so each bond is polar with chlorine partially negative. The four bonds point toward the corners of a tetrahedron, an arrangement in which four equal dipoles cancel exactly. The molecule has no net dipole and is nonpolar, which is why it does not mix with water. Replace one chlorine with hydrogen to make CHCl₃ and the symmetry is broken: the three C–Cl dipoles no longer cancel against a C–H, so chloroform is polar.

Worked example 4 — energy of bond breaking and forming. Hydrogen and chlorine react: H₂ + Cl₂ → 2HCl. Two bonds break, H–H and Cl–Cl, and two H–Cl bonds form. Breaking requires energy and forming releases it. The reaction is strongly exothermic, so the two new H–Cl bonds must release more energy than the H–H and Cl–Cl bonds absorbed; the products have stronger bonds in total than the reactants. The numbers, from a bond-energy table, are used in the rates-and-energy lesson; here the argument is qualitative, and it applies to any reaction: compare the total strength of bonds formed with the total strength of bonds broken.

Representation task: three bonding pictures and a continuum. Draw a labeled particle picture of (1) a small piece of a sodium chloride lattice showing alternating Na⁺ and Cl⁻ and the attractions in more than one direction, (2) a chlorine molecule with its shared pair and a hydrogen chloride molecule with its shared pair drawn closer to chlorine and partial charges marked, and (3) a metal lattice with positive ions and free electrons. Beneath them, draw a line from difference 0 to 3.2, mark 0.4 and 1.7 as the conventional guide boundaries, and place Cl–Cl, C–H, N–H, O–H, H–F and Na–Cl on it using the values in the lesson.

Independent performance: classify four unknowns. Carry out the household-solids investigation or use a teacher-provided set of results for four unknown solids. Success criteria: a complete four-test table for each solid; a bonding classification for each with a justification that names the particles present and the attraction between them; a statement of which single test was most decisive for each solid and why; one limitation of the test set; and a paragraph explaining why breaking bonds costs energy and forming bonds releases it, using the term bond energy correctly.

Extension: evaluate a commercial claim. Choose an advertised product whose marketing uses chemistry words, such as an ionic hair dryer said to emit ions that dry hair faster, a mineral bracelet said to release helpful ions, a shake weight dumbbell said to build strength through vibration, or a glue advertised as forming unbreakable bonds. Design an investigation that could test the claim fairly: state the claim as a measurable prediction, the variable you would change, the variables you would hold constant, and what result would count against the claim. Determine the function of the product, the intent of the wording and any potential bias in who makes the claim. Present your findings in multiple formats: a short written evaluation, a labeled diagram of the test, and a one-minute spoken summary.