Presentation on theme: "Topic 11: Chemical Bond Formation LECTURE SLIDES Valence electrons Ionic Bonding Covalent Bonding Lewis Structures Acid, Anion Relationships Resonance."— Presentation transcript:
Topic 11: Chemical Bond Formation LECTURE SLIDES Valence electrons Ionic Bonding Covalent Bonding Lewis Structures Acid, Anion Relationships Resonance Structures Bond Length vs Bond Order Octet Violators Formal Charge Kotz & Treichel, 9.1-9.6
CHAPTER 9: BONDING AND MOLECULAR STRUCTURE Now that we have examined the structure of the atom and the arrangement of its electrons, we are ready to turn to the molecules and compounds they form. Our next studies will center on the bonds that hold the atoms together in compounds, molecules and polyatomic ions. We will also examine the three dimensional shape of these species, and their polar or non polar nature.
The interactions between atoms which lead to bond formation are all centered around the electrons in incomplete subshells and in incomplete outer shells: the valence electrons... The atoms of the elements lose, gain or share these electrons to achieve, where possible, the noble gas configurations we have met. BOND FORMATION
For the “main group elements”, the s and p block members, the electrons available for bonding, the “valence electrons”are the outer shell s and p electrons (except those of the noble gases!) In forming compounds from these elements, only these electrons will be used, in two ways: They may be transferred to form ions so that incomplete subshells are completed or removed; They may be shared so that two atoms together have complete subshells
For the transition metals, the valence electrons include both the s electrons from their outermost shell and also electrons from their inner, incomplete d subshell. The PT column number of the main group and transition metals gives the sum of valence electrons for each element in the family. Note that the column number indicates the maximum positive charge (or oxidation state) these metals can achieve in a compound through loss of e’s in a chemical reaction.
For all main group elements (the columns 1A-8A), it is convenient for bonding purposes, to represent the elements as “Lewis Dot Symbols”, which include one dot for each of their valence electrons. The tendency of these elements to achieve an outer shell configuration of eight electrons (the “octet rule”) is easily visualized through use of these symbols. The valence electrons for transition elements (columns 3B-8B, 1B, 2B) are not represented by dot symbols.
LEWIS DOT STRUCTURES FOR PERIOD 2 All elements, same column: same dot structure
The elements come together to form compounds so that each element can achieve a more satisfactory outer shell electronic configuration. Elements may lose or gain electrons resulting in cation and anion formation and the attraction between the two which we call the “ionic bond” Elements may share one or more pairs of electrons. The attraction of both nuclei for the same pair of electrons results in the force we call the “covalent bond”.
The “ionic bond”: attraction of opposite charges when transfer of electrons cause formation of positive and negative species: cations and anions. The individual ions radiate charge in all directions and cluster in geometric patterns which are described as crystal lattices. Note in the following slide that each ion has many neighbors and the compound itself is not molecular in nature: no discrete “formula units” exist.
Ionic compounds are all solids at room temperature with elevated melting points. Their melting points reflect the very high degree of attraction exhibited by these fully charged particles, which depends on the magnitude of their charge and the ionic size: The larger each charge and the smaller each ion, the greater the attraction. Energy = n (+) X n (-) d n= magnitude of charge d=distance between ions
Ion formation and the resulting ionic bond occurs when metals of sufficiently low electronegativity (X) react with non-metals of sufficiently high X values. The most ionic of compounds are those formed between the active s block metals (X < 1)with the non-metals whose X values are 3.0 or larger. All compounds we have met containing “polyatomic” anions or ammonium are also of course truly ionic type compounds.
ELECTRONEGATIVITY VALUES, MAIN GROUP ELEMENTS METALS METALLOIDS Non-metals
The second mode of bond formation occurs when elements share one or more pairs of electrons to achieve where possible an outer shell octet. The attraction of both nuclei for the same pair of shared electrons is the basis of the covalent bond. THE COVALENT BOND
Covalent bonds are directed between two atoms, sharing together one or more pairs of electrons. This type of bonding leads to formation of discreet molecules, individual units made up of two or more atoms covalently bonded together. Any formula consisting solely of nonmetals and metalloids can assumed to molecular and covalent in nature. Covalent bonds hold together the atoms within a polyatomic ion. Occasionally, metals with higher electronegativity values will form a compound more covalent than ionic in nature.
Group Work, 11.1: Bond Type Ionic compounds, en >1.6 or 1.7
LEWIS DOT STRUCTURES: MOLECULES AND COMPOUNDS We are next going to use the Lewis Dot Symbols for the “main group elements” to represent the bonding and structure for various species, molecules, compounds and polyatomic ions. all valence electrons for every atom will be included all shared pairs of e’s will be indicated by a “dash” all unshared pairs of e’s will be indicated by a dot We will use the “octet rule” as our guiding principle.
OXYGEN ACTS LIKE THIS* CORRECT LEWIS STRUCTURE, Covalent double bond *Required whole new bonding theory to explain...
Lewis Structures: Compounds and Polyatomic Ions GUIDELINES Decide on arrangement of atoms. For most species, the element written first in the molecule or ion is the central atom and the remainder of the atoms are grouped around it. Hydrogen is a problem in “oxo acids” where it is written first in the formula. Ignore H, start with the next atom in formula and place the H or H’s on the O or O’s. First step:
Second Step Add up all available valence electrons. If species is cation, subtract positive charge from total. If species is anion, add negative charge to total. Divide total by two to determine available number of electron pairs Third Step Place a pair of electrons between each pair of bonded atoms to represent a single bond (use a “dash”!)
Fourth Step Place leftover electron pairs around “terminal” atoms to achieve their octet (except H). Do central atom last. Fifth Step Examine central atom to determine if a double or triple bond is required to achieve the central atom’s octet. Do so using unshared pairs, IF central atom is: C, N, P, O, S
GROUP WORK, 11.3: Lewis Structures # 2 Use 5 steps: Arrange; adds up e’s; draw bonds; assign unshared pairs; double bonds if needed to draw Lewis structures for following species: H 3 PO 4 NO 2 1+ ClO 4 1-
Let’s explore the relationship between various “oxo” acids (H, Non metal element, O) and the charge and formula of their anion relative. Recall that acids, by definition, ionize in water to lose one or more H’s as H +. The anion left behind is named according to the name of its “parent” acid. In an acid/base reaction, as we met last unit, acids(H + ) react with bases (OH - ) to form water, leaving behind the anion of the acid and the cation of the base to form a salt.
Recall that acids “ionize” in water, or react with a base to form water, in either case leaving behind some “anion”: H- “Anion” + NaOH H 2 O + Na + An - Acid: HCl, HNO 3 H 2 SO 4 etc... Cl -, NO 3 -, SO 4 2- etc... H- “Anion” H + + Anion - H2OH2O
In all three cases, O 3, NO 3 -, CO 3 2-, when forming a double bond from a “terminal oxygen” one has a choice of moving e’s from several different O’s to makeup the “central atom’s” octet. Examination of experimental evidence (x ray) shed an interesting light on this topic: When two atoms are bonded together, the distance between their nuclei, their “bond length,” depends on whether the bonds between the two are single, double, or triple.
TYPICAL BOND LENGTHS Note that triple bonds are shorter than double and also double shorter than single, as well as being characteristic between any two given atoms. X ray evidence of bond lengths in ozone, nitrate and carbonate ions should therefore prove interesting...
132 pm 121 pm Predicted, “usual” bond lengths: Instead of the predicted bond lengths observed in other compounds, both bonds in x ray showed identical lengths of 127.8 pm, close to an average of 1 1/2 bonds to each O.
Linus Pauling proposed the “theory of resonance” to describe this situation: When two or more equivalent Lewis structures can be drawn for a species, differing only in the position of electron pairs, then none are correct: The real structure is a hybrid of all structures drawn.
The Lewis structures drawn are called “contributing” or “resonance structures” needed to describe the makeup of the hybrid, which resembles all but is none of the above. A special double headed arrow is drawn between the contributing structures to indicate their hypothetical nature:
The hybrid structure, with two equivalent bonds to the central atom, are said to have a bond order of “1.5” or an average of 1 and 1/2 bonds between each O: THE HYBRID STRUCTURE OF OZONE
Bond Order describes the number of bonds between two atoms in a molecule. Normally, the bond number is 1 (a single bond) or 2 (a double bond) or 3 (a triple bond.) When hybrid structures and resonance situations exist, one must average the number of bonds between all atoms affected, and fractional values arise. In the case of the nitrate and the carbonate ions, the number of bonds to the central atom is averaged out over 3 atoms, and 4 bonds/3 atoms= 1.33 bond order. In both cases, x ray data confirms this theory.
The carbonate ion has three equivalent C-O bonds, of a length typical of 1 and 1/3 bond, for a 1.33 bond order.
The nitrate ion also has three equivalent N-O bonds, of a length typical of 1 and 1/3 bond, for a 1.33 bond order.
Group Work 11.5: Resonance Structure and Bond Order Draw two acceptable Lewis Structures for SO 2 and a resonance hybrid. What is the bond order for the bonds in this compound?
OCTET VIOLATORS Another aspect of drawing correct Lewis structures involves the handling of compounds that do not have an octet around the central atom. Three situations exist: 1. More than 4 e - pairs around central atom 2. Less than 4 e - pairs around central atom 3. Molecules with odd number of electrons In all cases we will handle, the irregularity occurs at the central atom; all “terminal atoms” will have normal octet.
EXAMPLE: Note: Only the central atom, P, is an “octet violator”
Note again: Only the central atom exceeds the octet rule.
Case #2: Less than 4 e - pairs around central atom This category specifically applies to the metalloid Boron, but also to metals that form salts that are more covalent in nature than ionic: Beryllium, Aluminum, for example. These elements use their valence e’s to form compounds but do not form an octet in the process and do not accept double bonds to compensate. These “octet deficient” species will react with other atoms however to form polyatomic ions or compounds which relieve the deficiency.
While B will not form a double bond to F to achieve an octet (F’s “don’t do” double bonds), it will accept electron pairs readily from other sources to do so: When one atom donates two electrons for a pair of atoms to share, the bond is called a “coordinate covalent bond” and introduces “charge buildup” in the species formed.
To keep track of this kind of charge within a molecule or polyatomic ion, the concept of “FORMAL CHARGE” is introduced. Formal charges look within a molecule or polyatomic ion and determine how the charges are distributed by considering for each atom: the number of valence e’s it started with the number of bonds formed the number of unshared electrons leftover
For each atom in species: formal charge = # valence e’s - (#bonds + #unshared e’s) FORMAL CHARGE: The “formal charge” system requires a Lewis Dot Structure and assigns an individual “formal” charge to each atom in the species. Formal charge is an alternate “bookkeeping method” for tracking electron distribution to the “oxidation number” system we met previously.
Now let’s return to the compound formed between ammonia and boron trifluoride, and determine formal charges: Formal charges
OXIDATION NUMBERS: “Ox #’s” are assigned or calculated based on known fixed positive and negative charges, and can be determined by examination of the formula for the species. Oxidation numbers are useful to identify how charges change in a redox (oxidation-reduction) reaction.
To see how both work, let’s look at chloric acid, HClO 3, and see how its charge distribution would be described using both the oxidation number and the formal charge systems.
Sum of all charges in compound = 0 Known ox #’s per atom
Finally, both oxidation numbers and formal charges must add up the same way: For compounds, which are always electrically neutral, the sum of all oxidation numbers or the sum of all formal charges must equal zero. For polyatomic ions, which always have a specific charge, the sum of all oxidation numbers or the sum of all formal charges must equal the charge on the ion.
1: Do Ox #’s from formula (same for both!) 2: Do formal charge from Lewis Structure for all atoms: formal charge = # valence e’s -( # bonds+ # unshared e’s) Group Work 11.7: Oxidation Numbers vs Formal Charge