4 Molecules and Compounds

Transcription

1 4 Molecules and ompounds 4 Molecules and ompounds Atoms that have assembled into substances (new or old) are bonded (glued) together. The way the atoms are bonded together will create different properties for the compound. Generally, there are two types of bonding: ionic and covalent. Ionic ompounds Atoms are held together by the attraction of opposite charges between a cation (often a metal) and an anion (often a nonmetal). No individual molecules, just an arrangement of ions in space where generally each cation (+) is surround by anions ( ) and vice versa. R Because there are no individual molecules, ionic formulas represent the simplest ratio of cation to anion that is needed to form the neutral compound. Subscripts are used to represent the ratio. Note that since the subscripts represent a simplest ratio, and not the actual numbers of atoms in the compound, all ionic formulas are empirical formulas. Ionic ompounds can be identified because they often contain a metal with a nonmetal. Most are binary (only two types of atoms). It is common for the ions to already exist and be attracted to each other. Na 1+ + l 1- Nal [ionic compound of sodium chloride] The ions could form from the neutral atoms via an oxidation reduction reaction (transfer of electrons) and then be attracted to each other ionically. Na + l 2 Nal [ionic compound of sodium chloride] Generally, the greater charges have a stronger attraction and thus are stronger ionic bonds. Na 1+ + l 1- Nal Mg Mg [stronger ionic bond] Formula Writing [forming ionic compounds from ions] Write the symbols side by side without the charges and determine how many of each ion are needed to create a neutral formula, write that number as a subscript. [Typically this is done by crisscrossing the charges and then reducing.] Note that if a polyatomic ion is used, the subscript is written outside of ( ) s Example: Ti Ti 2 Dissociating into Ions [breaking ionic compounds into ions] Mg 2+ + P 4 Mg 3(P 4) 2 1 st Split the formula into two halves. ften this is done by writing the first symbol as the first half and the remaining symbols as the second half. Do NT write any subscripts unless they are inside ( ) s or follow two capital letters. See the example. 2 nd The first ion is the cation (+) and the second ion is the anion ( ). See the example. 3 rd Now determine the magnitude of the charge. Do the opposite of Ions to formulas by crisscrossing the subscripts into the position of charges. See the example 4 th To make sure you are right (and often you are not) you need to check your charges with the periodic table or with the charges on the polyatomic ions you were to memorize.

2 4 Molecules and ompounds Example: For each equation, write the correct ions in the blanks. A bb a + or Ab(B x)a + 1 st A bb a A + B or A b(b x)a A + Bx 2 nd A bb a A + + B - or A b(b x) a A + + B x - 3 rd A bb a A a+ + B b- or A b(b x) a A a+ + B x b Example 1: Ti 2 + Ti 2 Ti Example 2: Mg 3(P 4) th Examples of hecking harges Mg 3(P 4) 2 Mg 2+ + P 4 u 2 u Note that the charge for matches the periodic table. So we assume all of the the charges are correct. K 2S 4 K 1+ + S4 2- Note the charge for K and S 4 matches the periodic table and the memorized charge, so we assume the charges are correct. Also note, the subscript 4 was not criss-crossed. The invisible 1 outside the invisible ( ) s was criss-crossed. The 4 follows two capital letters and stays put. u u Note that the charge for should be a 2- from the periodic table. So, we must double all charges to get it from 1- to 2-. So... u u Nomenclature Always name the ions not the formulas!!!! (cation then anion). The name tells the type of ions involved not how many of each ion so just name each type of ion. 1 st cations: use the elements name; if more than one oxidation state is possible (d-block) follow with the charge in Roman numerals in parentheses Example: Li 1+ lithium Al 3+ aluminum if d-block elements Ni 2+ nickel (II) Ni 3+ nickel (III) 2 nd anions: if monatomic then use the elemental name but with an ide ending if polyatomic then use the memorized name Example: F 1- fluoride (not fluorine) As arsenide (not arsenic) ombined Example: Mg 3N 2 magnesium nitride us 4 copper (II) sulfate List of ommon Polyatomic Anions phosphate P 4 sulfate 2- S 4 nitrate 1- N 3 phosphite P 3 sulfite 2- S 3 nitrite 1- N 2 hydroxide 1-1+ ammonium N 4 2+ mercury (I) g 2 carbonate 2-3 cyanide N 1- acetate

3 4 Molecules and ompounds ovalent Molecules Atoms are held together by covalent bonds. Each covalent bond is a shared pair of electrons, forming individual molecules. nly nonmetal atoms covalently bond. Metal atoms need to lose electrons to become stable so there is no driving force for a metal atom to share electrons. Poles are regions of charge within a molecule. In covalent compounds, poles occur because the electrons are unevenly distributed. Nonpolar covalent bond very low electronegativity difference between the bonded atoms, results in a nearly equal sharing of the electron pair and thus no development of a pole, (often - or same atoms, ex -) Example: nonpolar covalent bonds are found in methane, 4, and nitrogen, N 2. Polar covalent bond larger electronegativity difference, results in an unequal sharing of the electron pair and thus partial charge development Example: polar covalent bonds are found in water, 2. The covalent bond occurs due to overlapping of atomic orbitals between the atoms (one from each) forming molecular orbitals. The creation of electron orbitals shared/located between two atoms creates a strong bond which is difficult to break. Single bonds are sigma,, occur when one electron pair is shared and is located directly between the atoms. Double/Triple Bonds are pi,, bonds occurs when 2 or 3 electron pairs are shared. The first sharing is in a sigma bond. Because any orbital can hold a max of 2 e s, the second shared pair is located above and below the sigma bond. ne p-orbital from each atoms is used to reach around the sigma bond creating a bond. The third pair (triple) is located in another set of p-orbitals in front and in back of the sigma bond, creating a second bond. Structures and Structural Formulas The d arrangement of atoms in a covalent molecule causes many of the characteristics and properties of that substance. Many covalent molecules share the same molecular formula but have different structural formulas and are known as isomers. So the structure is very important. VSEPR [Valence Shell Electron Pair Repulsion] In molecules pairs of electrons due to their like charges will repel one another. So atoms in molecules typically spread-out evenly and as much as possible to reduce these repulsions. BUT not all electron pairs are alike (lone pairs vs bonded pairs) so not all repulsions are alike. Lone pairs repulse more than bonded pairs. So, in a molecule, lone pairs occupy more space and push bonded pairs together. So when drawing structures, lone pairs get more space than bonded pairs. Drawing Structures Structural formulas are drawings that show all thed arrangement of all the atoms in a molecular formula. Structural formulas are often called Lewis Structures or just structures. A Lewis Dot Diagram is the same as a Lewis Structure, but with all electrons shown as dots. Lewis Structure dimensional arrangement of atoms in a molecule An electron pair is drawn as either 2-dots or a straight line Bonded pairs are between atoms while lone pairs are only on one atom

4 4 Molecules and ompounds ow to draw a structure 1 st method alculate the total number of valence electron pairs in your molecule. Next bond all atoms to the least electronegative atom. omplete octets (except for atoms that form duets or sextets), beginning with the most electronegative atom, until all valence electron pairs are used. If there are any extra valence electron pairs place them on the central atom. If there are not enough pairs to complete octets then form multiple bonds until all octets are completed. r 2 nd method Draw all the atoms {the atom pieces} with valence electrons as dots. Bond {connect} all single electrons {dots}with single electrons {dots} on other atoms. Then redraw neatly and symmetrically. The atom pieces : N : : F Possible Geometries linear tetrahedral trigonal pyramidal bent linear trigonal planar linear linear linear Using the Families: Because every atom within a family has the same valence electron configuration, they all form the same number of bonds and are drawn the same way. Examples: 2 2S 2Se Example: Draw the structure of the molecule and state the geometry of 2 1 st method Valence e s: =1, =4 =6 so 2(1) pairs of electrons to use for the structure 2 1 st bond to central atom 6 prs- 3 bonds = 3 prs left 2 nd complete octets 3 prs- 3 prs = 0 prs left 3 rd Share to complete s octet Geometry is trigonal planar

5 2 nd method atom pieces : trigonal planar Nomenclature Unlike the name of an ionic compound, the name of a covalent molecule does indicate how many of each atom are present in the molecule. To write a name, name each element, typically with a prefix on the element denoting the number of that atom in the molecule List of Prefixes mono- one di- two tri- three tetra- four pentafive hexa- six hepta- seven octa- eight nona- nine decaten Electronegativity Difference and Bond haracter Electronegativity Difference, E, E = E 1-E 2, For example, each bond in 2 has: example: l 4 carbon tetrachloride P 2 5 diphosphorus pentoxide E = E - E E = E = 1.44 As the electronegativity difference grows from 0.00 to >3.00, the sharing of electrons becomes more skewed. First one atom becomes partially negative, -, while the other becomes partially positive, +, and the bond is polarized, a polar covalent bond. Then as the sharing becomes more and more unequal, the electrons reside only on one atom. The atoms are now ionized (+ or -), and the bond is called ionic. This is shown in Table 1. Table 1 nonpolar covalent polar covalent ionic Electronegativity Difference, E Percent Ionic haracter 0% 10 % 20 % 30 % 40 % 50 % 60 % 70 % 80 % 90 % Percent ovalent haracter 100 % 90 % 80 % 70 % 60 % 50 % 40 % 30 % 20 % 10 % For a shared pair of electrons, if one atom is able to attract the electrons to itself (more electronegative) that atom will begin to become negatively charged, -, (a negative pole) while the other atom (least electronegative) begins to become positively charge, +, (a positive pole). The two atoms become a dipole (meaning 2 poles), and the bond will become a polar covalent bond. [see Table 2] If the difference in attracting the electrons, E, is so great, then one atom may just take the electrons for itself. This stops any sharing of electrons, and the bond is an ionic bond. The atom that took the electrons is the anion (negative), and the atom that lost the electrons is the cation (positive). [see Table 2] If the difference in attracting the electrons, E, is very small, then the sharing remains relatively equal and no charges develop. No developed charges means there are no pole, which makes the bond a nonpolar covalent bond. In the special case that the electronegativity difference, E, is zero, then no atom attracts the electrons to itself and the sharing is perfectly equal. Such a bond is called a pure covalent bond and is nonpolar also. [see Table 2]

6 Table 2 Electronegativity Difference, E Polarity Bond Type E = 0.00 nonpolar pure covalent 0.00 < E < 0.65 nonpolar nonpolar covalent 0.65 < E < 1.67 polar polar covalent 1.67 < E ionic More on Dipoles An electric dipole consists of two opposite charges that are the same magnitude and separated in space. Unequal sharing of electrons (E > 0) results in one end of the bond being negative and one end being positive. This is a dipole (2 poles, one negative, one positive) The atom with the greater electronegativity becomes partially negative, -, The atom with the smaller electronegativity becomes partially positive, +. Drawing an arrow from + to - shows the electric field between the charges. Practice Problem: ydrogen fluoride, F o Draw the Lewis Structure o Determine the polarity and bond type for each bond. o Draw the dipole for each bond. F, would be written as F Since E = 1.10, the bond is polar covalent. E = = 1.10} F has the greater electronegative so it is partially negative, -, and with the smaller electronegativity is partially positive, +. Final Answer: : F + - polar covalent

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